! 7

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1

I DATE DUE

5

CAVLORD PftlNTSO IN U.S.A.



A11103 Qfifib75

NArt INST OF STANDARDS & TECH R.I.C.

A11 103088675

Workshop on Eloc1roc/Electrocata^ysis on QC100 .U57 N0.45S, 1976 C.2 NBS-PUB-C 19

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NBS SPECIAL PUBLICATION 455

U.S. DEPARTMENT OF COMMERCE / National Bureau of Standards

National Bureau of Standards

Library, E-01 Admin. BIdg.

OCT 1 1981

1038

SIS

on Non-Metallic Surfaces
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Electrocatalysis on Non-Metallic Surfaces

Proceedings of a workshop held at the

National Bureau of Standards

Gaithersburg , Maryland, December 9-12, 1975

Edited by

Alan D. Franklin

Institute for Materials Research

National Bureau of Standards

Washington, D.C. 20234

Sponsored by

National Bureau of Standards

National Science Foundation

Energy Research and Development

Administration

U.S. DEPARTMENT OF COMMERCE, Elliot L. Richardson, Secretary

Edward 0. Vetter, Under Secretary

Dr. Betsy Ancker-Johnson, Assistant Secretary for Science and Technology

NATIONAL BUREAU OF STANDARDS, Ernest Ambler, Acting Director

Issued November 1976

Library of Congress Catalog Card Number: 76-600062

National Bureau of Standards Special Publication 455

INat. Bur. Stand. (U.S.), Spec. Publ. 455, 372 pages (Nov. 1976)

CODEN: XNBSAV

U.S. GOVERNMENT PRINTING OFFICE

WASHINGTON: 1976

For sale by the Superintendent of Documents, U S Government Printing Office Washington, D C. 20402 - Price $4.45

otock No 003-003-01699-3/Catalog No. C 13.10:455

GENERAL ABSTRACT

This book presents the proceedings of a Workshop on Electrocatalysis on Non-metallic

Surfaces held at the National Bureau of Standards, Gaithersburg , Maryland, on December 9-

12, 1975. The Workshop was sponsored by the Institute for Materials Research, NBS, the

Division of Materials Research of the National Science Foundation, and the Division of

Conservation Research and Technology of the Energy Research and Development Administration.

The purpose of the Workshop was to review the most recent experimental and theoretical

investigations on electrocatalysis on non-metals and related topics, and to bring together

electrochemists , surface scientists, and solid state physicists and chemists involved in research related to this topic. A total of thirty-one invited and contributed papers are

here presented, together with a summary report of the final panel discussion. These

proceedings are arranged to reflect, with some small exceptions, the program of the

Workshop. The main groupings of papers are: Theory of Electrocatalysis and Related

Topics; Characterization; Electrochemical Processes on Non-metallic Surfaces (1,2, and

3); Electrochemistry at Solid Electrolyte Interfaces. In addition, a final panel discussion

considered major problems, advances, and opportunities revealed by the Workshop.

Key Words: Catalysis; characterization; chemisorption; electrocatalysis; electrochemistry;

electrode processes; non-metals; photoelectrolysis; solid electrolytes; surfaces

Disclaimer:

In order to describe experiments adequately, it has been necessary to identify commercial

materials and equipment in this book. In no case does such identification imply recommendation or endorsement by the National Bureau of Standards, nor does it imply that the material

or equipment is necessarily the best available for the purpose.

iii
FOREWORD

The dissemination of research results is almost as important as the scientific work

that produced the results in the first place. As part of its role in making available to the scientific and engineering community fundamental data on the properties of materials,

the Institute for Materials Research of the National Bureau of Standards from time to time

holds symposia and workshops on various aspects of materials research. Such meetings

provide opportunities for exchanges of ideas and results, and the proceedings, when published,

often become benchmarks in the development of the field. Other institutions, inside or outside of the Federal Government, often join in various aspects of these meetings.

This Workshop, sponsored not only by the National Bureau of Standards, but also by

the National Science Foundation and the Energy Research and Development Administration,

was concerned with the topic of Electrocatalysis on Non-metallic Surfaces. It brought

together electrochemists, surface scientists, and solid state physicists and chemists with

a co.nmon interest in electron transfer reactions at interfaces. Indeed, it was a major

purpose of the Workshop to provide a meeting ground for these somewhat disparate groups.

The field of electrocatalysis entails study of the electronic structure and transport in the solid, the nature and chemistry of the interface on both sides, including adsorbed

layers, the disposition of the electrolyte and its components at and near the interface,

and the nature of electron transfer reactions through the interface. It is a truly interdisciplinary field, hence the emphasis on representation among the participants from the

several relevant discipl ines

.

The National Bureau of Standards takes pleasure in presenting these proceedings of

the Workshop on Electrocatalysis on Non-metallic Surfaces. This, too, is an opportune

place to express our thanks to those who contributed to the Workshop and these proceedings,

including the participants from the United States and those who came from abroad, and our sister institutions, the National Science Foundation and the Energy Research and Development

Administration, whose support was essential to this successful outcome.

John D. Hoffman

Director

Institute for Materials Research

National Bureau of Standards

iv
PREFACE

Electrochemistry as a science dates back almost two hundred years, to the work of

Galvani, Volta, Faraday, and others. In the course of development of this science a great

i deal has been learned about electrocatalysis and the nature of electrochemical reactions

on electrodes. Most attention, however, has been given to the development of highly

i reproducible metallic electrodes, such as the dropping mercury electrode or platinum.

What work has been done on semiconductors has been concentrated on again the well-known,

I reproducible material s--sil icon, germanium, zinc oxide, etc.

I Recently, new impetus for the study of electrode processes and electrode materials

has come in the wake of renewed interest in fuel cells for electricity generation. While

I traditional electrode materials such as platinum and its alloys work well (but not perfectly)

I in fuel cells, they are so costly as to be discouraging to commercial development. One

j remedy is to find new, less costly, materials that perform as well, or better, as fuel

cell electrocatalysts. Non-metallic solids offer a promising class of materials.

j Current development of fuel cells is concentrated on only a few types. The alkaline

j cell, while efficient and reliable as a result of extensive development for the use in the

l! space program, is presently excluded in the competition for utility application by its

jj sensitivity to CO2. Closest to commercial realization is the hot (150-200 °C) phosphoric

j acid cell, with the molten carbonate cell some years away and the solid electrolyte cell

still further behind.

|i

i At the air electrode in both acid and alkaline media, present catalysts (including

Pt) reduce the oxygen at least partially to peroxide rather than to water. The working

voltage for the peroxide reaction is lower than that possible with the water reaction, and

therefore, the achievable efficiency of conversion of chemical to electrical energy is also lowered. The development of catalysts for oxygen reduction, particularly in the acid

I cell, that avoid the peroxide reaction or minimize it could result in significant increases

in efficiency which may be necessary in the second-generation cells and beyond. It is interesting to note that some highly-conducting metal oxides {e.g., Ir02, RUO2) appear to catalyze oxygen reduction directly to water. Unfortunately, they are very unstable in acids, but may ultimately be useful with neutral (buffered) electrolytes.

In a very practical sense, the appliaation of fuel cells in the extensive way needed

in the not- too-distant future may well depend upon progress in the science of electrocatalysis as applied to non-noble metal surfaces, with particular but not exclusive

reference to the oxygen electro-reduction processes. This is a field at the boundary

between two disciplines, electrochemistry on the one hand, and solid state physics and

chemistry on the other.

Semiconductors have long been used as heterogeneous catalysts, especially for oxidation

and reduction reactions. Early uses of such materials as electrocatalysts included the

oxides of Cu and Ag. Silver has found extensive use at the oxygen electrode in alkaline

fuel cells; probably as an oxide. Lithiated NiO at temperatures above about 100 °C has

also operated successfully as an oxygen electrode in alkaline media. More recently, work

has been done using spinels such as C0AI2O4 and perovskites such as Sr-doped LaCoOs.

The latter use is obviously related to the use of the perovskites ABO3, where A represents

rare earth, Pb, or alkaline earth ions, and B is Co or Mn, for oxidation of Co and reduction of NOx in automotive exhaust catalysts. Some of the early, exciting non-metallic

electrocatalysts, for both hydrogen oxidation and oxygen reduction were the tungsten

bronzes and related compounds. However, more recent work has shown that the observed high

catalytic activity probably depends upon the presence of Pt, which in some way activates

the oxide in a synergistic fashion. A fairly detailed body of theory has been worked out

for electrochemical processes on semiconductors and insulators.

With this a background, the time appeared ripe for a Workshop on Electrocatalysis on Non-metallic Surfaces to review the most recent experimental and theoretical investigations,

to consider what are the problems most important for scientific progress in this field,

and to strive for some common understanding on how to approach those problems. A fundamental purpose of the Workshop was to bring together electrochemists, surface scientists.

V
and solid state physicists and chemists to stimulate as much interdisciplinary interaction

as possible.

The organization and running of the Workshop and the publication of these Proceedings

required the cooperation and enthusiastic support of a number of people. The Steering

Committee, composed of A. J. Bard, L. H. Bennett, U. Bertocci, M. W. Breiter, R. J. Brodd,

S. Bruckenstein, E. Cohn, R. deLevie, A. Fickett, H. P. R. Frederikse, H. Gerischer,

J. B. Goodenough, and W. S. Norton, provided invaluable guidance in the planning of the

Workshop. The Institute staff, particularly R. B. Johnson, R. F. Martin, and their

assistants, and Mrs. Sara R. Torrence and her staff from the NBS Office of Information

Services, performed yeoman service in the actual detailed running of the Workshop.

Mrs. Rosemary Maddock, as Format Editor, deserves much of the credit for gathering the

Proceedings into a coherent whole, and the IMR Text-Editing Facility for correcting and

preparing some of the manuscripts.

Alan D. Franklin

Technical Editor

vi
CONTENTS

Page

General Abstract iii

Foreword iv

Preface v

THEORY OF ELECTROCATALYSIS AND RELATED TOPICS

Electron Transfer and Electrocatalysis (Invited Paper)

H. Gerischer 1 Surface States, Chemi sorption, and Catalytic Processes on Transition Metal Oxides (Invited Paper Presented by T. Wolfram)

T. Wolfram, F. J. Morin, and R. Hurst 21 Hartree-Fock Calculations of Hydrogen Adsorption on Nickel

Copper, and Magnesium Oxide

A. B. Kunz, M. P. Guse, and R. J. Blint 53

Stark Model of the Excitonic Surface States Observed on MgO

H. J. Zeiger, V. E. Henrich, and G. Dresselhaus 59

Quantum Chemistry and Catalysis

R. P. Messmer and K. H. Johnson 67

Relationship Between Electronic Structure and Catalytic

Behavior of Transition Metal Carbides

D. R. Jennison, D. L. Klein, A. B. Kunz, K. M. Hall, and

W. S. Williams 87

The Role of Intramolecular Quantum Modes in Electrode

Reactions via Intermediate States

Yu. I. Kharkats and J. Ulstrup 93

Comment on the Papers of H. Gerischer, and of Yu. I. Kharkats

and J. Ulstrup

C. B. Duke 103

Reply to Comments by C. B. Duke on the Paper of Yu. I. Kharkats

and J. Ulstrup

J. Ulstrup 105

CHARACTERIZATION

The Study of Electrode Surfaces by Electrochemical Methods

(Invited Paper Presented by B. Conway)

B. E. Conway and H. Angerstein-Kozlowska 107

Electroref lectance Study of the Defect States in Reduced n-Type

Titanium Dioxide (Invited Paper Presented by P. Raccah)

S. S. M. Lu, F. H. Pollack, and P. M. Raccah 125

Electron Energy Loss Spectroscopy of Surface States on Titanium

and Vanadium Oxides

V. E. Henrich, H. J. Zeiger, and G. Dresselhaus 133

Dislocations in ZnO as Channels for Electron Transfer

from Surface Species

S. R. Morrison and T. Freund 139

Characterization of Semiconductor Electrodes in Nonaqueous

Solvents

S. N. Frank, D. Laser, K. L. Hardee, and A. J. Bard 149

Comment on Paper by S. N. Frank, et al.

V. E. Henrich 153

vii
Page

ELECTROCHEMICAL PROCESSES ON NON-METALLIC SURFACES - 1

Chemical Theory of Chemisorption and Catalysis on the

Surfaces of Non-Metals (Invited Paper)

R. L. Burwell , Jr 155

Electrocatalysis at Transition Metal Oxide and Carbide

Surfaces (Invited Paper)

P. Stonehart 167

Adsorption and Oxidation of Bromide Ion on Tin Oxide

Electrodes

H. Yoneyama and H. A. Laitinen 175

Cathodic Pyrophosphate-Lead Dioxide Electrode Reaction

Properties

C. 0. Huber, T. A. Berger, and R. E. Reim 183

Electrolysis by Foreign Me-^al Monolayers

R. R. Adzic, A. R. Despic, D. N. Simic, and D. M. Drazic 191 Comment on Paper by R. R. Adzic, et al.

B. D. McNicol 201

ELECTROCHEMICAL PROCESSES ON NON-METALLIC SURFACES - 2

Mechanisms of Electrochemical Reactions on Non-Metallic

Surfaces (Invited Paper)

E. Yeager

Photoelectrolysis of Water

J. G. Mavroides, D. I. Tchernev, J. A. Kafalas, and

D. F. Kolesar

Electrochemical Processes on Strontium Titanate Electrodes

D. M. Tench and D. 0. Raleigh

Electrochemical Activity of Polymeric and Monomeric Iron Phthalocyanines as a Function of Their Spectral Properties

A. J. Appleby and M. Savy ....

Electrocatalytic Activity of Phthalocyanines for Oxygen

Reduction

R. J. Brodd, V. Z. Leger, R. F. Scarr, and A. Kozawa . . Continued Growth of Anodic Oxide Films at Platinum and

the Mechanism and Catalysis of Oxygen Evolution

A. Damjanovic

ELECTROCHEMICAL PROCESSES ON NON-METALLIC SURFACES - 3

Electrochemical Surface Reactions on Non-Metals

(Invited Paper)

R. Memming 267

Oxygen Evolution on Semiconducting Oxides

A. C. C. Tseung and S. Jasem . 281 Mixed Oxides as Oxygen Electrodes

G. Singh, M. H. Miles, and S. Srinivasan 289

Surface and Electrocatalytic Properties of Tungsten Bronzes

M. F. Weber and H. R. Shanks 297

The Electrochemistry of Molybdenum Oxides

J. Horkans, M. W. Shafer, and R. Peekema 305

Titanates as Oxygen Electrodes

U. Bertocci, 1^. I. Cohen, J. L. Mullen, and T. Negas 313

203

221

229

241

253

259

viii
Page

ELECTROCHEMISTRY AT SOLID ELECTROLYTE INTERFACES

Ionic Transport and Electronic Exchange at Solid Electrolyte

Interfaces (Invited Paper Presented by M. W. Brecter)

M. W, Breiter and G. C. Farrington 323

Influence of Water on Beta Alumina Interfacial Ion Transport

G. C. Farrington 341 Overvoltage Phenomena in Solid Oxide Electrolyte Cells At Elevated

Temperatures

W. L. Worrell 351

PANEL DISCUSSION

Major Problems, Advances, and Opportunities Revealed by the

Workshop

A. D. Franklin (Reporter for the Panel) 359

ix

THEORY OF ELECTROCATALYSIS AND RELATED TOPICS

National Bureau of Standards Special Publication 455, Electrocatalysis on Non-Metallic Surfaces

,

Proceedings of Workshop held at NBS, Gaithersburg , Md. , December 9-12, 1975. (Issued ^fovember 1976)

Electron Transfer and Electrocatalysis

H. Gerischer

Fritz-Haber-Institut der Max-Planck-Gesellschaft

1 Berlin-Dahlem, West-Germany

Non-catalysed electron transfer is reasonably defined as an electron

exchange "between an electron donor or acceptor and an electrode under the

condition of very weak interaction between the reactants and the electrode.

The case of strong interaction between reactants and electrodes is analysed

with respect to its catalytic effect on electron transfer. It is pointed

out that strong interaction may often inhibit electron transfer for a single

electron transfer step at its individual equilibrium potential. Electrocatalysis is however found in multistep electron transfer reactions when

the adsorption of intermediates in the reaction path brings the individual

standard potentials of the single electron transfer steps closer to that

of the overall redox potential. Examples for this effect of adsorptive

interaction are given for metal and non-metallic electrodes. It is further

analysed where the electron energy levels of the adsorbed reactants are located with respect to the characteristic energy levels in the metallic or non-metallic electrodes and in the non-adsorbed species in solution. The

important role of surface states at a non-metallic electrode for chemisorption and electrocatalysis is emphasized.

Key words: Chemisorption ; electrocatalysis; electron transfer theory;

multistep redox reactions; semiconductor surface states.

1 . Introduction

This lecture is supposed to act as a bridge between the "Workshop on the Electron Faci; tor in Catalysis on Metals" and the "Workshop of Electrocatalysis on Non-metallic Surfaces"

by discussing the problem of electron transfer. However, there is a fundamental difference

between electron transfer in chemical catalysis and in electrocatalysis. In chemical catalysis, electron transfer can occur only during a transitory state and is finally compensated

by the reverse electron transfer, since the reactants and the products are electroneutral.

; But, in electrocatalysis the final products are obtained by single or multiple electron

' transfer and differ in net electric charge. One can control direction and rate of the cata-

'I lysed reaction, at least to some extent, by an external variation of the driving force with

I' help of the so-called electrode potential.

|, There is only one particular situation where electrocatalysis is equivalent with chemir cal catalysis, namely when anodic and cathodic electron transfer processes just compensate

I each other so that the external current is zero and the driving force of the reaction is

solely the free energy difference between reactants and products as in chemical catalysis.

; One can consider this special case as an electrochemical mechanism of normal chemical catai lysis. For instance, H202-decomposition in electrolytes in contact with metals can be described [1,2]-'- in this way.

1 This relation between chemical and electrochemical heterogeneous catalysis is demonstra-

!i ted in Fig. 1 for one special type of reaction. The common feature of both types of catalyT7. ^ . . Figures m brackets indicate the literature reference at the end of this paper.

If 1
A

solid I gas phase solid

heterogeneous catalysis

A * B A -B

electrocatalysis

A% B*+ 2e-^ AB

A% B + e- AB

A + B% e- — AB

Fig. 1. Comparison: heterogeneous- and electrocatalysis.

sis is the strong interaction hetveen the reactants and the catalysts in the intermediate

stage of the reaction where the interaction species are in an adsorbed state. The way in

which the reactants reach and the products leave this state is, however, different. The

electric forces play a minor role in chemical catalysis hut a dominating one in electrocatalysis. All analogies can therefore be only very limited.

2. Non-Catalysed Electron Transfer at Electrodes

We cannot understand electrocatalysis without knowing what we mean with an uncatalysed

electron transfer reaction. Let us consider as a reasonable definition of an uncatalysed

electron transfer a redox reaction where the interaction between the redox species and the

electrode during electron transfer is so weak that its influence on the energy of those

electronic quantum states, which are involved in the electron transfer, can be neglected. A statistical theory for adiabatic electron transfer under such weak interaction conditions

has been developed by R.A. Marcus [3,4]. A. quantum mechanical treatment of similar character

with perturbation theory for describing the electron transfer has been worked out by Levich

and his co-workers [5-7] . Both theories give the identical result that the activation energy

for electron transfer in the case of weak interaction with the electrode is primarily con- trolled by the reorganisation energy which is needed to reach a structure of the solvent a- round the electron exchanging species which is identical before and after electron transfer.

This type of reaction can be interpreted [8] , based on the early work of Gurney [9] as an electron transfer between real or virtual electronic quantum states of equal energy

which are filled at one side of the interface and vacant at the other side ^. Such a process

is represented by the following equation

( acceptor )„ + E

t occupied state] ( donor )^ + (vacant state )„ E -ili (1

This model neglects electron transfer with energy dissipation by excitation of vibrations

in neighbouring molecules or by phononemission , processes which are important in reactions

with release of large energies.

2
The reaction rate is controlled "by the probability to find quantum states of equal

I energy E during an encounter "between the electron acceptor or donor components of a redox

couple and the electrode. Such electron transfer can occur in parallel at different energy

levels. The net electron exchange rate must therefore he obtained from an integration over

all electron transfer steps at various energies. The general formula for such an electron

transfer rate is given in the following 2 equations

I oo

j"*" - [ K(E) D (E) ^

• (E) dE (;

I don J vacant don

—00

oo

I f ~ c [ K(E) D (E) • W (E) dE (: acc J occup. acc

—oo

where D(E) represents the density of states in the electrode and c •¥ (E) the energy

distribution of donor states in solution which is proportional to the donor concentration

1 c . The energy distribution of acceptor states is represented by c 'W (E) and is

5 pJSportional to the acceptor concentration c^^^ . ^^'^ ^'^^

For a calculation of the electron transfer rates one needs information on the distribution of filled and vacant electron states in the solid and in the electrolyte . The formulation for the solid can be taken from band theory and Fermi distribution statistics. For

the electrolyte it has been derived that the distribution of electron energy states is of

the Gaussian type if one takes into account only the square terms [10] . The maxima of the

distribution functions represent characteristic energy levels for the acceptor and the donor

of every individual redox system [11]

.

The position of these characteristic energy levels can be derived from a thermodynamic

cycle. Fig. 2 shows in a very simplified way the influence of the interaction between the

I donor or the acceptor states and their surroundings in an electrolyte (the solvation shell

j primarily) on the position of the energy level for electron exchange. Assuming that elecedox

reorganisation coordinate reorganisation coordinate

Fig. 2. (a) Energy profile for electron transfer under Franck-Condon conditions between redox

species and electrons at vacuum level;

(b) the same for electron transfer between redox species and electrons at the Fermi

level of an electrode being at equilibrium with the redox couple

.

tron transfer occurs under Franck-Condon conditions, the energy level for electron exchange

is then given by the distance between the two energy curves of the acceptor and the donor

state at constant reorganisation coordinates. The redox reaction is formulated in part (a)

of the picture as an electron transfer between the redox species and the electron at the vacuum level. In part (b) the electron transfer occurs at equilibrium conditions between

the redox species and the electron in an electrode. This means the electron is at the Fermi

level of the redox system E ^ which is equivalent to the Fermi level of electrons in an electrode in equilibrium wi?H ?liis particular redox system. The correlation between the

electron energy scale related to the vacuum level and the electrochemical potential scale is

[12]:

E^ = Const. - U

F o

(3)
with Const. ^ - H.5 eV.

The interaction with the surroundings can "bring "both species to the same electron

energy as shown in part ("b) of Fig. 2 in the point of intersection. This solvation struc- ture characterizes the intermediate state which has to "be reached for adiabatic electron

transition. The energy necessary to reach this state is the activation energy for electron

transfer under Franck-Condon conditions. The result is that the minimum activation energy

between the redox species and their surroundings is, in the parabolic approximation of the

energy surfaces, one fourth of the reorganisation energy which is the energy released

after electron transfer under Franck-Condon conditions from the initial states with minimum

free energy (compare Fig. 2b). The distribution function of the electron exchange levels of

redox species in electrolytes has in this simplest case the form [10]

,

W(E) [k-n E kT) ^'^^exp K (E-°E)^

i+E^kT n

The translation of this model into the formulation given in equ. (2) for electron

transfer between quantum states at both sides of the interface is shown in Fig. 3. In the

vacant

filled

J = j = jc

D(E) W(E] j(E) «(D W)

Fig. 3. Energy distribution of electronic states at equilibrium between a metal electrode

and a redox couple ; energy spectrum of electron exchange rates

.

equilibrium situation, the electron transfer occurs preferentially in the close neighbourhood of the Fermi level with equal rate in both directions. If one applies another electrode potential the result is a net current in the one or the other direction. This is described in this model by a shift of all energy levels in the metal electrode relative to the

Fermi level in the redox system. It has the consequence that now the electron transfer

rates from the electrode to the redox system or reverse do no longer compensate each other.

If the shift of the Fermi level goes far enough, the electron transfer occurs only in one direction as shown in Fig. h. All these theories lead to the long known "Volmer-Butler equation for the rate of redox reactions in dependence on the applied overvoltage, n = U-U^ [8].

cathodicequilibr.

anodic

" ^Redox '^^'^

"E Redox -6o11q

D(E) W{E) j(E)

Fig. h. The same as in Fig. 3 for anodic (Ha) and cathodic (He) polarisation of the electrode

against the Fermi level of the redox couple

.

4
The characteristic kinetic parameter of an electron transfer reaction is the exchange

current density at the standard potential of the redox system. This is mainly controlled

by the reorganisation energy. A general experience is that one electron transfer steps

are normally fast, E heing in the order of 0.8 - 1 .6 eV if no kinetic complication arises.

This means one electron transfer reactions cannot he accelerated very much by catalysis.

We shall see later that in correspondence with this the main field of electrocatalysis is

multiple electron transfer reactions.

3. Electron Transfer with Strong Interaction

between Electrode and Redox Components

In the uncatalysed electron transfer reaction the electrode plays only the role of a source or sink for electrons. What happens with regard to the electron transfer rates when

now the reactants interact strongly with the electrode? We shall be inclined to consider this

situation as a case of catalysis and expect that the rate of electron transfer will increase.

But we shall see that strong adsorption of the redox species can in the range of the equilibrium potential also have an inhibiting effect on electron transfer.

For so close an approach to the electrode that strong interaction becomes possible, the

j redox species have to enter the region of the Helmholtz double layer. The interaction with

i the surroundings compared with the situation in solution is then drastically modified in two

ways

(a) A great part of the interaction which was prior directed to the solution

is now replaced by the interaction with the electrode. This applies as well to the outer solvation shell as in many cases to the interaction with

the inner solvation shell. Some ligands may be lost in compensation for

strong adsorption.

(b) By entering the inner region of the Helmholtz layer the electrostatic forces

become important for all redox species having an excess electric charge or a dipole moment. This makes the situation very complicated and difficult for

a quantitative treatment of electron transfer kinetics. Therefore, we shall

discuss the consequences only qualitatively.

One consequence is that the polarisation interaction is no longer the most important

coordinate for electron transfer. The distance from and even the particular site on the

surface become the more important the stronger the interaction energy. A multidimensional

representation of energy siirfaces becomes unavoidable. We shall discuss in the following

^ three different cases distinguished by the degree of interaction of the acceptor and the

; donor state of the redox system with the electrode.

Fig. 5 shows the situation where, at equilibrium potential, both redox species are adsorbed at the electrode with about equal adsorption energy. In part (a) the energy of both

sides of a redox reaction is plotted against the distance of the redox species from the electrode with the electron energy level at the Fermi-energy of the redox system (the equilibrium

condition). From this picture we obtain an energy for desorption and adsorption at equilibrium. For calculating the activation energy of electron transfer we need, however, a

I m\altidimensional representation, since we have to take into accoimt also the different inter1 action energy of the redox species with their surroundings. We must expect that the interi action with the surroundings in the adsorbed state of a redox system is much less different

(for the two components than in solution. This is indicated in part (b) of Fig. 5 where the

influence of the reorganisation coordinates on free energy is plotted for the redox compoI nents at 3 different distances from the electrode. The conclusion is that the activation

barrier for electron transfer in the adsorbed state is much smaller than in solution. The

electron exchange rate in the adsorbed state may often be so high that the different oxidation states become indistiguishable as has been pointed out already by Plieth and Vetter

[13] . The result is that in case of equal adsorption of both components the rates of desorption will become the rate controlling steps for the redox reaction, as Fig. 5a indicates.

A variation of the electrode potential changes the energy position of the reactants in

5
E

electrode

equilibrium

/ (-C + Bco

/ ^ Red ox

dad dad* e§oRedox

Helmholtz d electrolyte

double layer

acccni+epo

-Redox aCC (x<d) ac"c"ad-^e-E^edox. reorganisation coordinate

(c)

acc*+e(EF>E°i

^ cathodic

/^ dOHsol

- —: — anodic

-t- electrode Helmholtz d electrolyte x double layer

Fig. 5- Energy profiles for adsorption of both redox components at various electrode

potentials in dependence on the position in the double layer

(a) at equilibrium potential;

(b) energy profiles in dependence on the reorganisation coordinate for 3 positions in the double layer at equilibrium potential;

(c) for anodic and cathodic polarisation.

solution relatively to each other, and the energies for desorption shift in different direction. This is shovn in Fig. 5c for an anodic and a cathodic overvoltage. How much the

electrode potential influences the energy of both components in the adsorbed state depends

on the double layer structure, the charge of the adsorbed species and their position in the

Helmholtz double layer. It dependes also on the strength of the interaction vith the substrate. In case of very strong interaction, the electron exchange levels of the adsorbed

species can be treated like a surface states of a solid. This case is closest to normal chemical catalysis.

In most real systems with strong adsorption, the one or the other component of the redox

system has large preference in adsorption. This will usually be the electroneutral component

which is the stronger adsorbed one because electrostatic interaction would in any case prevent too high coverage of the surface with charged species. Strong adsorption of charged

molecules involves other complications connected with a change in double layer structure

which cannot be discussed here.

The situation for an uncharged donor species being stronger adsorbed is shown in Fig. 6. The

surface at equilibrium potential in this case is widely covered with the donor and the ex- change rate between the adsorbed state and the state in solution will be slow. Electron ex- change in the adsorbed state will not occur because of the too different adsorption energies.

If electron transfer is not possible via tunnel transitions between the electrode and the

redox components in solution, the electron exchange rate at equilibrium potential will be

very slow. It needs a high anodic polarisation to reduce the barrier for oxidative desorption

of the reduced species where the electron transfer occurs somewhere on the path of desorption.

This is shown in Fig. 6. If polarized in cathodic direction, the surface will be more and more covered with the reduced product, the desorption of which remains the rate determining

step. It is ass\imed in Fig. 6 that the reduced species has no electric charge and that the

desorption energy therefore does not depend on electrode potential. We fin^ such a situation,

for example, in the first step of the reduction of protons from solution, H + e —>

^ad' where the hydrogen atoms are strongly adsorbed at all electrodes and desorplion of H-atoms^

6
electrode

electron

-transition

acc*+e-(Ep>E°Redox) cathodic

equilibrium

anodic

double layer

Fig. 6. Energy profiles for preferential adsorption of the reduced component

of a redox couple at various electrode potentials.

does not occur. That in this system the reaction proceeds further is only due to consecutive

reactions generating the hydrogen molecule. We shall come back to this reaction later.

The opposite case v?here the acceptor is the uncharged species and much stronger adsorbed,

is shown in Fig. 7- At equilibrium the electrode is now covered by the adsorbed acceptor

component, the desorption rate of which is very slow. A variation of the electrode potential

causes in this case a drastic shift of the energy of the acceptor in the adsorbed state plus

electron in the electrode and also of the acceptor in solution plus electron in the electrode,

as is shown in part (b) of Fig. 7. Since the donor state has an electric charge its energy

varies also in the adsorbed state in correspondence with the local potential in the double

layer, but to a smaller extent than for the acceptor state.

By applying a cathodic voltage a reductive desorption of the donor can occur connected

with an electron transition occurring at some distance from the electrode . For anodic

polarisation, we see in Fig. 7b that the oxidation of the donor will occur very rapidly in

the adsorbed state but the desorption of the product is slow and rate determining. A situation

like this of Fig. 7 is found in the reaction OH < ^ OH + e , since the OH radicals are

3.CI

(a)-

Redox

double layer

acc + e"(Ep>E°)

cathodic

anodic

double layer

Fig. 7. Energy profiles for preferential adsorption of the oxidized component of a redox couple

(a) at equilibrium potential

(b) at anodic and cathodic polarisation.

7
strongly adsorbed on nearly all substrates. A reaction progress in anodic direction is again

caused only by consecutive steps. To complete our picture in terms of electronic energy levels, we have to discuss where

the electronic states in the adsorbed species at the electrode are found energetically. We shall begin with a thermodynamic argument and ask what is the standard redox potential for

the redox system in the adsorbed state (standard state: both components present at equal

concentration). If the interaction with both oxidation states of the redox couple is equal,

then the standard redox potential of this system in the adsorbed state will be the same as in solution. This means that the electronic energy levels for electron exchange must re- main in the range around the standard Fermi level Ej^^jq^j. defined for the system in solution. We have previously seen in Fig. 3 that the most probable energy level for the donor

state of the redox system is located at E

x~'^R most probable acceptor state at E°^^+Ep. We expect a narrowing of the distance between the most probable energy levels

for Doth oxidation states in the adsorbed position because the interaction with the solvent

has decreased. The consequences are shown in Fig. 8 as a comparison between the energy

distribution of the electron levels in solution and in the adsorbed state. If the electron

exchange in the adsorbed state is slow enough that two different oxidation states can be

distinguished, we shall still have two maxima for the adsorbed species in their different

state of charge (Fig. 8b). If the electron exchange with the electrode is so fast that both

states become indistinguishable, we shall find a situation as shown in Fig. 8c.

The broadening of the electronic levels in Fig. 8c for strongly adsorbed species is

mainly caused by the interaction between the manifold electronic states of the solid and

the electronic states of the adsorbate. This picture is equivalent to the assiamptions in

the Anderson model for chemisorption of atoms on metals [14,15]. To some extent the broadening into an "energy band" will also be caused by fluctuations in the remaining interaction

with the solvent. The occupation of electronic levels in this "energy band" of the adsorbed

species is controlled by the position of the Fermi level. Since the location of adsorbed

species extends into the region of the Helmholtz double layer, the position of the Fermi

level in this band will be shifted to some extent by a variation of the electrode potential,

which causes a variation of the average charge on adsorbed redox species with the applied

voltage.

W(E) W{E) W(E)

Fig. 8. Distribution of electron exchange energy levels of a redox couple with about equal

interaction between both redox species and the electrode

(a) in solution; (b) in adsorbed state with medium interaction energy;

(c) in adsorbed state with strong interaction energy.

The electron level distributions for the two cases with very different adsorption energie

of the donor or the acceptor species are schematically represented in Fig. 9 for electronic

equilibrium. In the case of strong adsorption the electronic energy levels of the adsorbate

are practically part of the electron system of the electrode. Whether they are occupied or not in the adsorbed state depends on their relative position against the Fermi level in the

electrode. If the reduced (donor) state is stronger adsorbed than the oxidized (acceptor)

state, the standard Fermi level for the adsorbed species will be far below the Fermi level
'^F= ^Redox'

tad

(a)

Fig. 9. Distribution of electron exchange energy levels in the adsorbed state of a redox couple with (a) preferential adsorption of the reduced state;

(b) preferential adsorption of the oxidized state.

of the redox system in solution and vice versa for a preferential adsorption of the

acceptor state at the electrode. It is obvious that in the latter case the electron state

of the acceptor species which controls electron transfer (the redox orbital) is not involved

in the adsorption interaction. This must be caused by other electron states at deeper

energies which intercombine with electron states of the electrode

.

Fig. 9 shows such a hypothetical energy term distribution for electrodes in equilibrium

contact with two different redox systems. In case (a) the stable adsorbed state is the occupied one, in (b)it is the vacant state. The standard redox potential in the adsorbed state of

the redox system may be defined as the energy level where the electron terms of the adsorbed

species would be half- filled. This is °'Eq^ in Fig. 9, which will shift nearly in parallel

with the electrode potential relative to tlie Fermi energy of the redox system in solution as does the Fermi energy of the electrode.

The small maxima above E in Fig. 9a and below in Fig. 9b indicates a slight

stabilisation of the differently charged states by the fluctuating interaction with the polar

solvent to which the adsorbed species are partially exposed. This contribution will, however

be of very little importance at strong adsorption.

k . Catalysis of Electron Transfer in Multi-Step Redox Reactions

We have explained that adsorption in a one-step electron transfer reaction usually does

not increase the reaction rate in the neighbourhood of the equilibrium potential but may

often cause the contrary. The interesting systems for electrocatalysis are therefore those

where the reaction is slow because more than one electron transfer step is needed in the over

all reaction and intermediates are formed which may interact strongly with the electrode.

A few examples may be given.

reaction

2H + 2e

possible intermediates:

H

+ 2H + he 20H

2NH. + 6h + 6e

CH^OH + HgO CO^ + 6H + 6e

Og, HO^, OH, 0 , 0

NH^, NH, N^H^, N^Hg, N^H, N, H

H, CH^O, CHgO, CHO, CHOOH, COOH

The common feature of all these reactions is that some intermediates with a high energy

have to be generated before the reaction can proceed to form the final products. An easier

generation of such intermediates due to interaction with the electrode will increase the

reaction rate in such multi-electron step processes, and this is the general reason of electrocatalysis in all systems of practical interest. For simplicity, we shall discuss this

9
for a reaction with two electron transfer steps. The conclusions are valid for all systems

independent of the numher of electrochemical and chemical reaction steps

.

Overall reaction: a"*" + 2e <

'' B standard potential: ^^"""/g" (5)

reaction steps: a"*" + e I

"^a"*"/!

I + e~ ^==^ b" E°^g- (513)

where the intermediate I acts as an electron donor in the first reaction and as an electron

acceptor in the second one. If we define the standard redox potentials for the two different

electron transfer reactions as usual (with standard concentrations of the species in solution] we have from thermodynamics a simple correlation between the redox potential of the overall

reaction and the redox potentials of the single steps (the principle of Luther [16]).

+ EI/B-) (6;

For a reaction with a chemical reaction step involved like;

overall reaction:

1st electrochem. step A + e

chem. step

2nd electrochem. step I + e

2

A + 2e + C ^==B

+ -

I."1

^A^, C/b"

AG

chem.

(7)

(Ta)

(Tb)

(7c)

this principle gives;

"'A ,C/B A /I

= E°+,^ + AG° + E°

chem -^2^^ (8)

A demonstration of this thermodynamic relation in terms of the position of electron-exchange energy levels for the two respecti;^e redox reactions of equ. (6) is given in Fig. 10a

and 10b. For starting the reduction of A in solution one has to apply an overvoltage in the

order of the difference between the standard redox potentials of the overall reaction and the

first electron transfer step. The second reaction will follow very fast because the acceptor

state of the intermediate is at a much lower energy level having a very high electron affinity.

If, however, the intermediate is strongly adsorbed, the energy levels in the adsorbed

state are shifted for I as a donor into negative direction of the electron energy scale and

for I as an acceptor into the opposite direction as shown in Fig. 10c.

This means, thermodynamically , that the standard redox potential of the first step is shifted in the adsorbed state to more positive values; that is a more negative Fermi level in our absolute energy scale of Fig. 10. The standard redox potential of the second reaction is -Shifted into the opposite direction (see Fig. 10a). The consequences of adsorption of such

an intermediate for the kinetics are obvious. The reaction can proceed now in cathodic and

in anodic direction at much lower overvoltages measured against the standard redox potential

of the overall reaction, because the intermediates can be formed with much less excess energy.

We see that the shift of electronic energy states due to adsorption can have an enormous catalytic effect on the reaction rate of electron transfer reactions with two or more one- electron transfer steps of very different standard redox potentials. This is a common situ10
-EAVI

E(AVl)ad

Eavb-

"En/B-)ad

/B'

'^ZZZ2ZZ2 Isol, donor

(b)

W(E) W(E)

Fig. 10. Two-step redox reaction A + 2e «=^B with intermediate I. (a) Standard redox potentials in absolute scale (= Fermi levels);

(Td) distribution of energy levels in solution;

(c) distribution of energy levels in the adsorbed state.

ation in the kinetics of organic electrode reactions , where radicals are formed as intermediates. The classical example, however, is the hydrogen electrode reaction which we now

shall discuss very briefly from this thermodynamic point of view.

Let us assume that the hydrogen evolution reaction follows the electrochemical desorption

path. In the absence of interaction with the electrode, we have the following situation for

the overall reaction and the single reaction steps:

2H + 2e sol 2, sol

H + e ^=^¥. , sol sol

H + H , + e ^==*H^ , sol sol 2, sol

+ 2.1 V

O 2

(9)

(9a)

(9b)

H , + H , = H- , sol sol 2, sol

with AE° = E° - E°^./j^^ (9c)

The two standard redox potentials of reaction (a) and (b) are shifted by adsorption of H-atoms at the electrode by an equal amount into opposite directions. They coincide when the

free energy of adsorption of H-atoms AG^ just compensates the differences of hE^+ This

is shown in Fig. 11 by plotting the standard Fermi energy of the adsorbed redox systems as function of AG° We see that the evolution of hydrogen by electrolysis would for AG^^ = 0 need an enofmous negative overvoltage to reach the first intermediate . Only because

hydrogen atoms are strongly adsorbed can the reduction start at much lower cathodic potentials.

The first step remains rate determining for the cathodic process as long as the second step

does not reach a more negative redox potential than the first one. This is seen in Fig. 11 where an inversion of E for the two partial reactions occurs when

h^H,ad h i I ^^2H-^hJ

•

11
EhVHsoI ,2 t AE

2H*/H2

sol "ad

Had*HVH2

Fig. 11. The hydrogen electrode; dependence of the standard redox potentials

(Fermi levels) for the partial reaction steps on free energy of

adsorption at the electrode

.

When this point is reached, the formation of the adsorbed H-atoms gets fast, but the desorption reaction becomes slow, a well known idea of the mechanism for the hydrogen generation

at electrodes with very high adsorption energies for hydrogen atoms [17,18].

A general conclusion is that the highest rates or the optimal catalysis can be expected

when the Fermi energies for all single electron transfer steps coincide approximately. The

real kinetic picture which has been given by Parsons and by myself at an earlier occasion

[17,18] is more complicated than this thermodynamic discussion can show. For a more specific

discussion of the kinetics one has not only to use the exact rate equations and their dependence on electrode potential but also to include the recombination step. The driving force

for this step can be taken from the scheme of Fig. 11 as the difference between the 2 Fermi

energies of the single redox reactions, but the kinetics of this step are quite different

from electron transfer steps. ^This means the optimum can be shifted to a somewhat different

adsorption energy of H , than 77 AG but this is only a minor correction,

ad. d dtr^n^ The reduction of oxygen via the hydrogen peroxide path or the oxide path and the electrolytic oxygen evolution can be discussed analogously [19] . In such a complex reaction with

four electron transfer steps, however, the adsorption of different intermediates can have a compensating effect. This is demonstrated in Fig. 12 for the peroxide path. One can derive

from such a figure what kind of interaction at the surface should be favourable and what not. The optimum will again be obtained when the standard potentials of all single electron transfer

steps approximately coincide with the standard potential of the overall reaction. Fig. 12 shows

that for this mechanism we would need the following free energies of adsorption to approach

this optimum;

OH

sol OH

ad' AG = -1 .6 eV; 0 0. 2, sol 2, ad

AG = -1 eV

HO ^ HO

2, sol 2, ad' AG = -0.6 eV

The oxygen electrode reaction can contain, however, different intermediates than here assumed,

for instance, the formation of 0 ^, O^H , which would need a different optimisation of adsorption energies if such a reac?'ion pal^ should be more favourable. For real kinetics, such

12
^Enhe

+ 1 -

02*e-= 0"

2 102,od

O2* H20+2e-=OH"* HO'z-—

H02 + H20 + e-=OH*2 0K|-__j_H02,ad

lOHad

O2 + 2H20Ae- = 4 0H" 02^H20 +e-= HO2 +0H"

HO2 H2O +2e-= 3 0H"

-1 -

OHod

-2- OH e-= OH"

Fig. 12. The oxygen electrode; dependence of the standard redox potentials

(Fermi levels) of four possible electron transfer steps on the free

energy of adsorption of the respective intermediates O^, HO^, OH,

assumed in one particular reaction path.

a thermodynamic consideration can again only give the general hackground, but no reliable

prediction of rates.

In o\ir previous discussion we have assumed a metal-like behaviour of the electrode with

the Fermi energy located within a continuous energy band. A variation of the electrode potential in this case changes the voltage drop in the double layer between the electrode surface and the electrolyte but has no effect on the charge distribution in the bulk of the electrode. For non-metallic electrodes this can be quite different. The electronic system there

contains a band gap which means that a smaller number of electronic quantum states will be

available for exchange interaction with adsorbed species. Therefore, the adsorption interaction will normally be weaker than on metals, if there is not a specific bond formation

due to interaction between localized electron states in the surface (surface states) and the

adsorbate

.

We shall discuss now a non-metallic electrode with the Fermi energy in the band gap, but

high enough conductivity to be usable for electrolysis. Non-catalysed electron transfer can here be described in the same way as for metals but with the restriction that all energy

levels of a redox system which correspond to energies in the forbidden gap of the semiconductor are excluded from electron exchange [20] . It is therefore very important to know where the

electronic energy levels of a redox system are located relative to the energy levels of the

band edges in the surface of a semiconductor. Fig. 13 gives a survey of typical situations

for different redox couples in contact with a semiconductor. For simplicity the semiconductor is assumed to be at a particular electrode potential, the so-called flat band potential,

where we have no excess charge in the bulk of the electrode. In Fig. 13a, the redox system I has a redox Fermi level above the flat band Fermi level E^g of the semiconductor , but below

the conduction band edge. The consequence is a slow rate of electron injection into the

conduction band from thermally excited levels of the reduced species which have reached the

energy range of the conduction band. The second redox system II of Fig. 13a has the redox

Fermi level below F^^, but above the valence band edge. The consequence is an electron transfer from the valence band to acceptor states of the redox system which have reached the

5. Electron Transfer at Non-Metallic Electrodes

13
Efb-

(a)

W(E)

electron

injection

1

V ///////.

(b)

1

e"

III

'ni

Red IV hole

inject. W(E)

Fig. 13. Energy distriljution of electron exchange levels for various redox systems in contact with a semiconductor at flat band potential. Indication of electron transfer rates

"by arrows

.

energy of the valence band by thermal activation,

into the semiconductor.

This is equivalent to a hole injection

In Fig. ISh we see the situation where the Fermi levels of the redox systems are either

in the range of the conduction band or in the range of the valence band. The consequence is a fast electron injection into the conduction band or a fast hole injection into the valence

band respectively. While the electron transfer rates in the situation of Fig. 13b are high,

comparable with metals or even higher, the rates of electron transfer in the situation of

Fig. 13a are much smaller than for metals. If the band gap is wide enough and the Fermi ener- gy of the redox system is located in the middle range of the band gap, electron transfer can practically be fully prevented.

The influence of the electrode potential on the relative position of the electron energy

levels can vary to a large extent. It depends on the concentration of mobile electric charge

carriers in the semiconductor and on the presence of surface states, their concentration and

energetic position^

g In^systems with a wide band gap and a not too high donor or acceptor

concentration (<10 cm ) and in the absence of surface states, a potential variation affects

primarily the charge distribution in the semiconductor and can be described by an equivalent

band bending. This means for a redox system with a Fermi level ^ closer to the conduc-

. Redox tion band than to the valence band that a more cathodic electrode potential will shift the

Fermi level in the semiconductor upwards and cause an increase of electron density in the

surface of the semiconductor. In this way the reduction via electron transfer from the con- duction band to the vacant levels of the redox system is accelerated as indicated in Fig. lUa

The concentration of electrons in the surface and the density of vacant electron states of

the redox system at the band edge control the rate.

(a) cathodic

I'E

W(E)

Ec

Ep

(b) anodic W{E)

Fig. 1^. Influence of cathodic and anodic polarisation on electron transfer

between a redox couple and a semiconductor with preference of electron

exchange in the conduction band.

14
Polarization in the anodic direction lowers the Fermi level in the semiconductor but leaves

the oxidation by injection of electrons into the conduction band at a nearly constant rate as long as no holes are generated in the surface of the semiconductor (compare Fig. 14b). Since

semiconductors with a wide band gap can usually not form an inversion layer, the oxidation by

means of holes normally plays no role at an n-type semiconductor with a wide band gap.

For p-type semiconductors we have the inverse situation and normally all electron transfer steps occur via the valence band. At anodic polarisation, the concentration of holes in

the surface increases and the current grows in parallel. At cathodic polarisation, the

current is limited by the injection rate of holes and remains constant as long as no electrons

are generated in the' conduction band.

6. Catalysis of Electron Transfer at Semiconductors by Surface States

Semiconductors with a wide band gap are normally uninteresting for catalysis because

they not only imply a high barrier for electron transfer, but they are normally also

inactive for chemisorption. This can be different by two reasons if surface states are present. The surface states can be energetically located within the range of the band gap

with the result that charge can be accumulated in the surface forming a variable countercharge for a Helmholtz double layer. The variation of charge in such surface states with

the applied voltage makes the semiconductor behave like at a metal electrode with regard to

the double layer behaviour. The other consequence is that surface states can specifically

form bonds with adsorbates , which can catalyse electron transfer.

If surface states interact with the components of a redox couple, charge can be easily

exchanged with the electrolyte by alternating adsorption and desorption of the redox species.

To complete the charge transfer with the bulk of the semiconductor, however, the coupling between bulk states and sirrface states must be strong enough. This condition is often not fulfilled for semiconductors with a wide band gap and a low concentration of donor or acceptor

states, because the electrons cannot pass the energy barrier between the electron surface states

and the bulk states if the barrier is high and too extended. With a high enough concentration

of mobile charge carriers, however, the space charge layer becomes thin and transparent for

electrons via tunnelling [21] . Electron transfer is also possible to a more or less extent if

the energetic position of the surface states is close to the band edges and therefore the

barrier height for electron transfer between siirface states and the bulk states remains low.

A variation of the electrode potential in these cases mainly changes the concentration of excess charge in surface states but only slightly the charge distribution in the bulk. This

has been shown in some experiments with artificially created surface states [22] .

Fig. 15 characterizes four situations of an n-type semiconductor with surface states

close to the conduction band edge being in contact with the electron states of a redox electrolyte. It shows their relative energy position and also schematically the energy spectrum for

electron transitions between the redox electrolyte and the electrode in these four cases. Electron transfer occurs in this system primarily via the surface states of the semiconductor and the relative shift between E^ and ^

is mainly caused by a variation of the potential drop in the Helmholtz double layer with the voltage applied.

For a p-type semiconductor with surface states around the valence band edge one finds

the analogous situation as shown in Fig. l6. Electron transfer is described in this figure

as hole injection and extraction into or from the semiconductor occurring via the surface

states.

Fig. 17 demonstrates the influence of surface states, with energies located in the

band gap, on electron transfer in the absence of strong interaction between the reactants

and the electrode. Continuous current flow is in such systems only possible if the potential

barrier between the bulk states and the surface states is so thin that tunnelling is efficient,

which needs a high donor or acceptor concentration [21] . The previous discussion of electrocatalysis has, however, shown that only strong interaction between intermediates of a multistep redox reaction and the electrode can make an important contribution to an increase in

overall reaction rate. Therefore, we shall now consider the consequences of strong interaction

between a non-metallic electrode and the intermediates of redox reactions.

15
(a)

- E Redox

-X W(E)

t E

7777777^

(b) -X W(E)

- X W(E)

Fig. 15. Correlation of electron exchange energy levels for redox couple in contact with an n-type semiconductor having surface states around the conduction

band edge. Situation at various potentials; (a) flat band (cathodic)

,

(b) equilibrium, (c) strongly cathodic, (d) anodic.

- X W(E)

i^/7/7//7W

(b) - X W(E)

W(E)

Fig. 16. Correlation of electron exchange energy levels for redox couple in contact

with a p-type semiconductor having surface states around the valence band

edge. Situation at various potentials; (a) flat band (anodic), (b) equilibrium, (c) strongly anodic, (d) cathodic.

16
tunnel

transition

cathodic (n -type)

Ef _ _

tunne

trans

cathodic (p -type)

tunnel

trans

anod ic ( n - type)

tunnel

trans.

anodic (p - type)

Fig. 17. Electron exchange levels of redox system in contact with an n-type

or p-type semiconductor having surface states in the band gap energy

range; various states of polarisation.

We discuss the hypothetical intermediate I of equ. (5) which can act as the electron

donor state in reaction (5a) and the electron acceptor state in reaction (5l3)- The electron

exchange term for the first step may have a rather negative standard redox potential E^"*" /j for the non-catalysed redox process in solution; the standard redox potential of

the second step may be rather positive in the electrochemical scale. Fig. 18 represents this

situation for a contact "between two semiconductor electrodes at equilibrium with the net redox system. For simplicity, this equilibrium potential is assumed to coincide with the flat

band potential. E.+ may be located in the range of the conduction band, E^ in the range

of the valence bana. This means, according to Fig. 13, that electron injection from I (oxidation) into the conduction band would be fast and also hole injection from I (reduction)

into the valence band. The intermediate is therefore not stable at equilibrium conditions

as already indicated by the large difference between E°+^^ and E^^.^^-. The overall reaction

(5) will therefore be blocked in both directions as long as one does not apply high

overvoltages

.

EcEf

Aad

'ad, donor — E A%2e-=B"

Ev ////////| lad. acceptor

Bad

Ec

Ef

;b)

A*ad

K~^AVlad

'ad, donor

\ ^ad, acceptor

lod/B

Bad

I/B'

Fig. 18. Electron exchange levels of redox couple with 2 electron transfer steps

and one intermediate broadened by interaction with a semiconductor surface.

(a) no surface states on the semiconductor.

(b) with surface states in the energy range of the band gap.

17
The semiconductor of Fig. l8a is supposed to have no surface states. It is to he ex- pected in this case that interaction is weak because the wave functions of the electronic

semiconductor hulk states decay very rapidly in front of the surface. A weak interaction

is indicated in Fig. l8a by a small broadening of the electron exchange terms for the intermediate acting as a donor state (reaction 5a) and as an acceptor state (reaction 5b).

In the presence of surface states the interaction will be much stronger and results in

a wide broadening of the electron exchange levels. This is schematically shown in Fig. I8b

where the surface states are supposed to form a continuous band over the range of the gap. The principal consequence is that the donor and the acceptor levels of the intermediate I come much closer together and approach the range of the standard Fermi energy of the redox

system. It has been assumed in Fig. l8b that this range is not yet reached, and therefore,

the exchange current at equilibrium potential of the overall reaction will still be negligible. If the interaction is so strong that the energy range of is reached in the adsorbed state, the exchange current starts to increase, and this means efficient electrocatalysis

.

In the absence of surface states, as in Fig. l8a, a polarisation of the electrode at medium overvoltages will only vary the band bending but not the relative position between the

band edges at the surface and the energy levels of the redox system. Only if degeneracy

of electron states is reached in the surface, the system' begins to behave like a metal

electrode, and the potential drop in the Helmholtz double layer increases until electron

or hole injection become possible. This means a very high overvoltage, much higher than

for metal electrodes.

The much more promising situation is to be expected in presence of surface states which strongly interact with the electronic bulk states and the intermediates . A slight

polarisation of the electrode is then sufficient to bring the Fermi level to the range of

electronic levels of the reactants, as is schematically shown in Fig. 19, because now the amount of charge accumulated in the surface varies with the applied voltage. Such a variation of the voltage drop between the surface and the electrolyte causes a relative shift

between the energy levels of the electrode and the solution similar to a metal electrode, as indicated in Fig. 19. Electron transfer starts in cathodic direction when the Fermi level

in the surface states reaches the acceptor levels of the oxidized species because the acceptor levels of the intermediate can then pick up electrons from the lower filled states of

the surface (compare Fig. 19a). The respective situation at anodic polarisation with electron injection into surface states is shown in Fig. 19b.

This discussion has shown that surface states can catalyse electron transfer at non-metallic electrodes in both cathodic and anodic directions for multiple electron transfer

reactions in the same way as electrocatalysis acts at metal electrodes. Is there to be ex- pected any advantage from the use of nonmetallic electrodes? One favourable possibility is

~ Ea*/i - ^AVIad

s-yiad.donor

^ad.acc.

anodic

Fig. 19. Electron exchange levels of a 2-step redox reaction with strong interaction

between the intermediate and surface states in the energy range of the band

gap. Demonstration of electron transfer catalysis in cathodic and anodic

direction by such an interaction.

18
that the interaction with particular intermediates may be more specific than on metals. A non-metallic electrode can have surface sites of very different character, especially if the

surface consists of different elements which can form specific chemical bonds. Structural

defects, special crystal faces or surface places may form favourable reaction sites. No generalizing prediction can at present be made for the usefulness of non-metallic electrodes

besides that the individual properties of the constituents, the variety of lattice structures and of the electronic properties of such materials present an enormous reservoir of

different possibilities. This seems to be the principal challenge in using non-metallic electrodes for catalysis. However, one should not forget that the necessary high conductivity

and strong coupling between electronic bulk and surface states implies serious limitations

to the use of such materials in electrochemistry. To some extent the conductivity problem

can perhaps be overcome by using very thin layers on highly conductive substrates. However,

the stability of such layers of different chemical composition will always remain a critical

problem. The conclusion is that one should not expect catalytic miracles from non-metallic

electrodes, but in some cases and for some specific reactions, as with the Ruthenium Oxide

on Titanium substrates for the Chlorine and Oxygen electrolysis (23, 2U), one might be es- pecially lucky with such materials.

7. References

[1] Weiss, J. Trans. Faraday Soc. 31, 15^+7 (1935)-

[2] Gerischer, H. and Gerischer R., Z. physik. Chem. (Frankfurt ) 6, 178 (1956).

[3] Marcus, R.A. , Can. J. Chem. 31, 155 (1959); J- Chem. Phys . h3, 679 (1965).

[4] Marcus, R.A. , Ann. Rev. Phys. Chem. 155 (196U).

[5] Levich, V.G. and Dogonadze, R.R., Dokl. Akad. Nauk SSSR, ^2h_, 123 ( 1959).

[6] Levich, V.G. , Adv. Electrochem, Electrochem. Eng., h_, 2h9 (I966).

[7] Dogonadze, R.R. in "Reactions of Molec\iles at Electrodes", ed. by N.S. Hush, WileyIntersciences , New York-London-Sidney-Toronto 1971, p. 135-

[8] Gerischer, H. , Z. physik. Chem. (Fraiikf\irt ) , 26, 223, 325 (i960); 27, ^8 (1961).

[9] Gurney, R.W. , Proc . Roy. Soc. (London) A 13^ , 137 (1931).

[10] Gerischer, H. , Photochem. Photobiol. J_6, 2h3 (1972).

[11] Gerischer, H. in "Adv. Electrochemistry Electrochem. Eng.", 1, 139 (l96l).

[12] Lohmann, F. , Z. Naturforschung 22a, 8U3 (1967).

[13] Plieth, W.J. and Vetter, K.J., Z. physik. Chem. (Frankfurt) 6±, 282 (1968).

[14] Anderson, P.W. , Phys. Rev. J2U, hi (1961).

[15] Gadzuk, J.W. , Surface Sci. 6, 133 (1967); h3, (197^+).

[16] Luther, R., Z. physik. Chem. 3^, 1+88 (1900); 36, 385 (1901 ). [17] Parsons, R. , Trans. Faraday Soc. 5ii, 1053 (1958).

[18] Gerischer, H. , Z. physik. Chem. (Frankfurt) 8, 137 (1956); Bull. Chim. Soc. Beige

67, 506 (1958).

[19] Appleby, A.J. Catalysis Rev. k (2), 221 (1970), Marcel Dekker Inc., New York.

[20] Gerischer, H. , Physical Chem. Vol. IX A (Electrochemistry) ed. by Eyring, Henderson,

Jost, Academic Press 1970.

[21] Memming, R. and Mollers, F. , Ber, Bunsenges. physik. Chem. 76_, k69 , ^75 (1972).

[22] Gerischer, H. , Pettinger, B. and Luebke, M. , Proc. Symposium on Electrocatalysis

,

p. 162 (ed. by M. Breiter), The Electrochemical Society, Princeton 1972.

[23] Kuhn, A.T. and Mortimer C.J., J. Electrochem. Soc. J120, 231 (1973).

[24] Grady, W.O. , Iwakura, C. , Huang, J., and Yeager, E., Proc. Symposium on Electrocatalysis, p. 286 (ed. by M. Breiter), The Electrochem. Soc, Princeton 1972.

19

National Bureau of Standards Special Publication 455, Electrocatalysis on Non-Metallic Surfaces,

Proceedings of Workshop held at NBS, Gaithersbiirg, Md. , December 9-12, 1975. (Issued November 1976)

SURFACE STATES, CHEMISORPTION AND CATALYTIC PROCESSES

ON TRANSITION METAL OXIDES

T. Wolfram, F. J. Morin and R. Hurst

Department of Physics

University of Missouri-Columbia

Columbia, Missouri 65201

A discussion of the bulk and surface electronic states of transition

metal oxides such as the perovskites is presented. The factors which

determine the s\irface electronic structure are described and several

examples of surface states and surface energy bands are presented using

both cluster and energy band models.

The role of d-electron surface states in chemisorption and catalytic

processes on transition metal oxides is considered. Several examples

of surface reactions are given using a cluster model for the substrate.

These examples include the chemisorption of OH ions, the dissociative

adsorption of H^ and the hydrogenation of ethylene.

A detailed description of a Green's function theory for surface

reactions is given in which both the reacting molecule and the solid

are described by local basis states. The concept of the spectral

weight function (local density of states) is used to describe

delocalized bonding of molecules to a solid surface. As an example

of the application of the theory the formation of localized electronic

states associated with the adsorption of OH ions is considered.

Key words: Catalysis; chemisorption; cluster states; electronic

structure; energy bands; perovskites; surface states; transition metal oxides.

1. Introduction

VThen a crystalline solid is terminated by a surface new electronic states appear,

associated with the surface and derived from the bulk band structure. The surface states have wavefunctions which can hybridize with the orbitals of reacting molecules to produce

a surface complex. Chemisorption involves the formation of a stable surface complex while

catalysis involves the surface complex as an intermediate or transient state which

facilitates a molecular transformation.

The transition metals and transition metal compounds dominate the science and

technology of catalysis and a great deal of effort is currently directed towards understanding the mechanisms by which molecules interact with the surface of such solids.

Models for chemisorption and catalysis have tended to emphasize one or the other of

two extremes. One approach is developed around the concept of a surface molecule in which

the effects of electron delocalization in the catalyst due to energy band formation are IPermanent address: Science Center, Rockwell International, Thousand Oaks, California 91360

21
regarded as secondary or negligible. The solid is approximated by a small cluster of atoms and the interaction of molecules with the solid is expressed in terms of molecular

bonding concepts. Such models are of considerable value in isolating and identifying

possible mechanisms of surface reactions. They cannot be employed to calculate heats

of adsorption or activation energy barriers because solid state effects such as band

formation, screening and polarization and delocalization of surface bonds significantly

alter the energetics of surface reactions.

The second type of model, based on energy band concepts, emphasizes the delocalized

character of the electrons of the catalyst. The interaction of reactant molecules with

the solid is expressed in terms of matrix elements between molecular orbitals and all of

the energy band states of the solid. In principle, such models are capable of describing

the formation of localized chemical bonds. However, until recently approximations were frequently employed which minimized or negated the role of surface states and which tended

to obscure local bonding mechanisms. This type of model has applicability to some metals

and has provided insight into mechanisms of screening by mobile carriers and delocalization

of chemical bonds.

The transition metal compounds present a particularly interesting and difficult class

of materials for theoretical models. Their surface chemical properties can not be

described successfully by either a completely localized or a completely delocalized

electron model. The challenge here is clearly the development of models which treat both

local and solid state properties on an equal footing.

In this paper we describe some initial efforts to integrate local and band concepts

into a model for understanding chemisorption and catalysis on the transition metal oxides.

We have selected certain materials as model compounds upon which to base a theoretical

program for the study of the mechanisms of chemisorption and the mechanisms by which the

interaction between two molecules is catalyzed by a solid surface.

Surface reactions involved in chemisorption and catalysis are so complex that no single

model or theoretical approach can be expected to correlate the wide diversity of real

situations. Even in the case of a simple, specific surface reaction understanding of the

actual mechanisms of chemisorption or catalysis is minimal. Nevertheless there are a number of factors which are known to be of primary Importance in surface reactions. In this paper we discuss some of the factors believed to be fundamental in considerations

of surface chemistry on transition metal oxides. These include:

1. a high density of d-electron and p-electron surface states located in the forbidden

energy gap.

2. the formation of localized electronic states involving electrons shared between

surface states and molecular orbitals.

3. the effects of surface perturbations such as changes in the electrostatic

potentials and coordinatively unsaturated surface ions.

4. the effects of energy band formation on the character of chemisorption bonds.

We begin in section 2 with a discussion of the bulk electronic states of the transition

metal oxides. For concreteness we present examples of the electronic properties of the

simple cubic perovskites. In this discussion we present both the localized cluster model

and the energy band model and also show relation between the two models. A discussion

of the various types of surface states is given and examples of surface energy bands on SrTiOg are presented. The importance of the electrostatic Madelung potentials in

controlling the energies of surface states is emphasized. Detailed calculations of surface

states derived from the cluster model and from energy band theory are presented. Finally,

some experimental results relating to surface states on TiO^ and Ti^Og are briefly described.

In section 3 qualitative examples of chemisorption and catalytic processes on transition metal oxides are presented. The localized cluster model is employed to show

how the electronic levels of surface ions can hybridize with the molecular orbitals of

reacting molecules to form stable and unstable surface complexes. As examples we consider

22
the chemisorption of OH ions, the dissociative adsorption of and the hydrogenation of

ethylene

.

In section M- we discuss an energy band model for chemisorption. The theory employs a Green's function formalism based on localized basis states for both the solid and the

reacting molecule. The concept of the spectral weight function (local density of states)

and its relation to chemisorption is described. A simplified model of OH adsorption is

described and the relation of the various chemisorption states to those of the cluster

model are discussed.

2. The Transition Metal Oxides

2.1. Recent Applications

The transition metal oxides have been of interest for many years because of their

catalytic and electrocatalytic properties . An abbreviated list of the oxides belonging to several structural families is presented in Table I. Some of these oxides have received

particular attention recently including the perovskites and their various mixed crystals,

and TiO^-

Table 1. Some examples of transition metal oxides.

ROCKSALT STRUCTURE

Antiferromagnetic Mott insulators

MnO (d^)

FeOx (d5)

CoO (d^)

NiO (d8)

Metals

TiO (d2)

VO (d3)

PEROVSKITE AND RELATED STRUCTURES

Wide band gap insulators

SrTiOg (dO)

BaTiOa (dO)

KTaOg (dO)

WO3 (dO)

PbZrOg (dO)

Magnetic semiconductors

LaCo03 (d^)

LaMn03 (d**)

RUTILE STRUCTURE

Wide band gap insulators

Ti02 (dO)

CORUNDUM STRUCTURE

Antif erromagnetic Mott insulators

a-Cr203 (d3)

a-Fe203 (d^)

Metals

Re03 (dl)

LaTiOs (d^)

KM0O3 (d^)

CaMo03 (d2)

LaNiOg (d^)

Metals

VO2 (dl) (S.C. Z metal)

Cr02 (d^) ferromagnetic

Metals

Ti203 (S.C. t metal)

V2O3 (d2) (S.C. Z metal)

Pedersen and Libby [1] suggested that the rare-earth cobalt oxides might provide

inexpensive catalysts for the oxidation of CO and the reduction of the oxides of nitrogen. Voorhoeve et al. [2,3] reported on the use of mixed crystals such as (La Pb )MnO and

^^l-x^^x^""°3 "^^"^^ "^^^ ^^i^ perovskite structure and Libby and Aegerter C4]^found that reduced SrTiO and BaTiO are both active catalysts for isomerization, hydrogenation,

hydrogenolysis synthesis and cracking.

2— ^ Figures in brackets indicate the literature references at the end of this paper.

23
In the area of electrocatalysis Tseung and Bevan [5] investigated various compositions

of (La Sr, )CoO^ as candidate oxygen electrodes. They reported that La„ r-Sr„ ^CoO^

performs as a nearly ideal reversible electrode in a 45% KoH aqueous electrolyte at 25 and

80°C.

The possibility of utilizing solar radiation for the production of hydrogen by

electrochemical photodecomposition of water has stimulated a number of studies on transition

metal oxide electrodes. Fujishima and Hondo [6] demonstrated the feasibility of such

photodecomposition using a TiO^ anode. More recently Mavroides et al. [7] reported

studies of the photoelectrolysis of water using TIO^ and SrTiO. anodes. Quantum efficiencies for band gap radiation (~3eV) were found to be near 100%. The search for a material whose band-gap is matched to the peak in the solar radiation distribution (~2.4eV)

is currently focused on the transition metal oxides because many of these materials are stable as anodes in an electrochemical environment.

2.2. Bulk Electronic Properties

The oxides include a wide range of different structures and electronic properties.

They include insulators, semiconductors, metals and superconductors. Examples of ferromagnetics, f errimagnets , antiferromagnets and ferroelectrics can be found. Some oxides

possess delocalized d-bands while others have localized d-electrons and transitions from

localized to delocalized behavior may occur and semiconductor to metal transitions are common. For the purposes of our theoretical investigation of d-band surface states and

catalysis, an ideal structural form is the perovskite. This is because the simple

structure makes it theoretically tractable and because essentially all of the electronic

properties found among the transition metal oxides may be realized in the perovskites

by metal ion substitution.

A large class of ionic oxides having the formula unit ABO^ crystallize in the

perovskite structure (the structural details of which are illustrated in fig. 1.).

as

(b) (7)

(c)

Fig, 1. (a) Cubic perovskite structure containing octahedral clusters; (b) BOg octahedral

cluster, B = transition metal ion, 0 = oxygen ion; (c) unit cell for the perovskite

structure; (d) symmetry points in the Brillouin zone for the cubic perovskite structure,

structure

.

24
(A= larger cation, B=sinall cation.) The stability of the structure derives primarily from

the Madelung energy achieved when cations occupy oxygen octahedra which share corners

.

This requires that the B cations have a strong preference for octahedral coordination. In the discussion to follow, we are interested in the case where B=transition metal ion. Also, we shall consider only the (001) surface which contains the B cations. Presumably this is a catalytically active surface and, due to its stability, the surface likely to be exposed

to a chemical reaction rather than the (001) surface which contains the A cation. Furthermore, for our purposes, the ABO^ oxides of interest are not only those having the cubic

structure, of which SrTiO^ is typical, but also those whose structure may be considered to

be distorted cubic such as LiNbO^ (rhombohedral) , NaTaO^ (orthorhombic )

, PbTiO

(tetragonal). Cation formal charge varies in these structures over the range A (3+ to vacant site), B (3+ to 5+) so that, when mixed crystals are included, a very large number

of oxide perovskites are possible. Their electrical properties include d-band insulators,

metals, Mott insulators and superconductors. A general survey of these materials has been

compiled by Goodenough and Longo [8].

The bulk energy bands of perovskites are reviewed in this section, using SrTiO^ as an example, in order to provide a basis for the description of the surface states to follow.

The properties^of the perovskites are dominated by the properties of the electronic states derived from 0 and the B ion. The A ion contributes strongly to the Madelung potentials

but has only a minor influence on the low-energy electronic and optical properties because

its energy levels lie much higher than those of the B ion. Therefore, for the purpose of

determining the structure of the d-bands , we ignore the A ion so that effectively we regard the perovskites as transition metal trioxides BO^. From this point of view ReO^ is a typical member of the d-band perovskite family. Figure 2 summarizes the major energy

parameters which influence the electronic structure. The energy levels of the free ions

relative to vacuum are shown on the left of figure 2a before the ions are combined to form

the crystal. These levels include the ionization_energy f^r the reaction Ti^t Ti^+ + e(-i4-3.2eV) and the electron affinity of oxygen, 0 + e -> 0 (+9.2 eV). When these ions

are brought together to form the perovskite structure large electric potentials arise which

shift the energy levels.

MADELUNG COVALENT CRYSTAL BAND DENSITY POTENTIALS MIXING FIELD FORMATION OF STATES

(a) (b) (c) (d) (e) Fig. 2. Schematic of the major energies involved in the electronic structure

of a perovskite.

25
The electrostatic field acting on any site in the lattice can be expressed in terms

of various electric multipoles. The monopole contribution produces large Madelung

potentials while the higher-order poles produce a crystal field effect. The effect of the

Madelung potential on the energy levels of the free ions is illustrated in figure 2a. The

negatively charged oxygen ions produce a repulsive Madelung potential at the B-ion sites.

These repulsive potentials decrease the effective ionization energy (raise the energy

levels) for electrons on the B ions. For example, in SrTiOs , the full Madelung potential

at the Ti^+ site is 45 eV. Similarly, the positively charged B ions provide an attractive

Madelung potential of -20 eV at the oxygen site which lowers the 02- level below the B-ion

levels. In the d-band insulating perovskites such as SrTi03 and LiNb03 the levels will

be filled and the B-ion levels will be unoccupied.

The perovskites are not completely ionic and there is some covalent sharing of

charge. This has the effect of altering the formal charge on the ions. This effect in

SrTi03, for example, reduces the formal charge on the oxygen ion from -2 to about -1.8 and

a corresponding reduction in the cation formal charges. Reduction of the formal charge

reduces the Madelung potentials, and changes the energy levels as illustrated in figure 2b. This in turn reduces the energy gap from 17 eV to 3.2 eV

.

Next we consider the qualitative effects of the crystal field. The B ion is surrounded

by an octahedron of oxygen atoms which produce a cubic crystal field. The d- levels are split by this field into the familiar t2g and Cg levels as illustrated in figure 2c. The

d-orbitals transforming according to the t2g representation are dj,y, dj^^s ^^d dy^,- The

doubly degenerate eg level includes the dj^2_y2 and d^2 states. The t^ states nave their

wavefunction lobes directed in between the repulsive oxygen-ions while^the lobes of the

Cg states point directly into the repulsive oxygen-ion sites. As a consequence, the states are lowered in energy while the e^ states are raised in energy. The "center of ^ gravity" of the d-level of course in unaffected by the higher order multipoles producing

the crystal fields. The oxygen-ions experience an axial crystal field due to the cations

and the p-orbitals may be classified as P and P according to the orientation of the

p-orbital lobes with respect to the cation-anion axis.

Now we proceed to a discussion of the energy band effects, as illustrated in figure 2d

for three points in the Brillouin zone. A number of studies, using a variety of techniques,

have been reported on the energy band structure. In terms of the simple LCAO model, the

largest energies which contribute to the formation of energy bands are:

1) The (pdo) transfer integral which leads to mixing of the e„ states with p-orbitals

(2-1+ eV).

2) The (pdir) transfer integral which leads to mixing of the states with the pi-orbitals

(1-2 eV). -L 3) The (ppo) and (ppir) transfer integrals which couple p-orbitals on different oxygen sites

(few tenths of an eV).

Spin orbit coupling, polaron effects, anion distortion effects, and more distant

neighbor interactions are not important in determining the qualitative features and will be

ignored in subsequent discussions. The (pda) interaction broadens the eg levels into a conduction band with width of the order of 3-4 eV. The (pdir) interaction produces a t2g

conduction band with width of the order of 2-3 eV. Similar band widths are produced in the

valence bands due to both (pdo) and (pdTi). Additional broadening of the valence band is also produced by (ppTr) and (ppo). At the point F, the t2g and e_ band wavefunctions have

pure d character, the mixing with the p-orbitals occurring away from F. Note that the

t2g-eg band edge separation at F is not equal to the crystal field splitting.

We shall be mainly interested in the lower or t2g conduction band. The surface states which derive from this band and exist in the energy gap between the valence and conduction

band have been shown to possess the t2g symmetry important in catalytic activity.

A feature to note in connection with the t2g-bands is the "flat-band" region between

F and X. This characteristic can be traced directly to the two-dimensional nature of the

(pd ) transfer integral. For example, in the two center approximation d^y interacts with

Pj^ and Py but not with p^; dj^^ interacts with p^^ and p^ but not with p and finally d interacts with p and p^ but not with p . As shown quantitatively in Xur previous ^

26
papers [9,10,11], this effect leads to three equivalent energy bands, E , (a6=sy,xz, and

yz) each of which depends upon two components of the wavevector (k^, ana k^) but not on the third. The flat band region results from the lack of dependence on the third component.

A more exacting band-theory calculation [12], of course, does not predict an absolutely flat

band but does predict very small dispersion and therefore effectively confirms features

exhibited here.

There are a great many perovskites, LaMn03 is an example, which are Mott insulators

because exchange and correlation energies are larger than the band width and thus delectron localization rather than energy band formation results. For these materials the

d-electronic structure is nearer to that illustrated by figure 2c than to the band

representation of figure 2d.

Figure 2e illustrates the density of electronic states expected for d-band perovskites.

The shaded portion is the density of states associated with the three t2 bands. The

abrupt jump in the density at the band edges and the approximate logarithmic singularity

in the center of the band are a result of the two-dimensional characters of the t2„

conduction bands. These properties are easily picked out in density of states histograms,

published by Mattheiss [12] based on APW calculations for ReO^ , SrTiO^, BaTiO^ , and KTaO^-

These features also produce characteristic structure in the optical properties of these

perovskites. In a previous paper [11] we developed approximate analytical expressions,

from a simple LCAO model, for the energy bands, wavefunctions , density of states » frequencydependent dielectric function and reflectivity. The structure in the energy dependence

of the dielectric function and reflectivity was shown to be in excellent agreement with the

experimental data of Cardona [13] on SrTiO . These analytic expressions provide the basis

for a description of the surface states [18].

Density of states diagrams for perovskite materials with localized d-electrons are not shown in figure 2e, but clearly they consist of delta function peaks at the t^ and e energies. ^ ^

We have emphasized the two-dimensional character of the t bands because this

characteristic is important in relating the d-electron states t§ the catalytic activity on perovskites. For example, the abrupt jump in the t -band density of states in the band

edge means that a high density of nearly pure d-ele8?ron states exist in a narrow range of

energies. This argument, for the perovskites having localized d-electrons, is even

stronger.

Where d-electron localization occurs, as in LaCO^ , conduction occurs by a "hopping"

process. In LaCoO^ the B ion configuration at low temperatures is thought to be

*2g ®g' small energy separation between the filled t2g and empty eg levels results

in a room temperature mixture of "low-spin" (t^g e^) and "hign-spin" (t2g ^g^ cobalt ions

(designated CoHI and Co3+ in the literature). Electron transport involves hopping of

electrons between B ions creating Co+2 ^nd Co"*"^ electron-hole pairs, p or n-type

conduction occurs in doped material. For example, when Sr+2 is substituted for La+3 in iha^ Sr"'"^)(Cot^ Co"*" )0. transport involves hopping of Co"'"'^ holes in the Co"*"-^ lattice.

1—X X 1—X X 3 + +4- +3 4-0 o Similarly in the doped (La Th )(Co Co )0 transport involves hopping of Co+2 elec-

_ r"^ . J_XX XXX O trons in the Co lattice.

In the case of perovskites with localized d-electrons the electronic structure is closer to that described by the crystal field picture, figure 2c . The "cluster" approach

is more appropriate than the energy band method. In the cluster model the electronic

properties of the perovskite are approximated by those of a BOg octahedral complex which

is shown in figure 1. The electronic states of similar clusters have been investigated

recently by Tossel, Vaughan and Johnson [14] using a self-consistent Xa approximation and

utilized to interpret the properties of rutile, wustite and hematite. In the application

of the cluster method to the perovskites, the effects of the Madelung potentials which are present in the bulk material must be included.

We have studied the character of the cluster states of the BO^ cluster and their

relation to band states in d-band perovskites. A schematic of some of the levels of a typical BO cluster spectrum are shown in figure 3a. The various states are labeled

according to octahedral group theoretical notation.

27
Fig. 3. Cluster and energy band states; (a) states of a BOg octahedral

cluster; (b) energy bands. The band states with equivalent

cluster symmetry are indicated.

The crystal field 3e and 2t^ states correspond to d-orbitals with a 20 to 30%

covalent admixture of oxyfen 2 p-o?bitals. The corresponding valence states, 2e and It^ are mostly oxygen 2 p-orbitals vjith a 20 to 30% admixture of d-orbitals. The ^2^'

^ and t^ states are essentially non-bonding states since they involve no hybridization^with

the ca¥ion orbitals. Additional cluster states arise when the B-cation Us and Up and

oxygen 2s states are included.

In addition to the states discussed previously new states appear because of the added

Up and 2s basis states. The group, 5a^ , le and '^'^-.^ ^re essentially oxygen 2s states and the 7t level is derived from the §p ca?lon states. The M-p states also enter the

lu

5t, levels so that these levels have some bonding character,

lu

There is a strong correlation between the cluster states and the energy band states

of the corresponding solid at points of high symmetry in the Brillouin zone [15]. The

correspondence for the perovskites is particularly simple. The group of the k-vector

possesses full cubic symmetry (Oj^) at the points T and R in the Brillouin zone. The

cluster wavefunctions and energies can be approximately related to the energy band wave- functions and energies at either T or R in the Brillouin zone. These relationships are indicated in figure 3b on the energy band diagram. For example, the upper valence band

state at T possesses the same symmetry as the 6t^ cluster state. One may imagine

constructing the band state wavefunction by repea'^ing a cluster wavefunction of Bt-,^ symmetry in every unit cell throughout the entire crystal. Similarly the pdir" band state

at R may be constructed by repeating a cluster wavefunction of 2t2g symmetry in each unit

cell with the phase reversed in adjacent cells. This process is illustrated in figure 4 for the band state F^^ in terms of a cluster state of Itj^ symmetry.

The cluster states are mixtures of particular symmetry coordinates which transform

according to the irreducible representations of the cubic group ^0 Specific relationships among the orbitals which make up ligand symmetry coordinates are fixed by symmetry.

28
2u CLUSTER STATE

Fig. 4. Illustration of symmetry relationship

between cluster and band wavefunctions

.

BAND STATE

In the band model the point group symmetry depends upon the wavevector k. For an arbitrary k-vector arbitrary phase relationships exist among the orbitals making up the energy

band wavefunction. As these phases vary with k so do the energies. The dependence of the

energy on k determines the energy bands of the solid.

In the insulating oxides such as SrTiO^, BaTiO^ , KTaO^ or TiO at the point T in the

Brillouin zone the It^ is the highest filled state and the 2t^ is the lowest unfilled

state. The energy sepiration between these levels approximates^the energy gap which tends

to be on the order of 3 eV. The energy difference between the 3e and the 2t- states is the crystal field splitting 10 Dq. ^ ^

The insulating materials can be made semiconducting or metallic by introducing oxygen

vacancies or by s\ibstitution doping. This results in partial occupancy of the 2t„ states.

In materials such as ReO^' ^'^2 °^ '^^2*^3 cationic configuration is d"*". This aa§itional

electron resides in the 2t„ state (or pdTT"-band) and these materials are metallic.

2g For magnetic materials the exchange and correlation interactions lead to a situation

in which the cluster states have different energies for different spin orientations as has been illustrated in recent calculations of Johnson and Messmer [16] for antiferromagnetic NiO. In this case the filling of the spin dependent levels with electrons would not

be equal and a net spin results on each cation.

The cluster model serves as a preliminary model for examining the local electronic

structure of the oxides and is reasonably appropriate for oxides in which the d- electrons

are localized as would be expected in, for example, LaCoO^. On the other hand, there are many oxides for which the d-electrons are delocalized, TiO^, Ti^O^, SrTiO^ , KTaO^ , and

KMoOg are examples, and the effects of energy band formation must be considered.

2.3. Surface States

Whenever a solid is terminated by a surface or interface, special electronic states occur whose wavefunctions are legalized at and near the surface. If a sufficiently large

portion of the surface (eg. 100 A x 100 8) is regular then the surface states will be

characterized by a two-dimensional (real) wavevector K and two-dimensional surface energy

bands E(K) will result. The wavefunctions of the surface state will be characterized by

K and by a complex constant y(K) which determines the decrease of the amplitude of the

wavefunction iiiK) normal to the surface. In most cases the surface state charge density

e|i|;l^ will decrease exponentially as e 2y(K)z ^ith increasing distance z into the solid as illustrated in figure 5

.

29
CHEMISORPTION INDUCED

9

Fig. 5. Schematic of the various types

of surface states and approximate

densities of states.

5 X lO'^ cm"^

A second type of surface state shown in figure 5 is associated with atomic edges or steps, grain boundaries, dislocations or other "line" discontinuities. These states are characterized by a one-dimensional wavevector K^^ and complex decay constants which

determine the manner in which the wavefunctions decays perpendicular to the line.

Finally atomic—like surface states occur at the site of a vacancy, intersitial,

impurity or chemisorbed atom or molecule. The intersection of a dislocation with a surface or an atomic corner may also produce localized surface states.

The densities of these various possible surface states can be estimated from general

considerations. A typical oxide crystal will have 10 d-electron states in a unit cell of

M- X on a side thereby producing 1.5 x 10^3 d-electron states /cm^. Then there will be a

6 x IqI^ states/cm^ for two-dimensional surface states.

o Assuming one line structure (step or-j^gdge) in each square 100 A on a side leads to a density of line states of the order of 10 cm ^ . An impurity content of lO^^ cm ^ would

produce 5 x 10-'-^ surface impurities/cm . All of these densities are significant in an electrochemical reaction. For example, if 6.3 x lO-'-^ impurities centers per square

centimeter are sites for a reaction in which the reaction time is on the average one millisecond an electrochemical current of one milliampere per cm^ would result. Chemisorption—induced surface states are of importance by definition and their density will be

equal to the surface coverage, 6. Approximate locations of different types of surface

states are indicated in figure 5 on the energy band diagram for SrTiOg.

The degree of localization of a surface state on a solid having energy bands

increases with the energy separation between the surface state and the corresponding bulk

energy. Highly localized surface states occur only in energy regions forbidden to the

normal bulk states. Thus if highly localized surface states occur they will be found in energy gaps. In this respect non-metallic materials such as insulators and semiconductors

provide more opportunity for the occurrence of well defined, highly localized surface

states than do metals. Insulators and semiconductors have a global energy gap in which

surface states can occur. By contrast a transition metal does not possess a global energy

gap in the energy range near the d-bands . Instead surface states can only exist in small

local energy regions provided by the sp-d hybridization gaps.

30
2.4. Origion of Intrinsic Surface States

Surface states which occur on a surface in the absence of foreign atoms are called

intrinsic surface states. The major factors which produce intrinsic surface states and

which influence the energy of such states are

(a) Missing neighboring atoms

(b) Changes in the electrostatic potentials at a surface

(c) Changes in the covalent bonding interactions at the surface

(d) Relaxation of the surface layers and changes in the interlayer spacings

(e) Reconstruction of the surface structure

In ionic crystals such as the oxides the electrostatic potentials acting on the surface

ions are substantially altered due to missing neighbors. This effect can be calculated

exactly [9]. Variation in the Madelung potentials at Ti or oxygen surface site is

illustrated in figure 7 for a type I configuration of a (001) surface (see fig. 8) of

SrTiOg. The Madelung potential at the surface Ti site is reduced by about 2 eV while that

for a surface oxygen is essentially unchanged from its bulk value. On a type II surface

(obtained by removing the top layer from the type I surface, see fig. 8) the oxygen and

strontium ions are exposed while the titanium ion is buried in the next layer. For the

type II surface the oxygen Madelung potential is increased (towards vacuum) by 3 eV while the Ti potential is unchanged from its bulk value.

31
TYPE 1 SURFACE (a)

(>

(011)- lURFACE -6

o- -6-

Fig. 8. Illustration of the type I and type II (001) surfaces of a perovskite structure. 0 = A ions; • = B ions; O = oxygen.

TYPE II SURFACE (b)

(001) SURFACE

A second feature of the surface is that due to the missing neighbors the electrostatic

crystalline field is no longer cubic. An axial field appears which can cause further

splitting of the levels. This is illustrated in figure 9.

Also from figure 7 it is clear that a large electric field will be present at the

surface. An estimate places the field in the 10^ volts/cm range at the Ti surface site.

32
BULK TYPE I SURFACE

-i—

6dq

4dq

Fig. 9 Crystal field splitting of

d-levels including an axial

field at a surface ion.

A large electric field will also act on the surface oxygen ion. These fields tend to move

the ions out of their normal surface positions. This tendency is countered by elastic

restoring forces. However, because for small displacements the electrostatic energy is

linear in the displacement while the elastic energy is quadratic, the electric forces

prevail as in the Jahn-Teller effect. It can also be shown that small displacements of the

surface ions leads to increased electronic bonding [10]. One possible result of these

forces is to produce "puckering" of the surface ions or an undulating surface. The precise

configuration of the surface ions is not yet known, but recent electron spin resonance

measurements suggest that the axial field component produced at the surface is only a few

percent of the crystal field energy (10 Dq) indicating that the electric field has been

greatly reduced. We shall ignore the possible Stark splitting due to surface electric

fields in this discussion.

It is also clear from figure 7 that the electrostatic Madelung fields extend several

angstroms out beyond the surface. The potential above a Ti ion, for example, exhibits a

3 eV attractive minimum at about 2 X above the surface. This attractive minimum is

believed to be important in attracting surface reactants to cation sites at the initiation

of a surface reactions. We shall return to this possibility later.

2.5. Cluster and Energy Band Surface States

As an introduction to surface states which occur on the transition metal oxides we consider a cluster consisting of a transition metal cation surrounded by an octrahedron of

oxygen ions. The cluster analogs of the surface states are defined as follows. The type

I states correspond to a BO^ cluster with a reduction in the electrostatic Madelung

potential acting on the cation. Type 11 cluster surface model corresponds to a BOg

cluster with the Madelung potential of one of the oxygen ions increased.

The surface cluster levels are shown in figure 10b and compared with the unperturbed

cluster levels. The center panel is the unperturbed cluster states. Figure 10c shows

the Type I cluster surface states and figure 10a shows the Type II cluster surface states.

For the discussions here the important cluster surface states are those derived from

9 9 9 the 2t^ and 3e states labeled xy, xz, yz and x'^-y and z in figure 10c. We shall also

discuss^the valince states e and a.2 derived from the It level. In order to simplify

the diagram we show on figure 11 only these levels. The^wavefunctions for two of the

cluster surface states are shown. The xz, yz surface states derived from the 2t d-states have d-orbitals pertruding normal to the surface which can react with mol^ules

coming in contact with the surface. The (doubly degenerate) e states have p^, orbital

extending normal to the surface.

33
-10 =

Fig. 10. Cluster states: (a) surface cluster states

for a BOg cluster corresponding to a type

II surface; (b) BOg cluster states corresponding to bulk levels; (c) BO5 surface

cluster states corresponding to a type I surface

.

When the d-electron surface states are occupied with electrons they provide orbitals

of the proper symmetry for strong interaction with the antibonding states of molecular

species such as H , 0^, and CO and other diatomic molecules. Transfer of electrons from

surface states into empty molecular antibonding orbitals by means of hybridization is

often important in facilitating dissociative chemisorption . These states also provide ideal

orbitals for adsorption of OH iofts and for interaction with the it electrons of hydrocarbons. We have previously described a model which illustrates the catalytic role of

d-band surface states in facilitating the concerted reaction of two ethylene molecules to form cyclobutane [17]. , .

34
BOe BOg BO5

3eg

Fig. 11. Cluster states and wave functions

of the d- and p-orbital surface states.

In an energy band model surface bands are found ivhich are the analogs of the cluster

states just described. The oxides possess a high density of localized d-electron and

p-electron surface energy band states in the energy gap between the valence and conduction

bands. In previous work [10,17] we showed that for a (001) surface the density of surface

states possessed square-root singularities at the band edges. This behavior is characteristic of a 1-dimensional energy band and is a direct consequence of the two-dimensional

character of the lower conduction band. The energies and density of states are shown in

figure 12 for SrTiO^. The dashed curves indicate the bulk energy bands while the surface

states are shown as solid curves. The cross-hatched regions of the surface density of states indicates a small energy range which each contain 10% of the total density of

NUMBER OF ElfCTRON STATES/Rd - SPIN - CELL

5 10 15 20 25 30 35 10 15

-5.C

C -8.0

-9.0

-10.0

-11.0

-12.0

-13.0

/ /

/ /

CONDUCTION LANDS

- SURFACE BANDS

-N ^ —

S ^

*

VALENCE BANDS

"I 1\—' r 1 1 r \ (BULK) s

10 20 30 to 50 60 70 80 90 NUMBER OF ELECTRON STATES/Rd -SPIN -SURFACE- CELL

HA

"lA"

I I X H

12. Bulk and surface energy bands for

SrTi03. The density of states for

the bulk and surface band are shown on the right hand side of the figure.

The hatched areas each contain 10 % of the total surface state density

(1.2 X IqI'* states/cm^). The lower

right portion illustrates one of the

two equivalent d-orbital surface

states near the bottom of the surface

band.

35
states. The states corresponding to the lower cross-hatched region are nearly pure

d-electron states having the t^ symmetry. This region contains about 1.2 x lOl^ states per cm2. The surface state wavi function for states in this region is shown on the right

hand side to consist of an array of dy^ orbitals. Equivalent surface states composed of

^xz orbitals are degenerate with the dy2 states

.

Surface states are also split-off from the pdo" bands and from the oxygen valence

bands. One type of surface band derived from the pdw valence band having p^ symmetry is believed to play an essential role in one mechanism for the dissociation of H2 . This will

be discussed in the next section.

TYPE II SURFACE

(a) Sr TiOj

In previous work we reported a study of the dependence of surface energy bands on surface perturbations. Some of these results are shown in figure 13. The different

surface bands correspond to different surface perturbation parameters . The details are described in reference 10. The point to be noted here is that both d-band and p-band

surface states occur in the energy gap and that on a non-uniform surface a distribution of

such bands may be anticipated corresponding to various crystallographic facets and surface

irregularities

.

A comparison of the d-electron surface bands for an ideal (001) type I surface of

SrTiOg with the corresponding cluster surface states is shown in figure 14. The cluster

energies are the same as those shown in figure 10. The unperturbed cluster levels 3eg and

2t^ are shown at R in the Brillouin zone and the cluster surface states are shown at

M i§ the zone for reasons described earlier. The solid curves are the bulk energy bands

and the dashed curves are the surface energy bands

.

36
Fig. 14. Comparison of cluster states with the

energy band states of SrTiOs. The BOg

cluster states 3eg and 2t2g are ' com- pared to band states at R in the

Brillouin zone. The cluster surface

states x^-y^, z^, xy, xz and yz are compared with band states at M.

The electronic structure of TiO^ and Ti 0^ surfaces has been studied recently by

Henrich, Dresselhaus and Zeiger [19] using electron energy-loss spectroscopy. They

observe a large peak at 1.5 to 2 eV in the electron energy-loss spectrum of freshly

broken samples of Ti^O^ but not for TiO^. After sputter-etching with Ar ions the TiO

samples also exhibited the 2 eV peak. Subsequent exposure to oxygen removed the peak in

TiO^ and diminished the intensity in Ti^O^. The sputter-etching by Ar ions was believed

on the basis of Auger studies to reduce the surface.

Henrich, Dresselhaus and Zeiger suggest that the 2 eV peak may be associated with

Ti (Sd-*-) surface ions on Ti^O^ or created on TiO^ by the sputter-etching process. The

2 eV energy is then attributed to transitions between the crystal-field-split levels of

the 3dl states of Ti ions

.

2 . . Transitions between crystal field surface states, XZ-»-Z may be involved. According to cluster calculations lODq is about 3 to 3.5 eV and therfore the transition for an octa- hedrally coordinated Ti ion would be too large to explain the 2 eV peak. However, as illustrated in figure 10, the splitting of the d-electron states of a TiO cluster

(E -E 2) is reduced to about 2.0 eV. The band model gives the same result.

XZ zFrank et. al [20] have obtained electrochemical data for the standard potential of a number of redox couples in nonaqueous acetonitrile from which they were able to deduce the

existence of surface states on n-type TiO and other semiconducting electrodes. Their

results indicate a density of surface state at energies about 1 eV below the conduction

band. This is in rough agreement with our theoretical results which show the XZ, YZ surface states lie about 1.5 eV below the 2t level (see figure 10).

3. Chemisorption and Catalytic Processes

In this section we discuss in a qualitative fashion some mechanisms for chemisorption

and catalysis which can be described utilizing the cluster and energy band surface states

from the previous section.

3.1. A Mechanism for Adsorption of OH Ions

It is well known that oxides have an affinity for adsorption of OH ions. The

surface coverage can be substantial, particularly when the surface is in aqueous solution.

37
It is obvious that a surface 0 anion can form an OH ion by accepting a proton from

a hydronium ion (H^O )^in solution. This process is energetically favorable since the

positively charged H^O ion will be attracted by the Madelung potential to a position

directly above the oxygen ion (see figure 8) and a-bonding to a p oxygen orbital can

occur, Electrostatic energy is also gained when a surface ion is oxidized to a OH

ion.

The process by which an OH ion becomes bound to a surface cation is less obvious.

One possible mechanism is illustrated in figure 15 which shows the molecular orbitals of

the OH ion together with cluster surface orbitals which bond strongly. For the

discussion here we treat the non-bonding OH states as simply the oxygen 2p and 2p

orbitals and the bonding state is an oxygen 2p orbital a-bonded to a hydrogen Is s¥ate. The OH ion is assumed to be attracted to the surface by the attractive Madelung potential

2^ directly above the cation site (see fig. 7). For discussion of this example we use the

TiO^ cluster results. The OH energy levels are determined approximately by several

estimates. The ionization energy and dissociation energies of OH are 2.1 and 5 eV respectively [21,22]. The separation between the non-bonding and bonding states is approximated by one half of the dissociation energy; 2.5 eV. The ionization energy

refers to the liberation of an electron from a non-bonding state. Thus we locate p and

p at -2.1 -2.9 = -5 eV below the vacuum level where the -2.9 eV term is the attractive

Madelung potential above a Ti surface ion. The bonding state then is located at -7.5 eV below the vacuum level.

In solution solvation of the OH ion and the Helmholtz layer will alter these gas phase energies in a manner which is not known precisely. It would be expected that the

energies will be shifted towards the vacuum level. However, the energies of the surface

states will also be shifted towards vacuum for a typical semiconducting electrode because

of band bending at the surface. In the absence of more information we assume these

shifts are comparable and therfore omit them in the energy level diagram. This

approximation should not invalidate the mechanisms of adsorption to be discussed, however,

the numerical results should not be taken too seriously.

The hybridized energy levels are determined by using the values of the two-center

LCAO transfer integrals determined by Mattheiss [12] or Wolfram, Hurst and Morin [15]

which both give (pdo)~2 eV , (pdTr)~l eV for SrTiO^ or TiO^.

From figure 15 it is clear that levels (a) and (b) represent the antibonding

states of an OH surface complex. Except for the proton the wavefunctions are very

similar to the bulk 3 e and 2t2 BO cluster states. Occupation of these antibonding

states by electrons wil? reduce fhe OH binding energy and therefore it follows that the

maximum bond will occur on an insulatin crystal such as SrTiO^ or TiO^

of OH will occur when the xz, yz and z'^ surface states are occupied. In LaCoO„ and its

Minimum bonding

3

-2

>D _4

>-o

cc^ -6

LU -8

yz,xz

a

b

c

d \n cr

OH

t

« % ^ CO Y\ OO oo

8 (a) (b) (c) S (d)

Fig. 15. Energy levels and wavefunctions for the

hybridization of 0H~ non-bonding states with d-electron surface states.

38
mixed crystals the bulk 3e states are partially occupied and on the basis of figure 15

a smaller binding energy w§uld be expected as compared with TiO^, for example.

Alternately the antibonding levels can be^occupied by application of a cathodic bias

which may also tend to move the OH towards OH . Thus with cathodic bias OH may be

discharged leaving an exposed cation. This cation may then serve as a catalytic site for

the associative desorption of H^ by combining hydrogen atoms attached to adjacent surface

oxygen anions. The inverse process, the dissociative adsorption of H^ , will be discussed

subsequently.

Finally we note that for OH attached to a surface cation it may be energetically

favorable for the proton to migrate to an adjacent surface anion and then to an interior

anion.

3.2. A Mechanism for Dissociative Adsorption of H^

We consider H^ in the gas phase interacting with a n-type reduced or doped oxide such

as SrTiOg or Ti02. The filled bonding states, oCH^), and the empty antibonding states,

oCH^)", of molecular hydrogen are located at about -15 and -10 eV respectively. The Fermilevel of the substrate is assumed to lie within the lower conduction band at about -5.5 eV

(xz state, figure 10) so that in the cluster picture the xz surface state is occupied

while the state is empty. When the H molecule approaches the surface as illustrated

in figure 16, hybridization of the xz and aiH^)'" and hybridization of the z^ and a(U^)

orbitals occurs. Because of the large energy separation between the and aiH^) states

the latter process is unimportant in this example. The hybridization of xz and aCH^)"

leads to electron transfer from the substrate to the molecule. The electrons residing in

the hybrid state have several effects. First, they constitute the formation of Ti-H bonds.

Second they weaken the H-H bond and third, they cause the H-H internuclear bond length to increase. Finally, they cause the oCH^)'* state to rise in energy (towards the affinity

level of H^ which is only -0.7 eV) due to the Coulomb repulsion among the electrons. As the state rises in energy the interaction between aiH^)* and xz is enhanced because these

levels move closer together in energy and because the molecule is being pulled in towards

the surface. Thus a positive feed back situation exists in which the transfer of charge

invites the transfer of even more charge. The process is terminated by two factors. The

first factor is that eventually the effective energy of oCH^)" would rise above the Fermilevel of the substrate due to Coulomb repulsion. Second, as charge transfer occurs the

H-H bond length increases eventually moving the hydrogen atoms away from the Ti ion as the molecule proceeds along the reaction coordinate. The maximum bonding to the Ti ion

39
Bond order calculation of the dissociative

chemisorption of H2 on SrTi03 showing (a)

the bond order and (b) the total bond

energies as a function of the reaction

coordinate

.

occurs when the H-Ti bond length is about 1.2 X as shown in figure 17 and the H-H bond

length has increased from 0.75 % to about 1.5 %. This means that the molecule is

effectively dissociated.

The formation of OH ions now proceeds by interaction of the individual atomic

hydrogen atoms with the e oxygen surface states (see figure 10) which have p orbitals

projecting normal to the surface. As can be seen from figure 17 (b) the process is

energetically favorable. As the hybridization with the e oxygen states proceeds the Ti-H

bonds are weakened. Conceptually the weakening of the Ti-H bonds means that electrons

are flowing out of the a(U^)" state and back into the xz surface state. This lowers the

coulomb repulsion energy and drops the hydrogen levels into registry with the e oxygen

states. Again a positive feedback situation exists so that as charge leaves the extended

hydrogen molecule the individual hydrogens become strongly bound to the surface oxygen ions

.

The protons may now migrate over the surface and into the crystal so that the same site

is again available for catalysis.

The results displayed in figures 15 and 17 were obtained by use of a bond-order model

similar to that described by Weinberg et al. [23] but modified to include the presence of a Fermi level.

According to figure 17 the process has a 1.1 eV activation energy in the gas phase.

This barrier may be greatly reduced for reactions in solution where the coulomb repulsion

energy is reduced by solvation effects and by mobile electrolyte ions in the Helmholtz

layer.

The inverse of the reaction described here would -lead to hydrogen evolution at an oxide cathode

.

Several other mechanisms for the dissociative adsorption of H are also conceivable.

For example, Burwell has suggested a "heterolytic" process in which one hydrogen atom

adsorbs to an oxygen surface ion while the other adsorbs on the adjacent surface cation.

The relative importance of various mechanisms is not well established.

3.3. Role of d-Electron Surface States in Concerted Reactions

We present here a schematic discussion of the concerted reaction of molecular hydrogen

with ethylene to form ethane. The purpose here is to illustrate qualitatively how the

40
hybridization of reactants with surface orbitals can lower an electronic activation energy

barrier.

Hydrogenation of an alkene such as ethylene is an exothermic reaction but it does not proceed in the gas phase without the aid of a catalyst such as a transition metal or transition metal oxide. a-Cr202 (chromia), for example, will catalyze the hydrogenation of

ethylene [24].

There are several different mechanisms that have been proposed for the hydrogenation

process. Most of these assume the dissociative adsorption of H on the catalyst as an initial step. Here we describe how the concerted reaction might be catalyzed by the

surface states of a transition metal oxide. The relative importance of this mode of

hydrogenation is not known and therefore this example should be regarded as illustrative.

Consider first the reaction of with C^H^^ in the gas phase without a catalyst. The

hydrogen and ethylene molecules are assumed to approach each other along the reaction

coordinate shown in figure 18 . On the right of the figure we show the approximate

locations of the bonding, a(e), and antibonding, a(e)", states of the C-H bonds which are formed in the reaction. On the left of the figure we show the approximate locations of the

relevant bonding and antibonding states of the non-interacting reactants. These energies

are based on experimental and theoretical results in the literature [25-27].

The reaction coordinate is the C-H distance of the incipient bonds and the reaction

proceeds from left to right on the diagram. The diagram is schematic but we have attempted

to indicate typical energies.

Fig. 18. Schematic of the correlation diagram for the

concerted reaction of H2 with ethylene to produce ethane. Arrows indicate electron flow.

The dotted curve illustrates an activation

energy of about 50 Kcal/mol per electron.

As the distance between the reactants decreases (left to right) various hybrid

orbital states are formed representing the intermediate states of a H . ...C H complex.

The strongly interacting combinations are: aiH^) with tt to form the bonding combination

designated as a(.H^)+-n and the antibonding combination ir-aCH^); a(H ) with ir'- to form the

bonding combination a(H )+7r" and the antibonding combination -rr-'-oCH ). Bonding here means

those combinations which concentrate charge in the region in which the new C-H bonds are

formed. Antibonding combinations deplete charge from the C-H regions. From such considerations it is evident how to connect the various branches to the product states a(e) and

0""(e). A schematic of the energies of different branches along the reaction coordinate is

given in figure 19

.

41
Fig. 19. Correlation diagram showing the d-electron surface

state level before hybridization. Strongly bonding

combinations of molecular and surface state orbitals are indicated at the top of the figure.

In the non- interacting condition the oCH^) and tt levels are each occupied by a pair of

electrons. Both of these pairs must be delivered to the a(e) levels of ethane for a ground state reaction to occur. The problem is that the occupied tt-oCH^) branch does not correlate with a(e) but rather with a'''(e). At cross-over between the tt-oCH^) and a"(H )+Tr" it could be assumed that the electron pair switches branches. Woodward and Hoffman [28]

have argued that such a transfer is symmetry forbidden because gross charge redistribution

would have to occur. Thus electron flow is constrained to the TT-a(H2) branch as indicated

by the dashed arrow. It can be argued that at some point of close approach this symmetry

selection rule must break down due to the imperfect alignment, asymmetrical vibrations and

electron-electron interactions. The result will be that the' pair of electrons in the

Tr-a(H2) branch will eventually switch to the a"(H2)+TT" branch whereupon they may be

delivered to the a(e) levels of ethane. The branch switching and electron flow as indicated

schematically in figure 19 by the dotted curve and the solid arrow.

There still remains a sizable energy barrier of the order of 50 to 100 Kcal/mole

which is a direct result of the failure of the hybridized reactant states to correlate

properly with the ground state orbitals of the product. Consideration of grossly asymmetrical reaction configurations does not improve the situation because then geometrical

barriers can arise which are larger than the activation energy for the symmetrical case

[29].

The assumption that the electrons do not switch branches near cross-over in figure 18 does not rest entirely upon symmetry considerations. Suppose that transfer from Tr-a(H2) to

o'''(H2)+T7" does occur. Because the crossing is so near to the a"(H2) and the tt energies the

approximate electronic configurations after electron transfer wouId be a(H2)2a'''(H2)2 for the

hydrogen molecule and ir^ for the ethylene molecule. These configurations correspond to ^2^'

and C2Ht^2+ ^j^^j thus a large energy would be required for the creation of such species

.

This is not apparent in a one-electron energy level diagram such as figure 18 because the

effects of changes in orbital populations are not included. If such effects were included the actual crossing of energy levels would occur near the apex of the dotted

correlation curve in figure 18. At the apex of this curve the molecules are sufficiently

close to each other that the energy loss in H-H bonding and in it bonding is compensated by

the formation of the new C-H bonds. The charge in the a"(H2)+7r" branch now resides in

orbitals which approximate those o^ an ethane molecule with extended C-H bonds rather than

in orbitals resembling those of H2 ~ or C^H^^t^

Now consider the reaction as it occurs on a catalytic substrate. When the reactants

come in contact with a solid surface the ionization potentials will be modified due to

42
molecular relaxation [30] and screening [31]. For example, DeMuth and Eastman found that

the effective ionization potential of the ir-level of ethylene on nickel was shifted by

2.1 eV towards the vacuum level. This type of shift is separate from bonding shifts due to chemisorption.

Similar relaxation shifts were also found for benzene and acetylene on nickel. As an estimate of these effects we shall assume that all of the relevant molecular levels are shifted by 2 eV towards vacuum as shown in figure 19. The xz, yz surface state of the

TiO^ cluster located at -5.5 eV are shown before interaction in figure 19a.

6 The strongly bonding combinations of molecular and surface orbitals are it and oCH^)

with d and tt" and a^CHo) with d . Next, assuming that the surface states interact with

xz V-Z the appropriate molecular states Through a two-center transfer integral with a maximum

strength of 1 eV we arrive at the diagram shown schematically in figure 20. There are two important effects. First, we note that if the d state is occupied with electrons then

hybridization with the TT"+a"(H2) branch provides^a symmetry-allowed path for transfer of

electrons from the surface state into the a(e) levels of ethane. Second, due to hybridization of Tr-a(H2) with d electrons in this branch can be transferred from the

reactants to the surface state. Along the reaction coordinate both transfer process

occur so there is little or no net charge transfer between the surface states and the

reactant molecules. To emphasize this aspect we use the charge transferred from d into

a^(H^)+-n* as the reaction coordinate in figure 20. The charge transferred from TT-Xf'H2) ^°

d IS assumed to be equal. At the completion of the reaction the surface state is unchanged as it must be in a catalytic process.

Fig. 20. The effect of d-electron surface Fig. states on the hydro genation of

ethylene correlation diagram. The

dashed line is a schematic of the

total energy of the surface com- plex as a function of the charge

transfer, q.

21. The effect of d-electron surface

states near the (H2)* level on

the correlation diagram for the

dydrogenation of ethylene.

The total energy for the electrons involved in the surface complex is shown schematically by the dashed curve in figure 20. The diagram suggests that the activation energy

can be substantially lowered by hybridization with surface states of the proper symmetry.

We note that this catalytic mechanism requires that the surface states be partially ' occupied with electrons and that the effectiveness is increased as the surface state energy approaches the a"'(H ) energy. In figure 21 a correlation diagram is illustrated for

surface states which are close in energy to the T7-level.

43
4. Energy Band Theory of Chemisorptlon

Many of the transition metal oxides of interest have d-electron energy bands. A more realistic treatment of the mechanisms of chemisorptlon and catalysis should include the

effects of electron delocalization. In this section we describe a band theory model for

investigating chemisorptlon and catalysis on transition metal oxides such as the perovskites

.

The cluster model approximation is valuable in understanding the local electronic

structure of strong chemisorptlon bonds, but the effects of energy band formation are not

included. When the states of the substrate form bands of substantial width (several eV),

the nature of the hybridization of surface orbitals with molecular orbitals can be

altered significantly from that suggested by a simple cluster model. The energy levels

associated with chemisorption bonds can be displaced and broadened into bands. A greater

variety of symmetries are represented by band wavefunctions so that a greater variety of

geometries for molecule-solid interactions are possible.

These effects mean that the chemisorption process involves the formation of virtual

localized electronic states in which the bonds are shared by many substrate atoms because

of electron delocalization. The chemisorption energy can no longer be computed from a "single bond" model.

Various approaches to an energy band model have been employed recently [3U-37] which

make use of Green's function methods. The solid is usually characterized by energy band

states and the chemisorbing atom by a single atomic state. The interaction between the

molecule and the solid is described by matrix elements between the atomic state and each

band state.

There are several problems encountered in a band theory. First, the interaction

matrix elements are complex function of the k-vector and their form is difficult to obtain.

Often, the k-vector dependence is ignored and an average matrix element is employed as a parameter. Such approximations are very poor and obscure the nature of molecule-solid

hybridization--that is the actual mechanisms of chemisorption. A second problem is that

surface states rather than bulk states are involved in surface reactions. In principal

surface states are included in the usual formalisms. In practice, however, surface states

are seldom included in evaluation of the model. The symmetry properties of the d-orbitals

are often neglected in applications of existing band models.

The basic problem in beginning with diates characterized by a wavevector for the solid

is that it is a poor starting point for describing localized surface interactions which

lead to the formation of localized electronic states (chemisorption bonds). A reasonably accurate picture of the localized interactions, hybridization, and bonding

in terms of molecular orbitals already exists. Therefore a suitable model for describing

chemisorption should take advantage of this knowledge. This can be accomplished by

beginning from a description in which the solid and reacting molecules are treated on an atomic orbital basis. In such a basis a large amount of existing quantum chemistry can be

inserted directly into the model.

4.1. Localized Orbital Model for Chemisorption on a Perovskite Surface

We describe a relatively simple model for the interaction of a molecule with a perovskite surface. The oxide is represented by the cation d-orbitals and anion 2porbitals and their interactions. The molecule is represented by the atomic orbitals of

its constitutent atoms and by their mutual interactions. Since we may wish to be able to describe a process such as the dissociative chemisorption of H2 we want to employ atomic

rather than molecular orbitals because when dissociation occurs molecular orbitals

become Inappropriate states. The interactions between the atoms of the molecule and those

on the surface of the solid are also expressed in terms of atomic orbitals. In this

representation we can make use of quantum chemical calculations of atomic overlap and

transfer integrals and their dependence on distance. The symmetry constraints which

dominate orbital interactions are also easily accounted for in this approach.

44
The Hamiltonian for the system is

(4.1.1)

The component H^^ describes the molecule, describes the substrate and H describes the

molecule-solid interactions. The Hamiltonian for the atoms of the molecule is

H = Ze n.+^ZU.n.n. +E V..„a'^.a„. (4.1.2)

M . a aia 2 . ai aia ai-a . . „ iiaS aia S>10 a^a aja iiag

where e is the effective ionization energy for an electron in an atomic orbital of the

molecule. The subscripts a,j,a designate the atoms (a=l,2,...), the orbital (j = s, p , p , p , etc.), and the spin respectively. Operators a"*" . and a . create and destroy

eiectfons in localized atomic states. The quantity U .°'-'is the 8iulomb repulsion integral

for electron-electron interactions and the operators n"' . are number operators for each

atom and spin-orbital and V. . are the interactions mairix elements between the atoms of

the molecule which produce molecular states. In the absence of substrate interactions H is an approximate Hamiltonian for the molecule. M

We also treat the solid in an atomic orbital representation in terms of linear

combinations of atomic orbitals or Wannier functions. For H^ we have:

H^ = H^ + H + H ^, (4.1.3)

S d p pd

H, = Z e c^, c„. + ^ U , Z n„, n^, , (4.1.4)

d dk Rka Rko 2 d Rka Rk-o

Rka Rka

H = Z e d^, d^, , (4.1.5)

p pm R'ma R'mo ^ R'ma ^A=k ^ c^, d^, + c.c.} (4.1.6)

2 RR'km Rka R'ma

e is the effective ionization energy of the d-electron in one of the orbitals including

the Madelung potential and electrostatic splitting (£^|^ = ^^^2 ^ °^ ^ ^

~ xy,xz, yz

or k = x^-y2 or z^) and the operator and c^^^ create and §estroy electrons respectively

in d-orbitals with spin a at a lattice site located at R. The Coulomb integral is and

"^Rka number operator for the d-ion at R. The Hamiltonian for the oxygen ions is given by (M-.1.5). Here e is the effective affinity energy for adding an electron to a 2p-orbital (m = p ,p and^p ) on an 0 ion to create an 0 ion and includes the Madelung

x v z potential and electrostatic static splitting; e = e or e for orbitals parallel and

perpendicular to the M-0 axis in a ABO perovsk?¥e. ?he operators d+, and d^^, ^ create

and destroy respectively electrons with spin a in 2p-orbitals at a lattice site located

at R' .

The quantities T^^dij^ two-center integrals for the interaction of d-orbital of

k-type at lattice site^ with a p-orbital of m-type at lattice site R'- The elements of

^R'km two-center transfer integrals such as pd-rr and pda. Cation-cation or anionanion interactions between atoms in different unit cells may also be added in a similar

fashion.

The eigenstates of H^ (in the absence of H^^^) are the energy bands of a finite solid

perovskite including surface states. Our previous work has shown that this model leads to

a reasonably good representation of d-band perovskites [10,11].

The interaction between the atoms of the molecule with those of the substrate is

written as

45
^

. (4.1.7)

+ E {T , d^, a . + c.c}

. R'amo R'ma ana

R'aman

s

The two-center integrals T^^ and R^,

^ correspond to the parameters for interaction

between the atoms of the moiecule ana ??ie cations and anions of the substrate. The sum over R and R' in eq . 4.1.7 refers to a sum which includes only a few substrate surface

s s atoms; that is, only those surface atoms which have a direct bonding to the atoms of the

interacting molecule. These parameters are strong functions of the reaction coordinate

since they depend on the atom-atom internuclear separation and on the atom-substrate

geometry. The results of a bond order model and the cluster calculations can be used to determine these parameters as a function of the bond lengths.

The operators a, c and d are Fermions and obey the usual anticommutator relations.

The effect of molecule-substrate overlap has been stressed by Schrieffer and Corner in

their "Induced Covalent Bond" model [38]. An approximate method for including overlap

has been suggested by Madhukar [39] in which the commutator relations are modified. In this

scheme

[a . , , ,] = 6 ,D ^ , . (4.1.8) aia R ka' aa' aR ik

s s .

[a . , d"!", ,] = 6 ,D „ . , (4.1.9) ana' R'ma' aa' aR nm'

s s

where the D „ and D are atom-substrate overlap integrals which vanish except for

aR ih aR'nm c- o t- s s-* those surface atoms which directly participate in the chemical bonding.

The information characterizing the molecular reaction is contained in the matrix of

Green's functions. These functions [40] are defined in general by:

G(t)„ „ = -ie(t) < [A(t), B(t)]^>, (4.1.10)

G(u))^^3 = /_.e-'"'G(t)^^gdt, (4.1.11)

where A and B are creation and destruction operators and e(t) is a unit step function in

t (time).

d Ox M A supermatrix of such functions with diagonal blocks G (w), G (w), and G (oj) are defined by the elements:

G'^(t)R^R2k^k2aa' = -ie(t) < ^cM^^^^^(0)^^^^^,1^>

' • (4.1.12)

G°''(t)^,^, , = -ie(t) < [d(t)^, d^(0)^, ,] > (4.1.13) R^R^m^m^aa' Rjm^a , n:^m^a' +

M + G (t) = -ie(t) < [a(t) . , a (0) ,., ] > (4.1.14)

aa' j

j

' ao aja a ' ]

' a +

J ^ J Ox-d M-d Ox-M M-Ox

d— Ox d—

M

The off-diagonal blocks are G (w), G (w), G (w), G (w), G (oj) and G (w). A typical

element of an off-diagonal is, for example.

46
g" '^(w) , = -ie(t) < [a . (t), c^(0)^, ,]^> (4.1.15)

A matrix of spectral weight functions (or local density of states) is derived from these

Green's functions. Spectral weight functions for an atom of the molecule are, for example,

p(a)) . = -Im G(a)) . . (4.1.16) aja aa]]aa

For the isolated molecule p . (to) will consist of a (normalized) weighted sum of deltafunctions at energies corresponding to the molecular states. When the molecule interacts

with the solid the character of p(a)) changes to reflect the formation of hybridized

molecule-solid bonds. The spectral weight function can possess both bands and deltafunctions which are peaked at energies different from those of the free molecule. The

broadening of a peak into a band occurs when a localized state is formed at an energy

within the energy range of the bands of the substrate

.

The bonding among substrate atoms is also altered by the chemisorption process. This

mation is contained in the

p ((»))„, , which are defined by R ma

information is contained in the spectral weight functions of surface atoms, p(a))j^j^^ and

p(to)„, = g'^(u)__,, (4.1.17) Rko RRkkaa

p(u))^, = G°''(a))^,^, (4.1.18) R'mo R'R'mmao

For the non- interacting substrate p(a))jy^ (and P^'^^j^ij,,^) have a distribution of spectral

weight characteristic of the admixture o? d-orbital (p-orbital) in the valence and conduction

bands including surface states. When interaction occurs the spectral weight of the surface

atoms is shifted. Some additional density now occurs at the energies of the chemisorption

bonding and antibonding states and some density reduction occurs at substrate band

energies. This shifting of spectral weight characterizes the participation of substrate

electrons in covalent bonds formed with the atoms of the reacting molecule. A schematic

of these features is shown in figure 22.

The energy of chemisorption (heat of adsorption) E^, can be determined from

E = E° + E° - E

where E° and E° are the total energies respectively of the non-interacting molecule and

substrate and E is the final energy of the interacting systems after chemisorption. This

quantity is determined by

E^„ = -Im/ d(o uTr{G(w)}, (4.1.19)

SM -°°

where E^ is the substrate Fermi-energy, Tr indicates the trace of a matrix and G(a)) is the

entire super matrix of Green's functions.

The formal application of the model is fairly simple, but obtaining self-consistent

solutions is a major problem. The Green's function matrix elements can be calculated if

suitable approximations are made to treat the electron-electron terms. Assuming suitable

approximations are made, the Green's functions may be calculated in terms of certain (unknown) quantities such as the electronic occupation of various eigenstates. The actual

energy of these states is dependent on a knowledge of their occupation because of the

electron-electron interaction terms . The occupations in turn are dependent upon the

47
o

Q.

ANTIBONDING

BONDING

/ STATE

La.

Fig. 22. Schematic of the local density of states function of a cation, (a)

the spectral weight function for

a state of an interior cation,

(b) the spectral weight function

for a surface dj^^ state before

interaction with a molecule, (c)

the spectral weight function for

a surface d^^ state interacting

with a molecule.

spectral weight functions. For example, the number of electrons in the jth orbital of

the a-th atom with spin a of the reacting molecule is given by:

^ja = L da)p„j^(u3) , (4.1.20)

where p . (w) is a function of n . . Thus an iteration procedure is required to obtain

self-consistent solutions for the-^Green's functions.

The approximation used in treating the electron-electron terms in most band theorymodels is the Hartree-Foch approximation. This leads to the replacement of the ionization

and affinity levels by a single average level. As discussed recently by Madhulkar and

Bell [41] the Hartree-Foch approximation can lead to an incorrect structure for the single

particle Green's function. A more complete treatment leads to the preservation of both

the effective ionization and affinity levels.

It is complicated but straight-forward to apply the formal theory to problems of

chemisorption of diatomic molecules on a perovskite surface. We are currently using this

model for the investigation of the dissociative chemisorption of and O^- These studies

will be discussed elsewhere.

4.2. Relation of the Band Model to the Cluster Model . .

The application of the theory to any realistic situation requires fairly complicated

numerical calculations. However, using a highly simplified version of the theory the

48
connection between the band model and local models such as were discussed m section 3 can be demonstrated. Our purpose here is to illustrate how localized chemical bonds appear in a band model. In particular, we consider the hybridized states which an OH molecule forms

with the t2g d-orbitals of the solid. These states are represented by the levels b and c of figure 15

.

We consider the simplified Hamiltonian:

H = H + H + H , (4.2.1)

s m sm

H = Z E,(R)c"^(R) c(R) + Z E d"^(R' ) .d(R' )„

= ct,R ^ " " BR- P ^

^

+ E {T(RR') ^c'^CR) d(R'). + H.C.}, (4.2.2)

otB a 3 ctBRR'

H = Z E a"^a (4.2.3) m Y m Y Y

H = Z V a"^c(0) + V"^ c"^(0) a (4.2.4) sm YC YQ' Y ot ya y a.

In eq. 4.2.2, E,(R) = E^ for interior cations and E, + AV„ for cations on the surface

. d d d M where AV^^ is the change m tne Madelung potential at the surface. The index a designates

xz, yz or xy d-orbitals and R runs over the cation sites. The index 3 designates p , p or p orbitals and R' runs over the anion sites and E is the anion energy including ^ eiectrostatic potentials. The interaction between^near neighbor cation and anions is

represented by T(RR' ) which is equal to the two-center transfer integral (pdTt) or vanishes

depending upon the relative orientation of the p and d-orbitals.

The non-bonding p-orbitals of the OH molecule are represented by y = P or p and

their energy before interaction with the surface but including electrostatic potentials is E . These non-bonding orbitals can ir-bond to a surface cation. The interaction , represented by V 5 is equal to V or vanishes depending upon y and a. The chemisorption is assumed to 2ccur at R = 0 which designates a surface cation site.

The details of the calculation of the various Green's functions will not be described

here. Instead we present some of the results and discuss the interpretation. The Green's

function for the molecular levels

G^Co)) = {oj-e -v^(u)-e )G°(w)} ^ m p

G0(<,) = ^f^ 271 -IT Ap(w)-e

-ie(a),k)

(4.2.5)

(4.2.5)

In equations 4.2.5 and 4.2.6 u is expressed in units of (pdw) and the parameters are defined by the following equations

e = E /(pdTT); v = V/(pdTr); Av^ = AV^/(pdTr)

p p

^ J'

' MM

e,dd^ = E-/(pdTr); -2cose = (w-e,)(a)-e dp)-4+2cosk; e mm*^ = E /(pdir);

Ap(w) = l+(e -u)Av,,. (4.2.7)

p M

49
The function G°(w) has several important properties. Consider the denominator of the

integrand in equation (4.2.5). We have shovm in our work on surface, states [10] that the

condition for the existence of a surface state with wavevector K is

Ap(a)) = e -i6((.,k) (4.2.8)

The function e

"'' is real in the energy gap between the d and p bands of the solid.

Solutions of equation (4.2.8) as a function of k determine the surface state energy band

w (K). Thus the poles of the integrand of G°(a)) correspond to the surface energy bands.

Since these poles form a continuum as a function of k the existence of a surface band

produces a real and imaginary contribution to G°(u)). The imaginary part is directly related

to the density of surface states and therefore vanishes for w outside of the surface band

frequency regime.

— iG

Second, the function e is complex for u inside the bulk energy bands. Consequently

G°(a)) will also have an imaginary part for oj coincident with a bulk energy. In this

energy range the imaginary part is related to the bulk density of states. Therefore we express G°(w) in the form:

G°(a)) = Rioi) + inioi) (4.2.9)

where N((ij) is related to the bulk density of states or the surface state density of states

in the relevant range of w . In the band gap but outside of the surface state bands

N(co) = 0.

Now we consider the molecular Green's function given by equation (4.2.5). It may be

written in the form

M -1 G (w) = {w-e -A(aj)-ir(a))} (4.2.10) m

where

A(u3) = v^ (oj-e )R(a)), (4.2.11)

P

r(w) = v^(a)-e )N(a)) P

M The density of states p (w) for the chemisorbed OH molecule is given by:

p^a)) = i ImG^a,) = ^ ^-i^^ ^ (4.2.12)

[co-e - A(a))] + r(a)) m

M The resonances of p (w) correspond to the formation of localized chemisorption states which

for this example are admixtures of the p and d orbitals of the solid with the non-bonding

orbitals of the OH molecule. In the non-interacting case v = 0 and p^(a)) 6(a)-e ). This

. m means a discrete level at the assumed molecular energy E . When interaction between the

molecule and the surface occurs, v ^ 0, then A(a))?!0 and ?he molecular energy is shifted.

The function A(a)) describes the hybridization of the molecular level with the continuum of

bulk and surface states. The new state will be discrete if T(ui) vanishes at the o) for which

w = £ + A(aj ) (4.2.13)

r m r -

50
If r(a) ) ?^ 0 then the localized state is broadened into a band whose width is governed by rCu). Since r(a)) is a strong function of w the resonance in pM((jo) does not have a simple shape.

We may conclude from equation (4.2.12) that localized hybrid states which resemble

those described for the cluster model in section 3 will occur if the states are formed

in the energy gap but not within the surface bands . When this condition is not met the

chemisorption bond is shared by many surface atoms--that is, delocalized. The degree of

localization depends on rCw). The smaller r(a)) is the greater the degree of localization.

M It should be noted that in general p (cj) will have several resonances corresponding

to bonding and antibonding combinations of molecular and band states.

The Green's functions in this example may be calculated and the various functions

evaluated. Typical solutions have the form shown schematically in figure 23 for the

parameters (pdir) = V = 1 eV and E = -5 eV corresponding to OH chemisorbing on a Ti ion

of SrTiO- on a (001) surface. ^

-4.0

-5.0 J

/ BAND

/

/ I /

-6.0

> -7.0-

>-o

liJ

Surface/

BAND

I

-8.0-

-9.0-

-10.0-

-II. 0--

DELOCALIZEO

ANTIBONDING

STATE

LOCALIZED

BONDING

STATE

DELOCALIZED

BONDING

STATE

(OJ)

Fig. 23. Schematic of typical solutions to the Green's function model for

chemisorption of an 0H~ molecule

on a surface cation, p^(cj) is the

molecular density of states

function.

The qualitative features of figure 23 are to be compared with those of figure 15. The localized bonding state corresponds to level c and the delocalized antibonding state

to level b of figure 15. It is clear that the hybridization mechanisms operating in the

energy band model are the same as those discussed in the cluster model. It is equally

clear that the effects of band formation can not be neglected in any quantitative

treatment

.

51
5. Summary

In this paper we have attempted to introduce the concepts which are believed to play

an important role in surface chemistry on transition metal oxide materials. The models

are crude and refinement is needed in order to produce quantitative results. Such calculations may be expected in the near future. A strong driving force for increased theoretical

effort will develop as the new experimental techniques of surface physics are turned to the investigation of transition metal oxide surfaces.

6 . References

[I] L.A. Pedersen and W.F. Libby; Science 176_, 1355 (1972).

[2] R.J. Voorhoeve, J. P. Remika, P.E. Freeland and B.T. Matthias; Science 177 , 353 (1972).

[3] R.J. Voorhoeve, J. P. Remika and D.W. Johnson, Jr., Science 18_, 62 (1973).

[4] W.F. Libby and S. Aegerter (unpublished report).

[5] A.C.C. Tseung and H.L. Bevan; J. Electroanal. Chem. and Interfacial Electrochem. M-5, U29 (1973).

[6] A. Fujishima and K. Honda; Nature 238 , 37 (1972).

[7] J.G. Mavroides, D.I. Tchernev, J. A. Kafalas , and D.F. Kolesar; Mat. Res. Bull. 10_, 1023 (1975).

[8] J.B. Goodenough; Progress in Solid State Chemistry 5_, 149 (1971) (ed. H. Reiss).

[9] E.A. Kraut, T. Wolfram and W.F. Hall; Phys. Rev. B6_, 1499 (1972).

[10] T. Wolfram, E.A. Kraut and F.J. Morin; Phys. Rev. B7, 1577 (1973).

[II] T. Wolfram; Phys. Rev. Letters 29_, 1383 (1972).

[12] L.F. Mattheiss; Phys. Rev. 181 , 987 (1957); Ibid B2_, 3918 (1970); Ibid B4, 3548 (1971).

[13] M. Cardona; Phys. Rev. A140 , 551 (1965).

[14] J. A. Tossel, D.J. Vaughan and K.H. Johnson; Am. Mineralogist 59_, 319 (1974).

[15] T. Wolfram, R. Hurst and F.J. Morin, "Surface States on SrTiO^ , BaTiO^ and TiO^" to

be published.

[16] K.H. Johnson and R.P. Messmer, J. Vac. Sci. Technol. 11, 235 (1974).

[17] T. Wolfram and F.J. Morin, Appl. Phys. 8_, 125 (1975).

[18] F.J. Morin and T. Wolfram; Phys. Rev. Letters 30_, 1214 (1973).

[19] V.E. Henrich, G. Dresselhaus and H.J. Zeiger; Thirty-Fifth Annual Conference on Physical Electronics, Penn. State University, June 2-4, 1975.

[20] S.N. Frank, D. Laser, K.L. Hardee and A.J. Bard, "Characterization of Semiconductor

Electrodes in Nonaqueous Solvents", presented at the Workshop on Electrocatalysis on Non-Metallic Surfaces, National Bureau of Standards, Gaithersburg , Md. , Dec. 9-12,

1975.

[21] F.H. Field and I.L. Franklin, "Electron Impact Pehnomena", Academic Press, New York,

1957.

[22] J.W. Maskowitz and M.C. Harrington; J. Chem. Phys. 42, 3512 (1955).

[23] H. Weinberg and R.P. Merrill; Surf. Sci. 33_, 493 (1972); H. Weinberg, M. Lambert,

CM. Comrie and J.W. Linnett, Surf. Sci. 30_, 299 (1972).

[24] S.W. Weller and S.E. Voltz, J. Amer. Chem. Soc. 75_, 4595 (1975).

[25] R. Hoffmann, J. Chem. Phys. 39_, 1397 (1963).

[25] H.A. Germer, J. Chem. Phys. 58_, 3524 (1972).

[27] J.W. Moskowitz, M.C. Harrison, J. Chem. Phys. 42, 1726 (1965).

[28] R.B. Woodward and R. Hoffmann; The Conservation of Orbital Symmetry ( Verlag-Chimie

,

Weinheim, F.R., Germany, 1970).

[29] See for example, hydrogen exchange reaction.

[30] D.R. Hamann, Chem. Phys. Lett. 22, 301 (1973).

[31] D.A. Shirley, Chem. Phys. Lett. 16_, 220 (1972).

[32] J.E. DeMuth and D.E. Eastman, Phys. Rev. Letters 32_, 1123 (1974).

[33] F.D. Mango and J.H. Schachtschneider , J. Am. Chem. Soc. 89^, 2484 (1967); F.D. Mango,

Adv. Catalysis 20_, 231 (1969).

[34] D.M. Newns, Phys. Rev. 179 , 1123 (1969).

[35] T.B. Grimley, J. Vac. Sci. Technol. 8_, 31 (1971).

[35] T.B. Grimley in "Molecular Processes on Solid Surfaces", McGraw-Hill,. New York (1969).

[37] J.W. Gadzuk, "Interaction of Atoms and Molecules with Surfaces" in Surface Physics of

Crystalline Materials , Academic Press, J.M. Blakely, Ed. [38] J.R. Schrieffer and R. Corner, Surf. Sci. 25_, 315 (1971).

[39] A. Madhukar, Phys. Rev. B8_, 4458 (1971).

[40] D.N. Zubarev, Sov. Phys. Uspechi 3_, 320 (1950).

[41] A. Madhukar and B. Bell, Phys. Rev. Letters 34, 1631 (1975).

52
National Bureau of Standards Special Publication 455, Electrocatalysis on Non-Metallic Surfaces,

Proceedings of Workshop held at NBS, Gaithersburg, Md. , December 9-12, 1975. (Issued November 1976)

Hartree-Fock Calculations of Hydrogen Adsorption on Nickel, Copper, and Magnesium Oxide

A, B. Kunz, M. P. Guse, and R. J. Blint

Department of Physics and Materials Research Laboratory

University of Illinois

Urbana, Illinois 61801

The potential energy curves for NiH, HNiH, CuH, and HCuH resulting from the ground and first excited state atomic configurations are calculated using the unrestricted Hartree-Fock and generalized valence

bond methods. NiH bonding in the 3d^4s2 configuration and CuH bonding

in the 3d^4s2 configuration is by means of an sp hybrid orbital which

comes from the 4s ^ shell leaving a singly occupied orbital free to bond to another hydrogen atom. The bonding in the NiH3d^4s and

CuH3d-'-*-'4s configurations is to the 4s orbital leaving in NiH a hole in

the nickel 3d shell which interacts with another hydrogen only weakly

and in CuH a closed 3d shell resulting in a repulsive potential energy

curve. The interaction of hydrogen with a cluster representing a magnesium oxide surface Vj center was calculated using UHF. The lowest

energy results when the Mg and 0 are singly ionized rather than doubly

as in the bulk resulting in a surface radical. The hydrogen is bound

over one of the oxygen atoms.

Key words: Adsorption; catalysis; copper hydride; magnesium oxide;

molecular structure; nickel hydride.

INTRODUCTION

The interaction of hydrogen with surfaces is important in a multitude of catalytic

processes. In the hope of learning something about this interaction for at least a few

cases we have studied the much simpler problem of hydrogen bonding to small molecules

approximating the essential features of the surface. We have used self consistent

Hartree-Fock methods to calculate potential energy curves for hydrogen interacting with

Ni, NiH, Cu, CuH, and a small MgO cluster.

Two methods of calculation were used in this work, unrestricted Hartree-Fock (UHF) and

perfect pairing generalized valence bond (GVB) . The UHF wavef unction is

- A[, 1* ^i+r-- n 1 1 1+1 (1)

where A is the antisymmetrizer , the ({is are the spatial one electron orbitals , and the as and Ss are the corresponding spin functions. The one electron orbitals need not be doubly

occupied so that the spin up and spin down electrons can occupy different spatial orbitals,

The ($s are varied self consistently with no restriction on their symmetry.

The perfect pairing GVB wave function

(2)

This research was supported by the National Science Foundation under Grant DMR 72-03026.

53
is in effect a two-configuration, optimized wavefunction where x is the spin eigenfunction

[1].^ All orbitals are restricted Hartree-Fock in nature except for the pairs which are GVB split. In both UHF and GVB methods, bonds are allowed to dissociate properly.

The valence shell of a nickel atom in its ground state ( F) has the configuration

Sd^As^. The first excited state (3d) is only about .03 electron volts (eV) higher in

energy and has the configuration 3d^4s . Our Hartree-Fock calculations find the Sd^As state

to be about 1.7 eV higher than the 3d^4s2 state. Consequently we have considered NiH and

HNiH resulting from both of these configurations.

The potential energy curves of Ni(3d^4s2)H were calculated using UHF and are shown in

Figure la. From Ni(3F) and H(2s) we can form NiH with spin 3/2 or 1/2. The high spin

state in Figure la is a nearly pure quartet while the low spin state is a mixture

. 8l| l/2>+. 58| 3/2> . At the potential minimum the energy of this state is lower than the pure

doublet state by about .07 eV.

-1505.71

--I505.72

a

- -1505.7 3

- -1505.74

--j505.75

-^505,76^

-1505.7?/

i\ rr

High Spin

-I505;79

Low Spin

Low Spin

1 1 1 1 1 , /H506.30-

1 1

y -1506.31 -

' b C

/ -1506.32-

/ -1506.33-

-1506.34-

-1506.35-

-1506.36-

-1506.37-

1 1

-1506.38-

1 1 1

1 1 2 3 4 5 6

RNiH(NiH)

3 4 5 6

RfJiH (HNiH)

oX

7 0 10 20 30

Bond Angle (degrees)

Fig. 1. In part a, the bonding curves for NiH in the doublet and quartet states are shown

as a function of Ni-H separation. In part b, the bonding curve arising from forming a linear HNiH molecule from the doublet NiH molecule is shown. The one H is

held at the NiH equilibrium separation and the other one is moved around. In part

c, the resistance to bond bending of the HNiH molecule is seen. The NiH distances

are held to those for minimum energy. Energies are in hartree (IHy = 27.2 eV)

,

length are in a.u. (1 a.u. = 0.529 X).

In order to determine the nature of the binding in NiH we examined the valence orbitals

of the molecule for large and small internuclear separations. The orbitals for the high

spin state are almost identical to those of the low spin state, which are shown in Figure

2. It is immediately clear that the nickel 3d orbital does not contribute significantly

to the bonding. This agrees with the findings of Blyh'older [2] and Melius [3] for the

adsorption of hydrogen on nickel. As the internuclear separation decreases the doubly

occupied nickel As^ orbital splits into two essentially sp hybrid orbitals. One of these

forms a covalent bond with the hydrogen atom while the other becomes a singly occupied non- bonding orbital directed away from the hydrogen. This is identical to the bonding we found

in ScH and MnH [4-5].

A second hydrogen atom interacts strongly with the NiH radical to form a linear HNiH

molecule. Figure lb shows the potential energy curve for H approaching the low spin

Figures in brackets indicate the literature references at the end of this paper.

54
.35

(A) 0

-.35

.70 -

(B)

-.70

.35

(C) 0

-.35

1 A

Ni H r . H. -4-2 0 2 4 6 8 -4 -2 0 2 4 6 8

Z Axis in bohr

Fig. 2. Amplitudes of the three valence orbitals as a function of inter-nuclear separation. A is the doubly occupied Ni4s^ orbital, B the Ni 3dz2 orbital and C the

hydrogen Is orbital. The origin is at the Ni atom and the plane includes the Ni

and H nuclei.

Ni(3d84s2)H radical calculated using UHF. The ground state is a nearly pure spin triplet

with the hydrogen binding to the NiH non-bonding Asp hybrid orbital. The energy required

to bend this second bond while keeping the Ni-H distance fixed is shown in Figure Ic.

The potential energy curves of Ni(3d^4s)H were calculated for the doublet state using

both UHF and GVB and are shown in Figure 3. The full curve is shown for the ^I"*" state while

the energies of the ^ir and GVB states are shown at 3.0 Bohr. Taking into account the

experimental fact that the Ni3d^4s and 3d^4s2 states are nearly degenerate we find the

ground state of NiH to be the state resulting from the Ni 3d^^s configuration. In each

of these states the bond involves the Ni4s and H Is orbitals leaving a single hole in the

nickel d shell.

The H Nl(3d^4s)H potential energy curve was calculated using GVB with both one and two GVB pairs. The initial NiH state is 2j^- with the second hydrogen Is orbital GVB paired with

the Ni 3d22 orbital to form the state of HNiH. Both curves are shown in Figure 4. It may be seen that hydrogen interacts very weakly with the 3d electron showing that the ground

state of HNiH is derived from the 3d^4s2 configuration of nickel. We are unable to find

solutions to H-NiH arising from the 3d^4s state of Ni for H-NiH separation less than shown

in Figure 4 because the solution jumps to the branch derived from the 3d^4s2 Ni configuration which has lower energy. We have also studied H adsorption by Ni bonded to a surface

of Si02 and find the chemisorption behavior in these cases to be similar to that for the

H-NiH system.

A similar study was done on copper using GVB. The CuH molecule may have both a -^l"^ and ^Z"*" state arising from the Cu 3d^^^s ground state. As is seen from Figure 5a only the

55
-1505.69

.70 „ .71 h r .72 oX

X.73

CI>

i5 -74

1.75

oI .76

.771-

(UHF Energy ot co )

Fig. 3. The potential energy curves for

Ni(3d94s). The energy is in

hartrees and the lengths are in

bohr

.

4 5 6

Atomic Separation (bohr)

-1506.24r

^ .25-

i.26

.27

5 .28

"5 .29

GVB Pair

4 5 6

(H Ni)-H Seporotion (bohr)

-2 GVB Pairs

Fig. The potential energy curves for

a second hydrogen on the 2^ state of NiH (Sd^As) in the per- fect pairing GVB method. The

energy is in hartrees and the

lengths are in bohr.

state arising from the Cu ground state is bonding. The ground state 3j;+ ig antibonding

and this is in contrast to the first period transition metals for which all states arising

from the atomic ground state are bonding. The first excited state of copper has the con- figuration 3d^4s2. The excited CuH constructed from this configuration can exist in the state. The potential energy curve for this state has a deep minimum. In making up Figure 5 the experimental energy difference for the Cu atom 3dlC)4s and 3d^4s2 ^j-g used to determine the infinite Rq^j^ energy separation, rather than the less accurate theoretical

value. When we bring a second hydrogen up to CuH we find that the linear HCuH molecule

exists only when it is derived from the Cu 3d^4s2 level as seen in Figure 5b. The form of

bonding is identical to that of Ni(3d^4s2)H. We see that a hydrogen atom approaching a copper surface must overcome an activation barrier in order to bond to the 3d^4s2 configuration. This is consistent with the experimental observation that Cu adsorbs dissociated H but is not able to adsorb H2 . ,

The above calculations suggest that hydrogen does not bind to a nickel surface by means

of the 3d-electrons but rather by means of the 4s-4p electrons. If a surface nickel atom

exists in the 3d^4s2 configuration one of the 4s electrons may be involved in the nickelsurface bond while the other lies in a nonbonding orbital forming a surface radical to which

a hydrogen atom may bind. The strength of this hydrogen-nickel bond will be just over half

that of the H2 bond in the Hartree-Fock limit (3.4 eV) so that an H2 molecule will tend to dissociate with each hydrogen bonding to a nickel atom. For a surface atom in the 3d94s

configuration the 4s electron is involved in the nickel-surface bond and the d electronhydrogen interaction leads to at most weak adsorption. We might then guess that the cata- lytic ability of a transition metal would correlate with the ability of its atoms to exist

in the d'^S^ configuration.

56
2 3 4 5 6

RCuH(CuH)

Fig. 5. In part a, the bonding curve for CuH

is shown for the singlet and triplet

state. In part b, the curve arising

from forming a linear HCuH molecule

from the ground state of CuH is

shown. The first H is held at the

CuH equilibrium separation and the

other one is moved around. Energies

are in hartrees and length in a.u.

RCuH(HCuH) RCuH(HCuH)

This mechanism is consistent with the constancy of relative catalytic activity across

the first transition metal series since the ground state for these atoms is already the

3d^4s^ configuration. The ground state of atoms in the second transition series is predominantly Ad'^Ss . We would then expect the catalytic activity of the second transition series

to depend on the ease with which the atoms are excited to the ^^6P~^5s^ configuration. The

minimum energy for this promotion occurs for the middle of the second transition series [6]

and the Boltzman occupation factor for the 4d'^~"'"5s2 state correlates directly with the relative activities for hydrogenolysis of ethane as found by Sinfelt [7]. This is seen in

Figure 6. The experimental data are due to Sinfelt and the specific activities are with

respect to Cu for the first transition series and Pd for the second transition series. In the theoretical determination, we assume the activity is directly proportional to the probability of occupation of the d^s^ state of the atom. As is seen from Figure 5b for Cu,

the actual barrier is less than the atomic ground state to d'^s^ state energy splitting. We

use a factor of 1/3 the atomic energy value in preparing this curve. This factor of 1/3 is our only parameter and is used universally for Cu and the second transition series elements.

T is 205°C as is the case experimentally. This then seems capable of predicting the strong

dependence of activity with element for the second transition series.

9 I

1 1 1 1 1 1 1 1 1 r

Column in Periodic Table

6. The experimental activities for hydrogenolysis of ethane are shown [2] . The theoretical relative activities are also given

on the assumption the activity is proportional to the occupancy of the d^s^ energy

state. This probability is proportional to C-A/kT_ yg 3 ^ which is 1/3 the atomic

excitation energy from the ground state and

T is 205°C as it is experimentally. The

theoretical rates are relative with that of Cu being 10 for the first transition series

and that of Pd being 10 for the second

transition series.

Magnesium oxide catalyzes the H2-D2 equilibration reaction at low temperatures. Boudart

has studied this reaction [8] and determined that the active catalytic site involves a paramagnetic triangular array of Q-ions situated on a (111) plane of MgO which he calls a Vj center. In order to study the interaction of hydrogen with this site we constructed a cluster consisting of an equilateral triangle of oxygen .atoms overlaying a similar array of

magnesium atoms rotated 60° (Table 1)

.

57
Table 1. The positions of atoms in the MgO cluster.

X y z (bohr)

Ol 00000 -3 24533 0 00000

02 2 81054 1 62266 0 00000

O3 -2 81054 1 62266 0 00000

Mgl 0 00000 3 24533 -2 29479

Mg2 -2 81054 -1 62266 -2 29479

Mg3 2 81054 -1. 62266 -2. 29479

We began with the system in the Mg++0- state as it is found in the bulk and calculated

the potential energy curve of H along the cluster axis using UHF. We found that this is an excited state. The state Mg"'"0 is over 3 eV lower in energy. This suggests that the

surface radical is energetically favorable. The potential energy curves for bringing H

4.0

3.0

o H above center of 0^ A H above O2 Midpoint H above 0 Atom

30"

30"

Fig. 7. Potential energy curves

for hydrogen on MgO Vj

cluster. The vertical

axis gives binding

energy in electron

volts. The horizontal

axis gives the height

of the H atom above

the O3 plane in bohr.

2.0 3.0 4.0 5.0 6.0 7.0 8.0

Height of H atom above O3 plane (bohr)

9.0 10.0

down over various points on the cluster are shown in Figure 7. The lowest energy results

when the H sits directly over one of the oxygens. This indicates that the V-j- site may be

able to accomodate three hydrogen atoms. This study seems to be fully consistent with the

results of Boudart. There is one important similarity between the behavior of the MgO

cluster and the supported Ni system. In both cases the H atom is found to adsorb into an existing "surface radical".

REFERENCES

[1] Hunt, W. J., Hay, P. J., and Goddard, W. A., J. Chem. Phys. 97, 738 (1972).

[2] Blyholder, G., J. Chem. Phys. 62, 3193 (1975).

[3] Melius, C. , 1975, Private Communication.

[4] Kunz, A. B., Guse, M. P., and Blint , R. J,, J. Phys. B: Atom. Molec. Phys. 8, L358

(1975) . [5] Blint, R. J., Kunz, A. B., and Guse, M. P., Chem. Phys. Lett. 36, 191 (1975).

[6] Moore, D. E., Atomic Energy Levels , (NSRDS-NBS Circular 35, Washington, D.C.),

Vol. I-III.

[7] Sinfelt, J., Advances in Catalysis 23^, 91 (1973).

[8] Boudart, M. , Delboville, A., Derovane, E. G. , Indovina, V., and Walter, A. B. J. Am. Chem. Soc. 94, 6622 (1972).

58
National Bureau of Standards Special Publication 455, Electrocatalysis on Non-Metallic Surfaces,

Proceedings of Workshop held at NBS, Gaithersburg, Md. , December 9-12, 1975. (Issued November 1976)

Stark Model of the Excitonic Surface States Observed on MgO

H. J. Zeiger, Victor E. Henrich, and G. Dresselhaus

Lincoln Laboratory, Massachusetts Institute of Technology

Lexington, Massachusetts 02173

Using the energy and angle -of- incidence dependence of the electron energy loss

spectra of MgO, we have separated the excitonic transitions from the Mg core levels

to the excited states into those of bulk and surface origin. The bulk transitions are very nearly those of the free Mg2+ ion. The (100) surface-state transitions can be

described by Stark- splitting of the energy levels of the surface Mg2+ ions in the intense Madelung electric fields at the crystal surface. The observations on other

crystal faces of MgO and in amorphous MgO, however, suggest the possibility of

the presence of surface Mg"*" ions. Key words: Bulk excitons; electron loss spectroscopy; ionic crystals; Madelung

potential; Mg"*"; Mg^"'"; MgO; Stark spectrum; surface excitons,

1. Introduction

While the electronic surface states of covalent semiconductors such as Si, Ge and GaAs have been

widely studied over the last few years, less attention has been paid to the surface electronic structure

of highly ionic materials. We have extended the technique of reflection electron energy loss spectroscopy (ELS) to the study of intrinsic surface states on MgO. By using both the angle-of -incidence and

primary -electron -energy dependence of the ELS spectra, we have been able to unambiguously separate transitions from the Mg(2p) and Mg(2s) core levels to excited states [1] into those of bulk and

surface origin. The energies of the bulk excitonic transitions are in good agreement with the optically

observed excited states of the free Mg2+ ion [2]. The surface-state transitions on a (100) face can be

described well by a simple, one-parameter model for the Stark splitting of the bulk Mg2+-like excitonic states in an electric field Fg at the position of a surface Mg^"^ ^'^^•o experimentally determined value of is 3.0 X 10^ v/cm, compared to the value of 3. 7 x 10 V/cm calculated for the electric field produced by the Madelung potential gradient. However, as we discuss below (see Section

4), observations on other faces of MgO and on amorphous MgO suggest the possibility that Mg"*" ions are present on the surfaces.

2. Experimental Observations

By varying the energy and angle -of-incidence of the primary-electron beam, a wide range of

effective depths can be sampled in ELS. The electron mean free path in a solid increases from 4-5 A

for energies between about 40 to 150 eV to 20-30 A for 2000 eV electrons [3]. Thus an ELS spectrum

taken with a primary -electron energy Ep ~ 2000 eV at normal incidence will be dominated by transitions of bulk origin, whereas a normal-incidence spectrum taken with Ep~ 100 eV will exhibit both

bulk and surface transitions. The contribution of bulk transitions to the spectrum can be virtually

eliminated by using a primary beam of ~ 100 eV at grazing incidence. We have used this approach to separate our spectra into structure of bulk and surface origin. By using a double-pass cylindricalmirror electron spectrometer and retarding the emitted electrons to a constant energy of 50 eV, an

energy -independent resolution of 0. 8 eV was maintained. Modulation (0.4 eV peak-to-peak) was

applied to the gun accelerating voltage; this eliminated Auger peaks from the spectra. Firstderivative energy loss spectra were obtained by recording dn(E)/dE, where E is the electron kinetic energy in vacuum, vs energy loss Ej^. The location of peaks in n(E) was determined from midpoints

of negative slopes in dn(E)/dE; most peaks could be located to within ±0.2 eV.

The MgO samples were (100) surfaces cleaved from a single crystal and etched in boiling HgPO^

(83%). The surfaces were then sputter-etched in an ultrahigh vacuum system with 260 eV Ar ions

IThis work was sponsored by the Department of the Air Force.

2 Figures in brackets indicate the literature references at the end of this paper.

59
incident at about 70° to the surface normal. Auger spectroscopy showed the surfaces to be atomically clean. Some samples were also annealed for 5 minutes at about 1200°C. The Auger spectra

for these surfaces showed traces of Ca. Both treatments yielded very good LEED patterns and gave

the same energy loss spectra.

Fig. 1. First -derivative energy loss spectra, dn(E)/dE vs energy loss El, for transitions

from the Mg(2p) (solid lines) and MgC2s) (dotted line) core levels in MgO(lOO).

(a) Ep = 120 eV, grazing incidence; (b) Ep = 120 eV (solid line) and Ep = 200 eV

(dotted line), normal incidence; (c) Ep = 2000 eV, normal incidence; (a) observed

Mg2+ free -ion spectrum from ref. 2. The Mg(2s) spectrum has been shifted down by

38.4 eV to align it with the Mg(2p) spectnjm. The positions of bulk peaks

through Bg and surface peaks through S5 are indicated in (c) and (a) , respectively.

60
The experimental ELS spectra for transitions from the Mg(2p) and Mg(2s) core levels to excited

states are shown in figure la-c. The solid curves are Mg(2p) spectra; they can be compared directly

widi the optically observed Mg2+ free-ion spectrum in figure Id [2]. The spectra for transitions from

the Mg(2s) core level, which lies 38.4 eV below the Mg(2p) level, show similar excited-state structure; one such spectrum is shown by the dotted curve in figure lb. The spectrum in figure Ic is for Ep = 2000 eV, normal incidence, and the positions of peaks in n(E) are labeled through Bg, indicating their bulk origin. The locations of these peaks, which are given in Table 1, are in good agreement with those for transitions from the Mg(2p) core level to excited states observed by high-energy

transmission ELS [4], x-ray absorption spectroscopy [5] and ultraviolet reflectivity [6], all of which

are basically bulk probes. The energies of peaks B^, B2, B4 and B5 are very close to the 3s, 3p, 3d

and 4s levels, respectively, in the Mg^"^2p - ni series in figure Id [2]. Additional peaks are seen up

to the ionization limit of the free Mg2+ ion, with no further structure until the onset of the Mg(2s)

spectrum at E-^ » 90 eV. Peak B3, which is only present for E > 500 eV, may be due to the excitation

of both a Mg2+ ion and an O^" ion.

Table 1

Observed Mg(2p) - Excited State Transitions in MgO

This Work (eV) Previous Experiments (eV)

SurBulk^Trans. X-Ray UV

face ELS'^ Abs.° Refl.^

51.3

53.7 54 53.5 53. 5 56.0 55. 8(vw)

58.0 57.8 57.9 57.4 57.3

60. 1 60. 5(w) 60. 5(w)

60.9

65.4 65.4

68. 6(w)

65 65.0

68.2

64

70.2 69. 5 70. 5 73.2

74.5

77.5

74

76. 5

a) Sj^ through S^ in Fig. la.

b) Bj^ through Bg in Fig. Ic.

c) From Fig. 1 of Ref. 4.

d) Ref. 5.

e) From Fig. 1 of Ref. 6. As the primary-electron energy is lowered, surface-state transitions appear in the energy-loss

spectrum, and at Ep = 120 eV, normal incidence (fig. lb), the bulk and surface contributions are of comparable magnitude. The Mg(2s) spectrum in figure lb, taken with Ep = 200 eV, normal incidence,

also contains peaks arising from both bulk and surface transitions. On going to grazing incidence at Ep = 120 eV (fig. la), only surface-state transitions should remain in the spectrum [7]; that is con- firmed here by the disappearance of peak Bi, and we have labeled the peaks S^ through Sg,. The

locations of these surface-state peaks in n(E) are given in Table 1. Both the bulk and surface peaks

listed in Table 1 were observed in our earlier ELS measurements on MgO [1], but we were unable to distinguish between them either experimentally or theoretically.

61
3. Stark Model of Surface States2+

The similarity between the excited-state spectrum of a free Mg ion and the ELS spectrum of bulk transitions from the Mg(2s) and Mg(2p) states in MgO suggests that an ionic description is also

applicable to the surface transitions. As discussed by earlier workers [8], an ion at the surface of a crystal does not experience the same Madelung potential as the same ion in the bulk. The Madelung

shifts of the initial and final states of a surface ion should be essentially the same, producing no net

shift in its transition energies. However, in an ionic insulator there is a large gradient of the Madelung potential at the surface, giving rise to a strong electric field that should produce a Stark

mixing of ionic states (see fig. 2). We have applied a simple model of this effect to the calculation of

the Stark- split spectrum of the Mg^"*" surface ions in MgO.

MgO (100) SURFACE

2+ 2 Mg 0

MADELUNG

FIELD LINES

Fig. 2. Schematic representation of the

(100) surface of an MgO crystal.

£5 IS MADELUNG ELECTRIC FIELD

AT Mg^^ ION DUE TO ALL OTHER IONS

We have formulated a ligand field description of this model of the MgO surface, seen by an Mg^"'' excited state electron is of the form,

The potential

Vg(r) = Vog(r)+V4g(r)[Y4(e,cp)]^^.^ +

in the bulk, and of the form.

Vg(7) = VQg(r) +V^g(r)Y°(e,ep) + V2g(r)Y2°(e,cD) +

+ V4g(r)[Y4(e,cp)]f^^^f^,, (1)

•' 1-^4^® '^^^cubi urfold respectively fourth rank spherical

ns 01 cubic a'm 'fourfold symmeWy,°and the symmetry axis for the spherical

at the surface. Here

harmonic combinations

harmonics is the normal to the surface. The terms Vob(i") and Vos(r) produce shifts of the 2s, 2p, 3s,

3p, .... levels. The terms proportional to V^^{x) and V45(r) produce ligand field splittings of the 3d, 4d, , , , levels, which are presumably less than the widths of the observed transitions (~ 1 eV).

The terms proportional to Vig(r), V2s(r), . . . split and mix the 3s, 3p levels, and the largest term is expected to be V^g(r)Yj^(G,cp).

If we assume that VQp(r) and VQg(r) produce essentially identical small

then the most important difference between the bulk and surface Mg^"*" ion p

the surface term Vis(^)Y|^(9. cp) -*5<^Stark ~ '^^s'^' where pg is the average electric field at the position of a surface Mg2+ ion, and is normal to the surface.

shifts of transition energy,

potential is the presence of

62
In evaluating the Stark effect we consider only the contribution of the one-electron 3s, 3p, 3d, 4s,

4p, 4d and 4f states of the excited Mg^"*" ion. Term structure due to the interaction of the core hole

and the excited electron is neglected. The Stark shifts of the 2s and 2p core holes are also neglected,

since they are calculated to be less than 3% of the 3s Stark shift. The matrix elements of the Stark

interaction for the Mg2+ ion are written, (n'^'m J^stark I '^J^"^^) ^ " (n';e'mj^| r

| nj^m_^), where

E' = e^ /^eff' ^eff effective charge seen fay the electron in its excited orbit, and the (n'^'m^l r |nlm£ ) are the matrix elements of r for the hydrogen atom [9]. The value for the 3s, 3p,

3d, 4s, 4p, 4d and 4f energy levels are taken from the observed bulk peaks (61,62,64,65,65,6^ and

6y, respectively), and each one lies roughly within the range of multiplet energies for the corresponding level of the Mg2+ free ion spectrum [2]. The identification of the peaks beyond 65 is not

critical to the location of the surface peaks S^ through S5.

The results of a Stark effect computer calculation including 3s through 4f levels are summarized

in figure 3, which shows the energy' levels of the Mg2+ excited states as a function of the parameter

e' expressed in atomic units (au). The levels above ~ 70 eV are not accurate because of the neglect of n = 5 and higher states. The best overall fit to the observed surface peak energies, indicated by the circles and triangle in figure 3, is obtained for g' ^ 0.035 au. For the Mg^+ free-ion spectrum, a rough fit of the energy levels for n = 3 -• » relative to the ionization limit is given by a hydrogenic

series with a screened charge of Zgff ~ 3, as expected. For p' = 0. 035 au, this value of Z^r^ gives

the average surface electric field seen by an excited Mg^"^ electron as Pg~ 0. 11 au (3.0 x 11)8 V/cm).

This value of Cg is in satisfactory agreement with the calculated Madelung electric field at the center

of a Mg2+ ion on the undisplaced (100) surface of MgO, = 0. 136 au (3. 7 x iqS V/cm).

Fig. 3. Energy of Stark-split Mg^"*" excited states

vs normalized surface electric field

(in atomic units) . Full lines represent

m£ = 0 levels, dashed lines are mi = ±3

levels, and dotted lines are m£ = ±2 levels. The m£ = ±3 level associated

with 4f is not shown. Squares are observed bulk peaks, dots and triangle are observed surface peaks. The triangular

point is inferred to be an m£ = ±1 level

in view of the selection rule discussed

in the text.

It should be noted that we have included the n = 4 levels in our Stark effect calculation, but have

neglected n = 5 and higher levels. While it is well known that the higher n-states all contribute to lowlying energy levels [ 10], it is also known that these levels rapidly broaden out with increasing n, because excited electrons "leak out" of the quasi-bound Stark levels. We must assume, in our interpretation, that the lowest n = 4 Stark level is barely resolved, and that for n = 5 and higher, the lowlying Stark levels are washed out.

Additional evidence for the Stark-effect model of the surface states is provided by the very small

amplitude of peak S2 in the 2s -• nl siorface spectrum at normal incidence (see dotted curve of fig. lb). 6ecause the interaction of the Mg^"*" surface ion with a normally incident electron beam has cylindrical

symmetry, the resulting transitions should have Am^ = 0 (assuming small angle scattering [ 11]) so

that the (2s, m^^ = 0) -> (nj^,m^ = ± 1) transitions are forbidden. When the electron beam is tipped away

from normal incidence, the symmetry selection rule breaks down. We do indeed observe a strong in- crease in the relative intensity of the (2s,mjj = 0) - {3p,m^ = ± 1) transition at grazing incidence. The

63
(2p,m^ = 0) and (2p,mjj_ = ± 1) states are almost degenerate, and therefore transitions from the 2p

levels to final states with = 0 or ± 1 are observed even at normal incidence,

4. Discussion

Although the phenomenological identification of the ELS spectra of the Mg ion in MgO with the spectrum of the free Mg^"*" ion appears satisfactory, a conceptual difficulty remains. For Zgff ~ 3, a rough hydrogenic value for the size of the Mg^"'", n = 3 orbit is r~ (3n2/2Zgff)agQ{^j. ~ 2. 25 A. Since

the distance between the Mg2+ ion and the ion in MgO is only 2. 1 A, there is a large overlap of the Mg2+, n =3 orbit with the O^" charge cloud, suggesting the importance of banding. Nevertheless, as demonstrated in ref. 1, we have found that the Mg excited states observed in MgO are in poor agreement with the conventional one-electron band picture of MgO.

Preliminary ELS spectra on MgO (110) and (111) and on amorphous MgO films showed small but

reproducible shifts (^ 0. 5 eV) of the surface peaks compared to the (100) surface. The above model

predicts larger Stark shifts on ideal (llO) and (ill) faces, however. Since the surface Madelung

potential is smaller than the bulk value, Mg+ and O" surface ions could be more stable than Mg^"*" and

O^' ions, and we should consider the possibility of such states. Our observations are consistent with

the presence of Mg"*" surface ions in a much reduced surface Stark field (see fig, 4).

Mg ION INFERRED

FROM OPTICAL SPECTRUM

(shifted to agree with S )

Mg ION, EXPERIMENT

(surface spectrum

)

< 2

Q Q- 0. c/i°

ro ^ —

"O Q. Q. ro ro ro

<" w <n '^in ro to po

^3

Mg ION, EXPERIMENT

( bulk spectrum

)

Mg^"" ION

OPTICAL SPECTRUM

( terms averaged

)

6d

7d Si

2 L

B B

5 4

4s

5d 4d 3d

80

B B

2 1

LOSS

60

(eV)

3p 3s

50

ELS SPECTRUM OF MgO (100)

Fig. 4. Comparison of the observed magnesium

bulk and surface transitions observed

by (LES) in MgO with the optical spec- trum of the free Mg2+ ion, and the

spectrum inferred for the free Mg"*" ion from that observed for the free

A£+ ion.

On the other hand, the (110) and (111) faces tend to form facets of the more stable (lOO) surface,

and had poorer LEED patterns than the (100) surface. Alternatively, then, the observations on (110),

(111) and amorphous MgO surfaces may be due to a restrucmring of the surface to produce a (100)-

like local environment about an Mg2+ ion. Additional experimental and theoretical work on the local- ized nature of the excitonic states in ionic crystals is necessary for a detailed understanding of this class of materials.

The authors thank B, J. Palm for assistance with the computer programming and B. Feldman for technical assistance.

64
References

[1] Henrich, V. E, and Dresselhaus, G. , Solid State Commiin. j_6, 1117 (1975),

[2] Moore, C. E. , Atomic Energ^^ Levels , Vol. 1, N.B. S. Circular No. 467, U. S. Government

Printing Office, Washington (1949).

[3] Tracy, J. C. , Electron Emission Spectroscopy (W. Dekeyser et al, , Eds. ) D. Reidel, Dordrecht

(1973), p. 305; Palmberg, P. W. (personal communication) finds escape depths in MgO no more

than 50% greater than in metals at energies larger than several times the band gap.

[4] Venghaus, H. , Opt. Commun. 2, 447 (1971).

[5] Fomichev, V. A. , Zimikina, T, M. , and Zhukova, I. I. , Soviet Phys. Solid State j^, 2421 (1969).

[6] Hanson, W. F. , Arakawa, E. T. , and Williams, M. W. , J. Appl. Phys. 43, 1661 (1972).

[7] Lucas, A. A. and Sunjic, M. , Prog, in Surf. Sci. 2, 75 (1972)

[8] Levine, J. D, and Mark, P., Phys. Rev. 44, 751 (1966).

[9] Condon, E. U. and Shortley, G. H. , The Theory of Atomic Spectra , Cambridge Univ. Press,

Cambridge (1951), p. 133.

[10] Rice, M. A. and Good, R. H. , Jr., J. Opt. Soc, Am. 52, 239 (1962).

[11] Mott, N. F. and Massey, H. S. W. , The Theory of Atomic Collisions , Oxford University Press,

Oxford (1959), Chap. XI.

65

.National Bureau of Standards Special Publication 455, Electrocatalysis on Non-Metallic Surfaces,

Proceedings of Workshop held at NBS, Gaithersburg, Md., December 9-12, 1975. (Issued November 1976)

QUANTUM CHEMISTRY AND CATALYSIS

R. P. Messmer

General Electric Corporate Research and Development

Schenectady, New York 12 301

and

K. H. Johnson

Department and Materials Science and Engineering

Massachusetts Institute of Technology

Cambridge, Massachusetts 02139

A number of recent theoretical calculations on model systems

relevant to chemisorption and catalysis is discussed. The approach

taken is basically chemical in nature, emphasizing the local active site

concept and employing a cluster of atoms to represent the substrate

which interacts with the molecular or atomic adsorbate. The computational procedure used in these studies has been the self-consistent-field

Xa scattered wave (SCF-Xa-SW) method.

Key words: Catalysis; chemisorption theory; metal clusters; molecular

orbitals; quantum chemistry; transition metal complexes.

1. Introduction

While the major focus of this workshop is on the subject of electrocatalysis on non- metallic surfaces, the science of catalysis, in the broadest perspective, is as diverse as chemistry itself. For example, it includes reactions catalyzed heterogeneously by metallic

and nonmetallic surfaces, reactions catalyzed homogeneously by transition-metal complexes

in solution, acid or base catalyzed polymerization of organic molecules, certain reactions

catalyzed in the upper atmosphere, photocatalytic reactions, and biochemical reactions

catalyzed by enzymes.

For the theorist who wishes to investigate catalysis from a fundamental electronic

point of view, the problems are enormous. For example, if one concentrates on heterogeneous catalysis and, in particular, on catalysis by metallic or bimetallic particles

dispersed on a support, should the metal particles be considered from the band-theory

point of view of the solid-state physicist or from the molecular point of view of the

chemist? Since chemisorption is a precursor in heterogeneous reactions on supported

catalysts, what are the most important adsorbate-particle configurations? What are the

effects of the supporting material on chemisorption and catalytic activity? Since the

theorist investigating electronic structure is accustomed to starting with the arrangement

and positions of atoms in a solid or molecule, structural information is crucial. Unfortunately, such inforTTiation is often lacking for heterogeneous catalysts. This is in con- trast to the isolated transition-metal complexes involved in homogeneous catalysis and

enzymes or proteins involved in biocatalysis , where the molecular structure and coordination chemistry are often known to reasonable accuracy.

67
Assuming that one has at least some idea of the molecular structure of the catalyst, a new set of questions can be asked. What is the mobility or lability of the adsorbates on the catalyst surface? What reaction intermediates or decomposition products of the

adsorbates, if any, are important after the initial step of chemisorption? How do these

intermediates convert or interact with other adsorbed species to yield the reaction

products? These questions are related to the reaction kinetics and, from a theoretical

point of view, are much more complex than those regarding simple chemisorption. In view of

the rather poor status of first-principles theories for the kinetics of even the simplest

chemical reactions, it is unlikely that accurate quantitative calculations for the

"potential surfaces" of catalytic reactions will be practical for many years to come. However, with the proper guidance from experiment, calculations on a few configurations and not the entire potential surface may be sufficient to elucidate key features of the reaction

kinetics

.

With these limitations in mind, one may take a more optimistic outlook and consider how

quantum chemistry can presently be employed in understanding aspects of catalysis at the

electronic level. Since it is clear that chemisorption is a precursor in heterogeneous

catalysis, it is important to try to understand the possible structures and bonding of

adsorbates on clusters of atoms which simulate the active centers of supported metallic and

bimetallic catalysts or the active centers of oxide and other nonmetallic catalysts. It is

also important, from the theoretical point of view, to try to understand the coordination

chemical bonding of ligands in isloated transition-metal complexes of the type important in

homogeneous catalysis and the electronic structure of the active centers of enzymes and

proteins involved in biocatalysis . As we will attempt to show in this paper, there are striking analogies between the chemisorption of molecules on the surfaces of transitionmetal aggregates and the chemical bonding of the same molecules as ligands in isolated

transition-metal coordination complexes, as well as equally striking analogies between the

active centers of certain metalloproteins and the active centers of supported metal

catalysts. Such analogies are probably not fortuitous. They should be investigated and

their common basis elucidated.

The self-conaistent-f ield-Xa-scattered-wave (SCF-Xa-SW) approach to molecular-orbital

theory is ideally suited for the quantitative investigation of such systems. In the

following sections, we will describe recent applications of this method to prototype

clusters and complexes of importance in heterogeneous, homogeneous, and biological

catalysis.

2. Discussion of Applications

2..1. Transition and Noble Metal Clusters

The electronic structures of small transition- and noble-metal aggregates (clusters

less than 10 A in size) are of considerable current interest and importance. For example,

the active centers of heterogeneous catalysts often consist of small metallic, bimetallic,

or multimetallic clusters, typically based on Group-VIII and IB elements, supported on a porous refractory material such as silica or alumina [1-3]-!-. The electronic structures of

small metal clusters are also intrinsically interesting in the ways they are related to the

bulk electronic band structures and surface states of the corresponding crystalline metals.

In comparison with the large number of band- structure calculations for crystalline

metals and alloys [4,5], relatively little fundamental work has been directed to the

electronic structures of small metallic clusters. Band theory, in its conventional form,

is based on the_^assumptions of long-range crystalline order, Bloch's theorem, and

reciprocal- or k-space representation, which do not apply to small clusters, where there is at most only short-range order.

Molecular-orbital (MO) theory, on the other hand, is well suited, in principle, for

describing metal-cluster electronic structure but is limited by the accuracy and

practicality of available MO techniques. The recently developed self-consistent-field

Figures in brackets indicate the literature references at the end of this paper.

68
X-alpha scattered-wave (SCF-Xa-SW) approach to MO theory [6] is based on the combined use

of the Xa density-functional approximation to exchange correlation [7] and the multiplescattered-wave method of solving the one-electron Schrodinger equation [8] . This approach

has been successfully applied to a wide range of polyatomic molecules and to clusters

simulating local environments in nonmetallic solids, as summarized in a recent review

article [9].

Small lithium clusters were the first metallic aggregates investigated by the SCF-XaSW method [10] . Results of these studies that are of possible general significance to the

properties of small metal aggregates are: (1) there is a gradual increase of cohesive energy

with increasing cluster size, approaching the value for the bulk solid; (2) there is a gradual decrease of first ionization potential (calculated by the SCF-Xa "transition-state"

method 16,7]) with increasing cluster size, approaching the bulk work function; (3) the

cluster with icosahedral geometry is the most energetically stable of all the clusters

considered.

Because the required computation time does not increase inordinately with the number

of electrons per atom, the SCF-Xa-SW method "is ideally suited for the study of small

transition- and noble-metal clusters. During the past year, we have been carrying out such calculations, with initial emphasis on elemental Cu, Ni, Pd, Pt, and Fe clusters and

Cu-Ni and Cu-Fe bimetallic clusters having cubic, cubo-octahedral , and icosahedral

geometries. Calculations have been completed and analyzed for Cu, Ni, Pd, and Pt clusters

and are in progress for Fe, Cu-Ni, and Cu-Fe clusters.

As an illustrative example, we may consider the results obtained for 13-atom copper

and nickel clusters having the cubo-octahedral geometry shown in figure 1, the structure

characteristic of the local arrangement of atoms in a face-centered cubic (fee) crystal.

First of all, the MO representation of the electronic wavefunctions of the cluster and the

use of orbital synmetries immediately allow one to distinguish between those orbitals

which have no components on the "interior" central atom of the cubo-octahedron and which

therefore are localized on the 12 equivalent "surface" atoms (hereafter called "surface

orbitals"), and those orbitals which have components on all 13 atoms or are predominantly

localized on the central atom (hereafter called "bulk orbitals"). Obviously, the

Fig. 1. Cubo-octahedral cluster containing

13 atoms.

distinction between "surface" and "bulk" orbitals of such a small cluster is academic since

most of the atoms are actually surface atoms. Nevertheless, we shall adopt this terminology

because it allows one to discuss the cluster results in relation to the bulk and surface

electronic states of the corresponding crystals. Indeed, the rotational equivalence of the

surface atoms of the cubo-octahedron in analogous to the translational equivalence of the

surface atoms on an fee crystal. However, the ctibo-octahedral cluster includes both "(100)"

69
and "(111)" type surface planes (see fig. 1), so that it is actually analogous to a faceted

or polycrystalline surface, rather than a single-crystal one.

Using the orbital symmetries associated with the electronic energy levels of Cu^^ ^.nd Ni^2 clusters, v/e can separate the "bulk" and "surface" orbitals, resulting in the energylevel diagrams shown in figure 2. The levels are labeled according to the irreducible

Cu

-0.3

-0.4

<o -0.5 Q

cc

>-

-0.6

-0.7

-0.8

-0.9

-1.0

13 Cu

13 Ni "bulk" "surface"

13 Ni

13 "BULK" SURFACE

Fig. 2. Separation of the SCF-Xa-SW electronic energy levels for Cui3 and Nii3

clusters into "bulk" and "surface" components.

representations of the cubic (0^^) symmetry group. The first thing to note is the high

density of surface orbitals in the vicinity of the Fermi level in the case of Ni^^- These

are antibonding and nonbonding d-orbitals which are spatially oriented away from the

surface of the cluster, as exemplified by the orbital contour maps shown in figures 3 through 5. In figures 3 and 4, respectively, the occupied orbital of Ni^ just below

the Fermi level is mapped in a plane passing through the four nuclei (numbered 1, 2, 3, and

4 in fig. 1) in a square face of the cubo-octahedron and in an equatorial plane containing

four surface atoms and the central atom (numbered 5 through 8 and 13, respectively, in

fig. 1). The equatorial map in figure 4 clearly shows the absence of any contribution of

the central atom to the wavefunction , consistent with the surface localization of the a orbital. ^

The unoccupied Ni]_3 orbital lying above the Fermi level in figure 2 is mapped in

figure 5 in the plane of a square face of the cubo-octahedron. The splitting off of this

d-orbital from the manifold of closely spaced d-levels around the Fermi energy may be

viewed, in the simplest approximation, as a "surface ligand-f ield" effect arising from the

repulsion of neighboring surface atomic d-orbitals. Because it is an empty "surface state" which is spatially directed away from the surface, it could possibly overlap and effectively

accept electrons from interacting adsorbate atoms and molecules. On the other hand, in

Cu^ the a^^ orbital is fully occupied and lies below the Fermi energy (see fig. 2). In fact , the entire manifold of surface and bulk d-levels in Cu lies well below the Fermi

energy, which itself coincides with an energy level (t. ) that corresponds to a highly

delocalized s,p-like orbital [2] . This is very similar to the well known band structure of

crystalline copper, where the d-band is completely occupied and lies well below the Fermi

level. Thus, it is not expected that the d-orbitals will play a dominant role in the interaction of adsorbates with a small copper aggregate or crystallite, as they do in the case of

the surface reactivity of transition metals such as nickel, where the density of surface

d-orbitals around the Fermi level is high. How this situation changes as a function of

"alloying" copper with nickel, to form a bimetallic cluster, is currently under investigate

ion.

70
Fig. 3. Contour map for the occupied anti- Fig

bonding "surface" orbital of a Ni]^3 cluster, corresponding to the energy

level -0.413 Ry shown in figure 2, plotted in the plane of the square

face containing atoms 1, 2, 3, and 4 of the cubo-octahedron illustrated in

figure 1.

4. Contour map for the occupied a2g antibonding "surface" orbital of a Nii3

cluster, corresponding to the energy

level -0.413 Ry shown in figure 2,

plotted in the equatorial plane con- taining atoms 5, 6, 7, 8, and 13 of

the cubo-octahedron illustrated in

figure 1.-

Fig. 5. Contour map for the unoccupied a^^

antibonding "surface orbital of a Ni^

cluster, corresponding to the energy

level -0.376 Ry shown in figure 2, plotted in the plane of the square

face containing atoms 1, 2, 3, and 4 of the cubo-octahedron illustrated in

figure 1.

71
The calculations also suggest a slight depletion of electronic charge and therefore a slight positive charge on the surface atoms of the cluster, compensated for by a buildup of excess electronic charge on the central atom of the cubo-octahedron to preserve total

electrical neutrality of the cluster. The effective surface charge (+0.06 e/atom) on Ni^^

is somewhat larger than the value (+0.04 e/atom) for Cu ^ and would be expected to be

especially significant on transition-metal aggregates where there is a large percentage of

surface atoms with relatively low coordination number, such as those at "corners" aad

"edges" of facets. The number of surface sites of low coordination number will tend to decrease rapidly with increasing particle size, for ideally shaped crystallites.

Further details of these results, including the effects of magnetic spin polarizatioh

on the electronic structure of nickel clusters and electronic structures for palladium and

platinum clusters, are described in a recent publication [11].

One frequently hears the statement that small metallic clusters are intrinsically

"different" in their catalytic behavior from larger particles or crystalline surfaces,

especially in regard to so-called "demanding" or structure-sensitive reactions [1] . The

higher fraction of corner and edge atoms on the surface of a small cluster, as compared

with an ideal "flat" crystal surface, and the presence of spatially directed surface

d-orbitals at these corners and edges (e.g. , see figs. 3-5) could possibly contribute to the structure sensitivity of catalytic reactions through the preferential chemisorption of

reacting molecules on these sites. Furthermore, the effective positive charges of these

sites and their associated electric fields could enhance the interaction of the reacting

molecules through electrostatic polarization or attraction of the adsorbate charge

distributions. Other things being equal, the relatively high ionization potentials {'^7 eV)

and electron affinities (^^4 eV) calculated for small transition-metal clusters by the

SCF-Xa-SW "transition-state" procedure [6,7] will also tend to promote electron transfer

from a nucleophilic adsorbate to cluster. The cluster ionization potential and electron

affinity gradually approach the bulk work function ('^>5 eV) with increasing cluster

size [11]

.

2.2 Bimetallic Clusters and Effects of Alloying

Bimetallic and multimetallic clusters have become increasingly important in the

design of active centers of supported heterogeneous catalysts, particularly those that are selective for specific reactions [2] . For example, the hydrogenolysis of ethane and

dehydrogenation of cyclohexane are reactions which are both activated by elemental nickel

catalysts. However, as copper is alloyed with nickel, the rate of hydrogenolysis decreases

precipitously with increasing copper concentration, while the dehydrogenation reaction

remains relatively constant over a wide range of composition [12] . Two factors of possible

relevance to this behavior are that the hydrogenolysis reaction is structure sensitive,

whereas the dehydrogenation reaction is normally structure insensitive, and there may be

surface segregation of copper in the Cu-Ni aggregates.

We have begun to investigate the effects of alloying on the electronic structures of

transition-metal clusters, with initial emphasis on bimetallic clusters of Group-VIII and

IB elements. This can be accomplished by systematically substituting the "impurity" atoms

in the host metal aggregate and then carrying out SCF-Xa-SW calculations for the electronic

structure of the composite system.

For example, suppose we wish to study a Cu-Ni cluster, and the host aggregate is a Cu^2 cluster having the cubo-octahedral geometry shown in figure 6. Then we can construct

the bimetallic clusters by selectively substituting Ni atoms for Cu atoms in the cubooctahedron, allowing for possible relaxations of interatomic distance. The effects of

surface segregation of Cu can be modeled by embedding small "islands" of Ni atoms, e.g.

,

Ni2 diatoms, in a "sea" of surface Cu atoms and then calculating the electronic structure

of the composite cluster. The effects of alloying and surface segregation on chemisorption

In this case, the t wavefunction s- and p-components on the outer 12 atoms are dominant

over the d- components on the central and outer atoms.

72
Fig. 6. 13-atom clusters: (a) Cubo-octahedral geometry; (b) Icosahedral

geometry.

can then be modeled, in turn, by allowing adsorbates to interact with the bimetallic cluster

and comparing this interaction with the results obtained for chemiso^^ption on pure Ni and

Cu^2 clusters. This procedure can, in principle, be carried out over a range of cluster

compositions cind geometries, including the effects of typical supporting environments.

Calculations for 13-atom Cu-Ni clusters are in progress.

It is possible to study the effects of magnetic spin polarization on the electronic

structures of transition-metal clusters and the changes in magnetic behavior with alloying,

using the spin-unrestricted SCF-Xa-SW method. For example, the spin-polarized orbital

energies for a cubo-octahedral Ni^^ cluster are shown in figure 7. The levels are displayed on two energy scales, one encompassing the entire d-band and overlapping s,p-levels,

and the other resolving the upper part of the d-band on a much finer energy mesh. The

"Fermi level" , separates the unoccupied spin orbitals from the occupied ones and passes

through a high density region of minority-spin levels. Note that the spin splitting of the

energy levels corresponding to orbitals which are predominantly s,p-like is smaller than

that of the d-levels. The spin magneton number per atom in Ni^ is 0.46, assximing the net

spin density, arising from the six unpaired spins in the topmost occupied t^ f (-0.425 Ry)

and t (-0.427 Ry) spin-prbitals , to be delocalized uniformly over all 13 a^oms. This

value Say be compared with the magneton number (0.25) per atom calculated for the smaller

cluster, Ni [11], and with the value (0.54) characteristic of ferromagnetic crystalline

nickel. Although we have thus far investigated the spin polarization in nickel clusters of

rather limited size and geometry, the results are consistent with the experimental

observation of Carter and Sinfelt [13] that the paramagnetic magneton number for small

nickel crystallites decreases continuously from the bulk value as the crystallite size is

reduced from 100 A to approximately 10 A. The so-called " superparamagnetism" of such metal

aggregates is often utilized at a measure of the particle size and dispersion of supported

metal catalysts [2].

Although our theoretical studies of bimetallic clusters have been initially focused

on the Cu-Ni and Cu-Fe systems, we plan to investigate other bimetallic systems as well. For example, a recent experimental investigation of Pt-Fe clusters supported on alumina,

using Mossbauer spectroscopy, chemisorption measurements, and kinetic studies of the

synthesis of hydrocarbons from CO and H (Fischer-Tropsch synthesis) indicate that Fe,

which is normally a good Fischer-Tropsch catalyst^ becomes catalytically inert in the CO-H^

synthesis when alloyed in the form of Pt-Fe clusters [14] . The inert behavior of Fe in the

Pt environment is attributed to a decrease in electron density on the Fe atoms, as suggested

by the Mossbauer isomer shift data [14] . We plan to carry out SCF-Xa-SW studies of Pt-Fe

clusters, using our recent results for elemental Pt clusters as the starting point [11]

.

Other bimetallic systems of catalytic importance which are readily within the scope of the

SCF-Xa-SW method include Pd-Ag and Pd-Au clusters.

73
0.2

-0.3

-0.4

t/5 to

cc -0.5

ujmo>-

en

>-


 -0.6

-0.7

-0.8

-0.9

eg

Ni,3

I -0.37r

sp

I -0.381-

-0.39

d -0.40-

'2g

-

S spd

- »d

sp

-041

-0.42 -

-0.43-

-0.44

-0.45-

-0.46

-0.47-

-0.48-

-0.49-

-0.50 -

'2u

e,

2g

Tin

4l!

i2u

:^2g

-Iig

fu :2g

'2u

'2g

1,

Fig. 7 . Spin-polarized SCF-Xa-SW electronic energy levels for a cubo-octahedral

Nii3 cluster with nearest-neighbor internuclear distance equal to that for bulk crystalline nickel. The levels are labeled according

to the irreducible representations of the symmetry group. The

"Fermi level" Ep separates the occupied levels from the unoccupied

ones.

2.3 Effects of the Supporting Environment

When utilized as the active centers of a heterogeneous catalyst, small metal clusters

are usually supported in a porous refractory material such as silica or alumina [3] . If the sensitivity of a catalytic reaction cannot be explained solely in terms of the surface

morphology or intrinsic electronic structure of the clusters themselves, then it is possible

that the supporting material itself may be involved [1] . Most supporting materials are oxides and may effectively withdraw electrons from the small metal clusters in contact with

them, much in the same way as the ligands of a simple transition-metal atom. The electron

depletion could then allow the small metal cluster to interact with adsorbates in a different manner from the elemental clusters themselves. On the other hand, for larger

metal aggregates or particles in a support, the electron withdrawing effect of the surrounding material may be largely screened out, by the intervening metal atoms, from the surface of

the particle. Thus for these larger particles the main effect on catalytic behavior might

be determined by the effective surface area.

For example, clusters of fewer than six platinum atoms contained in the supercages of

a Y-zeolite containing multivalent cations exhibit a catalytic activity per surface Pt atom

which is enhanced by a factor of about five over that of other platinum catalysts for the

hydrogenation of ethylene, a reaction that is normally structure insensitive [15] . They

74
also chemisorb less oxygen than larger platinum crystallites. This fact and the enhanced

catalytic activity have been attributed to electron transfer from the small platinum

clusters to the zeolite support, as if platinum were behaving more like its neighbor,

iridium, to the left in the periodic table [15]

.

We have begun to construct quantitative theoretical models for the effects of supporting materials such as Y-zeolite, SiO^ / and hl^O^ on the electronic structures of small

transition- and noble-metal clusters. For example, in considering a small platinum cluster

supported in Y-zeolite, the cluster is embedded in a local molecular environment simulating

that of the zeolite, and SCF-Xa-SW calculations will be carried out for the composite

system. The amount of effective charge transfer between the metal cluster and zeolitic

framework can then be determined from a comparison of the electronic structure of the

composite system and that of the isolated metal cluster. The effects of multivalent cations

(e.g., Mg^"*", Ca^"*") in the Y-zeolite on the metal clusters will also be investigated.

Recent applications of the SCF-Xa-SW method to SiO and Al^O^ [16] show that only

relatively small clusters are needed to simulate the electronic structures and properties

of these materials, and these results can be utilized in the modeling of Y-zeolite.

A possible effect of electron transfer from a metal cluster to its supporting environment is to increase the effective work function and electron affinity of the cluster, thereby promoting its interaction with nucleophilic adsorbates but diminishing its interaction

with electrophilic adsorbates. Such calculations as described above should provide an accurate check of this naive argument.

Somewhat related to the issue of metal-support interactions is the nature of the

electronic structure of the 4-Fe active sites in ferredoxin and "high-potential" proteins

which are important as electron-transfer agents in the metabolism of bacteria, plants, and

animals [30]. The active site, an Fe^S^ (S-Cys) ^ cluster, consists of an approximately

cubic array of Fe and S atoms on alternate vertices, with four outer sulfur atoms attaching

the cluster to the surrounding protein via cysteinyl groups. The sulfur atoms may be

viewed as providing a "support" for the Fe^ cluster, and therefore possible analogies with

the effects described in the preceding section are obvious. Indeed, the results of

preliminary SCF-Xa-SW calculations for [Fe^S^ (SCH^)^] 2-, a prototype model of oxidized

ferredoxin and reduced "high-potential" proteins, in which methyl groups have been

substituted for the cysteinyl moieties, suggest significant amounts of charge transfer and

orbital delocalization of the Fe4 electrons onto the surrounding sulfur ligands, with each

iron site in the effective Fe'''^-" valence state [17]. The distorted cubane geometry of the

active centers of oxidized ferredoxin, the reduced high-potential proteins, and their

analogs are explained in terms of the Jahn-Teller instability of the cubic (T^) symmetry,

and spin-orbital occupancy in the lower symmetry is consistent with measured

temperature-dependent magnetic properties tl7]

.

There are probably many useful analogies to be made between supported metal cluster

catalysts and the active centers of metallo-proteins and enzymes like ferredoxin. Such

analogies are probably not fortuitous. They should be sought after and the common basis

of understanding elucidated.

2.4 Models for Chemisorption and Reactivity on Metallic

and Non-Metallic Substrates

Let us now consider some theoretical models for chemisorption on Ni, Pd, Pt, Fe, and

Cu-Ni clusters, with emphasis on those molecular species and configurations which are believed to be important as reaction intermediates in the activation of the chemical

reactions

CO + CH^ + H^O

2 4 2 2 6 by supported metal and bimetallic heterogeneous catalysts.

75
In regard to the nature of chemisorption on transition-metal clusters, it is important

to note the occurrence in inorganic chemistry of metal cluster complexes [18] such as the

iron carbonyl carbide [19] illustrated in figure 8. This is a case where linear and

bridged carbonyl groups occur, both of which have been speculated to form in the chemisorption of CO on transition-metal surfaces and to be important as surface species in catalytic reactions involving CO. The existence of such cluster complexes suggests how

small metal clusters of the type that may be found as the active centers of heterogeneous

catalysts can form more or less normal chemical bonds on the cluster periphery and

particularly at corners and edges of a faceted cluster surface. The cluster complex shown

in figure 8, which is only one of many examples [18] , is a true molecular analogue of CO

chemisorbed on iron carbide, since there is a central carbon atom coordinated octahedrally

by Fe atoms at almost the precise Fe-C distance characteristic of bulk iron carbide [19]

.

Therefore, it may be no coincidence that iron carbide is an excellent catalyst for the

Fischer-Tropsch synthesis of hydrocarbons from CO [20]

.

Our investigation of CO chemisorption on transition-metal and bimetallic clusters draws

on the knowledge and experience gained in recent extensive SCF-Xa-SW studies of the molecular

transition-metal carbonyls, Ni(CO)^, Fe(CO)^, Cr(CO)g, W{CO) , Mo(CO)^, Mn (CO)^^, and

Fe^{CO)^, carried out in collaboration with W. G. Klemperer 121]. The systematic study of

these simpler mononuclear and binuclear metal carbonyls is an absolutely essential prerequisite before calculations on polynuclear carbonyls and cluster models for CO chemisorption are attempted.

It is also important to observe that the study of metal clusters "saturated" by

adsorbates, as shown in figure 8, is likely to be relevant to the high surface coverage of

adsorbates characteristic of catalytic reactions which normally occur on heterogeneous

catalysts at relatively high pressures, whereas theoretical studies which focus on a single

adsorbate bonded to an extended substrate are likely to be relevant to low-pressure chemisorption on a single-crystal surface. The SCF-Xa-SW molecular-orbital treatment of a metal

cluster more or less symmetrically covered with adsorbate atoms automatically includes the

effects of adsorbate-adsorbate interactions along with metal-adsorbate bonding.

O(l')

Feg(C0)|6c2-

To illustrate the type of information an SCF-Xa-SW molecular-orbital study of chemisorption provides, when implemented properly, we consider recent studies of prototype

transition-metal-olef in systems of catalytic importance, carried out in= collaboration with

N. Rosch [22-24] . Olefins are involved in many chemical reactions which are catalyzed

heterogeneously and/or homogeneously . For example, in figure 9 we schematically represent

possible catalytic precursors and reaction intermediates in the heterogeneous hydrogenation

76
of ethylene {C2H4) on a supported transition-metal cluster and homogeneous hydrogenation of

ethylene by a "square-planar" transition-metal coordination complex. The major steps in the

heterogeneous reaction shown in Fig. 9(a) are: the molecular chemisorption of ethylene and

dissociative chemisorption of hydrogen; transfer of chemisorbed hydrogen atoms to chemis--

orbed ethylene; desorption of ethane. The major steps in the homogeneous reaction

shovm in figure 9(b) are: the molecular addition of ethylene and dissociative addition of

hydrogen as ligands of the transition metal, with ligand exchange; the transfer of hydrogen

ligands to the ethylene ligand; dissociation of the hydrogenated ethylene ligand as ethane.

It is clear that there is a one-to-one correspondence between the chemisorption of ethylene

and hydrogen on the metal cluster and the addition of ethylene and hydrogen as ligands of

the isolated metal cc^plex. It has therefore been one of our principal objectives to determine the degree to which chemisorption and heterogeneous reactivity on the active sites of

supported metal or bimetallic catalysts are governed by the same principles of coordination

chemistry which underlie the chemical bonding and homogeneous reactivity of isolated

transition-metal complexes. Since transition-metal complexes tend to be more selective or specific with respect to the reactions they catalyze and are generally more resistant to

H2 H2C = CH2

catalyst L2

(b)

Fig. 9. Heterogeneous vs. homogeneous catalysis. (a) Schematic representation

of heterogeneous catalysis of ethlene hydrogenation on the surface of

a supported transition metal aggregate; (b) Schematic representation

of the homogeneous catalysis of ethylene hydrogenation by an isolated

transition-metal complex.

77
catalyst poisoning than are the extended surfaces of much larger metal aggregates [25] , a fundamental understanding of how the local site symmetry and coordination of a transitionmetal atom affect its chemical bonding could ultimately aid in the design of very small

metal or bimetallic cluster complexes which exhibit optimum catalytic activity, selectivity,

and stability.

As a first step toward this long-range objective, we have initiated a study of the

bonding of ethylene to very small aggregates and complexes of Ni, Pd, and Pt. The intent

of this study is to establish how the metal-olefin chemical bonding varies with the

composition and size of the metal aggregate, supporting environment, and presence of

coordinated nonmetallic ligands.

The simplest case to start with and the one that is also most directly comparable with

isolated transition-metal complexes that homogeneously catalyze the hydrogenation of

ethylene [25] is that of a single ethylene molecule bonded to a single atom of Ni, Pd, or

Pt. In figure 10 we show the resulting SCF-Xa-SW orbital energies for these systems,

calculated for a metal-ethylene bond distance equal to that used in earlier theoretical

studies of the platinum-olef in complex, Zeise's anion [Pt (C2H^)Cl2] ~, by the SCF-Xa-SW

method [22] . This bond distance was chosen for comparison of the bonding of C^H^ to Ni,

Pd, and Pt because many ethylene compexes of these Group-VIII metals have bond distances

o which are within 0.1 A of this value. The metal-ethylene orbital energies in figure 10 are labeled according to the C^^ symmetry group, and the highest occupied level in each case is

the one labeled 5a^. Also shown for comparison are the SCF-Xa-SW orbital energies for the

isolated ethylene molecule, labeled according to the symmetry group.

In figure 10, the related orbitals are connected by broken lines. The ethylene

orbitals from 2a up to lb undergo an almost uniform shift, with only a slight upward

shift on going f?om Pt to P§ tio Ni. The TT and TT* levels of ethylene on the other hand

0.0 3b,

Ni-C2H4 Pd-CgH^ Pt-C2H4 C2H4

-0.2

-0.4

.TT

4a, 207

60|

2b\

2b,

IT — -0.6 3a| 102

°3g

0.8

ui -1.0-

1.2 -

1.4

2a,

">2u

30fl

.2b,

1.6"- 2a,

lo,

Fig. 10. Comparison of calculated electronic energy levels for Ni-C2H4, Pd-C2H4,

Pt-C2H4 and C2H4. The same geometry was employed for all calculations.

78
experience a rather profound change on interacting with an atom of Pt, Pd, or Ni. Thus the

3a^ and 2b orbitals are in each case primarily responsible for the bonding, through the

overlap and interaction of metal d orbitals (d 2 and d , respectively, if the metalethylene and carbon-carbon bond directions are along tKI z and y axes, respectively) with

the ethylene u and tt* orbitals. The metal s and p orbitals also participate in the bonding

but to a lesser extent. The Sa^, ^a^, and 2b^ levels correspond principally to nonbonding

d , d 2 2, and d atomic orbitals, whereas the 5a^ level corresponds to a metal d^2

I orXitaJ tKat is slightly antibonding with respect to the ethylene tt orbital. These ^ characteristics of the orbitals are revealed in figures 11 (a) -11(d) and 12 (a) -12(d), where

contour maps of the 3a , 4a^, 5a^, and 2h^ wavefunctions are shown in the z-y plane

containing the metal atom and carbon atoms of ethylene. The overlap of the axially

symmetric metal d^2 orbital with the ethylene IT orbital in the 3a^ wavefunction is evident

in figures 11 (a) and 12 (a) , corresponding to the classic "O-donation" or ethylene tt electrons into the unoccupied metal d 2 atomic orbital as originally postulated for

transition-metal-olef in bonding by Dewar, Chatt, and Duncanson [26]. In figures 11(d) and

12(d) for the 2h^ wavefunction, we see the overlap of the four-lobe d orbital of the metal

with the ethylene tt* orbital, corresponding to the "TT-back donation" metal d electrons

into the unoccupied tt* orbital postulated by Dewar, Chatt and Duncanson [26]

.

The increasing stabilization of the 2b2 orbital and the increasing overlap of the metal

d ^ orbital with the ethylene tt* orbital in the 2b2 wavefunction as one goes from Ni to Pd

tX Pt [see figs. 10, 11(d), and 12(d)] suggest increasing metal-to-ethylene back donation

from Ni to Pt. In this regard, it is interesting to note that Pd, which exhibits a degree

of back-bonding to ethylene intermediate to that of Ni and Pt, is an order of magnitude

C Ni C Ni

Fig. 11. Contour maps of the (a) 3ai; (b) Aa-^; (c) 5ai; and (d) 2b2 orbitals

of Ni-C2H4. Nodes are indicates by dashed curves.

79
c Pt

Fig. 12. Contour maps of the (a) 3a]^; (b) 4a]^; (c) Sa-^; and (d) 2b2 orbitals

of Pt-C2H4. Nodes are indicated by dashed curves. more catalytically active for the hydrogenation of ethylene than are Ni and Pt [27] . It has been suggested that metal-olefin chemical bonding of "moderate" strength, rather than

bonding that is "too weak" or "too strong," is the optimum precursor in transition-metal

catalyzed olefin hydrogenation [28]

.

Of course, the above described study of the molecular orbitals for ethylene bonded to

a single transition-metal atom at fixed internuclear distance, while instructive as to the

principal components of metal-olefin bonding and their systematic variation from Ni to Pd

to Pt, is rather far removed from the interaction of olefins with metal clusters and

complexes of the type illustrated in figure 9 which are important in heterogeneous and

homogeneous olefin catalysis. As a preliminary study of the effects of gradually increasing the size of the metal aggregate [see fig. 9(a)] .and adding electronegative ligands

[see fig. 9(b)] on metal-olefin chemisorption , the following prototype examples have been

investigated: (1) adding another metal atom to the M-C^H^ complex and (2) adding three

planat chlorine ligands to the M-C^H^ complex. In the former case, the ^2~^2^4 complex has

been considered in two alternative configurations: (1) with the C^H^ molecule attached to the metal diatomic "end-on" (the "linear" configuration) and (2) with the carbon-carbon

bond direction parallel to the metal-metal bond (the "bridged" configuration) . Both

configurations are speculated to be important in the chemisorption and reactivity of

olefins on larger transition-metal aggregates and extended surfaces. In the case of

Ni^-C^H^, such calculations have also been made by others, using the SCF-Xa-SW method [29].

A direct comparison of the orbital energies calculated for Ni-C^H and ^i2~''2^4

the "linear" bonded configuration) is shown in figure 13. The broken lines in figure 13 connect orbitals which are related in the two cases and are of particular importance in

80
Ni-C2H4

0.0

-0.1

3b,

-0.2

0.3

-0.4

-3b2 6a,

-0.5

0.6 -

-0.7

Q,-

02-

30,

Fig. 13. Comparison of calculated electronic

energy levels for Ni-CaH^ and

Niz-CaHi^; the latter one from Rosch

and Rhodin [29]

.

understanding the effects of adding the extra metal atom on the bonding of ethylene a orbital, remains relatively unchanged. The 3a^ level, which corresponds to the ethylene

~ orbital that effectively "donates" electrons into the metal d ^ orbital, is slightly

lowered in energy in the case of ^^2~^2^4' ^^'^^'^^^^^'3 that the second Ni atom, although

not directly bonded to ethylene in the "linear" configuration, contributes to the

stabilization of this orbital. The 2b^ level, which corresponds to the metal d orbital

^ vz that effectively "back donates" electrons into the ethylene TT* orbital, is also somewhat

lowered by the presence of the second metal atom. However, the most dramatic change occurs

in the 5a level which, while slightly antibonding in the case of Ni-C^H^, is significantly

stabilized by the addition of the second Ni atom. This orbital is lowered in energy below

the nonbonding d orbitals and contributes to the metal-ethylene bonding in '^i2^2^4'

Similar results, differing however in important details, have also been obtained for

Mi -C^H^ in the "bridged" configuration. Contour maps for the most important molecular

orbitals of Ni^-C^H , in both the "linear" and "bridged" configurations, may be seen in

Ref . 29 and compared directly with the orbital contour maps for Ni-C H shown in figures

11 (a)- (d).

It is, of course, highly desirable to establish, from first principles, the relative

energetic stabilities of the linear, bridged, and other possible geometric configurations

in metal-olefin bonding, especially in regard to the precursors, reaction paths, and

kinetics of olefin catalysis. The same issues also arise in the chemisorption and

catalytic reactions of other m.olecules, such as CO, on transition metals. This requires

the calculation of total energies and their differences as a function of molecular geometry, often to an accuracy of a few kilocalories per mole. We are currently attempting

to improve our present computational procedure for calculating total energies, starting

with the recently proposed "overlapping-sphere" extension of the SCF-Xa-SW method [30-32]

,

which allows one to compensate for errors arising from the "muffin-tin" representation

without sacrificing the computational efficiency of the method. Indeed, overlapping

atomic spheres were used in the metal-ethylene molecular-orbital calculations described

above. Whether this approach is ultimately accurate enough to serve as the basis for

calculating catalytic reaction paths remains to be seen, particularly for larger metallic

and bimetallic clusters based on Ni, Pd, and Pt, where the total energies are very large

numbers and the pertinent energy differences are very small numbers.

81
Therefore, it is very important to attempt to develop alternative concepts of

catalytic reactivity, based purely on "one-electron" quantities that can be computed

accurately and efficiently with the SCF-Xa-SW method for both small and large systems. An example of the latter is Slater's "transition-state" concept [7], which has already

been successfuly applied to a wide range of molecules and clusters [8,9] to determine

electronic excitation energies, including orbital relaxation effects. Electronic

excitations of optical energies are, of course, crucial to photocatalytic and other photochemical behavior [33]. Such excitations redistribute electronic charge in the molecule,

often making it more susceptible to chemical reaction than the molecule in its ground

state. Electronic excitations of lower energy (e.g. , thermal energy) and those effectively

induced through charge transfer between adsorbate and catalyst probably underlie other

types of catalytic activity. For example, the combined effects of donating ethylene IT electrons into the d orbitals of a transition-metal catalyst and back donation of d

electrons into the ethylene TT* orbital can be viewed as an effective partial excitation

of electrons between the it and TT* orbitals, redistributing charge in ethylene so as to make it more susceptible for chemical reaction with certain other atoms and molecules.

Orbital symmetry rules, such as those introduced by Woodward and Hoffmann [34] and

Fukui [35] and further developed by Mango and Schachtschneider [36] and Pearson [37] can also be used to elucidate certain features of chemical reaction paths. The basic idea is

that, for a chemical reaction to occur with reasonable activation energy, there must be a low lying unoccupied molecular orbital (LUMO) for the reacting system having the same symmetry as the highest occupied molecular orbital (HOMO) , and the energy gap between the

HOMO and LUMO must be reasonably small (a few eV or less) . The requirement that the HOMO

and LUMO have the same symmetry is equivalent to the requirement that these two orbitals

have a net positive overlap. Implicit in the application of these rules is the notion of

electron flow between the HOMO and LUMO.

A catalyst' can often be viewed as providing a pathway for electrons to flow between

the reactants, where such flow is normally prevented to occur by orbital symmetry

restrictions. In other words, a primary role of the catalyst is to break the orbital

symmetry restrictions through the coordination chemical bonding of reactions and to permit

the reaction to occur at low activation energy. A good example is olefin hydrogenation

,

where the transfer of electrons directly from the filled TT orbital of the olefin to the

empty a* orbital of H is forbidden by the orbital symmetries. However, hydrogen atoms that have first been dissociated by a transition-metal catalyst can be readily transferred

to an olefin molecule that is d -TT* back-bonded to the metal resulting in the formation

and evolution of the hydrogenat^§ species (ethane). As shown in figure 9, this process

can take place heterogeneously on the surface of a supported metallic or bimetallic

aggregate or homogeneously through exchange of ligands by an isolated transition-metal

coordination complex. As emphasized, however, in the recent excellent review of heterogeneous catalysis by Madey et_ al. [38], the application of orbital symmetry rules to catalysis is based on the assumption of concerted reactions with high symmetry transition

states, whereas alternative mechanistic schemes are sometimes more likely [39].

In regard to the chemisorption and reactivity of hydrogen on transition-metal

aggregates, it has recently been suggested that delocalized metal s orbitals, and not

localized orbitals, are responsible for the chemical bonding and catalytic reactivity of

hydrogen [40] . While metal s orbitals may indeed contribute significantly to the chemisorption and lability of atomic hydrogen (SCF-Xa calculations indicate that metal s orbitals are important) , the d orbitals are essential to the activation of H^ dissociation

and, as we have shown above, to the coordination chemical bonding and reactivity of olefins

on transition metals. Preliminary theoretical studies of the interaction of CO with

transition metals also show that metal d orbitals are primarily responsible for the bonding, whereas metal s and p orbitals participate in the bonding to a lesser extent [21]

.

The SCF-Xa-SW method is particularly well suited for the quantitative implementation

of orbital symmetry rules in catalysis, since the ordering of both the occupied and

unoccupied molecular orbitals (HOMO and LUMO) and the corresponding orbital energies as a function of occupation number are automatic consequences of the self-consistent-f ield-xa82
density-functional representation of molecular-orbital theory [7] . In this regard, it

should be emphasized that the orbital energies in the SCF-Xa theory are rigorously equal

to derivatives of the total energy with respect to occupation number, i.e.,

£. = 3<E >/3n. ixa xa 1 and therefore should not be identified with Hartree-Fock orbital energies, which are equal

to the differences

£-,rr, =<E,^(n.=l)> - <E,„(n.=0)> iHF HF 1 HF 1 between total energies calculated when the ith orbital is occupied and when it is empty

(fixing the remaining occupied orbitals) [7] . In the conventional Hartree-Fock approach,

moreover, the unoccupied ("virtual") orbitals are represented less accurately than the

occupied ones, whereas in the Xa method the unoccupied and occupied orbitals are treated

to the same degree of accuracy. The Xa orbital energies, in fact, correspond closely to the "orbital electronegativities" defined by Hinze et^ al. [41] as a generalization of

Mulliken's definition of electronegativity. Therefore, the relative positions of the Xa energy levels for reacting molecules, as exemplified in figure 10 for the case of ethylene

interacting with Ni, Pd, and Pt, are an approximate measure of the relative amounts of

electron transfer between the reactants. Slater's transition-state procedure [7], applied

in conjunction with the SCF-Xa-SW method to the HOMO and LUMO of reacting molecules, leads

to remarkably accurate ionization potentials, electron affinities, and electron excitation

energies, including spin-orbital relaxation and the effects of averaging over many-electron

multiplet states which are usually beyond the scope of conventional molecular-orbital

theories [8,9].

The unique characteristics of the Xa orbitals described above greatly facilitate the

application of orbital symmetry rules and the concept of electron flow between reacting

molecules via the coordination chemical bonding of reactants to the catalyst.

As mentioned earlier, another important concern is to determine the effects on chemisorption and reactivity of adding nonmetallic ligands to small transition-metal aggregates,

in part, to model metal-support interactions in heterogeneous catalysis [see fig. 9(a)]

and, in part, to determine the electron-withdrawing effects of such ligands in homogeneous

catalysis by isolated transition-metal complexes [see fig. 9(b)].

As a simple prototype, we may consider the effects of adding planar electronegative

ligands to a platinum-olefin complex. This is just the situation in the classic platinumethylene-chlorine complex, Zeise's anion [Pt(C H^)Cl2]~, for which SCF-Xa-SW calculations

have already been completed [22] . To illustrate the effects of the chlorine ligands on platinum-ethylene bonding, we may compare the molecular-orbital energies for [Pt (C2H^)Cl2]

directly with those shown in figure 10(c) for Pt-C^H^. This comparison is made in

figure 14. The a orbitals (due to ethylene) of Pt-C H^, i.e., the levels labeled la, up to

la^, remain relatively unperturbed by the addition or the CI ligands. The localized CI 3s levels are found at approximately -1.5 Ry in [Pt (C^H ) CI ] ~ , as indicated in figure 14. A

very striking effect occurs, however, for those levels that are derived primarily from the

IT and TT* orbitals of ethylene. The levels in [Pt(C2H^)Cl ]~ which are derived from the

ethylene tt orbitals are labeled by "t" and those derived from the TT* orbitals are labeled

by "*". The broken line rectangle indicates the region in which the energies of predominantly nonbonding CI 3p orbitals are found. The strongly antibonding orbitals (not

shown) are found at approximately -0,3 Ry.

The highly electronegative chlorine ligands effectively withdraw electrons from the

central Pt atom, giving the latter atom a formal oxidation number of +2, and promoting the

bonding between ethylene and platinum by allowing the 4a and 5a^ orbitals, which

correspond respectively to nonbonding Pt d 2-v2 ^^'^ slightly antibonding Pt d 2 orbitals in

Pt-C H , to become significantly bonding in [Pt (C H^) CI ] ~ . Moreover, the 3a^ and 2b^

orbitals which are responsible for Pt (d o)-C^H,(ttT and Pt (d )-C^H^(Tr*) bonding in Pt-C„H^

z 2 4 yz 2 4 4 4 are more strongly bonding in [Pt (C^H^ ) CI^] ~ due to the influence of the chlorine ligands.

A classic example in heterogeneous catalysis of the stereochemical effects of

coordinating transition-metal atoms with electron-withdrawing ligands is the group of

transition-metal halide crystals, such as TiCl^, which catalyze the polymerization of

83
Pt-C2H4 PtC2H4Cl3 C2H4

O.Or 3b,

IT'

lb,

IT

{cL 3s}

Fig. 14. Comparison of calculated electronic energy levels for Pt-CzH

PtCaHitCla- and C2H1,. The level

for PtCaHijCls- are from Rosch,

Messmer and Johnson [22]

.

la,

olefins, as discussed by Ziegler and Natta [42,43]. The chemisorption and polymerization

of the parent olefin (e.g., ethylene), as originally postulated by Cossee and Arlman [44],

is envisioned to occur in the vicinity of a transition-metal atom, specifically titanium

in the +3 oxidation state, which is incompletely coordinated by chlorine ligands near the

surface of an a-TiCl3 crystal, as represented schematically in Fig. 15 (L=C1, R=alkyl).

R R

I I HjC HjC

vacant

site

Fig, 15, Proposed mechanism for the Ziegler-Natta catalytic polymerization

of olefins, [see Ref, 44],

Note that the geometry of the surface titanium-olef in complex is similar to that of the

isolated transition-metal complex shown in Fig. 9(b) which homogeneously catalyzes olefin

hydrogenation, although the Ziegler-Natta catalysis of olefin polymerization occurs

heterogeneously. We have recently carried out SCF-xa-SW studies (described in Ref, 24) of

titanium-ethylene complexes which simulate the first step shown in Fig. 15, These results

again suggest the crucial role of the electron-withdrawing ligands in determing the metalolefin bonding. Such effects are also probably important in chemisorption and catalytic reactivity on the surfaces of transition-metal compounds such as NiO, Ti02 , and

perovskites [38].

R

I vacant HjC~--pH \site —ti~^^l —Ti CH:

y ~CH: ^\ \ / ' CHj

84
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85
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.

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86
National Bureau of Standards Special Publication 455, Electrocatalysis on Non-Metallic Surfaces

,

Proceedings of Workshop held at NBS, Gaithersburg , Md. , December 9-12, 1975. (Issued November 1976)

Relationship Between Electronic Structure

^ and Catalytic Behavior of Transition Metal Carbides

Dwight R. Jennison, David L. Klein, A. Barry Kunz, Kent M, Hall, and Wendell S. Williams

Department of Physics and Materials Research Laboratory

University of Illinois at Urb ana-Champaign

Urbana, Illinois 61801

We have studied the electronic structure of TiC within the

unrestricted Hartree-Fock approximation. We find that strong covalent

bonding occurs between the carbon 2s and 2p and the metal 3d orbitals

.

An absence of strong metal-metal bonding was noted. We tested various

configurations on the metal atoms and found 4s°3d^ energetically

favored. Thus the formation of the carbon-metal bonds results in a loss of 4s character in the electronic structure. Cluster calculations

confirm that the strength of the bonds is sufficient to prevent the

reestablishment of 4s character on a surface. The recent studies of

Kunz, Guse, and Blint indicate that H bonds strongly to the 4s level,

but not to the 3d. Although Ti metal can dissociate H2.our results

explain why the carbide apparently can not. Similar results are anticipated on WC, suggesting a carbon deficient surface may have

properties intermediary to W and WC.

Key words: Adsorption; catalysis; electronic structure; surface

states; transition metal carbide.

1. Introduction

The transition metal carbides are not only useful for many technological applications,

but they also present a challenge to basic science in that their properties are difficult

to interpret theoretically [l]-^ Here we concentrate on some properties relevant to electrocatalysis, specifically the nature of the surface bonding and chemisorption

.

As Dr. Stonehart has pointed out in the previous paper, WC has been shown [2,3] to selectively chemisorb H2 in H2/CO mixtures. It is comparable to Pt as a fuel cell anode

catalyst as it is not poisoned by CO. Additionally, the specific activity of WC^ for

hydrogen oxidation is enhanced when x<l, i.e., when carbon vacancies exist in the material

and at the surface. The key questions are, what changes occur to the electronic properties

of a transition metal upon the introduction of carbon, and how do these changes depend on carbon concentration, x? We begin with one of the simpler transition metals, Ti. Our

results will then be generalized to other species.

TiCx is stable in the rocksalt fee structure when 0.5<x_<1.0. This great stability,

along with other unusual properties such as its great hardness and high melting temperature

[1], indicates that an unusual bonding relationship exists between the Ti and C atoms.

This research was supported in part by the U.S. Energy Research and Development

Administration under Contract E(ll-l)-1198, and in part by the National Science

Foundation under Grant DMR 72-03026.

Figures in brackets indicate the literature references at the end of this paper.

87
Ti readily dissociates hydrogen as does W. The probable limiting factor in the use of

Ti as a catalyst for H2 oxidation is that it bonds oxygen too strongly. This factor has

also been suggested for W [4]. However, TiC has been observed not to dissociate H2 [5].

Apparently the addition of carbon greatly weakens adsorbate-surf ace binding. This immediately suggests that a carbon-deficient carbide may have intermediary properties, accounting for the above mentioned observation on WCj^. Deficient TiC has indeed been observed to react with hydrogen [5].

We briefly describe our theoretical methods, then present the results for the bulk and

surface electronic structure. We close by interpreting our structures in terms of the recent calculations of Kunz, Guse, and Blint [14] on H-transition metal binding, and show

that their calculations are consistent with the above mentioned carbide catalytic behavior.

2. Theory

Our theoretical philosophy excludes the use of variable parameters which may prejudice

our results. We obtain accurate Hartree-Fock (HF) solutions for systems of interest. For those properties not well described by HF theory alone, such as the bulk optical and Fermi

surface properties, we then add electron correlation effects.

The HF solution for the bulk is obtained by utilizing a computationally advantageous

Hilbert space transformation given by the Adams-Gilbert equation [6]. This permits us to self-consistently solve for a set of localized orbitals from which the proper HF density

matrix, p(r,r'), may be found. Having found p, we may then construct Hamiltonian

matrices (in a Bloch representation) which differ only by a phase factor at two points in

reciprocal space. Thus a lengthy self-consistent procedure is avoided at each point, and we may economically use a dense mesh in the Brillouin zone. The local orbitals part of the

calculation is done by rigorously calculating single site parts (Including full non-local

exchange) while expanding the intersite parts in a rapidly converging series, using the

overl|g matrix, S^^^^j H<ai |bj > as a smallness parameter. We use the notation that |ai> is

the i^— local orbital on atomic site a_. Kunz has shown [7] that for a wide variety of

materials sufficient accuracy is obtained by taking the series to first order in overlap.

These materials are ionic and inert gas solids and some metals, such as calcium. Kunz,

Hall, and Jennison [8] have recently shown that excellent results may be obtained on a high -overlap solid (diamond) when the ground state is constructed from a linear combination

of Slater determinants, which maintains nearest-neighbor valence local orbital orthogonality

by virtue of spin. Thus we may still take the series to first order in overlap. This

method also permits "left-right" electron correlation and results in an unrestricted HF

(UHF) solution for the solid. Details of the TiC bulk solution will appear soon [9].

To calculate surface properties, we simulate the surface by a cluster of atoms with

dangling bonds tied off by hydrogens. A space-spin unrestricted HF solution is obtained

in terms of Gaussian basis functions.

Poor HF results may be obtained if the theorist unnaturally restricts the variational

freedom in the spin and/or spatial sense, e.g., if he requires the one-electron wavefunctions to have unnatural symmetry. Such restricted HF (RHF) calculations have occasionally

been unfairly cited as examples of poor HF results. Poor results may also be obtained if one attempts to use only HF theory to calculate properties greatly affected by electron

correlation, e.g., optical properties. One should not expect any one electron theory which

yields decidedly non-local, one-electron wavefunctions to correctly account for an optical

excitation which is decidedly a local event. However, the HF solution forms a proper many

body vacuum state. From this state one may calculate the correlation corrections in the

configuration interaction sense, which then must be added. On all materials so far

attempted [10,11], HF, when corrected for correlation, has consistently yielded better

agreement with experiment than any other method.

88
3. Results

In a UHF calculation, a strong covalent bond in the bulk is indicated by a deep interpenetration of the atomic-like charge densities, rather than by a localized increase in

charge density in the bond (i.e. above those produced by a superposition of atomic-like

densities). This point has been conclusively shown for diamond and silicon [8,12], A meaI sure of this interpenetration is the spatial local orbital overlap matrix which ranges from

-.5 in diamond to <.l in ionic solids.

Upon examining the bulk solution for TiC, we indeed find evidence for strong C-Ti

bonds of the pdir and sda types as shown in Fig. 1. Overlaps of ~.25 indicate that there is an optimal match between the size of the Ti3d lobes (expanded somewhat from free atom size)

and the carbon valence orbitals (about the same as the free atom sizes). We would not expect this match to be as ideal with transition metals having smaller and atomic-like

d-orbitals, as are found in Groups VII and VIII. Indeed, by Group VII (Group VI in the

first series), other phases, such as graphite, successfully compete energetically and prevent the formation of a mono-carbide.

Fig. 2. TiC electron

density along

.00 .27 .54 .81 1.08 1.35 1.62 1.89 2.16' (100).

R (ANGSTROM)

Evidence of the covalent bonding is also found in the position and magnitude of the

electron density minimum between nearest neighbors, Fig. 2. The approximate bands. Fig. 3,

found from the Coulomb potential of the local orbitals UHF calculation, are reproduced here

to illustrate the hybridization. It must be mentioned that, unlike simpler solids, there

is no immediate knowledge of the proper local orbitals (LO's) electron configuration (i.e.,

whether to assume Ti to be As^Sd^, 4s°3d^, etc.). However, within the fractional occupation sense given by the equivalence and symmetry restrictions of Nesbet and adapted to LO

theory by Kunz [13], the Hilbert space spanned by the LO's must be isomorphic to that

89
Fig. 3. Approximate RHF bands shown only to illustrate bonding. Exact bonds with calculated optical properties will be

presented soon [9]

.

spanned by the Bloch functions . At least in the complicated case of TiC, we must find the

occupied Bloch functions to determine whether the configuration we chose for the LO's is

self-consistent with the bands. The bands shown are not fully self-consistent, but are approximately so, Thus they should not be used except as a qualitative indication as to the nature of the bonding. Quantitative properties, e.g. optical and Fermi surface properties,

are calculated from the fully self-consistent bands which will be presented shortly.

The 1 owest band in Fig. 3 is predominantly C2s like with some d hybridization occurring

near L. Strong pd hybridization is found in the bands coming from the pure 3d T12 and r25' points and the pure 2p point. For example, the next lowest band begins pure d at but ends pure p at X4' . The reverse is noted along the next higher A]^ band. The highest Ti

band is Ti4s like and is obviously empty.

An explanation of the changes which occur to Ti metal upon the introduction of carbon

becomes apparent. The atomic levels Ti3d and C2p are about the same energy. Ti4s is

higher than these and C2s is lower. Due to the large overlap and the energy differences

between Ti4s and C2s, the 4s level in the compound is raised in energy, while the 2s is

lowered somewhat. This drains the atomic 4s^ level and the two electrons are shared by the

C2p-Ti3d bands. The C2s level remains full. The compound thus has a Ti4s°3d^, C2s22p°-n

bulk configuration with n being about 3-4.

Thus we may characterize the bulk solid as having strong, but unsaturated, sdo and

pdTT covalent bonds between the Ti and C atoms. The spatial orientation of these bonds

(see Fig. 1) suggests the stability of the fee structure.

Two characteristics of the solid indicate the surface will be electronically very

similar to the bulk. These are the great stability of the fee structure and the very high

activation energies for diffusion (-8eV for Ti and ~5eV for C [1]). Thus we expect the

surface geometry and composition to be the same as the bulk. TiC favors cleavage along

the {100} plane since this breaks only one bond per surface atom. We simulate this surface by a TiC5Hi3 cluster shown in Fig. 4. We obtain a full UHF solution, taking the

free atom configurations as an initial guess. The results are seen in Fig. 5 where the

90
£ -5 -

«>

Ulz

Ul

-2.0

UHF CLUSTER SPIN

Fig. 5. UHF eigenvalues from TiC^H3^3 cluster

presented with RHF band energies for

comparison. Eigenvalues are identi-ied

fied by spin and by dominant component. Note strong "left-right"

correlation and empty As-Ap^ levels.

RHF

BANDS

eigenvalues are identified by spin and by the dominant basis functions of the corresponding eigenvectors. The band eigenvalues are also shown for ease of comparison. Strong

"left-right" correlation is seen in the separation of the predominantly C2p and Ti3d eigenf unctions by spin. The same bonds as seen in the bulk are again indicated by the interpenetration and phases of the C2s , C2p and Ti3d cluster orbitals. In spite of our first

guess, the TiAs-Ap^ hybrid orbital at the surface is virtual and lies considerably above

the Fermi surface. This result strongly suggests that this orbital is not available to bond an adsorbate. However, a dangling lobe of almost pure 3d2^ type is occupied by

<le~ and is available to bond adsorbates. Other Ti3d orbitals are tied to the neighboring

C2p orbitals through the pdir bonds, which are strong and thus not easily disrupted. In summary, the independent UHF cluster calculation confirms the results of the local orbitals

bulk studies.

91
4. Discussion

One might expect that the TiC {100} surface with its dangling d-lobes would present

an attractive surface for catalysis. However, recent calculations of Kunz, Guse and Bllnt

[14] have shown that on both extremes of the first row of transition metals (Sc and Ni), H bonds well to the 4s level but not to the 3d lobes. They concluded that any bonding to the 3d level would be much weaker than to the 4s, and show that the H-4s bond is sufficient

to permit H2 dissociation. Their studies indicated this to be the case not only for isolated Ni atoms, but also for Ni clusters and Ni atoms bonded to oxygens and hydrogens

simulating a silicate substrate. Thus we would expect the unavailability of the 4s level

in TiC to greatly reduce the hydrogen-surface bond strength relative to the metal, as is

observed experimentally. One would also expect that the H-Ti bonding would be stronger

given C vacancies on a surface. The 4s level would lower energetically upon the progressive

removal of nearest neighbor carbons, until it would again be available for bonding.

Several calculations immediately are suggested. We will bring hydrogen down

upon our present cluster and determine the bonding curve. Then we will proceed to TiC^Hjj^

clusters with n<5 and m appropriate to tie off the dangling bonds. Thereby, we will investigate the defective carbide surface. Finally we will turn to WC and other carbides. These

results will be reported soon.

It is seen that our preliminary studies indicate a carbide behavior well consistent

with experimental observations to date. As the high activation energies for diffusion

keep surface composition constant, the possibility of tailoring a catalyst by varying the

carbon content in a carbide seems to exist.

5. References

[1] For a review article see Williams, Wendell S., in Progress In Solid State Chemistry

(H. Reiss, Ed.) 6, 57 (1971).

[2] Ross, P. N. Jr., Macdonald, J., and Stonehart, P., J. Electroanal. Chem. 63 , 450 (1975).

[3] Ross, Philip N. Jr. and Stonehart, Paul, J. Catalysis 39, 298 (1975).

[4] Levy, R. B. and Boudart, M. , Science 181 , 547 (1973).

[5] Goretzki, H. , Bittner, H. , and Nowotny, H., Monatsh 95 , 1521 (1964).

[6] Gilbert, T. L. in Molecular Orbitals in Chemistry, Physics and Biology (P-0 Lowdin, Ed.),

Academic, N.Y., 1964, p. 405.

[7] Kunz, A. Barry, Phys . Status Solidi 36, 301 (1969).

[8] Kunz, A. B. , Hall, K. M. , and Jennison, D. R. , to be published.

[9] Jennison, D. R. , Kunz, A. B., Hall, K. M. , and Williams, W. S. , to be published.

[10] For an independent comparison of theoretical results see Poole, R. T. , et al., Phys. Rev. Bll , 5190 (1975).

[11] On calcium see Mickish, Daniel J., Kunz, A. Barry, and Pantelides, Sokrates T.

,

Phys. Rev . BIO , 1369 (1974).

[12] Hall, K. M. , Jennison, D. R. , and Kunz, A. B. , to be published.

[13] Kunz, A. B., Phys. Stat. Sol . 46^, 385 (1971).

[14] Kunz, A. B., Guse, M. P., and Blint, R. J., these proceedings. -;

92
National Bureau of Standards Special Publication 455, Electrocatalysis on Non-Metallic Surfaces

,

Proceedings of Workshop held at NBS, Gaithersburg , Md. , December 9-12, 1975. (Issued November 1976)

The Role of Intramolecular Quantum Modes

in Electrode Reactions via Intermediate States

Yu. I. Kharkats and J. Ulstrup

Institute of Electrochemistry of the

Academy of Sciences of the USSR

Leninskij Prospect 31, Moscow V-71, USSR

Chemistry Department A, Building 20 7 The Technical University of Denmark

2 80 0 Lyngby, Denmark

The importance of intermediate electronic states

,

of energy-nonconserving (virtual) nature in various molecular

processes has been recognized recently. In homogeneous and

heterogenous electron transfer reactions such states participate effectively, when they provide good electronic and/or

Franck Condon overlap integrals with both the initial and

final manifold of vibrational states. For weak interactions

the electron transfer probability, averaged overall vibrational and electronic states, is calculated by second order

quantum mechanical perturbation theory. Account can thus be

taken of reorganization of both solvent (classical) and intramolecular modes (essentially quantum) resulting in physically

transparent expressions for the reaction probability. Also

adiabaticity of the overall reactions can be incorporated in

the theory. Explicit summation can be performed for various

model systems (Fermi distribution functions, harmonic or Morse potentials et . ) , which again leads to theoretical free

energy relationships (BraSnsted or Tafel plots) qualitatively

different from such relationships for direct reactions. The

electrochemical reduction of molecular oxygen on various

phthalocyanine catalysts is analyzed on the basis of the

formalism presented.

Key words: Bridge mechanism; classical modes; electrocatalysis

electron transfer theory; intermediate states;

intramolecular quantum modes; oxygen reduction;

phthalocyanines

.

1. Introduction

The theory of electron transfer reactions (ET) in polar solutions is now developed to such an extent that at least simple ET , in which no chemical bspnds are broken or established, enjoy a fairly complete descrip

tion [1,2] . Thus, within the dielectric formalism for the solvent outside the first coordination sphere of the reactants, the nature of the

activational process and the reorganization of the quantum subsystems are reasonably well understood. Furthermore, the same ideas have been exFigures in brackets indicate the literature references at the end of this

paper

.

93
tended to cover more complicated situations. This bears first of all to

electrode reactions [2], but also to reactions in homogenous solution

which are accompanied by substantial intramolecular restructuration

(ligand substitution [3] and inner sphere ET reactions [4]).

In the most important theoretical formulations the ET is considered

as the radiationless decay of a "supermolecule" consisting of the reactants and the surrounding solvent. The decay probability, W^^, is calculated by time dependent perturbation theory

W^. = 1^ Av. S N'F^It l4'.> l^6(E.-E.) (1) fi n

-"-f f''i' if

T, where Av. indicates averaging over all initial states, and ^ summation

over all final states, E the total energy, and . and the wave functions representing the stationary initial and final states, i.e. the

solutions of the equations

H.4'. = E.I', and H^-F^ = E^'i'^ . (2) 11 11 ff ff

The total Hamiltonian can be written -

H = H. + V. = + V^. (3)

1 1 r r

which then gives the form of the operator t

T = V. + V^(E. - H + iY)"\ Y ^ 0+ (U)

1 f 1

In the first Born approximation eq (1) takes the form

W.. =J^Av. I:|<1'.|V.Ih'.>|^6(E. -E.) (5) fi n 1^' f'l'i' 1 f

If the first order matrix elements in the expansion of t vanish, higher

order perturbation theory is required in the calculation of ^fj_- If

there are other states, d, for which V^^, V^^ i 0, the second order

matrix elements become

H. ^ = E ^fd ^di (6) i^f

d E. - E

,

1 d

This is usually interpreted as if the reaction proceeds from i to f via

the virtual intermediate states d, but really is based on the time evolution

of the system in terms of eigenfunctions of the unperturbed Hamiltonians

.

Virtual intermediate states have been considered in various physical

processes such as intramolecular radiationless transitions [5], transitions in molecular crystals [6], photochemical reactions [7], superexchange [8], and also in homogeneous and heterogenous ET [9,10]. Application of second order perturbation theory to ET was originally prompted

by a wish of formulating a theory for inner sphere ET. The intermediate

state here corresponds to the localization of the electron in an orbital

of a common ligand (M) for the donating (D) and accepting (A) metallic

centres. Figure la shows schematically the pathway of an electron in such

94
Fig. 1. a: Electron transfer pathway

in a collision complex of

bridge structure in homogeneous solution; b: the same

for an electrode reaction.

Fig. 2. a: Electronic levels in a bridge-assisted homogeneous ET;

b: electronic levels for an electrode reaction via discrete

surface states; c: electronic

levels for an electrode reaction

via surface bands

.

a collision complex. Figure 1b shows the analogous electrochemical reaction, in which e.g. the acceptor has been replaced by the electrode. M can here be a ligand on the depolarizer, or in a wider sense some discrete surface state, the nature of which might in principle be understood, e.g. from calculations of the kind presented in the previous workshop . Figure 2a shows schematically the three discrete electronic levels

on D, M, and A. For electrode reactions many levels in the electrode

must be included, and also, if the surface states interact, they may

form a band instead of discrete levels [11]. These two cases are illustrated in figures 2b and 2c.

For both homogeneous and heterogeneous reactions the system in each

electronic state can be divided into three subsystems, a) the electrons,

b) the intramolecular quantum modes, and c) the solvent. Due to the

large differences in characteristic frequencies, the Born-Oppenheimer

approximation can be applied twice, i.e. the wave functions of the total

system in each state can be written

"^i,d,f(>^'^'^) = *i,d,f(^'^'^^*i,d,f^"'^)x(q) (7)

where ^, (J), and x 3.ve the wave functions of the electrons, the intramolecular modes, and the solvent, respectively, and X, C, and q the corre- sponding coordinates. The electronic terms take the form

U. , . = U.° . + f. , . (q) + g. , . (0) (8) i,d,r i,d,r i,d,r ^ °i,d,f

where f and g are potential energy functions of the solvent (including

ion-solvent interaction) and intramolecular modes, respectively. In the

dielectric formalism [1]

95
i,d,f ^ , vk vk ^vk (9)

where the boson coordinate q^j^ is related to the polarization vector of

the medium; u^^^ are the associated frequencies, v and k refer to the

nature of the elementary excitation and the momentum states, respectively, and q°^]_, is the solvent coordinate shift by the interaction with the

reactants. The fundamental expression for the second order transition

probability is [12]

W, . = j:! Av . E I E f d d d 1 1 I 5(E.-E^), (10) ^ ^ f d E.-E^+iY ^ ^ id

Using the Condon approximation, and performing the summations with re- spect to the solvent system we can write eq (10) as

Wfi^l^dl'l^dil' ^int ' exp(-e^/kT) « n,m ^

(^^j

6l<*f l<l>d>|2|<*^!*.>12 exp (-E^'"^/kT)

v/here Lr:^ and L,^ are electronic matrix elements coupling the intermediate state with the final and initial state, respectively, Z^^^ is the

statistical sum of the high-frequency intramolecular modes, n and m the

quantum numbers of these modes in the initial and final states, Ej-^ the

corresponding energies, and E^'' an activation energy depending on the

nature of all the classical modes.

Equations (7)-(11) provide a general nonadiabatic theory for the

second order ET including account of all solvent and intramolecular

modes. It is noted that there are no restrictions in the nature of the

intramolecular modes , which means that effects such as frequency and

coordinate shifts, anharmonicity etc. can be included in the description. In order to see the physical meaning of the results we shall con- sider the four different cases described below.

2. Results

2.1 Reorganization of Classical Modes Only

If only classical modes (e.g. solvent modes) undergo coordinate

and/or frequency shifts during the reaction, then eq (11) contains only

one term in the sum over n and m. Depending on the character and number

of modes various expressions for Wf^ can be obtained, of which we only

present results for a one-dimensional harmonic representation with no frequency shifts. For the two relative positions of the various intersection regions of the terms shown in figures 3b and 3a Vlfj_ is found to

be [13]

TT3^2jL |2Il 12 exp[-(U , . ^ . , ,-U?)/kT]

W.. = ^'^ l/^, \ (12)

s s di rd

and

96
u

"A

AU-W

a:

U,\ /Uf b:

q

"fi

^fd' di ' exp [-(U^^-

' •!\!AUi2(E^^

s

One-dimensional electronic

terms for ET through an intermediate state of a: high

and b: low energy.

(13)

respectively, where are the reorganization energies (e.g. E = \ fiwCq^ - q^)2). In the case of a high-energy intermediate sfate

(figure 3a and eq 13) the system thus reaches the intersection region

between the intermediate and final state with a probability given by

the activational factor. In this region the electronic subsystem is

reorganized with a probability given by the pre-exponent ial factor,

and via the intermediate state which in this case is an energy nonconserving state. If the intermediate state has a low energy, the activation energy is determ.ined by the highest intersection region of the

intermediate state with either the initial or final state. In this case the system moves classically on the energy surface corresponding to the

intermediate state which is now converted to an essentially energy con- serving state. However, as opposed to two consecutive one-electron transfers, this description covers the situation in which the system does not

relax into the equilibrium configuration of the intermediate state.

2.2 Reorganization in Quantum Modes

If reorganization occurs in the high-frequency intramolecular modes

as well, it is convenient to introduce electron-intramolecular modes

terms. They represent three series of surfaces like the ones shown,

corresponding to all the possible vibrational quantum numbers for each

electronic state. The result for . has to be presented as the sum of

eq (11) with the appropriate expressions for E§'TM (eqs (12) and ( 1 3 ) ) and

the Franck Condon factors inserted. This sum expresses the independent

probabilities of reaction between each quantum state. Both the Franck

Condon factors [1'+] and the activational factors [15] have recently been

considered in some detail for various model systems. Thus, the energies

of the intersection regions of the intermediate terms with the initial

and final terms are (fig. 4)

U^. = (Egi + AU§i + 4 - 4)^/'+E?^ + + ,di (14)

and

97
Fig. 4. One-dimensional electronintramolecular modes,

terms and intervals showing the effective number

of participating vibrational quantum states.

^fd-

,fd iv2 ,fd

'fd (15)

to the exponentially rapid decrease of . with increasing n or dominating terms of the sum in eq (11) are those shown in figure

Due to the

the

Ua and 4b

m,

eq (11) are those shown in figures

For example, in figure Ua they are the ones lying between

the ground ^ initial state and the initial state which equalizes the

initial - intermediate intersection region with the intermediate - final

intersection region for the ground states. Furthermore, due to the rapid increase of the Franck Condon factors with increasing quantum numbers, the^highest of the latter within these intervals gives the main contribution to the sum. For example, for the position of the terms of figure 4a this number, n" , can be found from the equation

(E

di di

s (E

fd

+ AUfd) /4E fd

^di (16)

2 . 3 Electrochemical Reactions

If electrochemical reactions are considered, account must be taken

of the quasicont inuous electronic spectrum of the metal or semiconductor

electrode, and possibly of the set of intermediate states, by summing and

averaging over these states as well as over the vibrational states. Each

term of eq (11) therefore has to be modified by a factor which expresses

the probability that a particular level can donate or accept an electron.

In the one-electron approximation this gives for discrete intermediate

states, e.g. for an anodic reaction [10]

f 1 I

[1 n ,m

•n(e )]p(e . a a f 1

( e , en ) de

a a (17)

where en is the overvoltage , represents a particular metal or semiconductor electronic level, p(ec() the level density, nica^) the distribution function, and WQ^ni(ea,en) the transition probability to the level e^. Similarly, for a band of intermediate states [11]

98
^ ,e32,en)de^dEg^deg2 ^^^^

where and a^re the level densities of the electrode and the surface

band, respectively, n-j and n2 the corresponding distribution functions,

and Eg^ and "^^^ accepting and donating levels, respectively, of the

surface band. These integrals can be calculated by various approximations

depending on the concrete forms of the n's and p's, and for various kinds

of intramolecular modes. A general presentation of the results is

^fi

"fi ^N^(e;;)AN2(e=-^)AN3(e>-2)|L^^(e;Se=-^,e=-2)|^ B

l^di^^a'^Bl '"e2^l^^iit ^ S^^(n,m)S^.(n,m)exi [-E^(e;;,e==^ ,e=-2,n,m,en)/kr] n ,m (19)

where the e"'s are the electronic levels giving the main contributions to

the integrations, the AN ' s numbers of electronic levels effectively participating in the reaction, and the S's the Franck Condon factors. From

this equation theoretical Tafel plots can be calculated, which for suitable values of the parameters of the system are qualitatively different

from Tafel plots for direct (i.e. two-level) reactions.

2.4 Adiabatic Reactions

The formalism presented can be extended to stronger interactions

between the reactants , leading to com.pletely or partially adiabatic reactions [16]. Several different situations may arise, and no systematic

presentation of all the variants will be made here. However, it should be

emphasized that several important predictions, such as free energy re- lationships, are basically unaffected by the possible adiabaticity of the

reactions, i.e. the stronger interactions between the reactants. This is

also relevant to the analysis of the experimental data below.

For adiabatic as well as for nonadiabatic reactions the thermally

averaged transition probability is

^'fi

^ Av^{[T^(E)] %/^^(E)} (20)

where T^(E) is an oscillation period in the initial state, and W^^(E) is

the transition probability at given energy E. Considering classical modes

only, the reaction proceeds when E is sufficiently high that the system

can move classically over the barriers from the initial to the final

state. Depending on the slopes of the zero order surfaces in the intersection regions, different pathways become possible also giving different

expressions for W^j^- For example, the relative positions of the three

terms shown in figure 5 give for W^^ [16]

W^.(E) = (i-e"2^^di)(1-e"2^^fd)/(i_ie-2^Ydi) (21)

where y is the Landau-Zener parameter. For Ydi^Yfd'^^'l (21) reduces to

eq (12) corresponding to a totally nonadiabatic reaction. On the other

99
Ud Uf

u

Fig. 5. One-dimensional electronic

terms for ET in a case of

strong interactions between

reactants

.

hand, if one y is much smaller than unity, and the other one much larger,

a partially nonadiabatic reaction results. For example, if-^.j_<<1, and

Y£^>> 1 , the reaction is nondiabatic in the initial-intermediate step and

adiabatic in the intermediate-final step. Wf-i_ then becomes

W.. = 47rY^ • (Z- 2Tih/kT)"'^ exp[-(U. ,/2)/kT] (22) fi di ^ tdtd

where Z is the statistical sum of the classical modes, and ^E^^ the

splitting energy in the intermediate-final state intersection region.

Analogous expressions can be derived for other term positions [16].

3 . Application to Experimental Data

The formalism for second order ET presented has had a certain

success in the interpretation of various experimental data. This is so in

particular for homogeneous inner and outer sphere bridge-assisted ET [17]

but also for proton transfer reactions [15], some of which proceed

through intermediate states by polyfunct ional catalysis. We shall here

show briefly how a semiquantitative analysis of experimental data on the

electrocatalys is of the reduction of molecular oxygen on various phthalocyanines can be performed in an analogous way. The present formalism provides a theoretical justification of part of a recent empirical analysis

[18] but contains necessarily many reservations due to the limited number

of parameters available. Equation (19) is the basis of the interpretation,

which considers a) the logi-en relationship for a particular metal (i.e.

at fixed AUgj_, AUJ^); b) logi-AUgj_ relationships for different metals,

at fixed en; c) Catalytic activity, i.e. eri-AU|^j^ relationships for

different metals at fixed current density. Correlations between electrocatalytic activity and either magnetic moments or gas phase activity are not considered, since this would be reflected in both the electronic

matrix elements and the activational factor in a very inaccurately predictable way.

Certain simplifying assumptions about the system are necessary.

Thus , firstly it will be assumed that the reaction proceeds at discrete

surface sites (eq (17)). This should really correspond to a monolayer of

phthalocyanine , and the data of Savy et al. [19] on carbon-supported

polymeric phthalocyaninis are therefore used, since they seem to be most

closely related to the model. However, as indicated by eq (19), band

structure could also be incorporated. Secondly, it has been suggested [19]

100
that the 0^ reduction may proceed via excited states of the central metal

ion, induced by coupling with the unpaired spins of the 0„ molecule.

Inclusion of this point would require an additional summation in eq (19) over such states, but in view of the present uncertainty about the exact

role of the excited states in the process , this point will be omitted

here. Thirdly, the effective potential distribution in the catalyst layer

should be known. Two limiting cases were previously considered [11,12].

The whole electrode-solution potential drop could either occur in the

catalyst-solution or in the electrode-catalyst interface. For the

phthalocyanine systems the former situation is likely to prevail, as pointed out below. Finally, the ET may proceed in two ways. Either an electron is transferred from the catalyst to 0^, and the "hole" is filled

by an ET from the electrode (pull push mechanism). This is the pathway

implicit in previous interpretations and also supported by the considerations below. Or, the ET may proceed at first from electrode to catalyst

and then from catalyst to O2 (phali pull mechanism).

Since the first two correlations are not basically affected by the

presence of intramolecular quantum modes, we shall at first ignore such

m.odes. The activation energy for the pull pbsh mechanism, and the former

of the two potential distributions is then

E = (E^- + AU° + en)^/4E^^ (23)

a s di s

when the initial-intermediate intersection region is the higher, and

E = AU° + en + (E^^ + AU° )^/4E^^ (24)

a di s fd s

when the intermediate-final intersection region is the higher (figures 3 and 4). Assuming constant E values eq (23) gives an approximately 118 mv Tafel slope (normal region), whereas eq (24) gives a 59 mv slope (barrierless region); the second potential distribution would lead to

independence of logi on en in the low overpotential region, and 118 mv in

the high overpotential region. For the phthalocyanines of Ni, Cu, and Co

100-120 mv laws are observed experimentally [19], whereas 45-60 laws are observed for the Fe complex. Since furthermore the oxidation potential

is substantially lower for the Fe complex [20], the data are compatible

with the first step having the higher activation energy for Ni , Cu, and

Co, and the second step for Fe. Of course this does not exclude other

possible mechanisms for the reaction at the Fe complex, such as a preequilibrium in the first step. For the push pull mechanism the experimental Tafel plots are compatible with the second of the two limiting potential distributions. Strictly speaking this can not be ruled out, but it

would require that Fe(I) is the most readily formed, and that Cu(I) and

Ni(I) are about equally readily formed, which seems less likely.

Equation (23) also predicts a relationship between the current density and the oxidation potentials of the phthalocyanines as expressed by

AU§j_. Useful correlations can here only be made if the same step is rate

determining throughout the series, which is then restricted to Ni, Cu,

and Co. Since furthermore the potentials and catalytic activities of Ni

and Cu are very similar, the logi-AUg-i_ correlation is of limited applicability. A value of 81ogi/8 ( AU§j_) = 0.23±0.06 can be obtained in the over- voltage region of about -0.7 v.

A linear relationship of unit slope between the oxidation potential

and the catalytic activity (i.e. overvoltage at a fixed current density)

is predicted by eqs (23) and (24). The slope found experimentally is, however, less than unity and seems to increase with increasing oxidation

101
potential. This can be understood qualitatively on the basis of eq (19)

and can be explained by the excitation of intramolecular modes of relatively low frequencies in the cathodic overpotential region. The effective

energy of reaction is then AUg . + e^, where e^(>o) is the excitation

energy for the state giving the main contribution to the sum. For increasing potential (decreasing negative AU§i) m decreases, and the slope

thus approaches unity.

4 . Acknowledgement

One of us (JU) would like to thank Statens naturvidenskabelige

Forskningsrad for financial support.

5 . References

[I] Dogonadze, R. R. and Kuznetsov, A. M. , Fizicheskaya Khimiya.

Kinetika, Viniti (1973).

[2] Dogonadze, R.R., in Reactions of Molecules at Electrodes (N. S. Hush, Ed.), John Wiley and Sons (1971).

[3] German, E. D. and Dogonadze, R. R. , J. Res. Inst. Catalysis,

Hokkaido Univ. 20, 34 (1972).

[4] Dogonadze, R. R. Ulstrup, J. and Kharkats , Yu. I., J. C. S. Faraday

Trans. II 6_8, 744 ( 1 972 ). [5] Fischer, S. F. , Chem. Phys . Lett. 17_, 2 5 ( 1 9 72 ). [6] Jortner, J. and Rice, S. A., J. Chem. Phys. 44^ 3354 (1965).

[7] Grimley,T. B. , Faraday Disc. _58 , 7 ( 1 974 ). [8] Anderson, P. W. , Phys. Rev. 79_, 350 (1 9 50 ). [9] Vol ' kenstein , M. V., Dogonadze, R. R. , Madumarov , A. K. and

Kharkats, Yu. I., Dokl. Akad. Nauk SSSR, Ser. Fiz. lOiim. 199 , 124 (1971).

[10] Dogonadze, R. R. , Ulstrup, J., and Kharkats, Yu. I., J. Electroanal. Interfac. Electrochem. SJ.' (1972).

[II] Dogonadze, R. R. , Ulstrup, J., and Kharkats, Yu. I., J. Electroanal. Interfac. Electrochem. 4^, 161 (1973).

[12] Landau, L. D. and Lifshitz, E. M. , Quantum Mechanics ^ Pergamon

(1965).

[13] Levich, V. G., Madumarov, A. K. , and Kharkats, Yu. I., Dokl. Akad.

Nauk SSSR, Ser. Fiz. Khim. 2_03 , 1 35 ( 1 972 ). [14] Ulstrup, J. and Jortner, J., J. Chem. Phys. £3, 4358 (1975).

[15] Dogonadze, R. R. , Ulstrup, J., and Kharkats, Yu. I., J. C. S. Faraday Trans. II, 70, 64 (1974).

[16] Kuznetsov, A. M. and Kharkats, Yu. I., Elektrokhimiya , in press.

[17] Ulstrup, J., Acta Chem. Scand. 27_, 1 067 (1 973 ). [18] Randin, J. -P., Electrochim. Acta 1_9_, 83 ( 1 974 ).

[19] Savy , M. , Andro , P., Bernard, C. , and Magner, G. , Electrochim. Acta

1_8 , 1 91 (1 973 ). [20] Manassen, J. and Bar-Ilan, A., J. Catal. 12, 86 (1970).

102
Comment on the Papers of H. Gerischer, "Electron Transfer and Electrodatlysis Reactions,"

and of Yu. I. Kharkats and J. Ulstrup, "The Role of Intramolecular Quantum Modes

in Electrode Reactions via Intermediate States"

C. B. Duke Xerox Webster Research Center

800 Phillips Road

Webster, NY 14580

The equations for the electrode current utilized in both of these papers are of the

form

for a single-step reaction. In eq. (1), K(E) is a transfer probability as a function of

energy, E; D(E) is a measure of the occupied (cathodic) or unoccupied (anodic) electronic

states in the electrode; and W(E) is a measure of the corresponding states ("redox

orbitals") in the electrolyte. Because the electron-vibration coupling in the electrode

[D(E)] and solution [W(E)] is written in separable form as a simple product D(E)'W(E),

the validity of eq. (1) is highly circumscribed. In particular, it is not correct in some of the limits for which its use was cited in the papers referred to above.

A consideration of the limits in which formulae like eq. (1) are valid for electron

or exciton transfer processes has been given for conventional solid state tunneling by

C. B. Duke, "Tunneling in Solids," (Academic Press, NY, 1969), Chapter VII, and for

electron and exciton transfer between localized states by T. F. Soules and C. B. Duke,

Phys. Rev. 63^, 262 (1971 ). Specifically, the form of eq. (1) is correct if and only if

the electron is coupled to independent vibrational fields in the initial and final state

and the transfer operator is not a functional of either field. This limit corresponds to the Forster/Dexter model of resonant energy transfer (Soules and Duke, eqs. (3.10)-

(3.12)) and to the concept of electron interactions in the electrodes of both normal and

superconducting tunnel junctions (Duke, Sees. 20 and 21).

The coupling of both the initial and final states to the same polarization or vibration field {e.g., to the "medium" modes in the case of electron transfer reactions

in solution) invalidates eq. (1), as discussed in detail by Soules and Duke. In conventional

quantum-field-theory language, the propagator renormal ization diagrams which lead to eq.

(1) are in large measure canceled by vertex corrections in which the common ("medium")

polarization field links the initial and final states (see, e.g., Duke, pp. 244-245).

Thus, eq. (1) is not expected to be correct if the electron in both the initial and final states interacts with the same vibrational or polarization field.

A second process not described by eq. (1) is inelastic tunneling in which case the

electron actually loses energy during the electron- transfer process (Duke, Sec. 22;

A. J. Bennett, C. B. Duke and S. D. Silverstein, Phys. Rev. 176 , 969 (1968)). These

processes are known to be of importance for electron transfer in semiconductors and

insulators (A. M. Andrews, H. W. Korb, N. Holonyak, Jr., C. B. Duke, and G. G. Kleiman,

Phys. Rev. B5^, 2273 (1972)). Thus, one would anticipate that their omission from eq. (1)

is likely to constitute a serious flaw in the theory of electron transfer reactions.

Finally, the coupling of the transferred electron to low-energy particle-hold excitations or polar fluctuations in the solution can cause infrared divergences which modify

eq. (1). These are well known for piezoelectric semiconductors (C. B. Duke and G. D. Mahan,

Phys. Rev. 239, 1965 (1965)) and x-ray absorption and photoemission by metals

(J. J. Hopfield, Comments on Solid State Phys. 2^, 40 (1969)), and hence, would be expected

to be of relevance to the analysis of electron transfer reactions, especially for metal

electrodes.

oo

(1)

103
In summary, there are three phenomena which can lead to substantive conceptual

alterations in the derivation and interpretation of the formula, cited at this workshop,

for electrode currents associated with electron transfer reactions. These phenomena are

well known to be important in analogous areas in solid-state physics, and hence, should be

considered by electrochemists prior to accepting uncritically the analysis leading to eq.

(1) or its generalizations as discussed by Kharkats and Ulstrup.

104
Reply to Comments by C. B. Duke on a Paper by Yu. I. Kharkats and J. Ulstrup,

"The Role of Intramolecular Quantum Modes in Electrode Reactions

via Intermediate States"

J. Ulstrup

Chemistry Department A, Building 207

The Technical University of Denmark

2800 Lyngby, Denmark

The equation for the electrode current is not written as claimed by Dr. Duke, but has

the form

CO oo

i = jK(E)W(E)dE

J

D(e)de

^•o

where Uio is the minimum of the initial state potential energy surface, E refers to vibrational solvent and intramolecular levels, and e to electronic levels in the metal. The

separable form, W(E)D(£), is therefore related to the averagings over the different subsystems, and in fact, inherent in the Born-Oppenheimer approximation. For the analogous

homogeneous reactions, the integrations are carried out over E only. In this case, the

energy distribution is characterized by a single function of the medium energy states,

even though the part concerning the high-frequency intramolecular modes is written in separable form. However, this is fully justified in view of the nature of these modes.

Dr. Duke's comments, therefore, only relate to other kinds of processes, such as the

Forster/Dexter model of resonance energy transfer.

The theory outlined in the paper, as well as many descriptions of intramolecular

radiationless transitions, do not contain tunnelling in the Gamow sense, although other

formulations of electron transfer theory in the literature do. However, by Franck Condon

restrictions, both formalisms are incompatible with electron energy loss during electron

transfer in the saddle point region of the classical modes. Of course, energy transfer

occurs during the total act of electron transfer and is equivalent to the phonon emission

processes in the paper quoted by Dr. Duke in his second comment.

The third comment is, in fact, already answered by this. In summary, we do not believe that the comments have revealed a necessity of substantive reformulation or reinterpretation of the theory of thermal electron transfer reactions as outlined at this

workshop and other places. On the other hand, we think that a more detailed comparison

between such reactions and the processes mentioned by Dr. Duke would be useful.

105

CHARACTERIZATION

National Bureau of Standards Special Publication 455, Electrocatalysis on Non-Metallic Surfaces

,

Proceedings of Workshop held at NBS, Gaithersburg , Md., December 9-12, 1975. (Issued November 1976)

The Study of Electrode Surfaces by Electrochemical Methods

B. E. Conway and H. Angerstein-Kozlowska

Chemistry Department, University of Ottawa

Ottawa KIN 6N5 Canada

Electrochemical methods for the study of surfaces are reviewed.

Unlike high-vacuum techniques, they measure changes in the state of

surfaces by virtue of occurrence of electrode surface processes,

rather than giving information on the state of the surface itself.

However, electrochemical procedures have high sensitivity down to 2% of a monolayer and multiple states of chemisorption can be distinguished with an energy resolution of 100-200 small calories.

Interactions in the ad-layer can also be derived from the form of

electrical transients. Combination of electrochemical methods with

optical techniques, such as relative reflectivity and el 1 ipsometry

,

provides a new dimension for studies in electrochemical surface

science.

Keywords: Atomic layers; chemisorption; cycl ic-vol tammetry

;

electrochemical adsorption; monolayers; surface processes.

1. Introduction

In this review, it must first be pointed out that electrochemical methods, by their

nature, detect changes of state of metallic and non-metallic surfaces due to the passage

of charge. In these respects they are unlike diffraction techniques, or optical and

electron spectroscopy, which give information directly on the state of a surface itself,

apart from any modifications of the surface introduced by the necessary ion or electron

bombardment pre-treatment procedures reauired to produce a clean surface in the high- vacuum

envi ronment.

It must also be recognized that purely electrochemical techniques are not so generally

applicable to the study of non-metallic electrocatalyst surfaces as they are to metals,

unless the bulk conductivity is reasonably high. However, this condition is fulfilled for

some materials, e.g., mixed oxide el ectrocatalysts such as the La, Sr cobaltites, other

materials such as WC and tungsten bronzes, and also a number of almost metallic sulfides

such as PbS, Cu2S, etc. Hitherto, few electrochemical surface studies on such materials

have been published, partly because their surface stability is, in many cases, lower than

that of nob! e metal s

.

For non-metallic electrode materials that are semi-conductors, electrochemical

techniques combined with the study of photo-effects (e.g., as in the work of Gerischer[l ]^

)

have been most profitable. For such systems, however, both bulk and surface properties

are involved in the observed experimental behavior under photon irradiation.

The electrochemical procedures basically depend on counting charge for Faradaic

processes occurring in the formation or removal of surface layers, often monolayers or

less in extent. The relation between charge passed and electrode potential, measured

Figures in brackets indicate the literature references at the end of this paper.

107
with respect to some constant-potential reference electrode in the circuit, gives additional

information about the processes occurring in electrochemical surface studies as a function

of their relative free energy changes since distinction of processes occurring on an electrode-potential scale is equivalent to distinction on a scale of free energies. In electrochemical terms, this is hence analogous to spectroscopy.

Since surface processes of necessity are chemically and electrochemical ly limited in extent by the quantity of material that can be changed, or introduced or removed on a surface, e.g., deposition or removal of a layer of H atoms at Pt, electrochemical methods

for study of surfaces must employ transient procedures in which a time-dependent current i is registered, e.g., on an oscilloscope or recorder, in response to a change of potential

V or in which a time-dependent change of potential is recorded in response to an impressed

current. Because i and V are conjugate variables in electrochemical experiments, these

two types of procedures give the same information in different ways. In current work,

controlled-potential methods involving the use of an electronic potentiostat programmed

by means of a function generator are preferred.

In addition to observation of changes of state of a surface, electrochemical methods

also give information on the change of state of the electric double-layer at the electrode/

solution interface. Over some potential ranges at various metals, e.g., between 0.35 and

0.75 V on Pt or between 0.0 and 1.37 V on Au in aq. H2SO4 solutions, the electrochemical

transient methods give information on the double-layer charging processes alone, since no other Faradaic surface processes (e.g., deposition of H or oxidation of the surface) occur.

The transient electric currents i^ that pass under these conditions in response to a change

of potential dV/dt in time are "non-Faradaic" currents associated with changing the charge

on the double-layer capacitance C; in general, it = C(dV/dt). At semi-conductor electrodes

the distributed space-charge, wi thi

n

the electrode material, due to p or n charge carriers,

also contributes a capacitance which is measured in transient experiments. For quantitative evaluation of surface processes themselves, separation of the non-Faradaic from the

Faradaic component in the overall current or potential response of the system is necessary.

Since non-Faradaic processes are usually fast because they are associated with short-range

movements of ionic and electron charge near the electrode interface, they have relaxation

times of the order of 10"^-10"7 sec and hence can be separated by measurements on short

time scales. Alternatively, various semi-empirical procedures can be employed.

2. Conditions for Deposition of Atomic Layers at Electrodes

With a number of systems, an array of metal, H atoms or OH and 0 species, can be

deposited uniformly over a surface up to approximately a monolayer before deposition of the

species in bulk form occurs. This means that the initial ad-layer deposit is laid down at electrode potentials below the thermodynamic reversible potential for the substance in bulk

form. This phenomenon has been called "underpotential deposition" (UPD) and is to be contrasted with the normal mode of electrodeposition of most substances which occurs by

growth of nuclei deposited at significant overpotential . UPD can be studied most conveniently at noble metal surfaces as substrates where corrosion or dissolution currents for

the substrate itself are negligible. For an adsorbate A and a substrate M, UPD arises

when the bonding of A to M is stronger than that of A to A. This is expressed thermodynamically by the condition that AGj<0 for the process

AG° A-A + 2M 2 A-M, (I)

if coverage by the underpotential-deposited species A on S is to be significant (e.g., >2%,

a conveniently detectable limit) before the reversible potential for deposition of bulk A

in the process

AG°

2A ± 2e A^ (II)

o o is attained. Alternatively, the condition for UPD to be observed is AGj < AGH- In process

108
II, "bulk A" is represented as "A2", e.g., if bulk A were a diatomic molecule such as H2, O2. If A is a metal, bulk A would, of course, be a lattice Ap.

3. Basis of Electrochemical Methods for Study of Surface Processes

For quantitative studies of coverage of species on an electrode, it is necessary that

the species be "electroacti ve" i.e., the species can be adsorbed or desorbed with passage

of charge in a quantitative way according to Faraday's laws on a microscopic scale, as represented in process 11 above. When this is the case, electrochemical techniques provide

a very sensitive method for study of surface processes since e.g., formation or removal of

1 a monolayer of H atoms on a Pt (111) surface (PtHg^jg + H2O ^ Pt + H303q + e) requires a

i charge of ca. 220 y coulombs per real cm2 which can be accurately measured and in fact,

electrochemical methods are sensitive down to 2% of monolayer coverage. Also, passage of

I the charge during the transient can be followed on a potential scale with a practical

' accuracy of 0.1 to 1 mV and with a resolution, when multiple surface processes occur (see

below), of ca. 5-10 mV, i.e., 115-230 small calories in (free) energy for a le surface

J process. (It is to be noted that this is a much finer resolution on the energy scale for

' surface processes than that afforded, e.g., by current electron spectroscopic techniques.)

j A summary of the direct electrical transient techniques for electrochemical study of

surfaces is given below (Table 1). Some other indirect techniques which are required for

'; study of electro-inactive adsorbates will be described separately later in this review.

Table 1. Electrical methods^ for study of surface processes.

Method Transient measured Results Obtained

1. Controlled current pulse Time-dependent potential Charge and coverage for change

of state of surface.

2. Controlled current pulse

with electronically

differentiated output

3. Controlled potential

pulse (potential-step)

4. Controlled charge pulse

(coulostatic method)

5. Linear voltage sweep

V = Vn ± St

S

0dVdt

A.c. current or applied

a.c. potential (a.c.

impedance)

Electrical transient,

e.g., linear sweep method,

combined with reflectance

or ellipsometry technique; chronoel 1 ipsometry

Time-dependent potential

and dV^/dt resolved

Time-dependent current,

i ^

Time-dependent potential

Time-dependent current

i-t in cyclic mode

Periodic i or V response

i^ as f(Vt),AR/R or A and ip as f(Vt) or t

As in (1) but with better

resolution of surface processes.

As in (1) but with possibility

of more direct charge evaluation

as / i^ dt. Kinetic applications convenient.

As in (3). Kinetic applications

possible.

Direct evaluation of charge and

coverage from / itdt.

Excellent resolution of successive surface processes.

Impedance analysis, Ohmic and

capacitative components of a.c. can be analysed with equivalent

circuit. Kinetic information

obtai nabl e. Optical data on surface

refractive index changes

coupled with change of

surface coverage and charge

(Interpretation complex).

Some other advanced methods exist such as Faradaic rectification, harmonic analysis or timedomain reflectometry using the sampling oscilloscope, but they have not been applied to sur- face studies.

109
In addition to the purely electrochemical techniques listed in Table 1 (entries 1-6)

should be mentioned the use of optical techniques coupled with controlled change of the

electric potential. Methods based on measurement of relative reflectivity change d -^/dV,

in polarized light and on changes of the el 1 i psometric parameters ^ (the amplitude

change) and A (the phase shift) in response to change of potential are being profitably

used for study of electrode surfaces. They may also be applied to non-metallic surfaces.

The interpretation of the changes of optical properties with potential are complex [2,3,4]

as they involve the real and imaginary components of the refractive indices of the doublelayer, the surface film and of the underlying metal, and the thickness of the film. Also,

the electronic refractive index of the surface region of the electrode itself within the

Thomas-Fermi screening distance is modulated by the electric potential due to its effect on surface electron charge density. Further, the optical properties of the double-layer on the solution-side of the interface are modulated by potential due to changes of population

of anions and cations, and to orientation and electrostricti ve compression of solvent

molecules [3,5]. Corrections for these double-layer effects can be made semi -empi rical ly

[4] or they may be minimized by use of an appropriate angle of incidence [2].

The coupled optical electrochemical procedures for study of surfaces provide one of the

few methods whereby in situ "spectroscopic" information on electrode surfaces can be

obtained, in these cases, by measuring AR/R (or A and ip) as a function of wave-length over a range of potentials. Applications to semi-conductor electrode interfaces have been extensive [6] and information on inter-band transitions has been obtained.

4. Relation to High-vacuum Methods

In addition to the above methods, it is desirable to be able to characterize electrode

surfaces themselves by means of high-vacuum techniques such as LEED, HEED and scanning

electron microscopy, as well as by photo-electron spectroscopic methods such as Auger and

ESCA. Such approaches present serious experimental problems for el ectrochemists since

removal of an electrode from solution eliminates the potential control required for generation of a definite state of the surface. Only in the case where stable oxide films can exist on an electrode, or when the metal surface itself is to be examined, will characterization by high-vacuum techniques be meaningful. Also, pre-treatment techniques may introduce changes in state of the electrode surface to be examined. Some ESCA studies of oxide

films on electrodes, e.g., Pt and Au, have been carried out severally by Parsons, Dickenson,

Winograd and others, but it is uncertain if the el ectrogenerated films remain unchanged in the vacuum system.

5. Solution Purification

In all the electrochemical techniques referred to above, and in Table 1, it is essential that ultra-pure solutions be used. Special care must be exercised in all work to ensure that (a) depositable heavy-metal impurities are absent [7] and (b) that adsorbable

organic impurities that may block the surface or be themselves el ectroacti ve are not present

in the solution. A technique for pyrodistillation of water has been described by Conway

et al . [8] which gives very pure solutions with high purity grade H2SO4 or HCIO4 and criteria

for evaluation of solution purity have been established. Since coverage of an electrode by

an organic impurity of mw ca_. 100 requires only ca_. 1.7 x lO'^'^moles cm"2, it is readily

seen that meaningful work in electrochemical surface science can only be achieved by

working with ultra-purified solutions [7,8]. Unfortunately, this has not been the case with

much of the earlier published work in the literature, although Schuldiner's work, in many

papers, must be commended for its attention to the important question of solution purity.

Fortunately, in some cases, adsorbed impurities can be removed in situ at electrode surfaces

by appropriate electrochemical pre-treatment, e.g., in an anodic oxidation pulse or sweep[8],

but impurities tend to be readsorbed if the solution is insufficiently clean as soon as the

oxidized electrode-surface is returned to the metallic state in a following reduction cycle.

110
6. The Potential Sweep or Potentiodynami c Method for Study of

Electrode Surface Processes (cycl i c-vol tammetry ; method 5, Table 1)

One of the most useful techniques for study of electrode surface processes is the

linear potential sweep method or cycl i c-vol tammetry as it is called when applied in a repetitive manner with successive cathodic-and anodic-going changes of potential. Originating from a fast polarographic technique described by SevCik [9], the application of the

potential sweep technique to study of electrode surface processes was developed by Will and

Knorr [10] and employed in many papers by Breiter, Knorr and co-workers. Single-sweep

techniques were used by Gilman [11] and the current pulse method by Brummer et al . [12,13]

for characterization of chemisorbed organic residues which were not directly el ectroacti ve ^ee below)

.

The basis of the method is as follows: currents for an electrode surface process of

the kind A" + M ^ MAads ^ e ^re represented by the equation

and

i => z Fk^ (l-e^)C^- exp BVF/RT

i = z Fk_^ exp -(1-6) VF/RT

(1)

(2)

where V is the metal/solution potential difference and 3 (- '2) is a symmetry factor for the

charge transfer process. F^r a |^urface process such as deposition or desorption of A, to or from a monolayer limit, i or i are not continuous and can only be observed_^by s^me

transient technique; thus, T ^ 0 as 9a ^ 1 and T 0 as e/\ 0 . Generally, i or i is a function of time t in the sweep as V is varied linearly with time at a sweep-rate S=dV/dt.

The time-dependent current i or "i" we denote as i^. Then coverage e/\ at any point in the

sweep is

'A

i^dt

(3)

where Q is the charge for monolayer coverage. Introducing S = dV/dt for the sweep experiment enables eq (3) to be rewritten as

Q S (4)

Correspondi ngly

'A ''initial,

A

Q S (5)

Coverages are hence evaluated from the integral of the i-V-^ profile generated by application of the potential sweep to the electrode where

V, = V. , ± St

t initial (6)

Alternatively, since i = dq/dt in general, where dq is charge passed by the current i in time dt, i can be written

dq dq dV dq

dt dV dt dV

(7)

dq/dV is the capacitance associated with the electrode process, C, i.e., i = CS. C is made

111
up of a double-layer charging component and a component arising from potential dependence

of coverage Bi\ of the surface by A. The latter is the so-called adsorption pseudocapacitance for A, C/\ = dqp^/dy. Thus, C = Cd.i. + C^. Since 6^ and q/\ are related by 0^ = q/\/Q for any coverage < 1

,

i = C , , S + Q.-^ . S (8a)

dV

= Cj.l. S + C^S (8b)2

The profile of i with changing potential at the rate S (= dV/dt) hence gives de^/dV which

is the differential coefficient of the electrochemical isotherm for adsorption of A in the

process A" + M -> MAads + The double-layer charging contribution Cjj.S in eq (8) can usually be subtracted from i by calculation, by an empirical procedureof by high-frequency

measurements, as mentioned earlier.

The kinetic equations (1) or (2) can be expressed in terms of S under sweep conditions

by writing e.g., for (1).

^^-^-IFk^ n-/^]e.p6(V.^.^.^^ tSt)F/RT (9)

which can be evaluated numerically to give the i-V profile for the surface process as e/\ goes from 0^1, or vice versa, in the sweep.

7. Relation between i-V Profiles in Cycl ic-vol tammetry

of Surface Processes and the Adsorption Isotherm

When a process such as A + M ^ ^^ads ^ ""^ ^'^ equilibrium, i.e., when i = i in eqs

(1) and (2), a reversible cyclic voltammetry i-V profile is generated in a potential sweep

experiment and corresponds to the differential coefficient of the electrochemical isotherm

8/\ = f(C/\-, V, T) for el ectrochemi sorption of A on the substrate M. For the Langmuir case

iin[6y\/(l-9^)] = KC/\- exp VF/RT where K is an equilibrium constant kj/k.]. More generally,

when interaction effects are involved, the adsorption isotherm takes the form 2.«[e^/(l-e^)]

= K exp-ge/\*C/\"»exp VF/RT where g measures the change of binding energy with coverage.

When g is finite, the effect is to make the range of electrode potentials over which 9/\ goes from near zero to near unity increase in approximate proportion to g (Fig. 1). In fact, the hal f-wi dths of i-V profiles provide a good basis for evaluation of g [14].

The differential coefficient of the isotherm to which the current in cyclicvoltammetry of a surface process is proportional, is proportional to pseudocapaci tance

Ca [eqs (8a, b)].

Thus, for the general isotherm with interaction effects,

dq, FQ 6^(1-9,) ^

Cr, = • (10)

dV RT 1 + g 9^(1-9^)

applies [15] for reversible deposition or ionization of A.

^ Pure surface processes are distinguished by this relation from diffusion-controlled

processes for which currents are proportional to S'^. .

112
ISOTHERM AND ITS DERIVATIVE

ELECTRODE POTENTIAL (V)

Fig. 1. Relation between electrochemical isotherm and

pseudocapaci tance as a function of potential

for adsorption with (g >0) and without (g = o)

interaction effects [14].

The maximum value of C/\ corresponds to the current maximum in the sweep experiment and

arises, for a reversible surface process, at 9^ = h', i.e., (C^)niax = FQ/4RT • 1/(1+ g/4).

Also the potential range for appreciable C/\ is increased approximately proportionally to g for the same reason referred to in the discussion of the isotherm itself. The form of C/\ with potential in relation to isotherms with various g values is illustrated in Fig. 1.

When a surface process is irreversible, eq (10) no longer applies and the Ga or f(V)

relation, and the corresponding C/\. must be calculated [16] from eq (1) ojr (2) and not from

both at equilibrium. The i-V profile is then asymmetric and its peak shifts by 2.3RT/3F

volts for every 10-fold increase insweeprate S. This provides a convenient way of examining the kinetics of electrochemical surface processes [17,18]. Practically, the reversibility of a surface process in relation to the range of S values employed experimentally is characterized by the ratio, k/S of the rate constant to the sweep rate as discussed further

below. Reversible and irreversible i-V profiles are illustrated in Fig. 2.

POTENTIAL

CONTINUOUS FARADAIC REACTION

REVERSIBLE SURFACE PROCESS

POTENTIAL

IRREVERSIBLE SURFACE PROCESS

Fig. 2. Comparison of reversible and irreversible i-V profiles

generated in cyclic voltammetry of a surface process.

113
8. Multiple States of Chemi sorption below a Monolayer

Contrary to the current-potential response expected for a single chemical species

electrodeposi ted on, or ionized from, a surface, the experimental behavior of many systems

e.g., H, Cu, on Pt, Pb, Bi, Te, Cu on Au, OH and 0 on Ft and Au reveals that UPD of atoms,

with the eventual formation of a monolayer, occurs in a succession of four or more submonolayer states in a very reproducible way. For H on Pt [19,20], four states are resolved over an energy range of ca. 8 kcal . mol"^, while for Pb on Au, about 7 states are resolved over an energy range of £a. 11 kcal. moV

An obvious question arises [19,20] whether these states arise simply on account of

intrinsic heterogenity of the surface, e.g., exposure of various crystal planes of a polycrystalline surface. Two types of experimental facts indicate this is not the case:

(a) experiments on well prepared single crystal surfaces of Pt, e.g., in the (111) (110) or

(TOO) orientations, give [19] more rather than less resolution of the 4-peaked i-V profile

in cyclic voltammetry of H adsorption (Fig. 3); (b) changing anion of the electrolyte, e.g.,

from SO42- or CIO4" in acid solution to HCO3", CO3"" or OH' in alkaline solution, reveals a radical redistribution of adsorbed H amongst the 4 states, two being now filled to greater

coverages and two to less in comparison with the situation in acid medium. This cannot be

accounted for in terms of intrinsic heterogeneity since if a given peak in acid solution

conditions corresponded to filling of elements of a given orientation on a polycrystal 1 ine

surface, such elements could hardly become more filled simply by changing the anion of the

electrolyte. These conclusions apply to results obtained in di 1 ute solutions (0.02 N) when

specific adsorption of anion is not significant (e.g., the i-V profiles for H on Pt become

independent of H2SO4 concentration with increasing dilution and then become identical with

the behavior in dilute HCIO4). Strongly adsorbed ions such as CI , Br' etc. do have

redistribution effects, but this is due to blocking of the surface at the relative positive

end (0.25-0.35 V E^) of the potential range for H adsorption.

0.5M H2SO1, O.OIM H2SO4

273.7 K

1 1

1

\ 2974

1 i

325.2 K

348.2 K

366.9 K

3a. Resolution of 4 states of H adsorption at Pt in dilute acid solution

at various temperatures. These diagrams show the i-V profiles generated by cathodic and anodic sweeps

at the same dV/dt [19].

114
Fig. 3b. Resolution of 7 states of Pb atom deposition on Au with the isotherms.

The observation of multiple states of chemisorption of H on Pt (also observed on Rh and Ir) below a monolayer is paralleled by the results of thermal desorption experiments on dispersed Pt when several states of adsorption bonding of hydrogen are distinguished over

I an estimated energy range of c a

.

8 kcal . mol"l as in the electrochemical experiments.

I However, kinetic effects are involved in thermal desorption experiments while the electrochemical results for H on Pt correspond to equilibrium binding free energies since the

deposition and ionization of H at Pt is a very reversible process.

The observation of multiple states of adsorption in deposition (or removal) of atoms

or radical species is general and, for a given substrate metal, the general form of the

peaks is similar for various atom adsorbates.

I No present theories of chemisorption are able to account for the behavior observed, and

I structure in the electron densi ty-of-states function for various metals is on too coarse a scale [21] (electron volts) to be able to account for the relatively fine resolution demonstrated in electrochemical experiments. We may speculate that the behavior observed (in

the absence of major intrinsic heterogeneity effects) arises from "induced heterogeneity"

associated with one or more of the following effects [19]:

(a) use of various available hybrid orbitals on the surface as relative coverage approaches

an epitaxial or a geometric limit of 1; orbital energy splitting with change of 9a.

(b) modification of the electronic properties of the surface in a communal way by all the

adatoms at a given coverage as that coverage increases towards a monolayer,

(c) occupation successively of various different coordination positions on the surface (e.g. on substrate atoms, between pairs of atoms, in trigonal or square holes depending on orientation of the crystal face). This probably would be a result of (a) and/or (b)

above.

115
3 peaks for H oxidation

Fig. 3c. Resolution of 3 peaks in surface oxidation of Pt below monolayer coverage.

(d) long-range electronic interactions between ad-atoms on the surface (e.g., as discussed

by Newns, Grimley, Koutecky, etc.).

(e) reactive-reconstruction of the substrate/adsorbate layer as coverage by ad-atoms in- creases towards a monolayer, with resultant development of successive 2-dimensional

lattices in the ad-layer.

Probably [19], a combination of the above factors rather than any single one, determines the observed multiplicity of states of chemisorption of ad-atoms below a monolayer. The well-defined experimental behavior provides a challenge to surface chemists

and physicists for quantitative interpretation.

9. Information on the Kinetics of Surface Processes at Electrodes

from the Potentiodynamic Sweep Method and other Procedures

In principle, analysis of a transient which represents a kinetic relaxation response

to a stimulus applied to the reaction, (in this case, voltage) can give information on the

kinetics of the reaction. Potential-step, current-pulse and a.c. impedance methods, e.g.,

by Gerischer, Breiter, Bagotskii and others, have been applied for this purpose. Similar

information is conveniently afforded by the potential sweep method by recording i-V profiles

for a surface process over a Sufficiently wide range (2-3 decades) of sweep rates that the

process changes from a reversible to an irreversible one. Two kinds of kinetic information

are obtained [17,18]: (a) the Tafel slope RT/gF for the surface process and (b) the

limiting sweep rate,SQ, corresponding to the exchange current density, ig, below which the

process is thermodynami cal ly reversible, i.e., perturbation by a sweep at a rate S < Sq

does not significantly disturb the reaction from reversible behavior.

The Tafel slope is obtained [18] by plotting peak potential Vp vs^. log S for S > Sq.

116
The So is obtained by extrapolation of anodic and cathodic peak currents as a function

of log S to a common point with appropriate allowance for interaction effects [14,17]. The

required equations representing this situation have been developed by Angerstein-Kozlowska

and Conway [1 7]

. = + RI

ji„ BL. + BI s (11) P.irrev P,rev

^^^^ 3P

where RT/6F (= b/2.3) is the Briggsian Tafel slope, k is a standard rate constant [k =

k^ exp 6V° F/RT = k_^ exp -(1 -B)V°F/RT] for the standard potential V° for the surface process

defined as the peak potential (6/\ = h) when the process behaves reversibly. Equation (11)

follows from analysis of a surface process in terms of eqs such as (1), (2) and (9). Equation (11) leads to the useful relation

V„ . - Vn = -b log + b log S (12)

P, irrev P, rev ^ o

(of., the Tafel equation ri = -b log i + b log i). In eq (12), S has been defined from

eq (11) as ° °

= BF/k^RT (13)

It is seen that characterizes the kinetics of a surface process like Iq does for a regular continuous Faradaic process. It has the advantage (contrast ig) for kinetic evaluations

in electrocatalysis that the real -area of the electrode is not a required parameter. Hence

the general problem in evaluation of electrocatalysis of separating the real-area factor

from the intrinsic catalytic effectiveness of a surface is conveniently avoided in this

approach.

When lateral interactions on the surface are significant (g > o), experimentally

Vn . - Vn = -b log S ' + b log S (14) P,irrev P,rev ^ o

^

where in Sq = In % (experimentally evaluated by extrapolation) +

[correction term due to interaction effects] (15)

Since (for a le process) Vp = v° + RT/F g Op^^gy and Sp^pev ~ Angerstein-Kozlowska

and Conway [17] showed that'

Vp, irrev = f " ^ ^^^'^P, irrev)]

RT

+ gp

(1-B) g Qp^irrev (16)

which gives a final general relation for log S :

log S^ = log S' - log[l . f

(l - e^^.^^J . ^ (e^^.^^^^ - H) - VV2.3 b (17)

where ep^iYi-ey is [16] 0.63. Hence the kinetically significant Sq can be evaluated if g is known from the half-width [14] of the i-V profile.

The electrochemical kinetics of surface processes, including those for resolved submonolayer states, can hence be completely characterized since b and Sq as well as the

coverage and relative energy of the states can be measured.

An illustration of the method is given schematically in Fig. 4 and with respect to experimental results for two of the states of H deposition in Fig. 5.

117
Fig. 4. Schematic relation between peak currents and log S for transition

from reversibility to irreversibility in an electrochemical surface

process studied by cyclic voltammetry [19]. (a) without interaction

effects, 9=0. (b) with interaction effects taken into account,

g f 0.

0.6-

0.5

2 04

Q3

=1

0.2

01

-tQOl

—t-

+ + +-

A3,

4

Anodic hydrogen peak potentials

as a function of anodic sweep rate

at pH 11.1

0.1 I SWEEP RATE (Vs

10 100

-1

Fig. 5. Experimental relation between peak currents and log S for two of the states of chemisorbed H at Pt at 25 °C in 0. IN aqueous.

118
10. Resolution of Information on Co-adsorbed Species

which give Overlapping Peaks

In a number of cases, species are co-adsorbed and may react el ectrochemical ly on a surface over the same range of potentials. Their relative coverage and electrochemical

behavior cannot then be distinguished under equilibrium conditions. However, if the

kinetics of the surface reaction processes are different, they may be distinguished by

their kinetic relaxation characteristics, e.g., in terms of the current response to an impressed voltage sweep. This possibility arises if the rate constants, e.g., for two surface processes a^ and b^, are sufficiently different that for an impressed sweep at a rate

S, either

or

< S < S„ . (18a) o,a o,b

S„ , < S < ^ (18b) o,b o,a

Condition (18a) corresponds to process b^ remaining reversible while a_ is perturbed from

equilibrium by the sweep, while condition (18b) corresponds to process a_ remaining

reversible while b^ is perturbed.

Experimentally, if the rate S] of an initially slow sweep is suddenly increased say

10 or 100 fold successively at various potentials over the overlapping i-V profile for a pair of co-adsorbed electroactive species, the current increase will be immediate for the

fast process (say a) if lOSi or lOOSi < So^a "•'o'" the fast process but will be slower for

the slow process, b^, for which lOS-, or lOOS] > So,b- Hence, fast and slow processes a^ and b* can be distinguished in an overlapping i-V profile by the locus of fast current responses

for the more reversible process in the profile [22]. Similar distinction can be made by

using an a.c. modulated potential sweep [23].

These methods were applied by Angerstein-Kozlowska , Conway and MacDougall [22] to distinguish between extents of coverage by H and by chemisorbed acetonitrile at Pt which

reacts in the adsorbed state over the potential region where H is co-adsorbed. Theoretical

calculations were also carried out [23] to establish the basis of this method.

11. Initial Stages of Metal Oxidation

The sensitivity of electrochemical techniques allows the very initial stages of metal

oxidation, or passivation in corrosion, to be studied at coverages by 0 species down to 2%

of a monolayer[13] . For the noble metals, a rather general behavior is found initially,

0 species (or OH) are deposited as arrays which are determined by the crystal orientation of

the surface, and give distinguishable states of the diemisorbed film below a monolayer (e.g.,

see Fig. 3c for Pt). The first 15%, approximately, of the monolayer on Pt is reversibly

chemisorbed [13] but with progressively increasing extent of oxidation of the surface, the

oxidation becomes more and more irreversible in the sense that a more cathodic potential is required to reduce the oxide film. At Au, this effect already sets in at ca^. 2% coverage

by OH. The effect is due to rearrangement ("place exchange" [14]) of the adsorbed layer

from a chemisorbed film of OH or 0 species "on" the metal substrate surface to a two- dimensional lattice of OH and 0 species amongst metal atoms (or cations) in the metal surface; in fact, a surface compound is formed. The rearranged layer is thermodynamical ly more stable th.jn the chemisorbed ad-layer, so that hysteresis between formation and reduction of the surface oxide film occurs. The rearrangement is driven by (a) repulsive forces

in the initial ad-layer, (b) the field at the interface and (c) by the free-energy change

of the transformation ad-layer ->- surface compound. Experimental behavior for Au is shown

in Fig. 6.

119
Fig. 6. Experimental distinction between reversibly reduced chemisorbed

OH layer on Au and irreversibly reduced rearranged layer.

(IM aq. HCIO,,, 293 K, sweep rate 50 V-s-i.)

The rearrangement process and the consequent hysteresis can be looked upon in terms of

electrochemical kinetics as a post-electrochemical step that drains off the electro-active

chemisorbed species A on M and thus prevents the el ectrochemisorption reaction A- ± e + M

-y MA from being reversible. The situation may be represented by the scheme below:

Reversible Faradaic deposition process for

adsorption of A atoms on M. A + M ^, A - M + e

k -1

Reversible Faradaic deposition process with

reversible following step. A + M ^ A - M + e II

-2 (Place exchange)

M-A

Reversible Faradaic deposition process for

totally irreversible following stephysteresis.

A + M ^ M + e (III)

M-A

(Place exchange)

A combination of II and III can arise in practice at low coverages and in II there is probably some residual hysteresis component (-->•) under certain conditions. Resolution of

the reversible electrochemisorption process I in III from the irreversible place-exchange

will, in practice, depend on the sweep-rate in relation to the rate constants k] , k_] and

k2. The general behavior of chemisorbed atoms at electrodes, formed in a process such as

M + A- ± e * A - M. may be illustrated comparatively for various conditions by the

following scheme, where U's are energies of interaction between adsorbed atoms A and metal

substrate M; x terms in branch (iii) are electronegativities of A and M:

120
Nucleation and growth

of bulk A on M v/ithout

monolayer formation (OPD)

Chemisorbed A film

below a monolayer (,. (UPD)

Alloy (i)

Completion

of monolayer M-A

epitaxial

(ii)

non-epi taxial

iiy^^ Ordered Surface (iii)

Compound A-M

12. Indirect Studies on Chemisorbed Species

12.1. Cathodic and anodic sweep response

A number of substances chemisorb at electrodes but are not themselves directly "electroactive" in the same way as are H and metal atoms. In such cases, two related methods are

used (cf. Gilman [11], Brummer et al.[12,13]) to characterize the adsorbed species. First,

the substance is adsorbed at equilibrium at a potential where no surface processes can normally occur, e.g., at ca_. 0.5V on Pt. A cathodic sweep is applied and the accommodation of the surface for electrodeposited H is measured in relation to that for a clean

surface. This gives (by difference) the number of H-sites which are blocked by the chemisorbed species. Secondly, an anodic sweep is applied and the substance is oxidatively

removed (usually by reaction with anodically generated surface oxide of the metal). The

number of electrons required for oxidation of the adsorbate per H-site occupied can then be

calculated and the chemisorption characterized. This method has been used to study the

adsorption of many organic substances at Pt.

12.2. Anodic H-displacement currents

An effect discovered by Conway, MacDougall and Angerstein-Kozlowska [24] can be used

to evaluate coverage by species which chemisorb competitively with H. If an adsorbable

substance, e.g., thiourea, is added at a controlled potential in the H adsorption region at

Pt, an anodic current transient arises. It becomes larger the less positive is the potential at which it is adsorbed (Fig. 7). Hence, the current transient cannot be due to any

reduction process. It arises by displacement of the initially adsorbed H, anodically;

thus

A + MH , MA . + h"^ + e ads ads

Integration of the anodic current transient and measurement of the potentiodynamic i-V

profile before and after addition of the chemisorbing additive, quantitatively confirms the

interpretation of this effect. Some substances, e.g., nitriles, chemisorb by anodic H displacement and are themselves partially reduced. Study of the anodic H displacement

effect then gives a procedure for quantitatively following the surface electrochemistry of

the ad-species [22,23]. The effect is a general one and is observed with a variety of

Substances that cheinisorb at Pt.

121
Fig. 7. Anodic H desorption effect with (a) thiourea and (b) acetonitrile desorption

at Pt [24].

13. Applications to Non-metallic Surfaces [25,26]

In the case of non-metal surfaces where direct electrochemical changes can be introduced, studies similar to those described above for metals may be made, provided the ohmic

resistance is not excessive. Also, the optical methods may be applied. Examples where

electrochemical techniques have been applied are conducting oxides and sulfides (PbS, Cu2S),

carbon (very important in el ectrocatalyst technology), tungsten carbide and various tungsten

bronzes

.

122
14. Conclusion

It is convenient to conclude with a summary of information derivable from electrochemical methods of study of surfaces; this is given in tabular form below.

Table 2. Electrochemical studies of surfaces: information derivable.

1. Resolution of distinguishable bonded states below a monolayer, down to 2% coverage (monolayer of H on Pt is equivalent to 220 y coulombs

cm-2; 2.20 X 10-9 f cm-2).

2. State energies determinable to 5 mV (= 100 calories mol"^).

3. Reaction rates of distinguishable adsorbed species below a monolayer determinable.

4. Resolution of behavior of co-adsorbed species can be made by

kinetic relaxation studies.

5. Interaction forces in sub-monolayer states of adsorbed species

determinable from current-potential profile half-widths.

6. Electronic and optical data determinable by single reflection

measurements as a function of potential.

7. Surface space requirement determinable from extent of ads.

H-blocking.

8. Coverages and properties of electro-inactive adsorbates

determinable by H-blocking and indirect oxidation experiments.

15. Acknowledgement

Grateful acknowledgement is made to the National Research Council of Canada for

support of work referred to in this review.

16. References

[1] E.g., see H. Gerischer, Z. Phys. Chem. , 26, 223; 325 (1960); F. Beck and H. Gerischer,

Z. Elektrochem. , 63, 500 (1969); Record Chem. Prog., 23, 135 (1962); Z. Phys. Chem.,

N.F. 13, 389 (19571.

[2] J.D. Mclntyre, Adv. Electrochem. Electrochem. Eng., 9, 61 (1973); Surface Sci . , 37,

658 (1973).

[3] S. Gottesfeld and B.E. Conway, J. Chem. Soc. Faraday Trans. I, 69^, 1090 (1973).

[4] B.E. Conway, L. Lalibertg and H. Angerstei n-Kozlowska , J. Electrochem. Soc, 121 , 1596 (1974).

[5] B. Bewick and J. Robinson, J. Electroanal. Chem., 60, 163 (1975).

[6] E.g., see B.O. Seraphin, Chapter 1, in Semiconductors and Semimetals, Vol. IX, Modulation Techniques, Ed. R.K. Willardson and A. Beer, Academic Press, N.Y. (1972);

and D.E. Aspnes and N. Bottka, Chapter 6 in Semiconductors and Semimetals, Vol. IX, Modulation Techniques, Ed. R.K. Willardson and A. Beer, Academic Press, N.Y. (1972).

[7] A.M. Azzam, B.E. Conway, J.O'M. Bockris and H. Rosenberg, Trans. Faraday Soc, 46,

918 (1950).

[8] B.E. Conway, H. Angerstei n-Kozl owska , W.B.A. Sharp and E. Criddle, Anal. Chem., 45,

1 331 (1973).

123
[9] A. Sev?ik, Coll. Czeck. Chem. Comm., U, 349 (1948).

[10] F. Will and C.A. Knorr, Zeit. Elektrochem. , 64, 258 (1960).

[11] S. Gilman, J. Phys. Chem., 67, 78 (1963); Trans. Faraday Soc, 61, 2456; 2561 (1965).

[12] S.B. Brummer and A.C. Makrides, J. Phys. Chem., 68, 1448 (1964).

[13] S.B. Brummer, J.I. Ford and M.J. Turner, J. Phys. Chem., 69, 3424 (1965).

[14] B.E. Conway, E. Gileadi and M. Dzieciuch, Electrochim. Acta., 8, 143 (1963).

[15] B.E. Conway and E. Gileadi, Can. J. Chem., 42, 90 (1964).

[16] S. Srinivasan and E. Gileadi, Electrochim. Acta., 21, 321 (1966).

[17] H. Angerstei n-Kozlowska and B.E. Conway, J. Electroanal. Chem., (Barker 60th Birthday

Volume), in press, 0976).

[18] D. Gilroy and B.E. Conway, Can. J. Chem., 46, 875 (1968).

[19] B.E. Conway, H. Angerstei n-Kozl owska and W.B.A. Sharp, Symposium on Electrocatalysis

,

The Electrochemical Society, San Francisco, 1973, p. 94.

[20] F. Will, J. Electrochem. Soc, 112, 481 (1966).

[21] D.R. Penn, Surface Science, 39, 333 (1973).

[22] H. Angerstein-Kozlowska, B. MacDougall and B.E. Conway, J. Electroanal. Chem., 39,

287 (1972).

[23] B.E. Conway, H. Angerstein-Kozlowska, B. MacDougall, F.C. Ho and J. Klinger, Faraday

Disc, Chem. Soc. London, 56, 199 (1973).

[24] B.E. Conway, B. MacDougall and H. Angerstein-Kozlowska, J. Chem. Soc. Faraday Trans. I, 68, 1566 (1972).

[25] R. Woods, J. Phys. Chem., 75, 354 (1971); Proc. Aust. Inst. Min. Met., 241, 53 (1972).

[26] F.C. Ho and B.E. Conway, in course of publication.

124
National Bureau of Standards Special Publication 455, Electrocatalysis on Non-Metallic Surfaces

,

Proceedings of Workshop held at NBS, Gaithersburg, Md. , December 9-12, 1975. (Issued November 1976)

Electroreflectance Study of the

Defect States in Reduced n-type Titanium Dioxide

Samuel S. M. Lu, Fred H. Pollak and Paul M. Raccah

Belfer Graduate School of Science, Yeshiva University

New York, New York 10033

We have investigated the electroreflectance spectrum of reduced

n-type TiO , particularly below the fundamental absorption edge

(~ 3 eV) , m aqueous electrolyte, metal-semiconductor and metal-oxidesemiconductor configurations. We have observed for the first time in

electroreflectance features in the spectrum for ^ lu < 3 eV which are related to the defect states. It is found that this structure is strongly dependent on the applied bias electric field. One optical

peak appears at about 1.7 eV, an energy region in which structure has

been observed in various other measurements which are related to both

bulk (optical transmission, photoconductivity) and surface (electrochemical, electron loss spectroscopy) states. The intimate relationship between the bulk defect levels introduced by reduction and surface states will be discussed. Our results will be interpreted in terms of a recent model of Morin and Wolfram for _d-electron surface

states in transition metal oxides.

Key words: Defect states; d^-electrons ; electrolyte electroreflectance; surface states; titanium dioxide;

transition-metal oxides.

1. Introduction

Considerable attention has recently been given to the transition metal oxides because

of their possible application as catalysts. For example Ti02[l/2]2 and SrTi03 [2] have been

observed to catalyze the photoelectrolysis of water. This result has aroused interest in

the possible use of photoelectrolysis for the large scale production of gaseous hydrogen,

which can be conveniently stored and transported. It is generally believed that the catalytic properties of these materials are associated with the role of the ^-electrons in the surface states. With regard to the nature of the d^-electron states these materials may be

classified into two broad categories, i.e. intrinsic and extrinsic (or defect). In the for- mer class there exist occupied d^-electron orbitals (or bands) in the intrinsic, stoichiometric material (e.g., Ti202,Nb02) while in the latter the d^-electron states are empty but can

be populated by means of changes in stoichiometry (e.g., SrTiO , BaTiO^, TiO^). These latter

oxides are not good catalysts in their intrinsic state but do Become active when the d^-elec- tron states are populated by reduction or by the action of band-gap light. For example, it

is only in their reduced states that Ti02ri,2] and SrTi02[2] are active catalysts in the photoelectrolysis of water.

The process of reducing the intrinsic material so as to create those d^-electron surface states that appear to be related to the catalytic action also introduces a set of

"defect" levels and/or bands. These "defect" states then have a specific connection with

;Work supported in part by the National Science Foundation.

'Figures in brackets indicate the literature references at the end of this paper.

125
the d_-electron states and may, in some cases, actually consist of d-orbitals (or bands).

For example, the process of reducing TiO introduces oxygen vacancies as well as creating

Xi^ states. In fact the relation between defect states and catalytic activity may be very

general in the case of non-metallic catalysts. It has been pointed out that from the standpoint of widespread application it is extrinsic semiconductors that are of absorbing interest

[3]. It is in this spirit that we have undertaken an investigation of the defect states in

reduced n-type TiO^ by means of the electroreflectance (electric-field modulated reflectivity)

technique. Electroreflectance (ER) is particularly relevant in this situation since it can be performed in an electrolytic conf igurat ionr4 , 5 ] . This represents an in situ observation

of the properties of materials in relation to their catalytic activities, such as, for ex- ample, the photoelec trolys is of water.

We have investigated the ER spectrum of reduced n-type TiO^, particularly below the fundamental absorption edge (~ 3 eV) , in aqueous electrolyte, metal-semiconductor (MS) and

metal-oxide-semiconductor (MOS) configurations. We have observed for the first time in ER features in the spectrum for ^(ju < 3 eV which are related to the defect states. It is found

that this structure is strongly dependent on the applied bias electric field. One optical

peak appears at about 1.7 eV, an energy region in which structure has been observed in other

measurements related to both bulk (optical transmissi on[6 , 7 ] , photo-conductivity[8,9])and

surface (electrochemical[lO], electron loss spectroscopy [ll ]) states. The relationship

between the bulk defect states introduced by reduction and surface states will be discussed.

Our results will be interpreted in terms of a recent model of Morin and Wolframri2] for

d-electron surface states in transition metal oxides.

2. Exper im.ental Details

Modulation techniques, particularly ER, have been employed over the past several years in optical measurements of solids to achieve greater sensitivity and resolution [5]. They

have been extremely successful in enhancing small structure in optical spectra. Uniform

electric fields for ER can be applied to the surface of a solid by a number of different

techniques including the elec tr olyte [4 ] , metal-semiconductor (MS), or metal-oxide-semiconductor (MOS) methods [5]. There are similarities between these approaches since the interface between the solid and an electrolyte solution behaves very much like a metal-semiconductor diode, in which the applied field modulates the space-charge layer at the interface.

In an electrolyte-solid interface this is accomplished by means of the Gouy layer, which is typically about 5 A thick. In all these techniques large electric fields can be obtained at

the interface with relatively small modulating voltages (~ 1 - 10 V)

.

The experimental arrangement for measuring ER spectra has been described extensively in

the literature and hence will not be repeated herer4,5].

Pure, single crystal TiO^ was obtained from the Adolph Meller Company[l3] and was cut^

into parallelepipeds of dimensions 5x5x1 mm. ^The samples were reduced by annealing at 800 C for about 2 hours in a sealed, evacuated (10 Torr) quartz ampoule. The broad faces [(100)

face] were prepared for optical investigation by polishing with various grades of diamond

paste and finally by Siton[l4], the latter for a period of several hours. This procedure

resulted in a high quality optical finish on the surface. The TiO , which was originally

clear in color, became a deep uniform blue after this procedure indicating that the oxygen

vacancy concentration was rather uniform throughout the crystal. Ohmic indium contacts were

soldered onto the unpolished faces of the sample, which were also sandblasted and painted

black to avoid back surface reflection effects. The MS samples were prepared by evaporating

a semi transparent Igyer of about 70 A of Ni onto the polished surface of the Ti02 in a vacuum of about 10 Torr, thus giving rise to a Schottky type barrier. The MOS samples were prepared by first evaporating a dielectric layer of about 1000 A of SiO and then the semi trans parent 70 A Ni layer onto the Ti02 surface, also under high vacuum.

3. Experimental Results

Shown in figures 1 and 2 are the electrolyte ER spectra of reduced n-type TiO^ for

light polarized parallel and perpendicular to the c-axls, respectively, in the range 1.5 - 4.5 eV for different values of the dc bias. The sense of the bias is taken relative to a

126
o

<i 2^

REDUCED TiOg (100) Face E || c-axis KCI + H2O Electrolyte Configuration.

AC = DC = 6.4 V (peak to peak)

( 1 ) 0.02 V

(2) 5.0 V

(3) 7.03 V

(4) ll.lOV

2.5 3.0 3.5 ENERGY (eV)

4.0 4.5

REDUCED Ti02 (100) Face E±cKCI + H2O Electrolyte Configuration

^04

en<

axis

AC = 6.4 V (peak to peak) DC = (1 ) 1.95 V

(2) 4.02 V

(3) 8.03 V

(4) lO.OOV

2.0 2.5 3.0 3,5 ENERGY (eV)

4.0 4.5

Fig. 1. Electroref lectance spectra of reduced

n-type Ti02 in the electrolyte configuration for light polarized parallel

to the c-axis at various values of the

dc bias.

Fig. 2. Electroref lectance spectra of reduced

n-type Ti02 in the electrolyte configuration for light polarized perpendicular to the c-axis at various values

of the dc bias.

platimum electrode, i.e. positive bias indicates the sample is positive with respect to the

electrode. The electrolyte used was approximately I molar KC-E- in water[4]. At the lowest

value of the bias several features are observed above 3 eV and a peak is seen at about 1.7

eV. In general, the structure above 3 eV is similar to that which has been reported pre- vious lyC9 , 15 , 16 ] with only some minor differences. These peaks have been assigned to various interband optical transitions of the intrinsic material. We shall return to a discussion of these features shortly. What is of considerable interest is the optical structure around 1.7 eV, for both polarization directions, which has not been observed before in

ref lection[l7 ] or ERC9,15,16]. As the bias is increased there are two main effects on the

spectrum: (a) the 1.7 eV peak shifts to lower energies and (b) a new structure appears which

is split off from the 3.15 eV interband peak which also moves to lower energies. These two field dependent features appear to shift at the same rate with applied dc voltage. This

new peak is observed at the highest bias at about 2.4 eV in figure I, for which it is clearly resolved, and 2.6 eV in figure 2 (note that the biases are different for the two polarizations). The structure due to the interband transitions above 3 eV have been left relatively unaffected by the increase in dc bias. The dc bias voltages used in this study were kept

low enough to avoid oxygen evolution at the sample. Although we have not been able to com- pletely follow the shift of the field-dependent features we can estimate their change in

energy from the shift of the minima (~ 2.5 eV at the lowest bias) since the line shape does

not change with field. We have determined the shift of the minima for both polarizations

and found that in a certain bias range it is linear in /V, where V is the applied bias.

These results are shown in figure 3. For a fully depleted space change layer it can be

3.0-

> 2.5

CD

tr

UJ 2.0

1.5

° Ellc-Qxis

° E±c-axis

1 2 3 <V (volts)'^^

Fig. 3. Voltage dependence of the minima in the

electrolyte ER spectra for ^u) < 3 eV. This shift is representative of the

voltage dependent structure in the

spectra

.

127
demonstrated that the average surface field E is proportional to yv[4,18]. Therefore, the

results of figure 3 indicate that the shift of these structures is linearly dependent on the

surface field E . s The optical structure for ^ ud > 3 eV corresponds to that observed by previous workers

[9,15,16], the main difference being that their spectra shows a large single structure in

the range 3.5 < ^ uo < 4 eV, while in this region we find a doublet structure for this feature

for both polarization directions. The reasons for this difference are not understood at this

time

.

The ER spectrum for the MS configuration is shown in figures 4 and 5 for light polarized parallel and perpendicular to the c-axis, respectively, while in figures 6 and 7 are plotted to ER data for the MOS configuration for the two light polarizations. For both the MS and MOS cases the sample bias is taken in reference to the Ni electrode.

Fig. 4. MS electroref lectance spectra of re- Fig. 5. MS electroref lectance spectra of re- duced n-type Ti02 for light polarized duced n-type Ti02 for light polarized

parallel to the c-axis. perpendicular to c-axis.

REDUCED TiOg (lOO)Face E||c-axis

MOS Configuration

2 -

AC = ) 5 V (peak lo peak) DC Ml) -3.0 V

(2) 1,0V

(3) 7. IIV

(4) 10.19 V

1.5 2.0 2.5 3.0 3,5 4,0 ENERGY (eV)

Fig. 6. MOS electroref lectance spectra of re- duced n-type Ti02 for light polarized

parallel to the c-axis.

REDUCED TiOj (100) Face E-Lc-axis MOS Configuration

2 -

AC = 1 5 V (peok to peak) -2- DCMDl.OV

(2) 5.05 V

(3) 10.67V

I I

I - . 1

\ i 1 1.5 2.0 2.5 3,0 3,5 4.0 4.5 ENERGY (eV)

Fig. 7. MOS electroref lectance spectra of re- duced n-type Ti02 for light polarized

perpendicular to the c-axis.

128
For the cases of "dry" configuration these spectra show, in general, features at the same energy above 3 eV as for the case of aqueous ER (see figures 1 and 2), although there

are some differences in line-shape. Below 3 eV there is some indications of the 1.7 eV structure but it is not as clear as in the case of the electrolyte ER. For example, figure

4 shows that as the bias is increased an optical feature in the range of 1.8 eV begins to emerge. This can be seen clearly for the case of the 14.3 V bias. Figure 6 shows structure

in this energy range for bias values of 7.11 and 10.19 V although here the phase is inverted. No field-induced splitting of the 3.15 eV optical structure is observed for the dry configuration ER spectra. The differences between the aqueous and "dry" ER data is probably due

to the fact that higher surfaces fields are generally attainable in the electrolyte case.

No hysterisis effects were found for any of the three ER configurations employed.

After data at the highest bias was recorded the lowest bias spectrum could be recovered

without any compensation of the dc voltage.

The possible interpretation of our data for 'lou < 3 eV in terms of interference effects can be eliminated. For example, it might be argued that in the electrolytic configuration

it is possible to form a layer on the material due to electrochemical action. In the MS or

^fOS case a semi-insulating layer might be formed at the TiO^surface by the field-induced

migration of the oxygen vacancies near the interface. The f ie Id- induced modulation of the

thickness or refractive index of such a layer could produce interference effects. Such considerations can be ruled out since it has been clear Ly demonstrated that interference effects

of this nature produce oscillatory spectra with a well defined phase change [l9,20], i.e. the

spectra would swing positive and negative. We have not observed this pattern in any of our data for < 3 eV. In addition, the lack of hysteresis seems to rule out the formation of

any layer, particularly in the electrolyte case.

4. Discussion of Results

The main results of this work are the observation in the ER spectra of structure at 1.7 eV (below the direct gap), the field-induced splitting of the 3.15 eV peak and the

strong electric field dependence of these structures. We will now discuss the relationship of these results to the defect and surface states in reduced Ti02.

Intrinsic TiO^ has its lowest direct gap at about 3.2 eV and an indirect gap at a some- what lower energy of about 3 evCl5,16]. The direct transition is presumably due to 0(2p) - Ti(3d) transitions and is characteristic of the perovskite-type materials in which the Ti

ion is in a TiO^ octahedral environment. Therefore, in agreement with previous workers we

attribute our ER peak at 3.15 eV to the direct transition of the intrinsic material.

While optical absorption measurements of stoichiometric TiOo show no structure for h(j0<3ev,

after reduction optical transitions are observed in this region [6,7J. Of particular interest

is the fact that a number of investigators report structure around 1.7 eV[6-9]. The photoconductivity spectra also shows a peak in this general region[9]. These experiments reveal

that in the process of reduction levels are introduced in the band gap in the region around

1.7 eV above the valence band (and hence about 1.5 eV below the conduction band). Our data

reveals the first indication of these levels in a reflection or ER experiment and their

strong field dependence.

Surface state experiments have also found a response in this energy range. Using

energy loss spectroscopy Henrich et al[ll] have reported surface state transitions at about

1.5 eV. Frank et al. have investigated the electrode response of n-type TiO^ in acetonnitrile with a large number of electroactive compounds differing widely in their standard

potentials. They report a level located about 1.2 eV below the conduction band[lO].

The above mentioned studies show that there are levels in both the bulk and surface of

reduced Ti02 at about 1.5 eV below the conduction band. These results suggest that there

is a close connection between the bulk defect and surface states in this material. This is not completely unexpected if one considers the following situation. The bulk Ti atoms in

TiO^ are in a 4 + valence state and hence have no d^-electrons . Surface ions would have a 3+ valence due to the absence of some of the ligands that are present in the bulk. But

129
creation of such a surface is analogous to the process of reduction in which oxygen vacancies are created with the resultant change in valence of a neighboring Ti atoms. In fact,

Henrich et al only see a signal in their energy loss measurements on pure^^iO^ after Ar-ion

bombardment of the surface which sputters off oxygens thus leaving the Ti ions on the surface. It is therefore not surprising to find this intimate connection between bulk defect

states and surface states in reduced Ti02

.

We next turn our attention to the field-induced splitting of the 3.15 eV peak and the

strong field dependence of this and the 1.7 eV structure. Morin and Wolfram have calculated

d-electron surface states in this class of transition metal oxides by considerations of the

termination of some of the d_-electron orbitals of the intrinsic material. As has been discussed above there is an intimate relationship between the d^-electron surface states and

defect states introduced by reduction of the mater ial [l2 ] . Hence, we will use the model

of these authors to interpret our results and will henceforth talk about d^-electron surface/

defect states.

Morin and Wolfram report surface bands which are lowered by about 1.5 eV from the con- duction bonds at F in the Brillouin zone. In addition the shape of the surface/defect bands

is such that at the M point they are at the same energy as the intrinsic conduction bands at

r. On this basis the 1.7 eV peak is assigned to transitions from the top of the valence

band to the surface/defect states at F. There are two degenerate transitions at about 3 eV, one from the valence to intrinsic conduction band at F and another from the valence band to the surface/defect states at the M point. If we further assume that the effect of the dc bias electric field is to uniformly displace the surface/defect bands to lower energies the

field- induced-splitting of the 3.15 eV structure and shift of the 1.7 eV feature are explained. The former effect is accounted for by the removal of the energy degeneracy between the

intrinsic conduction band at F and the surface/defect states at M. Also the field-induced

lowering of these bands would explain the shift to lower energies of the field-dependent

s tructures

.

Such field-induced shifts are not completely unreasonable when it is considered that

the position of the surf ace /defect bands have been deduced from the changes in surface

electrostatic potential at the transition metal site. In addition, large field-induced

changes of ER positions (~ 0.5 eV) havg been observed in KTaO^ in the electrolyte configur^^ion^for surface fields of about 10 V/cm. If we assume an electron density of abgut

10 cm then for a bias of 10 V our surface fields are estimated to be about 2 x 10 V/cm

[4,5,16].

The intimate connection in TiO^ between defect states introduced by reduction and

d^-electron surface states has been strongly suggested by this study and other recent investigations. The results of this experiment can be qualitatively explained on the basis of

the Morin and Wolfram d_-electron surface/defect state model. These conclusions may also be

applicable to other transition-metal oxides of this class such as SrTiO^.

We wish to thank T. Wolfram for several very useful discussions.

References

[l] Fujishima, A., and Honda, K., Nature 238, 37 (1972). Hardee, K. L., and Bard, A. J.,

J. Electrochem. Soc. 122, 739 (1975).

[2] Marvroides, J. G-, Tchernev, D. I., Kafalas, j. A., and Kolesar, D. F., paper presented

at the workshop.

[3] Sastri, M. V. C, Srinivasan, V., and Viswanathan, B., in Modern Aspects of Solid State

Chemistry (Rao, C. N. , Ed.), Plenum Press, New York (1970), p. 425.

[4] Cardona, M. , Shaklee, K. L. , and Pollak, F. H., Phys. Rev. 154, 696 (1967).

[5] Cardona, M. in Modulation Spectroscopy (Seitz, P., Turnbull, D., and Ehrenreich, H.,

Ed.) Academic Press, New York (1969), and references therein.

130
[6] Cronemeyer, D. C, Phys . Rev. 87_, 876 (1952).

[7] Breckenridge, R. G., and Hosier, W. R., Phys. Rev. £1, 793 (1953).

[g] Addiss, R. R. Jr., Ghosh, A. K., and Wakim, F. G. , Appl . Phys. Letters 12_, 397 (1968).

[9] Ghosh, A. K., Wakim, F. G., and Addiss, R. R. Jr., Phys. Rev. 184, 974 (1969).

[10] Frank, S. N., Laser, D., Hardee, K. L. , and Bard, A- J., paper presented at this

workshop

.

[ll] Henrich, V. E., Zeiger, H. J., and Dresselhaus, G., paper presented at this workshop.

[12] Morin, F. J., and Wolfram, T., Phys. Rev. Letters 30. 1-214 (1973) and references there

in. Wolfram, T. , paper presented at this workshop.

[13] Adolph Meller Co., Providence, R.I.

[14] Manufactured by Monsanto Co., New York, N.Y.

[15] Frova, A-, Boddy, P. J., and Chen, Y.S., Phys. Rev. 157, 700 (1967).

[16] Vos, K., and Krusemeyer, H. J., Solid State Comm. 15, 949 (1974).

[17] Cardona, M. , and Harbeke, G. , Phys. Rev. 137, A 1467 (1965).

[is] See, for example, Sze, S. M. , in Physics of Semiconductor Devices , Interscience

,

New York (1969)

.

[19] Subashiev, V. K., Surface Science 37, 947 (1973).

[20] Aspnes, D. E., J. Opt. Soc. Am. 63, 1380 (1973).

131

National Bureau of Standards Special Publication 455, ElectTocatalysis on Non-Metallic Surfaces,

Proceedings of Workshop held at NBS, Gaithersburg, Md. , December 9-12, 1975. (Issued November 1976)

Electron Energy Loss Spectroscopy of Surface States on Titanium and Vanadium Oxides ^

Victor E, Henrich, H. J. Zeiger, and G. Dresselhaus

Lincoln Laboratory, Massachusetts Institute of Technology

Lexington, Massachusetts 02173

Using electron energy loss spectroscopy, we have studied tlie surface electronic structure of Ti02, Ti^Og, TiOg gy and V2O3. A surface -state transition in the region of 1. 5 to 2. 5 ev is present in all of these materials when the surface is re- duced by Ar-ion bombardment. Subsequent exposure of the surfaces to oxygen or water vapor reduces or completely eliminates this transition. The origin of the transition in terms of d-electron population of surface cations is discussed.

Keywords: Adsorption; electron spectroscopy; ligand-field splitting;

surface excitons; surface states; transition-metal oxides.

We are reporting here electron energy loss measurements made under ultrahigh-vacuum conditions on single crystals of Ti02, Ti203 and V2O3, and on polycrystalline TiO^, where x = 0. 87. Our interest in these materials stems from the recent observation that Ti02 and other transitionmetal oxides catalyze the photoelectrolysis of water [1]} and we are investigating the surface electronic structure of these catalysts.

In electron energy loss sfjectroscopy, a monoenergetic electron beam of energy Eq strikes the surface of a solid, and some of the electrons excite electronic transitions of energy A in the sample

and emerge with energy Eq - A. This is just the electron analog of Raman scattering. The extreme

surface sensitivity of this technique arises from the very short mean free paths of electrons in solids. Electrons with energies between about 40 and 150 eV have mean free paths of only 4 or 5 A [2], so surface-state transitions can contribute structure to an energy loss spectrum comparable in magnitude to that due to bulk transitions.

The samples were either oriented and polished or fractured to expose fresh surfaces before

mounting in the ultrahigh vacuum system. All samples were then Ar-ion sputter-etched to produce

atomically clean surfaces. This treatment produced enough disorder on the surfaces that no lowenergy-electron-diffraction patterns were observed, except in one case that will be discussed later. We will concentrate on the low-energy region of the energy loss spectra, and this is shown for Ti02 in figvure 1. The first derivative of the secondary-electron energy distribution n(E)was recorded

vs energy loss, so peaks in n(E) occur at midpoints of negative slopes. The resolution in these

measurements is about 0. 6 eV. The elastically reflected peak has been attenuated relative to the rest of the spectrum. Four peaks are seen in this region of the spectrum, and their locations are

indicated by the arrows. The method used for determining peak amplitude and width is also indicated.

Peaks B, C and D, which occur at 5. 5, 10 and 13. 5 eV, are present in all of the titanium oxides. We

attribute them to cross transitions from the ground state of the O^" ion to the relatively narrow, high

density-of-states band of titanium d-levels.

Peak A, which occurs at about 1.5 eV, is of primary interest here. It is due to a surface-state

transition, and its amplitude varies tremendously on different crystal surfaces and as a function of exposure to ambient gases. Figure 2 shows the energy loss spectra for Ar-ion sputtered surfaces of Ti02, Ti203 and TiOx. Peak A is much larger on the last two samples, while the triplet is similar

on all crystals. A very weak peak is sometimes seen between peaks B and C, and it is shown on the insert below the bottom trace. This happens to be for a TiOx surface that had been exposed to water

vapor, although water vapor is not necessary for the presence of this peak. We have not as yet

determined its origin.

This work was sponsored by the Department of the Air Force. Figures in brackets indicate the literature references at the end of this paper.

133
A

Fig. 1. Electron energy loss spectrum for

Ar-ion sputtered Ti02 • The derivaatlve of the secondary-electron

energy distribution, dn(E)/dE, is

plotted vs energy loss. Locations

of peaks in n(E) are indicated by

the arrows.

30 20 10 0 ENERGY LOSS (eV)

Fig. 2. Electron energy loss spectra Fig. 3. Electron energy loss spectra

for Ar-ion sputtered Ti02, for Ti02 after Ar-ion sputterTi203 and TiOo.87' The in- ing and subsequent exposure

sert shows a small peak that to O2

.

is sometimes present at about

8 eV. 134
Peak A is affected dramatically by exposure of the surfaces to either oxygen or water vapor.

Figure 3 shows the effect of oxygen exposure on the spectra for Ti02. As little as 1 Langmuir of oxygen begins to reduce the amplitude of peak A relative to the rest of the spectrum, and on heavy

exposure the peak can completely disappear. The exposures shown here are 10^, 10^ and 10^ Langmuir of oxygen. A similar reduction in the amplimde of peak A with oxygen exposure is seen

for Ti203 and TiOx, although on those surfaces the peak never completely disappears. The results

of measurements on several crystals are summarized in figure 4, where the amplitude of peak A,

normalized to that of peak C, is plotted vs surface oxygen concentration as measured by Auger electron spectroscopy. Peak A always decreases with increasing surface oxygen concentration, although

different surfaces have somewhat different functional dependences. The spectra from other surfaces

of Ti02 and Ti203 that we have measured, most of them fractured and thus not having high symmetry

orientations, he in the range of these data but do not lie right on these curves. This implies a dependence on crystal face symmetry as well as on oxygen concentration,

this dependence in more detail. We are currently examining

& 6

o o

TlOj (001)

TljOj (110)

TijOj (unari«nt«d)

TIOqj^ (polycryttallim) -i

Fig. 4. Amplitude of peak A, normalized to peak

C, vs surface oxygen concentration as determined by Auger spectroscopy for

several samples.

OM aee aea oro an

( o*<TI> ) ''E*'* HEIGHT

TIO2

(5.

0.69 0.70 a71 0.72

( o-t-<Ti> ) '^^''^^ ^^'^^ HEIGHT

0.73

Fig. 5. Effect of H2O and O2 exposure on amplitude

of peak A for TiOz0.74

135
Ar-ion sputtered surfaces of these materials always have the lowest oxygen concentration, with

the concentration increasing with oxygen exposure. It thus appears that ion bombardment is preferentially removing oxygen from the surface, leaving it titanium-rich compared to the bulk. This is corroborated by the absence of peak A on Ti02 surfaces that have been fractured in ultrahigh vacuum,

where the surface has essentially the stoichiometry of the bulk. Oxygen exposure restores the missing oxygen atoms, thus changing the environment of the surface titanium ions.

Exposure of the surfaces to water vapor oxidizes them in the same manner as does oxygen, and

figure 5 compares the effects of oxygen and water vapor on a surface of Ti02. The ordinate and

abscissa are the same as in figure 4. No differences are seen in the spectra for samples exposed to oxygen or water vapor.

Although the dependence of the amplitude of peak A on surface oxygen concentration is somewhat

different for different samples, the spectra from all of the oxides are indistinguishable for a given

amplitude of peak A. This is shown in figure 6, where we have plotted both the location and full width

of peak A versus its amplitude, normalized to peak C. These are data for all three compounds, and

for exposure to both oxygen and water vapor. The location of the peak varies from 1.4 eV when its amplitude is almost zero to 2.0 eV at the maxim.um amplitude, and its width increases from less than

0. 5 eV to about 1.4 eV.

2.0 r

>^ i.oi- >-o

UJz

PEAK A

• 990 A A A

r o AA

'^L • oaAo* A A

A *A

• «0 O OAA A

• • • ao O A

••

OA o o

• ao •• o

• • •

OAA o A AA

A> CO A A* ••

AE,

OA A A

A A

• TiOo

° ^'203 A TiO,0.87

10 AMPLITUDE OF / PEAK A VPEA

N K C /

Fig. 6. Energy loss, Ej^, and width,

AEj^, of peak A vs amplitude

of peak A for several

samples

.

The electronic structure of even a highly ordered crystal surface is invariably more complicated

than that of the bulk, and we do not yet have a clear picture of the origin of the surface states on the titanium oxides. However, one can consider some simple, purely ionic, models which are consistent

with most of the experimental observations. One extreme is to consider the ions to be essentially

free-ion-like. The bulk Ti ions in Ti02 will be in a 4+ valence state, but the surface ions, even on a cleaved surface, could have a 3+ valence due to the absence of some of the ligands that are present in the bulk. Neither the Ti3+ nor Ti4+ free ions have any observed transitions in the vicinity of 2 eV [3]. Removal of surface oxygen atoms, however, could create Ti2+ ions on the surface, and the free

Ti2+ ion does have optically observed transitions between 1.0 and 1. 8 eV [3]. Peak A in Ti02 would

then be attributed to these transitions. In Ti203, where the bulk Ti ions have a 3+ valence, the surface

Ti ions could be 2+, by the same argument used for Ti02. Peak A would then always be seen on Ti203, even when the surface had essentially the bulk stoichiometry.

Another approach to the sinrface- state transitions is to consider the ligand-field-split levels of the Ti d-electrons. If we assume that the Ti ions on a Ti02 surface that has not had any oxygen sputtered

off are in a 4+ valence state, they will have no d-electrons. Removal of some oxygen would create Ti3+ ions on the surface, and they have a 3dl configuration. The d-levels will be split by ligand

136
fields, and peak A would then correspond to transitions between these ligand -field -split 3d states. In 71203, even the bulk Ti ions have a 3d-'- configuration, so peak A should always be present. When

some ox-ygen is sputtered off of Ti203, Ti^"'" ions could then be created on the surface. The ligandfield-split spectrum for the resulting Sd^ configuration would be more complicated, with several

transitions over a range of 1 or 2 eV [4], As more oxygen atoms are removed from the surface of either oxide, more different valence states and ligand environments are possible for the surface

titanium ions, and this should cause peak A to broaden and shift to higher energy as its amplitude in- creases, as we have observed.

A more sophisticated approach to the problem of surface states on transition-metal oxides, which

allows for their partially ionic nature, is currently being taken by Wolfram [5]. He is performing

cluster calculations in the tight -binding approximation and finds surface states arising from d-to-d

transitions in the surface titanium ions. We are currently extending our energy loss measurements

to the atomic core levels of both surface and bulk ions in an attempt to determine which of these

models is applicable.

We have also looked at the low-energy region of the energy loss spectrum of V2O3, and this is shown in figure 7. Two peaks, B and C, occur within 0. 5 eV of the location of two of the peaks in the titanium oxides, but no third peak is seen. Peak C is quite a bit broader than in the titanium oxides,

however, and the third peak may just not be resolvable. These transitions presumably have the same

origin as those in the titanium oxides. Two peaks occur at very low energies in V2O3, peak 1 at

1. 3 eV and peak II at 2. 6 eV in this sample. Peak II has a sensitivity to oxygen similar to that of peak A in the titanium oxides, and figure 8 shows the decrease in amplitude of peak II with relatively

small oxygen exposures and its disappearance after large exposures. Peak I is essentially unaffected

by oxygen. V2O3 has the same crystal structure as Ti203, the major difference being one additional

electron on the V ion. Although we have not performed as extensive a set of experiments on V2O3 as

on the titanium oxides, we can say that peak II clearly arises from a surface-state transition, and its origin is probably analogous to that of peak A in the titanium oxides. Peak 1, on the other hand, could

arise from either a surface or a bulk transition.

I

Fig. 7. Electron energy loss spectrum

for Ar-ion sputtered V2O3(001).

Locations of peaks in n(E) are indicated by arrows.

30 20 10 0 ENERGY LOSS (eV)

We have just begun to study peaks B, C and D in Ti02 in greater detail. Excellent ( 110) surfaces

of Ti02 can be prepared by annealing to about 800°C in vacuum. We have made preliminary measurements of the energy loss spectra from this surface for 100 eV primary energy at both grazing and

normal incidence and at 2000 eV primary energy, normal incidence. The 100 eV, grazing incidence

spectrum should be dominated by surface transitions, while the 2000 eV spectrum should arise almost

entirely from bulk transitions. All three peaks occur in all of the spectra, but peak D is much weaker

at 100 eV, grazing incidence, than in the other two spectra, indicating that it has more bulk character

than the other peaks.

137
In summary, we have begun to extract information on the surface electronic structure of a class of materials that has great academic and practical importance. Surface states are present that depend on crystal face symmetry, cation valence and ligand environment. The presence of magnetic ions,

partially filled d-levels and phase transitions in these materials, together, of course, with their

catalytic activity, makes the study of their surfaces extremely interesting, both experimentally and

theoretically.

References

[1] Fujishima, A. and Honda, K. , Nature 238, 37 ( 1972); Mavroides, J. G. , Tchernev, D. I., Kafalas, J. A., and Kolesar, D. F. , Mat. Res. Bull. 10, 1023 (1975).

[2] Tracy, J. C. , Electron Emission Spectroscopy (W. Dekeyser et al. , Eds. ) D. Reidel, Dordrecht

(1973), p. 305.

[3] Moore, C. E. , Atomic Energy Levels , Vol. 1, N. B. S. Circular No. 467, U. S. Government

Printing Office, Washington (1949).

[4] McClure, D. S. , Solid State Phys. 9, 399 (1959).

[5] Wolfram, T. , paper presented at this Workshop.

138
National Bureau of Standards Special Publication 455, Electrocatalysls on Non-Metallic Surfaces

,

Proceedings of Workshop held at NBS, Gaithersburg, Md. , December 9-12, 1975. (Issued November 1976)

Dislocations in ZnO as Channels for Electron Transfer

from Surface Species

2

S. Roy Morrison and Thomas Freund

Stanford Research Institute

Menlo Park, California 94022

The results reported show that dislocations strongly affect the

properties of the ZnO electrode and that with the help of electrochemical

measurements a fairly complete picture of the electrical and chemical behavior of dislocations can be developed. It is shown that dislocations in

ZtiO act as hole or electron traps in the space charge region, as recombination

centers, as points of electric breakdown under anodic polarization, and as current paths leading to exchange current even in an indifferent electrolyte.

The dislocations can be passivated by annealing or by precipitation of

impurities. A qualitative model is presented that can account for the

observed affects of dislocations on the solid state and electrode

properties.

Key words: Dislocation; electrochemical; exchange current; semiconductor;

traps; zinc oxide.

1. Introduction

In studies of ZnO as an electrode in aqueous solution, we [1,2,3] and others [4,5]

have observed and reported that the dark anodic current is not associated with intentionally

added reducing agents of intermediate redox potential. For example, when measuring electron

exchange between ZnO and the redox couple ferro/ferricyanide , one finds that the cathodic

current is proportional to the ferricyanide concentration, but the anodic current is independing of the ferrocyanide concentration. We have suggested that the anodic current is

associated with flaws in the ZnO surface, and the current is limited by the rate at which

anodic current can pass through the flaws. If such a model is valid, the influence of this

behavior will be of significant interest in electrocatalysls at the semiconductor/electrolyte

interface. As will be discussed in more detail below, at flaws the presence of active energy

levels that are separate from the conduction and valence bands permits unwanted redox

processes to occur, and thus the selectivity of the electrochemical process will be lowered.

Such behavior will be deleterious also in basic studies of electrode reactions. For example,

tunneling has become a popular and powerful technique [6,7] for studying energy levels in

solution. However, to induce tunneling through a semiconductor surface barrier, very heavy

doping of the solid is necessary. Such doping will often lead to flaws, in particular dislocations, in the material, and much of the so-called "tunneling current" may in actuality be current through the flaws.

1^Supported in part by ARO-D

gNow at Xerox Corporation, Webster, New York, 14580

Figures in brackets indicate the literature references at the end of this paper.

139
One important type of flaw that is almost always present is dislocation. This type of

flaw is easily produced and removed. To test the Influence of flaws on the electrode

properties of ZnO (in particular the exchange current), we prepared a surface with an unusually high density of surface dislocation loops to determine the characteristics that dislocations imparted to the electrode properties.

A study of dislocations on the ZnO electrode was reported by Gerischer and his

co-workers [8] . They concerned themselves primarily with the bulk properties (hole trapping

in the space charge region) of the semiconductor, rather than the influence on electrode

reactions. In our studies we examined both aspects of the dislocation behavior. In addition,

we have been interested in possible techniques to remove the deleterious influence of dislocations, including high-temperature annealing and the precipitation of impurities at the

dislocations to determine the effect on the electrode properties.

2. Experimental

Most of our work was done on the (0001) face of ZnO crystals obtained from 3M

Corporation in St. Paul, Minnesota. Dislocations were introduced by lapping the crystals

using 15-micron grit. The properties of the surface so treated were examined, then the sam- ples were etched step by step until the effect of the dislocations was removed. The etch

used was 85% H P0^» which etches the (0001) face at a rate of about 1 micron per minute, but

does not etch ?he (0001) face. On the (0001) face it is found that only scratches and dislocations are etched.

Figure 1 shows a sketch to introduce the model assumed. For reference. Figure lb shows an edge dislocation, which is the end of an extra half plane. The "dislocation pipe" so formed, shown encircled, is a linear flaw extending through the crystal beginning and ending

on a surface. With our samples lapped on one side, the dislocation pipes must begin and end

on the same surface, and Figure la shows the type of dislocation loops expected. By removing

layers of the crystal surface in a stepwise fashion by etching, measurements can be made with

successively lower dislocation density in the space charge region.

Fig. 1. Dislocations at a lapped surface.

The density of the space charge near the surface was determined from the voltage

dependence of the capacity [9] using eq (1):

2 -2

Vg - kT/q = (1/2) qN^A Ke^C . (1)

where N is the density of immobile positive charges in the space charge region, A is the area of the electrode, C is the differential capacity of the space charge depletion layer.

140
< is the dielectric constant, g is the permittivity of free space. The parameter V is the

o s surface barrier potential given (in the depletion layer regime) by:

V = V + const (2)

s a

where V is the electrode potential of the ZnO (relative, say, to calomel). Most of the

a measurements reported below were made in one typical sequence. The exceptions are measurements of annealing and doping, where the effect of the dislocations was removed by a method

other than etching.

Figure 2 shows a plot of the current/voltage characteristics for the same (0001) surface

after (a) polishing, (b) etching 8 minutes, and (c) etching a total of 33 minutes. In all cases the readings labeled "dark" were taken after the sample had been left 16 hours (overnight) in the dark. The readings were all taken in the same order: beginning slightly

anodic, decreasing to 0.4 volt cathodic, then increasing the voltage in the anodic direction.

At each voltage the sample was held for 30 seconds before the reading. A definite procedure

was adopted since the current is time dependent. The zero current potential of the electrode

was close to the same value, (0.1 V relative to a saturated calomel electrode) in all three

cases.

The curves show a substantially higher current, both in the anodic and cathodic directions,

resulting from the presence of the dislocations.

Figure 3 shows various curves of C versus voltage, needed to determine N^^ by eqs. (1)

and (2) . The measurements were made on the same ZnO sample used in the experiments of

Figure 2. However, the solid curves represent measurements on the opposite face (from those

of Figure 2), the (0001) face. First, the (0001) face was polished and the capacity/voltage

dependence measured. Curve (a) shows the "dark" measurements, curve (b) the results when the

surface is illviminated by a tungsten lamp (the intensity will be discussed below). A much

141
Fig. 3. Determination of net charge density in the space charge region.

higher capacity is observed during illumination. After these readings the (0001) face was etched for 8 minutes in 85% H^PO^. As mentioned above, H^PO^ is a very slow etch on the

(0001) surface of ZnO, attacking primarily scratches and dislocations. After such a treatment, it is observed (curves c and d) that the capacity is at an intermediate value and is no longer sensitive to illumination.

We suggest that because the etch does not attack the (0001) face but only the dislocations, only a modest H^PO^ etch is necessary to excavate dislocation pipes a few

thousand angstroms and thereby remove dislocations from the space charge region. In the

analysis we will relate the N variation with illumination to the dislocation density and D „ ,, curves (c) and (d) can be considered as the blank , giving for a dislocation-free

sample

.

The contrasting etching behavior of the (0001) face, the faces used for the stepwise

etching experiments of Figure 2, is illustrated by curves (e) and (f) of Figure 3. These

measurements were made after the sample was lapped and etched for eight minutes, and so can

be compared directly to curves (c) and (d) . In contrast to the 0001) face, after etching

8 minutes the capacitance is still photosensitive, indicating the presence of dislocations.

This result is expected since H^PG^ does not rapidly and selectively etch dislocations on the

(0001) face, but very slowly etches the entire surface layer.

As will be discussed in the Analysis section, the results in Figure 3 suggest that

either electrons or holes can be trapped at dislocations. Clearly, if both electrons

and holes can be trapped, the dislocation should act as a recombination center and a generation center for carriers. To examine the characteristics of dislocations as centers

for electron-hole recombination, the photocurrent as a function of electrode potential was monitored

.

Figure 4 shows the photocurrent, the total current under illumination, plotted as a

f unction of inverse capacity. The light intensity wa^ 7.5% of that used in the experiments

of Figure 3. The surface area of the sample was 25 mm , the light source was a tungsten lamp.

The photocurrent for the well-etched sample provides a useful measure of the absolute light

intensity in these experiments. Later we will discuss the reason for the use of inverse

capacity on the abscissa, and the interpretation of these results. However, it is

142
E

RECIPROCAL CAPACITY, 1/c farads '

Fig. 4. Photocurrent vs reciprocal of the capacity for a specimen.

immediately apparent that the photocurrent at a given space charge thickness ( a given

capacity) is decreased by the presence of dislocations.

Summarizing ^Figures 2-4 make it clear that the presence of dislocations causes excess

exchange current at the electrode at zero net current, causes changes in N^, and causes a decrease in the photosensitivity.

Experiments were performed attempting to passivate the dislocation activity. Results

will be presented below showing the passivation when dislocations are annealed at high

temperature. Such annealing is to be expected, and the result of interest in the present

study is the temperature range on ZnO (T >900 C) at which it is observed.

The effect of a treatment in hydrogen in removing the deleterious effects of dislocations

is shown in Figure 5. The procedure is based on saturating the sample with hydrogen (10) at

a moderate temperature (300 C for the results of Figure 5) and passivating the dislocations

by subsequent annealing at a lower temperature where the hydrogen concentration is supersaturated .

Figure 5 shows a series of measurments analogous to those of Figure 4, showing the

j hydrogen passivation effect. The light intensity for the measurements of Figure 5 is

substantially higher than that for Figure 4. Curve (1) shows the photocurrent for a sample

o that was well etched, saturated with hydrogen by treatment before at 300 C in 1 atm H^, then

quenched to room temperature. Curve (2) shows the subsequent lowering of the photosensitivity

I by lapping to introduce dislocations, while maintaining the sample at room temperature to

! prevent hydrogen diffusion [10]. Curves (3), (4), and (5) are results after successive treatj ments at 250 C in helium. Curves (3), (4), and (5), compared to (2) show that helium

I annealing of the hydrogen-containing sample restores the photosensitivity to the polished

sample, presumably by permitting hydrogen to diffuse to and to precipitate at the dislocations. V/e believe that the lower photosensitivity of curve (5) as compared with curve

I (4) results from too much annealing time —the hydrogen has had an opportunity to move to

the surface and evaporate into the gas phase.

143
5. Photocurrent vs reciprocal of the

capacity for a specimen treated in

hydrogen before the introduction

of dislocations.

Fig.

100

INVERSE CAPACITY, l/C — cm2 ^uP-^

200

2.1. Analysis of the Results

The curves of Figure 2 show the influence of dislocations in causing a higher exchange

current at the surface of a sample with no intentionally added redox couple. Suggested

reasons for this activity will be discussed in the next section.

The results of Figure 3 suggest by the following argument that hole and electron trapping occurs when dislocations are present. The carrier dominant in the space charge region under

illumination will be holes, the carrier dominant in the dark will be electrons. If either

or both are trapped in the space charge region, the slopes in Figure 3 will differ depending on whether the sample is in the dark or illuminated. The curves for the (0001) face after

etching, curves (c) and (d) , have identical slopes, indicating neither hole nor electron

trapping. Tl^g sloge of curves (c) and (d) thus gives the true donor density in the sample

us^d, 5 X 10 cm . However with dislocations present, under illumination the slope of the

C /V curve is much lower than that of curves (c) and (d) , indicating a much higher concen- tration of positive charge in the space charge region, i.e., hole trapping. On the other

hand measurements in the dark with dislocations present show a much higher slope than curves (c) and (d) , indicating electron trapping (the effective is much less than the true donor density). Thus, dislocations can trap either electrons or holes.

The photocurrent data of Figure 4 are displayed in the form shown to permit calculation

of the minority carrier diffusion length L , using the following model: If we assume that

p the tungsten lamp produces hole-electron pairs uniformly in the surface region of the crystal,

and assume that holes produced in the space charge region will be swept to the surface by the

electric field, then the total photocurrent J will be determined by the number of holes

p produced in the space charge region plus the number produced within a hole diffusion length

of the edge of the space charge region. Thus:

J pop = Q'Kx + L ) (3)

where x is the thickness of the space charge region, the parameter I is the light intensity,

o

144
and is a proportionality constant. The parameter x is given by the parallel plate

capacity formula:

C = Kg A/x (4) o o

with K the dielectric constant, and the permittivity of free space. Then (3) becomes:

J pop = a^(AKe c""*" + L ) (5)

-1 A plot of J versus C , as shown in Figure 4, should be a straight line with the intercept

p

(at J =0) yielding -L . Now by definition:

P P

1 L = (Dt)^ (6) P

where D is the diffusion constant for holes, a constant in this study, and j is the lifetime

of photoproduced holes. Thus, a variation of the intercept of the Jp/C plot with dislocation density gives the variation in the minority carrier lifetime with dislocation density.

Three regions are apparent from Figure 4. Under highly cathodic conditions (large

capacity), the current is more negative than the straight line region that we attribute to eq (5 ) . At highly anodic voltages, the current is much higher than can be accounted for

by hole current as per eq (5). The case of high anodic voltag^ corresponds to electrical

breakdown. Thus, the display of the data according to J vo C permits very clear the

separation of breakdown current/photocurrent/cathodic region.

Table I shows the various parameters obtainable from Figure 4 where the breakdown field

is obtained by calculating the electric field E at the surface from the capacity and the

apparent donor density of the semiconductor using the Gauss relation

E = eN X /Kg (7) Do

together with eq (5) . The breakdown field E in Table I is defined as the value of E where the

B observed current is double that extrapolated from the straight line region.

In the presence of dislocations, the value of L decreases to the point where the

apparent L is negative —recombination is occurringPin the space charge region. The breakP down field also decreases with increasing dislocation density.

Measurements of L and E for a sample, which was first lapped to produce dislocations,

P B then annealed in air at temperature T to determine if they are removed, are shown in Figure

6 as a function of the temperature^T. It is observed that the dislocation loops appear to be removed by annealing at T > 900 C.

145
Table 1. Diffusion lengths and lifetimes of treated zinc oxide.

Treatment surface Ip (cm) (s) Eg (V/cm)

Lapped (0001)

Etched 3 min (0001)

Etched 8 min (0001)

Etched 33 min (0001)

Etched 3 min (0001)

("dislocation free")

-2 X 10

1 X 10

2 X 10

4 X lo"

-5b -6b

-5

3 X 10

1 X 10

4 X lo" 1.6 X 10

-12b

-10

1 X 10

-7

1 X 10C 1.8 X 10^ 2.4 X 10^ 7.4 X 10"

No breakdown

observed up

to lOOV.

Assuming D = 1 cm^/s

.

"The simple theory for Lp and T is not applicable, although qualitative information

is obtained from the calculated values of L„.

Fig. 6. Annealing of dislocations at high temperature, as observed

by the effect of annealing treatments in room temperature

measurements of the diffusion length and the breakdown

field.

146
3. Discussion

From the above measurements we can (a) conclude that dislocations can have a significant

effect on the electrode properties of a semiconductor electrode; (b) conclude that the dislocation effect can be removed either by a high- temperature anneal or by precipitation of a chemically active species on the dislocation. This section will be devoted to a discussion

of a model for the dislocations that can account for the observed behavior and suggest why

the dislocations should be so active at a semiconductor/electrolyte interface.

From modern theories of ionic semiconductors [11] as applied to states associated with

incompletely coordinated ions (usually assumed to be at the surface) and qualitatively from

the Madelung theory [12] , it can be concluded that an incompletely coordinated anion may

give rise to donor energy levels and that an incompletely coordinated cation to acceptor

energy levels in the bandgap. These levels have orbitals derived from the atomic orbitals

of the respective ions. At a dislocation, a line of incompletely coordinated anions and

cations exists which may be alternating or may be dominated by either anions or cations,

depending on the crystallographic direction of the dislocation as it wends its way through

the crystal. Thus, we Suggest that observed hole traps are associated with the anionic donors

and the observed electron traps are associated with cationic acceptors.

These electron and hole traps are in a line, and normally some overlap will be expected

between the orbitals of the anions, perhaps bridged by the cations or some overlap between

the orbitals of the cations, perhaps bridged by the anions. For example, in ^|ny senijg

g conductors, if the donor or acceptor impurity concentration rises to about 10 to 10 /cm , so the impurities are on the average 10 to 20 A apart, the orbital overlap is enough for

conductance through the "impurity band". In the present case we suggest that the injection

or extraction of electrons at the semiconductor/electrolyte interface is followed by "impurity

band conduction" along the dislocation pipe through the thin (about 1000 A) space charge region

of the semiconductor. The pipes need not have a high conductivity, but the conductivity

must be finite to explain the excess current in both the anodic and cathodic directions of

Figure 2,

The next problem is the identification of the electrode reaction. In the cathodic

direction we suggest that an electron migrating through the cationic levels to the surface

still has enough energy for reduction of traces of dissolved oxygen, the same reaction that must be associated with the cathodic reaction due to conduction band electrons. The anodic

reaction could conceivably be associated with the injection of electrons onto cationic

energy levels at the dislocation from some unknown impurity in the solution. It seems more

likely, however, that the anodic reaction is associated with the oxidation and dissolution

of incompletely coordinated lattice oxide ions from the dislocation pipe by a reaction such

as

2- + 20 + 2H _ H„0„ + 2e L 2 2 (7)

Such a reaction corresponds electronically to the injection of an electron e onto the cationic

acceptor level from the oxygen ion at the end of the dislocation pipe, and corresponds

chemically to preferential etching of the crystal at a dislocation.

The model provides an explanation for the action of the dislocations as recombination

centers and as centers of electrical breakdown. The action of the dislocations as recom- bination centers would correspond to capture of holes at the anionic donor levels, capture

of electrons at the cationic acceptor levels, and the movement of the trapped electrons or holes along the dislocation pipe until they recombine. The action of the dislocations as

147
centers of electric breakdown could be associated on the one hand with avalanche breakdown,

where free electrons and holes by impact ionization excite electrons from the anionic donor

levels to the valence band. On the other hand, it could correspond to Zener tunneling of

either electrons from the donor levels to the conduction band or holes from the acceptor

levels to the valence band. Either model would explain the breakdown data equally well.

The influence of hydrogen on the properties of the dislocations can be readily accounted

for by the formation °^_^^0^ groups (wher^^O^ is an incompletely coordinated oxygen ion at the

dislocation) or of HZn^^ groups (where Zn^ is an incompletely coordinated zinc ion). That

is, a proton (in the case of HOj^) or a hydrogen atom (in the case of HZn^ ) occupies a coordination position of the otherwise coordinatively unsaturated ions, forming a chemically

more stable complex. Such stabilization by complex formation would tend to increase the

iogization potential of the donor (O ) or decrease the electron affinity of the acceptor

Zn ^. In energy level terminology tSe complex formation would make the traps shallower in

both cases. Experimentally we find another donor (In) also tends to passivate the dislocations where an acceptor (Li) does not. The possible dependence gn the donor characteristics

is suggestive that the formation of the donor/acceptor complex Hzn^ may be the more important interaction.

4. Conclusions

We conclude directly from the evidence that electrode reactions normally not possible

on ZnO become possible when dislocations are present. In particular, we interpret the observations as suggestive that extra levels are in the bandgap region of the ZnO. Ions in

solution that cannot exchange electrons with the conduction band may be able to exchange

electrons with these levels associated with dislocations.

The most interesting new observations in this work regarding the "bulk" properties of

the electrode are that (a) the dislocations can act as channels for conduction of charge

through the insulating space charge region between the surface and the bulk semiconductor,

(b) dislocations are highly active as centers for electrical breakdown, and (c) the

deleterious effects of the dislocations can be eliminated by the precipitation of certain

impurities at the dislocation. In addition, the observations of others that dislocations

can act as traps for carriers and as recombination centers have been confirmed for ZnO,

5. References

[I] Freund, T. and Morrison, S. R. , J. Chem. Phys . 47_, 1543 (1967).

[2] Freund, T. and Morrison, S. R. , Surface Sci. % 119 (1968).

[3] Freund, T., J. Phys. Chem. 73_, 468 (1969).

[4] Berghe, R.A.L.V., Cardon, F., and Gomes, W.P., Surface Sci. 39_ 368 (1973).

[5] Berghe, R.A.L.V. and Gomes, W. P., Ber. Bunsenges. Phys. Chem. ]6_, 481 (1972).

[6] Hemming, R. and Mollers, F., Ber. Bunsenges. Phys. Chem. 76_, 475 (1972).

[7] Gerischer, H., B. Pettinger and Schoppel, H.R., Ber. Bunsenges. Phys. Chem. 7£,

451 (1974).

[8] Gerischer, H., Hein, F., Lubke, M. , Meyer, E. , Pettinger, B. and Schoppel, H. R., Ber. Bunsenges, Phys. Chem. 77_' 284 (1973).

[9] Dewald, J.R., J. Phys. Chem. Solids 14 155 (1960).

[10] Lander, J. J. and Thomas, D. G., J. Chem. Phys. 25^, 1136 (1956).

[II] Louis, E. and Yndurain, F., Phys. Stat. Solidi (b) , 175 (1973).

[12] Levine, J. D. and Mark, P., Phys. Rev. 144 , 751 (1966).

148
National Bureau of Standards Special Publication 455, Electrocatalysis on Non-Metallic Surfaces

,

Proceedings of Workshop held at NBS, Gaithersburg, Md. , December 9-12, 1975. (Issued November 1976)

Characterization of Semiconductor Electrodes in Nonaqueous Solvents

Steven N. Frank, Daniel Laser, Kenneth L. Hardee and Allen J. Bard

Department of Chemistry

The University of Texas at Austin

Austin, Texas 78712

The large number of simple and reversible redox couples available

in nonaqueous solvents allows convenient investigation of the band

structures of semiconductors using electrochemical techniques. Comparisons of the cyclic voltammetric behavior of these couples at platinum

and a semiconductor permits estimation of the relative energy levels of

the conduction and valence bands and provides a means of mapping the gap region for intermediate levels and surface states. This technique is

demonstrated with Ti02 and n-type Si semiconductors. In spite of the

large difference in band gaps between these semiconductors, their

electrochemical behavior is similar and is ascribed to a similarity in

the energies of their respective conduction bands and the energies of an intermediate energy level in Ti02 and the valence band in Si.

Key words: Acetonitrile ; cyclic voltammetry; electrochemistry;

nonaqueous; semiconductor; silicon; titanium dioxide.

1. Introduction

Although nonaqueous solvents offer a number of distinct advantages when compared to water for electrochemical investigations with semiconductor electrodes, few studies have

employed these solvents. Investigations of rather limited scope have been carried out with

Ge [1,2],-'- Sn02 [3], and more recently with GaP [4,5] and GaAs [5] in nonaqueous media.

Among the forementioned advantages of nonaqueous solvents are the availability of a large

number of kinetically simple and well understood charge transfer couples covering a broad

potential region, an extended potential range due to increased solvent stability and fewer

troublesome surface phenomena. Well-purified nonaqueous solvents are also expected to offer some relief from the undesirable electrode decomposition often found at positive

potentials with certain semiconductors, since water is often a participating reactant in the

decomposition reactions.

The current theory of the semiconductor-solution interface has been presented in

several recent reviews [6,7] and papers [s]. Semiconductor behavior differs from that of a metal in several important respects. Unlike a metal, modulating the potential of a semiconductor results in little modulation of the potential at the surface of the semiconductor

(unless the surface is degenerate) with most of the potential drop between the semiconductor and solution interiors occurring within the semiconductor space charge region, in

which the concentration profiles of the charge carriers and band energies vary in an exponential fashion from the field-free interior to the semiconductor surface; the surface

conduction and valence band energies remain essentially invariant with applied potential.

Changes of the rates of electrochemical processes with potential result not from a variation of the energies of the bands as in the case of a metal but from changes in the

Figures in brackets indicate the literature references at the end of this paper.

149
concentrations of the majority and minority charge carriers at the surface. The rather

sizable photocurrents observed when a semiconductor is illuminated with light of energies

greater than the band gap result from an increase in the number of minority carriers at the

electrode surface.

The oxidized and reduced forms of the solution species can be represented by two guassian energy distributions that intersect at an energy corresponding to the standard

potential of the redox couple [9]. Typically the widths of these distributions are 0.5 to 1.0 eV. Since the energies of these distributions are fixed relative to the bands of the

semiconductor and electron transfer must be isoenergetic , the observed behavior may be no electron transfer, only oxidations or reductions (rectification) , or both oxidations and

reductions (metal-like behavior) depending upon the extent of overlap of the energy distributions of the solution species with bands of the semiconductor, the type of doping

(acceptor or donor), the participation of minority or majority carriers, the band gap energy, the presence of surface states or energy levels in the forbidden band gap, and

whether the electrode is illuminated.

Below we present the partial results of our investigations [lO,ll] of two semiconductors, n-type Ti02, a wide band gap semiconductor (Eg = 3.2 eV) and n-type silicon, a narrow band gap semiconductor (Eg = 1.1 eV) . Using the above principles and the behavior

of the semiconductors in the presence of a number of solution redox couples, several

aspects of the electronic structure of the semiconductors in contact with the nonaqueous

solvent acetonitrile are explored. The location of the conduction and valence bands and

intermediate energy levels are deduced for Ti02 . The behavior of a wide band gap versus a narrow gap material is contrasted.

2. Experimental

Solvent and compound preparation, electrode pretreatments , materials handling, and

instrumental details are presented elsewhere [l0-12]. Single crystal Si and both single

crystal and chemically vapor deposited films of Ti02 were employed. Tetrabutylammonium

perchlorate was used as an electrolyte to increase solvent conductivity. The electrochemical cells were of a three compartment design that prevented mixing of the anode and

cathode solutions. The potential of the semiconductor and platinum working electrodes

was controlled versus a reference electrode. The cell voltage was applied between the

working and a platinum counterelectrode

.

The primary electrochemical technique used was cyclic voltammetry [l3]. In brief

this technique involves application of a triangular voltage wave form to the cell. For substances that are rapidly oxidized and reduced, so that the ratio of the oxidized to the

reduced form obeys the Nernst relationship, the separation of the resulting current peaks

for oxidation and reduction is about 58 mV for a one-electron process. The standard

potential lies midway between the two peaks on the potential axis.

3. Results and Discussion

Figure 1 shows the potentials of the current peaks for the oxidation and reduction of

a number of the substances used in these investigations both at platinum and the semiconductor electrodes. Also included are the estimated band energies for Ti02 and Si,

determined either from the behavior of the photopotential or the resistive component of the

electrode impedance or from the behavior in the presence of the various solution redox

couples [lO,ll]. The potentials given for the conduction bands in Figure 1 correspond

closely to the flat band potential for the semiconductors (absence of a space charge

region) . . ,

Since both semiconductors are n-type the majority charge carriers are electrons in the

conduction bands. Possible currents due to minority carrier generation in the space

charge region were calculated to be negligible. In order for there to be electron exchange

between the semiconductor and solution, the energy distributions of the solution species

150
must overlap an energy level in the semiconductor capable of accepting or donating electrons. For solution species with fairly negative standard potentials the energy distributions of both the oxidized and reduced forms will overlap the conduction band. In addition

at potentials just slightly negative of the flatband potential the semiconductor surface

becomes degenerate due to the large number of electrons at the surface and the semiconductor

behaves like a metal. This metal-like behavior was observed at both Ti02 and Si for sub- stances with standard potentials negative of the flatband potentials. For substances

with standard potentials a few hundred millivolts positive of the flat band, only the energy

distribution for the oxidized form has good overlap with the conduction band and only

reductions should be observed. As seen in the figure p-benzoquinone fulfills this criterion at Ti02 and Si. p-Benzoquinone can be reduced but not reoxidized at Ti02. The

reduction occurs with considerable overpotential when compared to platinum; rapid reduction

rates cannot occur positive of the flatband potential because of the depletion layer at the Ti02 surface. A similar overpotential is observed for p-benzoquinone reduction at Si,

although reoxidation is sometimes observed at Si with certain surface pretreatments which

activate the electrode surface.

Compounds with standard potentials near the flatband potential may behave either as at

a metal or with rectification depending on the particular shapes of their energy distributions. This transitional behavior was exploited to determine the location of the conduction

band in Ti02 . Ru (TPTZ)^^^ E° - -0.8 V (not shown), is reduced and reoxidized at Ti02 in a manner similar to its behavior at platinum. Anthraquinone , E - -0.9, on the other hand, is

reduced but its reoxidation is much slower than at platinum. On the basis of this behavior

the conduction band for Ti02 was assigned the level shown in Figure 1.

z0 -3 <

c

-2

T-

< -1 <

in > 0

>

< 1 z

UJh

0

Q.

\\\\\\\

Ec

Ev

777777

DP A

Ru^2

AO

BO — TMPDt

10 M P+

Th+.

TiO.

BZN

DPA

R+2 J

tpA

AO

BO

r

TMPDt

J

10 MP."

RifS

rBZN

Ru^2

1^ AO

BO

AO

BO

TMPDt

TMPD*'

////

Pt p-Si n-Sl

Fig. 1. Proposed band structures of Ti02 and n- and p-type Si and the electrochemical behavior of various redox couples at these semiconductors and

at Pt. Horizontal lines to the right of the vertical lines indicate

the cyclic voltammetric peak potentials for reduction while those to the left indicate the potentials for oxidation.

For couples with very positive standard potentials significant overlap of their energy

distributions with the conduction band is precluded. The reductions observed for these

compounds at Ti02 between -0.3 and -0.4 V must either occur by electron tunneling through

the space charge or by an intermediate energy level at this position. At the carrier con- centrations (~ lO-'-^) with which we are concerned tunneling seems most unlikely [l4] so that

the existence of an intermediate level appears probable. This level along with the

151
conduction band may also contribute to the reduction of TMPD"*". That this level can donate

but not accept electrons is demonstrated by the lack of any reoxidation of TMPD and p- benzoqinone

.

Illuminating the electrode with light of energies greater than the band gap creates

electron-hole pairs near the electrode surface. Production of holes in the valence band

provides a mechanism for removing electrons from the intermediate level. In line with this

reasoning, the anion radical of benzoquinone is oxidized when the Ti02 is illuminated (but not in the dark)

.

Although the band gaps of Ti02 and Si differ by a factor of 3, the electrode behavior

of the two is remarkably similar. (Extremely positive redox couples cannot be studied at

Si due to oxidation of the Si surface.) This similarity results from the conduction bands

of Ti02 and Si being at nearly the same location and the intermediate energy level observed at Ti02 occurring at almost the same place as the valence band of Si. In addition

to the previously mentioned similarities both display almost identical behavior with

anthraquinone . If one assumes that the valence band of Si and the intermediate energy

level of Ti02 can serve the same function with TMPD+, it is not surprising that its reduction occurs within =60 mV on the two electrodes.

The holes created in the valence band by illuminating the Si can accept electrons from

solution species. With illumination all of the compounds with standard potentials positive

of the flat band are reoxidized at Si. An interesting aspect of the photooxidations at semiconductors which is demonstrated at both Si and Ti02 is that the oxidations can occur

at potentials negative of their oxidations at platinum (i.e., with "negative over- potential") . TMPD is oxidized at potentials nearly 1 volt more negative at Si under

illumination, than at platinum. The main requirement for photoassisted oxidation is that

the electrode potential be positive of the flatband potential so that holes are forced to the surface and electrons to the interior.

4 . Acknowledgments

We wish to thank the National Science Foundation and the Robert Welch Foundation for

support of this work.

5. References

M. D. Krotova and Yu. V. Pleskov, Solid State Phys . , 3^, 2119 (1963).

H. B. Mark, Jr. and B. S. Pons, Anal. Chem. , 40, 110 (1966).

T. Osa and T. Kuwana, J. Electroanal. Chem., 2_2, 389 (1969).

R. Landsbergh, P. Janietz and R. Dehmlow, Z. Chem., 15^, 38 (1975).

R. Landsbergh, P. Janietz and R. Dehmlow, ibid . , 15 , 106 (1975).

H. Gerischer, in "Physical Chemistry: An Advanced Treatise," Vol. 9A, H. Eyring,

D. Henderson, W. Jost, eds . , Academic Press, New York, 1970.

H. Gerischer, in "Advances in Electrochemistry and Electrochemical Engineering," Vol.

1, P. Delahay and C. Tobias, eds., Inter science Inc., New York, 1961.

W. P. Gomes and F. Cardon, Z. Phys. Chem., 86, 330 (1973) and references therein.

R. A. Marcus, J. Chem, Phys., _43, 679 (1965).

S. N. Frank and A. J. Bard, J. Amer. Chem. Soc, Tl_ (1975) (in press).

D. Laser and A. J. Bard, J. Phys. Chem. (submitted).

K. L. Hardee and A. J. Bard, J. Electrochem. Soc, 122, 739 (1975).

R. S. Nicholson and I. Shain, Anal. Chem., 36, 706 (1964).

B. Pettinger, H. R. Schoppel and H. Gerischer, Ber . Bunsenges. Phys. Chem., 450

(1974).

152
Comment on Paper by S. N. Frank, et at., "Characterization

of Semiconductor Electrodes in Non-aqueous Solvents"

V. E. Henrich

Massachusetts Institute of Technology

Lincoln Laboratory

Wood Street

Lexington, MA 02155

It seems to me that your work on silicon might form the basis for a comparison between

the silicon surface states in an electrolyte and the large body of work that has been

done on the surface states on silicon under ul trahigh-vacuum conditions. This would be

a very important step in bridging the gap between the ul trahigh-vacuum environment and

the real world. Do you have sufficient data on silicon surface states, similar to that

you showed for TiOa, to permit such a comparison?

153

ELECTROCHEMICAL PROCESSES ON NON-METALLIC SURFACES - 1

National Bureau of Standards Special Publication 455, Electrocatalysis on Non-Metallic Surfaces,

Proceedings of Workshop held at NBS, Gaithersburg, Md. , December 9-12, 1975. (Issued November 1976)

Chemical Theory of Chemisorption and Catalysis

on the Surfaces of Non-Metals

Robert L. Burwell, Jr. Ipatieff Laboratory

Department of Chemistry

Northwestern University

Evanston, Illinois 60201

This paper surveys chemisorption and catalysis on nonmetals

primarily from the point of view of localized, coord inatively unsaturated surface (cus) sites and with use of analogies to the chemistry of

coordination and organometallic complexes. Topics covered include

chemisorption on oxides of main group elements and of transition

elements, heterogeneous catalytic hydrogenation and oxidation on oxides

of transition elements, and chemisorption on group IV elements and

group III-V compounds.

Key words: Chemisorption; heterogeneous catalysis; catalytic

oxidation; catalytic hydrogenation; surface

coordinative unsaturation; cus.

The usual heterogeneous catalytic reaction appears to proceed by chemisorption of one

or more reactants followed by surface reactions involving chemisorbed species and finally

by desorption of products and regeneration of the catalytically active sites. In its

present stage, the study of the mechanism of heterogeneous catalytic reactions is primarily

concerned with the determination of the sequence of chemisorbed intermediates which con- stitutes the catalytic reaction. Inevitably, then, the study of chemisorption and the

nature of the binding in the adsorption complex are essential to the study of mechanism.

The great majority of the more recent studies of chemisorption on non-metals has

involved oxides or mixed oxides and almost exclusively adsorption from the gas phase.

Studies of carbons, sulfides, silicon, and germanium constitute but a small fraction of

total studies. This paper will consider primarily chemisorption and catalysis on oxides

and usually on polycrystalline or amorphous ones. Oxides which have been investigated may

be divided into two groups.

(a) Oxides of main group elements: SiOg, AI2O3, SiOg-AlgOs, and zeolites. These are

jLmportant in technology both as adsorbents and as catalysts and supports for catalytically

active ingredients.

(b) Oxides of transition elements, for example, Cr^Os, NiO, C03O4, VgO^, BigOs' M0O3 , and also of post-transition elements like ZnO and also oxides of rare earth elements.

Studies of chemisorption on these materials have primarily been motivated by interest in

their catalytic properties.

In the 1950's there was a strong interest in interpreting chemisorption on oxides-- even AI2O3 and Cr^O^ at lower temperatures--in terms of the band theory of semiconductors.

In recent years, most interpretations have tended towards a localized bonding based on analogies with coordination and organometallic complexes. In this, the surface site is treated as a coordinatively unsaturated surface (cus) ion and the adsorbate as a ligand

which completes or partially completes the coordination sphere of the (cus) ion. Ligand

field theory provides a more elaborate extension of this approach. A survey of the

155
crystal field and ligand field aspects is provided by D. A. Dowden, Catal. Revs. , 5.» 1 (1971)' Some of the factors which led to this change of attitude are:

(a) The results of the earlier effort did not appear to be very useful.

(b) The earlier procedure ignored surface geometry, i.e., what crystal planes were

exposed, the edge/face ratio of surface atoms, surface defects, etc., and also such

chemical matters as the degree of hydroxylation of the surface.

(c) The earlier theory appeared to have no contact with the rest of chemistry.

(d) Few catalysts of any interest were typical semiconductors.

(e) The area of homogeneous catalysis by organometallic complexes developed and

provided many examples of catalytic reactions rather similar to those in heterogeneous

catalysis. Here a localized bond was necessarily involved. The existence of the homogeneous analogs proved nothing about heterogeneous catalysis, of course, but it did

establish that band theories were not essential to explain all catalytic reactions.

Presumably there is an exact quantum mechanical solution to the nature of binding in

any particular adsorption complex, but this solution is not accessible to present techniques and approximations must be used. The coordination complex approach might give a useful interpretation of the surface chemistry and catalytic behavior of an oxide or even

of germanium, whereas it might be useless in interpreting the effect of chemisorption upon

such bulk properties of the adsorbent as conductivity. That is, the effect upon the bulk

might be a relatively small perturbation upon the actual binding in the adsorption complex

and different approximations might, at this stage, be appropriate in discussing binding

and conductivity.

Measurement of the extent of or the rate of the chemisorption of various adsorptives

upon various adsorbents as a function of pretreatment, pressure, and temperature dates back more than fifty years. However, such investigations alone, although essential first steps,

give little information about chemical binding. Analogies with molecular reactions were

useful but most of the molecular analogs, particularly the organometallic complexes, have

been discovered only in recent years. Infrared spectroscopy of adsorbed species has

probably been the most important of the newer tools but electron paramagnetic resonance

has also been important, particularly on oxides. Optical spectroscopy has also been useful.

MUssbauer spectroscopy has been employed but it is limited by the small number of useful

isotopes. X-ray photoelectron spectroscopy is beginning to have an impact. Although there are problems with this technique, it should prove useful in the so far difficult task of

identifying the oxidation state of the metal ion in transition metal oxides and of changes

which occur consequent to chemisorption. The techniques of surface chemical physics have

so far been applied mainly to tungsten and a few other metals.

We present below a review of the kinds of chemisorbed species which have been proposed.

The review is necessarily selective and the references have been chosen primarily to serve

as leading references to the literature. It should be recognized that many of the species

proposed as catalytic intermediates are speculative and are derived only from studies of

catalytic reactions. In some cases additional support for such species has been provided by

the techniques of the preceding paragraph and by isotopic tracer experiments, stereo- chemical experiments, reactant structure variation, etc. applied to the catalytic

reaction [1].^ The studies leading to the proposed chemisorbed species have almost all

involved solid-gas systems, aqueous systems have rarely been examined, and, indeed, most of

these chemisorptions are strongly inhibited by water vapor. That is, water molecules are strongly chemisorbed on most of the sites for chemisorption to be discussed.

It should be recalled that the most strongly adsorbed species is probably not a good

catalytic intermediate. The very weakly adsorbed species may not be activated, the very

strongly adsorbed species must be so low in the energy scale as to be relatively unreactive.

Rather, the species of an intermediate energy of adsorption is most likely to be the most

reactive.

''"Figures in brackets indicate the literature references at the end of this paper.

156
Chemisorption on Oxides of Main Group Elements

Silica. Silica gel has been extensively investigated as an adsorbent. After preparation and drying at about I5O °C, the surface of silica gel is free of water vapor and

covered with about 5 silanol (surface SiOH groups) per nm^ [2]. Silica gel is a selective

adsorbent. For example, it adsorbs olefins, arenes, and ethers much more strongly than

alkanes because the silanol groups are strong hydrogen bond donors. However, unless one wishes to classify adsorption by hydrogen bonding as chemisorption, pure silica gel in its

usual form does not exhibit chemisorption. Commercial silica gel contains about 0. 05^ AI2O3

which gives it a little of the characteristics of SiOg-AlgOs described below [3].

Alumina. The cubic aliuninas ( , 6- , etc. ) which come in high area forms are also

covered with (H groups when equilibrated with water at room temperatures [k]. Upon heating,

dehydroxylation occurs to produce Al~^ (cus) and 0^" (cus).

A13+ A13+ Al3+ A13+

0^" (cus) may serve as a basic site to chemisorb acidic adsorptives [5]. Partially dehydroxylated aluminas chemisorb pyridine in two forms as shown by studies using infrared

and Raman spectroscopy.

Silica-aluminas. Silica and alumina in which the ratio Si/Al is greater than about 3 form amorphous solid solutions. One can look upon this material as a condensation polymer

of Si(0H)4 and A1(0H)30H2. Thus, each Al^"*" will be accompanied by one H . One might

schematically represent the undehydroxylated surface by

H

0

-0-SiH2

0

0-AlH

0

O-SiAs judged by indicator studies, some of the surface AlOHg groups are as strongly acidic as concentrated sulfuric acid [5]. Olefins, for example, propylene, appear to be converted to carbonium ions upon adsorption. Silica-alumina catalyzes reactions of hydrocarbons which

closely resemble carbonium ion reactions. The material was used for many years as the

catalyst in catalytic cracking of heavier petroleum stocks.

Zeolites . Zeolites are crystalline aluminosilicates with networks of small channels

whose diameters are in the general range of 1 nm. As in silica-aluminas, an extra plus

charge must accompany each Al-^"*" and, as the zeolites are synthesized, this extra plus charge

is Na"*" but other positive ions can be introduced by ion exchange against Na"*". For example,

Co^'*' can be introduced and this ion can be amminated by treatment with ammonia. The cobalt

ammine complex reacts with oxygen to form Co^"*" O-O" as Howe and Lunsford have shown by EPR

examination [6]. The Co^"^ ions are too far apart for there to be any question that the

binding is other than localized. Non-amminated ions of Co^'^ exchanged into zeolites chemisorb carbon monoxide, nitric oxide, and methyl isocyanide [7].

157
Chemisorption and Catalysis on Oxides of Transition Elements

Soft and hard acids and bases (SHAB) ., Pearson's concept of soft and hard acids and

bases [8] is qualitatively useful in considering chemisorption and catalysis. In general,

soft indicates large and polarizable^ hard , small and less polarizable. The prototype of

the hard acid is the proton, of the hard base, the hydroxide ion. The terms soft and hard

must be distinguished from the terms weak and strong. Other things being equal, a soft

acid prefers to combine with a soft base, a hard acid with a hard base. However, a strong

soft base may combine well with a hard acid, e, g. , CH3" with H

Among bases, the order of increasing hardness is

F~ > Cl" > Br" > l" - ,

0 » S > Se > Te as in H2O vs. H2S

N » P > As as in NH3 vs. PH3

Tetracyanoethylene, ethylene, carbon monoxide, H , and C2H5 are soft bases although

consideration as acids or bases in the first three is complicated by back-donation of

electrons.

Small, highly charged ions are hard acids: h"*^, Li"*", Mg^"*", Al^"^, Fe^''', and Cr^"'", also

COg and SO3* Cu"*", Ag"^, Pt^"*" are soft. The higher the charge on the ion of a given

element, the harder it is. Thus, Co^"*" is hard, Co^+ is intermediate, and Co(0) is soft. F" binds more strongly to Al'^'*' than does I" but the reverse is true to Ag"*", HgO binds more strongly than does HgS to Al'^"'" but the reverse is true to Ag"*", CgH4 and CH3" should

bind weakly to Cr~^"*" (cus), HgO should bind strongly and, indeed, water is a strong poison

for heterogeneous catalytic reactions on chromia. Poisoning is a pervasive problem in

heterogeneous catalysis and the SHAB concept is useful in predicting what compounds will

be poisons for a particular catalyst.

Chromia (O^^gOs of various degrees of hydration, crystalline or amorphous). Like

alumina, chromia exposed to water at 25 °C has a surface which is fully hydroxylated, i.. e. the surface is covered with OH". When such chromia is heated above about 200 °C, water is

lost by condensation of surface OH" to form pair sites, Cr^"^ (cus) and 0^" (cus) [9]. A

sample of amorphous chromia heated in an inert atmosphere (for example, flowing helium or vacuum) to kOO °C, contains about one pair site per nm^.

As could be predicted from the SHAB concept, the chemisorption of carbon monoxide on activated chromia is weak. At -78 °C, adsorption is complete and irreversible and the

extent of such chemisorption can be used to measure the content in Cr^"*" (cus). Warming to

0 "C liberates most of the carbon monoxide. One would hardly expect all Cr^"*" (cus) to occur in identical geometries or to have the same degree of coordinative unsaturation

Initial dehydroxylation would generate Cr®"^ (cus) in five-fold coordination (the coordination number in bulk is six) and 0^" coordinated to three Cr^"*" rather than four as in

bulk. More extreme degrees of dehydroxylation could generate Cr^"*" in four-fold coordination. There is catalytic evidence for non-uniformity of Cr^"*" (cus) [9] and Borello and

Zecchina and also Peri have provided evidence from infrared spectroscopy that Cr^"'" of

C.N. = k exists [10].

A major difference between chromia and alumina is the possibility of changes in the

oxidation number of chromium. Heated in oxygen at °C, chromia is covered with a monolayer of Cr(Vl) [11]. Heating in hydrogen or carbon monoxide at 500° appears to generate

Cr(ll) on the surface. Chromia is an oxidation catalyst and its action as such presumably

involves alternation in the oxidation number of chromium- Oxygen adsorbs strongly on activated chromia. On chromias of lower degrees of hydroxylation, the initially formed

adsorption complex is probably Cr'*+ Og" [11] but on highly dehydroxylated chromia, the

complex is probably Cr^"*" [10]. These structures are rather speculative.

158
Chromia is also a catalyst under reducing conditions and, for exanple, it gives

substantial rates in the hydrogenation of ethylene at -100 °C. There is still some question as to the relative degree of involvement of Cr^''" vs. Cr^"*" in chromias which have

not been exposed to reducing atmospheres at temperatures above 3OO °C or so. However, it

may be noted that Cr^^"*" (cus) should be more readily reduced than bulk Cr^"*" because octa- hedrally coordinated Cr^"^ is strongly stabilized by the crystal field stabilization energy

but Cr^"*" is destabilized in the crystal field of octahedral coordination. In bulk chromia,

^ach Cr^"*" is octahedrally surrounded by six 0^". In what follows we will use Cr^+ (cus)

as the example of the surface site.

H2O adsorbs on the pair sites Cr^"*" (cus) 0^" (cus) in what is the reverse of the

dehydroxylation shown in eq (1). Formally, the water is dissociatively adsorbed as H"*" and

oh" . This process is heterolytic dissociative adsorption. At temperatures of 25 °C and

above, hydrogen appears to adsorb heterolytically to give

H2(g)

Cr,3+ 0^- Cr3+ (2)

and at 200 °C methane also adsorbs by dissociative heterolytic adsorption.

CH,(g)

Cr3+ o2- H3CCr3+ (3)

The forms H and CH3 are the crystal field formalism of coordination chemistry. There is considerable charge redistribution and adsorbed species are much less ionic and basic than

NaH and NaCHs. It will be noted that heterolytic dissociative adsorption does not involve

oxidation-reduction reactions in the adsorbent. There are a number of organometallic

analogs to the adsorbed a Iky 1, one of which is £-tolylCr (lIl)Cl2L3 [12] where L is tetra- hydrofuran coordinated through itfe oxygen atom and £-tolyl is formally

H H

H H

Alkanes exchange with deuterium in the presence of activated chromia at about 200 °C to give primarily the monodeuteroalkane. This reaction appears to proceed via heterolytic

dissociative adsorption of alkane and deuterium as in eqs (2) and (3), followed by surface

migration of H"*" and D"*", and then associative desorption. Hydrogenation, at least at temperatures above 25 °C, appears to proceed by a related mechanism: reaction of, for

example, C2H4 with CrH to give CrCH2CH3 followed by reaction of the latter species with OH

to regenerate the pair site and to form CgHg. There are hydrogenations homogeneously

catalyzed by organometallic complexes which appear to proceed by a very similar mechanism

[9]. In addition, it has been proposed that alkylethylenes can adsorb alternatively by

heterolytic dissociation- With 1-butfine the product would be Cr (CHsrrr CH— CH-CH3) OH". Reversion to butylene could involve addition of H"*" to carbon-5 to form the isomeric 2-

butene. The proposed intermediates derive largely from experiments using deuterium as an isotopic tracer. They are reasonable but somewhat speculative.

Reasonable, but even more speculative, is the proposal of reductive adsorption as in

the following example.

159
HeCg)

H H~ Cr3+ 0^- (f- Cr3+ Cr2+ 0 0 Ct^+ . (^)

Such a form of chemisorbed hydrogen could lose water and generate coordlnatlvely unsaturated

and dehydroxylated surfaces containing Cr^"*". Continuing reaction and loss of water would

result in reduction of CrgOs to CrO. It should be noted that the species formed in eq ('0

is isomeric with that formed in eq (2). Probably the reductive adsorption form is favored

at more elevated temperatures. Reductive adsorption of carbon monoxide can be represented.

CO (g)

0

I 2- C 2 Cr3+ (f- (f- Cr3+ Cr2+ o 0 Cr^+

COf:^04. . This oxide is an active catalyst for a number of reactions under reducing con- ditions and it rather resembles chromia in its overall behavior. It has been particularly

studied by Ozaki [IJ]. The active sites are pair sites of cus cobalt and oxygen ions and

hydrogen appears to adsorb heterolytically.

Zinc oxide. The species outlined above which were first proposed for chromia have

been established more strongly as occurring on zinc oxide by R, J. Kokes who employed infrared spectroscopy as well as deuterium tracers [ 15]. Zinc oxide is particularly suitable

for such a study . Its optical characteristics facilitate infrared studies and it is a much

poorer catalyst. Thus, the surface intermediates are less reactive and more amenable to

study. Kokes clearly established the occurrence of heterolytic dissociative adsorption of

hydrogen and of propylene, the formation of a surface 77-complex of ethylene, and reaction

of adsorbed ethylene with Zn^"*" H" to give Zn^"^ C2H5". The nature of surface intermediates

for these reactions are better known on zinc oxide than on any other oxide.

Adsorption of Oxygen on Oxides of Transition Elements

Chemisorption of oxygen. The following forms of chemisorbed oxygen have been proposed,

0^" (oxide ion), Og^" (peroxide ion), 0", Og (superoxide ion), O3" (ozonide ion), and 04". Substantial progress has been made in studying adsorbed oxygen because all but the first two are paramagnetic and subject to study by EPR. Moreover, the existence of ^^0 of nuclear

spin 5/2 has greatly facilitated this work. The hyperfine splitting observed with ^^0 has

permitted clear assignments and provided structural data. This subject has been reviewed

by Lunsford [I5]. It should be recalled that the reaction

02-(g) 0-(g) + e"(g)

is very strongly exoenergetic in the gas phase and that 0^" is stable against this reaction

in solid ionic oxides because it is stabilized by the Madelung energy of the solid.

Presumably, then, the reverse of the reaction above is less favored at the surface than in

bulk and therefore, 0" would be relatively more stable at the surface than in bulk. Further, in the gas phase, NgO" decomposes to and 0~ when the ion contains very small

amounts of energy above its ground state. One might expect, then, that treatment with NgO

would be a method of choice in making 0" ads.

A number of studies of the adsorption of oxygen on reduced, supported molybdenum oxide

have been reported [15,16]. O2" is formed and the two oxygen atoms are not equivalent.

Further, there is some contribution from covalent bonding. At least with silica as the

support, the Oa" is stable to 100 °C< The EPR-invisible 0^ ion is probably also formed

[16]. The Os" ion appears to be located at a positive ion on the surface and the situation

is probably similar to that of a number of superoxide ions known in molecular complexes

160
such as the Fe(ll)-pcrphyrin complexes of myoglobin and hemoglobin. Molecular analogs to these are known in solution and they have been formed on the surface of imidazole—modif ied

silica gel [ITJ- The reaction with oxygen can be represented

L-Fe(ll) + Og —^ L-Fe(lll) 0

where four nitrogen atoms in the planar porphyrin molecule are coordinated to the iron atom

in a plane perpendicular to the plane of the paper. L is an imidazole group. Co(ll') complexes behave similarly. The reduced molybdenum catalysts exhibit the EPR signal of Mo(V')

which declines in intensity as oxygen adsorbs. The process can probably best be represented as

Mo^"'" (cus) + 02(g) Mo^+ O2"

with partial contribution from covalent bonding. However, ref. I6 considers that attachment occurs at surface Mo®"*" although no explanation is given as to how an electron is transferred to the molecule of oxygen.

Treatment with NgO leads to an EPR signal of 0 . This species reacts more rapidly

with reducing agents than does Og" (ads).

An adsorbent prepared as V20^/Si02 behaves rather similarly although some 0 appears

during treatment with Og. At 25 °C, 0"ads reacts rapidly with Hg, CH4, C2H4, and CO Og"

is much less reactive. As judged by hyperfine splitting, O2" is in a form in which there

is interaction with ^^V.

Little definite is known about adsorption of oxygen as what is formally , but, of course, it is invisible to EPR. It is well known in coordination complexes as, for

example, in Vaska's compound [I8].

CI PPhc, CI P .0 \ /

Ir + O2 —y *Ir' i CO OC P ^0

Ph3

where Ph represents phenyl and the oxidation state of Ir increases from I to III consequent to the 'adsorption' of oxygen. A similar reaction may well occur in some cases on

oxides or the like but most of the currently known examples in coordination chemistry

involve metal atoms with d^ and d^^° configurations. Such configurations are rare in oxides

of catalytic interest.

Cimino and Stone separately and jointly have reported a series of important studies of

solid solutions of oxides of transition elements in such insulating oxides as MgO and

^^2^3 [19,20,21]. This work has provided information on two items: the adsorptive and

catalytic effect of ions of different transition elements in the same environment, the

effect of the dilution of a given ion upon these properties. It has been established that

ions of transition metals have adsorptive and catalytic properties even in high dilution in

insulating oxides where semiconductivity theory is clearly inappropriate. However, there

is an effect of concentration. In particular. Stone [20] concludes that electronic interaction exists among neighboring Cr^^ ions in solutions of higher concentrations of Cr^"^ in

(Cr^"*", Al^"'')203. At least under oxidizing conditions the purely localized treatment of

adsorption and reactions on chromia given under Chromia is probably an oversimplification.

161
Adsorption of Oxygen on Oxides of Non- Transition Elements

Adsorption on these materials tends to proceed to a much smaller coverage than on the

materials of the preceding group. Furthermore, since oxidation states more positive than

Mg^"*", etc. are very high in energy, the formalism used to represent the adsorption of oxygen

in the preceding section is inapplicable. MgO has been extensively studied. The electrons

used in forming the EPR-visible species appear to come from traps and radiation either

before or during adsorption is often needed to effect adsorption. NgO under these con- ditions gives an O" ads in which the extra electron is largely localized in one of the 2p

orbitals of 0". The extent of covalency in the binding is low. Og" is formed by exposure

to oxygen, the two oxygen atoms are equivalent and the axis of Og" must be parallel to the

plane of the surface (unlike the situation on the transition metal oxides of the preceding

section). This is also true for TiOs, ZnO, and SnOg.

The ozonide ion O3 is formed by the following reaction on MgO

o"ads + Og(g) —^ O3 ads.

The species is not very stable and it slowly decomposes to Og or 0 . At present, O3

appears to be a relatively uncommon form of adsorbed oxygen.

Charge Transfer Adsorption of Organic Molecules

The surface of MgO can adsorb certain organic molecules by charge transfer adsorption

which is apparently of a different type than that involved in adsorbing oxygen. Thus,

MgO after thermal activation adsorbs nearly a monolayer of nitrobenzene (NB) to form the

nitrobenzene radical anion [22]. On partially dehydroxylated MgO, the process appears to be

0^"(cus) 0~(cus) + NB —>

oh" oh" • • • NB*

whereas on highly dehydroxylated samples no proton hyperfine splitting is observed and the

following process was suggested.

C^-(cus) + NB —> 0" NB" .

A number of examples of forming radical anions and also radical cations such as (phenyl)3N''"

has been observed on such insulating oxides as Alg03 and Alg03,Si02.

Heterogeneous Catalytic Oxidation

Many examples of heterogeneous catalytic oxidation exist. Some occur on metals, for

example, the oxidation of NH3 to NO (which is then converted to nitric acid), the oxidation

of SOg to SO3 on platinum, and the oxidation of hydrocarbons and carbon monoxide on platinum. Many examples in which oxides are the catalysts also exist, for example, the

oxidation of SOg to SO3 at about h'^^ °C on VgOs ,KgO/SiOg (an unusual case of heterogeneous

catalysis because under operating conditions the catalytic ingredient is a melt of vanadium

oxide dissolved in molten potassium pyrosulfate [23 1), the oxidation of naphthalene to phthalic anhydride on a similar catalyst but at about 350 °C at which temperature the

catalytic ingredient is probably solid, and the oxidation of methanol to formaldehyde on Fe203,Mo03 at 375 For a description of the catalysts and the operating conditions of

industrial oxidation processes, see ref. 2h.

162
It is not difficult to find catalysts which catalyze the oxidation of hydrocarbons to water and carbon dioxide and this can be accomplished even under the demanding conditions

of treating automobile exhaust. The problem is selectively to oxidize the hydrocarbon part

way. Ordinarily, the starting hydrocarbon is more resistant to attack than most imaginable

intermediates. Sometimes, however, an intermediate possesses a rather high resistance to oxidation and good yields of the intermediate can be obtained, as, for example, in the

oxidation of naphthalene to phthalic anhydride. Usually, the problem is more difficult and

somehow one must find a catalytic system which for some reason does not further oxidize

the desired product. This matter is very poorly understood but probably there is no general

method for stopping the reaction at a desired point and the problem is almost as broad as chemistry itself. Enzymes give selective oxidation even of alkanes to perfection.

One must presumably activate either or both the hydrocarbon and the oxygen or other

oxidizing agent. The chemical processes are usually complicated, the catalyst is often

complicated, and operating temperatures are usually high. Only a few systems have received

much fundamental scientific attention and we know something about the mechanism of the

reaction in only a few cases. Further, there are not many homogeneous catalytic oxidations

which are appropriate as analogs. The high operating temperature in most heterogeneous

catalytic oxidations militates against this. However, a low temperature example follows.

The reaction between tert -butyl hydroperoxide and cyclohexene at about 80 °C produces

cyclohexene oxide and tert -butyl alcohol.

t-buOCH + C=C —^ t-buOH + ^ C-C^

/ X - -yM0O3 is a heterogeneous catalyst for the reaction and MoOg(acac)2 (acac is a bidentate,

negatively charged ligand, the conjugate base of acetylacetone) and some other, but far

from all, molybdenum complexes are homogeneous catalysts [25]. The system presents the

advantage that the oxidation number of the molybdenum can be monitored by optical spectroscopy. Coordinatively unsaturated Mo species appear to be involved in both the heterogeneous and the homogeneous instances.

Where elemental oxygen is the oxidizing agent, either oxygen is incorporated into the

product (e.g. , CO + l/20£ —> CO2), or, more commonly when a hydrocarbon is a reactant,

oxygen appears both in an organic product and in water (CH3CH=CH2 + O2 —* CH2=CHCH0 + HgO),

or, in oxidative dehydrogenation (CH3CH=CHCH3 + I/2O2 CH2=CH-CH=CHs + HgO), oxygen

appears only in water. We can imagine two extreme types of mechanisms:

(a) The oxygen in the product comes from adsorbed oxygen.

(b) The oxygen in the product comes from lattice 0^~ which is replaced by adsorption

of oxygen in a subsequent step.

For a simple example, consider the oxidation of carbon monoxide. In case (a), the

sequence might be

m"+(cus) + I/2O2 M^"^^^"^ 0"

^(n+l)+ Q- + CO m''"^(cus) + CO2

In case (b), the sequence could be

m"+02-+C0 M^""2^+ + CO2

^(n-2)+ ^ ^/202 02-

163
Overall, 02(g) is the oxidizing agent, but in the step that forms CO2: in case (a), both

M(ri+l)+ and 0" are the oxidizing agents;, in case (b) only M'^'*' is the oxidizing agent since

the oxidation number of oxygen is the same in 0^ and in COg.

For an oxide which is reluctant to be reduced, such as CrpOs, one might expect to observe case (a) under conditions such that oxidation to higher oxides is relatively slow. On the contrary one might expect to observe case (b) for an oxide like M0O3 if reoxidation

is rapid. That is, case (a) is to be expected where under operating conditions the catalyst

is in the fully reduced form and case (b) is to be expected where it is in the fully

oxidized form.

In actual examples of heterogeneous catalytic oxidation, the matter is complicated but

it appears likely that case (a) often, probably usually, leads to non-selective, complete

oxidation, whereas, selective oxidations are more likely to be favored by case (b).

Oxidations on Bismuth Molybdate. Bismuth molybdate, typically in the form BigMoO-sj

gives three remarkably selective oxidations at temperatures of about ^4-50 "C [26,273.

CH2=CH-CH3 + O2 —^ CH2=CH-CH0 + H2O

CH2=CH-CH3 + NH3 + 5/2O2 > CH2=CH-C=N + 3H2O

CH3-CH=CH-CH3 + 1/2 Og > CH2=CH-CH=CH2 + HgO

From isotopic labelling experiments using deuterium and '-^C it has been established

that in the first reaction propylene is dissociatively adsorbed as some kind of allylic

species. This step is rate limiting.

A remarkable feature of these reactions is that one may omit oxygen and find the

reaction to proceed with the usual selectivity until the catalyst composition falls to BigMoO^g. During this process the electrical conductivity increases and the EPR signal of

Mo^'*' appears. Treatment with oxygen restores the original properties and composition.

Keulks has shown the acrolein formed initially from ^^0^ is nearly devoid of ®0. Further,

the oxidized catalyst does not exchange rapidly with ^^02- It appears very probable, then,

that the oxygen which is incorporated comes from the lattice.

The crystal structure of BigMoOg consists of sheets of 0^", (Bi^O^)^'^ , and(Mo02)^+

alternating in the following manner

(BisOs)^"^ 0^" (MoOs)^^ (f~ (Bi02)^+ (f~ etc.

The oxygen between the layers containing the metal ions is clearly very mobile and the

possibility exists that the oxygen which replaces the 0^" consumed may be added at some distance from the point of loss.

The initial reaction may be

C3H6 + Mo'3+ + (f- —» (MoC3H5)^+ + oh"

During the subsequent stages, it appears likely that the adsorbed organic intermediates are not exposed to adsorbed oxygen and this may be the key to the selectivity.

Another component of selective oxidation catalysts, V2O5, has a crystal structure

related to that of BigMoOg [28]. Here planes containing vanadium and oxygen are separated

by planes containing only 0^~.

164
One final connnent must be made with respect to heterogeneous catalytic oxidations

Free radicals may propagate from the surface of the catalyst into the gas phase and initiate

reaction sequences in the gas phase. This is common in the oxidation of hydrocarbons in the

liquid phase on oxides of transition elements [29], and it is likely that it is a side

reaction in oxidations on BigMoOsAdsorption on Group IV Elements and on Group III-V Compounds. There has been a number

of investigations of adsorption on elements of group IV and on group III-V compounds, many

on single crystal surfaces by the methods of surface chemical physics. Research on these

materials has not been motivated by interest in their adsorptive and catalytic utility (so

far they have none) but, of course, by interest in their semiconductive properties. The

conceptual approach to this work has often been different from that to the work considered

earlier and intercomparisons are not always easy.

If one were to crush a crystal of diamond in an ultrahigh vacuum, one would break

myriads of carbon-carbon bonds. One might expect that as a first approximation, one could

consider the surface to react as if it were covered with free radicals even though the

density of free radicals would be so high that there should be some kind of interaction

among them which might even lead to surface reconstruction. One would expect hydrogen to chemisorb with some rather substantial activation energy to convert the diamond surface to an alkane one. Oxygen should chemisorb with a very small activation energy and initially

to form alkylperoxy radicals, '^C- + Og —^ -C-O-O* .

Adsorption experiments have been reported on silicon and germanium cleaved in an ultrahigh vacuum [JO]. There is fast adsorption of HCl, HBr, H2O, NH3, and PH3 at 25 °C. The

adsorption is probably dissociative since one molecule of HCl adsorbs for every two surface

atoms, one of H2O for every four surface atoms. LEED studies have been reported for the

adsorption of oxygen on single crystal surfaces of silicon and germanium cleaved in ultrahigh vacuum [31] and other experiments have been performed on the (111) faces of silicon

crystals cleaved in an ultrahigh vacuum [32]. In the latter case, the sticking coefficient

is about 0. 001 and saturation corresponds to one atom of oxygen per surface atom. Results

of this type have generally been interpreted in terms of "dangling bonds" on the clean

adsorbent. The term must be a close cousin to "free radical. "

The idea of surface coordinative unsaturation is less applicable here although the

basic idea of unsaturation persists. Further, the materials are good semiconductors and

one might expect to observe more substantial interactions between localized surface states

and bulk than in most of the examples considered previously in this paper.

Oxygen adsorbs on the cleaved (110) surfaces of both n- and p-type GaAs to a saturation

coverage of one atom of oxygen per pair of atoms of gallium and arsenic. The sticking

coefficient is larger on n-type (j x 10"^) than on p-type (l.k x 10"®) [33]- It was sug- gested that the Oads is on As. However, an EPR signal of Og" has been observed on crushed

GaAs [3^]. The signal is due to a much weaker form of adsorbed oxygen, one which can be

removed by desorption at -100 °C. The Og" is not ordered parallel to the surface and the

signal intensity is ten-fold larger on n-type than p-type. On n-type AlSb 5fo of the

unpaired electron is in an ^^Al p-orbital.

Interpretation of chemisorption on single crystal planes of these materials may be

complicated by surface reconstruction of the adsorbent, particularly with germanium and

silicon. Further, there are apt to be problems associated with the role of surface defects

and steps in experiments using the techniques of surface chemical physics. Special problems

may be associated with adsorptives which generate adsorbates with a substantial and identical

charge. Thus, on polar III-V compounds the ionic adsorption of oxygen to high coverage

should be more difficult than that of hydrogen chloride which could adsorb heterolytically.

Summary

A simmary of chemisorption and catalysis has been presented largely in terms of

localized bonding and (cus) sites in analogy to inorganic and organometallic coordination

chemistry. Of the approximate treatments yet considered, the author believes this approxi165
mation to be usually the most useful for the purposes of heterogeneous catalysis and of

chemisorptlon itself. Presumably, in the future it will be replaced by a still better

treatment. However, it should be noted that not all workers in the area employ the theoretical framework employed here. Band theory [35 ,3^ '\ continues to be employed by some workers

[36,37] and, on some semiconductors, with little contribution from the chemical nature of

surface [38].

References

[1] M. Boudart and R. L. Burwell, Jr. , "Techniques of Chemistry," Vol. VI, Part 1, edited

by E. S. Lewis, John Wiley and Sons, New York, l9Jk, p. 693- [2] C. Okkerse, "Physical and Chemical Aspects of Adsorbents and Catalysts," edited by

B. G. Linsen, Academic Press, London, 1970, p. 213.

[3] P, B. West, G, L. Haller, and R. L. Burwell, Jr. , J. Catal. , 2^, k86 (1973).

[^1-] B. C. Lippens and J. J. Steggerda, "Physical and Chemical Aspects of Adsorbents and

Catalysts," edited by B. G. Linsen, Academic Press, London, 1970, p. I7I. [5] K. Tanabe, "Solid Acids and Bases," Academic Press, New York, 1970. [6] R. F. Howe and J. H. Lunsford, J. Amer. Chem. Soc. , 2.7, 51^6 (1975).

[7] P. Pichat, "Catalysis, Heterogeneous and Homogeneous," edited by B. Delmon and

G. Jannes, Elsevier, Amsterdam, 1975? P- I83. [81 R. G. Pearson, editor, "Hard and Soft Acids and Bases," Dowden, Hutchinson, and

Ross, Inc. , Stroudsburg, Pa. , 1973- [9] R. L, Burwell, Jr., G. L. Haller, K. C. Taylor, and John F Read, Advan. Catal.

Relat. Subj. , 20, 1 (I969).

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A. Zecchina, L. Cerruti, and E. Borello, J. Catal. , 2^, 55 (1972). J. B. Peri,

J. Phys. Chem. , j8, 588 (197^)-

[11] M. P, McDaniel and R. L. Burwell, Jr., J. Catal., 36, 39^^ (1975).

[12] J, J. Daly, R. P. A. Sneeden, and H. H. Zeiss, J. Amer. Chem. Soc. , 88, h28j (I966).

[13] H. Nihara, T. Fukushima, K. Tanaka, and A. Ozaki, J. Catal. , 2^, 28l (1971).

[Ih] R. J. Kokes, Accts. Chem. Research, 6, 226 (1973).

[15] J. H. Lunsford, Catal. Revs., 8, I35 (1973)-

[16] M. Che, A. J. Tench, and C. Naccache, J. C. S. Faraday Trans. I., 7O, 263 (197^).

[17] 0. Leal, D. L. Anderson, R. G. Bowman, F. Basolo, and R. L. Burwell, Jr. , J. Amer.

Chem. Soc. , 21, 5125 (1975)-

[18] F, A. Cotton and G. Wilkinson, "Advanced Inorganic Chemistry," Interscience

Publishers, New York, 1972, p. 772. [19] A. Cimino, V. Indovina, F. Pepe, and F. S. Stone, Gazz. Chim. Ital. , IO3 , 935 (1973),

[20] T. A. Egerton, F. S. Stone, and J. C. Vickerman, J. Catal. , 3^, 299 (197^).

[21] A. Cimino, Chim. Ind. (Milan), 27 (197^), a review of ions in oxide matrices.

[22] D. Cordischi, V. Indovina, and A. Cimino, J. C. S. Faradav Trans. I, JO, 2189 (197^)

[23] P. Mars and J. G. H. Maessen, Proc. Int. Congr. Catal. 3rd, 266 (I965).

[2i|] C. L. Thomas, "Catalytic Processes and Proven Catalysts," Academic Press, New York,

1970, chapters 18-20. [25] F. Triforo, P. Forzatti, and I. Pasquon, "Catalysis, Heterogeneous and Homogeneous,"

edited by B. Delmon and G. Jannes, Elsevier, Amsterdam, 1975, P- 509- [26] G. C. A. Schuit, J, Less-Common Metals, ^6, 329 (197^).

[27] D. J. Hucknall, "Selective Oxidation of Hydrocarbons," Academic Press, London, 19jh.

[28] W. M. H. Sachtler, Catal. Revs. , k, 27 (1970).

[29] L. Ya. Margolis, Catal. Revs., 8, 2kl (1973).

[30] A. J. van Bommel and F. Meyer, Surface Sci. , 8, 38I (1967).

[31] J- J- Lander and J. Morrison, J. Appl. Phys. , 2^, 2089 (1962)^ 3h, ikll (I963).

[32] H. Ibach, K. Dom, R. Dorn, and H. Llith, Surface Sci. , 38, k33 (1973).

[35] R. Dorn, H. Luth, and G. J. Russel, Phys. Rev. B. , 10, 50i^9 (197^).

[3h] D. J. Miller and D. Haneman, Phys. Rev. B, 3_, 29 18 (1971).

[35] A. Clark, "The Theory of Adsorption and Catalysis," Academic Press, New York, N. Y , 1970, chapter XV. [36] F. F. Vol 'kenshtein, "The Electronic Theory of Catalysis on Semiconductors,"

Pergamon Press, Oxford, I963. [37] A. C. Zettlemoyer and R. D. Iyengar, "The Solid-Gas Interface," edited by E. A. Flood,

Marcel Dekker, Inc. , New York, N. Y. , I967, Vol. 2, chapter 26. [38] T. A. Goodwin and P. Mark, Progress in Surface Sci. , _1, 1 (I97I).

166
National Bureau of Standards Special Publication 455, Electrocatalysis on Non-Metallic Surfaces

,

Proceedings of Workshop held at NBS, Gaithersburg , Md., December 9-12, 1975. (Issued November 1976)

ELECTROCATALYSIS AT TRANSITION METAL OXIDE

AND CARBIDE SURFACES

P. Stonehart

Advanced Fuel Cell Research Laboratory

Power Systems Division

United Technologies Corporation

Middletown, Connecticut 06457

Due to the cost of noble metal electrocatalysts, significant

research has been carried out to find non-noble metal electrocatalytic

materials. Transition metal oxides, usually doped tungsten oxides

(bronzes), have been postulated previously due to their favorable

electronic conductivity and stability in electrolytes. The surface

chemistry of such electrocatalysts has been largely neglected but

interesting parallels exist for the electron injection and extraction

with heteropolyanions in solution. The heteropolyanions exhibit

electrochemical properties of "bronze" crystals and can be considered

to be microcrystall ites. Optical studies show correlations between

absorbance in the visible region and the electron density per transition metal atom. By and large, the transition metal oxide electrocatalysts have not yet fulfilled their expectations.

A related research direction has examined the activity of transition metal carbides for hydrogen oxidation in acid. Here it was found

that tungsten carbide was the most active material but that it required

a pretreatment. Surface studies established that the active form of

the tungsten carbide electrocatalyst was slightly carbon deficient.

Key Words: Electrocatalysis; heteropolyanions; phpsphotungstate

complex ions; "spillover" effect; transition metal carbides; transition

metal oxides; tungsten bronze; tungsten carbide.

It is only recently that electrocatalysis, as a distinct sub-area in the field of

electrochemistry, has taken on increased emphasis. This has been due to more sophisticated

developments in energy conversion devices, such as fuel cells and new battery systems.

The field of gas-phase catalysis has advanced further than that of electrocatalysis,

principally due to an understanding of the surface features on gas-phase catalysts and the

importance of highly dispersed (small particle size) catalytic materials. Nevertheless,

the concepts developed for highly dispersed gas-phase catalysts can be applied equally

well to electrocatalytic materials [1,2]^ The singular difficulty in understanding the

operation of electrocatalysts is the lack of surface diagnosis techniques that are capable

of examining the electrocatalyst surface in the presence of the electrolyte. To overcome

this deficiency, one has to draw analogies and inferences from other sources, as this

paper will show.

Noble metal electrocatalysts mostly have been studied but due to the cost of these

materials, significant research has been carried out to find non-noble electrocatalytic

materials. Due to their stabilities in acid electrolytes, transition metal oxides (usually

of tungsten and molybdenum) have been postulated as suitable materials. Empirically, it

has been shown that the incorporation of transition metal oxides, together with platinum,

into porous electrode structures produces a higher reaction rate both for hydrogen oxidation

Figures in brackets indicate the literature references at the end of this paper.

167
and hydrogen oxidation in the presence of carbon monoxide compared to the reaction rate

for platinum alone [3]. It has not been determined whether this increased activity was due to an electrocatalytic effect or simply due to the formation of an improved electrode

structure. In addition, alkali doped tungsten oxides in conjunction with platinum have

been shown to be beneficial for carbon monoxide electro-oxidation [4].

It is well known that the incorporation of alkali metals into tungsten oxides produces

"bronzes" with very low electronic resistance (fig. 1). It is apparent that the conductivity of the "bronze" does not depend on the particular alkali metal, but does depend on

600r

x value

1. Dependence of electrical resistivity on composition for some alkali -metal tungsten

bronzes. Data from work at Iowa State

University (ISU) and Cornell University

(CU).

the doping level. At low electrode potentials, hydrogen atoms enter the WO3 structure to form the analogous "hydrogen bronze" [5,6], so that it is not necessary that an alkali

metal bronze be formed initially. Unfortunately, although there is some considerable

knowledge on the solid-state properties of these "bronzes," very little information exists

concerning the surface structure, especially in the presence of an electrolyte.

In order to identify the electrochemical processes occurring at the transition metal

oxide surface, one may examine the properties of heteropolyanions. These can be considered

as micro-crystallites of WO3 and as such, the WO3 groups are all surface entities. The

two principle structures of heteropolyanions are: a) dodecametal lo with the "Keggin"

structure, and b) octadecametallo with the "Dawson" structure. These two structures are

shown in figure 2. The former has an approximately spherical cage of 12(W03) groups

surrounding a single hetero ion group, whereas the latter has 18(W03) groups in an elongated

Fig. 2. (a) Dodecatungstophosphate ion,

[PWiaOi+o]^". ('Keggin structure')

showing coordination polyhedra; (b)

octadecatungstodiphosphate ion,

[P2Wig062]^"> ( 'Dawson structure'

)

showing coordination polyhedra.

168
cage surrounding two hetero ion groups. Since these ions have very few atoms, and as microcrystal 1 ites (clusters) the atoms are all surface entities, one is entitled to ask

whether the electrochemical properties are at all representative of larger "bronze"

particles.

An examination of the potentiodynamic electro-reduction of the [PWi20i+o]^~ ion in acid on mercury [7] shows that addition of the first electron to the cage occurs in the

vicinity of +0.2V versus hydrogen, shown as peak Ic in figure 3. Reoxidation of the

reduced ion is shown by peak la. Addition of the second electron to the heteropolyanion

occurs at more negative potentials, shown as peak lie and the consequent reoxidation

(removal of the second electron) as peak Ila. Two features may be noted. Firstly, the

potential differences between peaks la and Ic are much less than between the peaks Ila and

lie. This means that the kinetic rate constant for addition and removal of the first

electron to the WO3 cage is greater than the corresponding rate constant associated with

the second electron. Although only one electron is shared among 12(W03) groups to give

IIW^''' and one W-"*" (since the electron is not delocalized but probably hopping), it effectively hinders the addition of the second electron. Secondly, at the reversible hydrogen

potential, only one electron is added to the cage. This would mean that for a hydrogen

anode, either at or slightly positive to the reversible hydrogen potential, only a few

tungsten atoms (1 in 12) would be expected to be as W^"*". The remainder are hexavalent.

There is a real question, therefore, whether the surface of a "bronze" particle would

exhibit the commonly assumed metallic conductivity or whether it would be semiconducting.

The latter is more likely. In addition, although a variety of low surface area "bronzes"

have been prepared for el ectrocatalytic studies, it has not been possible to prepare these

compounds with useful high surface areas. McHardy [8] studied this problem and found that

on introduction of alkali doped tungsten bronzes into an acid electrolyte, rapid leaching

of the alkali metal occurred, producing a skin of WO3 to the depth of 0.5 ym. It is not surprising, therefore, that when one micrometer particles of alkali doped tungsten bronzes

are prepared, the only properties they exhibit in solution are those of stoichiometric WO3

particles

.

1

c

oj 0

sS-

-1

-0.2 j_

+0.2 +0.4 E vs H2|H+1N

Fig. 3. Cyclic potentiodynamic i-V curves for

[PWizOi^o]^- in 1 M H3PO4. V = 0.0925 V s"!.

Ig is the current peak associated with addition of the first electron and I, is the

current peak associated with oxidation of the

[PW12O40]'*" product. IIc and Ilg are the

corresponsi ng cathodic and anodic peaks asso- ciated with the second electron transfer.

Stationary mercury electrode.

Even though these heteropolyanions contain so few atoms, one may conclude that the

electrochemical properties are indeed representative of "bulk" materials. On comparing

the potentiodynamic scans in figure 3 to potentiodynamic scans on sodium tungsten bronze

crystals shown by Randin [9] in figure 4, the correspondence is obvious. Five different

potentiodynamic scan rates are shown, but it is clear that the potentials of the peaks

correspond to the first electron reduction and reoxidation (la and Ic) obtained with the

heteropolyanions.

169
There is additional evidence from absorption spectra of the reduced heteropolyanions

that these compounds can be considered as micro-crystallites of "bronzes." This subject

is covered more fully in a discussion of the electrochemistry of transition metal oxide

bronzes and related compounds published elsewhere [10]. It is somewhat surprising that so few atoms in a "cluster" exhibit properties manifest by bulk materials, except that is exactly what the theoreticians working on cluster calculations have been saying. Although

the electrochemist is denied many of those surface techniques that are being employed to unravel mechanisms in gas-phase catalysis, by judicious choices of ionic species, such as the various heteropolyanions, it is possible to identify electron exchange processes

occurring at the surfaces of solid electrocatalysts. Given the high sensitivity shown by

electrochemistry for measuring reaction rates at fractional monolayer levels, then electrocatalysis becomes a sensitive surface probe [2].

Examining the electrocatalytic oxidation of hydrogen on platinum, and on tungsten

oxide promoted by platinum, Hobbs and Tseung [3] showed an apparent maximum in the reaction

rate as a function of platinum loading, shown in figure 5. This maximum is indicative of

an electrocatalytic enhancement of the hydrogen oxidation rate at low platinum levels.

The effect is made pronounced by the method of plotting the data, since the effect is negligible at higher platinum loadings where the high rate of hydrogen oxidation on platinum

masks any enhancement due to the presence of tungsten oxide. A later paper by the same authors [11] showed an identical enhancement using nominally sub-stoichiometric tungsten

oxide. The enhancement in the reaction rate occurred at the same nominal platinum loading.

It is likely that the electrocatalytic materials were the same in both instances.

— THEORETICS CURVES O EXPERIMENTAL POINTS

Tkighly active support [based ON POINT AT 0.09 mg Pt cm-'

INERT SUPPORT BASED ON POINT' AT 0.09 mg Pt cm-^

0.02 0.04 0.06 Pt loading/mg cm-^

0.08 0.10

Fig. 5. Specific activity of supported Pt for

hydrogen oxidation (current per mg Pt at +0.05 V) as a function of Pt loading on WO3. The two theoretical curves were drawn assuming an inert

support ( ) and an active support

( — —), respectively. Poor activity

at low loadings ( —) was attributed

to ohmic losses.

170
Since it has recently been shown experimentally [12] that the rate-determining step

for hydrogen molecule oxidation on platinum in acid is dissociative chemisorption of the

hydrogen molecule

^ 2H (Tafel Reaction)

then two options are available for explaining the electrocatalytic enhancement by tungsten

oxide. Either the adsorption isotherm for hydrogen atoms on the platinum surface is shifted to lower coverages (since hydrogen atoms are "poisons" for the hydrogen molecule

dissociative chemisorption step), or the reduced tungsten oxide particles in the vicinity

of a platinum crystallite become weakly active for hydrogen molecule oxidation on the

oxide surface . The latter situation is considered more likely, although the reaction

mechanism may not be the same as on the platinum surface. There is one further interesting

point to note: the electrocatalytic enhancement of the hydrogen oxidation rate by tungsten

oxide with low platinum loadings is exhibited in sulfuric acid electrolytes but not in phosphoric acid electrolytes. It may be conjectured that a surface phosphotungstate

complex would be formed, preventing hydrogen molecule oxidation on the oxide. This is additional, although tenuous, evidence that the surface of tungsten oxide in sulfuric acid

is a weak hydrogen molecule oxidation el ectrocatalyst.

Extensions of the studies on transition metal oxides have been those studies on transition metal carbides, which are also stable in acid. Several international groups

have made contributions to the work, principally Bohm et al . [13], Mund et al . [14], Svata

et al. [15], Palanker et al. [16], and Baresel et al. [17]. A screening program for the

ability of transition metal carbides to oxidize hydrogen in sulfuric acid was performed by

Mund et al . [14] and is given in table 1.

Table 1. Comparison of electrocatalytic activity among

transition metal carbides.

Transition Metal H2 oxidation 22 °C carbide exchange current, io, A/cm^

TiC 8 X 10-8

ZrC 3 X 10-8

VC 2 X 10-7

^""3^2 8 X 10-^

M02C 4 X 10

^

WC 1 X 10-5

Pt 3 X 10-2

The exchange currents on tungsten carbide and molybdenum carbide are clearly superior

1 but are still three orders of magnitude lower than platinum [12], which is included in

Ij, table 1. A general consensus of this work on tungsten carbide el ectrocatalysts is that

for the best hydrogen oxidation activity, the surface of the tungsten carbide must be

slightly oxidized. This may be carried out either by gently heat treating in oxygen or by

electrochemical oxidation in the presence of hydrazine. Whichever method is used, it serves to remove excess surface carbon and is also believed to render the tungsten carbide

surface carbon deficient. Ross et al . [18] examined a variety of tungsten carbides and

compared the electrocatalytic activity to the surface structure probed by Auger electron

j'l spectroscopy. In addition, gas-phase volumetric adsorption of hydrogen was measured. The

1 results are reproduced in table 2.

171
Table 2. Correlation of gas-phase and electrochemical H2 adsorption with

the el ectrocatalytic activity of WC for H2 oxidation in acid.

Preparation Volumetric

adsorption/

pmole BET m"^

Capacitance

(0-0.4 V)/

pF cm"2

Current density/yA cm

at 0.2 V, 22 °C

1 M H2SO4

"2 Reference

AEG-Telefunken 7 475 2 ^ W2C 2.5 330 1.9 1 Ventron 0 62 0.2 1

) Ross et al Starck 0 < 50 0.4 1 [18]

nu n 9 WO3 + CO (600 °C) 0.5 << 50 0.09 ; W + CO (1100 °C) N/A 1100 (geo.) 7 (geo.) Sokol sky

et al . [16]

W + C (2500 °C) N/A 900 (geo.

)

8 (geo.) Mund et al

.

300 (real) 2.7 (real) [14]

It is clear that there is a correlation between hydrogen uptake and the specific

electrocatalytic rate for hydrogen oxidation, so there are special surface features required

to make tungsten carbide an active electrocatalyst. The surfaces of tungsten carbides,

depending on the mode of preparation and pretreatment, may range from having a large excess of metal atoms--in which case, irreversible surface oxide may be formed to act as a poison; to having a large excess of carbon, which is similarly inert for hydrogen oxidation.

At some intermediate point, the surface is slightly carbon deficient, active, and does not form an irreversible oxide. The exact surface structure is not known at this time.

By comparison with platinum, the best tungsten carbide electrocatalyst reported is still significantly inferior, as is shown in figure 6. This figure was constructed using

Fig. 6. Specific oxidation rate for hydrogen

molecules in acid at 0.1 V vs. rhe. Data extracted from references [12]

and [16].

172
rublished values of the kinetic rate constants for hydrogen oxidation of platinum in acid

^2] and for the oxidation of hydrogen on tungsten carbide [16].

One important operational advantage of tungsten carbide over platinum is that it is joisoned by neither carbon monoxide nor hydrogen sulfide. In this regard, the tungsten

carbide el ectrocatalysts when operating with contaminated hydrogen are superior to platinum

at low temperatures. Even with contaminated hydrogen, platinum is still superior at the

higher temperatures, since the adsorption isotherm coverages for carbon monoxide and

sulfur on platinum decrease and the higher reactivity of the platinum comes into force.

In conclusion, it is fair to state that although a start has been made towards our understanding of the factors controlling the operation of non-noble metal and nonmetallic

electrocatalysts, we are still far behind the comprehension that exists for gas-phase

catalysts. On the other hand, the field of electrocatalysis may well be developing faster

since we were further behind. There is no doubt that transition metal compounds, due to their stability, will continue to play an important role in this field. As the diagnostic

tools for probing the el ectrocatalyst surfaces are developed, together with the materials

developments for new electrocatalysts, then we can expect rapid progress towards lower

cost and more highly reactive electrocatalysts that are poison tolerant and decay resistant.

References

[I] Stonehart, P., Bett, J. A. S., and Kinoshita, K. , in Proceedings of Symposiim on EleatToaatalysis, San Francisco, M. W. Breiter, ed. p. 275 (The Electrochemical

Society, Princeton, New Jersey, 1974).

[2] Ross, P. N. and Stonehart, P., J. Res. I-nst. for Catalysis, Hokkaido University,

22, 22 (1974); Stonehart, P. and Ross, P. N., Catalysis Reviews, 12, 1 (1975).

[3] Ziering, L. K. , U. S. Patent No. 3,357,863; Hobbs, B. S. and Tseung, C. C,

J. Electroohem. Soa. 120 , 766 (1973).

[4] Niedrach, L. W. and Zeliger, H. I., J. Electroohem. Soc. ri_6, 152 (1969).

[5] Bagotsky, V. S. and Jofa, S. A., C. R. (Doklady) Acad. Sci. U.S.S.R. 53, 439 (1946).

[6] Glemser, 0. and Naumann, C, Z. Anorg. Chem. 265, 288 (1951 ).

[7] Stonehart, P., Koren, J. G., and Brinen, J. S., Anal. Chim. Acta, 40, 65 (1968).

[8] McHardy, J. and Bockris, J. O'M., J. Electrochem. Soc. 120, 53 (1973).

[9] Randin, J. P., Vijh, A. K. , and Chughtai , A. B., J. Electrochem. Soc. 120 , 1174

(1973).

'10] McHardy, J. and Stonehart, P., in Internatio-nal Reviews of Science, Phys. Chem.

Series 2, A. D. Buckingham and J. O'M. Bockris, eds.. Volume 6, p. 171 (Butterworths, London-Boston, 1976).

[II] Hobbs, B. S. and Tseung, A. C. C, J. Electrochem. Soc. 122 , 11 74 (1975).

[12] Vogel , W., Lundquist, J., Ross, P., and Stonehart, P., Electrochim. Acta, 20, 79

(1975). ~

n3] Bohm, H. Electrochim. Acta, 1_5, 1273 (1970).

[14] Mund, K. , Richter, G., and von Sturm, F., Coll. Czech. Chem. Communs. 36, 439 (1971).

[15] Svata, M. and Zabransky, Z., Coll. Czech. Chem. Communs., 39, 1015 (1974).

[16] Palanker, V. Sh., Baybatyrov, E. N., and Sokolsky, D. V., Electrochim. Acta, 20,

51 , 71 (1975). ~

173
Baresel , D., Gellert, W., Heidemeyer, J., and Sharner, P., Angew. Chem. 83, 213

(1971); U.S. Patent No. 3,902,917.

Ross, P. N., MacDonald, J., and Stonehart, P., J. Eleatroanal. Chem. 63, 450 (1975);

Ross, P. N. and Stonehart, P., J. Catal. 39, 298 (1975).

174
.ational Bureau of Standards Special Publication 455, Electrocatalysis on Non-Metallic Surfaces,

Proceedings of Workshop held at NBS, Gaithersburg , Md. , December 9-12, 1975. (Issued November 1976)

Adsorption and Oxidation of Bromide Ion on Tin Oxide Electrodes

H. Yoneyama and H. A. Laitinen

Department of Chemistry

University of Florida

Gainesville, Florida 326ll

Adsorption of bromide ion on a tin oxide electrode proceeded at such a slow rate that adsorption equilibrixam took longer than two days

to he reached. The electrode became active for oxidation cf bromide

ion depending on the extent of soaking in an acidic bromide solution.

It was found that hydration of the electrode is an important factor

to determine its activity for electro-oxidation of bromide ion. An

electrode having a high activity lost its activity owing to surface

oxidation upon polarization into the discharge potential region of

bromide ion.

Key words: Adsorption; anodic oxidation; electrocatalysis;

semiconductor electrode; tin oxide electrode; transparent

electrode.

1. Introduction

The tin oxide electrode has been used as a transparent electrode to enable speci^rochemical measurements to be made in combination with electrochemical measurements [l] . In addition to its optical transparency, this electrode is superior to noble metal electrodes

in that it can be used to a high anodic potential with a very low background current [2,3].

Redox reactions of several species have been investigated and it has already been clarified that a mechanism of electron tunneling through the space charge layer formed within the

electrode near its surface prevails in electrochemical reactions of species having a more

noble redox potential than the flat-band potential of the electrode, except for at least the

oxygen evolution reaction [U,5,6].

The position of the Fermi level of redox species of one-electron transfer systems relative to electron levels in a tin oxide electrode should be the same when the redox species

have the same redox potential. Inasmuch, as the width of gaussian energy distribution of a

species is larger than 0.5 V [7], no enormous difference in electrochemical behavior is to

be expected for species having redox potentials different by values of the order of O.OIV.

Comparison of the current-potential curves of iron (ll) phenanthroline and hydrobromic acid,

formal redox potentials of which are different by 0.05V, however, revealed an exception to

the above expectation. As shown previously [h] , the iron-phenanthroline redox couple shows

highly reversible behavior both at a highly doped tin oxide electrode and at a platinum

electrode, while the bromine/bromide couple shows highly irreversible behavior at a tin

oxide electrode. It was already speciilated in the previous paper that the abnormal behavior

of bromide may be brought about by some interaction between the surface and bromide ion.

This study was conducted to elucidate the interaction effect.

Figures in brackets indicate the literature references at the end of this paper.

175
2. Experimental

Tin oxide electrodes were prepared in our laboratory by a method similar to that already

reported [8]. An aqueous spraying solution was prepared from stannic chloride and hydrochloric acid containing the desired amount of antimony trichloride. The solution was repeatedly aspirated onto a quartz or glass plate heatexl at around 500°C.

After mechanical polishing with 0.051J alumina, the electrode was washed in a stream of

water by rubbing gently with paper towels. It was then dipped in isopropanol for one hour,

washed with de-ionized water for 30 minutes, soaked in nitric acid for 30 minutes, and finally washed with de-ionized water for 30 minutes. The chemical pretreatment was followed

by drying the electrode at 130°C for over 15 hours. Just before using the electrode, it was polarized twice by applying a triangular potential sweep at a sweep rate of 5 mV/sec between

0 . 2V/ ( vs .SCE) and a potential sufficient to give an anodic current of lOyA/cm^. The first

potential scan was aimed to remove impurities from the electrode surface and the second one

to verify the cleanness of the electrode. The design of the cell used was similar to that

reported earlier [3]. Electrical contact was made with a silver ring.

In order to monitor the amount of adsorbed bromide ion on the tin oxide electrode, con- ventional Y-counting was conducted using a sodium iodide scintillation crystal (Dumont 3292)

as a detector. The electrode was washed with several portions of de-ionized water for 30 seconds followed by two 30 second washing intervals with flowing water in the mounted cell.

The electrode was then de-mounted and washed for 30 seconds in each of three beakers of deionized water. The washing time (3 minutes total) was selected so as to achieve reproducibility with minimal loss due to desorption. A semilogarithmic plot of washing time against

apparent amount of bromide suggested a first order desojption rate-limited process. The rate constant determined for the process was 0.086 min . The amount of adsorption was determined from the counting results at 3 minutes, correcting to zero time by extrapolation

using the experimental desorption rate constant.

Current-potential curves were obtained by a triangular potential sweep method at a sweep

rate of 5niV/sec or lOmV/sec . The nominal carrier concentration of the electrode was determined by using Schottky-Mott plots of differential capacitances of the electrodes measured

at 1 KHz.

3. Resiilts

3.1 Radiochemical studies

Figure 1 shows the number of adsorbed bromide ions as a function of soaking time when an electrode (N=1.5 x lO^'^ carriers/cm3 ) was soaked in hydrobromic acid solutions. Notice that

the time scale is changed at a middle point of the abscissa. The figure shows firstly, that

adsorption progressed very slowly and secondly, that the number of adsorbed bromide ions was dependent on the concentration of bromide ion in the solution.

Figure 2 shows the number of adsorbed bromide ions as a function of polarization potential just after the electrode was subjected to the pretreatment. The electrode was held at each polarization potential for 15 minutes. Preliminary resiilts showed that 15 minutes of

polarization was enough to give a fixed amount of adsorption. The figure shows that potential of maximum adsorption was different for two electrodes having different carrier concentrations .

Figure 3 shows current-potential curves of tin oxide electrodes in one molar hydrobromic

acid. The potential necessary for discharge of bromide ion was different among the electrodes of different carrier concentrations on account of different tunneling probabilities.

It is found by comparing this figure with figure 2 that the potential of the m-aximum adsorption corresponds to the potential at the leading edge of the anodic curve, and that a negligible amount of adsorbed bromide can be expected under conditions of steady state discharge

of bromide ion.

176
r 5

-

i ; 0

-o

s_o

.JCr-

/

/ \ /

Fig.

E 0 1 2 W

Soaking time, h 30 50

1. Numbers of adsorbed bromide ion

as a function of soaking time. N = 1.5 X 1020 carriers/cm3

.

0 = IM HBr; « = 0.05M HBr;

• = 0.2M HBr.

0.3 04 05 0.6 Potential, V/SCE

0.7 0.8

Numbers of adsorbed bromide ion

as a function of polarization

potential. Polarization: 15 min

in IM HBr; 0, N = 5.4 x 1020

carriers/cm3

;

carriers/cm3

f, N = 1.5 10 20

0.3 0.5 0.7 0 9

Potential, V/SCE

Current-potential curves of tin

oxide electrodes having different

carrier concentrations.

Polarization: 5mV/sec in IM HBr; A:N = 8.2 X 10l8 carriers/7m3

;

B:N = 1.5 X 1020 carriers/cm3;

C:N = 5.4 X 1020 carriers/ cm^.

177
3.2 Electrochemical studies

_2 When a tin oxide electrode was soaked in IM HClOj^ containing 10 M HBr , the current-potential curves were gradually changed with soaking time. A typical result is shown in figure h. The potential at the leading edge of the anodic curve shifted in the negative direction and simultaneously a reduction peak of bromine became clear.

Figure 5 shows a comparison of current-potential curves between electrodes soaked under

different conditions. The electrode giving the solid line was soaked for 5 hours in IM HClOl^ containing 10~^M HBr, while the electrode giving the dash-dotted curve was soaked in the

same solution for a shorter time of 1.5 hours. However, only the latter electrode was soaked in IM HClOii for lUO hours before soaking in the acidic bromide solution.

CMe

u

lAC

0)

3 0

-

0 0.4 0.8 Potential, V/SCE

1.2

3.

0.4 0.8 Potential , V/SCE

1.2

Fig. 4. Effect of soaking time on currentpotential curves of a tin oxide electrode in 10~2m JBr/lN HCIO4 . N = 5.4

X 1020 carriers/cm3 ; sweep rate 5mV/

sec 0.25h; 19 h; 40h;

125 h.

Fig. 5. Effect of soaking in IN HCIO^ on current-potential curves of a tin

oxide electrode in 10-2m HBr/lN HCIO4.

N = 5.4 X 1020 carriers/cm3. sweep

rate 10 mV/sec soaked in IN HCIO4 for 140 h and then in 10~2m HBr/

IN HCIO4 for 1.5 h; soaked in

10-2m HBr/lN CIO4 for 5 h.

From the results in this figure, therefore, it is recognizable that soaking in an acid

solution, not necessarily in a bromide solution, is important for the electrode to increase

activity for oxidation of bromide ion.

Figure 6 shows current-potential curves of a tin oxide electrode in IM HBr when the

electrode was soaked in the electrolyte for 2h hours. The potential sweep was continuously

repeated. The first and third potential sweep started at OV, while the second at 0.7V. The

first potential sweep gave a sharp anodic peak, the area of which roughly corresponded to

178
seven times the numter of Sn-0 sites on the surface if the roughness factor of the electrode

surface was assumed to he two. Therefore, we cannot conclude that the peak corresponds

only to oxidation of adsorbed bromide ions. With repetitive polarization cycles, however,

I the peak was successively depressed, and finally completely disappeared. Once the electrode

1 had been brought into the stage, the anodic peak was never recovered by soaking again in the

solution for the same period as before. However, a tendency to recover the activity for

oxidation of bromide ion was noticeable by soaking the electrode for a period of the order

of a week or longer.

U. Discussion

Judging from a recent study by the transmission infra-red technique on surfaces of tin

:xide disks [9], prepared by pressing fine powders of hydrolysis products of tin chloride,

the surface of a tin oxide electrode of the present study might have many different consti-

! tuents such as coordinatively unsaturated tin site, hydroxyl groups, chlorine and coordina-

, tively saturated tin sites, when the electrode was prepared at around 500°C. Relative numj bers of the various sites would be dependent upon the exact conditions of preparation, and

modified by the pretreatment of the electrode although not in a known manner. The exact physical as well as chemical conditions of the pretreated electrode were, therefore, unfortunately unclear, though the electrode surface was judged to be electrochemically clean.

Oxide electrodes usually show dependency of their flat-band potentials on the pH values

of electrolytes. This is connected to dissociation of acidic surface hydroxyl groups [5,10].

The observation that the electrode pretreated in the manner mentioned in the experimental

section showed a similar dependency, as had already been reported [3,^], suggests that some portion of the electrode surface was already occupied by hydroxyl groups, just when the

pretreated electrode was soaked in a solution. Nevertheless, the present study can be sucoessfxilly interpreted by taking into consideration of promotion of hydration with soaking

time.

When a tin oxide electrode is soaked in an acid solution, the cyclic voltammogram of the

electrode gradually changes with soaking time. The main changes are in the back ground

current and anodic potential limit of the electrode. These changes are characteristic of

hydrated electrodes [ll] and were found to be removed by repetitive triangular potential

179
sweeps made in a manner used in the pretreatment of the electrode. Therefore, the hydration

structure seems not to be stable at a high anodic potential. Murray et . al . reported that

treatment by hot concentrated HCl or O.IM NaOH was necessary to impose immediate hydrolysis

on the electrode [12]. These findings in addition to the present results support the idea

that a tin oxide surface becomes hydrated at a slow rate.

It is clear from the result in figure 5 that the first important step to improve the

activity for oxidation of bromide ion is the hydration step with formation of -Sn-OH sites.

The results shown in figure 1 reflects the fact that the rate of hydration of the electrode

surface was low. Possible soaking time of the order of days is required in IM HClOj^ to com- plete this step.

The second step is adsorption of bromide ion on hydrated sites of the electrodes. In this process, adsorption must take place in such a manner to fulfill an equilibriinn relation

between the hydration sites and bromide ions in the solution, which is suggested in eq.(l),

-Sn-OH + h"*" + Br" ^ -Sn-Br+ H^O (l)

It will not take much time to complete this step which is analogous to complexation of aqueous tin (IV) ions by bromide.

The third step is oxidation of adsorbed bromide ions. Taking into account the fact that

the oxidation wave at the first potential seen in figure 6 was higher by far than that

expected for a monolayer coverage, a catalytic process possibly takes part in the oxidation

process

.

-Sn-Br -> -Sn^"^^ + 1/2 Br^ + e~ _ (2)

( + ) - Sn + Br Sn-Br (3)

From the point of view of semiconductor electrodes , appearance of the oxidation wave at

a highly hydrated electrode suggested the formation of surface states, promoted by soaking.

Surface states associated with the oxidation of bromide ion are -Sn-Br sites originating

from hydrated sites.

Loss of the activity for oxidation of bromide ion by repetitive polarization run was seemingly brought about by oxidation of previously hydrated sites of the surface.

-Sn^"^^ + H^O -SnO + 2h"^ + e~ (^)

This view is supported by the result in figure 2 that a negligible amount of bromide ions is

retained on the electrode surface when it is polarized under conditions of steady state discharge of bromide ions.

5. Conclusion

The present study revealed a case in which hydration of the electrode plays an important

role in affecting the kinetics of charge transfer between electrode and solution species.

Inasmuch as oxides are inevitably subjected to surface hydration though the rate and degree

of hydration will vary for different oxides and conditions studied, careful consideration

of the surface condition will be necessary when electrokinetic studies of adsorbable species

are conducted at oxide electrodes

.

180
6 . Acknowledgement

We would like to thank the National Science Foundation for support of this program and

so Professor M. L. Muga for kind advice in the radiochemical studies.

7. References

Kuwana, T., Ber. Bunsenges. Phys . Chem. , T]_, 858 (1973),

] Strojek, J. W. and Kuwana, T. , J. Electroanal. Chem., 16, ^+71 (1968).

] Laitinen, H. A., Vincent, C. A. and Bednarski, T. M. , J. Electrochem, Soc. , 113 , 102^; (1968).

] Elliott, D., Zellmer, D. L. and Laitinen, H. A., J. Electrochem, Soc, 117, 13^3 (1970)

] M8llers , F. and Memming, R., Ber, Bunsinges. Phys. Chem., 76^, h69 (1972T'!

] Memming, R. and MBllers, F. , ibid., 76, U75 (1972).

] Vanden Berghe, R. A. L. and Gomes, W. P., ihid. , 76, ii8l (1972).

] Kim,H. and Laitinen, H. A., J. Amer. Ceramic Soc, 58_, 23 (1975).

] Thornton, E. W. and Harrison, P. G, , J. Chem. Soc. Farady Trans. I, 71.,^6l (1975)-

0] Lohmann, F. , Ber, Bunsenges. Phys. Chem,, 70_, ^28 (1966),

] Kin, H. and Laitinen, H. A., J. Electrochem, Soc, 122, 5^+ (1975).

2] Moses, P, R,, Wier, L. and Murray, R. W, , Anal, Chem,, hj_, l882 (1975).

181

National Bureau of Standards Special Publication 455, Electrocatalysis on Non-Metallic Surfaces,

Proceedings of Workshop held at MBS, Gaithersburg , Md., December 9-12, 1975. (Issued November 1976)

Cathodic Pyrophosphate-Lead Dioxide

Electrode Reaction Properties

C. 0. Huber, T. A. Berger, and R. E. Reim

Department of Chemistry and Laboratory for Surface Studies

University of Wisconsin-Milwaukee

Milwaukee, Wisconsin 53201

In the vicinity of +0.4 volts vs_. S.C.E. in pH 8 solution pyrophosphate yields a cathodic current selective over orthophosphate at the lead dioxide electrode.

Several aspects of this response were examined using

a-el ectropl ated lead dioxide on a platinum base and

using mixed a-3 wax-bound lead dioxide electrodes prepared from the powdered material. Scanning electron

microscope results indicated the topographical nature

of the wax-bound electrodes. Cur ren t- vol tage plots

indicated an absence of a limiting current and relatively high residual currents as the potential becomes

more cathodic. Scanning cyclic voltammetry suggested

adsorption of reactive species. Capacitance and coulo me trie studies indicated clear evidence of adsorption and showed that the current was necessary to the

desorption step. Current vs_. concentration studies

allowed fits to el ectrosorpti on models within certain

concentration ranges. From these data estimates of some thermodynamic quantities were made. An electrode

reaction process involving quasi equilibrium adsorption

of the electron transfer product followed by desorption

of a pi umbous-pyrophosphate complex is discussed.

Key words: Amperometry; cyclic voltammetry; electrosorption; Langmuir; lead dioxide; pyrophosphate;

Temkin; wax-bound electrode.

1 . Introducti on

The conduction band of certain n-type metal oxide semiconductors involves the outer s orbital of the metal . Electron mobilities and electron

densities are several magnitudes larger than for the usual transition metal

oxides. Examples of this group are ZnO, SnOa and PbOz.

Lead dioxide exists in a and 6 crystalline forms [1].^ The structures

have been described by Thomas and are closely related with octahedral packing. X-Ray analysis has been used to distinguish the a and g polymorphs

[2,3] and to estimate individual concentrations in a mixture.

Resistivities of a and 6 lead dioxide at 25 °C are reported to be 1 x 10"5 ohm-m and 4 x iQ-s ohm-m, respectively. The temperature coefficient of

Figures in brackets indicate the literature references at the end of this paper.

183
resistivity is positive and relatively large, +0.07%/°C. Hall effect studies

have been reported to indicate concentrations of 10^^ electrons per cm^ [4].

Lead dioxide has been reported with s to i ch i ome t r i es from Pb0i_66 to PbOi.ge with the latter value apparently more representative. Evidence to support the non-stoi chi ometry as due to interstitial excess lead atoms has

been reported [4].

When lead dioxide is used as an electrode a cathodic current is obtained

in the presence of pyrophosphate in the vicinity of +0.4 v \j_s_. S.C.E.

Orthophos phate does not show a corresponding current. Some aspects of the

electrode reaction are presented and discussed here.

2. Electrode Preparation

In the work reported here, electroplated a lead dioxide was deposited

onto platinum from an alkaline lead acetate bath at 0.04 A/m^ [5]. All

plating solutions were stirred. At the pH values used in the experiments

reported here, the 3 crystal form would not be expected to persist [4], and

therefore was not used. Wax- or plastic-bound lead dioxide electrodes were prepared by simply mixing 855^ by weight of reagent grade powdered PbOz (as

purchased) with the liquid form of- the binding material. In the case of

paraffin wax, the partially-hardened slurry was pressed to form cylinders

6.4 mm or 3.2 mm in diameter and 2-3 cm in length.

In the pressing step a connecting metal terminal was incorporated into

one end of the electrode body. The other end was used as the electrode face.

Sides were covered with wax. Completed electrodes had resistances of 1000

ohms or less. Molded electrodes using polyester resin were also prepared.

No significant distinctions in the behavior based on the binder material

could be observed. Surface renewal of these electrodes was easily accom- plished vi

a

slicing or abrasion.

Scanning electron microscopy was used to examine the topography of the

surface face of the wax-bound electrodes. The scans obtained and the

associated data show that PbOa particles protrude from the surface (fig. 1),

that the surface of PbOz particles are porous and intricate (fig. 2), and

Fig. 1. Compositional scan of pressed Fig. 2. Compositional scan of a wax-bound lead dioxide elec- typical lead dioxide particle,

trode surface. Superimposed X 10,000.

line scan of total x-ray intensity. X 1000.

184
Fig. 3. Compositional scan of a polished wax-bound lead

dioxide electrode surface

with Pb x-ray line scan.

that attempts at polishing the surface generate a wax film which tends to

obscure the lead dioxide from exposure to the solution (fig. 3). These SEM

data and optical microscope studies showed that particle sizes remain un- changed during fabrication of the electrodes.

3. Results and Discussion

3.1 Current-Voltage Plots

Manually obtained current vs^. voltage plots show that near 0.4 v vs . S.C.E. at pH 8 the residual current is very low. At more cathodic voltages

other electrolyte constituents, and also orthophosphate when present, con- tribute to the observed current. A limiting current is not observed. The

Fig. 4. Scanning voltammetry curves. 3.84 volts/min. After 50

cycles and: A, allowed to stand at open circuit for one minute; b, C, and D, cycled successively.

185
pyrophosphate current diminishes as pH increases above 8 whereas in the

vicinity of pH 7 the selectivity to pyrophosphate over orthophosphate

diminishes, thus pH 8.0 was selected for these studies.

Scanning voltammetry experiments with a electrodes, illustrated in

figure 4, show that the electrode reaction is irreversible. A pre-peak

attributable to adsorption of product appears on the second and successive

scans. Both the parent peak and the post-peak (attributed to adsorbed

reactant) decrease somewhat on successive scans. As scan rate is increased

the height of the post-peak relative to the parent peak is observed to increase markedly, while that of the pre-peak decreases. These effects cor- respond to change in amounts of adsorbed electron transfer step reactant and

product respectively.

3.2. Capaci tance

The circuit described by McMullen and Hackerman [6] was modified

slightly to provide for automatic applied D.C. potential control. Solutions

were made from triple distilled, organic free water and reagent grade chemicals. Capacity measurements of a lead dioxide in ammonium nitrate buffer

120 -

(J

4.

100

Fig. 5. Differential double layer ca- pacity measurements on a-Pb02

electrodes: A, in O.IN^ NH4NO3

buffer of pH 8; B, same as A plus 1 .1 X lO-M P2O7'*; C, same

as A plus 3.1xlO-2M PaOy"*; D, reported for 0.18M KNO3 [7],

i. + 0.5 +1.00

VOLTS vs SCE

of pH 8.0 are presented in figure 5. Curves B and C show distinct enhancement of double layer capacity due to pyrophosphate. Comparisons of capacity

measurements in buffer at a mercury electrode at -0.25 to -0.75 v v_s_. S.C.E.

lead to a roughness factor for lead dioxide of 3.4, a value in good agreement with previous workers [7,8].

3.3. Coulometry

In these experiments the electrode was allowed to attain a stable

residual current in pyrophos ph ate-f ree buffer. The electrode was then

immersed in a buffered pyrophosphate solution and coulombs measured for a prescribed time, typically 4.0 minutes, with either open or closed (0.35 v

186
vs . S.C.E.) circuit. The electrode was then removed from solution, rinsed

with water, and placed in a new buffer containing no pyrophosphate at closed

circuit. A decaying current-time ( i.e . , stripping) curve was obtained

indicating depletion of the surface pyrophosphate. Some results are summarized in figure 6 in which a saturation effect appears to be indicated.

Fig. 6. Coulombs for adsorption and

desorption of pyrophosphate

(O) > Adsorption step; () , Desorption step.

1.0 2.0 PYROPHOSPHATE CONC M X lO'3.0

When the new buffer step was begun at open circuit for up to 4 minutes, the

number of desorbed coulombs were within 5% of the number obtained with the

circuit closed immediately. For the a electroplated electrode the time at

open circuit could be as much as 30 minutes with no significant decrease

in desorption coulombs. These results suggest that non-el ectroacti ve desorption is small, but that possibly the wax-bound electrodes are attacked

more readily by solvent and buffer species than are the electroplated

electrodes. The a-electroplated electrodes showed adsorption saturation

effects in shorter times and with lower numbers of coulombs than the waxbound version.

3.4. Amperometry

Potenti ostati c current vs_. concentration measurements allowed the best

quantitative data because a given electrode surface could be used without

disruption or disturbing while successive concentration increments were

introduced and because competing effects of other anions could be minimized

by maintaining a potential at which residual currents were very small.

With the wax-bound electrode a linear 1/i v^. 1/C relationship was

found for lower pyrophosphate concentrations observed (0.4 to 3 x 10"'* hi).

The electroplated a-form did not yield such linearity. At somewhat higher

concentrations both electrodes yielded linear i vs_. log concentration data

over at least a five- fold concentration range. The nominal (no roughness

correction) current densities were in the range of 0.10 and 0.005 A/m^ for

the cc-plated and the wax-bound electrodes, respectively.

A quasi equilibrium mechanism can be proposed

187
lattice Pb02 + HPzOy"^ lattice PbOz'PbPzOy + 2e + H (1)

k

- 1

k2 lattice PbOz-PbPzOy > lattice PbOz + PbPzOy (2)

where V.z << ki or k-i and i = "i 2" = k26. It can be shown [9] that applying a Langmuir model to this chemi sorpti on process results in

2F

1 _ ^ , J

i k2 KCk2^

where K = ki/k-i and C is pyrophosphate solution concentration. For small

values of n the linear reciprocal current v_s_. reciprocal concentration observed was used to estimate the apparent adsorption equilibrium constant

(Table I )

.

Table I

Thermodynamics Estimations

(AG° , AH° in kcal /mol

)

El ectrode Model Temp K AG° AH° AS°

Wax-bound Langmuir 30°C 4 ± 1 X 10^ 5

Temk i n 2.9° 0,,85 X 10^ -3 . 7

Temki n 23° 2.,6 X 10^ -4 . 6

Temki n 35° 5,,8 X 103 -5 . 3

9.9 42 e . u.

a-pl ated Temki n ^ 20° 0.,62 X 10^ -3 . 6

30° 1 ,,3 X 10= -4 . 3 40° 2 ,.0 X 10^ -4 . 7

10.5 49 e. u.

CI ausi us-Clapeyron 10.4

The Temkin isotherm adsorption model based on free energy of adsorbate

deposition linearly dependent on surface coverage, 6, yields:

^ ln(K°C) + ^ n

188
where K° is the apparent standard adsorption equilibrium constant and r is an empirical parameter [9]. The relationship applies to intermediate coverages.

The linear i vs_- concentration properties found for the system examined

here correspond to the assumption that i = kzQ, i.e ., that the desorption

step energetics are relatively independent of coverage. This situation

correlates to the product containing electrode material. The corresponding

values obtained are summarized in Table I. Each of the evaluations was performed at three temperatures and with these data AS and AH values were estimated.

The heat of adsorption at the a-plated electrode was also computed using

at constant apparent coverage (see Table I).

The free energy of formation of Pb(P207)~ in water solution at 25° is

8.6 kcal/mol based upon stability constant values reported by Costly and

Farr [10]. The data here suggest that the surface bond may resemble that of

the soluble complex with the pyrophosphate ligand satisfying two of the

octahedral coordination sites.

He i land, Gileadi, and Bockris [11] have examined the adsorption of

benzene on platinum and report an entropy change of 50 entropy units. They

suggest that displacement of solvent molecules from surface sites might

account for the relatively large positive entropy shift for larger molecules.

Similar reasoning may be applied to the case of the relatively large pyrophosphate ion.

The similarlity of results with both the a-el ectropl ated and the waxbound a-S mixture electrodes used here suggests that these surface reactions

are not specific for a particular crystal structure. Qualitative results

in this laboratory using 6-el ectropl ated lead dioxide electrodes support

this view.

[1] Thomas, U. B., Trans. Electrochem. Soc, 94 , 42 ( 1948).

[2] Brar, A. S. and Dupo, R. F., J. Electrochem. Soc, V\6_, 1 1 6 ( 1 969).

[3] Kordes, D., Chem. Ing. Tech., 38, 638 (1966).

[4] Mindt, W., J. Electrochem. Soc, 1_1_6 , 1 076 ( 1 969 ).

[5] Ruetschi, P., Angstadt, R., and Cahan, B., ibid . , 1 06 547 ( 1 959 ).

[6] McMullen, J. J. and Hackerman, N., ibid . , 106 , 341"Tr959 ).

[7] Carr, J. P., Hampson, N. A., and Taylor, R., J. Electroanal. Chem., 2 7 , 109 (1970).

[8] Hampson, N. A., Jones, P. C, and Phillips, R. F., Can. J. Chem., 45^,

2039 (1967).

[9] Gileadi, E., Ki rowa-Ei sner , E., and Penciner, J., Interfacial Electrochemistry, Addi son-Wesl ey Publ . Co., Inc. ( 1 975 ). [10] Costly, B. D. and Farr, J. P. G., Chemistry and Industry (London), 42,

1435 (1968).

Jl] Heiland, W., Gileadi, E., and Bockris, J. O'M., J. Phys. Chem., 7^,

1207 (1966).

AH° T2-T1

R T1T2

4. Ref e rences

189

::ational Bureau of Standards Special Publication 455, Electrocatalysis on Non-Metallic Surfaces,

Proceedings of Workshop held at NBS, Gaithersburg , Md. , December 9-12, 1975. (Issued November 1976)

EEECTROIYSIS BY FOREIGN METAl MONOLAYERS

R.R.Adzic, A.R.Despic, D.N.Simic and D.M.Drazic

Institute of Eleetrochemistry , ICITM and

Faculty of Technology and Metallurgy,

Belgrade, Yugoslavia

Catalytic effects of foreign metal ^^nol^^ers deposited at underpotentials have "been found in the case of Fe /Fe redox reaction on gold,

oxygen reduction on gold, and oxidation of some small organic molecules on platinum. It was found that the catalytic effects of foreign metal monolayers originate in the change of surface electronic properties and the

ch|j^ge energy of adsorption of oxygen or oxygen-containing species for

Pe /Fe and oxygen reduction respectively. In the case of formic acid

oxidation foreign metal monolayers decrease the hydrogen adsorption on platiniim thus preventing the formation of the main poisoning species.

2+ 3+

Key words: electrocatalysis; foreign metal monolayers; Fe /Fe

reaction; formic acid oxidation; oxygen reduction; underpotential deposition.

1. INTRODUCTION

Most of the work in electrocatalysis has so far "been concerned

with the effect of the nature of electrode material. However, it is quite

conceivable that the state of the electrode siirface can have profo\ind

influence on the rate of the heterogeneous reactions and one of the problems

in the study of electrocatalysis in the past was the lack of detailed

knowledge of the electrode surface layer and its relation to the hulk properties. Over the past few years various ultra-sensitive experimental

techniques of studying siirfaces have been developed that are capable of

detecting minor amounts of material on the surface and providing other information needed to characterize a s\irface. Thus, in situ application of

reflectance spectroscopy [1, 2, 3]1 recently revealed a large change of

siirface electronic properties of an electrode when foreign metal monolayers

are deposited at potentials more positive than the reversible one (underpotential deposition). This is a consequence of a strong interaction of the

valence orbitals of adsorbed atoms with a band structure of the substrate

[1] . There is no general agreement about the nature of such layers but a prevailing view is that these layers are partially charged adsorbed atoms

[1/ 4, 5]. The ir effective charge is presumably quite low [1, 5] . It seemed

that underpotential deposition (TJPD) of foreign metal monolayers is a con- venient way to modify electrocatalytic properties of surfaces and in general to study the surfaces effects in electrocatalysis.

The catalytic effects of foreign metal monolayers have been

established for several reactions. It was ^^cent^y shown that such layers

can substantially increase the rates of Fe /Fe couple on gold [6] and

^Figures in brackets indicate the literature references at the end of this paper.

191
oxygen reduction on gold [7] , and the rate of formic acid oxidation on platinum [8, 9]. The purpose of this paper is to review some of these

results and to report on more recent work on this subject,

2. EXPERIMENTAL

The studies have been carried out in 1 M HCIO solution prepared

from triply distilled water with the first distillation from alkaline permanganate and closed in a pyrex glass electrolytic cell. A polycrystalline

platinum plate to serve as the test electrode was polished to obtain a mirror finish. After polishing the electrode was washed with KOH and HNO^

solutions and stored in triply distilled water. The co\inter electrode was also a platinum electrode. The reference electrode was a saturated sulphate

electrode connected via Luggin capillary.

Nitrogen was bubbled through the electrolyte before meastirements.

It was purified by passing over copper turnings and then through molecular

sieves. All measurements were carried out at room temperature.

Potential sweep and steady-state polarisation measurements were

made first in the absence and then in the presence of known amounts of a cation in solution, whose catalytic activity was ^o be tested. The cation

concentration was varied in the range between 10 and 10-2

The standard circuitry included a Tacussel potent iostat, PAR

Universal programmer and a Hewlett-Packard XY recorder. Potential sweeps

have been made at a rate of 50 mV/sec. Steady-state data have been obtained

by holding the electrode at a constant potential and following the decrease

in cvirrent with time. The currents were read on a Data Precision digital

multimeter after 90 sec.

3. RESULTS AND DISCUSSIONS

2+ 3+

3.1. Ee /Pe redox reaction on gold

2+ 3+ The kinetics of Pe /Pe redox couple has been determined on gold

electrode in sulfuric acid solution in the potential range where the underpotential deposition of silver occurs. Subsequently silver was added to the

electrolyte as Ag^SO^, in the concentration of 1.10 ^to 1.10"^ M. An

immediate rise of current followed the addition of Ag which persisted

after sufficient time was allowed for adsorption of silver to attain

equilibrium. A quantitative investigation has been made using potentiostatic techniques and running the redox process at low overpotentials so that

the exchange current density could be determined from the linear part of

the current-overpotential function in the absence and presence of silver

in solution. Pig. 1 shows the effects of silver monolayer on the rate of

Fe2+/Pe5+ couple on gold. A two-fold increase in the exchange current

density is seen in both anodic and cathodic polarizations.

Several phenomena have been considered as possible causes of this

catalytic effect. It was shown that the double layer effect should decrease

rather than increase the reaction rate [6] . The silver monolayer can change

the electronic density-of-states at the surface which may be a possible

explanation of this effect. Although the analogy is somewhat distant it

appears that the resonance tunneling, observed with metal-adsorbed layer- vacuum interface [10] , should also be considered as a possible explanation.

192
100 200 300 400

i (mA/cm^)

Fig. 1. Steady-state current-potential curves of

gold in 2M H2SO4 containing 0 . IM FeSOit ^^i^

Fe2(S04)3. Curves 1, without Ag+; curves

2, with 1.10" Ag"'' in solution

3.2. Oxygen reduction on gold

Gold is known as a poor catalyst for oxygen reduction. This was

explained by its electronic structure which has no unpaired d-electrons

which could engage in honding with oxygen or oxygen-containing species

[11], or by the fact that gold is a poor peroxide-decomposing catalyst [12].

It was shown that the monolayers of Bi, Pb, Tl formed by UPD have pronounced catalytic effect on oxygen reduction. The monolayer of Cd showed very

small effect while Cu showed inhibitlve effect [7]. Fig. 2 shows the effects

of bismuth monolayer on oxygen reduction on gold in IM HGIO electrolyte.

Fig. 2. Steady-state current-potential

curve for oxygen reduction on bare gold, and gold covered

with a monolayer of bismuth

in IM HClOit .

^ Au \

•-<. \

\ \ \

\

1 11 1 1

\

1

1

I I I I I I I I I I I I

5 2 IQ-^ 5 2 5 2 W'' 5 ; /fl"^ i[Acill'^]

It is seen that the linear Tafel relationship existing in the absence of

Bi5+ turns into an almost horizontal line in the presence of the latter,

up to the current density region in which the diffusion limiting current

is expected.

In acid media significant UPD of Tl, Cd and Pb falls in the potential range where the oxygen reduction becomes controled by diffusion and

193
hydrogen evolution commences. This prevents the estimation of the full

merit of these metal adatoms as catalysts and the measurements at some more

convenient pH are warranted.

The effect of depositing foreign metals on the decomposition of

hydrogen peroxide as the intermediate in oxygen reduction has also been

tested for the case of two metals with opposing effects on the latter, i.e.

for Bi and Cu. Both metals exhibited similar and considerable electrocatalytic effects. Hence, the inhibitive effect of Cu concerns the initial

stage of oxygen reduction, i.e. the formation of hydrogen peroxide.

It was suggested that the rate determining step for O2 reduction

on gold is the exchange of the first electron [13, 14]. The same behavior

is found in this v/ork [7] . If the rate-determining step is

02(ads) + e 0^ (ads) (1)

the reaction is led into the "peroxide path" and the most probable further

step involves the reaction of 0^ with hydronim ion followed by another

electron transfer and formation of H-O^.

If in a peroxide free solution the first electron-transfer is

rate-determining, preceded by one or the other type of adsorption of the

oxygen molecules there are several possibilities of electrocatalysis [7]

.

However, the most probable possibility is the energy barrier for electron

transfer changed favorably by the change in the energy of oxygen adsorption.

In presence of a foreign adsorbed atom the potential at the boundary of

the metal sxirface is perturbed sufficiently as to produce a new surface

state. These surface states can play a role in determining adsorption properties of such electrodes, and lead to a more favorable adsorption of

oxygen on covered gold then on the bare gold surface.

3.3. Oxidation of formic acid on noble metals

It was recently shown that the monolayers of Cd, Tl, Bi and Pb

have striking catalytic effects on oxidation of formic acid on platinum

[9] . The UPD of these cations follows the reduction of platinum oxide.

Considerablly lower currents are obtained in the region of hydrogen adsorption indicating a supression of hydrogen adsorption by monolayers of

these cations [9]

.

The potential sweep diagram in the presence of formic acid,without

the addition of cations whose catalytic activity has been tested, is shown

in Pig. 3 (dashed line). The addition of Pb2+, which undergoes UPD in that

potential range, is shown to have a profound effect, causing a large catalytic effect. The peak current in anodic direction is about two orders of

magnitude larger than in pure formic acid solution. Similar, although

smaller effects were observed by Schwarzer and Vielstich [15] produced by

traces of lead in sulfuric acid solution. Other authors also have observed

the benefitial effects on oxidation of formic acid on platinum of the substances normally regarded as poisons. Breiter [16] and Brummer et al. [17]

have noticed the effects of copper monolayer and Angerstein-Kozlowska et

al, [18] of electrodeposited mercury and acetonitril. The effects however

were small and no systematic study was imdertaken.

Quantitative investigation has been carried out of the steadystate polarization of the platinum electrode in pure formic acid and in

the presence of different concentrations of Pb^"*" ions. The obtained Tafel

relationships are shown in Pig, 4,

194
i (mAcffT^)

Voltammetry curves of platinum electrode

in IM HClOit containing 0,265 M HCOOH in the

absence (dashed line) and in the presence

of I.IO-^M Pb2+ (full line). Sweep speed

5 0 mVs-1

.

E(V)

The curve in pure formic acid exhibits a linear Tafel region with

a slope of ~ 120 mV and other well known features at more anodic potentials of oxide formation.

^ Addition of very small amount of Pb2+ (5.10~ M) results already

in catalysis of HCOOH oxidation. The effects are considerable in the linear

Tafel region but even higher in the region of limiting current density.

The Tafel slope seems to be preserved for that concentration but decreases

somewhat as the concentration of Pb2+ is increased up to 1.10~%. Further

increase of the Pb2+ concentrations increases the Tafel slope back to the

value of 120 mVdec-''.

195
At higher anodic potential the situation is somewhat different.

The largest catalytic effect is obtained at the concentration of 5.10~"'m

With higher concentration there is a decrease of catalytic effect

and even inhibition at more anodic potentials. At high oxide coverages the

effects are small for all lead concentrations. However, only the effects

at most cathodic potentials in these plots are of technological importance.

It was shown [9] that the reaction order with respect to Pb^^ ions appears

to be ~ 0.5. However, the changing Tafel slopes with increasing concentration make the meening of such a reaction order questionable.

The reaction mechanism of oxidation of HCOOH on bare platinum has

been given very detailed consideration [of. 19]. The formation of several

poisoning species makes these studies difficult. It seems that majority of

authors agree that the main poisoning species is^^^H whose formation is

strongly dependent on the amount of hydrogen adsorbed on the electrode.

According Capon and Parsons [20] tha carboxyl ad-radical COOH , formed in

an initial reaction of HCOOH with Pt,

^

HCOOH > COOH + H"^ + e (2)

reacts immediately with H ad to form COH

XXX

XCOOH + 2Had ^ COH + H„0 (3) XXX 2

^

In the main reaction pathway the carboxyl ad-radical undergoes further

oxidation:

COOH —> CO^ + H"*" + e (4) X ^ The monolayers of metals investigated decrease the hydrogen adsorption on platinum to a negligible value and thus they could decrease

the rate of reaction. That this is the case was proved by the following

experiments.

In the absence of foreign metal cations the electrode was held at

the open circuit potential, which is in hydrogen adsorption region, for

5 min. The first sweep in the anodic direction (dashed line) shows considerably smaller current than that obtained under conditions of constant

cycling as shown in Pig. 5. The same experiment perfonaed in the presence

of Pb2+ in solution showed different picture (Pig. 6). The first sweep in

the anodic direction (full line) gives even higher current than the subsequent sweeps. This suggests that no or very litle poison was formed at

open circuit potential in the presence of lead monolayer in contrast to the

situation in the absence of lead. Desorption of lead during anodic sweep

opens a fresh platinum surface which can sustain the observed huge currents.

The difference in the currents associated with the first and subsequent

sweep suggests that at higher anodic potentials there is a foimation of

poisoning species. The reaction giving the poisoning species probably does

not involve adsorbed hydrogen in this case. This poison could be again cOH

formed by the disproportionation reaction, or intermediate ^(OHj ^ or dimer

of formic acid anhydride. The catalytic effects at higher anodic potentials

besides decreased hydrogen adsorption, if any at those potentials, may be

due to the "third body" effect [18] . The adsorbed foreign metal atoms may

block inhibitor adsorption or formation by blocking sites for adsorption.

Prom this considerations it appears likely that the investigated catalytic

adatoms have a dual effect on the oxidation of formic acid.

196
40 - i (mAcm"^)

Voltaitimetry curve of platinum electrode in

IM HClOi, containing 0.265M HCOOH and 1 mM

Pb2+. First sweep after 5 min at open

circuit (full line) , second sweep (dashed

line). Sweep speed 50 rnVs"^.

The experiments with Pd and Rh seem to confirm previous explanation. Palladium is known to form no poisoning species and in this case

lead monolayer inhibits the reaction (jPig. 7). On the other hand, rhodium

behaves similarly to platinum, and in the case of rhodium a large catalytic effects were again observed, as shown in Pig. 8. The effects of lead

on oxidation of HCOOH on Ir v/ere practically negligible.

197
i(mAcm-2)

Pd/Pbo

Fig. 7. Voltammetry curve of palladium

electrode in IM HCIO1+ containing 0.265M HCOOH in the absence

(dashed line) , and in the

presence of 1 mM Pb^"*" (full

line). Sweep speed 50 mVs"-^.

O.L 12 1.6

: (mAcm"2)

2.0

Rh/Pbad

1.0

--1.

/ \

Fig. 8. Voltammetry curve of rhodium

electrode in IM HClOi^ containing 0.265M HCOOH in the absence

(dashed line) , and in the

presence of 1 mM Pb^"^ (full

line). Sweep speed 50 mVs~-^.

E(V)

0.0 OA 12 16

It seems that the minimum polarity model which is used by

some authors to explain catalytic effects of alloys can be applied to

these systems. The model assximes a complete lack of interaction between

neighbouring atoms, the adsorption and catalytic properties being a function of the localized properties of the individual surface atoms. In

such a case one could regard the observed catalytic effects on Pt and Rh

as an indication of their behaviour if no poisoning species were formed.

If this approximate measure of catalytic activity of Pt and Rh

is taken, and that of Pd, Ir and Au obtained without foreign metal monolayers, and correlated with the metal, heats of sublimation, one obtains

the volcano type curve shovm in Pig. 9.

198
Capon and Parsons [19] recently obtained similar curve by plotting

the current at the cathodic peak vs. latent heat of sublimation. In this

case the correlation of the log i was used as the theory requires, and it seems that more quantitative significance can be attached to this result.

4. CONCLUSIONS

Catalytic effects of foreign metal monolayers have been found

for several electrode reactions. Pe^+/Fe5+ couple belongs to the so called

weak interaction case, while the oxygen reduction and formic acid oxidation

are representatives of strong interaction case. This demonstrates the

importance of surface properties of electrodes in electrocatalysis in general, and calls for a surface-oriented study of electrocatalysts.

The UPD of foreign metal monolayers was shown to be a quite convenient way of modifying surface properties of electrodes. By controlling

the electrode potential one controls both the coverage by foreign metal

adatoms and the state of their effective charge. However, this quantity

is not readily obtainable and one should obtain at least the electrosorption valency of such species. Such data could help the interpretation

of catalytic effects and eventually make possible the predictions of catalytic behavior of foreign metal monolayers.

It is likely that the coverages by foreign adatoms are different

in the presence and in the absence of reacting species. Therefore, a determination of the coverages during the catalytic reactions is warranted.

199
REFERENCES

1. R.Adzic, E.Yeager and B.D.Cahan, J.Ele ctrochem.Soc. , 121 (1974) 474.

2. T.Takamura, Y.Sato and K.Takamura, J.Electroanal.Chem. , 41 (1973) 31.

3. D.M.Kolb, D.Leutloff and M.Przasnysky , Surface Sci. , 47 (1975) 622.

4. J.W.Schultze and K.J.Vetter, J.Electroanal.Chem., 44 (1973) 63.

5. D.M.Kolb, M.Przasnysky and H. Gerischer, J.Electroanal.Chem., 54

(1974) 25.

6. R. R.Adzic and A.R.Despic, J.Chem.Phys. , 61 (1974) 3482.

7. R. R.Adzic and A.R.Despic, Ber.Bunsenges.Phys.Chem, , submitted for

publication

.

8. R. R.Adzic, D.N.Simid, D.M.Drazic and A.R.Despic, J.Electroanal.Chem.,

61 (1975) 117.

9. R. R.Adzic, D.N.Sirai6, A.R.Despic and D.M.Drazic, ibid., 65 (1975) 587.

10. J.W.Gadzuk, Phys.Rev. B1 (1970) 2110.

11. A.Damjanovic andV.Brusic, Electrochim. Acta, _12_ (1967) 1971.

12. J.P.Hoare, ibid., _M (1966) 311.

13. G.Bianchi,, F.Mazza and T.Mussini, ibid., _M (1966) 1509.

14. A.I.Krasilshchikov, Zh.Fiz.Khim. , 26 (1952) 216.

15. E.Schwarzer and W.Vielstich,C.R. , 3-ieme Journ.Intemat. d'Etude de

Piles a Combustible, Presses Academique Europeennes,

Bruxelles 1969.

16. M.Breiter, Electrochem.Acta, 8 (1963) 447.

17. S.Brummer and A.C.Makrides, J.Phys.Chem. , 68 (1964) 1448.

18. H.Angerstein-Kozlowska, B.MacDougall and B.E.Conway, J.Electrochem.

Soc, 121 (1974) 756.

19. A. Capon and R. Parsons, J.Electroanal.Chem., 44 (1973) 239.

20. A. Capon and R. Parsons, ibid. 45 (1973) 205.

200
Comment on Paper by R. R. Adzic, et al . , "Electrolysis

by Foreign Metal Monolayers"

B. D. McNicol

Shell Research, Ltd

Thornton Research Centre

P. 0. Box 1 Chester, CHI 35H, England

I have a comment and a question on the paper by Adzic, et al. Shell has been working

on Pt/Sn catalysts for methanol electro-oxidation for some time and the characteristics

of these catalysts are rather similar to those catalysts described by Dr. Adzic. Some

of our Pt/Sn catalysts were prepared by a special immersion technique developed by

M. M. P. Janssen (Koninkl i jke/Shell Laboratorium, Amsterdam), whereby the Pt catalyst is electrochemical ly pretreated prior to immersion in a Sn salt solution. Cyclic voltammetry

studies of such catalysts revealed that the hydrogen chemisorption regions of the profiles

were reduced in intensity while the oxide formation and reduction regions were shifted

significantly. The catalytic activity for methanol electro-oxidation on the Pt/Sn catalyst

is about 100 times greater than that of pure Pt. The effect of the Sn promoter on the Pt

has been interpreted as a ligand effect [1], thus modifying the Pt adsorption properties.

The method of preparation of the Pt/Sn catalysts seemed readily applicable to finely

dispersed Pt such as Pt on carbon. However, with certain Sn salts it was found that the Sn

was preferentially taken up by the carbon support. Therefore, I would like to ask whether

Dr. Adzic has looked at the influence of metal cations in solution on the activity of

supported Pt.

[1] Janssen, M. M. P. and Moolhuysen, J., manuscript submitted to J. Catalysis for

publ ication.

201

ELECTROCHEMICAL PROCESSES ON NON-METALLIC SURFACES - 2

National Bureau of Standards Special Publication 455, Electrocatalysis on Non-Metallic Surfaces,

Proceedings of Workshop held at NBS, Gaithersburg , Md. , December 9-12, 1975. (Issued November 1976)

Mechanisms of Electrochemical Reactions on Non-Metallic Surfaces

Ernest Yeager

Department of Chemistry

Case Western Reserve University

Cleveland, Ohio khlOG

Oxygen electrocatalysis has assTjmed new importance "because of the

involvement of the 0^ electrode in various fuel cells and electrolyzer

systems. For 0^ electrore duction to proceed at reasonahle current

densities requires the adsorption of the 0^ molecule or ion on the

electrode surface. Various models for the interaction of such 0^

species with surfaces and the corresponding pathways for the electroreduction are considered. Generally peroxide mechanisms are predominant in aqueous solutions on non-metallic as well as metal surfaces

hut this leads to less favorable operating potentials for 0^ cathodes.

A mimher of reasonably effective 0^ electroreduction catalysts have

been identified for alkaline solutions but so far only high area platinum appears to combine reasonable activity and stability in acid

electrolytes. Various electrocatalysts for which results have been

reported in the literat\ire are summarized. The electrochemical behavior of various carbons and graphites , lithiated NiO and various defect

metal oxides including RuO^/Ti is considered in some detail relative

to 0^ generation as well as reduction kinetics.

Key words: Carbon; electrocatalysis; nickel oxide; semiconductor

electrodes; oxide electrodes; oxygen electrochemistry.

1. Introduction

With the energy problems facing many nations, renewed interest has developed in

electrochemical energy conversion and storage, and among the various electrochemical

devices, particularly fuel cells and water electrolyzers . An important part of most

such devices is the oxygen electrode. Unfortunately it is one of the more troublesome

parts. The irreversibility of the oxygen electrode leads to excessive voltage losses and

restrictions on power densities of the cells. Fiirthermore, in an effort to achieve an acceptable level of performance, it has been necessary to resort to expensive precious

metal oxygen electrocatalysts in very high area forms, which are diffic\ilt to maintain over

long operating times.

In light of the importance of 0 electrodes and problems associated with oxygen

electrocatalysis, this lecture will be directed principeilly to oxygen electrode reactions

on nonmetallic siirfaces.

2. General features of oxygen electrode reactions

The pronounced irreversibility of the oxygen electrode reactions at moderate temperatures has severely con^jlicated mechanistic studies. The exchange current densities for the

oxygen electrode are very low—typically 10"-'-^ to 10""'""'" A/ cm on an effective catalytic

surface such as platimm. Consequently the current densities near the reversible potential

are generally too low to permit measurements under conditions where the kinetics are sensitive to the reverse as well as forward reactions. Further, the experimentally accessible

203
portions of the cathodic and anodic 'branches of the polarization curves are sufficiently

separated in potential that the surface conditions differ very substantially. Therefore,

the cathodic and anodic processes under these conditions are probably not the reverse of

each other. To complicate the situation further, the oxygen electrode reactions may proceed

through a large number of pathways. This explains why the mechanisms for 0 generation and

reduction are still not fully understood even on platinum, the most extensively used and

most studied 0^ electro-reduction catalyst.

Before discussing 0^ electrocatalysis on specific non-metallic surfaces, some of the

general features of the possible electrode mechanisms will be considered. 0 reduction in

aqueous solutions requires a strong interaction of 0^ with the electrode surlace for the

reaction to proceed at a reasonable rate. Three types of models for such interaction have

been proposed [l , 35]1 . These are shown in figure 1, together with the corresponding reaction

0 = 0 M

0 — 0 \ /

M

A. Griffiths

model ( 2)

M^-O

0 ^^ 120°

M ^

B. Pauling

model (10)

,0H

0-0

/ \

M M

C. Bridge

model

.z+2 0 2H

. —^ m'' + 2H 0

^OH

'

Pathway I

M^-^^- 0.

2e + HgOg (IIA)

Pathway II

2H+

1iH*"^m'^ + 2H-0 (IIB)

^0=

2E^ M^'*'-'--OH hE

Pathway III

Figure 1. Models for adsorbed 0^ and corresponding reaction pathways for 0^

e le c t rore duction.

pathways for 0^ reduction. The Griffiths model [2] involves a lateral interaction of the tt orbitals of the 0^ interacting with empty d 2 orbitals of a transition element, ion or metal

atom with back bonding from at least parti aily filled d or d orbitals of the transition

xz vz element to the tt* orbitals of the O^. This type of interaction should lead to a weakening

of the 0-0 bond with a corresponding lengthening of this bond. The Vaska complexes

[e.g., Ir(02)Cl(C0) (PPh-)^] appear to form such complexes with 0^ [3,^]. These compounds

are selective oxidation catalysts for cyclic olefins [5]. The formation of a strong metalto-oxygen interaction results in a weakening of the 0-0 bond and an increment in the length

of this bond [6]. Sufficiently strong interaction of this type may lead to the dissociative

adsorption of 0^ with probably simultaneous proton addition and valency change of the transition element m the manner represented by Pathway I in figure 1, followed by reduction of

the M(0H) to regenerate the catalyst site. Sandstede et al. [T ,8] have attempted to explain

oxygen reduction with square pyramidal Co(ll), Fe(ll) and Fe(lll) complexes as well as on

the thiospinels on the basis of such tt bonding. Tseiong, Hibbs and Tantram [9] have proposed

that 0 reduction on Li-doped NiO changes from a non-dissociative to dissociative mechanism

For reviews see ref. [1,35]. Figures in brackets indicate the literature references at the end of this paper.

204
above the Neel point (200°C for their 'vlO-atom % Li doped NiO) in order to explain the

increment in catalytic activity in KOH hydrate melts ahove the Neel temperature.

It is indeed likely that 0^ reduction does proceed hy a dissociative adsorption

mechanism on some electrocatalysts . Up to date, however, no unambiguous experimental

e^'/idence appears to have been presented for 0^ reduction by such a mechanism on either

metallic or non-metallic catalyst surfaces in aqueous solutions at moderate temperat\ires

.

With most transition metal catalysts, the most probably structure for 0^ adsorption

is the Pauling model [lO] in which sp orbitals of 0^ interact with d 2 orbitals of the

transition metal. The square pyramidal complexes of Fe(ll) and Co(ll}| which have good

activity for 0 reduction in acid solutions, appear to involve such an end-on interaction

on the basis or esr and other e^'/idence [ll]. This adsorption of 0^ is expected to be

accompanied by at least a partial and probably complete charge transfer to yield a superoxide and then peroxide state, as represented by Pathway II in figure 1. The adsorption of

the 0^ on the square pyramidal complexes of Fe(ll) and Co(ll) may lead directly to the

superoxide state. With somewhat similar oxyhemoglobin complexes of iron, various workers

have proposed that 0^ binding to the iron involves Op or Op states with Fe in the III valent state [l2-lU]. The change in valency state or the transition metal coupled with the

change in Op oxidation state during formation of the 0^ adduct corresponds in principal to

j the redox electrocatalyst concept proposed by Beck et_ al. [I5,l6].

The further reduction of the Op beyond the peroxide state requires rupture of the 0-0

bond. Such can occur in Pathway IIB through the formation of 0^ or HO" free radicals in

solution or the simultaneous reduction-bond cleavage (electrochemical desorption) to yield

E^O or 0H~. Neither of these processes are likely to be sufficiently fast at practical

operating potentials for Op cathodes but the electrochemical desorption is a better candidate. The free energies of formation of the 0~ and HO* free radicals in solution are just

too high to achieve sufficiently high concentrations for the subsequent homogeneous reactions to proceed at rates corresponding to reasonable ciirrent densities at acceptable electrode potentials. Substantial evidence exists for Pathway IIA yielding solution phase

peroxide for various metallic and non-metallic electrode surfaces. With non-metallic

electrodes such as carbon, graphite and lithiated NiO in aqueous alkaline solutions, significant amounts of peroxide are found in solution and the potential under open-circuit conditions follows the Nemst equation predictions for the O2-HO2 couple [see e.g. ref. 17, I8]

.

ij Eventually the peroxide is further reduced at practical O2 cathodes. The formation of soluI tion phase peroxide as an intermediate, however, is unfortunate since it leads to less

j favorable operating potentials. With carbon electrodes in alkaline solutions, it is necesj sary to incorporate a peroxide reduction or decomposition catalyst in the electrode to

i prevent the btiildup to high concentrations of the peroxide within the electrode structure.

' One possibility is to utilize the spillover effect with the peroxide produced at one site

I migrating on the surface to an adjacent or nearby site where the peroxide elimination step

j occurs. This indeed may occ\ir to some extent with high area carbon electrodes containing

peroxide elimination catalysts such as Pt or Ag,

I Under some circumstances it is possible that the superoxide species 0 * may desorb to

I yield the solution phase species. This ion is formed as a reasonably stable entity during

, Op reduction in aprotic solvents [see e.g., ref. 19-22] and probably in carbonate melts

[23]. The superoxide ion also has been proposed to be formed in aqueous solutions on Hg

[2k,2^], amalgamated gold [26] and carbon paste [27] cathodes in the presence of surface

> active agents. Under these circumstances, Divisek and Kastening [2^1,25] propose that the

surfactant molecules displace water molecviles from the surface, impeding access of water

to the adsorbed Op~, and thereby inhibiting further reduction. Dubrovina and Nekrasov [26] nave proposed that the 0^' radical may be stabilized through the formation of a complex

vith the surfactant molecule. In the absence of adsorbed organic species, however, it does

not appear that an O^* desorption mechanism contributes significantly to the observed

djrrent. Rotating disk experiments have rather clearly demonstrated this for graphite [28]

and eilso gold cathodes [29], as will be discussed for the former case later in this lecture.

Pathway III in figure 1 provides an alternate means for bringing about rupture of the

0-0 bond through the formation of an -0-0- bridge . Such a mechanism may come into play with

the proper surface spacing of transition metal atoms or ions in a metal, oxide or thiospinel or in a bimetal complex such as a macrocyclic. The formation of the bridge species

205
requires proper spacing of the two me^al species plus partially filled d or d orbitals

to participate in bonding with the sp orbitals of the oxygen. MacrocycJic transition

metal complexes of the type M-O^-M been synthesized [e.g., see ref. 30, 31, 36, 37] and

appear to occur naturally in hemeythrin.

For any of the mechanisms in figure 1, considerable questions exist as to the reversibility of the Og adsorption step at the rather high rates involved with practical 0^ cathodes. For Og to bond to will generally require the replacement of a water molecule or anion of the electrolyte—a situation which would normally be exgejted to be imfavorable

to 0^ unless the 0 adduct has a pronounced dipolar character (M 0-0") [32, 33]. The

reversibility of tne formation of such adducts in water and other solvent systems is strongly dependent on the ligands and generally favored in the "picket fence" porphyrins in which

the large rather bulky organic groups create an environment more conducive to 0^ uptake

[3!+]. Porphyrins of this type are probably responsible for the relatively good performance

of carbon cathodes as 0_ electrodes in blood, but it is doubtful whether they can be ren- dered sufficiently stable for prolonged use in the strongly acid and alkaline solutions of

most fuel cells above 100°C.

At steady state any peroxide formed on the electrode surface must be subsequently

further reduced or decomposed. When the 0 reduction proceeds entirely through a peroxide

state on the electrode surface and/or in the solution, the process is usually referred to

as "series" whereas when Op reduction proceeds simultaneously by a dissociation step without a peroxide state as well as through a peroxide state, the processes are described as "parallel". Peroxide has been detected in the solution phase during 0^ reduction on many

metallic and non-metallic electrodes ( e.g. , Pt , Au, Ag, Fb, Ni , NiO, carbon, graphite) in

various aqueous electrolytes at T < 100°C. Consequently little doubt exists that the

peroxide mechanism is often functional. Some authors have proposed that reduction also

proceeds simultaneously through a dissociative adsorption mechanism as a parallel mechanism.

Unfortunately the evidence is for the most part ambiguous.^

OH + OH-

. 2. Free energy relationships

for the reduction of

oxygen in basic solution

at 25°C. Standard reduction potentials in volts

are indicated on lines

designating steps.

Potential for O2/O2

couple from ref. [25].

Multiply ordinates by 23 to obtain free energy in

kcal/mol

.

'Damjanovic, Genshaw and Bockris [39] have proposed that their rotating ring-disk data provide evidence for the parallel mechanism on Pt. "Hieir treatment of the kinetics does not

consider the peroxide desorption step. This woi£Ld explain their results [kO] with.out a parallel mechanism.

206
0.0

-1.0

-2.0

o

m

u

c

V

-3.0

-I4.O

-5.0

0.00

°2 ^ \ ^\ \ "^^'^^ \ +0.68N. \

+0.72

OH + HgO \ +1.77^

+1.23\

—

+2.82

1 1 1

1

2H 0

1

1 1 2 3

I*

Fig. 3. Free energy relationships

for the reduction of

oxygen in acid solution

at 25°C. Standard reduction potentials in volts

are indicated on lines

designating steps.

Potential for O2/HO2

couple calculated from

data for O2/O2 (ref. 25)

and pK = 4.88 for O2-HO2

(ref. 38). (Multiply

ordinates by 23 to obtain

free energy in kcal/mol.)

Reduction state

The question of whether 0^ reduction can proceed by a parallel mechanism is more than

just academic. The free energy changes and corresponding standard electrode potentials

associated with the various 0 couples are shown in figures 2 and 3 for alkaline and acid

solutions. The standard electrode potential of the 0 - peroxide couple is far less anodic

than that for the four-electron 0^ reduction (by +0.h5 and -O.JSV in alkaline and acid solutions, respectively). Further the O^-peroxide couple is usually much more reversible than

the other steps shown in figures 2 and 3 with the result that the potential is often con-

*rolled predominantly by the O^-peroxide step.

In principle it should be possible for the potentisQ. of the Op-peroxide couple to approach the reversible potential for the four-electron reduction by depressing the peroxide

concentration to the equilibrium value for the reaction

2H0, = 0^ + 20H (1)

In practice this has not proved possible even with extremely effective peroxide decomposing

catalysts under open-circuit conditions because the equilibrium peroxide concentration is

extremely low ( i.e . , 10~-'-5m HO^" in IM caustic). With high area carbon electrodes on which

the 0 -peroxide co\3ple has a high exchange current density [ e.g. , 10 A/cm (true area)]

in alkaline solutions, it has been possible to improve very substantially the performance

through the incorporation of very effective peroxide elimination catalysts [see e.g.,

ref. i+1, 68].

207
3. Sununary of Materials Examined as 0 Electrocatalysts

Table I is a partial list of the various materials which have been examined as 0^

reduction electrocatalysts. Many of the compounds listed in this table have substantial

electrocatalytic activity for 0^ reduction in alkaline solutions but do not have sufficient

long-term stability. Only a few have significant activity in the concentrated acid solutions and only platinum in ultrahigh area form has combined sufficient activity and stability. A few of these systems will "be considered in some detail.

Table 1. Compounds considered for O2 electroreduction catalysts in the literature.

Metal chelates

Type Example Literature reference

N2S2 diacetyl-di(thiophenylbenzhydrazone)^ [DT] 7,42,43

^l^h bissalicylaldehyde-o-phenylenediaminea [Pfeiffer] 7,42,43

Nit tetraphenyl porphyrin^ ["TP] 7,42,43

phthalocyanines^ rp^j 7,42-44,47

84 bisdiphenyldithioethylene^ 42,43

poljrmeric versions of above 42,43

hemin 48

Order of activities:

for Ni^: Fe > Co > Cu > Ni '\- Mn

for N20i^: Co > Mn > Fe > Cu > Ni

for Co II: N2O1, > N4 > N2S2 > Oit ^ S4 for Fe II: >> N2O2 > N2S2 > O4 a, Si^ Thiosplnels

Me^Me2^Sit Me^ = Mn, Fe, Co, Ni, Cu, Zn Md^ = Ti, V, Cr, Fe, Co, Ni

also seleno- and telluro-spinels

activity: 00331^ > NisSi^ > FesSi^ Disulfides, diselenldes, ditellurides

Co compounds active; for Cu, activity is highest

with disulfide

Oxides

Spinels

Spinels MeaMe2b0it: Co^Oi^, CojNiOtt, C0AI2O4 active 52-55

Tungsten bronzes M5JWO3 57-60

Perovskites SrxCo03 active 9 55 Other oxides: NiO (lithiated) g'lg

RUO2 SI AU2O3 62 Metals

Ft and Ft family metals, alloys^ 65-68 83-85

Au and Au alloys 52 54' ^8 29*40,63

Carbons

Ion implanted carbon 17,18,28,29,69,86

^Cations: Ti, V, Cr, Mn, Fe, Co, Ni, Cu, Zn The work on Ft is very extensive. For a review see references 75-78,

208
It. 0^ Electrocatalysis on Carbon and Graphite

On both carbons and graphites of low mineral content, the reduction of molec\ilar 0^

proceeds to hydrogen peroxide in alialine solution almost quantitatively at low peroxide

concentrations. Typical current-voltage curves obtained with the rotating disk technique are

shown in figure h for the edge and basal planes of stress-anneaJ.ed pyrolytic graphite in

IM KOH. The hysteresis evident in these slow sweep curves (lO mV/sec) is caused by changes

-0.7 -0.5 -0.5 -0.4 -0.3 -0.2 -0.1 Potential re Hg/HgO, 1 M KOH (V)

Reduction of O2 on high

pressure annealed pyrolytic

graphite. edge orientation, cleavage orientation. Electrolyte: 1 M KOH,

pressure O2: 0.97 atm,

temperature: 22°C, electrode

area: 0.22 cm^. Rotation

rates indicated on curves in

rpm. Direction of sweep indicated by arrows. X-ray

rocking angle of graphite

.0k = 0.4°.

in the oxidation state gf various groups present on the edge orientation. Isotopic experi-

::ents [70] involving 0 and the rotating disk experiments [28] indicate the following

-echanism in NaOH and KOH solutions

Step I. S + 0^ -> S - 0, (2)

Step II. S - 0, + e -> S - 0^ (3)

Step III. S - 0.

S - 0,

HOH S + OH

s + 0„

+ OH

ih)

Step II is ordinarily rate controlling with a transfer coefficient of 0.50 but at high

current densities, the potential- insensitive Step I becomes controlling. Corroborating

evidence for such an 0^" adsorption step is to be found in the reversible chemisorption of

Op on carbon from the gas phase, presmably without bond cleavage, as revealed by esr

studies [71]. Step III is qmte surprising and is definitely a surface reaction rather than

a second order solution-phase reaction between superoxide radicals (O or HO^'). The fact

that Step III is second order in the surface concentration of adsorbea 0 ~ indicates that a direct interaction between two adsorbed 0^" is required for this step, while the kinetic

'measurements have been carried out principally on various types of graphites, the results

for various carbons [l7] are sufficiently similar to suggest the same mechanism.

209
Of particular note is that with single crystal graphite and stress-annealed pyrolytic

graphite, the 0^ reduction as veil as peroxide oxidation reactions are very inhibited on the basal plane with the exchange current at least three orders of magnitude less on the

basal plane than on the edge orientation. Imperfections in the basal plane set a lower

limit to the checking of how far the exchange current is depressed on this plane.

This depression of the exchange current on the basal plane may be caused by

1. anisotropic surface-semiconductor properties of the graphite

2. the lack of suitable functional groups or adsorption sites on the basal plane.

The carrier concentration in graphite is rather low near the Fermi level and hence some significant fraction of the applied potential may fall across a space charge region for

the basal plane. To check on this, a.c. impedance measurements have been carried out on this plane [72]. The differential capacitance curves and corresponding voltammetiy curves

are shown in figure 5 for both 1 K[ NaOH and 1 N H^SOj^. These curves are in strong contrast

to the corresponding curves for the edge orientation ' [73] which show much structure. The

capacitance on the basal plane is much lower than on the edge orientation or on typical

metals such as Hg or Au, suggesting that a significant potential drop occurs across a space

charge region within the graphite. The potential dependence of the capacitance, however,

is close to parabolic rather than the hyperbolic dependence expected for a semiconductor

electrodg. This may be caused by the high carrier concentration in the graphite; i.e.:

'^6 X 10 carriers/cm-^ [7^]. Further, the capacitance curves give no indication of specific

adsorption, even in the presence of relatively high concentrations of such anions as iodide [72].

4.5

E 4.0 u

3.

o

3.5

\

\

oV

V

o\

I

/

/

o

o

<

+ 0.5 0 -0.5

E (V vs NHE)

Fig. 5. Capacity-potential and current- potential curves for the basal

plane of stress-annealed pyrolytic graphite (AGi^ = 0.4°) in

1 N H2S0it and 1 N NaOH solutions at 25°C and 1000 Hz

(sine wave): x 1 N H2SO4;

o 1 N NaOH. Scan rate for the

current-potential curves 0.1

V/sec, direction of sweep indicated by arrows: A) 1 N

H2S0i,; B) 1 N NaOH.

210
To check further on the extent to which electrode reactions on the hasal plane may be

inrpeded by semiconductor effects, the behavior of the FeCCET)^-^ - Fe(CN)g redox couple has

been examined on this surface [28]^ The apparent exchange ciirrent density is appgomiately

1/3 of that on the edge (l.l mA/cm on the basal plane as compared with 3.0 mA/cm on the

edge in 0.005 M K3FeCCIl)g + 0.005 M Kl^Fe(CIT)g + 0.5 M KgSOj^ at pH = 3 and 26°C). This

difference may be caused by a difference in the ratio of true to apparent siirface area or

ionic double layer effects (different point of zero cheirge) as well as semiconductor effects

but is certainly far less than the two orders of magnitude difference in the exchange current densities for the reduction on the basal and edge planes.

The most likely explanation for the suppression of 0^ reduction on the basal plane is

the lack of suitable adsorption sites for O^Cads), 0^" (aas) and O^H'Cads) during the reduction. This raises the question as to what particular adsorption sites or functional groups

are involved in the 0^ reduction on the edge surface. A large number of functional groups

have been proposed to exist on carbon surfaces [79, 80] and some of these should persist on graphite as well as carbon at the potentials for 0^ reduction in alkaline solutions. Even

so, it is far from clear which, if any, provide the adsorption sites involved in the 0^

reduction.

Garten and Weiss [8l] some years ago proposed that quinone groiips are involved in the

following way:

0

The Og first adds to the olefinic bond associated with the qiiinone group to form the hydroperoxide or molecular oxide. Electron transfer then occurs to form 0 "(ads) and subsequently splits off as HO^" after a ftLrther electron transfer. Steps I and II of this mechajiism

a.re compatable with the kinetics but step III is not. The peroxide formation step is second

:rder in 02~(ads) and hence appears to involve interaction between two 0 "(ads) with simultaneous electron transfer either directly or through the carbon substrate. The nature of

this interaction is not clear. A similar second order step involving 0 "(ads) also appears

'0 be involved in 0 reduction on gold [29, 63].

5. 0 Electrode Reactions on Nickel Oxide

Oxygen reduction at room temperatures proceeds through a peroxide mechanism on lithiated

-iO in alkaline solutions [9, l8, 82]. In O^-saturated alkaline solutions of HO ", the

pen-circmt potential of anhydrous NiO produced by various methods corresponds to the

-.hermodynamic value for the 0^ - HO^" couple with the proper dependence on HO^" activity

Txie IJiO most extensively studied in the author's laboratory [l8, 82] is mosaic crystal-

-ine anhydrous NiO ([lOO] orientation) produced by hydrothermal growth on isomorphic MgO

£rom the NiBr^ - H^O vapor phase at 'v600°C, using the method of Cech and Alessandrini [88].

-ne mosaic ciystals are then doped with 0.1 - 1.0^ Li by high temperature infusion

l'+50-1550°C). This material is a p-type semiconductor and hence processes such as 0 -eduction would be expected to be inhibited. Examination of redox couples such as ferriferrocyanide with the rotating disk technique do indicate such cathodic inhibition in acid

solutions with the anode branch under combined kinetic and diffusion control [87]. Differ211
ential capacitance measurements in acid solutions indicate that most of the electrode potential change occurs across the space charge region within the NiO(Li) at electrode potentials cathodic to 0.8V vs. standard hydrogen electrode (SHE)[88]. In alkaline solutions,

however, the ferri-ferrocyanide couple is no longer cathodlcally inhibited and hoth the

cathodic and anodic branches are essentially diffusion controlled even at high rotation rates

[87]. The flat band potential shifts in alkaline solutions from a positive to quite negative value relative to the reversible potential of this couple with the result that the

potential changes occur principally across the Helmholtz layer. The reversible potential

for the O^-peroxide couple, however, is more negative and therefore 0 reduction even in

alkaline solutions is still qiiite inhibited [82]. The oxidation of HO^" to 0^ proceeds

more readily but is still much less reversible than on carbon or graphite. Similar results

have been observed on lithiated polycrystalline NiO produced by high temperature oxidation

of nickel metal in air or 0^ [18].

18

0 techniques have been used to demonstrate that the 0 liberated during the oxidation

of HO ~ to 0 at controlled potential on NiO in alkaline solutions originates only from the

HO ~ 182]. Thus the 0-0 bond does not appear to be broken during the oxidation process, ana presumably also during the reverse reduction of 0^ to HO^".

So far in this lecture, we have only considered the reduction of O^. Nickel oxide is

some interest from the standpoir . :trodes are used in water elects

oxide layers under such conditions.

of some interest from the standpoint of the anodic generation of O^, however, since nickel

electrodes are used in water electrolysis in caustic and these electrodes develop anodic

18

Og generation has been examined on mosaic NiO(Li) in acid solutions using 0 labelling

techniques [90]. The dissolution of NiO in concentrated acid solutions_is extraordinarily

slow and is under pure kinetic control. After correction for lattice 0 entering the solution through dissolution, a substantial fraction of the anodically generated 0^ has been

found to originate from lattice oxygen (10^ at 1. 8V vs„ SHE). The proposed mechanism is

the falling of p-carriers into surface 0 or OH traps. The subsequent formation of 0^ is

then coupled to the NiO dissolution process and may proceed as follows:

Step I: 2 0^" +

2p'^ -> 20" (6)

Step II: 2 0" + 2h"^ -> 2 OH (?)

Step III: OH + OH -> H^O + 1/2 0^ (8)

Step IV: 2Ni?'*'^^. ^ 2Ni^"^ (9) latti ce aq

2NiO + 2H"^ + 2p"^ ^ 2Ni^'*' + H^O + 1/2 0^ (lO)

aq 2 id

Steps I and II may be a single step or separate steps as shown. Step III undoubtedly involves several component steps. The overall process is independent of the steps. The

standard free energy change for the overall process relative to the reaction 2H + 2e H^

can be calculated from the standard free energies of formation as listed by Pourbaix [91];

NiO, AG° = -51.3 kcal/mole; Ni , AG° = -11.53 kcal/mole; H^O, AG° = -56.5 kcal/mole. The

calculated relative standard free energy change for the overall reaction is +23.1 kcal/mole

corresponding to a standard electrode potential of +0.5V.

Alte^atively the 0^ formation from lattice oxide may involve two p carriers falling

into a 0 surface state; i.e..

Step I: 0^" + 2p'^ ^0 (ll)

Step II: 0 ^ 1/2 0 X (l2)

2 .2+ „.2+

aq Step III: Ni, ^ Ni lattice

(13)

NiO, + 2p'^ ^ Ni^"^ + 1/2 0_ {ih] lattice aq 2

212
where Steps I and II are multiple step processes. For the overall process, AG° = +39.8

kcal/mole and E° = 0.8TV.

The potentials for toth of these overall reactions lO^^d ik are cathodic to that for

the discharge of 0 from water. Since the activity of Ni is far less than the unit

value, the eguilihrium potentials may he considerahly less*^anodic than the standard values.

This may he the explanation for the ohservation that at potentials even as small as +0.7V

vs. SHE anodic currents were ohserved for NiO (0.26 atom % Li) in 1 N acid solutions. In

the absence of such an anodic process a cathodic current is expected to attend the dissolution of the NiO due to the reduction of lattice Ni (resulting from the Li doping) to

Ni2+ [90].

The voltammetry curve for the [lOO] interface of the mosaic crystals indicates two peaks (fig. 6) in "both acid and alkaline solutions, one at 0.95V vs. SHE (Peak l) and a second at l.hY (Peak: II) . Both have a pH dependence of -60 mV/decade over the pH range 0-1^+. At slow sweep rates (<_ 60 mV/sec), the anodic and cathodic sweep peak potentials coincide,

indicating reversibility.

Yohe et al. [82] have compared the voltammetry resiiLts with the thermodynamic data

compiled by gjurbaix_^^91 ] ajid_|_gonclud^^ that peaks I and II probably correspond to the oxidation of Ni to Ni and Ni to Ni , respectively, coupled with the removal of protons

from adjacent oxygen ions. These reactions appear to occur only at the electrode siorface

since the total charge under both peaks corresponds to l^^s than a monolayer (65^) of the

charge for the (lOO) plane, assuming one electron per Ni ion for each peak, and an idealized ( 100

) 2^urface orientation with no surface roughness. The difficulty of converting all

surface Ni ions to a higher valency state is not surprising since complete conversion

would probably involve very large s\irface excess charge. It is even possible that both

peaks correspond to the Ni II - Ni III conversion with the marked difference in potential

due either to different siarface sites with differing charge compensation by surface 0~ and

OH ions or to induced heterogeneity arising from a large increment in the free energy of

oxidation for Ni II ions adjacent to an Ni III surface ion.

Another interesting feature is the extraordinarily slow rate of dissolution of the

mosaic NiO(Li) in acid electrolytes and the pronomced potential dependence of the solution

rate. At room temperatiires t^e dissolution rate in acid solutions (e.g., 6 M H^SOj^) is

typically a few micrograms/cm -hr or less, making it possible to make use of this oxide

as an electrode over very extended periods of time even in strong acids. To facilitate

dissolution rate studies, the dissolution kinetics have been studied in the author's

laboratory at temperatures just below the boiling points of the electrolytes. The dissolution rate is entirely under kinetic control [90]. In H^SO^^ solutions the dissolution rate

passes through a maximum as the potential is made increasingly anodic (fig. 7), in accord

with the predictions of Engell [92] and Vermilyea [93] when the dissolution is controlled

by the transfer of ions across the oxide-electrolyte interface. The maximum should corres- pond to a shift from control of the dissolution kinetics by the potential energy barrier

for cation transfer to control by that for anion transfer across the interface. The theory

213
1 1 r

E (V) vs SHE

Fig. 7. Dissolution of lithiated nickel oxide in helium-saturated

1 N HCl and 1 N H2SO4 solution at 95°C [89] . Symbols corre- sponding to different group of electrodes:

o • Group 1, (CLi)MS = 0 . 88% ; A Group 3, (CLi)MS = 0.26%;

OGroup 2, (CLi)i^s = 0-33%; A* Group 4, (CLi)MS = 0.12%.

Dashed line: calculated from eq (1) with the following

parameters: ot^z+ = 0.62, a_z_ = -0.62, r^ = 15,000 Ug/cm - h, (f)f = 1.68 V vs SHE. (Concentrations of Li in cation -%,

based on the slopes of the Mott-Schottky plots.)

predicts Tafel linearity on either side of the maximum. It is not possible, however, to

identify the specific ionic species being transferred over the potential energy barriers

from the observed Tafel slopes, which yield a n (a = transfer coefficient, n = charge of

species )

„

6. 0_ Electrocatalysis on other Oxides

The anodic generation of 0_ has been examined on a namber of defect metal oxides, prepared

hy thermally decomposing a metal salt (usually the chloride) with titanium as the substrate.

Potential-current curves obtained by 0' Grady and Iwakura in the author's laboratory are shown in figure 8 for a fe^ of these oxides [9^]. The current-voltage curves at lower ci:ir- rent densities ( <10 A/ cm ) are strongly influenced by the intrinsic changes in valency

state of the oxide even though the data were acquired with decreasing current with three

minutes at each data point. Since the current densities are apparent values, area differ- ences may have influenced the relative positions of the curves. Of particular importance

is the low Tafel slope ('^'0. OUv/decade ) for several of the defect oxides (e.g., RuO^, IrO^).

The behavior of the corresponding pure metals as 0^ anodes is distinctly different from

that of the defect metal oxides in figure 8 even though the metals are oxide coated at the

potentials where 0^ generation occurs. Ruthenivim metal is not stable as an 0^ anode in

U M KOH. Active dissolution begins already at '^^1.3V vs. RHE.

214
density ( A/cm^ )

Fig. 8. O2 generation on various metal oxides. Anodic polarization

curves on metal oxides on Ti substrate in 02-saturated

4 M KOH, obtained with 3 min/point after polarization with

increasing current up to 0.1 A/cm^ (dashed line represents

IR drop corrected curve for RuOx) [94]

.

The RuOjj catalyst appears to be similar to that used for the dlmensionally stable

anodes for CI2 generation [97]. This oxide has been investigated in some detail in the

author's laboratory [61]. The 0^ generation reaction on RuG /Ti is first order in OH over

the pH range 11 to ik. Unfortunately the stoichiometric number cannot be obtained because

of interference from the intrinsic oxidation-reduction reactions of the RUO2 at low current

densities near the reversible potential for the U-electron reaction.

O'Grady et al. [61] have found that the Tafel slopes are independent of the catalyst

loading over the ranges 10~" to 10 moles Ru/cm^ and the apparent exchange current densities

are approximately proportional to the catalyst loading over this range. Such a proportion is to be expected since the true-to-apparent area ratio is also approximately proportional to the catalyst loading. The Ti only acts as a substrate and does not have any directgeffect

on th^ activity of the RuO^ per unit true area. For RuO^ coverage greater than 10~ moles

Ru/cm , ESCA did not indicate any Ti in the RuO^.

Several of the reaction paths for 0^ generation tabulated by Damjanovic [T6] for acid

solutions can have Tafel slopes corresponding to dV/d In i = 2RT/3F or O.Oi+O V/decade when

modified to a form suitable for alkaline solutions. All have the common feature that the

rate controlling step is a second electron rather than first electron transfer. With judicious choice of HO vs. 0H~ as the reactant in the first and second electron transfer steps

or proceeding proton transfer steps, it is possible to have the kinetics also first order

in 0H~.

With electrode siarfaces involving species with ionic character, O'Grady et_ al. [61]

prefer to write the mechanism in such a way as to indicate the possibility of changes in

the effective valency state of the catalyst sites. A simple mechanism compatible with the

observed Tafel slope and reaction order is as follows

The next four paragraphs are quoted from reference 6I with only minor change.

215
_ k

Step I: + OH (SOH)^ + e (fast) (l5) -a

Step II: (SOH)^ (SOH)^'^-'- + e~ (slow) (l6)

+z followed by subsequent processes yielding finally 0^ and regenerating the site S ; for

example

,

Step III: (SOH)^'^''" OH" S - 0-H—OH" S^ 0 HOH

z+1 " - " ' - " z

I ^ (17)

(SOH) OH S - 0-H—OH S 0 HOH

The sites are probably Ru but also could be higher valency state Ru ions. These sites

are probably hydrated although the waters of hydration are not shown in the mechanism.

In Step I the site S has been proposed by 0' Grady et al. to be a Ru ion of the oxide

lattice. Step I involves the oxidation of the Ru to a higher valency state with partial

transfer of electron charge from the OH species to Ru forming a (RuOH) complex. Such a process should require much less energy than the formation of an OH free radical or a rather

weakly adsorbed OH radical. Step II involves the further oxidation to ( RuOH ) , possibly

followed by the foraiation of (RuO)^ before proceeding with Step III. The RuO electrode

surface may have some of the same features as the Ru(NH ) 0Ru-(NH2)^0Ru(NH ) T+ ion, re- ported by Earley and Razavi [95. 96] to oxidize OH" homogeneously to form Ogo These species

have the structure Ru-O-Ru-O-Ru Earley and Razavi have pointed out that Ru has a tendency to form species of coordination number 7 and that there are low-lying antibonding ir* orbitals which can accept electron charge from OH . Thus the (RuOH)^ intermediate (other

species in the inner coordination sphere not shown) seems reasonable."

From the applied standpoint, R'^^O^ does appear attractive for use as an 0^ anode electrocatalyst. While the exchange currents per unit true area are comparable to Pt in alkaline solutions, the much lower Tafel slope results in significantly lower overpotential at high current densities.

0^ reduction on RuO^ in alkaline and acid solutions appears to involve a peroxide

mechanism [6l] although it is possible that "parallel" mechanisms are involved. Pronounced

hysteresis and a lack of stability of the oxide have interfered with kinetic studies with

the rotating disk technique.

T. Acknowledgment

The experimental data obtained in the author's laboratory and presented herein, have

been obtained principally by the following colleagues and students:

Oxygen electrode studies on carbon: I. Morcos

A. c. impedance measurements on graphite: J. P. Randin

Redox couples on NiO(Li): D. Tench, D. Yohe

Dissolution kinetics and 0^ generation on NiO(Li): C. H. Lee, A. Riga

Og electrocatalysis on RuO^ and other defect oxides: W. O'Grady, C. Iwakura

and J. Huang

The author also acknowledges helpful discussions on 0^ electrocatalysis with his colleagues

R. K. Sen, W. E. O'Grady and F. Urbach.

This work has been partially supported through grants from the U.S. Office of Naval

Research and NASA-Ames.

216
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218
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The Dissolution Kinetics of Lithiated

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[91] Pourhaix, M. , "Atlas of Electrochemical Equilibria in Aqueous Solutions",

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[93] Vermilyea, D.A. , J. Electrochem. Soc 113 , 1067 (1966).

[9^+] O'Grady, W. , Iwakura, C. , Huang, J. and Yeager, E. , ^published data.

[95] Earley, J. andRajavi, H., Inorg.

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219
V
National Bureau of Standards Special Publication 455, ElectTOcatalysis on Non-Metallic Surfaces

,

Proceedings of Workshop held at NBS, Gaithersburg, Md. , December 9-12, 1975. (Issued November 1976)

Photoelectrolysis of Water^

J. G. Mavroides, D. I. Tchernev, J. A. Kafalas, and D. F. Kolesar

Lincoln Laboratory, Massachusetts Institute of Technology

Lexington, Massachusetts 02173

The photoelectrolysis of water has been investigated by experiments on cells consisting of an illuminated n-type semiconductor anode, an aqueous electrolyte, and a

platinized-Pt cathode. It has been found that such cells operate either in the photogalvanic mode (no H2 evolved) or in the photoelectrolytic mode (H2 evolved at the cathode

by decomposition of water), depending on whether or not the electrolyte surrounding

the cathode contains dissolved oxygen. In both cases, current flows through the exter- nal circuit and O2 is evolved at the anode. For operation in the photogalvanic mode,

maximum values of 80-85% for the external quantum efficiency (ti) for current production have been measured at hv 4 eV with both single-crystal and polycrystalline

Ti02 anodes. Similar results have been obtained with SrTi03 anodes. The internal

quantum efficiencies, corrected for reflection and absorption losses, are close to 100%, indicating that the band bending in both cases under photogalvanic conditions is sufficient to separate the electron-hole pairs generated by photon absorption and also

that the oxygen overvoltage for charge transfer at the semiconductor-electrolyte

interface is negligible for illuminated anodes. For operation in the photoelectrolytic

mode, 'H is only 1-2%, for Ti02 and ~ 10% for SrTiO^ if the anode and cathode are

shorted together, but the efficiency can be greatly increased by applying a bias voltage. All the observations, including the increased efficiency for SrTi03, can be

given a straightforward explanation in terms of the energy levels of the electrodes

and the electrolyte.

Keywords: Energy levels; photocatalytic electrodes; photoelectrolytic cells; photogalvanic cells; quantum efficiency; solar energy.

1. Introduction

About three years ago, Fujishima and Honda published a letter [1]^ in Nature on "Electrochemical

Photoelectrolysis of Water at a Semiconductor Electrode". In this report, these authors suggested

that an electrochemical cell consisting of an n-type single-crystal Ti02 anode and platinum black

cathode could be used with visible light to decompose water into H2 and Og. Aside from its purely

scientific interest, this report has aroused considerable interest in possible utilization of the phe- nomenon for obtaining fuel, Hg, from solar energy. In order to further evaluate the potential of photoelectrolysis for this application we have investigated the physics and electrochemistry of this

effect by experiments on ceUs with polycrystalline as well as single-crystal Ti02 anodes [2], This

work then led to studies with SrTiOo anodes [3] and we report some of these results as well.

^This work was sponsored by the Department of the Air Force.

^Figures in brackets indicate the literature references at the end of this paper.

221
2. Energy Diagram of a Photoelectrolytic Cell

Figure 1 is a simplified energy diagram of a photoelectrolytic cell consisting of an n-type semiconducting anode, an acidic electrolyte, and a metallic counter electrode. (An acidic electrolyte has

been chosen for simplicity. An entirely equivalent discussion could be given for basic solutions.)

The two electrodes are connected through an external circuit, so that the Fermi level is the same for both. Because of the difference in work function between the semiconductor and the electrolyte, the

energy bands of the semiconductor are bent at the surface, so that the analog of a Schottky barrier

exists at the semiconductor-electrolyte interface. The semiconductor surface is irradiated by photons of energy at least equal to its energy gap. The photon-generated hole-electron pairs do not re- combine but are separated by the electric field of the barrier; the electrons move away from the surface into the bulk of the semiconductor and then through the external circuit to the Pt cathode,

_____-rr:rL_ L^o("7^2^

SEMICONDUCTOR

1.23eV

ELECTROLYTE

COUNTER

ELECTRODE

Fig. I. Simplified Energy diagram of an ideal photoelectrolytic cell.

where they discharge by the reaction

2e" +2H"'' ^ H (1)

while the holes remain at the surface of the semiconductor anode, where they can interact with the electrolyte to produce by the reaction

2p ^ 1/2 + 2H

,+ (2)

The overall reaction is 2hv

H^O . +1/2 (3) From this brief description of the principle of operation, it is evident that a number of conditions are necessary for efficient photoelectrolysis: (1) The energy bands at the semiconductor- electrolyte

interface must be bent in order to separate the holes and electrons excited by the light: Since the bending is due to the difference in work functions (or electrochemical potentials) between the semiconductor and the electroljrte, this puts an upper limit (for an n-type semiconductor) on the work function (or

electron affinity) of the semiconductor; (2) the relevant electronic levels must line up, i.e. , the hole

states of the anode with e(H20/02) and the Fermi level Cp with £(H"VH2), in order to permit transfer

of holes at the anode and electrons at the cathode; (3) the semiconductor must be chemically inert serving only to absorb light and produce the holes and electrons which make the above reactions possible and (4) the semiconductor must have an energy gap between 2 and 2.5 eV so as to make the most

effective use of the sun.

222
At the present time it is not obvious that all these conditions can be met simultaneously. In particular, can we find a material with a small energy gap that does not dissolve under anodic polarization? All the common semiconductors, such as Ge, Si, CdS and GaP.dissolve under these conditions.

In this talk we will discuss our results onTi02 and SrTi03 with the above criteria in mind. It should

be pointed out at the onset that both of these materials are extremely stable and do not dissolve; they

have energy gaps of ~ 3 eV, however, so that they are not suitable for efficient solar photoelectrodes.

I The anode material used in most of our experiments was the rutile form of Ti02. Some Ti02

I anodes were fabricated from single crystals, but most were polycrystaUine disks prepared by hot- pressing powdered rutile at 750°C under a pressure of 10, 000 psi. To make the disks conducting

I [p ~ 1 - 10 Qm (10^ - 10^ n-cm)], they were heated in vacuum for three hours at~ 800°C, which causes

' a loss of oxygen that makes them n-type and gives them a blue coloration. Other anodes were prepared by heating Ti foils at~ 600° C in an oxygen atmosphere for periods of three hours or more to form polycrystaUine Ti02 films by thermal oxidation.

In initial experiments we used both the differential space-charge capacitance and AC photocurrent

techniques to determine the flat-band potential of Ti02 with respect to the saturated calomel reference

electrode. This potential, (ep)^^,, is the particular value of the Fermi level for which the energy

bands of Ti02 are not bent. Therefore the amount of anode band bending for a given Fermi level, Sp,

is tfie difference between Sp and the flat band potential. With increasing pH, (£p)fi3 was found to vary

linearly with a slope of 0. 059 eV/pH (i. e. 2. 303 e/kT), as expected if the principal effect of the change in pH is to vary the Fermi level in the solution. At pH = 0, (£p)^~ -0. 2 eV, which is close to £(H"'7H2) = "0' 24 eV. (By convention, in an energy diagram like figure 1, levels Ijrtng above the satu- rated calomel electrode are negative, while those lying below it are positive. ) This means that Ti02

is quite different from the hypothetical semiconductor of figure 1, which exhibits significant band bending when Gp = e(H+/H2). For Ti02, since (^p)f^ P,(H'VH2), it seen that the band bending is very

small. This situation, shown in figure 2, indicates a high electron-hole recombination rate. Thus

the quantum efficiency for photoelectrolysis, r], which is defined as 100 times the ratio of the number

of electrons flowing in the external circuit (N^) to the number of photons incident on the cell, Np, and

calculated from the measured short circuit current density, J , and light intensity, S, at each photon

energy hv, i. e.

3. Experiments with Ti02 Anodes

I

-zrrr

1.23eV

SEMICONDUCTOR ELECTROLYTE

COUNTER

ELECTRODE

Fig. 2. Energy diagram of a TiO^ electrochemical cell operating

in the photoelectrolj^ic mode.

223
N J

71 = 100 ,j^=100-^hv (4) P

is only 1 to 2% at most in our best samples.

If the cell containing the illuminated Ti02 anode and the platinized- Pt cathode is exposed to air so

that the electrolyte surrounding the cathode contains dissolved oxygen, then charge transfer takes

place at the platinum electrode and it assumes a potential close to the equilibrium potential of Thus the energy states of the counterelectrode are shifted downwards (relative to those of the electrolyte) by~ 0. 8 eV. This is indicated in figure 3. Now when photo-generated electrons are supplied to the cathode, the energetically favorable process is not the transfer of electrons from the cathode to

SEMICONDUCTOR

^(H^H^)

I.23eV

ELECTROLYTE

COUNTER

ELECTRODE

Fig. 3. Energy diagram of a TiO^ electrochemical cell operating

in the photogalvanic mode.

r the H /H2 level according to eq (1) but rather their transfer to the H2O/O2 level, which causes the reduction of oxygen. In an acidic solution, the resulting chemical reaction is

2e" +2H"'" +1/2 H^O (5)

Illumination of the anode thus causes the cell to function in the photogalvanic mode with significant

band bending, but the passage of current through the external circuit although producing O2 efficiently

is not accompanied by the production of H^. As will be discussed later, T) here is almost 100%.

The remaining experiments on Ti02 described here were designed to evaluate the performance of semiconducting anodes in the process of charge transfer across the semiconductor-electrolyte inter- face. To perform such an evaluation, it is necessary for the reactions at the anode, rather than those

at the cathode, to be rate- limiting. This was accomplished by using a platinized- Pt cathode with a geometric area about 50 times that of the anode and operating with the electrochemical cell open to air, so

that there was a plentiful supply of dissolved oxygen in the electrolyte surrounding the cathode. Under

these conditions the cell operates in the photogalvanic mode, so that the current flow in the external

circuit is determined by the photo-generation of electron-hole pairs at the anode, the separation of these pairs due to band bending, and the transfer of holes from the anode to the electrolyte. Anode

performance was characterized by measuring the current flow. The gas evolved at the anode, which was collected by displacement of the electrolyte solution from an inverted buret, was identified as O2

because it was completely absorbed by pyrogallol. The volume collected was found to be 80-85% of

224
the value calculated by assuming that the current flow was entirely associated with the generation of O2 according to eq, (2); we assume that the difference of 15-20% was due to collection losses.

The operating characteristics of a photogalvanic cell under constant illumination from a 150-W Xe

source were measured for various values of pH. Typical results, obtained for pH = 0, are plotted in figure 4. The upper curve gives the output voltage, V, as a function of the current drawn from the

Fig. 4. Operating characteristics of a Ti02 photogalvanic cell showing the variation

of output voltage (V) and power (P) as a function of current density (J). pH=0.

ceU per cm of illuminated Ti02 surface. The lower curve gives the electrical power, P, extracted

by the external load. The two curves can be fit to a fairly good approximation by a model that considers

the cell as equivalent to a constant voltage source in series with an internal resistance. Here, the open circuit voltage is 0.52 V and the internal resistance is ~ 350 Cl. For operation in the shortcircuit mode, the anode current density Jg^ is 19 A/m2 (1. 9 mA/cm^). When the pH of the electrolyte was increased, Jg^ initially increased quite rapidly but then reached a limiting value of ~ 25A/m^

(2. 5 mA/cm^) for pH values of about 5 or higher.

The external quantum efficiency for current production was measured as a function of photon energy

on an optimized photogalvanic cell for light intensities in the linear region. The anode was illuminated

by monochromatic light from a source consisting of a 1000-W Xe lamp and a grating monochromator.

The light intensity incident on the cell was measured as a function of hv by means of a thermopile

power meter that had a flat response down to 0. 25 ^m.

Figure 5 shows the values of r\ vs hv for Ti02 films on Ti foil that had been thermally oxidized under different conditions and measured in a non-buffered electrolyte of pH = 13. For film "a" (7 flm

thick), T) reaches a maximum value of 84% at ^. =0. 34 fxm (~ 3. 5 eV), while film "b" (<2 fjm thick) has

a maximum of 83% at 0. 30 /nm (~ 4 eV). The positions of the quantum efficiency peaks correspond to 3 major absorption peaks in Ti02 that occur at about 3. 5, 4 and 4. 7 eV (0. 34, 0. 30 and 0. 26 fxm) respectively. The relative intensities of the efficiency peaks depend on preparation procedure and oxide

225
hv (eV)

5.6 5.2 4S 4.6 4.4 4.2 4.0 3.8 3.6 3.4 3.2

UJ X(A)

Fig. 5. Variation of external quantum efficiency with wavelength of a Ti02 photogalvanic

cell for two differ ent thermally oxidized Ti02 foil anodes. pH = 13.

layer thickness. Equally high efficiencies are also obtained with single crystals and hot-pressed disks.

When the maximum measured r\ values are corrected by taking account of estimated reflection

losses at the interfaces and absorption losses in the electrolyte, we find that the internal quantum effi- ciency is close to 100%, for both single -crystal and polycrystalline TiOg anodes. This high quantum

efficiency shows that the band bending in Ti02 under photogalvanic conditions (~0. 8 eV) is sufficient

to separate all the hole-electron pairs generated by photon absorption. It also shows that the anodes

have hole states (probably surface states) which are close to the H2O/O2 level of the electrolyte.

Other evidence for such states, lying in the energy gap, has already been presented at this workshop:

Professor Bard's elegant work on the exploration of these states by means of solution redox couples,

the Lincoln Laboratory study of surface states using ultra high-vacuum energy loss spectroscopy and

Professor Wolfram's theoretical work on the origin of these states. The observation of such high

efficiencies implies that, for illuminated anodes of Ti02, the oxygen overvoltage for charge transfer

at the semiconductor-electrolyte interface is negligible.

4. Experiments with SrTiO^ Anodes

The SrTi03 anoctes used in the experiments reported here were n-type disks 5 to 15 mm in diameter that had been cut from a single crystal (Titanium Pigment Division of National Lead Industries).

Preliminary experiments with an anode prepared by hot-pressing powdered SrTiOg gave results similar to those obtained with single-crystal material. The material was reduced by annealing in H2 at atmospheric pressure for about 3 hours at 900° C. After reduction SrTi03 changed from insulating to conducting, with a resistivity ~ 0. 15 fim (15 Q-cm), and from, colorless and transparent to blue-black

and opaque. An ohmic contact was made to each disk by ultra sonically soldering In to its back surface

and attaching a Pt wire to the In. The entire disk, except for the front surface, was then coated with

insulating epoxy or stop-off lacquer.

Again photoelectrolytic cells were prepared by immersing a SrTiO^ anode and a platinized- ft cathode into an aqueous electrolyte contained in a fused-quartz vessel. Each electrode was capped with a inverted buret filled with electrolyte, so that the gas evolved could be collected by liquid displacement

for volumetric measurement and chemical analysis. Provision was made for flushing the electrolyte

with either N2 or Ar in order to remove dissolved so that the cell operated in the photoelectrolytic

mode. Measurements were generally made after the solution had been purged for several hours.

226
Figure 6 is a plot of the variation of small signal ac photocurrent with electrode potential at pH = 13. 3. One can use such a plot to estimate the flat band potential from the value of voltage at which the photocurrent goes to zero. Two curves are shown here_p- one for TiO^ and the other for SrTiOg. Also sho\vn is the equilibrium hydrogen potential P(H^/H ) for this pH. These results indi- cate that the band bending for photoelectrolysis is very small for TiO and~ 0.2 volts more for SrTiOg. This is consistent with the difference in work function, or electron affinity, of ~ 0. 2 volt

that we have measured by the Kelvin contact potential method in air.

(-z

UJ

(£

CCoo

H

o

X

a.o

<

-1.4 -1.3 -1.2 -1.1 |-1.0 -0.9 -0.8 -0.7 -0.6 -0.5 -0.4

V (relative to SCE), volts

Fig. 6. Influence of elect:rode potential on small signal ac photocurrent in an aqueous electrolyte.

This ~ 0. 2 volt larger band bending in SrTiOg results in a quantum efficiency for photoelectrolysis

about one order of magnitude larger than that found for TiO^; i. e. , T] = 10% at hu = 3. 8 eV. The gases

evolved at the electrodes were collected and analyzed. At high pH values (pH > 13) the volumes of gas

collected at the cathode and anode were found to be over 90% of the values calculated from the inte- grated electron current by assuming that only the reactions given by eqs. (1) and (2) are taking place.

These results are consistent with the conclusion that photoelectrolysis is the only reaction occurring.

Summary

In summary it is especially significant that we have been able to give a straightforward interpretation of our experimental results on photogalvanic and photoelectrolytic cells in terms of the energy

levels of the electrodes and the electrolyte. This application of theoretical energy-level concepts,

discussed earlier in this workshop by Professor Gerischer and derived from both solid state physics

and electrochemistry, should greatly facilitate the search for new electrode materials by making it possible to specify what properties are necessary for efficient photoelectrolysis.

References

[1] Fujishima, A and Honda, K. , Nature _238, 37 (1972).

[2] Mavroides, J. G. , Tchernev, D. I., Kafalas, J. A., and Kolesar, D. F. , Mat. Res. Bull. 10, 1023 (1975).

[3] Mavroides, J. G., Kafalas, J. A., and Kolesar, D. F. , Appl. Phys. Lett, (to be published).

227

National Bureau of Standards Special Publication 455, Electrocatalysis on Non-Metallic Surfaces,

Proceedings of Workshop held at NBS, Gaithersburg , Md. , December 9-12, 1975. (Issued November 1976)

ELECTROCHEMICAL PROCESSES ON STRONTIUM TITANATE ELECTRODES

D. M. Tench and D. 0. Raleigh

Rockwell International/Science Center

Thousand Oaks, California 91360

The electrochemical properties of n-type single-crystal SrTiO-^

electrodes (lO-'-^-lO-'-^ carriers/cm-^) were investigated in acidic sulfate

solutions, with special attention to the hydrogen evolution reaction.

AC impedance measurements, together with the voltammetric behavior of

the ferro-ferricyanide couple and the onset of anodic photocurrent

,

established that the flat band potential, U^-j^ , lies about 250 mV cathodic

to the reversible hydrogen electrode potential.

Slow-sweep cyclic voltammetry with a rotating electrode assembly

established several hydrogen discharge mechanisms in moderately acidic

solutions. Anodic to U^^ , chemical desorption (H^^g + H^^^ H2) appears

to be rate limiting. At more cathodic potentials, electrochemical desorption (H30"*' + e~ + H^jjg ^ H2 + H2O) dominates. At still more cathodic

potentials, where proton discharge becomes diffusion-limited, hydrogen

discharge from H2O appears to dominate. Anomalous current behavior at low pH values (~0.5) suggests possible proton diffusion into the electrode.

ESCA measurements showed that the first few atomic layers of the SrTi03

surface, when brought in contact with aqueous solutions, become depleted

in Sr.

Key words: Semiconductor electrode; hydrogen evolution; perovskite

surface composition; oxide electrode; flat band potential; rotating

electrode voltammetry.

1. Introduction

Transition metal oxides of the ABO3 perovskite structure have attracted considerable

attention as possible catalysts and electrocatalysts . Rare-earth cobaltites , manganites

and ruthenates have shown catalytic activity for olefin hydrogenation [1]-^, CO oxidation [2]

and NO reduction [3,4]. Strontium-doped lanthanum cobaltite, La-j^.j^Sr^CoO^ , has been reported

to show reversible oxygen electrode behavior [5,6]. SrTiO^ electrodes have been used for

the photoelectroly tic decomposition of water, with considerably improved quantum efficiency

over Ti02 electrodes [7].

Previous theoretical work at the Science Center has indicated that the d-band perovskites

are unique among transition metal oxides in having a high concentration (10 cm ) of electronic surface states of almost pure d-orbital character, with the appropriate symmetry for

catalytic activity [8,9]. Strontium titanate has been studied extensively as a representative

d-band perovskite [8-13]. The pure compound is a broad-band insulator with a 3.2eV band gap, but may be heavily n-doped (~10^^ cm~^) by hydrogen reduction [14], resulting in specific

conductivities up to ~1 QT'^ cm"-'-. The carriers populate a pdTr* conduction band, with which

are associated a high concentration of surface d-band states with the t2g orbital symmetry,

located in a double-peaked energy distribution in the band gap [8]. Since surface states of

this symmetry have been proposed to be instrumental in the dissociative adsorption of common

diatomic gas molecules [8], it was of interest to determine the properties of n-SrTiO^ electrodes with respect to the hydrogen evolution reaction in aqueous media.

Figures in brackets indicate the literature references at the end of this paper.

229
2. Experimental

Single crystals of SrTiO^ from NL Industries, Inc. were cut into samples 1 mm thick and

4 mm square, having a 100 orientation (111 faces gave similar results, but more hysteresis).

These samples were then lapped on 600 grit SiC paper and etched in 85% H3PO4 for 15 min. at room temperature. Various concentrations of n-carriers were introduced by quenching in a 4%

H2 (in N2) reducing atmosphere from elevated temperatures (900-1150°C) . Ohmic electrical

contacts were prepared by electroless deposition of nickel on one side of the crystals, heat

treating at 400°C for ~l/2 hr. [15], then attaching a Pt lead wire using Ag-epoxy resin.

Each crystal was mounted flush with the end of a Kel-F cylinder by compression molding

at elevated temperatures. Electrical contact was made through a threaded brass cylinder

molded concentrically into the other end. This configuration minimizes the possibility of

contamination, permits the use of strong echants , and is well-suited for rotating electrode

studies [16]. Before use, the electrodes were polished on 3 pm alumina powder in water and

etched for ~l/2 hr. in 85% H^PO^ at 100°C to remove mechanically disrupted surface layers.

The main compartment of the Pyrex electrochemical cell contained 50-100 ml of electrolyte, and had an optical glass bottom. All cell connections were made via a gas-tight Teflon

top. A controlled atmosphere (argon, unless otherwise specified) could be maintained even with rotating electrodes, since the shaft clearance was less than 1 mm. Separate compart- ments contained a Pt counter electrode and a commercial saturated calomel reference electrode,

which was positioned several mm from the crystal via a Luggin capillary. For a.c. impedance

measurements, a 30 cm^ mercury pool in the main compartment was used for the counter electrode [17]. The external Pt counter electrode, however, gave comparable results. All electrolytes were prepared with reagent grade chemicals and de-ionized water.

For the current-potential measurements, a Princeton Applied Research (PAR) Model 173

potentiostat/galvanostat with a PAR Model 176 current follower was used in the three electrode

mode to control the electrode potential. Data were recorded on a Hewlett-Packard Model 7004B x-y recorder. The electrode rotation rate was controlled and monitored with a Motomatic Model

E-550-M controller and motor. A.C. impedance measurements were made using a direct biased

General Radio Type 1650-A impedance bridge with an external polystyrene decade capacitor and

an external function generator. Identical results within 5% were obtained by measuring the

current response to an a.c, signal with a PAR Model HR-8 lock-in amplifier. In both cases,

a signal of 15 mV peak to peak was used across the cell. Photocurrents were measured using

monochromatic light chopped at 90 Hz and phase sensitive detection.

3. Results and Discussion

3.1. Potential Distribution

To identify the potential region where sufficient charge carriers are available at the

semiconductor surface to sustain electrochemical reactions, the a.c. impedance characteristics

of the semiconductor-electrolyte interface were first studied. The flat band potential, , and bulk carrier concentration, N, are obtained from the intercept and slope of a plot of

vs. U, according to the well-known Mott-Schottky approximation [18] for exhaustive depletion of charge carriers

where Cg^ = semiconductor space charge capacitance, EgQ = dielectric constant in space charge

region, e = electronic charge, and f = e/kT.

Plots according to this equation are shown in figure 1 at pH 3.0 for SrTi03 electrodes

containing various concentrations of charge carriers. Frequency dispersion of capacitance

was found to be less than 5% from 100 to 1000 Hz. The carrier concentrations shown were

calculated using the bulk value [19] of 305 at 25°C for Eg^. Considerable curvature is

evident at more anodic potentials, especially for large N. This presumably results from

dielectric saturation, as has been observed for KTaO-j electrodes [20]. For each carrier

concentration, there is an appreciable linear region which extrapolates to a value of

=-0.42 V vs. standard hydrogen electrode (SHE). This good agreement for Uf^ over the

230
range of 10 -lO-*-' carriers /cm-' suggests that the capacitance of the Helmholtz layer is

sufficiently large ( ~ 20 ycoul/cm^) that its effect on the total electrode capacitance is

small [21]. The pH dependence of U^^ is -70 mV/pH in acidic solutions, and is somewhat

more in basic media, as illustrated in figure 2. Thus, Hf^ always lies about 250 mV cathodic

to the potential of the reversible hydrogen electrode in acid solutions. In this case, an appreciable contribution to the overvoltage for hydrogen evolution can be expected from the

depletion of charge carriers in the semiconductor space charge region. A small but significant difference in the Mott-Schottky slope in IN H2S0^ should be noted, which seems to be

related to a change in the electrode surface, as will be discussed below.

The potential distribution at the interface was checked by studying the slow sweep

(10 mV/sec) voltammetric behavior of the f erro-ferricyanide couple on rotating SrTiO^ electrodes in pH 3.0 and 7.0 potassium sulfate solutions. As expected, the anodic branch is

inhibited and cathodic current does not begin to flow until 300-400 mV anodic to the U^^

values indicated by the impedance studies. With increasingly negative voltage, the cathodic

currents rise with negligible hysteresis to diffusion limits, i^, plots of which exhibit

the expected [22] linear dependence on the square root of the rotation rate, intersecting

the origin. The slopes of these "Levich plots" yield a value of 6.2 xlO"^ cm^/sec for the

diffusion coefficient of the Fe (CN)^ species, which agrees well with the literature value

of 7.2xlO~6 cm^/sec [23]. Figure 3 shows Tafel plots of log [i/(i-|^-i)] vs. U which are

linear, except for minor deviations in the more acidic electrolyte, and exhibit slopes of

60-70 mV/decade.

231
0.08 —

0.06 —

0.04 —

0.02 -

Fig. 2. Plots of C"2 vs. U for SrTi03

(1.2 X 10l5 carriers/cm3) at 1000 Hz in argon-saturated 0.5 M

sulfate solutions of various pH

values

.

U vs. SHE (V)

Fig. 3. Plots of log[i/(iL-i) ] vs. U for

reduction of ferricyanide ion at

2500 rpm and 10 mV/sec on SrTiOs

19 . ^. . - _ (1.2 X 10 carriers/cm-^) in 0.5 M K2SO4 (pH 3.0 and 7.0) containing

2.5 X 10" 3 M K3Fe(CN)5 and K4Fe(CN)g;

11 = 1.80 mA/cin2 (pH 3.0) and 1.82

(pH 7.0)

.

O.Ol -

-0.3

U vs. SHE (V)

232
When back reaction and double layer effects may be neglected and the symmetry factor is

0.5, the current-overpotential relationship for a first-order process involving a one-electron

transfer, like ferricyanide reduction, may be written as [16]

where n = overpotential, y = fraction of the total electrode potential drop occurring across

the Helmholtz region, i = current at overpotential n, i^ = diffusion limiting current, and

ijj = exchange current. This equation assumes that the charge carrier concentration at the

interface follows the Boltzmann distribution. If y is a constant, it is evident that a plot

of log [i/ (i^-i) ] vs. U will yield a straight line of slope 59/(1-y/2) mV/decade. Thus, when

all of the electrode potential change occurs across the Helmholtz layer, y=1 and the Tafel

slope is 118 mV/decade. For the other extreme, when all of the potential change is across

the semiconductor space charge region, y = 0 and the Tafel slope is 59 mV/decade. Thus, it

appears that practically all of the potential change occurs across the space charge region

for SrTi03 (lO-^^-lO-'-^ carriers/cm^) at potentials anodic to the Uf^, values determined in the

impedance studies. A final confirmation of Uf^ for SrTi03 was obtained by noting the onset

of anodic photocurrent. This was observed to occur in pH 3.0 potassium sulfate solution at -0.40V vs. SHE, in good agreement with the impedance value of -0.42 V.

3.2. Proton Reduction Anodic to Uf^

Figure 4 shows a typical linear sweep voltammetric curve covering the hydrogen evolution

region at 10 mV/sec for a freshly prepared rotating SrTiO^ electrode in argon saturated 0.5 M

^2^*^4 2.9). The current scale on the left side is magnified 100-fold to better illustrate

the relatively small currents in this voltage region. No dependence of the current on charge

carrier concentration was observed from 4^10^^ to 3 x lO-'-^ carriers /cm-^. Two cathodic waves

are evident.

U vs. SHE (V)

Fig. 4. Linear sweep cyclic voltammogram at 10 mV/sec for SrTi03 (4.0 x lol^

carriers/cm3) at 3600 rpm in argon-saturated 0.5 M K2SO4 (pH 2.9).

233
The first wave occurs anodic to Vf^ and reaches a plateau in the yA range. This current plateau is observed to be rotation-rate dependent and is decreased by a factor of 10 by

bubbling H2 through the solution. Although this wave is difficult to reproduce exactly from

experiment to experiment, it does not seem to depend on pH (from pH 0.5 to 5.0). The sulfate

and K"*" species are not directly involved, since comparable results were obtained in 1.0 M

NaC104 CpH 3.0).

Levich plots of limiting current density vs. the square root of rotation rate, oj, are linear for this first wave, but intercept the abscissa at = 30 rpm'S without passing through

zero. The latter effect is probably caused by interference from the much larger second wave

and does not preclude a first order diffusion process. Tafel plots of log [i/Ci^-i)] vs. U for the first wave are linear with slopes of 40-55 mV/decade, as figure 5 illustrates.

These data are consistent with hydrogen evolution rate-limited by a chemical desorption

step (Hg^ijg + ^ads ^ ^2)» with the plateau resulting from full coverage of the electrode [24].

The higher Tafel slopes than theoretical (30 mV/decade) probably result from the large

fraction of the total potential drop occurring across the space charge region in this potential range, but could also be the result of interference from the larger second wave.

3. 3. Surface Coverage

The coverage of the electrode surface with adsorbed hydrogen (H^^jg) was investigated by

chronopotentiometry and by moderately fast linear sweep cyclic voltammetry in pH 3.0 sulfate

solutions. For the chronopotentiometric experiments, the electrode was pre-polarized cathodically with a constant current, then switched to a constant anodic current of reduced

magnitude, and the electrode potential followed as a function of time. As shown in figure

6, an inflection point was observed, which presumably corresponds to the depletion of H^^jg

by re-oxidation at the electrode surface. For pre-polarization currents of -67 to -330

pA/cm^ and subsequent anodic currents of 3.3 to 10 pA/cm2, the charge passed up to the

inflection point is always 260-340 iJcoul/cm2. Approximate monolayer coverage by adsorbed

hydrogen would seem to be Indicated.

234
0 50 1 00

t (sec)

Fig. 6. Chronopotentiometric curve of U vs. t for SrTi03 (1.2 x lO^^

carriers/cm^) in argon-saturated 0.5 M K2SO4 (pH 3.0) at +3.3

ViA/cm^ after pre-polarizing at -67 yA/cm2.

235
Figure 7 shows representative linear sweep cyclic voltairanetric curves at 100 mV/sec. For each curve, a broad anodic peak is observed slightly positive to SHE. The integrated

areas under these peaks (to +1.7 V where the sweep was reversed) presumably correspond to the charge required to re-oxidize H^^jg on the electrode surface. When the sweep is reversed

at -0.44 V, this charge is 325 pcoul/cm^ and seems to increase only slightly at more cathodic sweeps. For sweep reversals at -0.34, -0.25 and -0.15 V, the respective charges passed

are 235, 150 and 115 ycoul/cm^.

Thus, about 300 ycoul/cm2 of charge is required to re-oxidize H^^g in the vicinity of

the first plateau, and at least this amount at more negative potentials. If we assume one Hgjg per 0 or Ti site on a 100 surface, about 200 ycoul/cm^ of charge would correspond to full coverage. Therefore, it is possible that adsorption occurs on both 0 and Ti sites or

that more than one H^^g can occupy the same site. Alternatively, a surface roughness factor

of 1.5 could be responsible for the additional charge. In any case, it seems reasonable to assume that full coverage obtains at potentials cathodic to the beginning of the first

plateau.

3.4. Proton Reduction Cathodic to U^^

At voltages cathodic to the first wave in figure 4 and to U^^, the current rises to a second plateau, which appears to be limited by diffusion of protons to the surface. In this

potential range, linear Levich plots which pass through the origin are obtained, as exemplified in figure 8. The diffusion coefficient for protons, DjjH-, was calculated from their

slopes. Because of the relatively high values measured at low pH (e.g., 4.5^10"^ cm^/sec

at pH 3.0), Dy+ was determined in 0.1 M K2SO4 (pH 3.2) and found to be 1.7x10"'^ cm^/sec,

in agreement with the value reported by Breiter and Hoffmann [25].

A typical Tafel plot for the second wave of figure 4 is shown in figure 9 for the

anodic sweep. Since this wave occurs cathodic to Uj^ , all of the potential drop is assumed

to be across the Helmholtz layer, so that the normal kinetic treatments used for metal

electrodes are applicable. At all rotation rates to 4900 rpm, the Tafel slopes are 120-130

236
Fig. 9. Plot of log[i/iL-i)] vs. U for

proton reduction at 10 mV/sec

on SrTi03 (4.0 x iqIS carriers/

cm^) at 3600 rpm in argon-saturated 0.5 M K2SO4 (pH 2.9);

It = 17.5 mA.cin2 (anodic sweep).

-0.6 -0.7 -0.8 -0.9 -1.0

U vs. SHE (V)

mV/decade, with the smaller slopes occurring for the cathodic sweeps. At the full electrode

coverage assumed here, the only rate limiting step consistent with a 120 mV/decade Tafel

slope is electrochemical desorption [24]. This mechanism involves the discharge of a hydrated proton in solution at the Helmholtz plane, in concert with coalescence with an

adsorbed hydrogen atom to form H2 gas: H30"'' + e~ + H^^g H2(g) + H2O. Since this process

requires an appreciable potential drop across the Helmholtz region, it would not occur

readily anodic to U^^. At still more negative potentials (cathodic to the second plateau

of fig. 4), one expects hydrogen discharge via proton abstraction from H2O molecules:

H2O + e" -I- Hg^g ^ HjCg) + OH".

3.5. Dependence on pH

Figure 10 shows steady-state voltammetric sweeps at pH values from 0.5 to 5.2 in the

voltage range of the second wave of figure 4. The results obtained in the pH range 2 to 5 are consistent with the previously discussed behavior at pH « 3.0. At pH 2.3 a currentvoltage curve corresponding to the "pre-plateau" region of the second wave at pH 3.05 is obtained, with higher currents, as expected at this increased proton concentration. Plots of

log i vs. U at this pH were seen to be linear over two decades, with slopes of 120-130 mV/

decade. At pH 5.2, the proton diffusion limited wave of pH 3.05 is seen to be absent, presumably because of the lower proton concentration, and the onset of appreciate current

corresponds to proton abstraction from H2O.

237
At low pH values, anomalous current-voltage behavior is observed. The current is con- siderably lower than expected, and there is a pronounced "aging" effect. Curve A of figure

10 is for a freshly prepared electrode immersed in the electrolyte at pH 0.5. After several

minutes, a steady-state is reached, given by Curve B. Tafel plots in this case are reasona- bly linear, but have slopes in the 180-190 mV/decade range. In addition, at higher pH (>2.0),

many electrodes showed an order of magnitude current suppression during the first few voltammetric cycles after immersion in solution, but then recovered with time and approached the

behavior of fresh electrodes. These effects would appear to be caused by some change in the

nature of the electrode surface. Because of the apparent suppression of discharge from H30'^ at these pH values, one might hazard a guess that proton diffusion into the electrode is

involved, which could electrostatically suppress the discharge mechanism involving the charged

electrolyte species H^G"*". This, however, is mere speculation.

3.6. Surface Studies

To determine whether the bulk composition SrTiO^ is indeed the composition obtaining at the crystal surface, the surface Ti/Sr ratio for several crystals was measured by the ESCA

technique before and after various treatments. The results are shown in Table 1. A pulverized single crystal was used as a standard, to calibrate the Ti/Sr ESCA peak height

ratio. After quenching in 4% H2/N2, the surface is seen to be severely depleted in Sr,

probably reflecting evaporative loss at elevated temperatures, resulting in a rutile (Ti02)

surface. Mechanical polishing with 3 ym alumina powder and water reduces the Tl/Sr ratio

considerably. Etching in 85% H-:,P04 at 100°C for 30 min. reduces this ratio still more, whereas cycling a SrTi03 (~10l8 carriers/cm3) electrode between +0.2 and -1.5 V re SHE in

0.5 M K2SO4 (pH 2.9) produced no further reduction. This suggests that a somewhat Ti-rich

surface might be stable in aqueous or at least acidic media.

238
Table !• Surface Ti/Sr atomic ratios determined by ESCA for single

crystalline SrTi03 ( ~ 10-'-^ carriers/cm^) after hydrogen

reduction and various surface pretreatments

.

Sample Preparation Ti/Sr Atomic Ratio

Pulverized single crystal 1.00

Quenched in hydrogen only 2.01

Mechanically polished only 1.56

Etched only 1. 32 Etched after mechanically polishing 1.38

Electrochemically cycled after

polishing and etching 1.36

Freshly cleaved 1.06

Exposed to water after cleaving 1.44

To test the effect of water on the surface, a crystal was freshly cleaved in air and

quickly transferred into the vacuum chamber. The Ti/Sr ratio was similar to those observed

for pulverized samples. After exposing this cleaved surface to de-ionized water for 2 min.

,

however, the Ti/Sr ratio (1.44) was found to be practically the same as for etched crystals

(1.3-1.4).

These results show that the SrTi03 surface becomes severely depleted in Sr when it is

brought into contact with water or aqueous electrolytes. If we consider an escape depth for

ESCA electrons from SrTi03 of 1.8-2.0 nm, the Ti/Sr ratio observed (1.3-1.4) would correspond

to total removal of Sr in the interface region, leaving two atomic layers of only titanium

and oxide ions at the surface, or some equivalent distribution. As a first guess, this

composition could reflect the good solubility of SrO in aqueous media. The resultant surface

composition might either reflect a Ti02 structure, or else the perovskite structure might be

retained, with protons replacing Sr ions.

The results raise serious questions about the validity of extrapolating bulk properties

to electrode surfaces, not only for SrTi03, but for perovskites in general. At least in the case of SrTi03, the nature and availability of surface states is uncertain. Also, as we have

Suggested earlier, the structural nature of the electrode surface may be potential dependent,

accounting for some of the anomalies reported in this work.

4. Summary and Conclusions

For SrTi03 electrodes, Uf^ is always about 250 mV negative of the reversible hydrogen

electrode potential in the same electrolyte. At potentials anodic to Uf^ , hydrogen evolution

proceeds only slowly, with chemical desorption (H^^g + H^jg H2) as the rate limiting step.

This would result from the fact that most of the electrode potential change occurs across

the semiconductor space charge region, so is not available to drive the electrochemical

reaction. At full electrode coverage by H^^^g , a current plateau is observed in the pA range.

This plateau is important for theoretical reasons, but has not previously been unambiguously

observed [26]. This is because, for metal electrodes, practically all of the electrode

potential is usually available to drive electrochemical reactions over the whole available

voltage range.

At potentials cathodic to Uf^ on SrTi03, electrochemical desorption (H"*" + H^^g + e~ ^ H2)

becomes the rate limiting step for the hydrogen evolution reaction. This change in mechanism

occurs because, cathodic to Uf^ , a sufficient fraction of the electrode potential appears

across the Helmholtz layer to drive the electrochemical desorption reaction.

Thus, it is clear that a perovskite with a 15 value at least 250 mV more positive is

required to lower the overpotential for the hydrogen evolution reaction. In this case, electrochemical desorption should become the rate limiting step at the reversible hydrogen

239
potential. A material on which hydrogen is more weakly adsorbed would also be desirable,

since this should increase the rates of both the chemical and electrochemical desorption

steps. Finally, since the surface structure of SrTi03 has been shown to be altered by

aqueous solutions, care should be exercised when dealing with other perovskites.

5. Acknowledgments

The authors express their appreciation to Mr. Cameron Ogden for performing many of these

experiments and to Dr. Leonard Szalkowski for carrying out the ESCA measurements.

6. References

Pedersen, L. A. and Libby, W. F. , Science JL76, 1355 (1972).

Voorhoeve, R. J.H. , Remeika, J. P., Freeland, P. E. and Matthias, B. T. , ibid. Ill , 353

(1972).

Voorhoeve, R. J.H. , Remeika, J. P. and Johnson, Jr., D. W. , ibid. 180 , 62 (1973).

Voorhoeve, R.J.H. , Remeika, J. P. and Trimble, L. E. , Mat. Res. Bull. 9_, 1393 (1974).

Meadowcroft, D. B., Nature 226, 847 (1970).

Tseung, A. C. C. and Bevan, H. L. , J. Electroanal. Chem. 45^, 429 (1973).

Mavroides, J. G., et al . , "Photoelectrolysis of Water", p. 221, These Proceedings.

Morin, F. J. and Wolfram, T. , Phys . Rev. Lett. 30, 1214 (1973).

Wolfram, T. and Morin, F. J., Appl. Phys. 8^, 125 (1975).

Wolfram, T. , Kraut, E. A., and Morin, F. J., Phys. Rev. B7, 1677 (1973).

Kraut, E. A., Wolfram, T. and Hall, W. F. , Phys. Rev. B6. 1499 (1972).

Wolfram, T. , Phys. Rev. Lett 29, 1383 (1972).

Morin, F. J. and Oliver, J. R. , Phys. Rev. B8, 5847 (1973).

Yamada, H. and Miller, G. R., J. Solid State Chem. 169 (1973).

Turner, D. R. and Sauer, H. A., J. Electrochem. Soc. 107, 250 (1960).

Tench, D. M. and Yeager, E., J. Electrochem. Soc. 121, 318 (1974).

Tench, D. M. and Yeager, E. J. Electrochem. Soc. 120, 164 (1973).

Schottky, W., Z. Phys ik 113^, 367 (1939); Mott, N. F. , Proc. Roy. Soc. (London)

A171, 27 (1939).

Linz, A., Phys. Rev. 91, 753 (1953).

Boddy, P. J., Kahng, D. and Chen, Y. S., Electrochim. Acta 13, 1311 (1968). De Gryse, R. , Gomes, W. P., Cardon, F. and Vennik, J., J. Electrochem. Soc. 122 , 711 (1975).

Riddiford, A. C, Adv. Electrochem. and Electrochem. Eng. 4_, 76 (1966).

Jahn, D. and Vielstich, W. , J. Electrochem. Soc. 109, 849 (1962).

Vetter, K. J., Electrochemical Kinetics, Academic Press, New York, p. 516-614, (1967).

Breiter, M. and Hoffmann, K. , Z. Elektrochem. 64_, 462 (1960).

Vetter, K. J., loc . cit., p. 521.

240
Hational Bureau of Standards Special Publication 455, Electrocatalysis on Non-Metallic Surfaces

,

Proceedings of Workshop held at NBS, Gaithersburg, Md., December 9-12, 1975. (Issued November 1976)

ELECTROCHEMICAL ACTIVITY OF POLYMERIC AND MONOMERIC IRON PHTHALOCYANINES

AS A FUNCTION OF THEIR SPECTRAL PROPERTIES

A. J. Appleby and M. Savy

Laboratoire d 'Electrolyse du C.N.R.S.

92-Bellevue, France.

Samples of monomer and polymer iron phthalocyanines , prepared both

in situ on carbon supports or in the form of thin films, using a gas-phase

synthesis method, and by dissolution of liquid-phase preparations in

sulfuric acid followed by reprecipitation on carbon, were examined by

optical and Mossbauer spectroscopy. Their electrochemical activities for

oxygen reduction were examined under steady-state conditions using the

ultra-thin pxjrous electrode technique, as a function of O2 partial

pressure, pH (11.8 - 16.4) and H2O2 additions. In addition, rotating

ring-disk electrode experiments were conducted. Analysis of reaction

orders and comparison of activities with those of the support have allowed

the most probable mechanism to be determined, in which the rate-determining step is a chemical dissociation of the 0-0 bond. Correlations between electrochemical activities and spectral properties indicate the role

of polymerization and of the spin configuration of the central iron atom

in promoting catalysis. These effects result from reoxidation of FelliS

to FemHS or iS, which is favorized by interaction between the polymer tt

I electrons and the d electrons of the iron.

Key words: Electrocatalysis; electronic structure; hydrogen peroxide;

iron phthalocyanines; Mossbauer; optical spectra; oxygen;

reaction mechanisms; ring-disk electrodes.

1. Introduction

Study of the influence of electronic parameters of the active surface in heterogeneous

tcatalysis is of great theoretical and practical interest. In electrochemistry, the major

I process of technical importance, and one of the slowest, is the reduction of molecular oxyr=v. to water in a four-electron process. This reaction has received extensive study on

:'.al surfaces, and, following previous work on hydrogen evolution [1]^, many correlations

C5tween properties related to substrate electronic parameters (e.g., d-orbital vacancies,

llheats of adsorption and sublimation) and reaction rates have been proposed. This is particularly true for oxygen reduction on the noble metals in acid solution [2] .

It is well known that most catalytic processes are characterized by relations of volcano

:/pe [3,4,5], in which reaction rates as a function of bond-strength reach a maximum value.

I In the most simple cases, involving reversible adsorption of reaction intermediates and no 'simultaneous irreversible adsorption of other species, the maximum corresponds to a site

-cupation of 0.5 by adsorbed reaction intermediates, either as reactants or products in the

- ite-determining step [4]. For such cases, it is easy to show that reduction of rate as ibond-strength increases (on the falling side of the volcano) results from a diminution of

available sites and from the fact 'that desorption becomes rate-determining [4]. On the rising side of the volcano, adsorption is rate-determining. Thus, in the most general case, a certain minimum bond strength for chemisorption of the reaction intermediates is required

tf or maximum rates

.

"igures in brackets indicate the literature references at the end of this paper.

241
Formation of this 'most favorable' bond between a reactant or product and the substrate

will require as a first condition sites presenting partially filled MO levels of the same energy as those of the intermediate species [6]. The existence of this favorable bonding

will also depend on the wave function and steric availability of the d-electron cloud at the

site, which may be modified by the electronic properties of the surrounding environment.

Such ideas are difficult to test in the case of metallic electrodes, owing to the problem of

the precise definition of surface states. However, stable chelate compounds of well-defined

structure and composition, whose electronic properties can be determined, are very suitable

for fundamental studies of catalysis. The effect of MO energy levels on catalytic processes

can be determined by examining the effect of the nature and degree of oxidation of the central ion; similarly, the effect of d-electron density may be demonstrated by changing the it electron properties of the surrounding organic structure. This paper describes the identification of the electronic characteristics of mono- and polymer iron phthalocyanines, and

their effect on the electrocatalytic reduction of oxygen in alkaline solution. A simmary of

previous work on oxygen reduction on these substrates has been given elsewhere [7] .

In previous work [8,9], it was found that the increase in the number of unpaired electtrons, detected by optical and ESCA spectroscopy, was accompanied by an increase in the

electrochemical activity for the reduction of oxygen on thin-layer phthalocyanine monomers

in both acid and alkaline solution. However, the techniques used in this work were not able

to differentiate between the activities of different lattice sites in terms of their oxidation states and spin configurations. In the present work, we have made these correlations

using a combination of optical reflectance spectroscopy, Mftssbauer spectroscopy, and electrochemical experiments on supported monomeric and polymeric iron phthalocyanines prepared by

gas-phase and liquid-phase methods [8-10]

.

The mechanism of the oxygen reduction process on the polymer iron phthalocyanine was determined using two types of electrochemical experiments on carbon supported material, viz.

a) - Steady-state measurements as a function of pH in the alkaline range and as a fimction

of oxygen partial pressure. The apparatus, ultra- thin porous electrode preparation and

measurement techniques have been previously described [8-10]

.

b) -_Rotating ring-disk electrode measurements, to detect the presence of free intermediates

(O^H ion) produced in the reaction.

A certain number of experiments were performed on Eastman Kodak monomeric iron phthalocyanine supported on carbon black, prepared by precipitation from concentrated sulfuric acid

solution [10] . Blank experiments on the support were also conducted.

2. Experimental

2.1 Optical Spectra

Three types of iron phthalocyanines were examined:

1) - Thin vacuum-deposited films of blue ( electrochemically inactive) and green (active)

forms of the monomer carried on gold foils. Preparation is given in Refs. 8,9.

2) - Iron polyphthalocyanines, consisting mainly of the dimer [9] in \Qffo by weight loading

on acetylene black (Lannemezan y), active carbon (Norit BEX) and high-surface area ground

graphite supports. The phthalocyanines were prepared in situ using a gas-phase technique

[10]. . . 3) - Pure mono (Eastman Kodak) and gas-phase and liquid- phase polymer [10] preparations in

the form of powders.

Transmission spectra (for the thin film materials) and reflection spectra (for the powders) were obtained both in the dry state and in contact with IN KOH solution, both in the

presence of air.

2.2 MBssbauer Spectra

MBssbauer spectra for monomeric and polymeric powders were obtained using a Co/Cu

source at 293°K, and two different absorber temperatures (293° and 77.4°K). In separate ex- periments, thin film monomers similar to those used in the optical spectra experiments, but

in this case vacuum-deposited onto Mylar instead of onto gold foils, were examined. For

242
these materials, spectra were obtained only at liquid nitrogen temperature, since at room

, temperature absorption proved to be insufficient. Experimental details have been given

elsewhere [7]

.

2.3 Electrochemical Measurements

Steady-state results: Steady-state measurements were performed on ultra-thin porous

electrodes of the type first describe^ by Vogel and Lundquist [11]. These electrodes, which

had catalyst weights in the 0.1 mg/cm range, allow pure kinetic measurements to be made on finely-powdered electrode materials without diffusion limitations [10]. Experimental techniques have been given elsewhere [10] . Experimental results obtained on phthalocyanines

supported on acetylene black and high-surface-area graphite materials were compared with

' those for the pure supports. All preparations (gas-phase or liquid-phase reprecipitated

from concentrated sulfuric acid) were in 10% by weight loading. Polarization - log cur^j^ent

1 density resiilts at 1 atm. 0^ partial pressure were obtained in IN K CO (pH 11 .8), IN K

; carbonate buffer (pH 13.0), IN EOH (pH 13.9), and in a series of KOH solutions varying from

2N to 12N concentration (pH 14.2 to 16. 4). pH values in the latter solutions were calculj ated from activity coefficient values given in Ref . 1 2^ Results were also obtained over a range of oxygen partial pressures varying from 5 x 10 atm. to 1 atm. in 2N KOH. Finally,

the effect of the addition of hydrogen peroxide on the rest potentials of the electrodes and

• on steady-state polarization results_'^as determined in t]jie presence of oxygen at 1 atm. partial pressure. Concentrations of 10 M, 10 M and 10 M were examined, ^^^sults were

compared with those obtained on a pure carbon material (Norit BRX, 10 mg/cm ). All experi-ents were conducted at 25° C.

Ring-disk electrode experiments: A ring-disk electrode (Tacussel) with a cavity shapi ed in the gold disk was used [13] . The ring was gold with a gold-black electrodeposit. The

; iron polyphthalocyanine on carbon black support was introduced into the cavity in the form

of a paste containing purified mineral oil [13] , after which the surface was accurately

smoothed. Collection efficiencv was determined to be 0.17. These experiments were conductI ed at ambient temperature (20°c).

3. Results and Discussion

3.1 Spectra

The optical spectra obtained are shown in Figs 1a,b,c. The most striking observation

involves the change in the structure of the spectra on wetting the materials with KOH in the

presence of oxygen. In the latter case, only two spectral peaks are observed, which are identical for all the phthalocyanines studied.

Curve 2 in Fig. lb represents the reflection spectrum obtained on the monomer gas-phase

: preparation powder. The position of the two main peaks on this curve shows^Jhem to be the

\ same as the two similar peaks on curve 1 . They show the presence of the Fe triplet state

( (see below). After drying in air at room temperature, the spectra were identical to those

:btained before wetting. Since a transformation involving the bulk phthalocyanine would be

r^likely to be reversible, and can in any. case be excluded for the reflectance spectrum

data, we conclude that wetting results in a modification of the oxidation state and spin

I' configuration of the outer monolayers of the phthalocyanine. Reflection spectra for monoeric powders (Eastman Kodak and gas-phase preparations) are shown in Fig. 1c. Spectra for

:oth monomers are identical. Globally, a shift of the principal absorption peaks towards

longer wavelengths is observed, in the order:

Liquid-phase preparation > Gas-phase ^ Monomer

From ESCA data [8,9] , the positij^^of the main absorption peak for the liquid-phase

polymer preparation^jorresponds to Fe HS. The corresponding peak in the case of the monomer is that for Fe iS. Since the sl^ft to lon^^j wavelengths increases with increasing

spin [8,9], we conclude that either Fe HS or Fe iS must be present in the gas-phase

preparation, in agreement with the observed order of absorption peak wavelength, i.e.

Fe"^ HS(5/2)>Fe^^ HS(2), Fe^^^ iS(3/2)>Fe" iS(l)

243
t a

3 4 5 6 7

3 I 5 6 7

3 X 5 6 7

. 1. Optical reflection spectra for phthalocyanines. a) Polymer gas-phase preparation: curve 1, air; curve 2, air + 1 N

KOH. b) Monomer Eastman Kodak and gasphase preparation: curve 1, air + 1 N KOH; curve 2, air; curve 3, thin layer

active form (transmission), c) Monomer

and polymer in air: curve 1, monomer

gas-phase and Eastman preparations;

curve 2, polymer gas-phase preparations;

curve 3, polymer liquid-phase preparation.

Ordinate: nanometres x 10~2,

Mttssbauer data (Quadrupole splitting, QS, and Isomer shift, IS^) are summarized in

Tahle 1 . QS values are oxidation-number independent and are normally on the order of 0.5

(in mm/s units) [14]: higher values indicate low ligand symmetry, e.g. C^-^- . The IS values

may he compared with literature data for identification of the species present. These are indicated in the Table, and are discussed in detaij^jn Ref. 7. The major conclusions are thjjjthe polymeric materials consist largely of Fe HS (liquid-phase preparation) and Fe iS (gas-phase preparation). The inactive monomer, on the other hand, consists of Fe (triplet state). These are in agreement with the tentative conclusions of the optical data.

However, comparison with the opticaj^spectra obtained after wetting shows that the presumed

inactive species in the monomer (Fe triplet) is also present at the surface (but not in

the interior) of the polymers.

3.2 Electrochemical Data

The effect of pH (v. NHE) is shown in Fig. 2. The polarization results obtained showed

no hysteresis and a reproducibility to within JmV at the same current density (outside of

the diffusion region) for different samples of the same preparation. Over most of the pH

range studied, the initial Tafel slope was close to 30mV/decade„ This was followed by a section of smoothly increasing slope leading into the diffuision region. In highly-concentrated KOH solutions, the Tafel-like regions were much shorter than in 1 or 2N KOH (for the

latter on the order of 2 decades), since oxygen solubility falls off rapidly with increasing concentration. In N K^CO , the slope observed was somewhat higher, perhaps suggesting a change in reaction mechanism at lower pH (see below).

Upper-case I is used to distinguish isomer shift (IS from intermediate spin (iS)

.

244
Table 1. Hossbauer data, 57(;o(Cu) source.

Monomer (90 K)

. a

= 1

0.18 1.16 0.16 B 0.54

15 0.12 0.23 0.01 1.62 -0.27 .83

Monomer (298 K) 0.14 1.24 0.10 0.57 0.07 0.25 -0.02 1 .76 0.39 1.09

Foil blue (90 K) 0.16 1.13 -0.11 0.70? -0.03 0.23 ?

Foil green (90 K) 0.24 1.22 -0.04

28 0.55 0.32 0.36

Triplet Fe2+ Fe3+ HS or Fe2+ LS

Fe-''" green HS Fe3+ LS

Polymer (298 K)

Polymer (90 K)

0.004

0.03

0.90 -0.10

0.83 -0.06

1.35 0.44

1.24 0.26

0.66

0.49

Fe^+ iS monomer

Triplet Fe^*

Fe5+ HS

Units 6. and a. are in millimeters.

Indicates percentage of each form present.

iRtential (v/n.KE)

10 wt % gas-phase polyphthalocyanine

on acetylene black: V—log i curves

as a function of pH (v. NHE—curves below are v. hydrogen in same solution), a-h: pH 11.8, 13.0, 13.9,

14.2, 14.8, 15.2, 15.8, 16.4 respectively.

The effect of oxygen partial pressure in 2E KOH is shown in Pig. 3a. Fig. 5b gives a log-log reaction-order plot for current densities in the Tafel range (925iiiV v. hydrogen in

the same solution). The experimental reaction order was 1.08. This value was used to correct the current density data at constant potential ir. Fig. 2 for the effect of oxygen solubility, so that the reaction order for 0H~ could be determined. This required a consideration of the change in water activity with KOH concentration [15], both on the kinetics of

the process (since one molecule of water is required for each 0H~ liberated), and on the H

potential scale used. A complete analysis has heen given (?). It_was concluded that the ^ most probable chemical reaction orders were +2 in HO and -^1 in OH . V - log i arations (monomer

plots at 1 atm. 0^ partial pressure in 6N KOH for all the phthalocyanine pre- gas- and liquxd- phase polymer), together with the untreated carbon supp245
Fig. 3. 10 wt % gas-phase polyphthalocyanine on acetylene black: left, V—log i curves as a function of O2 partial pressure (left to right, 0.005, 0.02, 0.053, 0.02 and 1 atm)

;

right, reaction order plot (925 mV) . Slope = 1.08.

ort, are shown in Fig. 4a. It is evident that the polymeric phthalocyanines confer activity

muck greater than that of the support in the non-diffusion-limited region (see Fig. 4b). In

contrast, the effect of the monomer is largely to mask the intrinsic support activity. The

liquid-phase polymer preparation shows an activity somewhat less than that of the gas-phase

species

.

T) (mV)

.000

800

600

AOO

Potential (mV/H.E.)

Log,o i(A/mg)

50

AO

30

20

10

i (pA/mg)

100 1000

Fig. 4. Oxygen reduction in 6N KOH: left, 10 % loadings on acetylene black: a, monomer;

b, support alone; c, liquid-phase polymer; d, gas-phase polymer; right, a, support

alone; b, gas-phase polymer (10 wt % on 400 m2/g graphite dust).

The effect of hydrogen peroxide additions in the presence of oxygen (1 atm.) is shown

in Fig. 5a for phthalocyanine electrodes, and in Fig. 5b for electrodes of Norit BRX carbon

black. In both cases, hydrogen peroxide is reduced on the material, and the current increases linearly with ^2^2 However, the fact that hydrogen peroxide additions

do not depress the rest potential of the phthalocyanine material is remarkable, and implies

that the O^/H^O^ couple is not reversible in this case (i.e., that ^2*^2 """^ oxidized on the phthalocyanine surface).

246
Fig. 5. V—log i curves in presence of H2O2 at 1 atm O2 . H2O2 concentrations are (left to right, in diffusion region): gas-phase phthalocyanine (left), 0.1, 0.01 M and zero (dashed curve is nickel support); BRX carbon (right), 0.01, 0.1 and 1.0 M.

Ring-disk electrode data in carbonate buffer and KOH solutions are given in Fig. 6a, b. Only negligible quantities {-C'\'f°) of HO^ are formed above 600mV v. hydrogen. The larger

quantities (increasing with pH up to ^'3.6fo at 200mV in 5N KOH) at low potentials are perhaps

due to direct reduction of oxygen on uncovered parts of the carbon black support. This is in accord with the horizontal disk current/ring current vs bT'^ plot shown in Fig. 7, which

indicates a reaction product that is stable in contact with the catalyst [16] . This cannot

be the case for phthalocyanines, which decompose HO^ rapidly. The half-wave potential in

N KOH is independent of rotation rate (Fig. 6b), so the reaction is activation controlled.

Ir (jjA)

Fig. 6. Ring-disk electrode data (collection efficiency = 0.17); Ir = ring current; Ij) = disk

current: left, results as function of concentration at 2400 rpm; a - c, 5, 2.5 and 1 N KOH; d, N carbonate buffer, pH 13.3; right, N KOH; a - d, 900, 1200, 1800 and 2400

rpm respectively. Gas-phase polyphthalocyanine on acetylene black, 1 atm O2.

1r

Fig. 7. Iq/Ir as a function of root rpm (oJ'S) Conditions as in figure 6 (left)

.

20 30 40

247
3.3 Reaction Mechanism On Polyphthalocyanines

The data in Pigs. 2 and 3» obtained on gas-phase iron polyphthalocyanine preparations,

show that the reaction is first order in 0^, and probably of order -1 in 0H~. An order of

+2 in H^O is probable. The Tafel slope is low: about 30mV/decade, or Rt/2F. The value of a, defined as in Ref. 21, is therefore close to 2.0. The reaction is not controlled by the

diffusion of a dilute product from the electrode surface, as demonstrated by the rotatingdisk electrode half-wave potential data, a will normally be given by the expression [21] :

a = n/v + 3n' 0n' (1

)

where n is the number of electrons reversibly transferred before the rate-determining step,

whose stoichiometric number is O, and for which the symmetry factor is 3* ^' values other

than 0 (chemical process) and 1 (electrochemical process) may be excluded on activation en- ergy grounds [22]

.

It is important to point out that Equation 1 will only apply to metallic or degenerate

semiconductor electrodes for 1 - 1 coupled process [17] , for reactions involving complete

transfer of charge across the double-layer in the rate-determining-step (r.d.s.). Three

other possible limiting cases have been discussed in the literature:

a) The presence of a purely resistive or barrier-layer film on the electrode surface [18]'

.

b) Charge-transfer across the wide-gap bulk semiconductor-solution interface [19] . c) Reactions involving adsorption of reaction intermediates at some point within the Helmholtz double layer, so that charge is transferred across less than the bulk material - solution potential difference in the r.d.s., in previous steps, or both [20]. This effect may

be in addition to those described in a) and b) above.

In case a), a second charge- transfer barrier occurs across the potential drop region in

the film itself. This leads to an effective 3 value whose reciprocal is equal to the sum of

the reciprocals for double layer and film transfer, so that very low apparent symmetry factors and high Tafel slopes result [18]. For case b), 3 will be effectively equal to unity

for a cathodic process involving electron transfer from the conduction band (normally the

predominant process at an intrinsic or n-type semiconductor), and zero for hole exchange to

the valence band [19]. In both cases, the initial term will remain the same in Equation 1 (for 1 - 1 coupled processes), since it follows from the change in electrochemical potential

in passing from the process reactants in their initial state (chemical species in the bulk

solution, electrons at the Fermi level) to the state following the r.d.s. In general* case

a) will always lead to higher Tafel slopes than those predicted from Equ. 1 , whereas case b)

will normally lead to slopes that are lower. As Gnanamuthu and Petrocelli have shown, the

slopes for case c) will always be higher than predicted from Equ. 1 [20] .

In the presen^ work, was determined to be close to 2 in the current density range between 10 and 10 A/mg, in particular in IN and 2N KOH. If for the moment we assume that

Equ. 1 applies (1 - 1 coupled processes under Langmuir conditions), then with a = 2, possible values of n/i7 are:

n/v = 2; n' - 0, for all 3 values (chemical r.d.s.; all cases except c)

n/o =1; n' =1, for 3 equal to 1 (semiconducting electrodes, case b)

.

Other possibilities certainly exist. However, processes involving cases a) and c) are improbable, owing to the very low observed Tafel slope: indeed, if case a) applies, electron

transfer of primary type vrill be expected to be rate-determining (slope >2RT/F) . Other

possible types of mechanism (e.g., processes other than those of 1 - 1 coupled type, such as mechanisms III - V in Ref. 20) can be eliminated on reaction order grounds. Reactions occurring under high-coverage Langmuir conditions can be similarly rejected, since they will nor- mally be zero-order in 0 [2] . Temkin-type processes are formally possible, and result in

higher effective values of 3(e.g, 3 = 1 on metallic electrodes, c.f. Refs. 2,22). However,

they imply fractional orders in 0H~[2,22i. We may therefore conclude that the most probable

mechanism at low current density in the strongly alkaline range involves two electron transfers before a chemical r.d.s., or one electron transfer before an electron transfer r.d.s.,

provided, in the latter case, that the substrate can be regarded as a typical bulk semiconductor.

248
We siaggest the following possible process:

°2 °2 ads ^^^P^^^ I

°2 ads =2° ^°2 ads ^^^P^*^)

ads ^ ^2° V2 ads ^

^2°2 ads 2 OH lY

in wMcti step (or possibly step III, for bulk semiconducting electrodes) is the r.d.s.

Electron microscope studies have shown [9] that in 10?^ loadings on acetylene black, phthalocyanines consist of relatively uniform 20-30i thick layers, whose surface area is somewhat

less than that of the support [10] . A semi-infinite linear surface model is therefore proposed, in whjjti elejjjon transfer from the carbon to the surface takes place via a hopping

process (Fe

^ Fe + e ), which, at least at lower current densities, is not rate-limiting. Case b) (step III rate-determining), which can only apply on bulk semiconductors whose

thicknesses are greater than the Debye length, can therefore be excluded. A rate-determining chemical step is therefore most probable. Since H^O^ in step III is not likely in

strongly alkaline solution, the most probable process xs I and II followed by:

^°2 ads + ^'-^ ^°2 ads _ ^^^P^^^ ^

2°2 ads V 2 OH + OH- (r.d.s.) VI

An alternative sequence may involve disproportionation as the r.d.s. (c.f. Ref. 23), i.e., V

followed by:

HO" + OH , » OH" + H0„ ^ VII

2 ads ads 2 ads

2 2°2 ads 2 OH + 0^ VIII

in which either VII or VIII is rate-determining, and in which 0 is recycled. At low pH

(11.8), the mechanism may change^^P^^^^Ps "to a second electron transfer step. Similarly, at higher current densities than 10 A/mg, the slow bending of the V - log i curve (compare

plot for active carbon. Fig. 4a) implies a change in r.d.s., which probably becomes electron

transfer through the phthalocyanine film.

Finally, it has been suggested iron phthalocyanines serve only to decompose peroxide

produced by reduction of oxygen on the carbon support [24] . This is unlikely, in view of

direct electron microscope evidence [9]/ and of the smaller limiting currents in the presence of phthalocyanine, indicating a masking effect [10] . As Goldstein and Tseung [25] have

indicated, a simple reaction scheme involving decomposition will give a maximum of only a factor of 2 increase over the intrinsic carbon activity at the same overpotential. This is

confirmed by a more complex model [26] , which also shows that Tafel slopes will be expected

to be similar on coated and uncoated supports. In practice, the enhancement factor will be

less than 2, due to the masking effect of the catalyst and the fact that the efficiency of

H^O^ production on carbon is not 100% [13]. Fig. 4b indicates a factor of ca 10 increase in

activity, together with a change in Tafel slope. The hypothesis given in Ref. 24 can therefore be rejected.

4. Correlation with Spectroscopic Data

As indicated above, the oxidation state of the iron and its spin configuration determine the rate of electrosorption of molecular oxygen and of subsequent electron-transfer

steps. Hoxever, reflection spectrj^results in KOH and oxygen show that the interface consists of a low-activity species (Fe triplet). If we compare the present data with that

previously given [10], it is clear that the polymeric forms are still much more active than

the monomer, which consists largely of the above iron species. MBssbauer and reflection

spectral data on the Jjy polymer show Fe iS to be the predominant species in the gas-i phase preparation (Pe HS in the liquid-phase material). These species can be presumed

to be present inside the wetted polymer. There is no contradiction between the catalytic

theory and the experimental data if it is assumed that initial absorption of oxygen molecules occu2?s within the bulk of the catalyst layer, producing oxygenated radicals which

249
react at the interface after migration. An initial reaction pa±ii»- consisting first of a chemisorption step activated by the iron spin configuration (Pe - 0^, iS or HS) in the

bulk material, followed by reaction with water at the interface (corresponding to steps I and II), can be proposed:

0 + PPc Pe-^-"-^ - 0" —> 0* - PPc Pe-"--^"^ - o" (in bulk)

TTT TT — PPc Pe - ^2 ^2° —^

^^'^ ~ (interface)

These are folljy^d at the injjjface by III and subsequent steps given above. The increased

activity of Pe iS over Pe HS (gas- and liquid-phase preparations) may perhaps be explained by the greater ease of electron transfer^j^ the absence of a change of spin. The

effect of polymerization is to stabilize the Pe species, via ir-d electron interaction.

This lowers the activation energy for electron transfer, as shown by the fact that oxygen

electrosorption appears to be rate-determining on the monomer [8]

.

5 . Conclusions

Iron phthalocyanines have been studied by MUssbauer and optical reflectance spectroscopy and by steady-state and rotating-disk electrochemical techniques. While reduction of

molecular oxygen on the monomer materials in alkaline solution is relatively slow, an increase in activity by over an order of magnitude is seen on polymer samples, particularly

those prepared in situ on carbon supports by a gas-phase method. The MBssbauer an^^yptical

data indicate that this increase in activity is associated with the presence of Pe iS or HS sites within the polymer. An increase in the electron transfer rate appears to occur in

parallel with an increase in oxygen adsorption rate on the polymer. Oxygen is reduced in

a direct or pseudo- 4-electron process, and the most probable r.d.s. is the breaking of the

0-0 bond, either directly or via a disproportionation process, taking place after two rapid electron transfers. Since optical reflection spectra in the^jresence of oxygenated

KOH are similar for both monomer and polymer, and correspond to Pe iS, we suggest that

adsorption occurs initially in the bulk material, the oxygenated radicals thereafter diffusing to the interface. This agrees with previous work showing an increase in activity

with increasing thickness for monomer films [8,9] , for which absorption is rate-determining.

The enhanced activity of polymers may be accounted for by the increased d-electron density

resulting from tt electron overlap in the conjugated rings. This increases both heats of

adsorption of oxygenated radicals and electron transfer rates.

6. References

Bockris, J. O'M. and Wroblowa, H, J. Electroanal. Chem. 2, 428 (1964).

Appleby, A. J., in 'Modem Aspects of Electrochemistry, Vol. 9' (J. O'M. Bockris, Ed.)

Plenum, New York 369 (1974).

Balandin, A. A., 'The Problems of Chemical Kinetics, Catalysis and Reactivity', Acad.

of Sciences, Moscow (l955).

Gerischer, H., Bull. Chim. Soc. Beige 61, 506 (l958).

Parsons, R. , Trans. Paraday Soc. (London) 1053 (l958).

Savy, M., Andro, P., Bernard, C. and Magner, G., Electrochim. Acta 18, 191 (1973);

Alt, H., Binder, H. and Sanstede, G., J. Catalysis 28, 8 (l973).

Appleby, A. J., Pleisch, J. and Savy, M. , J. Catalysis, to be published.

Savy, M., Bernard, C. and Magner, G., Electrochim. Acta 20, 383 (l975).

Savy, M., Andro, P. aad Bernard, C, Ibid. 12., 403 (197417

Appleby, A. J and Savy, M., Ibid., to be published.

Vogel, W. M. and Lundquist, J. L., J. Electrochem. Soc. m., 1512 (1970).

'Sprovochnik Khimika, Vol. 3', Khimiya, Moscow 595 (1964)

Appel, M., Thesis, Conservatoire National des Arts et Metiers, Paris (l974).

Greenwood, N. N. and Gibb, T. C, 'Mttssbauer Spectroscopy', Chapman and Hall, London

(1971).

Bro, P. and Kang, H. Y., J. Electrochem. Soc. 118 . 1430 (l97l).

Damjanovic, A., Genshaw, M. A. and Bockris, J. O'M., J. Chem. Phys. 4^, 4057 (l966).

250
i [17] Riddiford, A. C, J. Chem. Soc. 1175 (i960); Electrochim. Acta 1, 170 (l96l).

[18] Meyer, R. E., J. Electrochem. Soc. 107 . 847 (i960).

[19] Green, M., in 'Modem Aspects of Electrochemistry, Vol. 2' (j. O'M. Bockris, Ed.)

I Butterworth, London 343 (l959).

j [20] Gnanamuthu, D. S. and Petrocelli, J. V., J. Electrochem. Soc. 114 . 1036 (l967).

! [21] Parsons, R. , c.f. Hoar, T. P., in «Proc. 8th. Meeting C.I.T.C.E., Madrid, 1956', Butt1 erworth, London 439 (l958).

' [22] Damjanovic, A. and Brusic, B., Electrochim. Acta 12, 615 (1967).

[23] Hoare, J. P., J. Electrochem. Soc. 112 , 849 (196577

[24] Conference Proceedings: Fuel Cell Catalysis Workshop, Palo Alto, Ca., January 1975;

I E.P.R.I. SR-13, 89-90, August 1975.

j [25] Goldstein, J. R. and Tseung, A. C. C, J. Phys. Chem. 26, 3646 (l972).

' [26] Appleby, A. J. and Savy, M., to be published.

251

National Bureau of Standards Special Publication 455, Electrocatalysis on Non-Metallic Surfaces

,

Proceedings of Workshop held at NBS, Gaithersburg, Md. , December 9-12, 1975. (Issued November 1976)

Electrocatalytic Activity of Phthalocyanines

for Oxygen Reduction

R, J. Brodd, V, Z. Leger, R. F. Scarr, and A. Kozawa

Union Carbide Corporation

P. O. Box 6116

Cleveland, Ohio 44101

A survey of the activity of phthalocyanines as catalysts

for oxygen reduction has identified several metal phthalocyanines with good activity. The central metal ion evidently

controls the overall activity of the phthalocyanines for oxygen

reduction. Ferric phthalocyanine had the best overall activity

and approaches platinum in ability to catalyze the oxygen

reduction reaction. The structure also influences the phthalo

cyanic activity. For copper as the central ion the order of activity is: polycopper phthalocyanine > copper phthalocyanine > copper porphyrazine. The substrate also influences the activity

of the phthalocyanine as does solution composition. A mechanism for oxygen reduction is proposed which includes a hydrogen

peroxide adsorption -desorption step. Rotating ring-disk results

satisfactorily fit this model.

Key words: Ferric phthalocyanine; kinetics; oxygen

reduction; phthalocyanine catalysts.

1. Introduction

The oxygen reduction reaction on solid electrodes is one of the fundamental reactions in electrochemistry. The reaction proceeds at a useful rate only on "active" surfaces

as the reaction is highly irreversible (low exchange current). The search for catalysts for

this reaction - other than noble metals - frequently follows a path suggested by biological

compounds. Several biological agents, e.g., catalase, haemin, hemoglobin speed hydrogen peroxide decomposition and the transfer of oxygen in biological processes. The

solubility of most of the biologically active materials mitigates against their use in practical oxygen electrode systems. Structurally similar compounds, porphyrins and phthalocyanines, have been shown to have catalytic activity for the electrochemical oxygen reduction reaction in acid, base, and saline electrolytes [l -4]1.

2. Effect of Phthalocyanines

Using a rotating disk electrode technique, the activity for oxygen reduction of

phthalocyanine compounds with various central metal ions was determined on a pyrolytic

carbon electrode in isotonic saline solution. The phthalocyanines purchased from Eastman

Chemicals Co. were applied to the electrode surface from pyridine solution. The valence

state of the central ion was not verified experimentally. Typical results are shown in

Figures in brackets indicate the literature references at the end of this paper.

253
+ 0.2 0 -0.2 -0.4 -0.6 -0.8

Potential vs. S. C. E. , Volte

Figure 1. Ten successive cur rent -potential curves taken at a ferric

phthalocyanine -graphite electrode in 0. 05 M sodium acetate

(oxygen -saturated) solution at 23°C; electrode area: 0. 268 cm^.

Solution was renewed after third and seventh runs.

figure 1. The activity is judged by the potential at which the activity for oxygen reduction was noted as well as the shape of the reduction curve. The overall effectiveness of the

phthalocyanine for catalyzing the oxygen reduction reaction depends markedly on the central

metal ion. The most effective central metal ion was Fe"*"^. The relative order of effectiveness for central metal ions was found to be: Fe"*"^ > Fe"^^ > Co''"^ > Al"'"^ > Cr"*"^ > Cu"*"^ > Ni+2 > Sn + 2 with little or no activity for: UOz + ^, Ga+3, Sb+3, Na+i, Zn + 2, Ag+\ Ru+3.

Generally transition metals were most effective as catalysts. However, Al+^ and Sn^"^ exhibited some activity. The order of activity agrees fairly well with previous reports in view of the varying conditions. Savy [ 5-7] reported the activity of Fe^^ > Fe"*"^ but found

Ni+^ > Cu+^ in contrast to that reported here. T. Beck [8] reported essentially the same

order as Savy in acid solutions using gold electrodes. Alt, et al [9], however, report

that the activity of Fe"*"^ > Fe"*"^ based on theoretical arguments.

The activity of the phthalocyanine compounds resides in the ability of the central

metal ion to coordinate (react with) oxygen. One visualizes that the phthalocyanine compound orients itself flat on the surface, the conjugated benzene rings of the phthalocyanine

parallel to the hexagonal carbon structure of the base plane of pyrolytic graphite. Molecular oxygen coordinates with the central ion to form the Oj" ion. Subsequent electron

transfer may occur through the phthalocyanine molecule accompanied by proton addition.

Hydrogen peroxide forms then splits to produce hydroxide as the final product.

254
The influence of structure of the organic portion of the molecule is demonstrated

I with copper as the central ion.

decreasing

activity

Polycopper phthalocyanine

Copper phthalocyanine

Copper octaphenylporphyrazine

The polymeric form has the highest activity. Similar results have been reported for

ferric compo\mds in acid and basic solutions. Evidently the increased conjugated tt bond

character of the polymerized compound enhances the ability of the central metal ion to enter into the oxygen reduction reaction. Substitution on the ring structures also influences

the electron density on the central metal ion and thereby the catalytic activity.

3. Effect of Electrode Material

The substrate has a definite effect on the ability of ferric phthalocyanine (FePc) to function as a catalyst. Graphite seems to be specific in its ability to utilize phthalocyanines

as a catalyst. On graphite, the activity of FePc approaches that of a platinum electrode.

When deposited on platinum or nickel the activity of the phthalocyanine catalyzed electrodes

for oxygen reduction decreased. On some electrode materials, the FePc has essentially no enhancement of the oxygen reduction reaction.

The specificity of graphite may be related to the structure of the base plane. The

carbon atoms in graphite form a ring structure closely resembling the benezene ring. The

ir bond character of the phthalocyanine could interact with the carbon surface in such a way

as to enhance the catalytic activity of the phthalocyanine for molecular oxygen reduction.

4. Stability

To be useful in practical applications, the phthalocyanine catalyst must retain its activity over a wide range of potential and pH. In neutral and basic solutions the FePc is stable over the potential range from oxygen to hydrogen evolution. Repeated potential

excursions have essentially no effect on activity as shown in Figure 1 for saline solution.

Figure 2 for sulfuric acid solutions shows that an irreversible reaction occurs with FePc

when the potential approaches the -0. 3 V versus saturated calomel. Potential excursions

beyond this point destroys the catalytic activity of the FePc. At room temperature, the

FePc activity remains stable near the oxygen potential in acid? basic and neutral solutions.

In phosphoric acid at 125°C FePc quickly transforms into phthalic acid with no catalytic

activity.

5. Method of Catalyst Application

Four methods of catalyst application have been evaluated. Vapor deposition and

evaporation of a pyridine solution on the electrode surface generally produce electrodes of

equal catalyst activity. Due to the high resistivity of the phthalocyanine, care must be

exercised to prevent buildup of thick films which lower activity. Deposition of the catalyst

by slow dilution of a concentrated sulfuric acid solution produced the most uniform deposits

and electrodes with the highest activity. It was thought that a different modification of the

FePc crystal was produced by this technique but this could not be established experimentally. It was possible to deposit Ihin films of 1 0-1 00 jJ-g/cm^ of catalyst on graphite electrodes

with activity closely approaching that of platinum for oxygen reduction.

A variation of the acid deposition technique also produced electrodes of good activity.

The FePc was first precipitated at 0°C from concentrated sulfuric acid as described above.

This procedure produced extremely small catalyst particles which could be filtered out.

255
8.0

6.0 -

1

2.0

4.0 —

+ 0,4 0 Potential vs

.

-0. 4

S.C.E., volt

0. 8

Fi g. 2, Current-potential curves in 0. 05 M H2SO4 + 0. 15MNaCl, nitrogen- saturated;

electrode area: 0. 268 cm^. (1) Ferric phthalocyanine-graphite electrode (1st

sweep); (2) Ferric phthalocyanine-graphite electrode (2nd sweep); (3) Graphite

electrode.

These particles were suspended in acetone and sprayed onto electrode structures with an

air brush.

se use

60 Porous Oxygen Electrode Performance

" Teflon " -bonded porous fuel cell type oxygen electrodes were constructed using

:veral different active carbons [I0]o Table 1 lists several properties of the carbons

used in electrode fabrication. The most important parameter was the ability of the carbon

to decompose hydrogen peroxide. Heat treatment of the carbon in H2O and CO2 environments increased their ability to decompose hydrogen peroxide. Attempts to differentiate

between carbons using ESR were partially successful. The signal enhancement in air versus in vacuum qualitatively may be related to the strength of the carbon -oxygen surface

bond [11], However, a clear relationship between ESR results and carbon activity was

not established. Other factors which appeared to be of greater importance in choosing an

active carbon included surface area to resistivity ratio and carbon crystallite size from

x-ray line broadening.

256
Table 1

Effect of Carbon Support on Catalyst Activity in Saline Solution.

Electrode Potential

vs Calomel Reference q BET

Nuchar Decomposition Surface

Type QCV ImA/cm^ Rate Constant S. A. /pXlO'^ Area mV^

C-190-N 0.212 -0.246 8 1.05 800

CN 0.202 -0.275 11 1.00 960

WV-L 0.168 -0.215 45 1.38 890

WV-H 0.160 -0. 224 27 1.57 1100

The carbon was catalyzed with FePc before and after electrode fabrication with

little effect on the overall electrode performance. Over 3700 hours of continuous operation were attained for several electrodes in saline solution. Failure was associated with delamination of the carbon matrix from the nickel conductor mesh, rather than deterioration

of the catalytic activity.

7. Comparison of Ft and Carbon Surfaces

Using a rotating ring-disk technique, the activity of platinum and carbon surfaces

for oxygen electroreduction was compared in alkaline solution. Various reaction models were used to fit the results from the experiments. This type of analysis permits a direct

evaluation of the effect of various catalysts on each specific kinetic step of oxygen reduction.

In order to satisfactorily fit the experimental results, the following model was used:

Ki K, (1)

O2 ^ HO2" (ads) -* OHK.

HO2" -* ring.

An interesting feature of the model is the inclusion of an adsorption -desorption for hydrogen peroxide, steps K_4 and K4. The platinum res\ilts could be fitted satisfactorily without

the adsorption step in the model. However, for graphite the inclusion of a slow desorption

step was essential in order to account for the observed behavior. Apparently the polar

H02~ ion is more strongly adsorbed on the carbon surface than on a platinum surface. This

may lead to increased surface coverage on the carbon and may account in part for the

slower peroxide reduction rate. The results in Table 2 suggest that platinum surfaces are

10 to 30 times as active for oxygen reduction as are carbon surfaces. This difference

arises primarily from the ability of platinum to support a higher rate of hydrogen peroxide

reduction and the slow initial oxygen reduction step on the carbon surface. The FePc catalyzed carbon surface approaches the activity of platinum. The overall reaction rate is about the same for both electrodes as indicated by the disk current, but the rates on the

various steps in the reduction process differ. There is more peroxide reduction on the

FePc catalyzed surface which closely parallels a platinized platinum surface activity. The

rate constants for the FePc catalyzed surface were fit for a specific set of experimental

data and will vary slightly with catalyst loading.

257
Table 2

Rate Constants for Oxygen Reduction Reaction on Platinum and Catalyzed

and Uncatalyzed Pyrolytic Graphite Electrodes in 1 N KOH.

Platinum

Pyrolytic

Graphite

FePc on

Pyrolytic

Graphite

Kj (cm/s )

K2 (cm/s)

K3 (cm'^/mil s)

K4 (cm/s)

K_4 (cm/s)

0. 15

1. 0

10'

4

3. 6 X 10-2

0, 013

0. 03

10^

0. 08

3. 0 X 10-^

0. 10

5. 4

10^

0, 013

3. 9 X 10-2

8. Acknowledgement

The authors gratefully acknowledgement the partial financial support from the

National Institute of Health, Artificial Heart Program, Contract No. PH-43-68-1391.

References

[1]

[2]

[3]

[4]

56

78

9

[10]

[11]

Zilionis, and R. J,

Zilinois, and R. J.

Brodd, J. Electrochem. Soc. 117 , 1474

Brodd, J. Electrochem. Soc. 118, 1705

A. Kozawa, V. E. Zilionis, R. J. Brodd, and R. A. Powers, "Artificial Heart

Program Conference Proceedings" National Heart Institute, June 9-11/ 1969

Washington, D. C. , Chap. 71, p. 849, A. Kozawa, V. E. Zilionis, and R. J. Brodd, J. Electrochem. Soc. 117, 1469

(1970).

A. Kozawa, V.

(1970) . A. Kozawa, V. (1971) . M. Savy, C. Bernard, and G. Magner, Electrochim. Acta 20_, 383 (1975).

M. Savy, P. Andro, C. Bernard, and G. Magner, Electrochim. Acta 18, 191 (1973).

M. Savy, P. Andro, and C. Bernard, Electrochim. Acta, 19, 403 (197T). F. Beck, Ber. Bunsenges Physik Chem. 77_, 353 (1973).

H. Alt, H. Bender, W. Linder, and G, Sandstede, J. Electro Anal. Chem. 31,

App. 19 (1971). R. J. Brodd and A. Kozawa in Methods in Electrochemistry, A. Salkind and E. B. Yeager, Eds. , J. Wiley & Sons, in press.

L. S. Singer in Proceeds of the 5th Carbon Conf. , Pergammon Press, New York

1963, Vol 2, p. 37.

258
National Bureau of Standards Special Publication 455, Electrocatalysis on Non-Metallic Surfaces,

Proceedings of Workshop held at NBS, Gaithersburg , Md. , December 9-12, 1975. (Issued November 1976)

Continued Growth of Anodic Oxide Films at Platinum

and the Mechanism and Catalysis of Oxygen Evolution

A. Damjanovic^

It is shown on the example of oxygen evolution at platinum

that thin anodic oxide films may profoundly affect kinetics and

mechanism of an electrochemical reaction. At any potential above

l.OV an "insulating" oxide film grows over the electrode surface.

The potential difference across the anode is then divided into the

potential difference across the oxide film and that across the

electrolyte double layer. The film is a barrier to electron

transfer and together with the reaction across the double layer

controls the kinetics and mechanism of the reaction. The exchange

current density of the reaction decreases exponentially with

thickness of the film. Only when the distribution of potential

across the oxide film and double layer is taken into account,

it is possible to explain the observed fractional reaction order

with respect to activity of hydrogen ions and to propose a mechanism

of the reaction that explains this reaction order. In the proposed

mechanism a chemical step that follows first electrochemical

step controls the rate in the double layer.

Key words: Anodic films; dual barrier, electrocatalysis; oxygen

evolution; platinum anodes; reaction mechanism.

I. Introduction

Thin surface oxide or hydroxide films form at sufficiently high potentials even at anodes of noble metals such as gold or platinum. These films may affect kinetics of an electrode reaction; usually they decrease the reaction rate. For instance, the rate of

anodic oxidation of hydrogen [1]^ and hydrocarbons [2] decreases by orders of magnitude

once an electrode has reached the potential at which an oxide film initiates and grows over

the electrode surface. Cathodic reactions too, e.g. of oxygen reduction [3], are affected

by surface oxide films. Though it has been evident for some time that a thin anodic film

decreases the rates of various electrode reactions, little attention has been paid in the

past to analyze in some details its effect on the kinetics and mechanism, — and hence

electrocatalysis, — for these reactions. The knowledge of the nature and properties of the

anodic oxide films, — and in particular of the kinetic character of film growth, — that is essential for such an analysis was frequently lacking. It will be illustrated here by the

example of anodic evolution of oxygen at platinum that only when the presence and kinetic

character of film growth is taken into account one can avoid erroneous conclusion about the

reaction mechanism and factors affecting the catalysis.

Address: 809 Harbor Square, Rochester, New York 14617.

Figures in brackets indicate the literature references at the end of this paper.

259
II. Oxide Film Formation

At a platinum anode an oxide film grows at any constant current density (e.g. 10"^ -

10"* Am"^) whenever the electrode potential exceeds l.OV (vs HE). As shown by a number of

workers, — and notably by Vetter and Schultze [4], — the potential, V, at a constant current

density increases linearly with time or integrated charge, q, as shown in figure 1. Significantly, the dV/dq slopes are proportional [5] to log current density (fig 2). As shown with a rotating disk-ring electrode, in the linear V-time region and before a critical

potential is reached (cf. N in fig 1) current is used up entirely for the formation of an

20 40 60 80 100 TIME(sec)

Fig. 1. Change of potential with time at Pt anodes at a constant current density

of polarization (i = 2xlO"^Am"2). The linear region is followed by a quasi-steady state. Oxygen starts to evolve after point N [5].

-6 -5 -4 -3 -2 -I

LOG I n in Acm-^l

Fig. 2. The dependence of (9V/3q)i on current density of polarization [5].

260

:
anodic oxide film (with no O2 evolution though the electrode potential exceeds the reversible

O2 potential by as much as 0.3 to 0.4V). After the critical potential, — which increases as the applied constant current density increases, — oxygen starts suddenly to evolve and soon becomes predominant reaction. The rate of oxide growth is then substantially reduced. These

changes in the anodic processes are reflected in the V-t and V-q diagrams: the potential

'flattens off and assumes a 'steady state' value. As will be discussed below, in this

region too, neither potential nor film thickness remains constant; they continue to increase

with time of polarization but now very slowly.

The rate of growth of anodic films in the linear potential region is well accounted for

by the formalism of the Mott-Cabrera model of growth for high field assisted ionic diffusion.

For this model the rate equation can be written as:

- zeNv exp [- exp l^^A,

Here, Vq is the potential at which a film starts to growS, AVof is the potential difference

across the film, d is (average) film thickness (Vgf/d is hence field within the oxide film),

z is the charge and X 'half jump' distance of migrating ions. N represents a concentration

term and AG* is zero field activation energy. Other symbols have their usual significance.

From figure 2 and with a reasonable value for the conversion factor to relate q to d, the

transfer coefficient, a, has been found [5] to be 16nm V"^. This leads to A s:0.2nm and

z = 2. Anodic film is, therefore, a divalent oxide phase in accordance with a previous

ellipsometric study [6].

Experiments at various temperatures gave AG=j= 25 kcal mol"^, and Nv a^3.10^^ s"^ m"^.

This product is a characteristic of the rate determining step. With the mean frequency

V = 3.10^^ s"^, — calculated from the Debye temperature, — N is very closely equal to the

density of atoms in the surface of the metal. This remarkable agreement strongly suggests

that a process at the metal/oxide interface, — rather than within the oxide phase, — controls

the rate of oxide film formation'+»

The most important outcome of these studies of oxide film growth is the realization

that a strong field ( = 5.10® Vm"M must exist within a growing oxide film [4,5,8]. Hence,

the electrode potential difference, AV, must be distributed between the potential difference

across the oxide film itself, AVof, that across the oxide film — electrolyte double

layer, AVd]. Such a potential distribution is expected to affect both kinetics and mechanism

of an electrode reaction that occurs at oxide covered electrodes [9].

Ellipsometry has shown that after the critical potential (N in fig 1) at any constant

current density of polarization, in, film thickness increases logarithmically with time

(cf. fig 3) in such a way that [10]:

' f(ip) (2)

and

(TltTp), = "-Mf(t) (3)

The rate of this slow, extended growth is described by equation:

3 Vq is close to 0.9 V vs. HE. Contrary to some literature readings, the same rate equation as above describes initial

stages of oxide film growth irrespective of the mode of polarization [7].

261
Fig. 3. Thickness of anodic oxide films at platinum as function of

time at constant current densities of polarization [10].

that is diametrically different from the equation of initial growth (cf. eq (1)). Similar

relationship is obtained for the extended growth at a constant potential mode of polarization

[10]. Apparently, relationship given by eq (4) does not reflect any particular change in AVq^ which is now controlled by the major ongoing reaction, the oxygen evolution reaction,

and which even at a constant mode of polarization is not constant any longer. Rather, this

relationship reflects a genuine change in the mechanism of growth. The change occurs only

after O2 starts to evolve and has become a major reaction. This is an unexpected result,

the meaning of which is as yet unclear. The growth according to eq (4) has previously been

observed at other metal electrodes, e.g. at iron electrodes [11]. However, in the latter

case oxide growth is the only reaction (no O2 evolution) that occurs at the electrode and

the change of potential with time reflects the mechanism of oxide growth only.

III. Oxide Films and Oxygen Evolution

It has recently been shown [12] that both the el 1 i psometri c thickness of anodic films

and electrode potential change linearly with logarithm of current density, ip, with which

a prereduced electrode is anodically polarized for a given time. For all times of polarization the dV/d log ip slopes are close to 170 mV and dd/d log ip to 0.1 nm (figs 4 and 5).

After an electrode has been polarized for a given time with a constant current density and

the current density, i, is decreased in steps, no change in thickness, d, is observed (from

A to B in fig 5). In contrast to this, AV decreases as i is decreased (fig 4). The

dV/d log i slopes are now close to the frequently reported value of 115 mV or to 2.3 (2RT/F).

Thus, the linear V-log i relationship with the slopes log i = 120 mV, — observed for over several decades of current density [13, 14], — is obtained only at electrodes that were anodically pretreated at a higher potential or current density to form a stable oxide film,

and when it was ensured that the thickness of the film remained unchanged during

measurements. The slope is then independent of film thickness [9], i.e.:

/ 9V \ 2RT , .X (5)

262
Fig. 4. Electrode potentials after 125, 750 and 1800 seconds at a given

current density of polarization (circles). Squares represent

subsequent potential --current density relationships at decreasing

current densities [12].

Fig. 5. Thickness of anodic oxide films after 750 seconds at a given

current density of polarization (circles). Squares represent

subsequent thickness--current denstiy relationships at decreasing current densities [12].

263
It has been shown [12] that log i is linearly related both to V and to d, and that the

function^

= C exp (-

-J-) exp (^pj) (6)

gives the dependence of the rate of oxygen evolution on electrode potential and film

thickness. For constant thickness, current density is given by:

i

with exchange current density equal to

i 0

Thus, the exchange current density, — a measure of the catalytic activity, — decreases

exponentially with the thickness of the oxide film. Unless the effect of film thickness is taken into account, any correlation of the catalytic activity with a particular physical

parameter at various anodes (e.g. with bond strength of adsorption of reactants or reaction

Intermediates or with work function of anode material) is, therefore, of dubious validity.

As a consequence of the increase of film thickness with logarithm of time, potential

too at a constant current density increases with the logarithm of time (cf. fig 6). The

V-log i lines at various constant current densities are, therefore, parallel between themselves and equally spaced for an equal increment of log i.

I 10 100

TIME (sec)

Fig. 6. Potentials at platinum as function of times at different constant

current densities of polarization [10].

^The same relationship was obtained by Schultze and Vetter [15] from different experiments

and with different arguments.

,FAV,

^0

C exp (- f)

(7)

(8)

264
IV. Mechanism of Reaction

In some earlier studies of O2 evolution, either the presence of oxide films at electrodes was ignored or insufficient care was taken to keep film thickness constant in solutions

of different pH. This then led to erroneous conclusion that the reaction was pH independent

and that the first electrochemical step was rate determining [13]. Recent experiments have

shown, however, that for the same thickness of the oxide film the rate is represented by

equation

i = k [H^]-i/^ exp [i^] (9)

i.e., tTie reaction order with respect to is negative and fractional (not zero!). The

fractional reaction order precludes the first electrochemical step as rate determining.

Moreover, fractional reaction order could not be accounted for by any simple analysis either

under Langmuir or Temkin conditions of adsorption. It is explained [9] in terms of a dual

barrier model [16, 17] according to which there is a potential difference both across the

oxide film and across the oxide/electrolyte interface, i.e. it is explained only after the

presence of anodic oxide films is considered in the analysis. According to this model, current densities across the oxide film and across the double layer are given respectively

by [12]:

''of = ^0 (-"^^^ 1^^^^

and [9]:

id! = A [H""]"' exp [Mdl] . (11)

Here, Xq and A are constants independent of pH. Evidently, only reaction across the double

layer is pH dependent but with no fractional reaction order. Analysis has shown further that

the observed current density is given by [9]:

T = = = (AX^) [H ] exp (- "2-) • exp [^^y] (12)

which for a constant thickness reduces to eq (9) and thus accounts for the unusual pH

dependence of the reaction. It is suggested that the rate determining step for the reaction

across the double layer is a chemical step which follows the first, fast electrochemical

step. The rate across the oxide film is most likely controlled by the quantum mechanical

tunneling process, but as yet there is no satisfactory account for this equation.

V. In Conclusion

In conclusion, it is illustrated here on the example of oxygen evolution at platinum

anodes that a thin, 'insulating' anodic oxide film may profoundly affect the rate of an electrochemical reaction. The observed change in the catalytic activity of the oxygen

electrode is related to the change in thickness of the oxide film. Oxide film affects also

the mechanism of the reaction. Hence, whenever an electrode surface is covered by an oxide

film, its presence has to be taken into account in analyses of mechanisms and catalysis of

electrode reactions. Of course, an anodic oxide film at an electrode does not have to be

'insulating' in which case the rate of a reaction should not be greatly affected by the

thickness of the oxide film. Oxygen evolution at nickel anodes is perhaps an example for

this. However, even at platinum,-- when an oxide film is formed at very high current

densities (e.g. i > 10^ Am"^) for a longer time, --there is experimental evidence [18] that

a 'conductive' oxide phase eventually forms. Thus, at a high current density electrode

potential first increases with time, as expected when an 'insulating' oxide film forms, but

then it decreases substantially with further polarization (reactivation of the electrode)

though film thickness continues to increase [18]. Apparently, divalent oxide is transformed

to a higher valence 'conductive' oxide phase. The latter experimental evidence suggests one avenue for a revived search and 'tailoring' of new catalysts for oxygen electrode reaction.

265
VI. References

[1] Schuldiner, S. , J. Electrochem. Soc. rj5^, 362 (1968).

[2] Wroblowa, H. , Piersma, B. J. and Bockris, J. O'M., J. Electroanal. Chem. 6, 401 (1963).

[3] Damjanovic, A. and Bockris, J. O'M., Electrochim. Acta Vl[, 376 (1966).

[4] Vetter, K. J. and Schultze, J. W. , J. Electroanal. Chem. 34, 141 (1972).

[5] Damjanovic, A., Ward, A. T. , Ulrick, B. and O'Jea, M. , J. Electrochem. Soc. 122^, 471 (1975).

[6] Visscher, W. , Optics 26, 402 (1967).

[7] Harris, L. B. and Damjanovic, A., J. Electrochem. Soc. 122, 593 (1975).

[8] Ord, J. L. and Ho, F. C. , J. Electrochem. Soc. 118, 46 TT971).

[9] Damjanovic, A. and Jovanovic, B., J. Electrochem. Soc. 123 , 374 (1976).

[10] Ward, A. T. and Damjanovic, A., Submitted to J. Electrochem. Soc. [11] Sato, N. and Cohen, M., J. Electrochem. Soc. VM, 512 (1964).

[12] Damjanovic, A., Ward, A. T. and O'Jea, M. , J. Electrochem. Soc. 12]_, 1186 (1974).

[13] Bockris, J. O'M. and Hug, A. K. M. S. , Proc. Roy. Soc. (London) A237 , 277 (1956).

[14] Damjanovic, A., Dey, A. and Bockris, J. O'M., Electrochim. Acta ri_, 791 (1966).

[15] Schultze, J. W. and Vetter, K. J., Electrochim. Acta 1^, 889 (1973).

[16] Meyer, R. E. , J. Electrochem. Soc. 107 , 847 (1960).

[17] MacDonald, J. J. and Conway, B. E., Proc. Roy. Soc. 269, 419 (1962).

[18] Vissher, W. and Blijlevens, M. , Electrochim. Acta 1_9, 387 (1974).

266
ELECTROCHEMICAL PROCESSES ON NON-METALLIC SURFACES

;.irional Bureau of Standards Special Publication 455, Electrocatalysis on Non-Metallic Surfaces,

Proceedings of Workshop held at NBS, Gaithersburg , Md. , December 9-12, 1975. {Issued November 1976)

Electrochemical Sirrface Jleactions on Non-metals

R. Memming

Philips GmbH, Forschungslaboratorium Hamburg

2000 Hambiirg 54, Germany

Various surface reaction of semiconductors during

electrochemical polarization are investigated which are

of interest in connection with catalytic and photo-electrocatalytic processes. Examples for the anodic and

cathodic decomposition are given and discussed with re- gard to the stability of electrode materials. In addition

the formation of semiconducting layers on metal electrodes

is described and their analysis and characterization investigated as far as electron transfer is concerned.

Reactions at bronze electrodes leading to semiconducting

surface films are included.

Key words: Anodic decomposition; cathodic decomposition;

electrochemistry; inorganic bronze electrodes;

inorganic bronze electrodes; metal electrodes;

semiconductor electrodes; stability, surface

layers; surface reactions.

1. Introduction

Since there is such a great demand for less expensive electrodes for

fuel cells and suitable materials for photoelectrolytic devices the electrochemical properties of quite a large number of non metallic electrodes have

been studied in the recent past. Unfortunately, not only oxidation and re- duction of hydrogen and oxygen, respectively, were observed but also other

processes such as electrochemical decomposition or film formation have been

detected. These surface reactions could be divided into several processes

such as

semiconductor — decomposition

semiconductor formation of insulating film

semiconductor » formation of metallic film

metals formation of semiconducting film

(or metal-like

conductor)

metals

(or metal -like insulating film

conductor)

In the present paper examples of these processes are given and the stability of the electrode material correlated to the band gap of the semiconductor. The main emphasis is put on the analysis and characterization of the

semiconducting properties of semiconductor electrodes, of semiconducting

films on a metal substrate and of bronze electrodes.

267
2. Electrochemical Decomposition and Stability

of Semiconductor Electrodes

Most semiconducting materials can be decomposed in aqueous solutions byanodic oxidation or cathodic reduction. It is useful to distinguish between el

ement and compound semiconductors since in the first type the bonding is

pure covalent whereas in compounds a certain degree of ionic bond exists. In

all cases the electronic states at the surface have a key function in the

reaction mechanism as demonstrated schematically for the dissolution of an element semiconductor in fig. 1 . In this figure A represents a surface atom

of an element semiconductor to which two hydroxyl groups are bonded. Principally the anodic dissolution occurs in two steps by which the chemical bond

between two neighbours has to be broken under the influence of an applied

electrical field. The latter process, however, is only possible if the surface density of holes is sufficiently large. In the first reaction step a surface radical is formed. Since the chemical bond between the surface atom

A and the crystal is now much weaker the energy level of the surface radical

should not be located anymore within the valence band but it should occur a- bove the valence band as indicated in the lower part of fig. la. The next

reaction step (complete separation of the surface atom) is characterized by an electron transfer from the surface level E, into the valence or into the

conduction band. In the first case also holes are required at the surface. A

conduction band process, however, is only possible for a rather small energy

distance of relative to the conduction band, i.e. such a process can oc- cur only with small band gap semiconductors. It has been observed with germanium electrodes [1]^ (E = 0.8 eV) but with no other semiconductor.

element semiconductor

anodic dissolution /OH . ^OH

^OH ^ ^OH

.a: ,0H

'OH

reduction

b) ^A-OH

^

^A-.OH

-P

Fig Elementary process of anodic

dissolution and reduction of

element semiconductors.

As mentioned above the anodic dissolution mechanism represented in fig.

1a contains only the most important reaction steps. In this example two

charges are required for the dissolution process as it was found e.g. for

Si-electrodes in hydrofluoric solutions [2]. In most other cases four charges are required. Then the unpaired electron remains at the crystal surface

and has to be neutralized either by electron transport to the back side con- tact or by a further hole from the valence band [1,2]. The crystal surface

Figures in brackets indicate the literature references at the end of this paper.

268
is restored by chemical reaction with the reacting molecules such as 0H--

molecules. It should be emphasized that the dissolution current i+ is proportional to the density of holes at the surface? Pg , i.e.

i"-P3 . (1)

Assuming Boltzmann distribution then p is related to bulk hole density p

by so/ eU \

Ps = Po Tfj ' (2)

in which Ugc represents the potential across the space charge layer. Since

in most cases the externally applied potential only occurs across the space

charge layer the anodic current is given by

i"^ ~ exp ^eUg^/kT) = exp^eAUg/kT^ (Ug = electrode potential). (3)

In this case a Tafel slope of 50 mV/decade is expected as it was proved for

various semiconductors [2,3].

The same principle scheme is valid for cathodic processes (fig. 1b). In

this case an electron transfer from the conduction band across the interface

leads to a separation of an OH~-group from the surface. A radical is formed

again and the subsequent step may occur either by charge transfer via the

conduction or the valence band [4,5].

These basic reaction steps also occur in the decomposition process of

compound semiconductor electrodes. The main difference from element semiconductors is due to possible differences in the bonding character. The bonding

could be of covalent or of ionic nature. Two examples are GaP- and ZnO-electrodes. For instance in the case of GaP in which the bonding is mainly covalent, the anodic dissolution in acid solutions is given by [6]

GaP + 3H2O + ep'^-^Ga^^ + H^PO^ + 3H'^ (4)

whereas in alkaline solutions:

GaP + 40H" + 3p"^-»Ga02 + P + ZE^^O . The surface structure is shown in fig. 2. Although the chemical bonding is more covalent differences occur in acid and alkaline solutions. If both atoms are dissolved electrochemically then six charges are necessary for the

dissolution of one GaP-molecule as observed in acid solutions. Since in alkaline solutions only three charges are consumed it may be concluded that

one electron of each bond remains at the phosphorus atom, i.e. neutral phosphorus is formed which is dissolved slowly in a subsequent step. Breaking

one bond, for instance by trapping a hole at the surface, the corresponding

phosphorus surface atom is left with an unpaired electron (fig. 2a). The

energy of this unpaired electron does not coincide with the valence band,

its position is somewhat above the valence band, as indicated by the energy

diagram in fig. 2. Principally the unpaired electron can be transferred into

the valence band. On the other hand the transition probability from this

energy state into the valence band must be very low in alkaline solutions

because only three charges are consumed. The different behavior in acid and

alkaline solutions are rather difficult to interpret. A more speculative ex- planation could be the assumption that the phosphor radicals are stabilized

in acid solutions by H30+-ions bridged to the surface atoms as shown in

iig. 2b leading to a dovmwards shift of the radical energy level.

The situation is somewhat different for compounds of more ionic bonding

such as ZnO and other Il-VI-compounds . It is.asstuned that hydroxyl groups

are bonded to the more positively charged Zn '''-atom. According to Gerischer

[7] the anodic dissolution can be described as follows:

269
(a)

Ga Ga Ga Ga

Ga Ga"

OH ^ OH

<in>

P.

OH

step

in NaOH

(b)

i I I K P P Ga Ga Ga

P P Ga Ga

(*) ()

<111>

I

Ga H^ ^H

H

^H

in H2SO4

EcFig. 2. Surface structure of GaP

lattice during anodic

dissolution and energy

diagram, (a) in NaOH;

(b) in H2SO4.

„2- + slow ^-

O3 + p ^ 0^

r,- o2- + slow f ^\2- O3 + Og + P (0-0)

(0-0)^" + 2p+ £^ 0. (6)

2Zn^+ + aq 2Zn^^ + 2 aq

2ZnO + 4p^

^2Zn^'^ + 0^

In this case the hole p produced by light excitation is trapped at a more negatively charged oxygen atom at the surface (0§~) and oxygen is evolved in

two subsequent steps leading also to a dissolution of Zn. According to Justice and Hurd [8] the anodic current is quantitatively due to dissolution.

During cathodic polarization the reverse process was observed. In this

case electrons are trapped at a more positively charged Zn-atom at the surface resulting in the reduction of Zn2+ to Zn and of oxygen to H2O according to the mechanism given by

-7 2 + Zn + e

s Zn

Zn + e Zn (7)

0^" + 2H^

s H2O

ZnO + 2e' 2H Zn + H2O

270
These reactions are typical for various other Il-VI-compounds . For instance anodic polarization of CdS results into the dissolution of cadmiiam.

I In this case, however, sulfur remains as an insulating layer at the surface [9]. On the other hand cathodic polarization leads to the formation of

•a metallic film on the surface.

In table 1 the decomposition products of various semiconductor elecitrodes during anodic and cathodic polarization are summarized.

Table 1. Decomposition products o£ various semiconductor electrodes

during anodic and cathodic polarization.

Electrode E , eV Anodic Cathodic

material ^ decomposition products decomposition products

^e(n,p) u . o Ge^- [1] "2 r 1

1

[

' J

, 3i(n,p) 1 1 Si^'',Si^^ (in HF)[2] "2 r ?i

{jaAs (n,p

)

Ga^+.A^^ [10] AShl^ J n2

;("dSe (n,p

)

1 .75 Cd^-'.Se^ [9] Cd, (H2Se) [9]

(-aP(n,p) 2.35 Ga3^P° or P^-' [6] E^i (PH^ < 0.1%) [6]

K:dS(n) 2.4 cd2^s° [9] Cd,H2S [9]

|W03(n) 2.5 02 H^W03,H2 [11]

:SiC(n,p) 3.0 Si^^,C (in HF) [12] H2 [12]

1 ZnO(n) 3.2 Zn^"^, O2 [8] Zn,H20 [8]

:i02(n) 3.2 (Ti'^^),02 [13] Ti,H2 [14]

iK^TaO^Cn) 3.5 O2 [15]

Sn02(n) 3.8 O2 [16] H2,Sn (in HCl

)

[16]

According to this table the most stable or inert electrode materials

exhibit a large band gap. This result seems to be quite reasonable since

jivarious authors have found a linear relationship between band gap Eg and av-

'terage bond energies of binary compounds [17]. Materials of large bond enerIjgies are expected to be relatively stable. This result is important especially

for the selection of electrode materials in photo-electrocatalytic devices

;in which mainly Ti02 has been used. In connection with these problems it

!ishould be mentioned that several materials show a quasi stable behavior in

I'the dark, e.g. n-type electrodes in the anodic range where the current is

limited to a very low value due to the absence of holes. Any excitation of

tides by light absorbed by the electrode, however, may lead to anodic disso-

|lution or film formation. Whether dissolution or formation of an insulating

:.film occurs depends mainly on the electrolyte used. For instance SiC seems

\to be a very stable electrode in aqueous solutions. However, this stability

lis only due to the formation of an oxide film which also limits the electron

(transfer [12], Only in hydrofluoric acid anodic dissolution was observed.

271
3. Characterization o£ Electrode Surfaces

Some information about the electrochemical behavior can be already obtained from current-potential measurements. Taking n-Sn02 as an example the

anodic current is limited to a very low value up to large anodic potentials

(fig. 3). This current rise is due to oxygen evolution caused by electron

tunneling through the space charge region[l6,18,20] . This tunneling process

only occurs with heavily doped electrode materials, For lower dopings this

current rise does not occur or only at much higher potentials. In the lower

potential range the current is limited to a very low value because holes are required for this process. Light excitation leads to current increase

(fig. 3). The corresponding photocurrent sets in at the flat band potential

and is proportional to the light intensity. Its absolute value is also dependent on the recombination of electron-hole pairs. In some cases such as Sn02 and Ti02 the recombination depends on the doping and on the crystal

structure. For instance much larger photocurrents were found with single

crystals of these compounds than v^ith polycrystalline layers [16,19].

electrode potential vs. s.c.e

Fig. 3. Current vs. potential for SnOo - electrodes

in 1 N H2SO4 (vs . SCE) .

The flat band potential Ufb and the donor density Nd are usually obtained from measurements of the space charge capacity Csc vs. electrode potential Ue. In the Schottky-Mott form the capacity is given by

Lou

The donor density Nd can be derived from the slope of the curve provided

that any externally applied potential occurs only across the space charge

layer.

272
4. Electrochemical Process at Oxide Covered Metal Electrodes

4.1 Characterization of oxide layers

In many electrochemical experiments semiconducting oxide films are

formed on metal electrodes. In some cases such as platinum these oxide films

are very thin and the current is mainly determined by electron tunneling

through the oxide layer [21] (< 20 ^). These thin layers cannot really be

characterized by semiconducting properties because the lattice periodicity

is too short. Thicker oxide layers (> 100 S), however, exhibit semiconducting properties. These properties are mainly derived from measurements of the

photocurrent or of the space charge capacity in the oxide. Usually the latter method is used and the results are plotted as l/C^ vs. electrode potential Ug (Schottky-Mott ) . In many cases a linear dependence has been obtained

and from the slope of such a curve the donor density has been derived. It is

rather surprising, however, that this linear relation could be verified for

.itany oxide films formed on metal electrode although the composition should

vary from the metal-oxide to the oxide-electrolyte interface considerably as schematically shown in fig. 4b. Accordingly the donor density should decrease from the metal towards the oxide surface resulting in a corresponding band bending shown in fig. 4c. In fact different slopes and consequently

I Me.Oy

metal

a)

oxide

x=0

electrolyte

x = d

c)

Fig, 4. Schematic presentation of a)

metal/oxide/electrolyte interface; b) oxide distribution;

c) potential distribution.

different donor densities were obtained for different oxide thicknesses on niobium electrodes [22,23] as demonstrated by Schottky-Mott-plots of the

capacity in fig. 5. According to Heusler and coworkers [23] the donor density Nj) decreases inversely proportionally to the square of the thickness d,

i.e. Nd = kd~2 in which k is a constant. This relation was not only found

for Nb20^ but also for Ti02 titaniiam [24]. The question arises, however,

v/hether under these circtimstance the slope of the Schottky-Mott-plot still

273
U. V

Fig. 5. Schottky-Mott plot of space charge capacity for

niobium with oxide films (NboO^) in 0.1 M oxalic

acid, 15 nm () and 28 nm (oj oxide thickness

(taken from K. E. Huesler [22]).

determines the donor density in bulk of the oxide. This problem can be solved by an approximate solution of the Poisson equation which is given for n- type semiconductor by

^ = aCN-RoB-y) ~ aN (9) dx2 ^

in which a = N donor density, no electron density in the bulk, y is

the potential o in the space charge region in kT-units. The above equation

is valid for upwards band bending, i.e. for an exhaustion region (fig. 5b).

Assuming a donor density of

N = ^

(see fig. 6a) (10a)

(a-x)2 . the donor density at the surface is given by . "

•

N„ =^~-| (10b)

a d

'

Integrating the Poisson equation one obtains a potential distribution of

274
ad 0

Fig. 6. Schematic presentation of a) donor

distribution, b) energy diagram at

metal/ oxide/electrolyte interface

.

y = - k ln(1 - |) + Kx + (11

)

or after expanding the first term into a series:

y = k(| + ^ + ...) + Kx + y^ for I

< 1 . (11a)

Terminating the series after the square term one obtains for the SchottkyMott-equation z^, _ d^dy/dx\

- ^^o dy^ ^

1 ^ 2 kT

? ee e^N

0 o y, (12)

i.e. the slope is determined by the donor density Nq at the oxide -electrolyte interface. A similar result is obtained assuming an exponential depend- ence between donor density and thickness and subsequent series development.

In a certain potential range in which the thickness of the space charge lay- er is still much smaller than that of the oxide layer this approximation is

valid and the surface donor density Nq may be derived from a Schottky-Mottplot. It must be emphasized that these capacity-potential dependences are not sensitive to the donor distribution within the oxide.

4.2 Surface Reactions

The electrochemical processes at oxide film covered metal electrodes can

be determined by the semiconducting properties of the oxide film in the same way as found for single crystals. For instance the current potential curve

I of polycrystalline Ti02 on Ti is identical to that of single crystals, i.e.

I the anodic current is limited to a very low level due to the lack of holes

required for oxygen evolution [13,19] (fig. 7) whereas the anodic current is

increased by illumination. The photocurrent depends considerably on the pretreatment of the oxide layer [25]. Other electron transfer processes such as

275
0 +0.5 +1.0 +15 V +2.0 ^ electrode potential

Fig. 7. Dark and photocurrent vs. electrode potential for Ti02

layer (d = 0.4 ym) on Ti substrate in 0.1 N H2SO4.

redox reactions proceed at Ti02-layers in the usual way although the ex- change current via the conduction or the valence band is rather small due to

the large band gap [19,21]. On the other hand the spectral distribution of

the photocurrent may depend on the substrate used as a film support as it was also found for Ti02-layers. In fig. 8 the spectral distribution of the

anodic photocurrent is shown, as measured with Ti02 deposited by pyrolytic

decomposition on Sn02. According to this figure the spectral distribution

depends very much on the thickness of the Ti02-layer if the electrode was illuminated through the electrolyte whereas it was independent of the thickness illuminating the system through the backside. Light was not absorbed by

Sn02 because the band gap (3.8 eV) is larger than for Ti02 (3.1 eV). In the

first case the maximum occurred at rather short wavelengths using 0.I [i thick Ti02-layer. This maximum occurs at the same wavelength as found for

backside illumination (compare fig. 8a and b). This result indicates that

the highest yield for the photocurrent was obtained if electron-hole pairs

were created near the interface Sn02_-Ti02, i.e. the charge carriers were separated by the electric field at this interface causing hole injection into Ti02 which are then consumed at the Ti02/electrolyte interface. On the

other hand illumination of a thicker Ti02-layer (> 1 |i) by light passing

through the electrolyte leads to a charge seperation at the Ti02/electrolyte

interface due to the electric field of the space charge layer below this

surface. In this case the spectral distribution of the photocurrent is identical to that found for single crystals [25]. This substrate effect is ex- pected to depend also on the relative position of the Permi level in Ti02

and consequently on the doping and preparation technique of the oxide.

The question arises, however, whether always semiconducting properties

of an oxide layer determines the charge transfer process across the interface. Mostly capacity measurement were performed in order to prove whether

an externally applied potential occurs across the space charge. As outlined

in the previous section this type of measurement yields reasonable Schottky276
A cm

Acm

3

u

_o

"o

Q.

electrolyte

hv

b)

350 ^00 X[nnn]

wavelength

c)

electrolyte

Fig. 8. Spectral distribution of photocurrent for Ti02 on Sn02 - substrate

a) illumination through electrolyte; c) illumination through

substrate; b) and d) excitation process.

Mott-dependences in a limited potential range for oxide films grown on a metal electrode. This has been proved also for Ti02 on Ti-electrodes [19,24]

and the flat band potential could easily be determined. In some other cases,

however, rather porous oxide films are formed on the metal electrodes such

as for instance Ag20 on silver. Ag20 is readily formed on a Ag-electrode in

alkaline solution above around 0.3 V (vs. SCE) (fig. 9). The further oxidation to AgO occurs above 0.65 V. Ag20 and AgO are also expected to show

semiconducting properties. However, they are difficult to characterize since

no reasonable Schottky-Mott-plots were obtained from capacity measurements.

On the other hand, after having formed Ag20 on silver by anodic polarization

according to the reaction

2 Ag + 2 OH"—^2 Ag20 + H2O + 2e" (13)

anodic photocurrents were observed due to light absorbed by AgpO. Gohr [27]

interpreted this photoeffect by assuming a photodissociation. According to our ov/n measurements [28], however, the photocurrent could be related quantitatively to the further oxidation to AgO, as also confirmed by Conway [29].

277
It is interesting to note that the photocurrent sets in at a potential

(0.4 V) 0.25 V lower than that where the oxidation occurred in the

dark (fig. 9b). This is a typical semiconductor effect and could be interpreted by assuming the flat band potential to be around 0.4 V so that the

energy bands are bent upwards above this potential. The corresponding electric field forces the holes created by light excitation towards the surface

where they are consumed for further oxidation. Accordingly the second oxidation may described by

Ag20 + 2 p"^ + 2 0H~— 2 AgO + H2O . (14)

This description is oversimplified since also ionic conduction is required

for this process.

It should be mentioned that the photocurrents caused by light absorption in Ag20 are partly also due to a bulk photoeffect since the conductivity of this oxide is very low. This assumption is supported by the result

that AgO formed during illumination is not reduced at the usual potential

(around 0.5 V) but in the potential range where Ag20 is reduced to Ag. In

this range also a cathodic photocurrent was observed (fig. 9b). The reduction of AgO is inhibited since AgO formed during light excitation is still

separated from the metal electrode by low conducting Ag20. Further details

are given in a previous paper [28].

278
5. Electrochemical Processes at Tungsten Bronzes

Sodium tungsten bronzes have been studied extensively because this material seemed to be quite promising for oxygen reduction reaction [30,31].

This bronze of the composition NaxWOj with x > 0.25 shows metallic conduc-

' tion. Different models for the conduction band of bronzes have been postulated and have been reviewed recently [32,33]. The electrical properties

change for low x-values and it seems that below x = 0.25, sodium tungsten

bronzes are semiconductors. At least pure WO3 (x = O) is a semiconductor

s with a band gap of 2.5 eV.

Although the sodium tungsten bronzes are very resistant towards acids

it has been shown [30,31] that on anodic polarization sodiiim diffuses out of

the bronze surface according to

1 Na^WO^— Na^ + ^ ^°3 + ^~ (''S)

leaving a semiconducting surface on the metallic bronze electrode. This was

I proved by ion probe analysis [30]. Indications for a semiconducting surface

j layer were already obtained by Hardy and Bockris [30] who observed an an-

! odic photocurrent after having polarized the bronze electrode anodically.

The photocurrent occurred above light energies of about 2.2 eV. Since this

energy corresponds roughly to the band gap of ¥0^ it has to be assumed that

a WOj-layer exists on the bronze electrode. The photoprocess initiated by

light absorbed by WOj must be a valence band process according to

WO^ + 2 OH" + 2

p"^ »W0^ + H2O+I/2 O2 (16)

This is in agreement with current -potential measurements at WO-^-electrodes

I performed by Hobbs and Tseimg [11] who found a small anodic current in the

j-j dark which is typical for n-type semiconductors.

Randin [34] has analyzed and characterized the bronze electrodes by

capacity measiirements . He obtained very nice Schottky-Mott-plots after anj,j odic treatment of the electrodes and derived from the slope of the curves

donor densities of the order of 10 - 1 020 cm3. Moreover, he correlated

the donor density to the concentration of sodium (N(j^ = Nj^^+). It was concluded that the sodium concentration must be much smaller than that obtained by ion probe measurements.

According to our analysis of Schottky-Mott curves for an inhomogeneous

distribution of donor states (section 4.1") only the density of donor states

at the very surface can be derived from the slope of the capacity curves

which may be indeed very low. As mentioned before the Schottky-Mott-plot is

' insensitive to the distribution of donor states. Since in some cases a linear relation between l/C^ and Ug was found in a potential range of about

( 7 V (yo = 300) the thickness of the semiconducting surface layer can be p„

estimated by using eqs (10) to (12) and one obtains d ~ 300 S for Nq = 10

cm"5 which is in agreement with the ion probe measurements.

6. Concluding Remarks

It has been demonstrated that many compound semiconductors undergo

electrochemical decomposition. Only few of them, such as oxides, are relatively stable. However, these materials have a large band gap. This fact

leads to problems as far as photogalvanic systems are concerned. In order

to obtain a large quantum efficiency for conversion of solar energy into

chemical energy semiconductor electrodes with a band gap much smaller

than that of Ti02 are needed. The second condition, that the flat band

potential is more negative than the hydrogen potential, required for photo-

|i galvanic system, is fulfilled for most n-type semiconductors [12].

279
7. References

1] Beck, F. and Gerischer, H. , Ztschr. f. Elektrochemie 63, 500 (1959).

2_ Memming, R. and Schwandt, G. , Surf. Sci 4, 109 (1966).

3' Mehl, W. and Lohmann, F. , Bar. Bunsenges. Phys. chem. 71, 1055 (196?)

4' Gerischer, H. and Mindt, W. , Surf. Sci. 4, 440 (1966).

5\ Memming, R. and Neumann, G. , Surf. Sci. 10, 1 (I968).

6' Memming, R. and Schwandt, G. , Electrochim. Acta 13, 1299 (1968).

7 Gerischer, H. , J. Electrochem. See. 113, 1174 (1966).

8" Justice, D. D. and Hard, R. M. , J. Electrochem. Soc. 118, 1417 (1971),

9] Williams, R. , J. Chem. Phys. 32, 1505 (i960).

10] Gerischer, H, and Mattes, J., Ber. Bunsenges. Phys. Chem. 69, 578

(1965).

[11] Hobbs, B. S. and Tseung, A. C. C. , J. Electrochem. Soc. 120, 766

(1973).

[12] Gleria, M. and Memming, R. , J. Electroanal. Chem. and Interfac.

Electrochem. 65, I63 (1975),

13] Boddy, P. J., J. Electrochem. Soc. 115, 199 (1968).

14] Gissler, W. , Lensi, P. L. and Pizzini, S. , J. Appl. Electrochem.,

in press.

[15] Boddy, P. I., Kahng, D. and Chen, Y. S. , Electrochim, Acta 13, 1311

(1968) . [16] Mollers, F. and Memming, R. , Ber. Bunsenges., Phys. Chem. 76, 469

(1972).

17] Vijh, A. K. , J. Electrochem. Soc. 117, 173C (1970).

18] Memming, R. and Mollers, F. , Ber. Bunsenges., Phys. Chem. 76, 475

(1972).

[19] Mollers, P.. Tolle , H. J. and Memming, R. , J. Electrochem. Soc. 121,

1160 (1974).

[20] Memming, R. , Proc. of Symposium Electrocatalysis , Electrochem. Soc.

Princeton 1974, p. 178, ed. by M. Brieter.

21] Schultze, J. ¥. and Vetter, K. J., Electrochim. Acta 18, 889 (1973).

22' Heusler, K. E. and Schulze , M. , Electrochim. -Acta 20, 237 (1975).

23] Stiitzle, D. and Heusler, K. E. , Ztschr. Phys, Chem. N.P. 65, 201

(1969) . [24] Allard, K. D. , i^Jirens , M. and Heusler, K. E. , Werkstoffe u. Korr. 26,

694 (1975).

25] Gissler, W. , submitted to J. Appl. Electrochem.

26' Hardee, K. L. and Bard, A. J., J. Electrochem. Soc. 122, 739 (1975).

27\ Gohr, H. and Breitenstein, A., Electrochim. Acta 13, 1377 (1968).

28] Memming, R. , Mollers, F. and Neumann, G. , J. Electrochem. Soc. 117,

451 (1970).

[29] Tilak, B. v., Perkins, R. S. , Kczlowska, H. A. and Conway, B. E.

,

Electrochim. Acta 17, 1447 (1972).

30] Bockris, J. O'M. and McHardy, I., J. Electrochem. Soc. 120, 61 (1973).

31" McHardy, H. and Bockris, J. O'M., J. Electrochem. Soc. 120, 53 (1973).

32' Hagenmuller, P., Proj. Solid State Chem. 5, 71 (1971).

33" Dickens, P. G. and Whittingham, M. S. , Quart. Rev. 22, 30 (I968).

'34] Randin, J. P., Electrochim. Acta 19, 83 (1974).

280
National Bureau of Standards Special Publication 455, Electrocatalysis on Non-Metallic Surfaces

,

Proceedings of Workshop held at MBS, Gaithersburg, Md. , December 9-12, 1975. (Issued November 1976)

Oxygen Evolution on Semiconducting Oxides

A.C.C. Tseung and S. Jasem

Department of Chemistiy

The City University

St. John Street

London, E.C.I

The oxygen evolution reaction is of particular interest to secondary metal air batteries and water electrolysis plants. However,

most of the earlier work has been on precious metals and there are no

guidelines for the choice of semicondicting oxides as oxygen evolving

electrodes. In this study, the role of the metal/metal oxide or the

lower metal oxideAugher metal oxide coiple in determining the minimum

voltage required for the evolution of oxygen is enphasized, together

with other essential requirements such as electrical resistivity

electrode micros tructure, corrosion resistance and catalytic properties.

A survey of various metal oxides based on the above criterion suggested

thatNiCo20it is of particular interest and teflon bonded electrodes based

on this material gave over 13,000 A/m^ at 1.63 V vs DHE, 70°C, 5N KOH.

Key words: Nickel cobalt oxides; oxygen evolution, Pourbaix diagram;

semiconducting oxides; spinels; teflon bonded electrodes

1. Introduction

The oxygen evolution reaction on semiconducting oxides has not been studied in detail.

Hoare [l]-"- suggested that the mechanism of O2 evolution in alkaline media could involve:

(O)ads ' (O)ads

^

(02),ds

However, this does not indicate whether such reactions occur on metal oxide surfaces.

In the case of oxygen on Pt^, it has been shown that unless the Ft electrode is anodically

polarized beyond the equilibrium potential of the PtO/Pt02 couple, oxygen evolution cannot

proceed. Similar arguments apply to all other metals, since oxygen evolution can only

proceed when the potential of the electrode is higher than the metal/metal oxide coiple

or the lower metal oxide/higher metal oxide coiple if there is more than one form of oxide.

'''Figures in brackets indicate the literature references at the end of this paper.

281
2. General Considerations

2.1 Metal/Metal Oxide and Lower Metal Oxide/Higher

Metal Oxide Couples

Since Hoare [1] has shovm that in all the metals studied, the oxygen produced initially

reacts with the bare metal surface to form surface oxides and hence the potential at which

oxygen can start to evolve from the oxide surfaces is related to the metal/metal oxide

couple or if there is more than one form of oxide, the lower metal oxide/higher metal oxide

coiple. Thus, the ideal coiple will be lower or similar to the theoretical potential of

the oxygen electrode.

The data for the various oxide couples can easily be obtained from the Pourbaix

Diagrams [2]

.

2.2 Electrical Resistivity

It is impractical to use solid semiconducting oxide electrodes since the ohmic losses

will be too high. However, it has recently been shown [3-5] that teflon bonded porous

electrodes, sipported on nickel screens, are far more efficient than solid electrodes for

oxygen evolution. Since the catalyst/teflon mixture is in close contact with conducting

nickel wires, the likely ohmic losses in such electrodes can easily be estimated if we

assume that almost all the ohmic losses occur in the catalyst/teflon mixture [6] . If we

assume that current is being collected on four sides of the square holes in the nickel

mesh, we have

Where I is the total current, Amps; R is the total resistance of the electrode, ohm ; i, current density, A/on^; L, length of each side of the square hole in the nickel mesh,

cm; p, resistivity of the catalyst/teflon layer, ohm-cm, t, thickness of catalyst/teflon

layer. Taking the typical case of teflon bonded porous electrode supported on 100 mesh

nickel screen, L is 0.04 cm and t is 0.025 cm. Hence if the catalyst/teflon layer has a resistivity of 100 ohm cm, and operating at 0.5A/cm^, the IR loss will be about 0.2 V. However if the resistivity p is 10^ ohm cm, the IR loss will be over 2 V. Hence the

resistivity of the catalyst/teflon layer should be below 100 ohm-cm. In actual practice,

taking into account the increased resistivity due to the presence of teflon particles and

contact resistance between the catalyst powder, the maximum allowable bulk resistivity of

the semiconducting oxide for use in such electrodes will be about 10-15 ohm-cm.

2.3 Corrosion Resistance

It is inportant to know the stability of the candidate oxide materials under the

operating conditions (i.e. pH 14, 1.3-2V). In the absence of detailed experimental data,

the best guide is the Pourbaix diagram [2], whereby stability is defined as when the

equilibrium concentration of the metal ion in solution is less than 10"^ ions/1.

2.4 Catalytic Properties of Semiconducting Oxides

It is difficult to be certain which of the available data on semiconducting oxides

would be of relevance in predicting which oxides are of interest in oxygen evolution

reactions. In any case, it should be able to desorb the oxyeen molecule from the surface

relatively easily. Since the O2 reduction reaction on most of the common oxides has been

282
studied, their activity could be used as a guide-high activity iirplies that oxygen is not

j too strongly chemisorbed on the surface.

2.5 Electrocatalyst Microstructure

In order to maximize the area available for 03^gen evolution reaction, it is necessary

to use teflon bonded porous electrodes [5], Therefore, the electrocatalyst aggregates must

be strong enough to withstand the gas pressure built ip inside the pores; hence, the semiconducting oxides should be prepared by chemical methods to ensure there are solid bonds

betu'een the particles inside each electrocatalyst aggregate [4] . Of the numerous methods

available for the preparation of high surface area semiconducting oxides, freeze drying [7] and co-precipitation methods [8] are particularly useful.

2.6 List of Candidate Materials

Table 1 lists the conductivity, corrosion resistance, oxygen reduction performance as well as the equilibrium metal/metal oxide or lower metal oxide/higher metal oxide couples

of a selected number of candidate materials. The physical properties listed are for the

Table 1. Properties of semiconducting oxides.

Oxide Electrical Corrosion Oxygen Potential vs NHE conductivity resistance

at pH = 14 reduction volt

La/La2O

3

Poor Good Poor -2.069

TizOs/TiOa Poor Good Poor -0.556

Poor Poor Poor -0.666

CraOa/CiOa Poor Poor Poor +1.284

M0O2/M0O3 Poor Poor Poor -1.09

W2O5/WO3 Poor Poor Poor -0.029

MnaOa/MnOa Fair Doubtful Fair +1.014

C02O3/C0O2 Poor Good Poor +1.447

CozNiOi^/CoaNiOit Good Good Good +1.4

Ni203/Ni02 Poor Good Poor +1.434

Ru02/Rl£)^ Good Poor Good +1.387

Rh203/Rh02 Good Doubtful Good +1.73

Pd02/Pd03 Good Good Good +2.03

OsOa/OsO^ Good Poor Good +1.00

IraOa/IrOa Good Poor Good 0.93

PtO/Pt02 Good Good Good +1.7

CU2O/CUO Poor Doubtful Poor K).667

AgO/Ag203 Good Poor Good +1.57

AU203/AU32 Good Doubtful Poor +2.63

higher metal oxides. Obviously, if the oxides whose equilibrium potential is lower than or close to that of the oxygen electrode possess adequate conductivity, corrosion resistance and

ability to reduce oxygen (implying low heat of oxygen adsorption) , they will start to evolve

O2 sooner than other oxides and be expected to give higher anodic current at a given overpotential. An examination of Table 1 shows that it is difficult to get suitable oxides

which can satisfy all the above requirements, but the oxide couples on Ni and Co seem to satisfy most of the requirements except for their high electrical resistivity. However, the

spinel C02 NiOi,, [9] possess relatively low resistivity ("^ 10 ohm cm) and is therefore chosen

for detailed study.

283
RuOa-TiOa would seem to be the next best candidate material. This was substantiated

by recent studies of Srivinsan et al. [10]

.

3. Materials and Methods

The NiCoaOij samples were prepared by freeze drying [9] of mixed nitrate solutions, followed by vacuum decomposition at ZSQOC and subsequent heat treatment at 400OC for 10 hours

in air. The BET area was 41m2/g and x-ray diffraction confirmed that only spinel NiCo204 v/as present. The resistivity of the powder was less than 10 ohm cm. The powder was mixed with

teflon dispersion (ICI Fluon) and painted onto either 100 mesh nickel or platinum screens, dried at 110°C for 1 hour and finally cured for an hour at 300°C to form teflon bonded

electrodes. The catalyst/teflon ratio was 10/3 in all cases. The electrode was immersed in

a thermostated glass cell containing 5N KOH. A piece of 5 x 5 cm gold foil was used as the

counter electrode. The potential of the working electrode was measured against a dynamic

hydrogen electrode (IHE) and the potential was controlled by a 3A Chemical Electronics

Potentiostat. An interrupter technique [11] was used to measure the ohmic drop between the

Luggin capillary and the working electrode.

4. Results and Discussion

4.1 Oxygen Evolution on Metal Substrates

The anodic performances of a 1 cm^ nickel screen and a porous nickel plaque electrode

at 1.638 V vs IHE are shown in figure 1. The performances of both electrodes decreased from

20 X 10^ A/m^ to less than 7 x 10^ A/m^ over a period of 40 hours, suggesting that the growth

of high resistance nickel oxide films has greatly reduced the activity of the electrodes.

0 40 80 120 160 200

Time (h)

Fig. 1. Oxygen evolution on 100 mesh Ni screen and porous nickel

as a function of time.

284
In addition, it is ivXJrth noting that despite the very much higher nickel area in the poroiis

plaque, the activity is similar to that of the nickel screen electrode, suggesting that when

oxygen is produced inside pores, it will have to build up a high enough pressure to expel

the electrolyte from the pores before oxygen could escape; hence only the mouth of the pores

could fuiction^. Thus, it is necessary to use teflon bonded porous electrodes to maximize

the area available for oxygen evolution. Furthermore, the results confirmed our earlier

hypothesis - oxygen evolution on nickel began at about 1.45 V, just above the potential of

lower nickel oxide /higher nickel oxide coiple.

06

X

>>

C

<uctz

o

25 °C 5 N KOH

Electrode held potentiostatically

at 2.0 V vs DHE

_L

0 2 4 6 8 10 12 Time (min)

Fig. 2. Oxygen evolution on Pt screen as a function of time.

The results for oxygen evolution on Pt screen (fig. 2) showed a similar trend - though

the performance decreased to a low^er, steady level after only 20 minutes. The results

indicated that initially sane of the anodic current is used to build up the PtOa layer,

t Once the process has reached an equilibrium state, oxygen evolution proceeds on the PtOa

surface. However, it is worth noting that the over-voltage is very much higher than on

nickel, and this is in agreement with the higher potential of the PtO/PtOa couple.

4.2 Oxygen Evolution on Teflon Bonded NiCo20it Electrodes

The results obtained on nickel screens suggest that they can give a significant amount

I of current initially and cannot be neglected in our study. On the other hand, Pt screen is

far more passive and the oxygen evolution reaction reaches a steady value within 20 minutes.

Figiire 3 shows the performance of teflon bonded NiCozO^ electrode supported on 100 mesh Ni

screens. It is interesting to note that the iR corrected results showed that over 13 x 10^

A/m^ could be obtained at 1.63 V vs THE and to 70°C (fig. 3). The corresponding result for

XiCoaO^ supported on Pt screen is 10.5 x 10^ A/m^ at 1.71 V (fig. 4), since NiCoaOi, alone

contributed to most of the activity in this case. In this study, no attempt has been made

to study the Tafel relationships since it is meaningless to apply the Tafel equation to gas- evolving porous electrodes- -we cannot ascertain the contribution to the overall electrode

polarization from mass transfer effects (emptying of electrolytes inside the pores, blockage

of electrode area by gas bubbles, etc). We plan to study the problem in more detail by

using pKDtentiostatic pulse techniques [12] to eliminate mass transfer effects at a later date.

285
5 N KOH

(Ni screen as current collector) . (Pt screen as current collector)

4.3 Stability of NiCo20i^ Electrodes

Figure 5 confirmed that the performance of teflon bonded NiCoaOi, electrode supported on

Pt screen did not change with time. In practice, Ni screens would have to be used as the

current collector and the temperature of operation would definitely be higher than 250C.

Figure 6 showed that after an initial decay in performance at SQOC, the performance of teflon

bonded NiCoaOi, on Ni screen remained stable and that the initial decay was mainly associated

with the growth of insulating nickel oxide film on the Ni screen. Over a total test period

of 300 hours there were no visible signs of physical deterioration in the electrode and

atomic absorption analysis of the KOH solution after the endurance test did not reveal the

presence of Ni or Co ions (sensitivity of apparatus 2 ppm)

.

286
5^

(V

S25.

20

15

10

25 °C 5 N KOH

Electrode held potentiostatically

at 1.69 V vs DHE (ir corrected)

0 50 100 150 200 250

Time (h)

Fig. 5. Long term test for oxygen evolution on NiCo20i, (Pt screen as as current collector) 5 N KOH, 250C.

1.54 V

~ NiCo20;t (on Ni screen)

(ir corrected) —-i

•% • •

1 .50 V (ir corrected)

5 N KOH 80 °C

Electrodes held potentiostatically

(potentials vs DHE)

1 .58 V

, ,. ,, (ir corrected) Ni screen

1.54 V (ir corrected) ^ , ^ . . . ^ A i t 4 tr-^ A A A A A A

100 150

Time (h)

200 250

Fig. 6. Long term test for oxygen evolution on NiCoaOn (Ni screen

as current collector) 5 N KOH, SQOC.

287
5 . References

T.P. Hoare. The Electrochemistry of Oxygen. Interscience (1968).

M. Pourbaix, Atlas d'Equilibres Electrochemiques ^ 25°C. , Gauthiers-Vi liars, Paris

[1963). P.R. Vassie, A.C.C. Tseung, Electrochimica Acta, 20, 759 (]975). P.R. Vassie, A.C.C. Tseung, Electrochimica Acta, 20, 763 (1975).

A.C.C. Tseung, P.R. Vassie, Electrochimica Acta (in the press).

A.C.C. Tseung, unpublished results.

A.C.C. Tseung, H.L. Bevan, J., of Mat. Sci. ,5, 604 (1970) A. C.C. Tseung, J.R. Goldstein, J. of Mat. Sci., 7,1383 (1972) W.J. King, A.C.C. Tseung, Electrochimica Acta, 19,485 (1974)

S. Srivinsani et al. Paper presented at the NBS Meeting on Electrocatalysis on non- metallic surfaces, Washington, Dec, 1975. K.R. Williams, An Introduction to Fuel Cells, Elsevier, Amsterdam (1966)

B. S. Hobbs, P.R. Vassie, A.C.C. Tseung, in Electrochemical Engineering Symposium,

Inst, of Chem. Eng., Edit. J.D. Thornton, 1.123 (1971).

288
National Bureau of Standards Special Publication 455, Electrocatalysxs on Non-Metallic Surfaces,

Proceedings of Workshop held at NBS, Gaithersburg , Md., December 9-12, 1975. (Issued November 1976)

Mixed Oxides as Oxygen Electrodes

G. Singh, M. H. Miles and S. Srinivasan

Department of Applied Science

Brookhaven National Laboratory

Upton, New York 11973

The irreversibility of the oxygen electrode reaction on metals

and alloys is the main cause of efficiency losses in fuel cells,

water electrolysis cells and some battery systems. Oxygen evolution

always occurs on oxide covered surfaces while the reduction reaction

generally takes place on bare metallic surfaces but can do so on oxides. Since the most likely candidates for reversible oxygen electrodes are oxides, mixed oxides or compound oxides, three types of

oxides - NiOjj, NiCo20^ and Ru02|Ti02 were examined as oxygen electrodes in KOH solution. When these oxide electrodes were prepared by

thermal decomposition of their salts on a substrate, ratios of true to apparent surface areas of the electrodes were quite high (over 200)

and the transfer coefficients, obtained from the Tafel lines for

oxygen evolution in the temperature range 25 - lOO'-'C, were 1.5, as compared with about half that value on nickel electrodes . Above 100°C

NiCo20^ and Ru02|Ti02 electrodes corroded in the strong (50%) KOH

environment. Mixed oxides are more stable than simple oxides in the

region of potential where oxygen reduction occurs. From the values of

the Tafel slopes and exchange current densities for oxygen evolution

and reduction on NiCo204 (a spinel) and Ru02|Ti02, it is clear that

reversibility of this electrode reaction was not obtained. Other mixed

oxides (including spinels and perovskites) will have to be investigated with the aim of developing reversible oxygen electrodes.

Key words: Electrocatalysis ; nickel-cobalt oxide; nickel oxide;

oxide electrodes; oxygen electrode reaction; ruthenium

dioxide/ titanium dioxide; spinel.

1. Introduction

The oxygen electrode reactions are of special importance in water electrolysis, fuel

cells and some battery systems (e. g., zinc-air, aluminum-air, and iron-air) because

activation overpotential at this electrode is the chief cause of efficiency losses in these

devices. Oxygen evolution always takes place on oxide covered metallic surfaces whereas

oxygen reduction occurs on bare metallic surfaces but in some cases on oxides or mixed

oxides. Thus, the most likely candidates for reversible oxygen electrodes are oxides, mixed

oxides or compound oxides. Mixed oxides, including spinels, appear to be better catalysts

for several reactions than simple oxides. One particular feature of mixed oxides is their

increased stability over the simple oxides. In spinels, the mixed valences of the cations

help in the reversible adsorption of oxygen by providing donor-acceptor sites for chemisorption.

With a view to understanding and developing reversible oxygen electrodes, the electrocatalytic activities of a simple oxide (NiO^,) and of the mixed oxides NiCo204 and

Ru02|Ti02 for oxygen evolution and of the mixed oxides for oxygen reduction were determined

289
in the present work. There is some evidence that the spinel NiCo204 is a good electrocatalyst for oxygen reduction [I]-'-. The mixed oxide Ru02|Ti02 has been extensively used in

the chlorine industry [2]. It also has shown promise as a good electrocatalyst for the

oxygen electrode reaction [3].

2. Experimental

The nickel-cobalt oxide electrodes were prepared by two methods. In the first [4],

the two nitrates, Ni(N02)2' 6H2O and Co(N03)2. 6H2O were weighed in the exact proportion of

Ni:Co = 1:2 and dissolved in water. The solution was evaporated to dryness until there were

no more NO2 fumes. The black powder was heated in an electric furnace in air for 48 hours

at a temperature of 325 - 350°C. The spinel, NiCo204, was mixed with Teflon 30 emulsion

and shaken well. Teflon 30 and NiCo20^ were in the proportion of approximately 20:80 by

weight. A nickel screen was dipped in the above solution and a coating of NiCo204 was deposited on it. The nickel screen was heated in an electric furnace for about an hour at 310°C. The electrode thus prepared was used for determining its electrocatalytic activities

for the oxygen electrode reactions.

In the second method, nickel oxide (NiO^^) and nickel-cobalt oxide (NiCo204) were prepared by thermal decomposition on a substrate. Details of the preparation of the nickelcobalt oxide electrode are as follows: Nickel nitrate and cobalt nitrate, mixed in

stoichiometric amounts were dissolved in 1-butanol. A clean piece of a metal (Ni or Ti) or stainless steel was spot welded to a wire of the same metal or alloy and washed. It was heated for five minutes in an oven at 250oc and then dipped in the solution of the nitrates.

It was placed in the oven again at 250°C for 2-3 minutes and this process was repeated

about 10 times. At the end of this operation, the temperature of the oven was raised to 350°C and the electrode heated in it overnight. The electrode was cooled and washed. It had about 1 - 5 mg of NiCo20^ per cm^ of the metal plate. The nickel oxide electrode was prepared in a similar manner. The ruthenium dioxide-titanium dioxide electrodes for the

oxygen evolution and reduction studies were obtained from T. Bomaraju of Hooker Chemical

Company, Niagara Falls.

The electrochemical measurements, to determine electrocatalytic activities, consisted

of steady-state potentiostatic measurements using a PAR (Model //173) potentiostat . The

electrolyte used was 30% KOH. The oxygen evolution studies were conducted at 25, 50, 80 and 100°C while the oxygen reduction studies were carried out at temperatures of 25, 50

and 80°C. For comparison of electrocatalytic activities, the oxygen reduction studies were

also done on platinum at the same temperatures. The true surface areas of the electrodes were determined using cyclic voltammetric techniques. For this purpose, the electrodes were maintained at the highest potential used for oxygen evolution studies. The cyclic

voltammograms were then recorded in a small range (about 100 mv) , starting at a much lesser

potential (say 1 volt vs. RHE) , as a function of sweep rate. As a check, the surface areas

in one set of experiments were also obtained by BET measurements.

With NiCo20, and Ru02|Ti02, preliminary electrocatalytic studies for oxygen evolution were also carriea out at temperatures up to 150°C

.

3. Results

The Tafel lines for oxygen evolution at the four temperatures on nickel-cobalt oxide,

prepared by the thermal decomposition method, are shown in figure 1. From figure 2, it is clear that the electrocatalytic activities of the three oxide systems for this reaction

at 80°C show little difference. Figures 3 and 4 show the methodology of obtaining the true surface areas of the electrodes. By limiting the cyclic voltammogram to a narrow range of

potential, the double layer charging currents are constant, with the anodic current being

equal in magnitude to the cathodic current at each sweep rate (fig. 3). A plot of the

charging current versus the sweep rate is linear, as expected, and it passes through the

origin (fig. 4). Similar results were obtained on the nickel-cobalt oxide electrodes.

The performance of the Teflon bonded nickel-cobalt oxide electrode was much less

satisfactory than that of the electrodes prepared by the thermal decomposition method.

Figures in brackets indicate the literature references at the end of this paper.

290
b

o

>

o

w

<

I-z

UJ

I-o

Q.

UJOO

o:O

UJ

_l

UJ

1.0

0.8

0.6

0.4

0.2

0.2

1 1 1 M M 11 1 1 1 1 M M 1 1 1 1 1 1 1 1 1

1 1 1 1 1 1 111

0

80° C

V I00°C

^ -

IT ^^^^^

1

1 1

1 1 1 1 1 1 M 1 1 1 1 1 1 1 1 1 1 1

0.1 1.0 10 m A/cm^

100 1000

Fig. 1. Potential-current density relations for oxygen evolution

on NiCo-O^, deposited thermally on nickel plate, in 30%

KOH at 25, 50, 80 and 100°C,

CO

I-

_lO

>

1.0

, 0.8-

LlJ o

CO 0.6

UJ 0.2

~i I I I I iii| I I I I 1 1 1

1

1

1—I I II I III

1—I I I 1 1 II

o NiC0204

^ NiOx

Ru02/Ti02

J I

I I I 1 1 1

1

J I

I I

I I

I I 1 1 1 1

1

I I I I I 1 1

1

10 mA/cm100 (000

Fig. 2. Potential-current density relations for oxygen evolution on

NiO^, NiCo^O^ and RuO^ | TiO^ electrodes, prepared by the

thermal decomposition method, in 30% KOH at 80°C.

The stability of the electrodes was considerably better in the latter case.

Though the NiCo204 and Ru02|Ti02 electrodes showed a good performance for oxygen

evolution, and were stable at temperatures belox^ 100°C, they were quite unstable at higher

temperatures. This is probably due to the formation of cobaltates and ruthenates in the

strongly alkaline environment.

291
100

mV vs SCE

150

Fig. 3. Cyclic voltammograms on RUO2 | Ti02 electrodes in IM nitrogen

saturated KOH at four sweep rates. Designations 1, 2, 3 and 4 represent sweep rates of 10, 20, 50 and 100 mVsec"-'-.

SCAN RATE, mV/sec

Fig. 4. Double layer charging current (from fig. 3) versus sweep rate.

Results of oxygen reduction at 80°C on Pt, NiCo204 and Ru02|Ti02 from 30% KOH solution

are graphically represented in figure 5. The current is partly due to the reduction of the

lattice oxide, which is quite evident in figure 6 which shows the current-potential

relations on a nickel-cobalt oxide electrode in 30% at 80°C with oxygen or argon being

bubbled through the electro],yte . The spinel NiCo204 appears to be less stable than the

mixed oxide Ru02|Ti02 in the region of potential where oxygen reduction occurs at appreciable rates.

292
0

-0.2

-0.4

>

-i -0.6 <Z

uj -0.8

I-o

CL

UJ -1.0

o

o

tr

I- -1.2 o

UJ

_lu

0.1 10 100 1000

1 1 1 1 M 1 l| 1 1 1 1 1 1 1 1 11 1 1 1 1 1 1

1

1 1 1 1 1 1 1 1

1

~ — --^Sz;?—^ -

• Pt

O-io——!_ •

\

\

> —

0 Ni COgO^

\

-

° RuO,/TiO„

_ 2 2 \ :

1 1 1 1 M 1 1 1 1 1 1 1 1 1 1 1 1 1 1 1 1 1 1 1 1 1 1 10 100

fi A /cm'

1000 10000

Fig. 5. Potential-current density relations for oxygen reduction

Pt, NiCo^O^ and RuO^ | Ti02 in 307o KOH at 80 C.

P 0.0

0.2

0.4

0.6

0.8

O

a

liJ - 1.0 Q

o

01 1.2

UJ

_iUJ

1 1 1 1 M 1 l| 1 1 1 1 1 1 1 l| 1 1 1 1 1 1 1 1{ 1 1 1 1 1 1

1

• \

1 1 1 1 1 1 1 ll 1 1 1 M 1 III 1 1 1 1

1 10 100 1000

Fig. 6. Potential-current density relations on NiCo20^ in 307o KOH

with oxygen (o) or argon (•) bubbling through the electrolyte.

4. Discussion

4.1 The Oxygen Evolution Reaction on the Oxide Electrodes

The kinetic parameters for the oxygen evolution and reduction reactions on the oxide

electrodes in 30 wt % KOH are presented in Table 1.

293
Table 1. Kinetic parameters for the oxygen evolution

and reduction reactions in 30 wt % KOH.

Electrode Temp (OC)

NiCo20i^ 2550

80

100

Ru02/Ti02 80

NiOj^ 80

Pt 80

NiCo204 80

Ru02/Ti02 80

Tafel slope (V)

Evolution

0.040 1.48

0.040 1.60

0.042 1.67

0.046 1.61

0.045 1.56

0.047 1.49

Reduction

0.033 2.1

0.060 1.2

0.070 1.0

io(A/cm2)b AE^^c

X 10^ Kcal/mol

1.78 9.2

5.31

29.9

54.7

194.5 9.2

71.0 8.2

8.7x10"'* 6.8x10"!

1.0

^Transfer coefficient.

Exchange current densities are given in terms of geometric areas. The roughness factors of various electrodes are, Pt = 2, NiCo20it = 400, and

Ru02/Ti02 = 700. Arrhenius activation energy.

One of the most striking features is that the electrode kinetic parameters for all the

three oxides are quite similar (cf. fig. 2 and Table 1), even though their structural

and electronic properties are quite different. The transfer coefficient (a) for the

oxygen evolution reaction on nickel electrodes (in the form of wires, foils or screens)

is about 0.7 at temperatures below 100°C [5], while in the present study on nickel oxide

prepared by the thermal decomposition method it is 1.5. This value of a on NiO^^,

NiCo204 and Ru02|Ti02 indicates that the rate determining step for oxygen evolution on these electrodes is the second electron transfer step, as for example in the reaction

sequence [3]

+ oh" t (SOH)^ + e~ fast (1)

o

(SOH)^ ^ (SOH)^'^-'- + e" slow (2)

o 7+1 - - 7 (SOH) + OH -> S-0-H...OH S +0 HOH

7+1 - _ 7 (SOH) OH S-0-H...OH S 0 HOH

(3)

where S represents the substrate,

The exchange current densities on these oxides, based on the true surface areas, are

lower than on nickel electrodes, because of the lower Tafel slopes on the former systems.

However, at desirable current densities (say 200 mA cm"^ -geometric) the overpotentials

for oxygen evolution on the oxide electrodes prepared by the thermal decomposition method

are at least 150 mV less than on nickel (wire, screen or foil electrodes) . This result has

practical implications in developing more efficient (from energy considerations) water

electrolysis cells.

294
Ruthenium dioxide on titanium dioxide is only a slightly better electrocatalyst than

, nickel cobalt oxide or nickel oxide. Both these mixed oxide systems are not stable at temperatures above 100°C in 50% KOH. It has been previously suggested 16] that the lower

overvoltage for oxygen evolution on ruthenium in acid solution is due to the presence of

ruthenium ions in solution produced by the anodic dissolution of the metal. In alkaline

solution, it is very probable that the ruthenium ion is not involved in the overall rej action path due to the fact that the other oxide systems as well show similar electrode

, kinetic parameters. It is very probable that in this medium, it is necessary to create

I the conditions for a lower activity of adsorbed hydroxyl radicals, which causes the

i oxygen evolution to be controlled by the second electron transfer step (cf . reaction path

; represented by eqs . 1-3). In practically all water electrolysis cells, nickel is used

I as the oxygen electrode and is in the form of nickel clad iron electrodes, sintered nickel

plaques or screens . Since the present work shows that nickel oxide prepared by the thermal

decomposition method has electrocatalytic activities similar to nickel-cobalt oxide as well

as ruthenium dioxide but considerably superior to nickel in the above forms, the present

approach to preparation of oxygen electrodes, using nickel, is suggested.

4.2 The Oxygen Reduction Reaction on the Oxide Electrodes

The Tafel slopes for oxygen reduction on NiCo20^ and in Ru02|Ti02 are about twice that

on platinum. Based on the true surface area, the exchange current densities for oxygen re- duction on the oxide electrodes are only slightly higher than on platinum. In a previous

study [7], it was shown that oxygen reduction to hydroxyl ions on the spinel NiCo20^ does

not proceed through a peroxide intermediate. The present work (fig. 6) confirms another

investigation [8] that this spinel is not stable in alkaline environments under conditions

of oxygen reduction.

From the values of the transfer coefficients and exchange current densities in Table 1,

it is also apparent that oxygen evolution and reduction occur on these mixed oxides by

different mechanisms. The exchange current densities are considerably lower for oxygen

reduction than for oxygen evolution. It is planned to investigate other oxide systems with the hope of developing reversible oxygen electrodes for regenerative fuel

cells and battery systems.

5. Conclusions

In alkaline solution, nickel oxide (NiOj^) , the spinel NiCo204 and the mixed oxide

Ru02|Ti02 are good electrocatalysts for oxygen evolution but not for oxygen reduction.

The superior performance of these oxides, as compared with that of conventional metal

electrodes (e. g., nickel), for oxygen evolution is due to the method of preparation of

high surface area electrodes. Oxide electrodes prepared by the thermal decomposition

method have a higher electrocatalytic activity and stability than the Teflon bonded electrodes. The mechanism of the oxygen evolution reaction on oxide electrodes, prepared by

the former method, involves a rate determining electrochemical desorption of the adsorbed

hydroxyl radical while on nickel electrodes in the form of wire, foil or screen, the r.d.s.

is probably a chemical desorption of the adsorbed hydroxyl radical. Mixed oxides are more

stable than simple oxides in the region of potential where oxygen reduction occurs. Other

mixed oxides (including spinels and perovskites) will have to be investigated with the aim

of developing reversible oxygen electrodes.

6. Acknowledgements

This work was performed under the auspices of the U. S. Energy Research and Development Administration, The authors wish to thank Miss Patricia J. Dwyer for carrying out

the experimental studies on oxygen reduction and Mr. F, J, Salzano for his interest and

encouragement of the work,

7 , References

[1] King, W, J. and Tseung, A, C, C, Electrochim. Acta 19:485, 1974.

[2] DeNora, 0,, Chem. Ing,-Techn,, 42:222, 1970,

[3] O'Grady, W. , Iwakura, C, Huang, J. and Yeager, E., Proceedings of the Symposium

on Electrocatalysis , M. Breiter Ed., 286, The Electrochemical Society, Princeton,

295
New Jersey, 1974.

Knop, 0., Reid, K. I. A., Sutarno and Nakagawa, Y., Can. J. Chem., 46:3463, 1968.

Miles, M. H., Kissel, G., Lu, P. W. T., and Srinivasan, S., J. Electrochem . Soc,

123:332, 1976.

Burketz, L. D., and O'Meara, T. 0., J. Electroanal. and Interf. Electrochem.,

36:App. 25, 1972.

Shumilova, N. A., and Kruscheva, E. I., paper presented at 10th Intersociety Energy

Conversion Engineering Conf., Wilmington, Delaware, Aug. 18-22, 197 5. King, W. J., and Tseung, A. C. C, Electrochem. Acta, 19:493, 1974.

296
National Bureau of Standards Special Publication 455, Electrocatalysis on Non-Metallic Surfaces,

Proceedings of Workshop held at NBS, Gaithersburg, Md. , December 9-12, 1975. (Issued November 1976)

Surface and E lectrocatalyt ic Properties of Tungsten Bronzes

M. F. Weber and H. R. Shanks

Energy and Miner'al Resources Research Institute

Iowa State University, Ames, Iowa 50010

Tafel plots have been obtained for the oxygen reaction in acid

solution on several alkali tungsten bronzes having different crystal structures. Platinum doped sodium tungsten bronze crystals were

studied and the results compared with those of platinum free crystals

of the same composition. In both cases sodium tungsten bronzes were

found to be poor electrocatalysts for the cathodic reduction of oxygen.

Similar results are reported for other alkali tungsten bronzes and for

tungsten trioxide. Anodic treatment of the crystals affected the

electrocatalyt ic activity of only the sodium tungsten bronze and the

effect was a negative one. Cyclic voltammetry was employed to study

the effects of the anodic treatment which created a sodium depletion

layer on the sodium tungsten bronze surface. The existence and depth

of the sodium depletion layer was determined by an Auger Electron

Spectroscopy depth profile.

Key words: Tungsten bronzes; oxygen reduction; electrocatalyst;

Auger Electron Spectroscopy; sodium depletion; Tafel plots.

1 . I ntroduct ion

Tungsten bronzes are nons toch iomet r i c compounds of the form M^WO^ , where M is a metal

ion and x is a number between 0 and 1. The limiting composition "MWO," has the perovskite

structure (ABO^). As the x-value or metal ion content of the bronze decreases, the structure passes through several different phases. For example, in the sodium tungsten bronze

system (Na WO,) a cubic structure is found for x in the range 0.49(x(0.93, a tetragonal

phase for 0.25(x(0.49, a second tetragonal phase for 0.07(x(0.25, an orthorhombic phase

for 0.01(x(0.07 and the monoclinic phase of WO^ for very low x values [1]^. In general,

the crystal symmetry decreases as the x-value decreases. Most of the tungsten bronzes are resistant to acids. In particular, sodium tungsten bronze is not visibly attacked even

by aqua regia over very long periods of time. However, all tungsten bronzes are attacked

by alkaline solutions.

In recent years, reports [2-5] of high electrocatalytic activity for oxygen reduction

by NaQ yWO^ crystals containing traces of platinum stimulated a large research effort to investigate this effect and substantiate these reports [6-9]. The later research results

differed from the earlier reports, in that no exceptional electrocatalytic properties of

NaQ yWO^ for oxygen reduction were found, even with crystals doped with up to 1000 ppm

pi at I num.

^This work was supported in part by a grant from the National Science Foundation.

^Figures in brackets indicate the literature references at the end of this paper.

297
2. Experimental

The alkali tungsten bronzes were prepared by the electrolysis of a fused salt of the

appropriate alkali tungstate and tungsten trioxide as described by Shanks [1]. The crystals

of WOt were prepared by vapor transport, following the techniques of Sawada and Danielson [13]. Since Bockris and McHardy [5] have suggested that AI2O3 and platinum together as impurities

affect the el ect rocatal yt i c properties of the sodium tungsten bronzes, a special effort was made to eliminate any uncontrolled source of these impurities. High purity sodium carbonate

and tungsten trioxide were used as starting materials for the growth of undoped and platinum

doped (up to 800 ppm) high purity sodium tungsten bronzes. Typical mass spectrographic

analysis results for crystals grown from this material and for the starting materials are shown in Table 1. From these results it appears that Mo, Fe, and Ti tend to concentrate in the crystals during electrolysis since the concentration found in the crystals were greater

than in the melt. In addition, the impurity concentrations decreased

Table 1, Mass Spectrographic Analysis of Starting Material and Na WO, Crystal

(concentrations in ppm). ^

Impurity WO3 Na2C03 NaQ 7WO3 Impurity

Source

C 30 S^

Al 2 0.2 0.2 s Si <2 Electrodes

P 500 <0.2 0.2 S

CI 10 60 9 S K 10 2 70 Crucible

Ca 1 <3 100 Crucible

Ti 0.7 3 50 S Fe 3 5 20 S Mo 30 <5 100 S Au 0.3 0.5 <0.1 Electrode

Pb 3 <1 <0.3 S V <0.06 <2 1 S Co 0.1 13 <0.1 S Ni 0.3 2 0.9 S Pt <10 <3 <0.6 S

^Starting material.

as successive crystals were grown from a given melt, indicating depletion of the impurity

in the melt. Platinum doping of the Naj^WO^ crystals was achieved by using a platinum

anode in the melt, and vanadium doping by adding V2O5 to the melt.

The crystals used in the experiments were all mounted in teflon and received a metallographic polish with different sizes of alumina down to 3 Mm. The sample and sample

holder were then cleaned with isopropanol and then with boiling conductivity water. Electrical connection was made to the back of the samples by means of a stainless steel

2/56 threaded screw which made physical contact with the crystal as shown in figure 1. All data are given with respect to the apparent surface area of the electrode.

All electrochemical studies were done in a standard three compartment cell of pyrex

construction, using triple distilled water and ultra pure acid. The electrolyte used was

0.1 NH2SO/1 in all cases. The reference electrode was a palladium bead saturated with

hydrogen and immersed in an electrolyte having the same composition as that in the main

compartment of the cell. This reference electrode was constructed using the design of

McHardy [1^] and is shown in figure 2. The reference electrode was separated from the

main compartment by means of a Luggin capillary and closed stopcock. The counter electrodes for all experiments was either a sodium tungsten bronze crystal with an exposed area

298
6mrr, PYREX

TUBING

-6 mm PYREX TUBING

TEFLON

MOUNT

SAMPLE

-Pt-Pd JUNCTION SEALED

IN GLASS

/Pd WIRE

Pd PELLET

Fig. 1. Sample mounting and electrical

connection

.

Fig. 2. Palladium-hydrogen reference

electrode

.

of about 1 cm^ or a spectroscopic grade graphite rod. Before taking data the solution was pre-electrolysed potent iostat i cal ly with a gold electrode held at 1.3 V NHE. In this manner, no impurities were introduced into the solution by possible dissolution of the pre- electrolysis electrode.

Extensive efforts were made to purify the gases bubbled into the cell. The starting

material for the oxygen gas was Air Products "zero oxygen" which was further purified by

passing it over hot platinized asbestos. The trace amounts of COj generated by this

process were removed by then passing the gas over moist KOH and then through two traps

containing Linde type 13x molecular sieve. The gas used to provide an inert atmosphere

in the cell whenever needed was Matheson "zero helium" which was further purified by

passing it over heated oxide free calcium and magnesium chips. The helium was then passed

through two liquid cold traps, one containing activated charcoal and the other, Linde

type 13x molecular sieve.

The Tafel plots v/ere obtained by the constant current method, starting at 10 A/m^

and decreasing the current in steps to 10"^ A/m^. Individual data points were taken at intervals of two hours or greater to allow the potential to reach equilibrium. The currents

were generated and the potentials measured using a PAR model 173 potent iostat/gal vanostat

in the galvanostatic mode. The cyclic voltammetry curves were obtained by using the same

apparatus in the potent iostat ic mode with a triangular wave input from a Wavetek model 111 signal generator. All measurements were made at room temperature.

3. Results

Tafel plots have been obtained for oxygen reduction in acid solution on various alkali

tungsten bronzes with different crystal structures. The bronzes studied included hexagonal

^^0.33^^3^ tetragonal I NaQ ^^WO, , high purity cubic Hbq -^WO^ undoped and doped with up to

800 ppm platinum, cubic Na« 7W0-3 doped with 0.12% vanaciium, cubic Na^^ qnWO-, . cubic and

299
tetragonal II Li WO^. and single crystal monoclinic WO^. The rest potentials varied greatly

from sample to sample, with potentials (vs. NHE) ranging from 0.62 V for tetragonal II Li^W03 and 0.45 V for WO3 to -0.17 V for hexagonal Kg 33WO3.

Cathodic measurements were performed on freshly polished crystals both before and

after prolonged anodic polarization, since anodic potentials were thought to alter the

x-value of a bronze surface. Typical cathodic Tafel lines obtained on the crystal before

any anodic treatment are shown in figure 3. Data for all the samples are not shown because

many samples exhibited similar Tafel lines. The Tafel lines for NaQ gWO, and Nag yWO, are the same, and remain unchanged when doped with platinum (8OO ppm) or vanad i urn* (0.12%)

.

All noncubic samples and cubic Li^WO^ exhibited higher overpotent i al s than Nag yWOo , and are the same as the values for WO3 , except the hexagonal Kg 33WO3 crystal which exnibited

still higher overpotent i a 1 s than WO3 . Very high overpotent I a1s were observed in the current range 10" A/m^ to 10 A/m^ on all samples including Pt-doped (8OO ppm) Nag^yWO^

and very high purity Nag yWO^ . Below 10"^ A/m^ the potentials drifted to the rest potentials listed above suggesting the Tafel relationship no longer holds. Tafel lines

for all samples extended below zero volts NHE, the lowest curve being the one obtained

for Kg 33WO3 , which had a rest potential of -0,17 V. For this sample, the Tafel line for

oxygen' reduct ion was only 100 mV above the Tafel line for hydrogen reduction.

Anodic current-voltage measurements were performed on all samples, but no Tafel lines

as such, could be obtained. The overpotent ials increased rapidly with increasing anodic

current and, for many samples, exceeded five volts at a current density of 10 A/m , However, at low current densities the potentials dropped well below 1.0 V NHE. On the

Kg 33WO3 sample the potential dropped to 0.5 V NHE at the rather high current density of

lOA/m^. These low potentials may be due to a corrosion mechanism.

In the following discussion the anodic current-voltage measurements will be referred

to simply as an "anodic treatment," and the samples treated in this manner as "anodized"

samples. The maximum anodic current used on all samples was 10 A/m^. The anodic treatment

had no effect on the cathodic Tafel lines of any crystals except cubic Na^WO^ crystals, and even there the effect was not large. The Tafel lines for the cubic Na^WO^ crystals, before

and after the anodic treatment, are shown in figure k. The overpotent ials of anodized cubic

Na^^WO^ samples are close to those of WO3 and the noncubic bronzes.

~I

o PLATINUM

Fig. 3. Representative Tafel plots of unanodized tungsten bronzes compared

to a plot for pure platinum ( represents WO3 and all noncubic bronzes

except Kq 33WO3)

.

300
CURRENT A/cm

Fig. 4. Tafel plots of anodized and unanodized cubic NaxW03 compared to a plot for pure platinum. The curves represent both undoped crystals,

and crystals doped with Pt and V.

Cyclic vol tammetry was performed on samples of each crystal structure. Very different

results were obtained for the cubic Na^WOj crystals when compared to the noncubic crystals.

A typical curve for an anodized sample of NaQ gWO^ is shown in figure 5. A well defined

oxidation-reduction peak occurs near 0.2 V NHE as has been reported in the literature [12].

However, the curve for the same crystal before the anodic treatment is considerably different. The peak current at a given sweep rate is an order of magnitude smaller before the

anodic treatment is performed, as shown in figure 6. This oxidation-reduction peak has been

attributed to the movement of hydrogen ions in and out of the crystal lattice [12], The

301
data presented here are consistent with this idea because the anodic treatment removes many

of the sodium ions from the lattice, allowing more hydrogen ions to enter the lattice.

For the noncubic crystals the oxidation peak broadens and the current density is much

larger at higher potentials than for cubic NaxW03 crystals. The prolonged anodic current is nost likely a corrosion current, and is quite large on the Kg j^WO^ crystal. On the noncubic samples the reduction peak became only a slight shoulder superimposed on the hydrogen

evolution peak. The cyclic voltammetry curves for a WO^ crystal, shown in figure 7, are representative of the curves of most of the noncubic bronze crystals and cubic Li^WO^. The

only difference on some crystals was an even broader oxidation peak. The peak currents were not affected by the anodic treatment except on Kg 33WO2 which had slightly higher peak currents after the anodic treatment. All the cyclic voltammetry curves, both before and

after the anodic treatment, are stable curves. For example, the curves for the unanodized

Nag gWOj sample are from approximately the 100th cyclic sweep.

0.0 0.4 0.8 1.2 VOLTS (NHE)

Fig. 7. Cyclic voltammetry curves for WO3 (also representative of curves for noncubic bronzes)

.

The extent and depth of sodium depletion from a cubic sodium tungsten bronze crystal,

with an x-value of 0.70, has been determined by Auger Electron Spectroscopy. Depth profiles

of the sodium-tungsten ratio were obtained on both a freshly polished crystal and on a

^ polished and anodized crystal. The latter sample was anodized with a current of 10 A/m at h V for 2 hours. The results of the Auger depth analysis are shown in figure 8. The

freshly polished crystal contained a constant sodium x-value of 0.70 at all depths. The

1.0

0.8

5 0.6-

g

§ 0.4

0.2

«HPo--o-o- -O —o —o —o—o-

/

-a—0—0'

Na^WO,

o POLISHED

o POLISHED a ANODIZED

I \ \ I L

Fig. 8. Auger electron spectroscopy

depth profile of the sodiumtungsten ratio for cubic

Nao.7W03.

400 800 1200 DEPTH (&)

1600

302
i surf ace layers of the anodized crystals were not pure WO, as expected, but contained some

I sodium. The x-value at the surface of the anodized crystal was about 0.25 and increased

'to 0.33 at a depth of 50 8. The x-value was constant at 0.33 to a depth of 500 8 and then

I increased gradually to the bulk value of 0.70 at a depth of 1000 8.

4. Conclusions

I From the evidence presented here in the Tafel plots it can be concluded that doping

I of cubic sodium tungsten bronzes with platinum does not affect their el ect rocatal yt i

c

I activity for oxygen reduction in acid solutions. Although none of the other tungsten

'bronzes studied here were doped with platinum, the results are expected to be the same

due to the similarity of their Tafel plots with those of WO^ and undoped cubic Na^WO^.

Although a more extensive Auger analysis of anodized cubic Na^WOj crystals is needed,

it can be concluded that the crystals of cubic Na^WO^ do not have a layer of WO^ on their

I surf aces , even when moderately anodized. The x-value of the sodium depletion layer, as determined by the Auger analysis, is near a critical x-value for lattice stability and the

i metal -semiconductor transition. A further decrease in the x-value (<0.25) would require

la major change in the lattice to the tetragonal II structure. The energy requirements of

this transition may prohibit the formation of a lower x-value bronze or WO^ on the crystal

surface.

5. References

_1] Howard R. Shanks, Journal of Crystal

Growth I3/iit. ^33 (1972)

l2] D. B. Sepa, A. Damjanovic, and J. O'M. Bockris. Electrochim. Acta, 1

2

, Ihe (1967)

[3] A. Damjanovic, D. Sepa, and J. O'M. Bockris, J. Res. Inst. Catalysis,

Hokkaido Univ. , _I6, 1 (1968).

[4] J. McHardy and J. O'M. Bockris, In "From Electrocatalys Is to Fuel Cells,"

G. Sandstede, Editor, p. 109, University

of Washington Press, Seattle, Washington

(1972).

Jl J. O'M. Bockris and J. McHardy, J. Electrochem. Soc, , ]20_, 61 (1973)

'6] B. Broyde, J. Catalysis, _l_q, 13 (1968)

,7] R. D. Armstrong, A. F. Douglas, and

D. E. Williams, Energy Conversion,

JI, 7 (1971)

'8] J. Heffler and H. Bohm, Metaloberf 1 aeche-Angew. Electrochem., 27 , 77 (1973)

[9] Jean-Paul Randin, J. Electrochem. Soc,

121, 1029 (1974)

[10] M. V. Vojnovic, D. B. Sepa, and D. S. Ovcin. Croat. Chem, Acta, M+, 89 (1972)

[11] J. McHardy and J. O'M. Bockris, J. Electrochem, Soc, 120, 55 (1973)

[12] J. P. Randin, A. K- Vijh, and A. B. Chughtai, Ibid, 120, 117^ (1973)

[13] S. Sawada and G. C. Danlelson, Phys,

Rev, _M3, 1005 (1959)

[14] J. McHardy, Ph.D. Thesis, University of

Pennsylvania, Philadelphia, Pennsyl\ania

(1972)

303

National Bureau of Standards Special Publication 455, Electrocatalysis on Non-Metallic Surfaces

,

Proceedings of Workshop held at NBS, Gaithersburg , Md. , December 9-12, 1975. (Issued November 1976)

The Electrochemistry of Molybdenum Oxides

Jean Horkans, M. W. Shafer, and R. Peekema

IBM, Thomas J. Watson Research Center

Yorktown Heights, New York 10598

The general electrochemistry of four oxides of mylybdenum, M0O2,

^°4'-*ll' ^°8^23j ^'^^ ^09026' ^^'^ reduction of oxygen on these oxides,

in 1 N H2SO4 have been investigated on natural faces of single crystals.

These compounds are found to undergo electrochemical oxidation and re- duction in the potential range between 0.13 and 0.73 V vs NHE. Oxygen

does not influence the oxidation of M0O2 and Mo^O]^^, but promotes the

anodic dissolution of Mog023 ^'^'^ ^°9*-'26 • Oxygen reduction is supported

on all these oxides at low currents and high overpotentials . Large

effects of crystal orientation are observed. The chemical composition

of the crystal has a much smaller effect. This is apparently due, at least in part, to a change of surface composition with changing electrode

potentials. Thus, at a given potential, the compositions of the surface

layers of all samples may be similar.

Key words: Electrochemistry; molybdenum oxides; orientation effects;

oxygen reduction.

1. Introduction

Stable metal oxide electrodes offer a possible alternative to the use of metals as cata

lysts for oxygen reduction. The use of an oxide may allow the minimization of corrosion and

offer the possibility of investigating both the oxygen reduction and oxygen evolution on sur

faces of the same chemical composition. One such material to be investigated was sodium

tungsten bronze [1] , but it was disappointing as an oxygen catalyst [2]. A logical group o materials for a systematic investigation as oxygen electrodes is the transition metal oxides

which offer a range of chemical and electronic properties. Of these, the oxides of molybdenum and tungsten are particularly attractive because of their good electronic conductivity

and their stability.

I Although in a real catalyst it is necessary to have high surface areas to maximize the

t; current, a basic understanding of catalysis and the nature of interactions at catalyst surfaces may be advanced by an investigation of electrode processes at single crystal faces.

Single crystals of M0O2 , Mo^O^^, Ho^O^^, and

^'^^^2S ^^^^ been grown and investigated as

electrodes. In this paper we discuss tne ability of these surfaces to sustain the 0^ reduction reaction as a function of both chemical composition and crystal orientation.

2. Experimental

Single crystals of the four oxides studied were grown by vapor transport using Mo metal

and MoO^ powder as starting materials and I^ as the transporting agent. Crystals of the

desired compositions were obtained by varying the Mo/MoO^ ratio and the growth temperature,

which was between 790°C and 950°C [3]. Crystal habit could also be changed slightly by adjusting the growth conditions. The size and quality of the crystals varied with composition,

Numerals in brackets indicate the literature references at the end of this paper.

305
Mo^O^^ being the largest, with faces up to several cm2 and Ko^O^^ being the smallest (3-4 mm^

and poorest quality. Crystal morphology also changes markedly with composition. The crystals were identified and oriented using X-ray diffraction methods.

Contact was made to the crystals by attaching a Au wire with Ag-filled epoxy. The Au

wire, the contact, and all of the crystal except the face of interest were then masked off

using Miccrostop (Michigan Chrome and Chemical Co.), a polyvinyl ketone used as a plater's

mask. The inertness of this material under the conditions of the investigation was established by partially masking off wires of Pt and Au and running their cyclic voltammetry in 1 N

H„SO, . No anomalous features attributable to contamination were observed. In the case of

2 4 several of the molybdenum oxide crystals it was possible to examine several natural faces of

the same crystal by cleaning the crystal and remasking.

Measurements involved potentiostatic linear sweep and steady state measurements with a three-electrode system. All measurements were done in 1 N H2S0^ through which either He or

N2 was being bubbled. A three compartment cell was used, the compartments being separated

by glass frits. The counter electrode was a Pt foil and the reference electrode was Hg/Hg2S0^

All potentials are quoted vs. the normal hydrogen electrode.

The steady state data were taken with the electrode potential being changed in the negative direction. Each point represents 10-15 minutes. Data are not always reproducible when

the direction of the potential change is reversed, as will be discussed below.

3. Results and discussion

The molybdenum-oxygen system contains a number of compounds, of which MoO^ and MoO^ are best known. Between these compositions, five other distinct phases have been identified and

their crystal structures have been determined by Magneli [4,5] and Kihlborg [6]. The

structural studies have shown that all the compounds can be classified as belonging to one

of several basic structural types. The four compounds we studied are of either the rutile

or ReO^ type. The M0O2 is a rutile type, which is common to many of the transition metal

dioxides and has the molybdenum in only one coordination, i.e. MoO^ octahedra, which share

both corners and edges. The high conductivity of this compound is due to the short Mo-Mo

distance along the direction in which the MoO^ octahedra are joined by edges (c axis).

The basic structural unit of the other oxides, ^o^'-^j^^' '^°8^23' ^°9^26' °^ ReO^ type. In this structure the MoO^ octahedra are linked exclusively at the corners.

The oxygen deficiency of these compounds, i.e. the deviation from pure MoO , can be accounted for by connecting MoO^ units in other ways. In ^o^O^^ and Mo the appropriate fraction share edges, and in Mo^O^^ there are MoO, tetrahedra joining the basic ReO -type slabs. A detailed description of these structures and their relationship to each other is given by

Kihlborg [6].

The transport properties of the four oxides have been measured by the four-probe method

as a function of temperature. A summary of their room temperature resistivities and the

pertinent structural data is given in table 1.

As expected, the resistivities increase as the composition approaches the stoichiometric

MoO^, which Is a rather good insulator. The conduction mechanism of the ReO -type oxides is

complex and presently not understood.

The cyclic voltammetry (CV) of two faces of MoO in 1 N H2SO is shown in figure 1. The

solid line indicates the behavior in the absence of 6^ and the dashed line that in 0 - saturated solution. Even in the absence of O2 , currents flow over the entire potential

range shown. Although the general appearance of the CV for the two faces shown here is

similar, the currents are more than an order of magnitude larger for the (101) face than

for the (100) face. The (211) face exhibits currents intermediate to the two faces of figure

1. A reduction of the surface occurs cathodic to 0.4V and oxidation occurs beyond

0.65V in the anodic sweep. In the absence of O2, the total charge_consumed in oxidation is

essentially equal to that consumed in reduction on the (100) and (2ll) faces. However, on the (101) face the cathodic charge is substantially larger than the anodic charge. The shape

of the voltammograms is quite dependent on the potential range scanned.

306
Table 1. Structural data and resistivities of molybdenum oxides.

Compound Symmetry Lattice constants

(X)

Resistivity (300 K)

(ohm cm)

MoO.

^°4°11

^°8°23

^°9°26

monoclinic^ a 5.61

b 4.85 1.2

c _ 5.62

orthorhombic^ a 24.49

b 5.45 2 X

c 6.75

monoclinic'^ a 16.88

b 4.05 6 X

c 13.39

monoclinic^ a 8.14

b 11.89 4 X

c 21.23

.-4

,-1

Brandt and Skapski [7] Kihlborg [6].

When O2 is present, the anodic current is basically the same as in He-saturated soIut

tion, but larger reduction currents are observed due to oxygen reductions. The proportional

increase in current on addition of 0„ is of the same order on the (101) , (100) , and (2li)

faces but the absolute magnitude of the increase is quite different.

MoO

Fig, 1. Cyclic voltammetry of the (101) and (100)

faces of M0O2 in He-saturated ( ) and

02-saturated ( ) 1 N H2SO4. Sweep speed

10 mV/s.

eo

0.6 04 02

E(Vvs NHE)

Steady state measurements for 0„ reduction on MoO„ are summarized on figure 2. The

background currents in the absence or 0^ have been subtracted from the total currents in

i 02-saturated solution to obtain the data of figure 2. There is no substantial linear region

and a limiting current is reached at quite low current values. Cathodic to O.lOV the

background currents become larger compared to the O^-reduction currents. Anodic to 0.45V

an anodic background current flows.

307
Reproducibility of the steady state measurements is poor due to irreversible changes

in the oxidation state of the electrode surface with changes of potential. The curves in

figure 2 will retrace when the direction of the potential change is reversed until a potential is reached where large anodic currents associated with oxidation of the M0O2 surface

begin to flow.

Figure 2 thus indicates that the 0^ reduction reaction can be observed on these sur^

faces, uncomplicated by reactions of the electrode surface itself, only between about 0.30

and 0.45V. The overpotential for 0 reduction is high. Both the magnitude of the current

and the overpotential are of the order of those reported by Randin [2] for the Pt-free

Na WO^. X 3

i(a/cm2) E(VvsNHE)

Fig. 2. Steady state masurements for O2 Fig. 3. Cyclic voltammetry of the (100) and

reduction in 1 N H2S0^ on several (111) faces of Mo402]^ in He-saturated

faces of a M0O2 single crystal. ( ) and 02-saturated ( ) 1 N H2SO4. Sweep speed 10 mV/s.

A complicating factor in the interpretation of the steady state measurements is the

fact that, due to the oxidation and reduction of the surface which are apparent from the

CV, the nature of the surface at any given potential is ill-defined. It cannot be assumed

to consist only of Mo CIV). A rather large effect due to orientation is observed' O2

reduction currents are more than an order of magnitude higher on the (101) face than on

the (100) face. All of the faces studied support only small 02-reductlon currents.

The CV of two faces of Mo^O^^ is shown in figure 3. As in the case of M0O2 , significant currents flow over the entire potential range studied, both in the presence and absence

of O2. These currents are larger for the (111) face than for the (100) face. Oxygen causes larger cathodic currents to flow on the (111) face but has little effect on the (100) face.

The steady state data on these two faces are shown in figure 4. The results are similar to those obtained on M0O2. Small O2 reduction currents are observed between about 0.25

and 0.40V. These are larger on the (111) face than on the (100) face. Cathodic to 0.15V

the background current becomes large compared to the O2 reduction current.

The higher oxides, MOoO„- and Mo„0„,, show somewhat different behavior in the cyclic

o y ZD voltammetry (fig. 5) than M0O2 and Mo^O . (Only one face of each of these crystals was

examined.) Oxygen increases not only the cathodic current, but also the anodic current.

In the absence of 0^, the anodic and cathodic charge consumed during one complete cycle are

equal. When O2 is present, the anodic charge is significantly larger than the cathodic

charge. Apparently O2 promotes the dissolution of ^'^o'^23 ^^'^ ^°9^26 ^^'^ increased cur- rents observed are apparently due to dissolutions. The steady state data for MOqC- and

308
Fig. 4. Steady state measurements for O2

reduction on 1 N H2S0^ on the

(100) and (111) faces of a Mo40j^3^

single crystal.

Fig. 5. Cyclic voltammetry of unidentified

single crystal faces of Mog023 and

M09O25 in He-saturated ( ) and

02-saturated ( ) 1 N H2SO4.

Sweep speed 10 mV/s.

^°9^26 same general trends as for MoO. and Mo^O^^, but they apparently are less

reliable because of reduction of the oxide surface.

As mentioned earlier, the shape of the CV depends strongly on the potential range over which the electrode is cycled. Figure 6 shows the CV of a number of samples, which have

been cycled only between 0.43 and 0.73V. These potentials are not sufficiently cathodic

to cause significant reduction of the surface. Comparison with figures 1, 3, and 5 shows

that the current-voltage characteristic in this potential region is substantially different

than when the cathodic potential limit of the scan is 0.13V. Anodic currents occuring in

figure 6 are apparently due to the anodic dissolution of the oxides. This is found to be

highly dependent on both the chemical composition of the oxide and the crystal face being

studied

.

The higher oxides MOgO^^ and MOgO^g show no tendency to dissolve anodically in this

potential range. Substantial dissolution currents flow on '^o^'-'ii above ^ 0.65V. These are highly dependent on orientation. The (100) face of Mo^O^^ exhibits an unusual oscillatory

current which Peekema and Shafer [8] have attributed to dissolution of the 12A layers perpendicular to the [100] direction which make up the Mo^O^^ structure. Since the [111]

direction has no simple relationship to the periodicity m the structure, no structure in

the CV is seen for the anodic dissolution of the (111) face.

There is no periodicity of this type in the M0O2 structure. All faces show CV in this

region similar to that for (100) as seen in figure 6, although the magnitude of the current

varies.

The question of the nature of the surface reactions which occur in the absence of 0^ in

solution has been addressed only in a preliminary way to date. The large number of possible

oxidation states which can be assumed by molybdenum complicate this process. Pourbaix [9] indicates that at the more cathodic end of the potential region scanned, it may be possible to form Mo (III), which would probably exist in solution. The formation of Mo (III) is unlikely

in our solution because sulfate is not a good complexant for this species: its absence has

been confirmed by chemical analysis for molybdenum in solution after severe cathodic treat309
10 r 0

-10^ 400 r

-400 -

Eo

100

-10 -20

-30 -40

E ( V vs NHE

)

Fig. 6. Dissolution of molybdenum oxides. Cyclic

voltammetry of M0O2, Mo^On, Mog023, and

Mog025 samples in He-saturated 1 N H2S0^

which have never been cycled into potential

regions where cathodic currents flow.

MoO,

Mo^^ (3d 5/2)

1/

Mo^^(3d 3/2)

240.4

!

>-CO

240.6

220.4

M0O2 AS GROWN (100)

Mo''M3d 5/2)

0 2p BANDS

-Mrf'^(3d 3/2)

Mo 4d BANDS

2206

M0O2 ELECTROCHEMICALLY TREATED

ENERGY 2206

Fig. 7, ESCA data for as-grown MoO-j and

for M0O2 as-groiim and after

electrochemical cycling.

310
ment of a sample of Mo^O^ . Another possible oxidation state of molybdenum is +5, which is

stable in the absence of 0^. An EPR study of Mo.O^^ in H2S0^ showed Mo (V) formation only

when the sample was being actively dissolved anodically. Thus, in a potential range in which

anodic dissolution does not occur, it appears that only the +4 and +6 oxidation states of

molybdenum need to be considered in dealing with the electrochemistry of its oxides.

Qualitative information about the effects of electrochemical cycling on the electrode

surface was obtained using ESCA, as illustrated in figure 7. At the top of the figure are shown the Mo (VI) 3d peaks and the oxygen 2p bands of a sample of as-grown MoO^. Beneath

this are the same spectral regions of two M0O2 ClOO) samples, as-grown and after extensive

cycling between 0.13 and 0.73V, (The last potential seen by the electrode before removal

from solution was 0.43V.) The as-grown M0O2 sample is seen to have a considerable amount

of Mo (VI) on the surface. After cycling, the amount of Mo (VI) is reduced and the Mo (IV)

has increased, as indicated by the growth of the Mo (IV) 3d peaks relative to the Mo (VI) 3d peaks and the increase of the molybdenum 4d band. This surface was in contact with air for

several hours before measurement, so the actual amount of Mo (VI) at the time of removal

from solution may have been even less. Similar results were obtained for Mo^Oj^^ although

the relative increase of Mo (IV) is less dramatic. More quantitative information of this

kind is certainly desirable. It would require instituting better handling techniques in

transferring the sample from the solution to the vacuum chamber and more careful potential

control

.

4. Conclusions

The electrochemistry of the molybdenum oxides shows many broad similarities among MoO^

,

Mo^O^^, HOgO^ , and Ko^O^^. In the potential range between 0.13 and 0.73 V in 1 N H2SO

these compounds undergo a highly irreversible reduction and oxidation of the surface. These

reactions lead to high background currents in the steady state measurements and interfere

with the measurement of O2 reduction except in the narrow range between about 0.2 and 0.4 V.

Oxygen reduction occurs on these surfaces at a high overpotential . Only small 0„ reduction currents are sustained. Oxygen promotes the dissolution of Mog022 and MOg02^. All

results are highly dependent on crystal orientation.

The fact that all these oxides exhibit similarities in their electrochemistry may not

mean that chemical composition and electronic conductivity do not play in important role.

Instead, it may mean that the surface oxidation and reduction reactions that occur in this

potential range cause all the surfaces to have similar composition at a given electrode

potential, independent of the composition of the bulk. The lack of specific information

about the surface composition makes the modeling of the interactions at the surface dangerous.

5 . Acknowledgement

The authors gratefully acknowledge the assistance and support of R. Figat for crystal

growth, T. Penney for resistivity data, R. Pollak and H. Hughes for ESCA data and interpretation, R. Title and W. Fitzpatrick for EPR studies, B. Gilbert for wet chemical

analysis and J. Angilello for crystal orientation.

6. References

[1] For example, D. B. Sepa, A. Damjanovic, and J. O'M. Bockris, Electrochem. Acta , 12 , 746 (1967); J. McHardy and J. O'M. Bockris, J . Electrochem. Soc

.

, 120 , 53 (1973);

J. O'M. Bockris and J. McHardy, ibid. , 120 , 61 (1973).

[2] J. -P. Randin, J. Electrochem. Soc , 121 , 1029 (1974).

[3] M. W. Shafer, R. A. Figat, and T. Penney, to be published.

[4] A. Magneli, Acta Chem. Scand. , 2, 501 (1948).

[5] A, Magneli, ibid. . 2, 861 (1948).

311
[6] L. Kihlborg, Arkiv fHr Kemi , 21^, 471 (1963).

[7] B. G. Brandt and Skapski, Acta Chem. Scand

.

, 21 , 661 (1967).

[8] R. M. Peekema and M. W. Shafer, Abstract no. 59, Electrochemical Society Meeting,

Dallas, October, 1975.

[9] M. Pourbaix, Atlas of Electrochemical Equilibria (Pergamon, Oxford, 1966) p, 275.

312
National Bureau of Standards Special Publication 455, Electrocatalysis on Non-Metallic Surfaces

,

Proceedings of Workshop held at NBS, Gaithersburg , Md. , December 9-12, 1975. (Issued November 1976)

Titanates As Oxygen Electrodes

U. Bertocci, M. I. Cohen, J. L. Mullen and T. Negas

Institute for Materials Research

National Bureau of Standards

Washington, D.C. 20234

Single crystal or polycrystalline compacts of five, Ti02-based,

oxide materials were investigated as oxygen electrodes. Potentiostatic

and potentiodynamic measurements were accomplished at room temperature

using neutral borate or 5M H3P0^ and inert (Ar) or oxygen environments.

Observations of electrochemical behavior during illumination also were

made.

Materials display slow, but continuous, changes in features of

potentiodynamic scans, as well as changes in current with time at constant potential. Failure to attain steady-state conditions is

attributed to specimen mechanical/chemical inhomogeneities and changes

prior to and during measurement, variable surface conditions, together

with electrode pre- treatment. For example, cathodic currents for

O2 reduction tend to increase after cathodic polarization, while H2-

evolution is hindered by dissolved O2. Illumination increases the

current density for O2 reduction at very negative potentials while

current decreases are apparent during H2-evolution in Ar.

Results indicate that these oxides are not promising as electrocatalysts for 02-reduction, although for a definitive assessment, welldefined and closely controlled surface conditions are necessary.

Key words: Electrocatalyst ; hydrogen evolution; lanthanum titanate;

oxygen electrode; photo-assisted electrolysis; rutile.

1. Introduction

Transition metal oxides have shown interesting catalytic properties. Some of these

materials, particularly those with perovskite structure, have been found to be effective

catalysts in gas phase reactions [1]-'-, but they have been little studied as electrodes [2]. It was felt, therefore, that it would be interesting to examine these materials as electrocatalysts, particularly as oxygen electrodes in electrolytes, such as concentrated phosphoric acid, which are proposed for the operation of low temperature fuel cells [3]

.

Due to its high resistance to acid attack and to the fact that its conductivity is

controllable by doping or reduction [4], Ti02 was chosen for these investigations.

Reports of electrochemical behavior are few, primarily being limited to photoassisted

decomposition of water by Ti02 [5-10] , although some information as to the general electrochemical behavior of rutile is available [11]

.

In the lanthanum titanate series, it appears that the conductivity can be varied from

metallic to semiconducting to insulator, by varying the lanthanum oxide/titanium oxide

ratio [12]

.

Figures in brackets indicate the literature references at the end of this paper.

313
In this paper we will report on the behavior, as oxygen electrodes, particularly on

the cathodic side, of five materials: reduced Ti02, Nb and Ta doped Ti02, Lao 75'^^'^3 ^^"^

2. Experimental

2.1. Preparation

Reduced rutile was prepared by heating a Ti02 single crystal [13] at 1200 °C in a atmosphere. After reduction, the material had 6 x 10^-^ charge-carriers/cm-^ at room temperature. The Nb doped Ti02 (0.1%Nb) was also a single crystal [13] . Ta doped Ti02 (95:5) was prepared by heating at high temperature (1600 °C) a mixture of Ti02 and Ta205. The material

was further homogenized by grinding the resultant mass and pressing the power into pellets

?

followed by further sintering at the same temperature. The two lanthanum titanates were

made by melting in gettered argon a mixture of the starting materials (Ti02, Ti and

La (OH) 2) , followed by extensive annealing of the compound in gettered argon. Materials prepared in this manner were polycrystalline

.

The samples were mounted in epoxy resin after indium soldering to a copper block. The

part exposed to the solution, having surface areas ranging from 0.3 to 0.5 cm^ , was ground

flat and then polished. 1 ym diamond powder was the final polishing stage. Electrical

contact was made in a conventional way by a screw into the copper block [14] . The electrode

surface was, in general, smooth although there were some small holes and cracks.

2.2 Electrochemical Measurements

Investigations were carried out at room temperature (23 °C) in a buffer solution made

of a 1:1 by volume mixture of saturated H2BO3 and Na2B^0y (pH ^^7.5), as well as in 5M H^PO^

.

A saturated calomel electrode was used as reference. The reversible hydrogen electrode

potential was -685 mV vs. SCE in the neutral solution and -240 mV vs. SCE in phosphoric acid.

Measured values and those calculated from pH measurements agreed within 10 mV.

The measurements were taken either by slow potentiodynamic sweeps or by stepwise potentiostatic runs. Corrections for the ohmic drop in the neutral borate solution (where the

resistance was measured as 30fi*cm^)2 ^g^^g calculated and applied to the plots. In acidic

solution the ohmic resistance, measured by applying a square wave current signal to the

electrode and observing the voltage response, was less than'5Q-cm^ and its effect was negligible, since the current density was seldom more than 1 mA/cm^. Measurements were taken in a closed cell, either under an inert atmosphere (Ar) or with oxygen. Both gas bubbling and

stirring with a magnetic stirrer were used for O2 elimination if an inert atmosphere was desired and to enhance oxygen transport to the electrode when in a 0^ atmosphere.

2.3. Effect of Light

Observations have been made also on the effect of illumination on the electrochemical

behavior of the materials. The light source was a 200 W quartz-iodine lamp, whose beam was focussed on the electrode surface by means of mirrors. Since the beam, however, entered

the cell through a Pyrex glass window, most of the energy in the short wavelength region

did not reach the electrode.

3. Experimental Results

3.1. Doped TiO^

The results of the stepwise potentiostatic runs, on the cathodic side, for the three

doped Ti02 electrodes are given in figures 1, 2 and 3. The data, all referred to the reversible hydrogen electrode (RHE) , cover both neutral and acidic electrolyte in the presence

and in the absence of oxygen. The primary non-steady-state feature observed was that polarization of the electrode to a certain potential tended to increase, at least temporarily,

the cathodic current density which could be obtained at more positive potentials. The effect

2 The unit Q-cm^ is the ratio of voltage to current density.

314
Fig. 1. Cathodic current density-potential curves for reduced rutile in

acidic and neutral solution.

Fig. 2. Cathodic current density-potential curves for Ta-doped Ti02 in

acidic and neutral solution.

315
600

400

200

i

a: 0 >

E -200

I

"5c -400

«

o

-6001-

TiOg: 0.1% Nb Zi'C

10"

Atmosphere Acid Neutral

Active A

02

S.S. O

Inert A

10' 10 Cathodic current — A /cm'

10'

Fig. 3. Cathodic current density-potential curves for Nb-doped Ti02 in

acidic and neutral solution.

1.0

100 _ 1 1 1 I*HVH2 ' —Og/HgO

0

1

1 r

i -100 TiOg (Reduced) ~ w

E

u<

-200

1 \l

I r

I

|:| sat. H^BOj-NOgB^O.^

scon rate 4mV/s ~

1 / /

! O2 atmosphere c -300 Inert atmosphere —

01

l_i- Stationary values (in Og)

3

O

-400

-500 - 1 ' 1 1 1 r

-0.5 0 0.5

Potential - Volts vs S.C.E.

1.0

Fig. 4. Potentiodynamic scan for reduced rutile in 1:1 saturated

H3B03-Na2B[^07 solution. Scan rate 4 mV/s.

316
was particularly pronounced if the pre-treatment consisted of polarizing to a negative value

such as to obtain vigorous (50-100 mA/cm^) hydrogen evolution. Accordingly, as shown in the

figures, the current density was in general larger during the "return" runs, when the potential was increased (in a positive direction) stepwise, than when the potential was decreased

stepwise. A number of points are marked "active" and "steady state". The "active" values

were obtained after strong cathodic polarization, and tended to decay towards the "steady

state" value in times of the order of several minutes. It should be noticed that activation

was particularly effective in increasing the rate of oxygen reduction.

The potentiodynamic scans in an O2 atmosphere have shown the presence of a cathodic

peak located about 200 mV more positive than the potential for steady state O2 reduction.

This peak, which is shown in figure 4, decreases by repeated cycling and reappears only

after long times (of the order of one hour) at open circuit. However, larger cathodic

currents in the same potential region (approximately +300 mV vs. RHE) are so persistent

that they appear also in the potentiostatic stepwise runs, as shown in figure 5.

1 1 1 1 1

Ti02 f Ta doped

1000 1:1 sat. HjBOj - sot. Na^B^Qj

/

/ / pH • 7.5 / / O2

UJ / 500

0 Dark / cn

in>

0

Light /

/

; >

E

UJ -500

-1000

1 .1 1 1 1 -100 -80 -60 -40 -20 0 20

i (/iA/cm^)

Fig, 5. Current density-potential curves for Ta-doped Ti02 in 1:1 saturated

H3B03-Na2Bit07 solution. H2 and O2 refer to equilibrium potentials

for the two reactions.

The effect of illumination on the electrochemical response of doped Ti02 electrodes

was similar for all three. Figure 5 shows the effect for Ta doped Ti02 and figure 6 shows

the light effect in more detail for the Nb doped Ti02.

Microscopic observations of the electrode surface before and after the measurements

gave no indication that these materials were being attacked.

317
15

10

-10

-15 -

CME

5 / o< 1 4. 0 1

-Dark

-Light

TiOg 01% Nb

l:i sat. HjBO^- NOgB^O^

O2 atmosphere

-0.4 -0.2 0 0.2 0.4 0.6 0.8

Potential - V vs S.CE.

1.0

Fig. 6. Effect of illumination on currentpotential curves for Nb-doped Ti02

in 1:1 saturated H3B03-Na2Bit07

solution.

3.2. Lanthanum Titanates

The current-potential curves for cathodic reduction on the more metallic of the two titanates investigated, LaQ_g2Ti03, are given in figure 7 and those for Lag y^TiO^ in

figure 8. The effects of "activation" of the electrode surface are similar to those observed

on doped TiO, As shown in figure 8, hydrogen evolution in acidic solution was found to

10® 16^ lo** 10^ id

Cathodic current — A/cm''

Fig. 7. Cathodic current density-potential curves for LaQ,92Ti03,

in acidic and neutral solution.

318
10^ 10^ 10^ 10^ lO'^ Cathodic current — A/cm^

Fig. 8. Cathodic current density-potential curves for Lag 751103,

in acidic and neutral solution.

occur on Lag -^gTiO^ much more easily than for any other condition, including the same electrode in neutral borate. Therefore, the occurrence of oxygen reduction on this electrode in

phosphoric acid could not be established, since O2 bubbling in the cell did not make any

appreciable difference in the current density at constant potential.

The reason for this conspicuous lowering of the H2 overvoltage in H^PO^ might be etching of the surface, which was observed on both titanates in acid. Figures 9 and 10 show the

surface of the two electrodes after the measurements. The overvoltage decreases, therefore,

might be due to production of a more active surface as well as an increase in real surface

area due to etching.

Fig. 9. Photomicrograph of a LaQ^gaTiOs elec- Fig. 10. Photomicrograph of a 130.75X103 electrode surface after measurements in trode surface after measurements in

E^FOl^. Nomarski interference contrast. HsPOi^. Nomarski interference contrast.

319
The attack seems to involve anodic dissolution: in figure 11 are some anodic currentpotential curves, showing that in H3PO4 a non-negligible anodic current, of the order of

several microamperes per square centimeter is recorded at potentials well below that necessary for oxygen evolution in the dark.

4. Discussion

The results presented here show that none of the materials tested in this work seem to be promising as electrodes for the electrochemical reduction of oxygen. The best performance

was observed on Nb doped Ti02, but it was only a current density of the order of 10"^ A/cm^

for potentials less than 200 mV positive with respect to the RHE . It is true that in a practical fuel cell the higher operating temperature should decrease the kinetic hindrance,

but during the course of this work it was observed that in acid electrolyte the electrocatalytic properties with respect to oxygen reduction were, if anything, worse than in

neutral solution.

The problem of obtaining reproducible current-potential curves even in the absence of

detectable surface attack, might be linked to inhomogeneities such as compositional changes

(inferred, for instance, from variations in points to point light absorption properties) or voids and cracks, and point to the desirability to work with good single crystals of known

orientation. The beneficial effect of cathodically polarizing the electrode for the subsequent oxygen reduction seems to indicate that on these materials, as is the case with Pt and

Ir, oxygen adsorption is a hindrance for the oxygen reduction reaction [15,16]. The slow

changes in current with time at constant potential, as well as the slow changes in the

features of the potentiodynamic scans, show that any adsorption and desorption processes,

if present, have long relaxation times.

\

The difficulty in reaching steady-state conditions, prevents an attempt to investigate

the kinetics in greater detail. The slope of the semilogarithmic plots seem to change in a continuous fashion, or, at least, are broken into several segments. In general, the slopes

are fairly high (200 to 300 mV/decade of current density) in all conditions, and this holds

for O2 reduction as well as H2 evolution. This behavior, not dissimilar to that observed

320
for ©2 reduction on noble metals, might be due to a potential dependent adsorption

isotherm [17] . The exchange current densities obtained by extrapolation have to be considered

I only of orientative value: for H2 evolution i^ should be less than 10~^ A/cm^ and for O2

reduction less than 10~10 A/cm2 in the most favorable cases.

The effect of light in producing an anodic current on doped Ti02, caused, presumably,

by photoassisted oxygen evolution has been reported by several authors [5-10] . As a consequence of this fact, there is a potential range (see fig. 6) where illumination shifts

the current from anodic to cathodic. However, it is interesting that light increases,

however slightly, the cathodic current density for oxygen reduction at very negative potentials, as indicated in figure 5. In inert atmospheres, during H2 evolution, light illuminaj tion, on the contrary causes a transient corresponding to a current decrease.

! 5 . Acknowledgements

We should lilce to thank the Environmental Protection Agency for its partial support of

this program; also A. D. Franklin, W. S. Horton and H. P. R. Frederikse, who provided much

useful discussion. In particular, we would like to thank N. L. Loeffler of the University

of Tennessee Space Institute who provided us with La^TiO^ samples and with much useful

information.

6. References

[1] R. J. H. Voorhoeve, J. P. Remeika, P. E. Freeland and B. T. Matthias, Science 177

353 (1972)

.

[2] P. J. Boddy, D. Kahng, Y. S. Chen, Electrochem. Acta 13_, 1311 (1968).

[3] A. E. Gillis, Advanced Development. In: Seventh Status Report on Fuel Cells, J. R, Huff (ed.) U. S. Army Mobility Equipment Research and Development Center, Ft. Belvoir,

Va. Report 2039 NTIS No. AD-755 (1972).

[4] H. P. R. Frederikse, JAP Supplement, 32^, 2211 (1961) . [5] A. Fujishima, K. Honda, S. Kikuchi, Kogyo Kagaku Zasshi, 72, 108 (1969).

[6] A. Fujishima, K. Honda, Nature, 238, 37 (1972) . [7] J. G. Mavroides, J. A. Kafalas, Lincoln Lab., MIT, Solid State Research, 2 (1974).

[8] M. S. Wrighton, D. S. Ginley, P. T. Wolezanski, A. B. Ellis, D. L. Morse, A. Linz,

Proc. Nat. Acad. Sci. USA, 72_, 739 (1975).

[9] K. L. Hardee, A. J. Bard, J. Electrochem. Soc , 122 , 739 (1975).

[10] A. Fujishima, K. Kohayakawa, J. Electrochem. Soc. 122 , 1487 (1975) . [11] J. W. Higgle, Oxides and Oxide Films, M. Dekker, Inc., N. Y. Vol. I (1972) and Vol. II

(1973)

.

[12] N. L. Loeffler, University of Tennessee Space Institute, Private Communication.

[13] Crystals Grown by Flame Fusion Technique. Dopants Added to Starting Materials.

[14] M. Stern, A. C. Makrides, j. Electrochem. Soc. 107 , 782 (1960).

[15] J. P. Hoare in "Encyclopedia of the Electrochemistry of the Elements", A. J. Bard,

Editor, Chapter 4, p. 307. M. Dekker, N. Y. (1974).

[16] J. P. Hoare, J. Electrochemical Chem. 18_, 251 (1968).

[17] A. J. Appleby, J. Electrochemical Chem. 27, 335, 364 (1970).

321

ELECTROCHEMISTRY AT SOLID ELECTROLYTE INTERFACES

National Bureau of Standards Special Publication 455, Electrocatalysis on Non-Metallic Surfaces

,

Proceedings of Workshop held at NBS, Gaithersburg, Md. , December 9-12, 1975. (Issued November 1976)

Ionic Transport and Electronic Exchange at Solid Electrolyte Interfaces

M. W. Breiter and G. C. Farrington

General Electric Research and Development Center

P. 0. Box 8 Schenectady, New York 12301

Interest in highly conductive solid electrolytes,

compounds that display conductivities comparable to

aqueous electrolytes at room temperature, is great because

of their many potential applications. This paper reviews

conduction mechanisms and charge transfer at solid

electrolytes having fast ionic transport. For the

discussion, solid electrolytes have been grouped into

stoichiometric compounds, doped compounds and non- stoichiometric compounds; interfaces are classified

as two-phase and multi -phase boundaries.

Key words: Charge transfer processes, conduction

mechanisms, interfacial phenomena, solid electrolyte

1 . Introduction

Solid compounds with greater ionic cponductivities than electronic

conductivities have long been known. [1] Usually, ionic mobility in such

compounds is appreciable only at elevated temperatures. In recent years, new

materials which in certain temperature ranges display ionic conductivities

comparable to aqueous electrolytes at room temperature have been found.

Interest in these solid electrolytes is great because of their potential

application in fuel cells and batteries for the production and storage of

electrical energy, in electroly zers for the production of certain gases, and

in sensors for monitoring gases or dissolved species in liquids.

Conduction mechanisms in highly conducting solid electrolytes have been

extensively reviewed [2-5] , but less information is available [4] on electrochemical processes occurring at interfaces where one boundary is established

by the solid electrolyte. A previous review of this area was written by

Raleigh [4] . The emphasis of the present paper is on conduction mechanisms

and charge transfer processes at interfaces with solid electrolytes of fast

ionic transport. •

2. Characterization and Conduction Mechanisms of Solid Electrolytes

Solid electrolytes with fast ionic transport fall in three broad

categories

Figures in brackets indicate the literature references at the end of

this paper.

323
1. stoichiometric compounds, for instance, Agl

2. Doped structures, like calcia stabilized zirconia.

3. Non-stoichiometric compounds, such as beta alumina.

Ionic conductivities of examples for each class are plotted versus reciprocal

temperature in figure 1. For comparison, the conductivity of a "regular"

ionic conductor, LiF, is also shown.

1000 500 300 200 100 0 "C

Fig. 1. Plots of the product of

conductivity and absolute

temperature, T, versus I temperature for different 1.2 2.0 2.8 3.6 solid electrolytes. 1000 /T CK)

The shape and region of high conductivity of each curve in figure 1 reflect its underlying mechanism of conduction. The curve for LiF, ror example, consists of two parts. Ionic conductivity is appreciable only

within several hundred degrees of the melting point and results from thermal

generation of current-carrying defects in the crystal lattice. At temperatures below the point where the concentration of thermally generated defects

is comparable to the impurity concentration, conduction is dominated by fixed

impurities

.

The curves for ionic conductors of class 1 consist of two parts separated

by a sharp transition region in which a phase change occurs. Relatively large

conductivity is observed above the transition region and normal, thermally

induced conductivity in the lower temperature phase. In the high conductivity

phase, a large number of interstitial sites is provided for the small cations

in the rigid anion structure. The cations may be considered to be in a molten

state, migrating through the fixed anion sublattice in a disordered fashion.

In contrast, the phase below the transition temperature is cation-ordered.

Compounds of class 2 are mixtures of oxides and display large oxygen ion

mobilities at high temperatures. Doping of a suitable oxide with another

oxide in which the cation has a lower valency leads to formation of a cubic

phase with a large anion vacancy concentration. The extent of doping is

considerable. Calcia stabilized zirconia is a classic example for group 2.

The conductivity of a class 2 conductor is largest at one particular doping

level, for instance, at about 12 mole % CaO for calcia stabilized zirconia.

Usually this doping level is not identical with the maximum miscibility of

the two oxides. A detailed understanding of conductivity in compounds of

class 2 is complicated by problems of metastability and partial ordering. [6]

Sodiiim beta alumina, formulated as (1+x) Na20 • 9AI2O3 , is the best known

example of a class 3 compound. The normal beta-phase contains 15-30 percent

excess sodium ions which are compensated either by aluminum vacancies or oxygen inters titials . All the sodium ions are located in planes perpendicular

to the c-axis which are separated by four oxygen-aluminum planes. Rapid ion

migration occurring in these planes is two dimensional, in contrast to three

dimensional conductivity in compounds of class 1. As shown in figure 2,

there are three sites, d, b, and MO available for Na"*" ions in the conducting

324
Fig. 2. Ionic arrangement in

conduction plane of

sodium beta aliimina. OOXYGEN

(CONDUCTION PLANE)

OXYGEN

(CLOSE PACKED) O

SODIUM

plane. The distribution of the Na"^ ions among the sites changes with their

concentration and temperature, becoming more uniform with increasing

temperature.

Sodium beta" alumina differs somewhat in composition from sodium beta

alumina and contains a Na"*" ion excess of about 50 percent. These ions reside

in nearly equivalent sites. The beta" phase has higher conductivity, 0.3 cm

-1 vs . 1.2 fi~lcm~l at 300°C [7] , in comparison to beta alumina.

The preceding discussions of conduction mechanism are strictly valid only

for ideally pure single crystals. Additional effects, especially grain

boundary resistance, must be considered for polycrystalline samples of classes

2 and 3. Grain boundary resistivity is not significant for class 1 compounds

in which appreciable conduction is only found at temperatures one to several

hundred degrees below their melting points where single crystal and polycrystalline specimens display identical behavior. In fact, grain boundary

resistivity may actually be smaller [4] than bulk resistivity for some class

1 compounds

.

Grain boundary impedance has been measured for stabilized zirconia [8,9] and sodium beta alumina. [10,11] The frequency dependence of the impedance

was analyzed to separate bulk resistivity, grain boundary resistivity and

electrode polarization. Figure 3 shows results obtained with (Sc20^ 0 . 1 (2r02) o .

9

The ratio between grain boundary resistivity) Rj-,, and* between 550° and 571°C.

Fig. 3. Ratio of grain boundary resittivity to crystal resistivity

as function of the average grain

size for two different samples

of stablized zirconia.

325
that of the crystal interior, R^, is plotted as a function of average grain

size for two specimens sintered at different temperatures. The relative

contribution of grain boundary resistance increases as the grain size decreases

and the number of grain boundaries becomes larger. The ratio R^/Rc still

< 1 at the smallest grain size for (30203)0 i(^^^2^0.9' demonstrated [10]

for Na"*" beta aliimina specimens containing small amounts of additives, Rj^/Rc

also depends upon temperature because the two conduction processes have

different activation energies. More recent results for beta alumina will be

discussed later. Other factors affecting solid electrolyte impedance have

been discussed in reference 12.

3. Characterization of Interfaces

Interfaces between a solid electrolyte and other compounds may be

separated into two classes:

1. Two-phase boundaries

2. Multi -phase boundaries

Class 1 represents the simplest case of direct contact between a solid

electrolyte and another compounds Two-phase boiindaries may be further

subdivided:

- Interfaces with blocking electrodes

- Interfaces with non-blocking electrodes

a. Electrode reversible to conducting ion

b. Electrode not reversible to conducting ion

- Interfaces between solid electrolytes and other electrolytes.

Multi-phase interfaces, at which at least two additional phases contact

a solid electrolyte within the same area, are more difficult to classify.

The simplest case is a three-phase boundary between a solid electrolyte, an electronic conductor, and a third phase which is either a liquid or gas.

Although individual interface behavior is discussed, experimental

interfacial studies require at least two interfaces. To isolate the behavior

of one particular interface, the cell must be constructed to eliminate effects

of the secondary interface on an experimental or theoretical basis. As with

aqueous systems, the simplest approach is to minimize the contributions of

the secondary interface compared to those of the interface under study.

Practical preparation of interfaces of type 1 and 2 involves numerous

considerations. The reader is referred to the excellent discussion of this

subject by Raleigh in chapter IIIA of reference 4.

If inert electronic conductors are employed as contacts, the potential

range of study may be chosen so that faradaic reactions do not occur. The

electronic conductor then behaves as an ideally polarizable electrode. Many

examples can be found in the literature in which Pt or Au/ionic crystal

interfaces are used in studying ionic mobility and intrinsic polarization

effects within crystals. In many, the electrode contacts are only assumed to

act as blocking electrodes even though this point is crucial to successful

interpretation of results.

But, blocking behavior cannot be assumed without proof. Charge transfer

reactions may no more be absent at an interface such as Pt/beta alumina

merely because of the unlikely role of Pt as parent of ions injectible into

beta alumina than at a Pt/non-aqueous electrolyte or Pt/aqueous electrolyte

interface at which Pt is an equally unlikely source of ions. A Pt/electrolyte

interface can, of course, approximate an ideally polarized electrode; but,

regardless of the electrode composition or physical state, that behavior must

be established experimentally, taking into consideration such factors as the
decomposition potential of the electrode and the electrolyte, and the presence

of tertiary substances capable of redox behavior at the electrode. Only then

can the electrical response of the electrolyte be interpreted on the basis of

models assuming blocking behavior.

In theory, a Pt/beta aliimina interface or its equivalent should approximate an ideally polarizable electrode over a wide potential range in the

absence of species capable of bringing about charge transfer such as adsorbed H2O or H2 . Beyond this, the range of ideal polarizability is

limited by decomposition of the beta alumina, which involves Na"*" reduction

at one potential extreme and 0 oxidation at the other. This represents a range of ideal behavior much wider than typical liquid electrolyte systems.

To date, no detailed experimental exploration of blocking electrode/beta

alvunina interfaces has appeared.

General models for the AC behavior of solid electrolyte interfaces have

been developed for different conditions by Macdonald. [13,14] Simple models

for interfaces with blocking electrodes without specific ion adsorption, for

interfaces with blocking electrodes with specific ion adsorption, and for

interfaces with non-blocking electrodes were proposed by Armstrong [15]

.

The models refer to solid electrolytes of class 1 or class 3 with high cation

conductivity. The solids are assumed to consist of cations which are all

equivalent and all mobile and of immobile anions. Because of their large

mobile ion concentrations, diffuse double layer contributions at such solid

electrolyte interfaces are considered negligible. This is a generalization

of a concept proposed previously by Hull and Pilla [16] for the graphite/

silver iodide system. Only the concentration of cations in the first layer

differs from that of the bulk, resulting in a compact double layer. The

situation is similar to that in concentrated aqueous solutions where the

diffuse double layer is negligible in comparison to the Helmholtz layer.

Analogue circuits [15] describing some of the main features of solid

electrolyte/electrode interfaces are shown in figure 4a and 4b. Both

circuits consist of two separate circuits in series, reflecting the interfacial impedance and the impedance of the solid electrolyte between the

interface and measurement probe. If the frequency is below ca. 100 kHz, the

capacitance, CeI, of the solid electrolyte is negligible. The circuit for

an interface with blocking electrodes (fig. 4a), then, simplifies to a series

circuit of double layer capacity C,^ and electrolyte resistance R^i if

specific adsorption of ions, represented by C^^ and Ra^, and the Warburg

impedance, W, are absent. According to Armstrong [15] , the Warburg impedance

only exists at blocking interfaces if two mobile cations are present, one in excess of the other. In the circuit for non-blocking electrodes (fig. 4b)

,

a

Fig. 4. Analogue circuits for the interface

electrode/solid electrolyte:

a. for blocking electrode

b. for non-blocking electrode

327
Ret represents the charge transfer resistance. It contains no Warburg

component if only the reaction M^"*" + e~ with one cation occurs. The possibility of an adsorbed intermediate leads [15] to introduction of Rg^^j and C^jj

into figure 4b. The applicability of these interfacial models and their

respective analogue circuits will be discussed on the basis of experimental

results for different interfaces in subsequent chapters.

4. Electrochemical Processes at Interfaces with Stoichiometric

Compounds (Class 1)

Interface studies with stoichiometric compounds may be grouped into two

broad categories

- electronic conductor/solid electrolyte

- aqueous electrolyte/solid electrolyte

Investigations of the second category are less common because of the problem

of solid electrolyte dissolution in water.

4.1. Studies of Electronic Conductor/Solid Electrolyte Interfaces

The transient behavior of graphite/AgI and Pt/AgI interfaces was investigated by Hull and Pilla [16] in cells of the type Ag/Agl/M at room temperature. Cells were assembled by pressing the powdered constituents together

with 40,000 psi into pellets. A small silver wire was embedded in each solid

electrolyte as a reference electrode. Ohmic potential drop inside the solid

electrolyte was electronically compensated.

Interfacial behavior during a voltage sweep of 5 mV/sec, starting at

0.46 V vs. Ag/AgI , is represented in figure 5. No appreciable Faradaic

current is observed to flow across the graphite/AgI interface until a potential

of about -0.03 V is reached during the cathodic sweep. Silver deposition

commences abruptly at this point. After sweep reversal, silver deposition

occurs at lower potentials than during the preceding cathodic sweep. This

implies that silver is more easily plated on silver islands than on free

graphite sites. Removal of the silver deposit occurs in the broad anodic

peak between 0 and 0.2V. Contrasting behavior occurs at a Pt/AgI interface.

Silver is continuously plated out during the cathodic sweep, starting with the

formation of a monolayer at underpotentials as in aqueous solutions.

I/S(/iA) + 200

328
Iodine is produced at about 0.5V for graphite and 0.46V for platinimi

during the anodic sweep. Graphite and platinum display similar electrocatalytic behavior as in aqueous solutions

.

Both voltage sweep and galvanostatic studies have demonstrated that

graphite/AgI approximates an ideally polarized electrode between 0.5V and

0.35V. Such a potential range does not exist for platinum. This behavior is

very similar to that of graphite/silver bromide and platinum/silver bromide

interfaces at higher temperature (244° - 292°C) . A more detailed study of

halogen discharge on graphite electrodes from three different halide solid

electrolytes was recently reported by Raleigh. [17]

Detailed studies of the kinetics of Faradaic interfacial processes have

been carried out for Ag/Ag-jSI [18,20] and Ag/Ag4Rbl5 [21,25], Anodic currentvoltage curves [18] of Ag m the cell Ag/Ag3SI/ (C) I2 at 25°C and 45°C are shown in figure 6. Although these curves and similar cathodic currentpotential curves in reference 19 have shapes usually found when diffusion is

rate-determining, analysis under this assiimption leads [18] to incorrect

values of the diffusion layer thickness. The interpretation that homogeneous

reaction between interstitial silver ions, ion vacancies, and ions in regular

lattice positions is rate-determining yields a theoretical expression [18]

which adequately describes the experimental results over about a decade of

current density. But the sharp increase of potential with current density in

figure 6 must be explained [18] by an additional process, namely precipitation

of a substance of lower conductivity, such as Agl , at higher currents. Since

linear plots of current density versus potential are not given in reference

18, it is difficult to decide if, as suggested [18], the shajrp current increase

represents transition to a limiting current.

Interfacial polarization is eliminated [18] by amalgamation of the silver

electrode. It was proposed [22] that this implies a heterogeneous rather than

homogeneous reaction is rate-determining. Further results [19,20] are free of

this objection, but lead to similar conclusions as reference 18. Other interpretations for the behavior of amalgamated silver electrodes must be found.

A linear plot of overvoltage versus current density revealed the

existence of a limiting current during cathodic deposition of silver onto

silver in the cell Ag/Ag^SI/Ag. The interpretation [17] that the ratedetermining step is transition of silver ions from regular lattice positions

to interstitial positions leads to an adequate theoretical description of the

observed phenomena.

Impedance measurements [20] of the interface Ag/Ag3SI between 100 Hz and

100,000 Hz in a three-electrode cell without polarization and at -0.046 V led

329
to a Warburg-type frequency dependence of the Faradaic impedance components,

Rg and l/^jjCg, in a series circuit after correction for ohmic potential drop

inside the solid electrolyte and for double layer capacity. The latter value was determined experimentally according to the procedure in reference 26. Rg vs. a)"-'-/^ goes through zero and l/coCg vs. a)~l/2 intersects the ordinate as a positive value corresponding to about 100 yF/cm^. This is attributed to an adsorption capacity. It was proposed [20] that the rate-determining transition

[19] of an ion from a regular lattice position to an interstitial position is so slow that it does not affect the impedance in the given frequency range.

Two conclusions may be drawn from the results with Ag^SI:

a. The assumptions made [15] in modeling interfaces between

solid electrolytes and non-blocking electrodes are inadequate in this case.

b. The charge transfer step is rapid.

The first conclusion is based on results from DC and AC studies, the second

conclusion on AC measurements.

The anodic behavior of the interface, Ag/Ag4Rbl5, was studied [21,22] in

the cell, Ag/Ag^Rbl^/Ag, with a silver reference electrode. Figure 7 shows

current-potential curves at 0°C and 25''C. The IR drop is eliminated. At

small current densities, a Tafel line with a slope of 20 mV is found. A

second Tafel line with a slope of about 110 mV appears to occur [21,22] at

larger currents at 25°C. However, the strong temperature dependence of this

second region, the gradual increase [21] of overvoltage with time, and the

curvature of the Tafel plot [22] raise serious doubts to the conclusions [22]

that either ion-transfer or surface diffusion of adatoms is rate-determining

there. Impedance measurements [23,25] of the Ag/Ag4Rbl5 interface also lead

to ambiguous conclusions

.

Since the analogue circuits contain more than two elements, different

circuits [23,25] allow the same degree of approximation to the experimental

results. In general, the large number of parameters determined from experimental data in the description of an electrode/solid electrolyte interface by

analogue circuits poses serious difficulties for reliable interpretation.

4.2 Studies of Solid Electrolyte/Aqueous Electrolyte Interfaces

Solid electrolyte/liquid electrolyte interfaces in the system Ag3SI/

AgN03+KISI02 have been studied at room temperature by DC [27,28] and AC measurements [291. The contributions of current lead contacts to Ag3SI were minimized

330
by a special cell construction [27] . The potential drop across the interface,

measured against a calomel electrode, varied [27] with silver ion activity in

a Nernstian relationship. Cathodic polarization curves of silver deposition

displayed limiting current behavior identical [28] for rotating discs of Ag3SI

and of Ag. Mass transport of ions in the liquid solution is rate-determining.

If the electronic conductivity of Ag-^SI is sufficiently small, silver ions

migrate into interstitial lattice positions. If not, silver deposition occurs

on the solid electrolyte surface in a parallel reaction.

In contrast to the cathodic process, transfer [28] of silver ions from

the solid electrolyte into the liquid electrolyte requires a larger overvoltage

than the comparable anodic process on silver metal (see fig. 8) . The currentpotential curves do not depend upon rotation speed and silver ion concentration in the liquid electrolyte. Above 4 mA/cm^ , the anodic current-potential

curves become Tafel lines, n = a + b logi, with b equal to 0.3 8 V for Ag3SI

and 0.28 V for Ag. It was suggested [27] that transfer of ions from interstitial positions into the liquid electrolyte is not rate-determining but

that supply of silver ions to the interface in Ag^SI is slow.

Fig. 8. Dependence of rate of silver ion

transfer from Ag3SI (curve 1) and Ag

(curves 2 to 4) into liquid electrolyte upon electrode potential at

rotating disc electrode: • 0.01 M

AgN03 + 0.09 M KNO3 , 0 ro tations/min;

9 0.05 M AgN03 + 0.09 M KNO3, 60

rotations/min; o 0.05 M AgN03 + 0.09

M KNO3, 100 rotations/min.

Further experiments examined the capacity of the interface, Ag4SI/AgN03+

KNO3, and found it much smaller than that of the interface, Ag/AgN03 + KNO3,

at 100 kHz. It was concluded []9] that the double layer capacity is determined

by the double layer inside the solid electrolyte. Cathodic polarization

decreased the double layer capacity of an Ag^SI/AgNO^ + KNO, interface and

anodic polarization increased it. These results were attributed to a change

of the space charge inside the solid electrolyte. The effective thickness of

the space charge layer becomes smaller by the introduction of more interstitial silver ions under cathodic polarization. The reverse is true for

anodic polarization.

Studies [27,29] of solid electrolyte/liquid electrolyte interfaces are not in agreement with the predictions of the models. [15]

5. Electrochemical Processes at Interfaces with Doped Compounds

Although many investigations of two-phase boundaries have been carried

out with electrolytes of classes 1 and 3, few have appeared for doped

compounds of class 2. The primary reason is that interest in multi-phase

boundaries has been greater with doped compounds because of their potential

use as oxide ion conductors in such devices as high temperature fuel cells.

The simplest use of multi-phase interfaces with doped conductors is in

331
potentiometric study of oxide ion activity. For example, the cell, Pt/Fe,

FeO/ (Zr02) 0 . 85 (^^'-'^ 0 . is/'^^ '^^^/•^'*^' been used to determine the free energy

of formation of NiO. [30] Similar cells have been used as current sources in

coulometric titrations [31,32] and in oxygen transfer studies. [33,34] The

mixture of metal and oxide serves both as an electronically conducting contact

and as an oxide ion reservoir. The interface is of the multi -phase type. In

a coulometric titration, oxide ion depletion at the negative interface is

compensated by metal ion reduction.

The three-phase interface which has received major attention in fuel cell

research is the combination of an oxide ion conductor, electrocatalyst , and

molecular oxygen. As with analogous three phase boundaries in aqueous electro

lytes, the electrocatalyst strongly influences interfacial behavior.

Current-voltage curves obtained [36] for a Pt/ ( Zr02) q . 9 (CaO) q j^/Pt cell

at 1000°C are replotted in figure 9. The electrodes were adherent^ porous,

fine-grained platinum deposits on both sides of closed, stabilized zirconia

tubes. The inner electrode was exposed only to pure oxygen gas and served as both counter and reference electrode. The composition of the gas mixture at

the outer platinum electrode was varied. Current-voltage curves remained

linear up to 250 mA/cm2 at 1000° and the DC resistance taken from the slope

of the i-U curve was equal to the AC resistance, determined separately wh.en

both electrodes were exposed to pure oxygen. Polarization effects were absent under these conditions implying that the inner electrode could be used

as a reference electrode. The influence of the oxygen partial pressure on oxygen evolution is negligible up to 0.5 8% O2 . In contrast, oxygen reduction

is accompanied by considerable polarization m argon-oxygen mixtures dilute

in oxygen. The shapes of the cathodic i-U curves in figure 9 indicate that a limiting current which depends upon the oxygen content is reached. A more detailed study reveals 137,38] that the appearance of the limiting current

becomes less pronounced and desappears when the required potential approaches

and_ surpasses 2V, e.g., at relatively high ohmic resistance of the electrolyte

At currents above the limiting current, an additional process of direct

exchange between the electrode and the electrolyte occurs, [37,38] leading to

electronic conductivity. The electrolyte itself is reduced.

Ippm O2

100 TO 0.58% 0

2

-200

100 %02

i(mA/cin ) Fig. 9. Current-voltage characteristics

for Pt/(Zr02)o 9(CaO)o i/Pt at

1000 °C for different gas

mixtures

.

Since polarization effects are absent, mechanistic conclusions for the

separate steps in the process of oxygen evolution cannot be drawn at present.

Mass transport [36,39,40] in the gas phase at 1000°C, diffusion of oxygen

atoms in the adsorbed phase [37] at 810°C, and diffusion through the platinum

electrode [38,41] at 520° to 560°C have all been suggested as rate-deteirmining

steps for oxygen reduction at currents below the limiting current. These

deductions are based on interpretation of DC measurements. It was pointed

[42] out recently that a transition in the rate-determining step from oxygen

diffusion through the platiniam to mass transport in the pores, should occur

as temperature, oxygen pressure, and porosity change. This may explain the

discrepancy between various hypothesized rate-determining steps. It is clear

332
that at high potentials either oxygen transfer [38,41] into the electrolyte

or electronic conduction [38] within the electrolyte is controlling.

One kinetic step in the net reaction was postulated [8] as ratedetermining on the basis of impedance masurements on Pt/(Zr02)Q 9(^2^3^0 l-^-^^ over a wide range of oxygen pressure between 400° and 800°C. The analogue

circuit for interpreting the impedance data is that in figure 4b with the

Warburg impedance taken as zero and a third ohmic resistance added in series.

The components on the right side of the analogue circuit are proportional [8]

to the length of the solid electrolyte and thus reflect [8,43] contributions

from grain boundary and crystal impedance. The values of 1/C^ and R^^ are proportional [8] to the geometric surface area of the platinum electrode.

They result from interfacial processes. The resistance, Rcf increased by

about a factor five at 400°C when the porous platinum electrode was replaced

by a nonporous one. The latter result was attributed [8] to a diminished

number of three-phase boundaries. Direct evidence for such an effect was previously given [44] by DC measurements. In general, R^-,^ became [8,43] very

small at temperatures above 800°C. The latter result is in agreement with

the conclusion [36,39,40] from DC measurements that mass transport in the gas

phase, which is not reflected in the impedance, becomes rate-determining at

1000°C. It was suggested [8] on the basis of the dependence of R^^ upon

oxygen pressure that the electronic transfer step between adsorbed oxygen atoms

and oxygen ions or the dissociation step of molecular oxygen is rate-determining at the three-phase boundary. The latter interpretation which should

hold for temperatures below 800°C is inconsistent with the conclusion from DC

studies that diffusion of oxygen in the adsorbed phase [37] or through the

platinum electrode [38,41] is the slowest step. However, recent AC studies

[45,46] of the system indicate that the simple analogue circuit is not

adequate and suggest [45] that the formation of a platinum oxide affects the

impedance

The behavior of the interface, metal/stabilized Zr02, was also investigated [47-49] in the presence of an inert gas in which the content of

molecular oxygen had been reduced as far as experimentally feasible. Under

this condition, solid platinum [48] or liquid silver [47,49] or liquid tin

[48] approximate ideally polarizable electrodes. Since the resistance and

capacity in a series circuit did not change [50] significantly with frequency

above about 100 kHz, the capacity at 400 kHz was taken as the doiible layer

capacity. Plots of this capacity versus potential, referred to a platinum

electrode in air, had [48] parabolic shapes with sharp minima for liquid

metals at 975°C. The double layer capacity curve with platinum was less

symmetrical, [47] but also displayed a minimum. The minimum was attributed

to the point of zero charge, assuming a space charge layer inside the

stabilized zirconia. As expected, the capacity-potential curves became [49]

wider with increasing yttria content, that is, increasing defect concentration.

The described results [47-50] are not in agreement with a model for the interface in which only a compact double layer is considered.

6. Electrochemical Processes at Interfaces

with Non-Stoichiometric Compounds

These studies fall in two groups:

- M/M'*'-beta alumina

- M'''-beta alumina/liquid electrolyte

M/M'*'-beta alumina represents a broad group of beta alumina interfaces in

which M is th^ reduced form of the primary mobile ion within the beta alumina

structure. M can be Na"^, Li"*", K"*", Rb"*", Tl"^, or Ag in beta aliimina; only

Na"*" has been reported to be mobile within beta" aliimina.

Not all substituted beta aluminas are stable in contact with their

parent alkali metals. For example, molten Li metal reduces [51] Na"'"-beta

alumina at 450°C. Na/beta alumina and Na/beta" alumina are believed to be

stable at least to 400°C based on results of Na/S battery research. Ag and

Tl should also be stable in contact with beta or beta" alumina.

333
The M -beta alumina/M -liquid electrolyte interface presents an unusual

opportunity to study ionic transport between dissimilar electrolytes over a wide range of chemical conditions because of the chemical stability of beta

alumina. M"*" can be any of the ions mobile within beta alumina such as Na"*", Li"^, K"*", Tl"*", or Ag"*". Only Na''"-beta alumina/Na"*" containing electrolyte has

been discussed in detail in the literature.

6.1 Studies of M/M -Beta Alumina Interfaces

M/M -beta alumina approximates an ideally non-polarizable electrode.

Deviation from ideality can arise from space charge effects within the solid

electrolyte, which should be insignificant because of its high charge carrier

concentration, and from transition and crystallization overpotentials

.

Transition overpotential , or energy required to bring about ionic transition

from the solid electrolyte structure to the metallic electrode, should be the

primary component of overpotential at a liquid metal/beta alumina interface.

Crystallization overpotential may occur in addition to transition over- potential at a solid metal/solid electrolyte interface. Published work has

so far only described charge transfer at Na/beta alumina and Na/beta" alumina

interfaces

.

Studies [52] of interfacial transport in symmetric cells of the type Na/

beta" alumina/Na from 150-350 °C at frequencies between 10"^ and 10^ Hz

revealed no evidence of charge transfer polarization on finely polished discs.

In contrast, a marked effect of surface roughness on the impedance exists as expected from theoretical predictions [26] for aqueous systems. Wetting of

beta alumina with sodium is also strongly affected by surface preparation.

This agrees with the observation [53] that Na begins to wet monofrax beta

alumina only above 300°C.

Significant interfacial polarization was observed [54] in a freshly

prepared Na/beta alumina/Na cell with two probe conductivity measurements at

150 °C from 1 to 140 mA/cm^ using single crystal boules of beta alumina. No

polarization was detected at 150°C after the cell was heated to 300°C for one

hour. This appears to be another manifestation of Na wetting of beta alumina.

No interfacial polarization could be observed above 150°C after wetting, but

it was apparent at temperatures below 150°C, especially below 97°C, the

melting point of Na. The experiments with solid Na appear to be the first

AC examination of polarization at a solid Na/beta alumina interface. As

expected, polarization increases with decreasing temperature.

Other Na/beta alumina interfacial data are sparse. Will [55], in

examining the performance of the Hg(Na)/beta alumina/Bro battery, reported

there is insignificant polarizatioh at the liquid Hg(Na)/beta alumina interface. In contrast, polarization effects become noticeable [56] at temperatures below about 150°C in symmetrical cells of the type Hg(Na)/beta alumina/

Hg(Na). Armand [57] briefly discussed the difficulties of preparing an oxidefree solid Na/beta" aliraiina interface and found at an imperfect Na/beta"

alumina interface, held under compression, a current of approximately 0.1 mA/

cm''2 can be drawn for an average time of 30 min. before contact is lost due

to Na depletion. This corresponds to the removal of less than five monolayers

of Na.

Whittingham and Huggins [58] studied the interfacial response of solid

Ag/Monofrax beta alumina contacts prepared by painting the solid electrolyte

with Ag resinate solution and annealing for 12 hours at 800°C. No frequency

dependent impedance attributable to interfacial polarization was observed from

10^ to 105 Hz at 25°C.

Liquid Na/beta alumina is the inverse of most electrochemical interfaces

at which the electrode is solid and the electrolyte liquid. Just as meticulous

334
electrode preparation is essential in a study of solid electrode/liquid

electrolyte charge transfer kinetics, similar care must be applied to electrolyte preparation in solid electrolyte investigations. The few studies

available show that the exchange current at a carefully prepared Na/beta or

beta" alumina interface is very large above 150°C. Between 150°C and 97°C,

interfacial polarization becomes apparent, although insufficient data are available to identify its origin. Below 97°C, polarization increases steadily

as temperature decreases. Few data are available in this temperature range

where effects of electrode depletion, poor electrode/electrolyte contact, and

inhomogeneous current density should be most pronounced.

6.2 Studies of M'^-Beta Alumina/lyt-Liquid Electrolyte Interfaces

Na -beta aluinina/Na liquid electrolyte is a two-phase system with a single mobile ion common to each phase.

Kinetic considerations predict [59] that the i/U relationship between

interfacial polarization, r\ , and current density, i, is

. r azFn -(1 - a) zFnl (i\

in which a is the transfer coefficient, Iq, the equilibri\im exchange current

density, and z the charge of the mobile ion.

The i/U relationship for Na"*" entry into and exit from beta al;jmina across

a propylene carbonate interphase should be described by equation (1) in which

ion exit from the solid electrolyte is considered a positive current occurring

at a positive overpotential. This assumes that transport is hindered

primarily by transition overpotential.

Voinov and Tannenberger [60] examined interfacial transport across a beta alumina/propylene carbonate + Nal + I2 interface using a polycrystalline

beta alumina tube which dipped halfway into and was halfway filled with the

propylene carbonate solution. Four electrodes were used, one current-carrying

electrode in each electrolyte compartment and two reference electrodes. One

reference electrode was positioned halfway between the outer current carrying

electrode and the beta aliimina surface and the other was wrapped around the

outside top of the tube and wetted with electrolyte solution. This electrode

placement insured that only interfacial polarization from the outside interface was measured. It did not, however, eliminate ohmic polarization from

bulk propylene carbonate resistivity. Current through the outer electrodes

was stepped while monitoring the potential difference between the reference

electrodes

.

The results showed Na"^ exit from beta alumina into propylene carbonate

to be a slow process with a high activation energy. Extended Na+ injection

into beta alumina was accompanied by steadily increasing interfacial polarization for both ion entry and exit. This was attributed to irreversible

accumulation of H+ ions in the beta alumina and a concurrent decrease in

local conductivity. Since neither the beta alumina nor the propylene carbonate were specially dried, water was surely adsorbed at the interface and

could serve as a source of H"*".

It is reasonable that a hydrogen ionic species can be injected into beta

alumina from a water contaminated interface. The hypothesis that this is

irreversible, however, is inconsistent with the mechanism which produces it.

Any ion diffusing into the conduction plane in an electrid field should

diffuse out in a field of opposite polarity. ^^0+ is also more likely to

diffuse into the conduction plane in the presence of excess than H"^

335
because it is a larger ion and should have a larger diffusion coefficient

within beta alumina.

These minor objections do not diminish the broader conclusion of Voinc

and Tannenberger that water significantly influences beta alumina interfaci il transport. Because of water's uncertain effect, the results reported do no reflect intrinsic polarization at a beta alumina/propylene carbonate interface.

Farrington [59,61] also studied Na"*" transport across a polycrystalline

beta alumina/propylene carbonate interface. In his procedure, a polished

disc of polycrystalline beta alumina, dried at 10~^ torr and 400°C for 24

hours, was held by silicone rubber gaskets between two chambers of a Teflon

polymer electrolysis cell. The cell compartments were filled with a carefully dried solution of 0.01 M NaClO^ + 0.10 M tetrabutyl ammonium tetrafluoborate in propylene carbonate. The apparatus incorporated four electrodes,

an outer current-carrying pair of Pt foil, and an inner reference pair of Tl

(Hg) dissolved in Au wire. One reference electrode dipped directly into the

propylene carbonate near one disc face, and the other was inserted in a hole

drilled into the beta alumina halfway between each face. Interfacial over- potential at the surface between the reference electrodes was separated from

bulk propylene carbonate and beta alumina resistivity by a galvanostatic

transient technique, which took advantage of the longer time constant (10"^

sec.) of interfacial polarization compared to bulk resistivity (10~ sec).

Constant current maintained unidirectional ion flow to differentiate ion entry

into the solid electrolyte from ion exit.

2

. No interfacial polarization overpotential is apparent from 30 yA/cm to

10 yA/cm^ at 23°C for Na transport across a propylene

carbonate/beta alumina interface that has been scrupulously dried.

Interfacial polarization is apparent for Na"*" exit from polycrystalline

beta alumina previously exposed to 100 percent relative humidity as figure 10

shows. In this plot, bulk iR drop has been subtracted. Overpotential follows

a Tafel relationship for which the transition coefficient (a) is 0.24 at 23°C

and 0.22 at 62°C and the exchange current (i ) is 3.0 x 10"^ at 23°C and 2.4

X 10-5 A/cm2 at 62°C. °

Ion entry into beta alumina from propylene carbonate proceeds without

observable interfacial overpotential across a carefully dried interface at

current densities less than 10^ yA/cm^ . Entry across a hydrated interface

700

CURRENT DENSITY (/iA/cm^

Fig. 10. Current voltage relationship for Na"*" exit from a hydrated beta alumina interface into 0.01 M Na+ in propylene carbonate; (•) experimental; (—) calculated; at

25 °C, a = 0. 24, i = 3. 0 X 10-6 A/cm2; at 62 °C, a = 0.22, io = 2.4 X 10-5 A/cm2

.

336
should follow tJie same Tafel relationship that describes ion exit under

identical conditions. But, as ions are injected, large increases in interfacial voltage drop occur, the same phenomenon observed by Voinov and

Tannenberger [50] . The increased resistivity can be reversed by extended

electrolysis of ions out of the interface. ^3'-'"^ injection into the beta

alumina is a reasonable explanation for the phenomenon.

Farrington's work suggests that Na"*" ion transport across a beta alumina/

propylene carbonate interface occurs through an adsorbed intermediate state

in which the ion is neither within the beta aluirdna conduction plane nor solvated in the liquid electrolyte. At 1310 yA/cm^ interfacial overpotential

remains unchanged when 9 coulombs of Na"*" passed out of the interface.

However, at 329 0 yA/cm^ a steep voltage transition occurs after only two

coulombs have passed. This increase is repeatable, initial behavior being

restored at zero current. As the interface is increasingly hydrated, the

height of the transition and the current density above which it occurs

decrease

.

These data are consistent with the transport mechanism shown in Eqn. 2 According to this

Nat t Na^^,^ t Na"^^^ (2)

p ads pc

mechanism, interfacial transport is primarily influenced by species which

compote with Na"*" for surface adsorption sites, of which water is one, or influence the rate of Na"^ adsorption and desorption. Although the results

reported were for Na"*" transport, similar behavior can be expected with other

ions mobile within beta alumina such as Ag"^, Li"*", and K"*". The work summarized here clearly indicates that Na^ beta alumina/

propylene carbonate interfacial transport can be described in low current

densities by a Tafel relationship. This implies that the space charge region

within beta alumina at the interface is insignificant. However, the Tafel

relationship has been experimentally seen only at interfaces first exposed to

water vapor. Similar Tafel polarization surely occurs at a dry interface,

but its magnitude is small in comparison to the voltage drop arising from

current flow through the beta alumina and propylene carbonate and is,

therefore, not detected.

6.3 Studies of Multi -Phase Interfaces

The interface between Na -beta alumina and sulfur or a mixture of sulfur

and polysulfides , contained in a carbon fiber matrix, is the well known

example of a multi-phase boundary in the sodium-sulfur battery. This cell

operates at temperatures between 300°C and 350°C. Measurements of the cell

voltage as a function of the cell current revealed [62] an ohmic relation

between these two parameters, independent of the state of charge and of the

rechargeability of the cell. These results and impedance measurements [63]

lead to the conclusion that a hindrance of kinetic steps at the said interface does not exist at current densities of up to 100 mA/cm , referred to the

geometric area of the beta aliamina tube. Electrochemical processes involving

the sulfur electrode are responsible for the behavior of the cell impedance

towards the end of the discharge or charge cycle.

Acknowledgement

The support of the work on interfacial phenomena in sodium-sulfur cells

by EPRI is acknowledged.

337
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.

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,

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[19] S. V. Karpachev, Yu M. Ovchinikov, and L. D. Yushina, Elektrokhim. , 8_, 788 (1972)

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Electroanalyt. Chem. , 29, 301 (1971)

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[26] R. DeLevie, Electrochim. Acta. , 10 , 395 (1965).

338
[27] Yu V. Shirokov, V. S. Borovkov, P. D. Liikovtsev, Elektrokhim. , 5_, 12U (1969) .

[28] Yu V. Shirokov, P. D. Lukovtsev, V. S. Borovkov, Elektrokhim. , 1_, 575

(1971) . [29] Yu V. Shirokov, P. D. Lukovtsev, V. S. Borovkov, Electrokhim.

,

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(1971)

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[30] K. Kiukkola and C. Wagner, J. Electrochem. Soc , 104 , 379 (1957) .

[31] C. B. Alcock and T. N. Belford, Trans . Faraday Soc

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, 6 0, 822 (1964) .

[32] T. N. Belford and C. B. Alcock, Trans. Faraday Soc , 61 , 433 (1965).

[33] W. L. Worrell and J. L. Iskoe, Fast Ion Transport in Solids , W. van

Gool, ed., North-Holland, Amsterdam, 1973, p. 513.

[34] L. F. Donaghey and R. Pong, J. Electrochem. Soc , 121 , 506 (1974) .

[35] H. H. Mobius and B. Rohland, Z. Chem. Lpz. , _4, 15 8 (1966) .

[36] T. H. Etsell and S. N. Flengas , J. Electrochem. Soc , 118 , 1890 (1971).

[37] M. Kleitz, J. Besson and C. Deportes , Rev. Energ. Primaire , 2^, 35 (1966).

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[39] S. V. Karpachev and Yu M. Ovchinnikov, Elektrokhim. , 5_, 200 (1969).

I [40] C. S. Tedmon, Jr., H. S. Spacil, and S. P. Mitoff, J. Electrochem. Soc , 116 , 1170 (1969) .

[41] H. Yanagida, R. J. Brook, and F. A. Kroger, J. Electrochem. Soc. , 117 , 593 (1970) .

[42] F. A. Kroger, J. Electrochem. Soc , 120 , 75 (1973) .

[43] See also M. V. Perfilyev, Elektrokhim. , 7, 792 (1971)

.

[44] S. V. Karpachev and A. T. Filiayev, Elektrokhim. , 2_, 1330 (1966) .

^ [45] E. Schouler, M. Kleitz and C. Deportes, J. Chim. Phys. , 70 , 923 (1973).

[ [46] E. Schouler and M. Kleitz, J. Electroanalyt . Chem. , 64 , 135 (1975) .

I [47] S. V. Karpachev and A. T. Filiayev, Elektrokhim. , 2^, 617 (1966).

[48] A. T. Filiayev, I. D. Remez and S. V. Karpachev, 7, 257 (1972).

i [49] I. D. Remez, S. V. Karpachev, L. M. Solov'eva, V. N. Chebotin, and

1 A. T. Filyaev, Elektrokhim. , 11 , 292 (1975)

.

, [50] A. T. Filiayev, S. V. Karpachev and I. D. Remez, Elektrokhim. , 283

j (1970)

.

[ [51] G. C. Farrington and W. L. Roth, unpublished.

[52] R. D. Armstrong, T. Dickinson, and J. Turner, J. Electroanalyt . Chem. , 44, 157, (1973)

.

339
[53] A. Imai and M. Harata, Jap. J. Appl. Phys , 11 , 180 (1972).

[54] W. L. Fiedler, H. E. Kautz, J. S. Fordyce, and J. Singer, J. Electrochem

Soc. , 122 , 528 (1975).

[55] F. G. Will and S. P. Mitoff, J. Electrochem. Soc , 122 , 457 (1975).

[56] R. D. Armstrong, T. Dickinson, and J. Turner, J. Electrochem. Soc.

,

118 , 1135 (1971)

.

[57] M. B. Armand, Fast Ion Transport in Solids, W. vanGool, ed., American

Elsevier, New York, (1973), p. 670.

[58] M. S. Whittingham and R. A. Muggins, J. Electrochem. Soc , 118 , 1 (1971).

[59] G. C. Farrington, Report No. 75CRD193, GE CR&D, Schenectady, New York,

(1975); J. Electrochem. Soc , in press.

[60] M. Voinov and H. Tannenberger , Electrochim. Acta. , 19 , 959 (1974)

.

[61] G. C. Farrington, Report No. 75CRD146, GE CR&D, Schenectady, New York

(1975) ; J. Electroanalyt . Chem . submitted for publication.

[62] Sodium-Sulfur Battery Development for Bulk Energy Storage, EPRI 12 8-2,

Interim Report, September 1975.

[63] M. W. Breiter, Proc. 4th Australian Electrochem. Conf .

, in print.

340
National Bureau of Standards Special Publication 455, Electrocatalysis on Non-Metallic Surfaces

,

Proceedings of Workshop held at NBS, Gaithersburg, Md., December 9-12, 1975. (Issued November 1976)

Influence of Water on Beta Alumina Interfacial Ion Transport

Gregory C. Farrington

General Electric Research and Development Center

P. 0. Box 8 Schenectady, New York 12301

Water, adsorbed on a beta alumina surface, significantly influences the nature of Na"*" ion transport across

a non-aqueous electrolyte/beta alumina interface and the

behavior of a Pt/beta alumina interface. Na"^ interfacial

exchange current at 23°C between beta alumina and

propylene carbonate decreases with increasing surface

hydration. Adsorbed water undergoes a number of

faradaic reactions at a Pt/beta alumina interface,

which otherwise should approximate an ideally polarizable or "blocking" electrode. This paper discusses

the effects of adsorbed water on ion and electron

transport at these two interfaces

.

Key words: Beta alumina; blocking electrode; interface;

solid electrolyte; water.

Introduction

Beta alumina (ca. 1.2 Na20«9Al202) is a solid ionic conductor in which a variety of monovalent cations, including Na"*", Ag"^, H"*", and H3O+, are mobile.

Its structure [l]-*- and conductivity [2,3] have been previously discussed.

Breiter and Farrington [4] have reviewed the electrochemistry of charge

transport at a beta alumina interface. Farrington [5,6] has described

experimental investigations of Na"*" transport at liquid electrolyte/beta

alumina interfaces. Selected portions of these data will also be presented in

this discussion of the influence of adsorbed on charge transfer at

propylene carbonate/beta alumina and Pt/beta alumina interfaces.

and Na"*" both adsorb on an alpha alumina surface. Raman analysis of

water adsorbed on alpha alumina concludes that Na"*" ions and water molecules

compete for Lewis acid sites [7] . _^Na^ adsorption is thought to cause alphaalumina catalyst "poisoning" in Na -containing environments. Infrared spectroscopic analysis of pyridine adsorption on alpha alumina concludes that

pyridine adsorbs at non-protonic Lewis acid sites in the absence of water.

Proton activity is observed in the presence of water. Similar adsorption

behavior may be expected for beta alumina.

The interaction of molecular water and beta alumina may not be confined

to surface adsorption. Kline et ai. [8], in broadband NMR measurements,

found that the ^-^Na spectrum in beta alumina, which is sharply narrowed by

rapid localized ion motion at room t emperature , broadens in the presence of

water vapor. Heating restores a normal spectrum. Weight change measurements

corroborate the reversibility of the exchange [9] . This has been interpreted

Figures in brackets indicate the literature references at the end of this paper.

341
as evidence that molecular reversibly diffuses directly into the beta

alumina conduction plane, decreasing local Na"*" motion and conductivity.

Water is also a source of ionic hydrogen species such as 1130^ and h"*", both of which can replace_|_Na"'' in beta alumina. H^O^-beta alumina has been

prepared by exchanging Na beta alumina in concentrated H2SO4 [10] . H"'"-beta

alumina is synthesized by exposing Ag"*" beta alumina to a reducing atmosphere

of H2 [11]. Both H+ and H3O ions, therefore, are capable of diffusing

spontaneously or under the influence of a field into the beta alumina

structure

.

The energy required to bring about ion transfer between two dissimilar

chemical environments such as between solid and liquid electrolytes appears

as an interfacial polarization. Kinetic considerations [6] predict the i/V

relationship relating interfacial polarization, x] , and current density, i,

to be

in which a is the transfer coefficient, i^, the equilibrium exchange current

density, and z the charge of the mobile ion. The i/V relationship for Na"*" entry into and exit from beta alumina across a propylene carbonate interface

should be described by eq. (1) for which ion exit from the solid electrolyte

is considered a positive current occurring at a positive overpotential

.

Results of a preliminary examination [12] of beta alumina/propylene

carbonate interfacial transport at 25°C performed in this laboratory revealed

that the interfacial impedance for Na entry and exit increases after exposure

of the beta alumina to water vapor. Similar effects were noted by Voinov and

Tannenberger [13] who speculated that protons may be irreversibly driven into

the beta alumina under the influence of an electric field, increasing its

resistivity.

The following discussion describes the use of a galvanostatic technique

to distinguish interfacial polarization from effects of bulk resistivity

during Na"'" ion transfer across a propylene carbon ateA>eta alumina interface.

The influence of adsorbed on interfacial transport is described.

Cyclic voltammetry was used to examine the electrochemical behavior of

adsorbed H2O at a Pt/beta alumina interface. A Pt/Na-beta alumina interface

should (approximate an ideally polarizable electrode over a potential range

limited only by Na"^ reduction and 0 oxidation. This is a broader range of

ideal polarizability than is encountered with typical Pt/liquid electrolyte

interfaces. Adsorbed H2O should limit the range of ideal polarizability of a Pt/Na"'"-beta alumina interface since H2O can be oxidized to 1130'*", an ion mobile

within the beta alumina structure.

Experimental

Beta alumina ceramic discs were prepared from Alcoa 'XB-2' beta alumina

powder. The powder was ground with Zr02 grinding media, dried under vacuum,

screened, and hydrostatically pressed into rods approximately 1.6 cm. in

diameter and 15 cm long. The rods were cut into 1.6 cm segments and fired at

1700°C. The resulting ceramic was single phase, polycrystalline beta aliimina

of 9 8 percent theoretical density, containing approximately one percent MgO

and one percent Zr02

.

Rods were sliced into samples 1.3 cm in diamter and 4 mm thick. A 0.6 mm hole was drilled 3 mm into each sample from its circumference. Sample

faces were polished through 600A grit SiC paper, washed in methanol, and

fired for 24 hours at 700°C. Between individual experiments, samples were

342
wet polished with methanol through 600A grid SiC paper, rinsed with methanol,

and baked at 400°C at 10~^ torr for seven days or at 10"^ torr for six hours.

They were trans fered in vacuo to a Vacuum Atmospheres Ar drybox, polished

briefly on dry 600A grit SiC paper, and mounted in the experimental cell

shown in figure 1 between silicone rubber gaskets. In the following discussions beta alumina prepared by this procedure is referred to as "dry" or

I "anhydrous." "Hydrated" beta alumina is that which has been treated as above

but exposed to 100 percent relative humidity at 23°C for 20 minutes before

moiinting.

Burdick and Jackson "Spectroscopic" propylene carbonate was dried by

multiple percolation through a 16 cm x 2 cm coliimn of alumina which had been

activated at 500°C in vacuo for 24 hours. Baker anhdrous reagent Nal and

Southwestern Analytical anhydrous tetrabutylammonium tetraf luoborate (TBABF^) were baked 48 hours at 110°C in vacuo before dilution. Two electrolyte

solutions were prepared: 1) 0.010 M Nal + 0.100 M TBABF^ and 2) saturated

Nal and TBABF^. Solution 1 was the test solution and is referred to in this

I paper by its Na"*" concentration. The second was used in the unmonitored

compartment of the experimental cell.

The assembled cell (fig. 1) incorporated four electrodes, an outer

current-carrying pair of Pt foil (1 and 4) and an inner reference pair (2 and

j 3) . Au wire dipped into Tl amalgam was used for reference electrodes. A

small amount of electrolyte was injected into the sample reference electrode

hole to provide satisfactory ionic contact to the beta alumina.

0

b:

2

0

1

1

Fig. 1. Experimental cell; solid dots,

Teflon; open circles, silicone

rubber; slashes, beta alumina.

PL

5cm A

^r.r. ^° Study the behavior of a Pt/beta alumina interface, Pt electrodes,

1000 A thick, were sputtered onto both faces of a 1.2 cm diameter, 4 mm thick, polished disc of polycrystalline beta alumina. A reference electrode hole,

0.6 mm m diameter, was drilled 3 mm into the disc from its edge. Cyclic voltammetric results were carried out in vapor using a Au wire dipped in Tl (Hg) as reference electrode. A small amount of propylene carbonate

electrolyte was injected into the reference electrode hole to insure good electrode/beta alumina contact. Cyclic voltammetry results, thus, are referred to the T1°/T1''' couple in anhydrous propylene carbonate. The Tl°/Tl"'" couple has a standard potential of -0.34 V vs. N.H.E. in acidic aqueous media.

The technique of galvanostatic analysis has been described previously

[3,6]. In brief, it consists of applying single, fast-rise-time

constant current square waves of varying widths to electrodes 1 and 4 in figure 1 while monitoring the differential voltage response across electrodes 2 and 3. A typical result showing voltage vs. elapsed time during

a constant fliix of Na ions exiting hydrated beta alumina into propylene

343
1^

10" 10" 10',-3

I 1 1

1

10'

.

I

10" 10'

ELAPSED TIME (sec)

Fig. 2. Voltage (Vj-, + Vj^p) vs. elapsed time; Na+ exit

from hydrated interface at 1.38 x 10"'* A/cm^ ; Vj-) is bulk voltage drop, Vip is interfacial

polarization, 23°C.

carbonate is shown in Fig. 2. The time constants for bulk propylene carbonate

and beta alumina resistivity (Vj^) are much shorter than the rise time of the

square wave (tj- 10"^ sec) , so Vj^ is constant from 10"^ sec to dc. Interfacial polarization (Vj^p) , in contrast, has a time constant of approximately

10"-^ sec. It is this difference between bulk and interfacial time constants

which permits isolation of interfacial polarization by this technique.

Results - Interfacial Ion Transport

Figures 3_j_and 4 present plots of total voltage (Vj_ + Vj^) V£. applied

current for Na entry and exit across a dry beta alumina/propylene carbonate

interface. V- is undetectable within measurement accuracy. Both plots are ohmic over the current ranges shown and have identical slopes corresponding

to the resistance expected for bulk beta alumina and propylene carbonate.

Fig. 3. Voltage (Vb + Vip) vs. current

density; Na"*" entry into dry

beta alumina; slope = 1.00;

23°C.

Fig. 4. Voltage (Vj^ + Vip) vs. current

density; Na+ exit from dry

beta alumina; slope = 1.00;

23°C.

344
In contrast, significant charge transfer polarization is detectable

during ion transport across a hydrated interface. Typical resistivity/time

response is shown in Figure 2. The potential attributable to interfacial

polarization is readily separable from bulk resistivity • Figure 5 shows

I characteristic plots of Vj_ vs. current density for Na exit from a hydrated

interface at two temperatures.

700

CURRENT DENSITY (/lA/cm^)

Interfacial polarization for Na"*" exit from a hydrated interface clearly

follows the Tafel relationship described in ^q. (1) for which the transition

1; coefficient (a) was fo\ind to be 0.24 at 23°C and 0.22 at 62°C by plotting

1

1 In i vs . n . The interfacial exchange current (i^) is 3.0 x 10"^ A/cm^ at

l' 23°C and 2.4 x 10"^ A/cm^ at 62°C. Tafel behavior implies that any space

charge region within the beta alumina is insignificant, a result that was expected for such an electrolyte having a large density of charge carriers.

Non-ohmic transition polarization must also arise during ion transport

across an anhydrous interface. However, the Na"^ exchange current appears to

be much larger at a dry interface and polarization is_^too small to be

detected. A lower limit on i^ for dry interfacial Na transport, estimated

from figure 3, is 2 x 10""^ A/cm^ , at 23 °C.

These results illustrate the dramatic effect of water on beta alumina

interfacial transport. Adsorbed water decreases the equilibrium exchange

current for Na^ at a propylene carbonate/beta alumina interface. Previous

results of Farrington [5] indicate that this decrease is a function of the

I extent of surface hydration, the exchange current increasing with decreasing

|| hydration.

Tafel-type polarization observed for ion exit from a hydrated interface

j, should also occur during ion entry through such an interface. However, the

I' interfacial voltage drop is not constant but steadily increases during ion

flux into beta alumina through a hydrated interface. This effect is

significant over the entire range of current density examined in this work,

10"^ to 10~2 A/cm2 . Initial interfacial behavior is restored by extended

electrolysis of ions out of the solid electrolyte.

These results are consistent with the hypothesis that reversible H3O''' substitution into beta alumina is the crucial factor behind the interfacial

voltage drop increase during ion entry. Water, adsorbed on the beta alumina

interface, is a source of H30'*" ions which can siibstitute for Na"^ in the beta

alumina structure. Voinov and Tannenberger [13] suggested and Farrington [5]

experimentally corroborated that H30'^ or a similar ionic hydrogen species

competes with Na"^ for injection into beta alumina at a hydrated interface.

345
H3O"'" substitution is accompanied by an increase in electrolyte resistivity

which is manifested as a gradual increase in interfacial voltage drop during

ion entry

.

Further studies suggest that interfacial Na^ transport occurs through an intermediate state in which Na"*" is adsorbed on the beta alumina surface.

Under anhydrous conditions (fig. 6) , the total voltage during ion exit undergoes a sharp increase after approximately 1 coulomb has passed at a current

density of 329 0 yA/cm^ (curve A) . At a lower current, 1310 yA/cm^ (curve B)

,

no transition is observed even with passage of 10 coulombs.

Hydration of the interface decreases the current density at which a voltage transition is observed and decreases its sharpness as figure 7 demonstrates. Curve A was observed with a beta alumina disc first exposed to

100 percent relative humidity for 20 minutes then baked in vacuo for six hours

at 400°C and curve B with a disc that had been hydrated without bakeout.

Fig. 6. Voltage (Vj-j + Vip) vs. coulomb

density; Na+ exit from dry interface; curve A, 3.290 x 10"^

A/cm2; 23°C; curve B, 1.310 x 10" 3 A/cm2.

Fig. 7. Voltage (Vj_, + Vj^p) vs. coulomb

density; Na+ exit from beta

alumina; A, partially dried

beta alumina; B, hydrated beta

alumina; both at 3.290 x 10~3

A/cm2 and 2 3°C.

All three plots are consistent with the hypothetical mechanism for Na

transport acros^ a beta alumina/propylene carbonate interface shown in eq.

(2) in which Na exists in a surface adsorbed state intermediate between bulk

beta alumina and propylene carbonate solvation.

N^%Al203 - N-^ads - ^-'pc

According to this mechanism, ion solvation occurs as a step in maintaining

adsorption equilibrium and is independent of ion transport into or out of the

beta alumina structure so long as the transport rate is less than the

adsorption/desorption rate. Solvent structure and dielectric constant should

exert their most significant effects on interfacial transport in influencing

adsorption/desorption rate and equilibrium.

346
Pt/Beta Alumina Interface

The effects of surface adsorption and H3O injection into the

conducting planes of Beta aliunina can be independently observed by cyclic

voltammetry at a Pt/beta alumina interface. Results are shown in figure 8. Little cell current is observed upon initial polarization at 0 V vs.

Tl°/Tl"'". Positive potential sweep encounters several oxidation peaks (A-C)

which can be attributed to oxidation to H30'''. Sweep reversal produces

one significant peak at D which can be associated with reduction of a product,

presumably O2 , formed at A-C.

At potentials less than -0.40V, a second major reduction peak appears

(E) . The current increases dramatically beyond -0.75V. Sweep reversal

reveals peak F, not observed on initial cell polarization. The rapid decay

of current in F suggests it corresponds to depletion of a particular species

previously produced in peaks D or E.

The characteristics of peaks E and F are consistent with the hypothesis

that E corresponds in part to reduction of adsorbed with simultaneous

Na"^ migration out of the solid electrolyte to the surface. The overall

reaction is

4H„0 + 4Na"'", ^ + 4e 2H„ + 4NaOH (3)

2 beta 2

Peak F, then, represents reoxidation of H2 and reinjection of Na^ into beta

alumina. Both H2 and Na"*" exist at the interface only as the result of

electrochemical production at E.

As the complete sweep cycle shown in figure 8 is repeated, all peaks

decrease in magnitude, particularly A-C. Evidently, H^O^ ions gradually

substitute into the beta alumina adjacent to the interface and significantly

decrease its conductivity. This observation is consistent with the results

previously described which indicate that extended Na"*" entry into beta alumina

across a hydrated interface leads to a progressive increase in interfacial

voltage, presiimably from buildup of 1130^ ions in the beta aliimina.

347
Summary and Conclusions

No interfacial polarization can be detected at an extremely dry propylene

carbonate/beta alumina interface from 3 x 10"^ to 1 x 10 A/cm^ at 23°C.

If the beta alumina is first exposed to water vapor, significant interfacial polarization is apparent. Polarization follows a Tafel relationship

for which the transition coefficient (a) is 0.24 at 23°C and 0.22 at 62°C.

The interfacial exchange current (Iq) is 3.0 x 10~° A/cm2 at 23°C and 2.4 x 10-5 A/cm2 at 62°C.

Na^ ion transport across a beta alumina/propylene carbonate interface

occurs through an intermediate adsorbed state in which the ion is neither

within the beta alumina structure nor solvated in the non-aqueous electrolyte.

Interfacial Na"*" ion transport is primarily influenced by those species which

either compete with Na"*" for surface adsorption sites, of which water is one, or influence the rate of Na"*" ion adsorption and desorption. This latter step

should be affected by interfacial ionic composition and by the solvent and

its dissolved secondary components

.

adsorbs at a Pt/beta aliimina interface and can be oxidized to ionic

species, presumably H3O"'", mobile within the beta aliamina structure. The

possible presence of trace must be considered in any experimentation

assuming blocking behavior at a Pt/beta alumina interface or its equivalent.

Meticulous interface preparation including isolation from normal atmosphere

must be exercised if an inert metal/beta alumina interface is to approximate

an ideally polarizable, or blocking electrode. The presence of trace water

is surely responsible for the satisfactory behavior of Pt and Au contacts in

low frequency impedance measurements on Na'^-beta alumina [14,15,16] and may

also account for unexpectedly large capacitance observed in several investigations [14,17] of Pt/beta alumina/Pt or Au/beta alumina/Au cells.

Acknowledgements

Dr. R. W. Powers kindly supplied samples of beta alumina used in this

work.

References

[1] W. L. Roth, J. Solid State Chem. , 4, 60 (1972)

.

[2] M. S. Whittingham and R. A. Muggins, J. Chem. Phys

.

, 54 , 414 (1971),

[3] G. C. Farrington, Report No. 75CRD193, GE CR&D, Schenectady, New York

(1975); J. Electrochem. Soc , in press.

[4] M. W. Breiter and G. C. Farrington, this volume.

[5] G. C. Farrington, Report No. 75CRD146, GE CR&D, Schenectady, New York

(1975) ; J. Electrochem. Soc , submitted for publication.

[6] G. C. Farrington, J. Electroanal. Chem. , submitted for publication.

[7] G. Careri, V. Mazzacurati, M. Sampoli, and G. Signorelli, J. Catal. , 26^, 494 (1972) .

[8] D. Kline, H. S. Story, and W. L. Roth, J. Chem. Phys. , 56, 5180 (1972).

[9] W. L. Roth, General Electric Research and Development Center, Report

No. 74CRD054, (1971)

.

[10] H. Saalfeld, H. Matthies , and S. K. Datta, Berichte Deutsche Keram. Ges . 45, 212 (1968)

.

343
[11] G. P. Kummer, J. Solid State Chem. , 7, 141 (1972).

[12] G. C. Farrington, abstract, Symposivim on Interfacial Phenomena, Amer.

Inst. Chem. Eng., 67th Meeting, December 1974, Washington, DC.

[13] M. Voinov and H. Tannenberger , Electrochim. Acta. , 19 , 959 (1974)

.

[14] A. Imai and M. Harata, Jap. J. Appl. Phys

.

, 11 , 180 (1972)

.

[15] J. H. Kennedy and A. F. Sammells, J. Electrochem. Soc. , 119 , 1609 (1972).

[16] R. W. Powers and S. P. Mitoff, J. Electrochem. Soc , 122 , 226 (1975).

[17] S. P. Mitoff and R. J. Charles, J. Appl. Phys. , 43 , 927 (1972).

349

National Bureau of Standards Special Publication 455, Electrocatalysis on Non-Metallic Surfaces,

Proceedings of Workshop held at NBS, Gaithersburg , Md. , December 9-12, 1975. (Issued November 1976)

Overvoltage Phenomena in Solid-Oxide Electrolyte Cells at Elevated Temperatures

W. L. Worrell

Department of Metallurgy and Materials Science

University of Pennsylvania

Philadelphia, Pennsylvania 1917^

To develop a more quantitative understanding of the factors which

establish overvoltages in solid-oxide electrolyte cells, constant

currents have been passed through two types of symmetrical cells. At temperatures "between 800 and 1000°C, constant and reproducible cell

overvoltages are obtained when the imposed current is less than 50 ]iA. The cell overvoltage (ri) is directly proportional to the cmrrent (l) , and the interfacial resistance, defined as equal to r\/l, is constant

for each cell at constant temperature. Assuming that oxygen concen- tration gradients in the metal portion of the electrodes establish the

cell overvoltages, a quantitative expression has been derived to calculate the interfacial resistance for each cell. Reasonable

agreement between calculated and experimental values of the interfacial

resistance is obtained for cells having two-phase metal-metal oxide

electrodes. This agreement indicates that oxygen is transferred

between metal particles and the electrolyte rather than through oxide

particles in the electrodes. For cells with oxygen- saturated copper

electrodes, the solubility-diffusivity product of oxygen in solid

copper has been calculated from the overvoltage data.

Key words: Copper electrodes; interfacial resistance; metal-metal

oxide electrodes; oxygen transport; steady-state overvoltages.

1. Introduction

An understanding of the factors which establish cell overvoltages is important because

large over^raltages are detrimental to the useful application of many electrochemical cells.

The chemical potential difference between two electrodes establishes the open-circuit or thermodynamic cell voltage. Under closed-circuit conditions, the cell voltage is usually

less than the thermodynamic value, and the difference between the thermodynamic and the

measured voltage is the overvoltage. Phenomena which are responsible for the overvoltage can be conveniently divided into interfacial and transport effects. Interfacial effects

include slow electron transfer and chemical reactions at the electrode-electrolyte interface. Slow transport in the electrolyte or in the electrodes causes concentration

gradients in the respective phases.

Electrochemical cells using zirconia-based electrolytes have been used in a number of

coulometric titration and oxygen diffusivity studies [l].-'- Raleigh has reviewed studies of

electrode processes in solid-state cells [2]. Except for ovcc work with two-phase electrodes

[3,^], previous investigations have been conducted using single-phase solid electrodes. In

''rigures in brackets indicate the literature references at the end of this paper.

351
this paper, our experimental measurements of steady-state overvoltages in cells using a Zr02(CaO) electrolyte at elevated temperatures are summarized, and the method used to quantitatively calculate such overvoltages is described.

2. Experimental Aspects

In our overvoltage investigations using a zirconia-hased electrolyte and solid

electrodes, two symmetrical electrochemical cells were employed [3-5]. Two-phase metalmetal oxide electrodes were used in the first cell.

M, M0|Zr02(Ca0) |m, MO, (A)

where M, MO was Cu, CU2O; Fe, FeO or Ni,NiO. These metal-metal oxide electrodes are most

commonly used in thermodynamic and kinetic cell studies [l]. In cell (A) the zirconia

electrolyte was sandwiched hetween two identical electrode pellets. The electrolyte was a commercial-grade Zr02 (T.5 'wt. pet CaO) disc, 0.25 to 0.6 cm thick and ~1.2 cm diam.

,

obtained from the Zirconiiom Corp. The preparation and characterization of M, MO electrode

pellets have been described [3,^].

In the second cell,

Cu, Cu20|Cu|Zr02(Ca0) |Cu|Cu, CU2O (B)

the electrodes were oxygen-saturated copper foils of precisely known thickness. Copper

foils 3.2 X 10~3 cm in thickness were used in one series of experiments, while foils having

a thickness of 1.27 ^ 10"^ cm were used in another series. To saturate the copper

electrodes with oxygen, cell (B) was heated to a temperature of 1060°C and was held for

approximately two days before beginning experimental measurements. To maintain oxygen

saturation in the electrodes, the copper foils were sandwiched between the electrolyte and

Cu, CU2O pellet.

Two critical experimental precautions were necessary to obtain reproducible results.

Although measurements were made imder an atmosphere of purified argon, metal foils of

copper or iron were placed within 0.5 cm of the cell to minimize any harmful reactions

between impurities in the argon and the electrodes. Intimate electrode-electrolyte contac was also essential for accurate measurements. Preliminary measurements showed that

excellent interfacial contact could be obtained by using an alumina push-rod to press the

electrodes and the electrolyte together. The alimiina rod could be moved by adjusting a brass screw at the top of the reaction tube. Using this design, the measured steady-state

cell voltages were stable and independent of the argon gas flow rate. After an experiment

the interfacial adherence was so strong that the electrolyte would fracture internally

rather than separate from the electrode.

Detailed descriptions of our experimental apparatus and technique are available [3-5]j

and only a brief summary is necessary here. The experimental technique was a galvanostati

one in which constant currents were imposed across cells (A) and (B), and the cell voltage

was measured as a function of time at temperatures between 8OO and 1000°C. The voltage

drop across a standard decade resistance box provided a highly precise measurement of the current through the cell. Cell voltages were measured (±0.02 mV) using a Keithly 63O high

impedence (lO-'-^ ohms) potentiometric electrometer. When the imposed current was above

50 liA, the cell voltage continuously increased with time and never reached a steady-state

value. Figure 1 shows an example of this behavior for cell (A) with Cu, CU2O electrodes.

At lower currents, steady-state cell voltages were obtained as shown in figure 2. Only

steady-state values were used in our quantitative analysis.

352
10

tlJ o

Cu, CU2O

CELL (A) 900 °C 120 /xA

t, TIME (min)

Fig. 1. Typical cell coltage versus

time behavior for cell (A) at currents above 50 \\A. The electrodes are Cu, CU2O

mixtures

.

14

O

>

0.4 O

2.42 /iA

J L

Cu, CU2O

CELL (A) 900 t

Fig, 2. Typical cell voltage versus

time behavior for cell (A) at currents below 50 JjA. The electrodes are Cu, CU2O

mixtures

.

6 8

t, TIME (mm)

10 12 14

For cells (A) and (B), the steady-state cell voltage (V) is the sum

across the electrolyte and the overvoltage (r|); thus.

of the in drop

V = IQ + T). (1)

To determine the overvoltage, the dc resistance (fi) of the cell must be known. With the

Ni, NiO electrodes, the steady-state cell voltages were very large, and the ac cell resistance

could be used for f2 in Eq. (1) with negligible error. For these cells, the term was less

than 1% of the cell voltage. However, a direct measurement of the dc cell resistance was necessary for the other cells. Therefore after the steady-state voltage measurements, the dc

resistance was measured directly by taking oscillographs of the change in cell voltage during

the first tenths of a second after current was started or stopped.

353
3. Results

Values of r) calc\ilated using eq.(l) for cell (A) at 900°C are shown in figure 3. Eesults are shown on a log ri versus log I plot so that results from the three different

electrode cells can he compared. Figure 3 is not a Tafel plot. Over the experimental

range of 0.5 to 50 UA, a linear relationship is ohser^ed between log ri and log I. The

slope of each of the three lines in figure 3 is unity; thus, a constant interfacial

resistance (Rn) is observed, where

= n/i. (2)

At 900°C Rr^ is 12±2 ohms for the Cu, Cu20 cell, 120±20 ohms for the Fe, FeO cell and

1200±200 for the Ni , NiO cell. Similar results have also been obtained for cell (A) at 800 and 1000°C. However, values for R^ at 800°C are higher, while values of Rj^ at 1000°C

are lower than those at 900°C.

Steady-state overvoltages determined for cell (b) at constant temperature and constant

electrode thickness also exhibit the behavior shown in figixre 3. Experimentally

determined values of the interfacial resistance for cell (b) are tabulated as a function of

temperature and electrode thickness in Table 1.

100

> 10

UJ<

I-

_jo

>

>o

0.1

0.01

1 900°C

1

^ Ni.NiO

D Fe, FeO — o Cu, CU2O

1 1

Fig. 3. Overvoltage versus current for cell

(A) at 900 °C. Results are shown

for three different two-phase electrode cells; the three electrodes

are Cu, CU2O; Fe, FeO; and Ni, NiO.

1.0 10 100 1000

I, CURRENT (/lA)

Table 1. Values of for cell (B)

.

3.2 X 10"^ cm Cu Electrode 1.27 X 10"^ cm Cu Electrode

T, °C R^ , ohms T, °C , ohms

800 1I+5.5 796 625.0

8^1 853 255.0

900 21.7 .903 106.0

9\\ 12.0 951 U9.2

1000 ^.5

354
h. Discussion

Only oxygen ions move through the zirconia electrolyte, and the cell current density

(i) is directly related to the oxygen flux (J) in the cell by

i = 2FJ, (3)

where F is the Faraday constant. Because the metal particles or foil at each electrode are saturated with oxygen, oxide should he reduced at one electrode and formed at the other when

oxygen ions are transported across the electrolyte. Assuming all the oxygen which enters

the sink electrode forms oxide, one calculates a maximum formation rate for the oxide layer

of 0.1+ to 0.7 yA~-^iiin~"'". Th\is , for a typical experimental current of 10 \iA for 10 min,

a 1;0-T0 X thick layer of oxide is formed.

k.l Calculation of Interfacial Resistance.

The basic question is what establishes the constant interfacial resistance (Rr|) when

the cell current is less than 50 ]iA. The most likely cause of R^^ is either a slow chargetransfer reaction at the electrode-electrolyte interface or slow transport of oxygen in the

electrodes. It should be noted that oxygen-ion transport in the electrolyte determines the

ohmic resistance (fi) of the zirconia electrolyte and makes no contribution to the interfacial

resistance. Overvoltage decay times of 5 to 150 min and the absence of a Tafel slope indicates that the observed overvoltages are not caused by a slow charge- transfer reaction.

Thus, slow oxygen transport in the electrodes must be responsible for the observed overvoltages and for R^j

Oxygen transport in the two-phase electrodes of cell (A) could be through either the

metal or oxide particles. However, for cell (B), oxygen transport through the oxygensaturated copper foil electrodes is the only possibility. Assuming that oxygen concentration gradients in the metal (either particles or foil) electrodes establishes the cell

overvoltage, the following eq.\iation has been derived [^,5J to calculate the interfacial

resistance for cells (A) and (b):

RTAx » nF A CqDo

where R is 8.3lU J mol~''"K~"'", F is 96,1+87 J V~"'"eq~''", n is 2 eq mol"""" for oxygen, T is in K,

A(cm2) is the electrolyte-metal electrode interfacial area, Ax(cm) is the oxygen diffusion

distance in the metal electrode, and CqDo (mol cm'-^s"-*-) is the solubility-diffusivity

product for oxygen in the metal.

2 Estimates of the metal particle-electrolyte interfacial area (A 0.5 cm ) and of the

average metal particle size (2 x 10~3 cm) have been made using photomicrographs of the

electrode-electrolyte interfaces of cell (A). The average oxygen diffusion distance (Ax)

in the metal particles is estimated to be 1/3 of the average particle size. With these

assumptions and available data for CqDo, values of Rr| calculated from eq.(U) are compared

with the experimental values from cell (A) at 900°C in Table 2.

As shown in Table 2 , calculated and measinred values are in good agreement for the

Cu, CU2O electrodes and in fair agreement for the Fe, FeO and Ni, NiO electrodes. Some

uncertainty in the calculated values is due to the oxygen solubility-diffusivity product.

Although CqDo for copper and iron is fairly well established (within a factor of two), the

data for nickel are very imcertain and could be in error by as much as a factor of 10 to 50

[11]. Another major uncertainty in the calculated values shown in Table 2 arises from the

estimation of Ax, which is probably accurate only within a factor of four.

A major advantage of cell (B) is that the diffusion distance (Ax) is simply the thick355
Table 2. Comparison of calculated and experimental values of R-^ at 900 °C for cell (A)

.

Electrodes (CqDo)^ in Metal Rri (cal) R,-| (meas)

Cu, CU2O 1.5 X 10 LoJ 5 12

T.l IQ-"^"'" [5] lU 12

Fe, FeO

-12

2.5 X 10 [6] 279 120

Ni, NiO l.k X lo""''-^ [7] U988 1200

Units are mol cm~ls~l.

ness of the copper foil electrodes, which can he precisely fixed and varied experimentally.

At constant temperature, the interfacial resistance measured from cell (B) should he

directly proportional to the electrode thickness. As shovn in Table 1, this is ohserved

within the experimental uncertainty. The measured values of Rf-| tabulated in Table 1 can he

used with eq.(^) to calculate values of CqDo for copper. For this calculation, the interfacial area for cell (B) is O.785 cm^. In figure h, calculated values of CqDo are compared with results from three other investigations [8-10] in which different experimental

techniques were used. The good agreement shown in figure k indicates that eq.(ii) is valid

and that oxygen concentration gradients in the metal electrodes do estahlish the over- voltages in cell (B).

9,0

9.5

° 10.0 oo

o

10.5

.0

11.5

^ 0.0032 cm Cu ELECTRODE

O 0.0127 cm Cu ELECTRODE

«^ y Pastorek and Rapp [8]

7.5

— Ramono, Rao and Tore [l5]

8.5 9.0 I/T X 1 0"^

Fig. 4. The solubility-diffuslvlty

product of oxygen in solid

copper as a function of reciprocal temperature. The points

and solid line represent values

for CqDo calculated using eq.

(4) and overvoltage data from

cell (B) . The three dashed

lines are experimental results

from other investigations

18-10]

.

10.0

k.2 Further Comments Concerning Equation (^4)

A numher of ancillary experiments [h] showed that the overvoltage of cell (A) is

independent of the direction of current and is produced equally at each electrode. During

current flow, oxygen enters the sink electrode hy diffusing through the oxygen-saturated

metal and forming oxide at an oxide-metal interface. During a typical experiment,

calculations using eq.(l) indicate a UO-70 2. layer of oxide should form, which is a negligible thickness compared to the estimated diffusion distance of 6.7 ^ IQ-^ cm for cell

(A). At currents below 50 yA the diffusion distance is essentially constant, and constant

values of are obtained. At higher currents, the oxide layer thickness could become

356
significant, and the oxygen diffusion distance coiild vary with time. For oxygen to enter

the metal electrodes, there must be a small supersaturation of oxygen in the metal at the

electrode-electrolyte interface. For an overvoltage value of 10 mV, the ratio of the

supersaturated oxygen content to the solubility value is only 1.2. Currents above 50 yA

produce a large oxygen flux into the metal electrode, and the supersaturation oxygen content

of the metal could become so large that oxide layers are nucleated at the electrodeelectrolyte interface. Under these conditions steady-state overvoltages for cell (A) would

never be observed.

The agreement shown in Table 2 requires that the major oxygen transport across electrode-electrolyte interfaces in cell (A) is between the electrolyte and metal particles

in the electrodes. There is evidently negligible oxygen transport between the electrolyte

and metal oxide particles in the electrodes. Ductile metal particles in the electrodes can deform under cell pressure to fit into microcracks and crevices on the electrolyte surface,

which results in intimate contact between the metal particles and the electrolyte. Our

initial results were a function of cell pressure until press\ires were high enoiagh to insure

good electrode-electrolyte adherence.

In the derivation of eq.(U), short circuit diff\ision paths for oxygen such as grain

boimdaries and siirfaces of the metal particles have been ignored. Experimental temperatures

are close to the melting points of the metals (particularly copper), which minimizes the

influence of short-circ\iit diffusion.

Our experimental approach can be used to investigate overvoltage phenomena in other

electrochemical cells. However, eq.(l+), or one similar to it, applies only to those cases

in which cell overvoltages are established by concentration gradients in the electrodes.

One example is a study of the oxygen transfer kinetics between Ga, Ga203 electrodes and a Zr02(CaO) electrolyte at temperatiires between 800 and 900°C [12]. The experimental results

followed eq,.(U), and it was concluded that oxygen transport thro\igh liquid gallium in the

two-phase electrodes established the cell overvoltages.

5. Conclusions

Assuming that oxygen concentration gradients in the metal portion of the two-phase

electrodes establish the overvoltages in cell (A) , calculated values of Rf) are in reasonable

agreement with experimental values. This agreement indicates that the major oxygen

( transport across the electrode-electrolyte interface is between the electrolyte and metal

j particles in the electrode. As shown in figure 3, overvoltages produced by Cu, Cu20

ji electrodes are ten times less than those produced by Fe, FeO electrodes and 100 times less

; than those produced by Ni, NiO electrodes. Thus, Cu, Cu20 electrodes are the preferred

! two-phase electrodes for use in coiilometric titration and diffusivity studies.

Overvoltage data from cell (B) have been used with eq.(l+) to determine values for the

1^ oxygen solubility-diffvisivity product in solid copper. The good agreement between these

values and other data shown in figure k indicates that steady-state overvoltages in cell (B) are established by oxygen concentration gradients in the copper foil electrodes.

Equation (k) can be used to calculate and overvoltages for any solid-oxide electrolyte

^ cell in which the overvoltages are established by slow oxygen transport in oxygen-saturated

{ metal electrodes.

6. Acknowledgment

The financial support of this program by the National Science Foundation and the

Laboratory for Research on the Structure of Matter at the University of Pennsylvania is

^ grate fxilly acknowledged.

357
7. References

Worrell, W. L. , Amer. Ceram. Soc. Bull. 53, ^+25 (197^+).

Raleigh, D. 0., Electroanal. Chem. 6, 87 (1973).

Worrell , W. L. and Iskoe , J . L. in Fast Ion Transport in Solids , Solid State

Batteries and Devices , (W. van Gool, Ed.) North-Holland Publ. Co., 513 (1973).

Iskoe, J. L. , Ph.D. Dissertation, University of Pennsylvania (1972).

Ramanarayanan, T. A. and Worrell, W. L. , Metal. Trans. 5_> 1773 (l97^+).

Swisher, J. H. and Turkdogan, E. T. , Trans. TMS-AIME 239, ^26 (1967).

Barlov, R. and Grundy, P. J., J. Mat. Sci. £, 797 (1969).

Pastorek, R. L. and Rapp, R. A., Trans. TMS-AIME 2h^, 1711 (1969),

Wood, S., Adamonis, D., Soffa, W. A., and Meier, G. H., ONR Report NRO 31-7^9 (1973).

Ramana Rao, A. V. and Tare, V. B, , Z. Metallkund. 63, 70 (1972).

Kerr, W. R. , and Rapp, R. A., to be published.

Donaghey, L. F. and Pong, R. , J. Electrochem. Soc. 121 , 506 {l9lh)

.

358
PANEL DISCUSSION
i

i

ij

I
Panel on Major Problems, Advances, and Opportunities

Revealed by Workshop

Chairman: Prof. Dr. Heinz Gerischer, Fritz Haber Institute

der Max Planck Gesel 1 schaf t, Berlin

Panelists: Dr. William Gadzuk, National Bureau of Standards

Dr. Jerome Kruger, National Bureau of Standards

Prof. Robert de Levie, Georgetown University

Dr. Ralph J. Brodd, Union Carbide Corp.

Dr. Ugo Bertocci , National Bureau of Standards

Prof. Allen J. Bard, University of Texas at Austin

Prof, A. C. C. Tseung, The City University, London

Prof. Norman Hackerman, Rice University

Dr. Sigmund Schuldiner, Bureau of Mines

Prof. J. Ross MacDonald, University of North Carol

i

Dr. Douglas 0. Raleigh, Rockwell International

Reporter: Dr. Alan D. Frankin, National Bureau of Standards

Panel on Major Problems, Advances, and Opportunities

Revealed by Workshop

A. D. Franklin

Reporter for the Panel

Institute for Materials Research

National Bureau of Standards

Washington, DC 20234

Summary of Discussion

Theory

The theoretical papers read at the Workshop dealt mainly with chemisorption rather

j than with electrocatalysis as such. Because of the great complexity of real catalytic

systems, it cannot be expected that theory at present will be able to deal in detail with

catalytic problems. Rather, theory can develop ideas and concepts which can guide researchers

I in developing strategies for investigating and understanding real catalytic systems,

'i

j Two trends in chemisorption theory are evident. On the one hand, calculations are

I being made of the electron states of clusters of atoms that represent small particles,

including adsorbed species. Molecular orbital theory plays a dominant role here. Electron

energy spectra can be calculated and compared to experimental observations, especially

for energy differences. Of particular interest would be calculations on alloy clusters,

to explore size effects on alloy catalytic properties, and on clusters containing lattice

defects, although clearly the cluster size would have to be large enough to make a lattice

defect meaningful

.

This kind of theory is closely related to the chemists' view of specific chemical

bonds between adsorbate molecules and atoms of the catalyst surface. Such theory is useful to understanding mechanisms and reaction paths, but is not reliable for predicting

energy levels of electrons in surfaces and chemical bonds because of neglect of delocaliz-

« ing effects of the electron states in the solid.

On the other hand, band theory has been extensively used to describe electrochemical

reactions on semiconductors based on experimental analyses on the energetic position of

band edges, at the interface between semiconductors and electrolytes. Specific adsorption

sites are introduced as surface states, and the nature of these states is the probable

meeting ground for the molecular orbital and band approaches. Indeed, Wolfram showed

that the two descriptions converge when considering surface states in the energy gap of

the semiconductor, both giving localized electron states capable of interacting strongly

with anti-bonding states of adsorbed species. That this kind of theory is applicable to electrochemical situations is demonstrated by the observation in Wolfram's calculations

of a group of surface states about 1.5 eV below the conduction band in n-type Ti02,

agreeing well with a peak in the same energy in the electron energy loss spectrum for

reduced Ti02 in vacuo, reported in the paper by Henrich, Zeiger, and Dresselhaus, and

with a peak found by Lu, Pollak, and Raccah near 1.7 eV in the electro-reflectance spectrum

of reduced Ti02 in contact with 1 M aqueous KCl solution. In addition, Frank, Laser,

Hardee, and Bard in their paper report a concentration of surface states about 1.2 eV below the conduction band in n-type Ti02, from redox behavior of a number of redox couples

in acetonitrile with Ti02 as the working electrode.

Characterization

Adequate characterization of electrocatalyst specimens, like the characterization of

materials in many other kinds of studies, involves determining a number of things. The

specimen purity and stoichiometry must be determined, as the potential role of lattice

oxygen in the activity of WC for the electrochemical oxidation of H2 shows. The paper of

Morrison and Freund demonstrates that lattice defects can be important, and Tseung

and Jasem showed that activity for O2 reduction may correlate with magnetic susceptibility.

359
Analytical chemistry techniques for measuring very low concentrations of reactants

and products in the electrolyte near the electrocatalyst can be helpful in sorting out

the reactions taking place and the concentration levels at which impurities might become

important. An example is found in the oxygen electroreduction, where a variety of intermediate oxygen-containing species occur, and the problem arises of assessing the balance

between the reduction paths leading to H2O and H2O2.

Techniques for probing the nature of the surface and the species formed on it by

chemi sorption fall into two classes. On the one hand, there are very sensitive and

powerful methods for studying surfaces under high vacuum or other ultra-clean conditions,

such as LEED, ESCA, Auger spectroscopy, etc., while on the other hand, there are now a number of techniques capable of studying the surface in situ, in contact with an electrolyte and with an applied potential. These include optical techniques such as internal

reflection absorption and Raman spectroscopy and electroref lectance, and also include

electrochemical methods like linear sweep voltammetry and a.c. impedance. The paper by

Conway reviews the tremendously useful information available from linear sweep voltammetry,

and Bard and coworkers, in their paper quoted above, use linear sweep voltammetry with a series of redox couples in a non-aqueous solvent to probe the surface states of the

electrode material.

The problem with applying the ultra-clean techniques to understanding electrochemical

processes is how to assess the correspondence between the nature of the surface as seen

by the measurement and the surface as it exists in the electrochemical situation, in contact with the electrolyte and under the influence of the applied potential. In general,

changes undoubtedly take place on the surface as the sample is removed from the electrolyte

and prepared for and transferred to the ultra-clean environment. However, examples do

exist in which good correspondence has been seen between measurements made under electrochemical conditions and those on ostensibly the same surfaces under high vacuum or controlled gas environments.^ for instance, the phenomena of under-potential deposition

discussed in the paper by Adzic, and coworkers arises because the substrate-adsorbate

interaction is energetically stronger than is the self-interaction of the adsorbate. In effect, the heat of adsorption of the first monolayer is larger than the heat of vaporization of the bulk. It is therefore expected, provided that the surfaces are essentially

the same, that the amount by which the under potential for monolayer deposition differs

from the thermodynamic potential for bulk deposition should correlate with the difference

between the temperature for flash desorption of bulk deposited layers and that for the

last layers of adsorbate to vaporize, and indeed such a correspondence is observed for

metals on metals.

Optical techniques, such as internal reflection infrared or Raman spectroscopy,

offer the hope of probing in a fundamental way the nature of surface species in situ in the electrochemical environment. A cautionary note must be sounded, however. Even such

a seemingly simple system as CO adsorbed from the gas phase on Ni surfaces has yet to be

unscrambled in spite of work over three decades using highly sophisticated optical

techniques.

Characterization of nonmetallic electrocatalytic surfaces and materials may require

measurements not always used in electrochemical studies. An example arises from the work

reported in this Workshop by Morrison and Freund, pointing out the role dislocations may

play in determining the electrochemical behavior. Characterization with respect to dislocations and similar lattice defects may be very important in some cases. Similarly,

the discussion brought out that the actual composition of WC electrocatalysts may be

quite complex, involving perhaps excess W, substitutional oxygen, or dissolved gases, and

that these compositional anomalies may be very significant in determining electrochemical

behavior. The paper by Tseung and Jasem brought out a possible correlation between

electrocatalytic activity and magnetic susceptibility in some of the transition metal

oxides, opening up the possibility of using magnetic measurements as a tool of characterization.

From another standpoint, very sensitive analytical chemistry techniques for the

quantitative assessment of very low concentrations of reactants and products can be

valuable in sorting out actual reaction mechanisms at various potentials. Such chemical

studies in situ during the course of electrochemical reactions on the surface can help

360
elucidate the role of impurities. Because of the variety of oxygen species possible,

they may be particularly valuable for studying oxygen reduction.

In situ studies under electrochemical conditions are particularly important for oxide

semiconductors or metallic compounds. These materials are apt to be relatively strong

Lewis acids or bases and, as mentioned above, to interact strongly with H2O. Furthermore,

as often demonstrated {e.g., the paper by Weber and Shanks on sodium tungsten bronzes in this Proceedings), the nature of the surface is influenced by the electrochemical reactions

being catalyzed. Extrapolation from the bulk properties to the surface is seldom possible,

and the surface must be studied under dynamic conditions to understand how it influences

the electrochemical reactions.

Influence of the Support on the Catalytic Activity of the Electrocatalyst, and Vice Versa

The paper by Stonehart reviewed some of the evidence for the so-called "spill-over

effect," in which the support appears to take a more active catalytic role with the

catalyst present than it is capable of in the absence of the catalyst. Several mechanisms

for this effect appear possible. Examples are Pt on graphite and the earlier work of

Hobbs and Tseung on Pt on WO3.

One possibility is that the support acts as a cocatalyst. In this mode, the support

possesses some catalytic activity that operates in a way synergistic to that of the

catalyst itself. By shifting concentrations of reactants or products critical to the main

reactions, the support may be able to affect the overall reaction rates in a way not possible in the absence of the supported catalyst. A possible example of this mode is the

enhancement of the oxygen reduction activity of graphite by the addition of phthalocyanines

on the surface. One suggestion is that the role of the phthalocyanine may be to decompose

H2O2 produced by the reduction of oxygen on the graphite surface, but note evidence to the

contrary in the paper of Appleby and Savy. This catalytic activity for oxygen reduction

is observed when the phthalocyanine is added to a graphite surface, but not when it is

; deposited on surfaces of NiO or Au. The presence of the phthalocyanine also enhances the

life of the graphite, which again may be due to the removal of H2O2.

A second mechanism possible for this apparent synergistic effect between catalyst and

support is the modification of the catalyst properties by the support. For instance, it

is well known that the different crystallographic faces of a given substance possess

different catalytic activities. Deposition of the catalyst on a given support may result

in a preponderance of particularly favorable crystal faces exposed to the electrolyte and

therefore catalytic activity enhanced over that of the same catalyst on other supports or unsupported. Similarly, favorable structural defects may be stabilized by deposition of

' the catalyst on a given substrate.

A third mechanism that might in some cases explain an apparent influence of the

substrate on the catalyst is the storage or enhanced dissemination of reaction products.

By reducing the back reactions, the rate of the overall catalyzed reaction can be

increased.

Finally, there may be some similarity between the interaction of, e.g., Pt and

graphite and that of Pb adsorbed on Pt. The paper by Adzic and colleagues shows that in the latter case, the adsorption of H on the Pt is inhibited, and reactions such as the

oxidation of formic acid, with which adsorbed H interferes, are enhanced.

Photoelectrolysis

Since the work of Fujishima and Honda (Bull. Chem. Soc, Japan, 44, 1148(1971 );

• Nature, 238 , 37 (1972)) demonstrating direct decomposition of water in a photocell, there

I has been interest in the possibility of energy conversion using the electrochemical

activity induced by the absorption of light at the electrocatalytic surface. Scientifically,

there is interest because the phenomena appear to be amenable to treatment by rather wellestablished solid state concepts and it appears to be a reasonable testing ground for the

applicability of such theory to electrochemical reactions. As a practical source of

hydrogen or a practical means of converting solar to electrical or chemical energy, it is still uncertain. Too much energy is wasted with such a wide band gap semiconductor as

361
Ti02 for practical application, but other, better materials may be found.

"Engineered" Surfaces

A concept that is being actively explored is to attach molecular compounds of known

catalytic activity to substrates capable of acting not only as supports but also as electronic connectors with external circuitry. Not only must the substrate be a good

conductor, but the fixing of specific functional groups must allow rapid electron transfer

to the substrate. Graphite has been a favorite substrate to which to fix various polycyclic compounds with controlled functional groups, particularly in a search for oxygen

reduction catalysts. One apparent problem is to maintain stability in the presence of the

reactive species formed during the electrochemical reactions. A three-dimensional structure

with self-renewing surface might work.

Catalysts for the direct oxidation of hydrocarbons and for the reduction of oxygen to water rather than hydrogen peroxide might be sought in this way. The suggestion was made

that the splitting of the oxygen-oxygen bond needed for reduction to water might go well

with coordination compounds involving low-valence states of second and third row transition

elements, e.g., Ir"*". In such studies, the mode of adsorption should be determined,

whether it involves isolated molecules, clusters, or monolayers of adsorbate. By observing

the change in activity of the adsorbate in going from the solution to the adsorbed state, one might obtain some information on the mutual interaction between the adsorbate and

substrate.

Interfaces Involving Solid Electrolytes

In a very general way, electron exchange of reactants/products with the electrodes

and ion exchange with the electrolyte appear to be very rapid, and early fears of large

activation barriers, particularly with respect to the ion exchange with the electrolyte,

appear to be unfounded. Indeed, examples of slow processes of this kind are hard to find.

The electroreduction of CO2 with Pt on a stabilized zirconia electrolyte is one possibility. The paper by Farrington in this Proceedings suggests that the transfer of Na+ between 3-AI2O3 and propylene carbonate when H2O (or other protonated species) is present

may be another.

Most theories of the development of electrode polarization involve the concept of the

blocking electrode. Like the perfect insulator, the completely blocking electrode is an idealization, not really reproduced by any real electrode. A very useful form taken under

simplified conditions by the theories of electrode polarization is the equivalent circuit,

and it must be remembered that these are idealizations so that complete correspondence

with experimental data must not be expected.

As work on solid electrolytes and their interfaces with various electrodes has

proceeded, it has been found that the concepts and instrumental methods of aqueous electrochemistry have been very useful.

362
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4. TITLE AND SUBTITLE PROCEEDINGS OF WORKSHOP ON ELECTROCATALYS I S ON NON-METALLIC

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5. Publication Date November 1976

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16. .ABSTR.ACT (A 200-word or less (actual summary of most significant information. If document includes a significant bibliography or literature survey, mention it here.)

This book presents the proceedings of a Workshop on E 1 ect rocata 1 ys i s on Nonmetallic Surfaces, held at the National Bureau of Standards, Ga i thersburg , Maryland,

on December 9-12, 1975. The Workshop was sponsored by the Institute of Materials

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