NBS hiNI

453

Electrochemical

Analysis Section

Summary of Activities

July 1967 to June 1968

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U.S. DEPARTMENT OF COMMERCE

National Bureau of Standards
NATIONAL BUREAU OF STANDARDS

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UNITED STATES DEPARTMENT OF COMMERCE

C. R. Smith, Secretary

NATIONAL BUREAU OF STANDARDS • A. V. Astin, Director

NBS TECHNICAL NOTE 453

ISSUED JULY 1968

Electrochemical

Analysis Section:

Summary of Activities

July 1967 to June 1968

Edited by Roger G. Bates

Electrochemical Analysis Section

Analytical Chemistry Division

Institute for Materials Research

NBS Technical Notes are designed to supplement the Bureau's regular publications program. They provide a means for making available scientific data that are of transient or limited interest. Technical Notes may be

listed or referred to in the open literature.

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

Price 55 cents.
FOREWORD

The Analytical Chemistry Division was established as

a separate division at the National Bureau of Standards on

September 1, 1963* and became part of the Institute for

Materials Research in the February 1, 1964, reorganization.

It consists at present of nine sections and about 100 technical personnel encompassing some 57 different analytical

competences from activation analysis and atomic absorption

to vacuum fusion and x-ray spectroscopy. These competences,

and in turn the sections which they comprise, are charged

with research at the forefront of analysis as well as awareness of the practical sample, be it standard reference

material or service analysis. In addition it is their

responsibility to inform others of their efforts.

Formal publication in scientific periodicals is a highly

important output of our laboratories. In addition, however,

it has been our experience that informal, annual summaries

of progress describing efforts of the past year can be very

valuable in disseminating information about our programs.

A word is perhaps in order about the philosophy of these

yearly progress reports. In any research program a large

amount of information is obtained and techniques developed

which never find their way into the literature. This

includes the "negative results" which are so disappointing

and unspectacular but which can often save others considerable work. Of importance also are the numerous small items

which are often explored in a few days and which are not

important enough to warrant publication—yet can be of great

interest and use to specialists in a given area. Finally

there are the experimental techniques and procedures, the

designs and modifications of equipment, etc., which often

must be covered in only a line or two of a Journal article.

ii
Thus our progress reports endeavor to present this

information which we have struggled to obtain and which we

feel might be of some help to others. Certain areas which

it appears will not be treated fully in regular publications

are considered in some detail here. Other results which are

being written up for publication in the journal literature

are covered in a much more abbreviated form.

At the National Bureau of Standards publications such

as these fit logically into the category of a Technical Note

In 1968 we plan to issue these summaries for all of our

sections. The following is the fourth annual report on

progress of the Electrochemical Analysis Section.

W. Wayne Meinke, Chief

Analytical Chemistry Division

iii
PREFACE

This is the fourth annual progress report of the

Electrochemical Analysis Section of the Analytical Chemistry

Division. It covers the fiscal year 1968, which began on

July 1, 1967, and ended on June J>0, 1968.

The Section concerns itself particularly with ionic

processes occurring in solution, with areas of analytical

measurement where ionic equilibria play a part, and with

the explanation of solution behavior in terms of the interactions of ionic solutes with solvent molecules. The

special interests of the members of the Section lie for the

most part in solution electrochemistry, electrochemical

thermodynamics, and ionic transport phenomena. In the

context of modern analytical chemistry, their competences

bear most directly on potentiometry, electrometric titrations, and conductometric analysis. In line with a uniform

policy of the Division, the Section's programs have both

research and sample aspects. .During the fiscal year just

ending, the total effort of the Section was divided almost

equally between research and the development of standard

reference materials.

The permanent staff of the Section was further reduced

during the year with the resignations in July of Hannah B.

Hetzer and Dr. R. A. Robinson. Dr. Robinson had served as

assistant chief of the Section for six years, and with his

departure the project on isopiestic vapor pressure measurements was terminated. During the summer months of 1967,

however, the staff of the Section was augmented by the

presence of Dr. Karl H. Pool of the Department of Chemistry,

Washington State University at Pullman, and of Edgar S. Etz,

a graduate student at Clarkson College of Technology.

iv
Dr. Donald Rosenthal, professor of chemistry at Clarkson,

was named a consultant to the Section. During their stay,

Dr. Pool and Mr. Etz made material contributions to our

studies of acid-base phenomena In mixed solvents. The

arrival In October of Dr. Marinus Alfenaar of the University

of Utrecht, Holland, to fill a one-year appointment, enabled

us to initiate a project on standards for ion-selective

electrodes. This work was continued throughout the fiscal

year with the welcome financial support of the Foxboro

Company, of Foxboro, Massachusetts.

With a permanent staff of only four, it has obviously

been impossible to conduct a broad program in the electrochemistry of solutions. For this reason, the activity of

the past year has been limited to emf studies and conductance measurements. Unlike the formal publications that

appear in the scientific journals, this report summarizes

the total program of the Section, incomplete studies as well

as finished work, and stresses its relationship to the missions of the Division and Institute.

In order to specify adequately the procedures, it has

been necessary occasionally to identify commercial materials

and equipment in this report. 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 identified is necessarily the best available for

the purpose.

The assistance of Mrs. Rosemary Maddock, Mrs. Marguerite

Raudenbush, and Robert J. Boreni in the preparation of this

report is gratefully acknowledged.

Roger G. Bates, Chief

Electrochemical Analysis Section

Washington, D.C.

June 30, 1968
TABLE OP CONTENTS

PAGE

1. FACILITIES AND EQUIPMENT ..... 1

2. ACIDITY MEASUREMENTS , v **'***>> y* ' 4

Concepts and Program ....... 4

Surface Potentials and Single Ion Activities . 5

Objectives . 5

Instrumentation ........ 7

New pH Standards 8

Potassium DIhydrogen Citrate 9

Carbonate Buffer Solution ... 10

Internal Consistency 14

pD Standardization 15

Response of the Glass Electrode

in Deuterium Oxide 15

Standards for pD Measurement 15

Potassium Dihydrogen Citrate .... 16

Carbonate Buffer 17

Internal Consistency of th«.

pD Scale ..>... 19

Certification of Standard Reference

Materials ..... 22

Measurement of Acidity In Sea Water 24

3. SOLVENT EFFECTS ON ELECTROLYTE PROCESSES ... 25

Ion-Solvent Interactions ..... 25

Water Structure and Medium Effects in

Mixed Aqueous Solvents ..... 26

Structural Effects in N-MethylpropionamideWater Mixtures 30

Emf Studies in Methanol-Formamide Mixtures . . 32

Emf Studies In Methanol-Propylene Carbonate

Mixtures 34

Emf Studies in N-Methylproplonamide-Hydrocarbon

Mixtures 35

vi
PAGE

4. STANDARDIZATION OF ION-SELECTIVE ELECTRODES . . 37

General Principles 37

Definition of Activity Coefficients ... 38

Liquid-Junction Potentials 39

Electrode Response , 40

Measurement of Chloride Ion Activity 4l

Measurement of Sodium Ion Activity 4j5

5. EQUILIBRIUM DATA FOR IONIC SYSTEMS 46

Thermodynamics of Hydrochloric Acid

in 92.41 Wt. Percent Ethanol 46

Second Dissociation Constant of Carbonic Acid

in Deuterium Oxide ............. 50

Properties of Acids and Bases in HydrogenBonded Mixed Solvents 53

6. CONDUCTANCE AND TRANSPORT BEHAVIOR OF

ELECTROLYTES 57

Oceanographic Conductance Standards ...... 57

Apparatus 57

Characteristics of "Azulene Periodate" .... 60

Hydrodynamic Properties of Ions ........ 6l

7. PREPARATION AND PROPERTIES OF SOLVENTS .... 64

Propylene Carbonate 64

Cryoscopic Behavior of N-Methylpropionamide . . 64

8. FOREIGN- CURRENCY PROGRAMS 66

9. REVIEWS, CONFERENCES, AND COMMITTEE ACTIVITY. . 67

10. IMPACT OF SECTION'S ACTIVITIES 68

11. PERSONNEL 69

12. COMMITTEE MEMBERSHIPS 70

13. PUBLICATIONS AND MANUSCRIPTS,

JULY 1967 TO JUNE 1968 71

14. TALKS, JULY I967 TO JUNE I968 73

15. LIST OF REFERENCES ,...,,, 75

vii
LIST OF FIGURES

FIGURE NO. PAGE

1. Laboratory for the certification of

pH standards. ....... 2

2. Equipment for dielectric constant

measurements 2

3. Cell for the measurement of the surface

potentials of electrolyte solutions ... f

4. pH(s) as a function of temperature. .... 12

^. Values of pD(S) for two primary standards in

deuterium oxide as a function of

temperature ........ 20

6. Solvent effect on the dissociation of

p-nitroanilinium ion in methanol-water

solvents compared with that in methanolpropylene glycol solvents . 28

7. Molar volume of transfer from pure organic

liquid to aqueous solution at j50 °C . . . j50

8. Pseudo-Kirkwood correlation parameters for

binary mixtures of water and organic

JLiquiQs ................. _p2

9. Response of a chloride ion-selective

electrode in solutions of sodium chloride

at 25 °C 42

10. Response of a chloride ion-selective

electrode in equimolal mixtures of sodium

chloride and potassium nitrate at 25 °C . kj>

11. Equipment for emf measurements with

ion-selective electrodes. ........ 45

12. Variation of log y+ for hydrochloric acid

with the square root of the molality in

92.41 wt. percent ethanol compared with

that in water and in 50 wt. percent

ethanol .. .......... 49

13* Determination of pK2 for deuterocarbonic

acid at 25 °C ............. . 52

14. Standard thermodynamic functions for the

transfer of hydrochloric acid from water

to ethylene glycol-water solvents

at 25 °C 55

viii
LIST OP FIGURES (con't)

FIGURE NO. PAGE

15. Standard thermodynamic functions for the

transfer of acetic acid from water to

ethanol-water solvents at 25 °C 56

16. Janz-Mclntyre conductivity bridge 58

17. Demountable conductivity cells ...... 59

LIST OF TABLES

TABLE NO. PAGE

1. Standard reference values of pH(S) for

0.05m potassium dihydrogen citrate ... 10

2. Standard reference values of pH(s) for

the solution 0.025 molal with respect

to both sodium bicarbonate and sodium

carbonate 11

5. Homogeneity tests on carbonate salts.

Acidity functions for solutions of the

composition: NaHCO^C 0.025m),

Na2CO;3( 0.025m), KCl(O.Olm) at 25 °C . . . 13

4. Inter comparison of new reference buffers

with the phosphate buffer standard

from 5 to 45 °C .... 14

5. Standard reference values of pa-p for 0.05m

KI>2 citrate in deuterium oxide 17

6. Standard reference values of pa-p for the

solution 0.025 molal with respect to

both NaDCO-2 and Na2C0^ in deuterium

oxide 19

7. Intercomparison of reference buffers in

deuterium oxide at 25 °C measured against

the phosphate buffer standard

[pD(s) = 7.429] . 21

8. Emf of the cell: Pt;H2

, HCl(0.02m) in

solvent a or b, AgCljAg at 25 °C

(in volts) 34

9. Experimental determination of -log 7^=+

in solutions of sodium salts at different

ionic strengths from measurements with a. sodium- selective NAS 11-18 glass electrode

at 25 °C 44

ix
LIST OF TABLES (con't)

TABLE NO* PAGE

10. Determination of the standard emf ( E°)

for the cell: Pt;H2, HCl(m) in s 92.41 wt. % ethanol, AgCljAg at

5, 25, and 50 °C . . . . 48

11. Comparison of L2 and J^ for hydrochloric

acid (0.1 molal) in water and in 92.41

wt. percent ethanol at 25 °C 50

12. Dissociation of the deuterocarbonate ion

(DCO-x") in deuterium oxide from

5 to 50 °C. . . . 53

13. Freezing point and cryoscopic constant of

N-methylpropionamide .......... 65

x
ELECTROCHEMICAL ANALYSIS SECTION: SUMMARY OF ACTIVITIES,

JULY 1967 TO JUNE 1968

Edited by Roger G. Bates

ABSTRACT

This survey of the activities of the Electrochemical

Analysis Section, Analytical Chemistry Division, covers

the period July 1967 to June 1968. An attempt is made to

summarize a year's progress on the technical projects of

the Section in such a way as to stress the program and

capabilities of the organizational unit as a whole. Brief

summaries of several lines of work under way are given.

The review describes briefly new equipment and facilities

acquired during the year. A major effort devoted to acidity

measurements has led to new pH standards and new reference

solutions for pD in heavy water. Solvent effects on acidbase phenomena in mixed solvents have shed some light on the

factors, including solvent structure, governing ion- solvent

interactions. A start was made on the study of ion- specific

electrodes and the establishment of standard scales for

ionic activity. Equilibrium data for hydrochloric acid in

92.41 wt. percent ethanol and for the ionization of carbonic

acid In heavy water have been obtained. The establishment

of conductance standards for oceanographic research was

begun. The survey concludes with lists of the members of

the Section staff, publications and manuscripts produced

during the year, talks given by the staff, and committee

assignments.

Key Words

Acidity, Analysis, Conductance measurements, Electrochemical

analysis, Electrochemistry of solutions, Ionic activity,

Ion-selective electrodes, pD measurements, pH measurements,

Medium effects, Solvent effects

xi

1. FACILITIES AND EQUIPMENT

Excellent facilities for the Section's programs are

provided on the second floor of the Chemistry Building at

the Gaithersburg site of the National Bureau of Standards.

These laboratories have now been occupied for more than

two years, and the Inevitable minor changes and adjustments

necessary to adapt a general- purpose laboratory to our

special needs have been completed.

The space devoted to research consists at present

of a conductance laboratory and four well-equipped emf

laboratories. The Instrumentation room has been outfitted

to accommodate the recording spectrophotometer, the instrument for measuring dielectric constants of solvents, the

automatic titrator, microbalance, and several pH meters.

One of the emf laboratories is shown in figure 1, and

the dielectric constant instrument is shown in figure 2.

Two items of capital equipment were acquired during

the year. One was a vibrating-reed electrometer needed

for the measurement of the potentials of glass electrodes

and ion-selective electrodes. A commercial impedance

comparator that utilizes closely matched transformer ratio

arms was also purchased. This instrument, which indicates the difference between two impedances with a specified accuracy of ^0 ppm, is especially useful for precise

measurements of the conductance of highly conducting

solutions.

One of the four emf laboratories is now devoted to

studies of aqueous solutions and to certification of pH

standards, although the cells and measuring equipment are

so designed and arranged that studies of nonaqueous and

mixed solvents can be conducted when time permits. A second

emf laboratory was devoted almost exclusively to work in
Figure 1. Laboratory for the certification of pH standards.

Figure 2. Equipment for dielectric constant measurements
deuterium oxide with deuterium gas electrodes and glass

electrodes during the current year. Investigations of

ion-selective electrodes were conducted in a third, while

a fourth, in use only during the summer months, was given

over to studies of acid-base processes in nonaqueous mixed

solvents.

This installation is sufficiently large to warrant

the inclusion of an automatic data read-out center capable

of recording automatically and simultaneously the emf of

all the cells under study in each of the laboratories.

A multichannel digital voltmeter with an accuracy of 10 |j,V

is needed. It would be coupled to a print-out unit from

which a record of all the emf data could be obtained and,

when desired, a tape could also be produced for computer

analysis. Measurements of all cells at desired intervals

could be programmed in advance.

Because of staff limitations, it has not yet been

possible to initiate work on the accurate measurement of

pH at high temperatures and pressures or to supplement

our present capability with calorimetric studies of ionic

processes in solution. When these types of work are begun,

additional facilities and equipment will be needed.

(R. G. Bates)
2. ACIDITY MEASUREMENTS

A. Concepts and Program

The measurement of acidity continues to be one of the

major concerns of the Section. The concepts and approaches

of our efforts in this area were set forth In the annual

progress reports for 1966 [l] and 19&7 [2].

The NBS pH scale Is fixed by a series of standard reference solutions chosen to cover uniformly that region of

the scale in which the emf of the pH cell Is a linear function of the pH of the solution in which the electrodes are

Immersed. Corrections are made, If necessary, for imperfections in the response of the glass electrode. Unlike

the pH scale of the British Standards Institution [;5]> which

recognizes only a single primary standard, the NBS approach

permits a series of primary reference solutions. The search

for new and better standards is going on continually. A

new standard may be selected because it extends the useful

range of the standard scale or because It shows a stability

superior to those of existing standards In its pH range.

During the past year, work was continued and extended

on the measurement of acidity in deuterium oxide (heavy

water). The pD scale in heavy water bears the same fundamental relationship to the deuterium gas electrode that the

pH scale bears to the hydrogen gas electrode. Basically,

the phenomena which are regarded as a manifestation of

"acidity" in these two media can be attributed to two distinctly different entities, namely the deuterium ion and

the hydrogen Ion. For this reason, the pD scale Is not

related to the pH scale. The same methods can nonetheless

be utilized for establishing acidity standards in deuterium

oxide as have been used to set up pH standards In ordinary

water. During the past year an operational pD scale fixed

by three standard reference solutions has been established.
In amphiprotic (water-like) solvents of moderate to

high dielectric constant it is useful to correlate acidbase behavior with the concentration or activity of the

acidic elements hydrogen ion and deuterium ion. This

approach is a worthwhile one not only for water but also

for aqueous mixed solvents and for some nonaqueous media.

The extension of the pH concept to media other than pure

water has received considerable attention in the Section

during the past few years. At the time of issuance of the

new pH and pD standards, a review article on this subject

was published [4],

A truly general scale of acidity as envisioned by

Brjzfasted [5] will become a reality only when an acceptable

scale of ionic medium effects or ionic transfer free energies has been determined or defined by convention [6].

Although the medium effect for a single ionic species cannot

be derived by thermodynamic means, there is a growing belief

that it may often be a real and useful physical concept.

A valid procedure for separating medium effects for electrically neutral combinations of ions into their cationic

and anionic constituent parts is still being sought. These

efforts continue to claim the interest and attention of

members of the Section.

(R. G. Bates)

B. Surface Potentials and Single Ion Activities

1. Objectives

It has been shown that single ion activities play an

important role in the establishment of useful pH scales [7].

Nevertheless, single ion activities cannot be measured.

They have an arbitrary meaning derived mainly from a convention which regards the Inner potential (the Galvani

potential tp ) as a macroscopic property of the medium.
Thus, no net electrostatic work is performed when an

electrically neutral combination of ions is removed from

the medium, and the total work can be defined as the sum

of the chemical potentials of the ions.

The activity of an ion i is formally related to the

chemical potential p.. by

RT In a

±

« \jl±

- |ij (1)

where p,f is the chemical potential in a selected standard

reference state. Heretofore, all attempts to estimate p,. for single ionic species have been based on an arbitrary

splitting of chemical potentials for electrically neutral

combinations of ions or on similar nonthermodynamic procedures. In principle, however, it is possible to measure

the electrochemical potential jr. = p.. + P^» of a single

ionic species. Then if <p is known or can be estimated, one

can obtain an estimate of p.. . When v is small, a rather

inaccurate estimate of its value is sufficient to obtain

relatively accurate values of p...

It is the purpose of the present study to obtain more

information on the Galvani potential <p , For this, measurements will be made on cells of the type

AgjAgCl, MCl(m

1

) in H

2

| air | MCl(m

2

) in H

2

0, AgCljAg

with different cations M

+

and different ionic strengths.

Starting at very low concentrations, where it can be assumed

that single ion activities for different univalent ions do

not differ appreciably, it is hoped to find some relation

to describe the surface potential at the solution-gas

interface. This relation, we hope, can be extended to

higher concentrations so that an estimate of p.. in that

range will also become possible.
2. Instrumentation

For the solution of this problem, emf measurements

have to be performed with high precision (0.1 m"V or better)

on cells with very high Impedance. For this purpose, we

have at our disposal a vibrating-reed electrometer. It Is

expected that the measurements to be performed will fall

just within the capability of this instrument. The cell

that has been constructed for this study is shown in

figure 3« The design is that of Randies [8].

Figure j5. Cell for the measurement of the surface

potentials of electrolyte solutions.

(M. Alfenaar)
C. New pH Standards

During the past year, research on two pH standard

solutions proposed as new reference points on the NBS

standard pH scale has been completed. The solutions are

0.05 molal potassium dihydrogen citrate and a mixture of

sodium bicarbonate and sodium carbonate, each 0.025 molal.

The precise emf measurements leading to the assignment of pH(s) values to these reference solutions over the

temperature range to 50 °C were completed. The procedures

used for this assignment have already been adequately described in previous reports [1,]]. An over-all view of

the equipment used is given in figure 1.

A few noteworthy alterations of the usual procedure

were made. The calculation of the ionic strengths of the

buffer solutions, for example, was adapted to an iterative

computer technique which utilized all of the appropriate

dissociation constants for the acidic groups governing the

pH of the buffer, namely the three dissociation steps of

citric acid and the two steps of carbonic acid.

Computer techniques were applied extensively in the

analysis of the emf data. In fact, by the use of several

computer programs it is now possible to start with the

measured emf and the corresponding barometric pressure and

arrive at the pa^, value of the solution entirely by computer. The procedure enables one not only to apply precise

corrections for barometric pressure, etc., to each individual data point but also minimizes human computational

errors. Each of the three applicable programs produces a

tape of results which can then be fed directly into the

subsequent program or modified to include only selected

data, if desired.

8
1. Potassium Dlhydrogen Citrate

Inasmuch as each of the present standards which fix

the lower end of the pH scale has certain disadvantages

under various conditions, potassium dlhydrogen citrate was

chosen as an addition to the list of primary pH standards.

The 0.05 molal solution has a. pH(S) of 3.776 at 25 °C.

Two of the present reference solutions have pH(s)

values in the same region as does this citrate solution.

They are the tartrate and phthalate solutions. However,

the saturated solution of potassium hydrogen tartrate

[pH(S) = 3.557 at 25 °C] is subject to mold growth after a

few days, and the pH may change by as much as 0.1 unit.

Furthermore, the widely-used phthalate reference solution

absorbs strongly in the ultraviolet region of the spectrum;

hence, this solution is of limited application in many

spectrophotometric procedures for the study of dissociation

constants and other acid-base interactions.

The emf measurements on the solution of potassium

dlhydrogen citrate were completed and pH(s) values assigned

by the usual methods. The pH(S) at 11 temperatures from

to 50 °C is listed in table 1. A large lot of the salt

has been ordered and will be tested and certified upon

receipt. It goes without saying that minor variations in

the pH(s) of different lots may occur, but these are not

expected to exceed a few thousandths of a unit.

The standard solution of potassium dlhydrogen citrate

is easily prepared by weighing a single salt. Furthermore,

the citrate salt can be obtained readily in highly pure

form. Its usefulness for spectrophotometric measurements

in the ultraviolet and visible regions and its high stability both in the solid form and in aqueous solutions in

contact with the hydrogen electrode recommend it as an

additional standard reference material for pH measurements.
Table 1. Standard reference values of pH(S) for 0.05m

potassium dihydrogen citrate.

t, °C pH(S) t, °C pH(S)

o 3.865 30 3.766

5 3.840 35 3.759

10 3.820 40 3.753

15 3.802 45 3.750

20 3.788 50 3.749

25 3.776

2. Carbonate Buffer Solution

The addition of the carbonate buffer solution, 0.025 m

in both sodium bicarbonate and sodium carbonate, to the

list of pH standards likewise offers several advantages.

This reference solution extends the pH range beyond pH 10

at 25 °C, into a region where the uncertainties in the

response of glass electrodes often become large. The pH(s)

of the carbonate buffer is 10.012 at 25 °C. The materials

from which the solution is prepared are obtainable in high

purity and are very stable when kept dry. On the other

hand, the borax used for the present standard of highest pH

(9.180 at 25 °C) may lose a part of its water of crystallization during prolonged storage. Furthermore, the ionic

strength of the borax standard is not as well defined as

might be desired, due to a tendency of boric acid to

polymerize in solution. For this reason, the concentration

of borax in the reference solution has to be kept low, with

a consequent sacrifice of buffer capacity. The carbonate

buffer does not suffer from this difficulty.

10
The emf measurements needed to assign conventional

pa^ values to the carbonate standard have been completed.

These values, which are identified with pH(s) in the

operational definition of pH, are listed in table 2. The

variation of pH(S) for the citrate and carbonate standards

with changes of temperature is illustrated in figure 4.

Table 2.

t, °C

Standard reference values of pH(s) for the

solution 0.025 molal with respect to both

sodium bicarbonate and sodium carbonate.

PH(S) t, PH(S)

5

10

15

20

10.517

10.245

10.179

10.118

10.062

10.012

30

35

40

45

50

9.967

9.926

9.889

9.857

9.828

In making the measurements, it was found necessary

to purge the water with C0

?

-free air or nitrogen, or by

boiling, in order to remove carbon dioxide prior to preparing the solutions. It was also necessary to protect the

solutions from atmospheric carbon dioxide after preparation.

In addition, best results were obtained when the flow of

hydrogen through the cells was interrupted during the night,

In order to minimize loss of carbon dioxide from the cell

solutions as a result of the continued bubbling of the

hydrogen over long periods of time.

H
pH VS TEMPERATURE

1

20 30

TEMPERATURE (°C)

Figure 4. pH(s) as a function of temperature.

The assay of the carbonate salts was made by weight

titration with a hydrochloric acid solution that had been

standardized by the gravimetric silver chloride procedure.

The sample was acidified with an excess of the acid, boiled

under reflux, and back-titrated with a. solution of sodium

hydroxide that had been standardized by titration to the

same end point (pH 7) with the standard solution of hydrochloric acid.

12
New lots of sodium "bicarbonate and sodium carbonate

have been acquired. They have been assayed and tested for

homogeneity and, having been found acceptable, were certified for issuance as Standard Reference Materials for pH

measurement. The homogeneity of the lots was demonstrated

by a comparison of the acidity functions p(a„Yr,.

] ) for solutions of the composition; NaHC0^( 0.025m), Na2

C 0,(0, 025m),

KCl( 0.01m), prepared with different carbonate samples.

The salts used were removed from the five different containers in which each of the salts was supplied* In the

first run of six cells, a. composite sample of sodium

carbonate was used throughout, with six different samples

of the bicarbonate, including a composite sample. Five of

the samples of the bicarbonate were withdrawn from the five

different containers; equal portions of all five were mixed

to form the composite sample. In the second run, the composite sample of bicarbonate was used in each of the six

cells, together with six different samples of sodium carbonate (including the composite sample). The results of

the homogeneity tests are shown in table 3«

Table 3. Homogeneity tests on carbonate salts. Acidity

functions for solutions of the composition:

NaHCO-J 0.025m), Na2

C0,( 0.025m), KCl(O.Olm) at 25

'

Sample No. p(aHycl ) p(aHYcl ) Run 1 Run 2

1 10.117 10.117

2 10.121 10.116

3 10.117 10.116

4 10.117 10.118

5 10.117 10.115

6 10.118

13
3. Internal Consistency

The several primary standards that fix the NBS standard

pH scale must display a satisfactory consistency when intercompared in the pH meter with liquid junction. Otherwise,

experimental pH values for "unknown" solutions will depend

on the reference solution chosen to standardize the measuring instrument. No exact consistency can be expected,

however, when a liquid- junction potential is involved, as

the standard solutions have different ionic strengths and

necessarily contain different types of ions whose mobilities

are not the same.

The new citrate and carbonate standards were compared

with the equimolal phosphate standard buffer solution

[pH(s) = 6.865 at 25 °C] in a cell in which two hydrogen

electrode compartments were separated by a bridge of saturated potassium chloride solution [9]- The results given

in table 4 show that the pH indicated by the cell with

liquid junction is higher than the assigned pH(S) for the

citrate solution and lower than that assigned to the carbonate solution. The differences are less than 0.011 unit

between 5 and 45 °C.

Table 4. Inter comparison of new reference buffers with the

phosphate buffer standard from 5 to 45 °C.

t, °C pH(obs.)-pH(S) pH(obs.)-pH(S)

Citrate soln. Carbonate soln.

5 0.011

10 0.010 -0.011

25 0.007 -0.010

45 0.008 -0.007

(B. R. Staples)

14
D. pD Standardization

1. Response of the Glass Electrode In Deuterium Oxide

The studies of the response of the glass electrode

in heavy water described in the last report have been published during the year [10], giving firm evidence that this

electrode, when suitably conditioned, responds to deuterium

ions in heavy water as well as to hydrogen ions in ordinary

water. If one uses a glass-calomel pH assembly standardized

with a reference solution made up in ordinary water, an

empirical correction factor of 0.45 unit (molal scale) for

the pD range 2 to 9 must be added to the pH meter reading

to convert this "operational pH" to a pD value.

This correction need not be applied if the pH meter is

standardized with a reference buffer solution in heavy water

to which a pD(S) value has been assigned. Measurements with

nine buffer solutions in heavy water indicate that no correction factor is necessary when a buffer mixture containing

KD

2

P0^ and Na^POu (each 0.025 molal) in heavy water is

used to standardize the measuring equipment. The differences between the operational pD values and the assigned

pa^ values of these nine buffer solutions were negligible.

2. Standards for pD Measurement

As the result of earlier work in the Section, pD(S)

values were assigned to two solutions suitable as reference

buffer solutions for pD measurements in heavy water [11,12].

The two buffer solutions were l) a mixture of potassium

dideuterium phosphate and disodium deuterium phosphate

(each 0.025m) and 2) a mixture of acetic acid and sodium

acetate (each 0.05m). These solutions have pD(s) values

of 7.429 and 5.230, respectively, at 25 °C. The proper

behavior of glass electrodes in heavy water has indeed been

affirmed, but, since the response of individual electrodes

is not constant over the entire useful pD range, two new

additional primary standards have been proposed to fix the

15
practical pD scale. The new buffer materials are potassium

dideuterium citrate ( 0.05m) and a. mixture of sodium deuterium carbonate and sodium carbonate (each 0.025m).

a. Potassium Dihydrogen Citrate

This particular buffer material was selected as one

of the new primary reference standards for two reasons:

first, because of the low pD value of its 0.05m solution

(4.29 at 25 °C), and second, because of the relative ease

with which its solutions can be prepared. Only one solid

material has to be weighed, thus making this buffer preferable to the acetate buffer as a reference solution.

The preparation and purification of potassium dihydrogen citrate were described in the last report [2].

The standard values of pa~ were determined by a procedure

which follows closely the method for establishing primary

pH standards in ordinary water. The emf of the cell

Pt;D2 (g), KD2 Cit (0.05m), KCl(m 2) In D2

0, AgCljAg

was measured for buffer solutions containing three different

molalities of potassium chloride (nu = 0.005, 0.01, and

— l 0.015 niol kg ) over the temperature range 5 to 50 °C.

The KDpCit salt was formed in solution by exchange with

DpO when KHUCit was dissolved in heavy water.

The measured emf (e) of this cell, together with the

standard emf (E°), was used to calculate the acidity function p(aDYc

, 1

). The limiting value of the latter P( ar) Yp-|)

at zero concentration of added chloride was obtained by

least- squares extrapolation.

Conventional pa^ values were computed by the equation

paD

= p(aDycl

)° + log <ycl

- (2)

where the single ionic activity coefficient for the chloride

ion was defined by the Bates -Guggenheim convention [13] in

16
the form

AIl/2

- lQs yci- = rz

^

H4.565BX '

In equation 3, A and B are the Debye-Huckel constants

adjusted to consistency with the dielectric constant and

density of heavy water. The ionic strength I was estimated

from the three ionization constants of citric acid [14,15]

corrected for isotope effects. The pa^ values obtained by

this method were identified with pD(S) values for these

reference buffer solutions j they are listed in table 5«

Table 5. • Standard reference values of pa-* for 0.05m

KDp citrate in deuterium oxide.

t, °C paD

t, °C paD

5 4.378 30 4.279

10 4.352 35 4.268

15 4.329 40 4.260

20 4.310 45 4.254

25 4.293 50 4.250

b. Carbonate Buffer

The second solution proposed as a new primary reference

standard is a mixture of two components, sodium bicarbonate

and sodium carbonate, each 0.025m. This buffer mixture has

a pD value of 10.74 at 25 °C, thus permitting the behavior

of the glass electrode to be checked at relatively high

alkalinities.

17
The sodium bicarbonate and the anhydrous sodium

carbonate were commercial samples. The first compound

was used without any further treatment. It assayed 99-98

percent (standard deviation 0.03) when neutralized with

excess standard hydrochloric acid and back-titrated with

standard sodium hydroxide solution after boiling to expel

the carbon dioxide formed. Anhydrous sodium carbonate was

heated at 275 °C under an atmosphere of carbon dioxide.

It assayed 100. 0^ percent (standard deviation 0.02) when

titrated by the same procedure. Coulometric titrations

gave an assay of 99.9^9 percent (standard deviation 0.004)

and 99.965 percent (standard deviation 0.02), respectively,

for the two salts. The coulometric titrations were performed by C. E. Champion of the Micro chemical Analysis

Section.

The standard values of pa~ were determined by the

method outlined in the previous section. The emf of the

following cell was measured:

PtjD

2 (g)> NaDC 0^(0. 025m), Na

2C0( 0.025m), Nad(m ), AgCljAg

where m, was 0.005, 0.01, or 0.015 mol kg" . Precautions

were taken to avoid contamination with carbon dioxide

by preparing the cell solutions under an atmosphere of

nitrogen. The psu values obtained are listed in table 6.

Equations for the variation of pa- with temperature were

found with the aid of the CEIR computer, using Dartmouth

BASIC Language, and checked by 0MNITAB. The smoothed

curves of pa^ v_s. temperature are shown in figure 5.

18
Table 6. Standard reference values of pa- for the

solution 0.025 molal with respect to both

NaDCO-, and NaoC0-, in deuterium oxide.

3 2 3

t, °C pap t, aC paD

5 10.998 30 IO.685

10 10.923 35 IO.638

15 10.855 40 10.597

20 10.793 45 10.559

25 10.736 50 IO.527

c. Internal Consistency of the pD Scale

The internal consistency of the pD scale fixed by

these three buffer solutions was examined by the intercomparison of standards in cells of the types

PtjD2

(g), Standard soln. (S) | KC1 (satd.), calomel

reference electrode

and

PtjD 2

, Unknown soln. (x) | KC1 (satd.),, calomel reference

electrode

where the vertical line represents a liquid junction.

The operational pD of the unknown solution is given by the

following equation

E-jr-Eo

pD(X) = PD(S)+- — W

(RT In 10)/F

where pD(s) is the assigned value of the standard reference

19
solution and E

x

and Eg are the emf values for the cells

containing unknown and standard solutions, respectively.

If the reference electrode potential and the liquid- junction

potential remain constant when the unknown solution is

substituted for the standard solution, pD(x) will fall on

the standard scale of pa^ on which pD(s) is based.

i r^

0.025m NC1DCO3,

0.025m Na 2 C0 3

in D2

Figure 5. Values of pD(S) for two primary standards

in deuterium oxide as a function of

temperature

20
The 0.02$m phosphate buffer solution was selected as

the standard solution (s) and the citrate and carbonate

buffer solutions (unknown soln., X) were compared with It.

The cell vessel used for the comparison permitted two

capillary liquid junctions to be formed [9]« The experiment was carried out with two types of reference electrode,

one a calomel electrode prepared in heavy water and the

other a calomel electrode in ordinary water. The results

are given in table 7« They indicate that it makes little

difference whether the reference electrode and salt bridge

are prepared in ordinary water or in heavy water; the

residual liquid- junction potential is the same, within

experimental error, as long as both the standard and unknown solutions are made up in heavy water.

Table J. Intercomparison of reference buffers in

deuterium oxide at 25 °G measured against the

phosphate buffer standard [pD(S) = 7-429].

Reference pa„

Buffer

D

2 ca

pD

( ob s .

)

lomel

pD~paD

HpO calomel

pD pD-paD

( ob s .

)

Citrate 4.293

Carbonate 10. 736

4.304

10.729

+0.011

-0.007

4.305

10.728

+0.012

-0.008

21
j5. Certification of Standard Reference Materials

Pour new lots of buffer materials to be issued as pD

standards have been acquired. The new Standard Reference

Materials have been given the following numbers:

Potassium dihydrogen phosphate SRM No. 21861

Disodium hydrogen phosphate 2186II

Sodium bicarbonate 2191

Sodium carbonate 2193

Standard reference values of pD(s) were assigned to buffer

solutions prepared from these materials. The method of

assignment was an abbreviated version of the procedure for

determining pa^ values described in the previous sections

of this report.

The phosphate buffer materials were each supplied in

a single container. An estimate of the homogeneity of the

contents of each container was first made. Two individual

samples were removed from each lot and tested for adherence

to the purchase specifications. No difference was found.

The following tests were performed; hydrogen ion test,

ignition test, and acid-base titration.

Only a single lot of each of the carbonate salts was

acquired for issuance as Standard Reference Materials for

use in both heavy water and ordinary water. The tests for

homogeneity and for assay on these materials are described

in Section 2.C of this report. The emf measurements with

the phosphate and carbonate buffer solutions in heavy water

were performed in the usual manner, already described.

One added complication, however, was encountered.

After the values of the acidity function p(a-r)Yc

-i ) had

been obtained for each buffer solution, it was noted that

they were consistently higher by 0.00J to 0.006 pD unit

over the entire range than those obtained earlier in the

22
Section. This was thought to be caused by a difference

in the heavy water used. On examination, the new lot of

heavy water was found to be slightly radioactive (about

0.4 (iCi/cnr). However, no activity beyond that of the

background was observed when two samples of heavy water

used earlier were examined. The radioactivity measurements were performed by Dr. H. L. Prush. The amount of

tritium oxide indicated is very small. It is estimated to

be about 10 mol kg" . Unfortunately, no non-radioactive heavy water was

immediately available. Hence, the emf measurements were

repeated, for one selected ratio of buffer to sodium

chloride, in heavy water that had a relatively low activity

(about 0.14 |j,Ci/cnr). The values of the acidity function

p(a_Ycl ) found were identical with those previously obtained with inactive heavy water, within experimental error

limits. A set of pD(s) values was thus calculated for each

buffer mixture, combining the results reported by Gary

et al. [11] (phosphate) or those given in Section 2.D.2

of this report (carbonate) with the data obtained from

measurements in the mildly radioactive D^O (0.14 |j.Ci/cnr).

On the average, the certified values for the new Standard

Reference Materials differ but slightly from those found in

the extensive work done earlier by Gary et al. and those

given in table 6. The average difference is less than

0.001 unit.

(M. Paabo)

23
E. Measurement of Acidity In Sea Water

This work has as its objective the development of pH

standards for media approximating sea water in composition.

These reference standards would be useful in adjusting pH

equipment used for accurate determinations of hydrogen ion

concentration or activity in oceanographic studies. They

would be made available through the Standard Reference

Materials program.

During the reporting year, electrochemical studies

of the behavior of hydrochloric acid in simulated sea water

(SSW) were continued at Clarkson College of Technology under

a cooperative arrangement with Professor Donald Rosenthal.

The medium is an aqueous mixture of the chlorides of sodium,

magnesium, calcium, strontium, and potassium. It matches

closely the ionic strength and composition of natural sea

water. When completed, the data will yield the standard

potential of the cell: Pt;H" 2

, HC1 in SSW, AgCljAg over

the temperature range 5 to 35 °C

A computer program facilitating the correction of

the observed emf to the standard hydrogen pressure, as

well as the calculation of acidity functions p^a-vy^n ) F^01®- the corrected emf, has been written. New cells and a new

hydrogen supply line were constructed. Difficulties were

experienced in obtaining reproducible data at 45 °C; since

this temperature is of limited oceanographic interest, it

was decided to place an upper limit of 35 °C on the temperature range studied. The planned experiments were

continued but are not yet completed,

(J. T. Clark, Jr.)

24
3. SOLVENT EFFECTS ON ELECTROLYTE PROCESSES

A. Ion-Solvent Interactions

In its research interests, the Section is intimately

concerned with the relationship between the behavior of

electrolytes in solution and the properties of the solvent

medium. It is well known that relatively small changes in

the composition of the solvent often have a very large

effect on the ionic processes taking place. Nevertheless,

there is very little understanding of the mechanisms of

these profound effects and of the complex interactions

which bring them about.

The interactions of ions which lead to the formation

of ion pairs and higher aggregates can be investigated by

a variety of physicochemical techniques, including conductance, potentiometry, and spectrophotometry. The key

to an understanding of solvent effects, however, appears

to lie in the interactions between ions and solvent molecules, and these processes are much less accessible to

detailed study. In a mixed solvent, it may be particularly

difficult to identify the factors responsible for acid-base

behavior. In these situations, the solvent not only furnishes the medium for the reaction of an acid and a base:

in addition, the two types of solvent molecules may interact

differently with the reactants and products of the primary

acid-base process.

In a general way, ion-solvent processes provide a unifying theme for a large part of the research activity

of the Section. Solvent effects on ionic equilibria, and

medium effects for individual chemical species are being

determined. The results may be expected to shed some

light on the nature of the parameters governing the interaction of solute species with solvent molecules. This

subject is also being attacked from the point of view of

25
transport processes such as conductance and diffusion.

Progress in this area, we believe, is an essential prerequisite to advancements in the electrochemistry of nonaqueous solutions. Furthermore, numerical scales of ionic

transfer free energies "based on either experiment or theory

will provide the basis not only for a unified scale of

single electrode potentials but for a single acidity scale

for all solvents in which the proton is the acidic element.

(R. G. Bates)

B. Water Structure and Medium Effects in Mixed Aqueous

Solvents

In an important paper published in 19^8, Braude and

Stern [16] reported the measurement of Hammett acidity

functions for hydrogen chloride in aqueous solvent mixtures

and in mixed solvents containing no water. The differences

between the results in these two groups of solvents, one

containing water and the other completely nonaqueous, were

striking. When the solvent contained water, the acidity of

the hydrochloric acid solutions appeared to decrease as the

mixture was enriched with the organic constituent (ethanol,

acetone, or dioxane) to a minimum value in the vicinity of

a composition of 50 mole percent. Above this composition,

it increased again, and, when the proportion of water in

the solvent mixture approached zero, the apparent acidity

might exceed that of the pure aqueous solution. On the

contrary, the addition of dioxane or acetone to an ethandle solution of hydrogen chloride was almost without effect

on the apparent acidity up to a composition of 50 mole

percent . At higher acetone concentrations, solutions of

hydrochloric acid displayed a gradually increasing acidity.

When dioxane was added above the 50 mole percent level, the

acidity gradually decreased. The rate of decrease accelerated as the last of the ethanol was removed,

26
Evidently the presence of water in binary mixed

solvents has a. rather profound and anomalous effect on

acid-base behavior in these mixtures. Braude and Stern

attributed the contrasting behavior to a quasi-crystalline

tetrahedral structure which conditions the unique properties

of water [see 17]. According to this view, the proton

affinity of the individual water molecules is partly satisfied by hydrogen bonding with neighboring molecules. On

the addition of an organic solvent to the purely aqueous

medium, this structure is gradually broken down, and

hydrogen bonding between neighboring water molecules is

partly replaced by much weaker bonding with organic solvent

molecules. The result is that the effective proton affinity

of the water molecules increases.

As the proportion of organic solvent is raised, the

aqueous solvation shell is gradually dispersed and the

aqueous solvates (protons associated with more than one

molecule of water) are replaced by hydmnium ions H^0+ . It was suggested that this condition is reached at the

minimum in the acidity curve. When still more organic

solvent is added, the hydronium ions are gradually replaced

by organic oxonium ions of the type EtOH

p

+ . The behavior

of water in mixed solvents is thus conditioned by the

unusual structure of water and by its exceptionally high

proton affinity.

For an improved understanding of the structural properties of alcohol-water mixtures we are indebted to a

recent review by Pranks and Ives [18]. These authors have

shown that the thermodynamic excess functions, partial

molar volumes, densities, sound absorption, spectroscopic

properties, and dielectric constants of alcohol-water

mixtures are qualitatively consistent with the views given

above. One conclusion stands out, however. When alcohols

27
and certain other hydrogen-bonding organic liquids are

added to pure water, an initial increase in structure may

occur, due to the enhanced possibilities for hydrogen bonding. This effect persists only to a composition of a few

weight percent of the organic constituent, after which

breakdown of structure ensues.

One of the remarkable consequences of this change in

solvent structure when water is present is the existence of

a. minimum in the pK value for a positively charged acid

(such as ammonium ion and substituted anilinium ions, for

example) as the aqueous solvent is enriched with an organic

constituent. When the solvent mixture contains no water,

this minimum is absent, as shown in figure 6. The pK values

100 80 60 40

Wt. PERCENT METHANOL

Figure 6. Solvent effect on the dissociation of

p-nitroanilinium ion in methanol-water

solvents compared with that in methanolpropylene glycol solvents.

28
plotted in the upper portion of the figure were obtained

recently by Kundu, De, and Das [19].

If the initial effect of adding the organic component

to an aqueous solution is to increase structure, it is

possible that the pK values of an acid of this charge type

may actually increase initially before falling off. It is

uncommon to measure pK between and 10 wt. percent of the

organic constituent, but a careful examination of the best

pK data reveals a flattening of the curve in this region,

that is, a. delayed drop in pK [20-22].

According to the evidence presented by Franks and Ives

[18], _tert-buta.no 1 is particularly effective in bringing

about increased water structure. It seemed of interest,

therefore, to examine the change in pK for a. cation acid in

water- tert -butanol mixtures containing between and 10 wt.

percent of the organic component. The dissociation constant

of m-nitroanilinium ion can be obtained readily by spectrophotometry measurements in solutions of a strong acid;

consequently, this system was chosen for study. Within the

uncertainty of the measurements (about 0.006 in pK), no

increase was found, nor did the addition of tert-butanol in

amounts of ^ to 5 percent produce a significant decrease

in pK. Although these results are not highly conclusive,

they suggest that increases in the structure of an already

highly structured water result in only relatively small

changes in the proton affinity of the solvent medium.

(M. Alfenaar, R. G. Bates)

29
C. Structural Effects In N-Methylproplonamide-Water

Mixtures

Densities and dielectric constants of mixtures of

N-methylpropionamide (NMP) and water were presented in the

annual progress report for 1967 [2], A paper in press

discusses these results in relation to the structure of

solutions by comparison with the systems consisting of water

and each of three other organic liquids j N-methylacetamide

(NMA), ethanol, and acetic acid.

In the essentially aqueous range of compositions, an

informative parameter is the apparent change in volume of

a mole of the organic component on being transferred from

the pure organic liquid to a dilute aqeuous solution. This

change, AV, as shown in figure 7, is negative for all the

Figure J.

0.1 0.2 0.3 0.4 0.5 MOLE FRACTION ORGANIC COMPONENT

Molar volume of transfer from pure organic

liquid to aqueous solution at J>0 °C.

30
systems, corresponding to a contraction in volume for the

solutions, but it shows a distinct minimum for ethanol

and NTVLA. at 10 to 15 mole percent of the organic solute.

The minimum has "been interpreted [18] as resulting from

competing effects of the polar and non-polar ends of the

organic molecule upon the structure of the water. NMP, in

contrast to its homologue NMA, appears more like acetic

acid in showing the greatest contraction in the infinitely

dilute solution. The measurements were not carried to

sufficiently low concentrations of NMP, however, to establish this conclusion reliably.

In the composition range where water is the minor

component, a more revealing structural insight is given by

the dielectric measurements. A pseudo-Kirkwood correlation

parameter, g, was calculated on the basis that the square

of the effective dipole moment and the high-frequency dielectric constant of the binary mixture can be expressed as

linear functions of the mole fractions of the components.

As shown in figure 8, the two amides resemble each other

quite closely. NMP has the higher g value in the pure

liquid, corresponding to greater parallel orientation of

the dipoles in chain-wise association [2^], but the g value

for NMP also decreases more rapidly than that for NMA with

the addition of water. The decrease indicates the effect

of water in breaking up the amide-amide associations. As

has been noted [18], ethanol and water form a mixture

in which g is virtually constant over the whole range of

compositions. The acetic acid system is very similar to

the ethanol system up to a composition of 50 mole percent

acetic acid. With increasing acid content, g decreases,

becoming less than unity when the mixture contains more

than 90 mole percent of acetic acid. The latter fact reflects the anti-parallel orientation of dipoles in acetic

acid dimers.

51
0.2 0.4 0.6 0.8 MOLE FRACTION ORGANIC COMPONENT

Figure 8. Pseudo-Kirkwood correlation parameters for

binary mixtures of water and organic liquids.

(T. B. Hoover)

D. Emf Studies in Methanol-Formamide Mixtures

As pointed out in Section 3.B^ there is good reason to

regard water to be anomalous in its influence on acid-base

behavior. It is therefore instructive to study the solvent

effects produced by nonaqueous binary solvent mixtures.

To preserve the general similarity to mixtures containing

water, one component should have hydrogen-bonding capabilities j the other might well be aprotic. Methanol is useful

32
as the hydrogen-bonding component. Formamide or propylene

carbonate can be added to it to provide mixtures covering

a considerable range of dielectric constants.

Exploratory emf measurements in formamide and in a mixture of formamide and methanol containing 50 wt . percent

of each component were made. Hydrochloric acid was added

to each cell solution in a concentration of about 0.05

mol kg" . The experiments were designed to test the

stability of the solutions and of the hydrogen and silversilver chloride electrodes in these solvent systems.

The observed emf of each solution increased steadily

with time, and the drift was more rapid for the methanolformamide mixture than for the pure formamide solvent.

Over a period of six hours, the emf for the mixed solvent

increased from 0.6271 V to 0.6517 V. The hydrogen flow

was continued overnight, and on the following day the emf

had risen to 0.6429 V. In the same time intervals, the

cell containing pure formamide showed an increase from

0.6592 to 0.6606 V, and further to O.6616 V on the next day.

This instability of the mixed solvent is too great to

permit solvent effects on the properties of strong electrolytes, and on the pK of weak electrolytes, to be studied

with the accuracy desired. Moreover, formamide is difficult

to purify, and its instability demands that it be freshly

distilled before use. Hence, further work with this solvent

was abandoned in the hope that mixtures of methanol and

propylene carbonate would prove more suitable.

(B. R. Staples)

55
E. Emf Studies in Methanol-Propylene Carbonate Mixtures

Preliminary emf measurements were made on solutions

of hydrochloric acid (0.02m) In "both pure propylene carbonate and a mixture containing 50 wt. percent of methanol

and propylene carbonate. The objective was to ascertain

the feasibility of using these solvent systems for potentiometric studies, particularly for the determination of pK

values for comparison with solvent effects found in the

methanol-water system.

The uncorrected values of the emf recorded over a period of several days are listed in table 8. Initial

equilibrium was achieved rapidly. It is evident that the

methanol-propylene carbonate system is much more stable

than the propylene carbonate solution.

Table 8. Emf of the cell: Pt;H

2

, HCl( 0.02m) in solvent

a or b, AgCljAg at 25 °C (in volts).

Days after a b

preparation Propylene carbonate 50 wt. percent MeOHpropylene carbonate

1 0.02390 0.19743

2 0.0167 O.I9665

3 0.0103 0.19563

6 0.011* 0.19534

unstable emf

Propylene carbonate has many advantages over formamide

for studies of the type we wish to perform. Among them are

a greater stability and greater ease of purification (see

Section 7. A). Although the emf is not as constant as might

34
be desired, it seems feasible to use mixtures of methanol

and propylene carbonate in the study of solvent effects

on acid-base processes.

Two series of measurements of hydrochloric acid solutions directed toward a determination of the standard emf

of the cell in 50 wt . percent methanol-propylene carbonate

were carried out over the temperature range to 40 °C.

Further information needed for the evaluation of the results includes vapor pressures of the solutions and die-

.lectric constants of the solvents over the temperature

range studied.

(B. R. Staples)

F. Emf Studies in N-Methylpropionamide-Hydrocarbon Mixtures

Electromotive-force studies in N-methylpropionamide

were begun by D. Rosenthal in the Section and were described

in last year's report [2]. This study has been continued

with further work at NBS between June and September 19^7

and, in cooperation with Professor Rosenthal, at Clarkson

College of Technology during the academic year 1967-68.

An extension of the measurements of the cell

Pt;H2 (g), HCl(m) in nonaqueous solvent, AgCljAg

with N-methylpropionamide (NMP) as solvent was first made,

to complete the molality range 0.005 to 0.1 mol kg at

temperature intervals of 10 °C from 5 to 55 °C. The experimental data have been obtained, and an analysis of the

results is under way. The data will yield the standard emf

of the cell, the mean molal activity coefficient of hydrochloric acid at various round molalities, as well as the

relative partial molal enthalpy and heat capacity for HC1

in this nonaqueous medium. Computer programs have been

written to facilitate the calculations.

35
In new work undertaken during the year, an attempt

is being made to clarify the nature of solvent effects

through the study of "isodielectric" solvents. N-methylpropionamide has a very high dielectric constant (175.6 at

25 °C). Consequently it is possible, by adding inert

nonpolar organic liquids to NMP, to form mixtures having

the same dielectric constant as water at 25 °C (78.5)

As a first approximation, ions dissolved in such isodielectric solvents should be subject to identical electrostatic

effects in each medium; hence, observed solvent effects on

acid-base processes in these solvents might reveal other

ion-solvent interactions that are sometimes obscured by the

large influence of unequal electrostatic forces.

Two nonaqueous mixed solvents with the same dielectric

constant as water are being studied. These are NMP-cyclohexane and NMP-decalin (decahydronaphthalene) . In the

cyclohexane system, emf measurements have been made from 5

to 35 °C on solutions of hydrochloric acid with molalities

from 0.005 to 0.1 mol kg" . Independent determinations of

the appreciable solvent vapor pressure were necessary in

order to interpret the results.

In the decalin system, emf measurements have been made

from 5 to 55 °0 on solutions of hydrochloric acid with

molalities from 0.005 to 0.05 mol kg" . The vapor pressure

of this solvent mixture is negligible, and no corrections

were necessary to obtain the effective partial pressure of

hydrogen. Experimental measurements have been completed,

and the analysis of the results is in progress.

(E. S. Etz)

36
4. STANDARDIZATION OP ION-SELECTIVE ELECTRODES

A. General Principles

Several different Ion-selective electrodes have been

developed in recent years and have found many applications

in the determination of ion concentrations and activities.

With the widespread use of these new tools, problems

concerning uniform standardization procedures have arisen.

At the request of several persons and organizations

concerned with construction, development, and application

of these electrodes, a program has been started in the

Section to develop standardization methods and standard

reference materials for use with ion-selective electrodes.

In response to this demand a standardization procedure

analogous to that generally accepted for the measurement

of hydrogen ion activities (pH) is being developed. Thus

far, the usefulness of this method has been demonstrated

for chloride-selective and sodium-selective electrodes.

In the approach chosen, an operational definition of

the negative logarithm of the activity of an ionic species

A (designated pA) is formulated as follows:

pA(x) = pA(s) +

EX

" E

S

(nRT In 10)/F (5)

in which EL. and E„ are respectively the emf values for the

cells X and S and n is the number of charges borne by the

ion A:

Ion-selective Soln. X

electrode, ( pA unknown)

and

Salt Reference (X)

bridge, electrode

Ion-selective

electrode,

Soln. S

(pA known)

Salt Reference (s)

bridge, electrode

37
Three considerations are important for the interpretation of pA(x) in terms of the ionic activity in solution X.

1. The activity assigned to the ionic species in the

standard solution (S).

2. The liquid- junction potentials "between the salt

bridge and the solutions in which the measurements are made.

3. The response of the ion-selective electrode with

respect to ionic activity, ionic compositions of the solutions, time variations, and the like.

1. Definition of Activity Coefficients

As is well known, the activity of single ions is a

quantity that is not accessible to direct measurement. It

can only he estimated, and for this reason any pA value

assigned to standard solutions is devoid of exact fundamental meaning. In the establishment of a standard pH

scale, the following convention for the activity coefficient

of chloride ion was adopted [13] *

_AI l/2

los yci"

= "jT^

^

1+1.51

for use at ionic strengths no greater than 0.1. In this

equation, I is the ionic strength (molal scale) and A is

the Debye-Huckel slope constant, having a different value

at each temperature and dielectric constant of the solvent.

In accord with this convention, the pGl of a standard

chloride solution can be defined by

pCl = -log acl- = -log(mcl _ . 7

C1-) = -log m

cl_ Ml/2

+ 1/2- (?)

1+1.51 '

38
For the standardization of ion-selective electrodes, it

is proposed that equation 6 be applied at all temperatures

and at all ionic strengths.

It should be emphasized that a conventional definition

of the activity coefficient is needed for only one ionic

species. With its adoption, the activity coefficients of

all other ions, in theory, become fixed by their interrelationships to the standard ion through mean activity

coefficients or other measurable combinations of ionic

activity coefficients. Thus, this convention for Yr1 - makes it possible to estimate the activity coefficients of

cations from the mean activity coefficients of their chlorides. The proposed definition has the advantage that it

interrelates the activity scales of all ions and incorporates the already existing and generally accepted pH

scale into a single conventional array. This approach was

chosen as the most suitable one after several different

means for the estimation of activity coefficients were

examined and compared.

2. Liquid-Junction Potentials

Like the single-ion activity, the single-electrode

potential and the liquid- junction potential are also inaccessible to direct measurement. Once the activity scale

is defined, however, the liquid- junction potential likewise

becomes fixed. The magnitude of this diffusion potential

(liquid- junction potential) can be obtained from measurements of cells of the type

Ag;AgCl, Soln. of known pCl I Salt bridge, Reference

electrode

in which the location of the liquid junction is indicated

by the vertical line. The emf, E, of this cell, for example.

39
is given by

RT

E = Kef - E

lgCl + Y ^ a Cl" + K

J

(8)

where E° is the standard electrode potential and E. is the

liquid- junction potential.

By varying the composition and ionic strength of the

solution, the effect of these changes on the liquid- junction

potential can be found. Only when E. remains unchanged by

the replacement of the standard solution S by the "unknown"

solution X does the experimental pA(x) obtained from equation 5 correspond closely with -log a, in solution X.

This problem is identical with that which complicates the

useful interpretation of experimental pH numbers [24].

3. Electrode Response

In general, ion-selective electrodes belong to the

class of membrane electrodes. The electrode response is

primarily determined by the ratio of the activities of the

selected ionic species on the two sides of the membrane

and by the relative mobilities of all ions through the

membrane. For the perfect electrode, the transference

number of the selected ion (A) will be 1. In this ideal

situation, the potential of the electrode (E -, ) obeys the

Nernst equation

RT

E = ± — in a (9) ei nF A

No membrane is completely selective, however, and one

of the objectives of this study is to ascertain the magnitude of the effect on the electrode response of "inert"

salts or ions present in the solution of the ionic species

to be determined. The desired information can be obtained

by comparing the response of the ion-selective electrode

40
with that of a "norm" electrode in suitable electrolyte

solutions. For instance, the silver-silver chloride

electrode is a. useful "norm" for the study of chloride

ion response. Ions known to react chemically with the

electrode materials were excluded.

The time response is another important characteristic

of the electrode. In this connection, it is important to

know not only the length of time required for a constant

potential to be reached but also the reproducibility of

this equilibrium or steady- state value.

(M. Alfenaar)

B. Measurement of Chloride Ion Activity

The behavior of two types of chloride-selective electrodes was examined. One electrode was of the "solid state"

design and the other utilized a membrane of the liquid ionexchanger type. A plot of some of the results obtained

with one electrode of the former type in sodium chloride

solutions is shown in figure 9. Similar results for equimolal mixtures of sodium chloride and potassium nitrate

are shown in figure 10.

The experimental values of -log Y^-i- at different

ionic strengths and in different electrolyte solutions were

obtained with electrodes standardized in a sodium chloride

solution of molality 0.01 mol kg" , for which pCl was taken

to be 2.044 at 25 °G in accordance with equation 7. The

curves in figure 9 represent three different "theoretical"

values of -log Yr

-,-» Curve 1 and the curve in figure 10

were derived from equation 6, while the other two curves

show the result obtained when yr n- is set equal to the mean

activity coefficient of sodium chloride (curve 2) and the

mean activity coefficient of potassium chloride (curve J>)

in their pure aqueous solutions. The latter convention is

41
that suggested by Garrels [25]. Studies such as these are

useful in selecting procedures for the standardization of

these electrodes and in making predictions concerning the

precision of experimental measurements of ionic activities.

O

0.30

0.20

0.10

?

-0.I0+-

+

o

2

log I

Figure 9. Response of a chloride ion-selective electrode

in solutions of sodium chloride at 25 °C.

Circles, ion-selective electrode; crosses,

Ag;AgCl electrode. Curves calculated from

theory, see text.

42
o

-0.10

-0.20 -

Figure 10. Response of a chloride ion-selective electrode

in equimolal mixtures of sodium chloride and

potassium nitrate at 25 °C.

X, triangles: ion-selective electrode vs . two different reference electrodes.

+, filled triangles: AgjAgCl electrode vs . two different reference electrodes

Curve from theory, see text.

(M. Alfenaar)

C. Measurement of Sodium Ion Activity

In table 9, values for -log y„ + in aqueous solutions

of different sodium salts, measured with a commercial

sodium-responsive glass electrode fabricated from NAS 11-18

glass and standardized in 0.01m sodium chloride (pNa = 2.044

at 25 °C), are listed. Two different reference electrodes

gave results agreeing, on the average, to 0.01 unit in

pNa. The equipment used for the studies of ionic activity

described above is shown In figure 11.

4?
Table 9. Experimental determination of -log y + In

solutions of sodium salts at different ionic

strengths from measurements with a sodiumselective NAS 11-18 glass electrode at 25 °C

"Observed" -log yM + in solutions of

I NaCl NaNO^

3

Na2S0

2|

NaCl-KNO^

3

Calculated

"log 7Na+

0.0001 -0.25 -0.23 _ 0.005

0.00025 -0.04 - -0.11 - 0.01

0.0005 -0.02 - - - 0.01

0.001 0.02 - 0.01 - 0.02

0.0025 0.03 - 0.02 - 0.02

0.005 0.03 - - - 0.03

0.01* 0.044* o.o4 o.o4 - 0.044*

0.02 - - - 0.03 0.06

0.025 0.06 - 0.06 - 0.07

0.04 - - - 0.06 0.08

0.05 0.08 - - - 0.09

0.1 0.08 0.08 0.11 - 0.10

0.2 - - - 0.12 0.11

0.25 0.10 - 0.10 - 0.12

0.4 - - - 0.13 o.i4

0.8 - - - 0.15 0.15

1.0 0.08 0.09 0.23 — 0.16

Standard solution

44
Figure 11. Equipment for emf measurements with

ion-selective electrodes.

In the last column of table 9 are given standard values

of -log Y-xt + obtained by combining Yp-i- (equation 6) with

the mean activity coefficients of sodium chloride in its

pure aqueous solution. Strictly speaking, these should only

be compared with the "observed" -log -y,, + in sodium chloride

solutions. The rather satisfactory agreement with "observed"

values for other sodium salt solutions, however, suggests

that the ionic strength rule is at least approximately valid

from I = 0.001 to I = 0.8.

Our planned future programs include the testing of the

complete line of available ion-selective electrodes and the

development of standardization procedures and standard

reference solutions for their uniform calibration.

The work described was made possible by the financial

support of The Fbxboro Company of Foxboro, Massachusetts.

(M. Alfenaar)

45
5. EQUILIBRIUM DATA FOR IONIC SYSTEMS

A. Thermodynamics of Hydrochloric Acid in 92. 4l Wt. Percent

Ethanol

The past two decades have seen an increasing interest

in the properties of electrolytes in mixed solvents, especially those solvents of which water Is a constituent.

There is now a considerable body of information on the

halogen acids, limited in general to media, containing at

least 40 wt. percent of water. Most of this information

has been derived from emf measurements of suitable cells.

Measurements of the solubilities of the alkali halides in

mixed solvents have also been made; these are, by their

very nature, less informative than the emf studies which

cover a range of solute concentrations.

The most extensive of these investigations have dealt

with solvents of moderate to high dielectric constant where

It is safe to assume that the electrolyte, strong in water,

is completely dissociated In the mixed solvent as well.

For this reason, a study of the thermodynamic behavior of

hydrochloric acid in 95 vol. percent (92.41 wt . percent)

ethanol by careful emf measurements in the temperature

range 5 to 50 °C was undertaken. The dielectric constant

of this medium is rather low (27.2 at 25 °C), and it may be

expected that the hydrochloric acid will not be fully dissociated. One of the aims of the study was to ascertain

the effect of this ion-pair formation on the accuracy with

which the standard emf of the cell and the thermodynamic

constants of the acid could be obtained.

Electromotive force measurements were made of the cell

Pt;H

2 (g), HC1 (m) in 92.41 wt. % ethanol, AgCljAg

where m represents molality. The expression relating the

emf E and standard emf E° (based on the standard state in

s

46
92. 4l wt. percent ethanol) can be written

E =

S

E° - 2k In (m

±

. s

<y ±

) =

g

E - 2k In {m l±

• sy

l± ) (10)

where m, and y, are the stoichiometric molality and

± s ' ±

J activity coefficient of hydrochloric acid, m+ and y\ are the

corresponding quantities for fully dissociated hydrochloric

acid, and k is written for RT/E. In order to derive the

standard emf from measured values of E over a range of m+, one can utilize the Debye-Huckel formula, for y

l

, , but only if

the ionic concentration m.*+ is known. To obtain the latter,

it was necessary to know the ion-pair dissociation constant

in 92.41 wt. percent ethanol.

Values for the ion-pair constant in ethanol-water

solvents of various compositions are available in the literature, but these are, in general, only known at 25 °C. To

obtain values at the other temperatures from that at 25 °C,

an equation proposed by Puoss [26] was used. In order to

calculate the actual ion concentrations and the real activity coefficient most accurately, it was found desirable to

employ the "extended terms" form of the Debye-Hiickel equation. The manner in which the extended terms are applied

is such that the calculations are very laborious. This

difficulty was circumvented by fitting the extended terms

contributions to a power series in Ba/m^, where B is the

second Debye-Hiickel constant and a is the ion-size parameter.

For a selected value of the ion-pair dissociation constant,

it was then possible to program the entire calculation for

solution with the aid of the time-shared computer.

By the procedure outlined above, "apparent" values of

the standard emf were obtained and found to vary in nearly

linear fashion with m+, making it possible to extrapolate

readily to mj. = 0, Linearity was improved by varying the

value of the ion-size parameter (a) chosen, and the best

47
value was taken to "be that which yielded the minimum

variance from a straight line. The values of a so obtained

were very reasonable (a has been found to be 4.J5A in water

at 25 °C [27] ), and the fact that the variation of a with

temperature is only slight lends confidence to the estimates of the ion-pair dissociation constant over the range

of temperatures studied. The values of K., the ion- pair

constant; a; and E° at 5, 25, and 50 °C are summarized

in table 10.

Table 10. Determination of the standard emf ( E°) for

the cell: PtjHp, HCl(m) in 92.41 wt . f ethanol,

AgCljAg at 5, 25, and 50 °C.

K

d

a

sE

°

oA V

5 0.04^6 4.0 0.10607

25 O.0325 4.1 O.07185

50 0.0214 4.3 0.02269

The standard emf having been determined, it was possible to calculate the stoichiometric activity coefficients

for hydrochloric acid in 92.41 wt . percent ethanol by

1/2 equation 10. The variation of log y+ with m is shown

in figure 12 and compared with that found earlier for

water [27] and for 50 wt. percent ethanol [28]. The lower

values in 92.41 percent ethanol reflect the increased

interionic forces at the relatively low dielectric constant

as well as the incomplete dissociation of the solute.

48
-0.3

-0.4

-0.6

Figure 12. Variation of log y+ for hydrochloric acid with

the square root of the molality in 92.41 wt

.

percent ethanol compared with that in water

and in 50 wt. percent ethanol.

dotted line, Debye-Huckel limiting slope

for 92.41$ ethanol

The change of log y with temperature furnishes a

means of evaluating the relative partial molal enthalpy

(TL) and the relative partial molal heat capacity (T2

) for

hydrochloric acid in this mixed solvent. The values at

25 °C for the 0.1m solution in water and in 92.41 wt.

percent ethanol are compared in table 11,

49
Table 11. Comparison of Lp and J

2 for hydrochloric acid

(0.1 molal) in water and in 92.41 wt. percent

ethanol at 25 °C.

Lp Jp

J mol~ J K~ mol

Water 742 13

92 .41 wt. percent ethanol 8590 157

(K. H. Pool, R. G. Bates)

B. Second Dissociation Constant of Carbonic Acid in

Deuterium Oxide

The dissociation constants of several weak acids in

deuterium oxide [11,12], as well as the ionization constant

of deuterium oxide itself [29], have been determined in the

Section by the emf method. These studies have now been

extended to include a measurement of the second dissociation constant of deuterocarbonic acid in heavy water.

This acid-base equilibrium has not been thoroughly investigated, and existing values for the dissociation constant

at 25 °C differ considerably [30,31],

The dissociation constant was determined from 5 to

50 °C by measuring the emf of the cell

Pt;D

2 (g), NaDCO^m), Na^CO^(m), NaCl(m) in D

2

0, AgCljAg

50
p

The molality (m) of each carbonate was kept equal to that

of the sodium chloride in all of the cell solutions; m was

varied from 0.00^ to 0.03 mol kg

-1

An expression relating the "apparent" dissociation

constant Kp to the emf E of the cell was obtained by combining the mass law with the Hernst equation for the emf:

E-E* ^)C0- + [°D-J 2AI1/2d

Q

l/2

PK* = + log m

pn

_ + log =2 + yyx—T/ (11)

k

01

m

co

= -[CD

-

] l+BaI1/dd 1/d

3

where E° is the standard emf of the cell, k is (RT In 10)/F,

A and B are the Debye-Huckel constants in deuterium oxide,

I is the ionic strength, a is the ion-size parameter, and d

is the density of heavy water. The small correction for

hydrolysis was made by estimating the deuteroxide ion concentration from the emf and the self-dissociation constant

of deuterium oxide.

The value of pK

?

, that is, the limit of pKp at 1=0,

was obtained by fitting the data to an equation linear in

I by the method of least squares. A plot of the data at

25 °C is shown in figure 1J5. The values of pK?

from one

series are consistently higher than the results from the

other two series. This is presumed to have been caused by

a small amount of tritium in the heavy water used for this

run, as this lot of solvent had a radioactivity of about

0.4 u-Ci/cnr5

. The results of this series were not used in the final

calculation of pKp

. Instead, another run was made with

heavy water that had a relatively low activity of about

0.14 [iCi/cnr (see Section 2.D.J5). The values of pKp found

were identical with those of two earlier runs within the

limits of experimental error.

I 51
Q.

0.05 0.10

IONIC STRENGTH

0.15

Figure lj5. Determination of pK? for deutero carbonic

acid at 25 *C

Several different values of the ion-size parameter were

used in the calculations of pK', and that which gave the

smallest standard deviation in pKp

was selected. In this

way, the ion- size parameter was found to vary from 6.0 K at

5 °C to 7.2 A at 50 °C. The pK2 values (molal scale) found

are listed in table 12. The value at 25 °C ( 11.077 ) is to

be compared with 10. 85 found by Curry and Hugus [30] and

with 11.00 of Glasoe and Long [31].

52
Table 12. Dissociation of the deutero carbonate ion (DCO-, ) 3

in deuterium oxide from 5 to 50 °C.

.t, °C pK2

t, °C pK2

5 11.332 30 11.028

10 11.259 35 10.984

15 11.192 40 10.945

20 11.132 45 10.911

25 11.077 50 10.882

(M. Paabo)

C. Properties of Acids and Bases in Hydrogen-Bonded Mixed

Solvents

Binary mixed solvents which contain water as one component react protolytically both with hydrogen acids and

with their conjugate bases. Although the acid-base behavior

of many of these solvent mixtures resembles that of water

in kind, wide differences of degree are observed.

During the past year a comprehensive review of the

equilibrium properties of acids and bases in amphiprotic

mixed solvents was prepared. The effect of adding an

organic constituent to aqueous solutions of acids and bases

was examined through summaries of data for standard electromotive forces, equilibrium constants, medium effects, and

thermodynamic functions for the transfer of weak and strong

electrolytes from water to mixed solvents of various compositions. Systems consisting of strong and weak electrolytes in mixed solvents were considered.

53
The solvent effect on an acid-base equilibrium depends

strongly on the charge type of the acid and its conjugate

base. The change in dielectric constant produced by altering the composition of the solvent also has a large influence. Other factors tending to shift the equilibrium are

more difficult to identify. In general, they may be classified as secondary solute-solute interactions and solutesolvent interactions. Any side reaction capable of altering

the free energy of one or more of the participants in the

primary equilibrium will evidently contribute to the solvent

effect. Ion-solvent interactions, in particular, are influenced by hydrogen-bonding capability, the formation of

Ion pairs and higher aggregates, sorting of the two types

of solvent molecules by the ions, and alterations of solvent

basicity. In the water-like region of relatively high dielectric constant, solvent effects show a considerable uniformity that disappears as the solvent becomes rich in the

organic constituent.

There is a growing realization that the structure of

water and of the mixed solvents themselves plays a primary

role in altering the equilibrium properties of acids and

bases in mixed solvents. In evaluating solvent effects,

attention is usually focused on the change in free energy

and its relation to the solvent composition. It is now

evident that free energy changes of this sort are the resultant of large changes in enthalpy and entropy. This may

be seen in figure 14, where the changes in AG°, AH°, and

TAS° for the transfer of a mole of hydrochloric acid from

the standard state in water to the standard state in

ethylene glycol-water solvents are plotted as a function of

the composition of the solvent mixture. It is also true

for the dissociation process of a weak acid, as shown in

figure 15. In this figure, the thermodynamic functions for

the process,

54
H

+

(w) + A<r(w) + HAc(s) = H

+

(s) + Ac~(s) + HAc(w)

in which acetic acid (HAc) is transferred from water (w) to

a series of ethanol-water mixtures (s), are plotted as a function of the solvent composition. The values of these

energy quantities are shown in kilocalories; 1 kcal = 4.184 kj.

kcal

mof1

I -

-I -

-2 -

1 1

HYDROCHLORIC ACID

i i

ag;/

^Nah?

'

i

\tas; _

20 40 60

WT. PERCENT GLYCOL

80 100

Figure 14. Standard thermodynamic functions for the

transfer of hydrochloric acid from water

to ethylene glycol-water solvents at 25 °C

55
kcal

mol

2 -

2 -

-4 -

1 1

; ACETIC ACIO

1 A

/AG?

.AH?

i i

\TASt

i 20 40 60 80 100

WT. PERCENT ETHANOL

Figure 15. Standard thermodynamic functions for the

transfer of acetic acid from water to

ethanol-water solvents at 25 °C.

These conclusions were presented in a plenary lecture

at the Symposium on Equilibria and Reaction Kinetics in

Hydrogen-Bonded Solvent Systems held at the University of

Newcastle upon Tyne (England) in January. This international meeting was arranged in recognition of the

retirement of Lord Wynne-Jones from the position of head

of the department of physical chemistry at Newcastle upon

Tyne.

(R. G. Bates)

56

fl^H
v

6. CONDUCTANCE AND TRANSPORT BEHAVIOR OF ELECTROLYTES

A. Oceanographic Conductance Standards

The oceanographer uses salinity determinations for two

main purposes: to identify and trace water masses and as a. starting point for computing other properties such as the

density and sound velocity. At present, salinity determinations are made almost entirely by conductometric methods,

and for maximum usefulness the conductance should be known

to 50 ppm [32]. This requirement places the oceanographic

needs at the very forefront of developing techniques, since

ery few research measurements of conductance have been

made with a standard error less than 200 ppm. Equipment

with the requisite sensitivity has been developed for salinity determinations, and high precision is obtained either

by frequent calibration with a standard or by a differential

method in which the unknown is compared directly with a

standard.

The internationally accepted standard for this purpose

is natural sea water that has been carefully certified for

chlorinity. Standards corresponding to other salinities

are urgently needed, and it seems advisable to relate them

to a single pure electrolyte so that there is a direct

relationship between conductivity and a chloride assay or

the weight of solid electrolyte. For the above reasons,

the Section has initiated a program to develop as standards

solutions of 0.5, 1, 2, and 4 percent sodium chloride in

water. The specific conductance of these solutions at 15

and 25 °C will be certified.

B. Apparatus

The major new item of equipment is the transformer

ratio-arm conductivity bridge shown in figure 16. This

instrument is based on the design of Janz and Mclntyre [33]

and is built around a commercial impedance comparator.

57
Figure 16. Janz-Mclntyre conductivity bridge.

The cabinet shown on top of the comparator in the figure

contains switches for use with four-terminal cells in order

to eliminate the effects of lead resistance, and it also

contains the variable capacitors needed to equalize the

phase angle of the standard and unknown impedances.

The major departures from the recommendations of

Janz and Mclntyre are the selection of a more sensitive

model of comparator and a narrower range of frequencies

This instrument can indicate an unbalance of impedance of

10 ppm, and the specified accuracy is 3>0 ppm. When highquality resistors were interchanged between the Reference , and Standard connections of the bridge, no inequality of

58
the ratio arms was detectable. The comparator was specially modified to provide frequencies of 1, 2.5, 5, and 10 kHz.

Within this range,, the a-c impedances of the coils used as

standard resistors differ negligibly from the d-c calibration values.

Some conductivity cells of a demountable design are

shown in figure 17. The heart of the design is the quartz

body, which has optically flat and parallel ends and a precisely-ground cylindrical interior. The cells are

closed by plane electrodes made of platinum sheet soldered

with pure tin to quarter-inch stainless steel plates for

rigidity. Tests are in progress to determine the reproducibility of the cell constants as the cells are taken

apart and reassembled, but it is hoped that the constant

can be determined absolutely from the dimensions of the

quartz cylinder with an uncertainty less than 50 ppm.

Figure 17. Demountable conductivity cells

59
The parts shown in figure 17 provide cell constants of

: a

-1

0.02, 0.03, 0.48, and 0.51 cm . Additional bodies are

being prepared to give constants of 7.6 and J>0.6 cm"

Experience thus far indicates that aqueous solutions and

many nonaqueous solutions can be contained in the cells

without the use of gaskets or seals. One interesting application of one of these cells is described in the next

section of this report. The glass cells of conventional

design and large constant (10, 3>5.> and 100 cm ) have been

constructed for use in the work on oceanographic conductance

standards.

(T. B. Hoover)

C. Characteristics of "Azulene Periodate " A conductometric characterization of an azulene derivative, believed to be a polymeric electrolyte, was

requested by the Organic Chemistry Section. Only 3 rng of

the material (in the form of a nearly saturated solution

in 5 cm of dimethylformamide) could be supplied. Less

3 than 2 cnr of this stock solution was actually used for the

following measurements.

The specific conductivities of the stock solution and

of three dilutions were measured in the smaller assembled

cell shown in figure 17. This cell has a constant of

-1 "5 0.509 cm and a volume of 0.314 cm . The major limitation

of the measurements was the relatively high specific conductance of the solvent used to prepare the original solution. The solvent correction amounted to 10 percent in

the case of the stock solution and 25 percent for the third

dilution.

A "solute specific conductance" was calculated for

concentration units in grams per liter and was extrapolated

as a function of the square root of the concentration to

60
give a. limiting value not less than 0.3 1 g" 1

ft

"" « This

latter value established an upper limit of ^20 for the

equivalent weight of the unknown material, since all electrolytes whose conductances have been measured in dimethylformamide have limiting equivalent conductances less than

95 1 mol" 1

, The slope of the plot of the "solute

specific conductance" vs. the square root of the concentration was also much greater than reasonable estimates of the

limiting Fuoss-Onsager slope for the same solvent. The

experimental results were not consistent with the hypothesis

that the specimen was a strong electrolyte nor that a strong

electrolytic impurity was being observed. They conformed,

however, fairly well with those expected for a weak electrolyte having pK, between 3 and 4. The pK, for the dissociation of ion pairs in dimethylformamide is near 2 for several salts [3^-], so "that it may be concluded that the

azulene derivative is a moderately weak electrolyte.

(T. B. Hoover)

D. Hydrodynamic Properties of Ions

The parameters which determine the mobility of an ion

in an electric field have been of concern almost since the

beginnings of ionic theory. Ideally, one would like to be

able to separate the factors into those characteristic of

the solvent and others characteristic of the ion. The

extremely useful Walden rule, in combination with the Stokes

law, states that the limiting equivalent conductance is

inversely proportional both to the solvent viscosity and to

the ionic radius. Fuoss observed, from extensive studies

in binary solvent systems, that the Walden product (viscosity times equivalent conductance) varies systematically

with dielectric constant and offered the explanation of

dielectric friction [35]. This friction is the retarding

61
force on an ion, in addition to Stokes friction, that results from the energy dissipation in reorienting solvent

dipoles around the moving charge. By heuristic arguments,

Fuoss arrived at an equation of the form

l/0?oO * R

o + A/D

( 12 )

where rj and D are the viscosity and dielectric constant

of the solvent, X is the limiting equivalent conductance

of an ion, and R is its "Stokes" radius. He suggested

empirically that the constant A was approximately 3>6/R

in aprotic solvents and nearly 4R when one component of

the solvent is hydrogen-bonding.

A rigorous derivation of equation 12 for a dielectric

continuum model of the solvent was given by Zwanzig [36] • The term A turned out to be proportional to t/R , where t is a third solvent parameter, the dielectric relaxation

time. When applied to experimental data, the Zwanzig equation runs into inconsistencies that may be best expressed

by saying that most solvents have dielectric relaxation

times differing by more than an order of magnitude from the

values required by the conductance data.

Since the weakness of the Zwanzig treatment may derive

from the dielectric continuum model, the application of the

Eyring reaction rate theory to conductance is being explored

in the Section. Eyring' s model considers discrete solvent

molecules which are comparable in size to the ions, A key

parameter in his treatment of conductance or viscosity is

the characteristic jump distance of the solvent. Following

the explanation of Fuoss, it is plausible that a major part

of the activation energy for conductance should be the

electrostatic energy of the dielectric atmosphere, which

must be reoriented at each jump of the ion to a new vacancy.

62
This step leads directly to an exponential form of equation

proposed by Mezhennii [37] for the variation of the Walden

product with dielectric constant:

%K * Ro exp(-S/D) (13)

Empirically, equation 12 and equation 13 seem equally

satisfactory.

Rough estimates have been obtained for the parameters

of equation 13 for six electrolytes whose conductances have

been measured in six or more pure aprotic solvents. The

values of R were all smaller than the crystallographic

radii of the ionsj S values for the anions were all negative but probably not significantly different from zero,

while the corresponding parameters for the cations showed

larger, positive values and appreciable differences. The

dielectric energy of the region surrounding an ion is independent of the sign of the charge on the central ion, so the

distinction between anions and cations may have its source

in an entropy of activation. Mine and Werblan [38] have

measured the entropy of activation for conductance for several electrolytes in acetonitrile and have found that this

quantity is correlated with solvation energies, being positive for small cations and negative for perchlorate and

tetraethylammonium ions.

The numerical results obtained thus far are so tentative that very little importance can be attached to them,

but the approach gives promise of providing a fresh insight

into the hydrodynamic behavior of ions in various solvents.

(T. B. Hoover)

63
7. PREPARATION AND PROPERTIES OF SOLVENTS

A. Propylene Carbonate

Due to its relatively high dielectric constant (64.4

at 25 °C) and good solvent properties, propylene carbonate

is rapidly gaining favor for electrochemical studies. The

following method of purification was devised by combining

useful features of several proposed methods [~59], at the

same time eliminating unnecessary steps.

The commercial propylene carbonate was dried overnight

with molecular sieves and then distilled twice. The initial

150 to 200 cm^ of distillate from the first distillation of

about 3 1 was discarded. The major fraction, distilled at

106 to 109 °C under a pressure of 7 to 8 torr, was collected

for use. The product had a specific conductivity of between

3x10"^ and 8x10"^ n"

1

. A second distillation at 114 to

116 °C and 4 to 5 torr yielded a product having a specific

o _o _-i

conductivity of 2x10 to 4x10 fi . It was believed to

be of sufficient purity for emf studies.

Further work on the characterization of pure propylene

carbonate is under way. In addition, the properties of the

50 wt. percent mixture of methanol and propylene carbonate

will be determined over a range of temperatures. These

properties will include dielectric constant, vapor pressure,

and density.

(B. R. Staples)

B. Gryoscopic Behavior of N-Methylpropionamide

Because of Its remarkably high dielectric constant

(about 175 at room temperature) and useful solvent properties, N-methylpropionamide is of interest in the Section

for electrochemical and kinetic studies. As part of a characterization program, the freezing point and cryoscopic

64
constant of this solvent were determined. A 50~g sample

was frozen, with mechanical stirring, in a dewar flask

while protected from atmospheric moisture by a stream of

dry nitrogen. The temperature was measured with a calibrated platinum resistance thermometer and Mueller bridge. The

unbalance of the bridge was amplified and continuously

recorded on a chart recorder. Heat was abstracted at a rate of approximately 2 J/s and the run was continued until

nearly one-third of the charge had frozen, when the mixture

became too stiff for effective stirring. This limitation

prevented a cryoscopic determination of sample purity, but

the specimen was representative of the best material that

had been used for conductance and dielectric constant

measurements.

The cryoscopic constant was determined by measuring

the lowering of the freezing point produced by known additions of N-methylacet amide. This solute has a freezing

point about 60 °C higher than that of the solvent, but it

should form nearly ideal solutions with N-methylpropionamide

because of the close chemical similarity of the two compounds. The results are summarized in table 13 . The cryoscopic constant of 4.68 deg kg/mol found for N-methylpropionamide corresponds to a heat of fusion of 9*2 kJ/mol.

This quantity apparently has never been measured calorimetrically.

Table 13 . Freezing point and cryoscopic constant of

N-methylpropionamide

.

Molality of N-methylacetamide - 0.0425 O.I658 0.2216

Freezing point, -30.917 -30.920 -351.152 -31.695 -31.958

°C.

Cryoscopic 4.681 4.674 4.685

constant,

deg kg/mol.

65

(T. B. Hoover)
8. FOREIGN-CURRENCY PROGRAMS

The two projects sponsored by the Electrochemical

Analysis Section at universities in India and financed

through the provisions of Public Law 480 continued active

during the reporting year. A brief description of these

programs follows.

Properties of Electrolytes in Nonaqueous Media, Jadavpur University, Calcutta, Dr. M. N. Das and Dr.

K. K. Kundu . (a) Dissociation constants and related thermodynamic

quantities for acetic acid in ethylene glycol and propylene

glycol over the temperature range 5 to 45 °C were determined.

(b) The standard potential of the silver- silver bromide

electrode in propylene glycol was determined from 5 to 45 °C.

(c) Standard potentials of the silver-silver bromide

electrode in isodielectric media formed by mixing methanol

with propylene glycol were determined at 25 °C.

(d) The standard free energy change on the transfer

of sodium bromide from water to isodielectric mixtures of

methanol and propylene glycol was measured at 25 °C.

(e) The acidic dissociation constants of ammonium ion,

diethyl ammonium ion, and triethylammonium ion were determined in ethylene glycol over the temperature range 5 to 45 °C.

Physicochemical Studies in Nonaqueous Solvents, PanJab

University, Chandigarh. Professor R. C. Paul . The current work under way in Professor Paul's laboratory includes studies of ion-solvent processes and the

behavior of Lewis acids in media of limited hydrogen-bonding

capability. Studies of ion-pair formation and the dissociation of weak electrolytes using conductance techniques are

also of interest.

(R. G. Bates)

66
9. REVIEWS, CONFERENCES, AND COMMITTEE ACTIVITY

Three reviews on specialized topics in solution chemistry were prepared, during the year. The first, co-authored

with W. J. Waddell, dealt with intracellular pH and was

prepared for publication in Physiological Reviews . The

second was based on a lecture given at an international

symposium on intracellular glass microelectrodes (Montreal,

May 1967) and bore the title "inner Reference Electrodes".

The third was prepared for presentation as a plenary lecture

at an international symposium on "Equilibria and Reaction

Kinetics in Hydrogen-Bonded Solvent Systems" (Newcastle,

January 1968) [see Section 5.C]. A monograph "Acid-Base

, Behavior in Aprotic Organic Solvents" by M. M. Davis will

be published shortly in the NBS Monograph Series. Three

articles on pH measurement were prepared for forthcoming

handbooks and encyclopedias.

Considerable work on nomenclature, terminology, and

units was done under the auspices of the U.S. National

Committee for ISO/TC-12, the IUPAC Commission on Symbols,

Terminology, and Units, and two committees of the U.S.A.

Standards Institute. The IUPAC activity culminated at the

Prague Conference (August 1967) where a draft of the new

manual of symbols and units was completed.

A revised draft of an ASTM standard on pH measurement

was prepared.

(R. G. Bates)

67
10. IMPACT OF SECTION'S ACTIVITIES

The impact of the Section's activities can best be

judged in terms of the distribution of Section publications,

invited talks, and requests for advice and consultation.

In this connection, the following statistics, compiled for

the calendar year 1967 unless otherwise indicated, are

pertinent.

Reprints requested 88l

Manuscripts reviewed for

outside journals ^0

Project proposals,

books reviewed ............. 5

Significant consultations ........ J50

universities 7

industries 8

federal and

state agencies 15

Committee memberships 18

Publications (7/67 to 6/68) 12

Talks (7/67 to 6/68) . 15

scientific societies 6

universities 7

industries 2

68
11. PERSONNEL

Roger G-. Bates, Section Chief

Marguerite Raudehbush, Section Secretary

Conductance Measurements

T. B. Hoover

Electromotive Force and

Acidity Measurements

Aqueous solutions,

mixed solvents

Heavy water,

mixed solvents

Nonaqueous solvents

Ion-specific electrodes

Consultant

B. R. Staples

M. Paat>o

K. H. Pool (July to Sept.)

from Washington State

University, Pullman

E. S. Etz (July to Sept.)

from Clarkson College

of Technology

M. Alfenaar (Oct. to June)

from University of

Utrecht, Holland

D. Rosenthal, Clarkson

College of Technology

69
12. COMMITTEE MEMBERSHIPS

R. G. Bates

Member, Subcommittee on Hydrogen Ion Determinations,

ASTM Committee E-l, Sub. 22.

Councilor, American Chemical Society,

representing the Division of Analytical Chemistry.

Member, NRC Committee on Symbols, Units, and

Terminology for the Physical Sciences

(U.S. National Committee for ISO/TC-12).

Member, U.S. National Committee for IUPAC.

Member, Commission on Electrochemistry, IUPAC.

Member, Commission on Symbols, Terminology, and

Units, IUPAC.

Associate Member, Commission on Electroanalytical

Chemistry, IUPAC.

Member, Subcommittee on Physicochemical Symbols and

Terminology, Committee on Physical Chemistry, NRC.

Member, Committee Yl, Abbreviations for Science

and Engineering; U.S.A. Standards Institute.

Member, Committee Y10, Letter Symbols, U.S.A.

Standards Institute.

Member, Board of U.S. Civil Service Examiners

(Chemistry Panel).

Member, Advisory Board for Analytical Chemistry . Member, Board of Editorial Advisers, Analytica

Chimica Acta . Member, Washington Editorial Review Board (NBS)j

deputy chairman.

Member, Museum Committee (NBS).

Member, Advisory Committee on the Senior Lunch Club

(NBS).

T. B. Hoover

Member, Community Services Committee, Chemical Society

of Washington.

Member, Audit Committee, D.C. Chapter, Society of the

Sigma XI.

70
13- PUBLICATIONS AMD MANUSCRIPTS, JULY I967 TO JUNE 1968

1. Bates, R.G., (Ed.)

Electrochemical Analysis Section: Summary of

Activities, July 1966 to June 1967.

NBS Tech. Note 423 (1967).

2. Robinson, R.A.

The Dissociation Constants of Some Substituted

Nitrophenols in Aqueous Solution at 25 °C.

J. Res. Nat. Bur. Stand., 71A, 385-389 (1967).

3. Hetzer, H.B., Robinson, R.A., Bates, R.G.

Conventional pa^j Values for Buffer Solutions of

Piperazine Phosphate from 0° to 50° C.

Anal. Chem., 40, 634 (1968).

p.

4. Bates, R.G., Covington, A.K.

Behavior of the Glass Electrode and Other

pH-Responsive Electrodes in Biological Media.

Ann. N.Y. Acad. Sci., 148, 67-80 (1968).

5. Bates, R.G.

Book Review—"Titrations in Nonaqueous Media"

by W. Huber.

Anal. Chem., 40, 111A (1968).

6. Covington, A.K., Paabo, M., Robinson, R.A.,

Bates, R.G.

Use of the Glass Electrode in Deuterium Oxide and

the Relation between the Standardized pD (pa^N

Scale and the Operational pH in Heavy Water. ' Anal. Chem., 40, 700-706 (1968).

7. Bates, R.G.

Standardization of Acidity Measurements. Extension

of the pH Concept to Mixed Solvents and Heavy Water.

Anal. Chem., 40, 28A-38A (May 1968).

8. Schindler, P., Robinson, R.A., Bates, R.G.

Solubility of Tris(hydroxymethyl)aminomethane in

Water-Methanol Solvent Mixtures and Medium Effects

in the Dissociation of the Protonated Base.

J. Res. Nat. Bur. Stand., 72A, l4l-l48 (1968).

9. Hetzer, H.B., Robinson, R.A., Bates, R.G.

Dissociation Constants of Piperazinium Ion and

Related Thermodynamic Quantities from to 50 °C.

J. Phys. Chem., £2, 2081-2086 (1968).

71
10. Bates, R.G.

Measurement of pH

Handbook of Biochemistry, Chemical Rubber Co. (1968)

11. Robinson, R.A., Covington, A.K.

The Thermodynamics of the Ternary System:

Water-Potassium Chloride-Calcium Chloride at 25 °C.

J. Res. Nat. Bur. Stand., 72A, 239-245 (1968).

12. Bates, R.G.

Book Review—"Glass Electrodes for Hydrogen and

Other Cations" by G. Eisenman (Ed.)

The American Scientist, 56, 148A (1968).

13. Davis, M.M.

Acid-Base Behavior in Aprotic Organic Solvents.

NBS Monograph 105 . (in press).

14. Bates, R.G.

Medium Effects and pH in Nonaqueous Solvents

Chapter for "Solvation Interactions"

(C. D. Ritchie and J. P. Coetzee, Eds.)

Marcel Dekker, publisher. (in press).

15. Bates, R.G.

Hydrogen Ions

Encyclopaedia Brittanica. (in press).

16. Bates, R.G.

Inner Reference Electrodes and Their Characteristics

Chapter in "intracellular Glass Microelectrodes"

(M. Lavallee and N. Hebert, Eds.)

John Wiley and Sons, New York, N.Y. (in press).

17. Bates, R.G.

Equilibrium Properties of Acids and Bases in

Amphiprotic Mixed Solvents.

Chapter in "Equilibria and Reaction Kinetics in

Hydrogen-Bonded Solvent Systems"

(A. K. Covington and P. Jones, Eds.) (in press).

18. Hoover, T.B.

The N-Methylpropionamide-Water System. Densities

and Dielectric Constants at 20 to 40°.

J. Phys. Chem. (in press).

19. Waddell, W.J., Bates, R.G.

Intracellular pH.

Physiological Reviews. (in press).

72
14. TALKS, JULY I967 TO JUNE I968

1. Hoover, T.B., "N-Methylpropionamide-Water Systems.

Densities and Dielectric Constants",

American Chemical Society, Chicago, Illinois,

September 14, 1967.

2. Bates, R.G., "Equilibrium Properties of Acids and

Bases in Amphiprotic Mixed Solvents", in Symposium

on Equilibria and Reaction Kinetics in HydrogenBonded Solvent Systems, University of Newcastle

upon Tyne, England, January 11, 1968.

3. Bates, R.G., "Meaning and Standardization of pH

Measurements", Unilever Research Laboratory,

Port Sunlight (Cheshire), England, January 15, 1968.

4* Bates, R.G., "Meaning and Standardization of pH

Measurements", Unilever Research Laboratory,

Bedford, England, January 16, 1968.

5. Bates, R.G., "Acidity Measurements in Mixed Solvent

Systems", Chemistry department seminar, Drexel

Institute of Technology, Philadelphia, Pa.,

January 25, 1968.

6. Bates, R.G., "Acid-Base Behavior in Hydrogen-Bonded

Mixed Solvent Systems", Iowa State University,

Ames, Iowa, March 7, 1§68.

7. Bates, R.G., "Acid-Base Behavior In Hydrogen-Bonded

Mixed Solvent Systems", Chemistry department

seminar, University of Arizona, Tucson, Arizona,

March 28, 1968.

Bates, R.G., "Acid-Ba.se Behavior in Hydrogen-Bonded

Mixed Solvent Systems", Chemistry department seminar,

University of California at Santa Cruz, California,

March 29, 1968.

9. Staples, B.R., "The Kinetics of Proton Transfer

Reactions Determined by the Second WIen Effect",

(paper by Bert R. Staples and Gordon Atkinson),

American Chemical Society Meeting, San Francisco,

Calif., April 2, 1968.

10. Paabo, M., "Standardization of a pD Scale", (paper by

Maya Paabo and Roger G. Bates), American Chemical

Society Meeting, San Francisco, Calif.,

April 3, 1968.

73
11. Staples, B.R., "Two New Standards for the pH Scale",

(paper by Bert R. Staples and Roger G. Bates),

American Chemical Society Meeting,

San Francisco, Calif., April 3, 1968.

12. Pool, Karl H., "Thermodynamics of Hydrochloric Acid

in 95 Vol. Jo (92.41 Wt. %) Ethanol from E.m.f

Measurements at 5 to 50 °C", (paper by Karl H. Pool

and Roger G. Bates), American Chemical Society

Meeting, San Francisco, Calif., April 5., 1968.

13. Alfenaar, M., "Standard Chemical Free Enthalpy,

Enthalpy, Entropy and Heat Capacity of Hydration

of the Hydrogen Ion at 25 °C", Chemistry Department,

University of Oregon, Eugene, Oregon, April 9 f 1968.

14. Bates, R.G., "Acid-Base Interactions in Alcohol-Water

Media", Maryland Section, American Chemical Society,

Towson, Maryland, April 17, 1968.

15. Bates, R.G., "Meaning and Standardization of pH

Measurements",, Undergraduate Research Participation

Seminar, University of Maryland, College Park,

Maryland, June 26, 1968.

74
15. LIST OF REFERENCES

[1] Electrochemical Analysis; Studies of Acids, Bases,

and Salts by Emf, Conductance, Optical, and

Kinetic Methods. July lgb5 to June 19ob . (R. G. Bates, Ed. J, NBS Technical Note 400

(September 6, i960).

[2] Electrochemical Analysis Sectiont Summary of

Activities, July 19bb to June 19b7 « "R. G. Bates, Ed. J, NBS Technical Note 423

October 1967).

[3] British Standard 1647: Specification for pH Scale , British Standards Institution, London (19bl).

[4] Bates, R.G.

Anal. Chem., 40, 28A (May 1968).

[5] Brszfasted, J.N.

Rec. Trav. Chim. Pays-Bas, 42, 718 (1923).

[6] Bates, R.G.

"Medium effects and pH in nonaqueous solvents".

Chapter in Solute-Solvent Interactions

(J. F. Coetzee and C. D. Ritchie, Eds.).

Marcel Dekker. In press.

[7] Bates, R.G.

Chem. Rev., 42, 1 (1948).

[8] Randies, J.E.B.

Trans. Faraday Soc, 52, 1573 (1956).

[9] Bates, R.G., Pinching, G.D., and Smith, E.R.

J. Res. Nat. Bur. Stand., 45, 4l8 (1950).

[10] Covington, A.K., Paabo, M., Robinson, R.A., and

Bates, R.G.

Anal. Chem., 40, 700 (1968).

[11] Gary, R., Bates, R.G., and Robinson, R.A.

J. Phys. Chem., 68, 3806 (1964).

[12] Gary, R., Bates, R.G., and Robinson, R.A.

J. Phys. Chem., 6£, 2750 (1965).

[13] Bates, R.G. and Guggenheim, E.A.

Pure Appl. Chem., 1, 163 (i960).

75
[14] Bates, R.G.

J. Amer. Chem. Soc, 7£, 1579 (1948).

[15] Bates, R.G, and Pinching, G.D.

J. Amer. Chem. Soc, 71, 1274 (1949).

[16] Braude, E.A, and Stern, E.S.

J. Chem. Soc, 1976 (1948).

[17] Bernal, J.D. and Fowler, R.H.

J. Chem. Phys., 1, 515 (1935).

[18] Pranks, P. and Ives, D.J.G.

Quart. Rev. (London), _20, 1 (1966).

[19] Kundu, K.K., De, A.L., and Das, M.N.

Private communication (1967).

[20] Sager, E.E., Robinson, R.A., and Bates, R.G.

J. Res. Nat. Bur. Stand., 68A, 305 (1964).

[21] Paabo, M., Bates, R.G., and Robinson, R.A.

J. Phys. Chem., £0, 247 (1966).

[22] Schlndler, P., Robinson, R.A., and Bates, R.G.

J. Res. Nat* Bur. Stand,, 72A, l4l (1968).

[23] Bass, S.J., Nathan, W.I,, Melghan, R.M., and Cole, R.H.

J. Phys. Chem., 68, 509 (1964).

[24] Bates, R.G.

Determination of pH, Theory and Practice , Chap. 3

and 4. John Wiley and Sons, Inc., New York (1964).

[25] Garrels, R.M.

"Ion-sensitive electrodes and individual Ion activity

coefficients". Chapter 1J> in Glass Electrodes for

Hydrogen and Other Cations (G. Eisenman, Ed.)

Marcel Dekker, New York (I967 )

[26] Puoss, R.M.

J. Amer. Chem. Soc, 80, 5059 (1958).

[27] Bates, R.G. and Bower, V.E.

J. Res. Nat. Bur. Stand. , 53, 283 (1954).

[28] Oiwa, I.T.

Sci. Rpts. Tohoku Univ. [1], 4l, 47 (1957).

76
[29] Covington, A.K., Robinson, R.A., and Bates, R.G,

J. Phys. Chem., JO, 3820 (1966).

[30] Curry, J. and Hugus Jr., Z.Z.

J. Amer. Chem. Soc, 66 651 (1944)

[31] Glasoe, P.K. and Long, P. A.

J. Phys. Chem., 64, 188 (i960).

[32] Cox, R.A.

Progr. Oceanogr., 1, 243 (1963).

[33] Janz, G.J. and Mclntyre, J.D.E.

J. Electrochem. Soc, 108, 272 (1961).

[34] Prue, J.E. and Sherrington, P.J.

Trans. Faraday Soc, 57, 1795 (1961).

[35] Puoss, R.M.

Proc Nat. Acad, Sci. U.S., 45, 807 (1959).

[36] Zwanzig, R.W.

J. Chem. Phys., 38, I603 (1963).

[37] Mezhenni, Ya.F.

Ukr. Khim. Zh., 24, 703 (1958).

[38] Mine, S. and Werblan, L.

Rocz, Chem., 40, 1989 (1966).

[39] Jasinski, R.J. and Kirkland, S.

Anal. Chem., 39, 1663 (1967).

77

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