(1 Bureau ol

.brary, <

'

OCT 4 1965

NBS

Reference Desk not to

taken iry.

technical ^2©te 271

ELECTROCHEMICAL ANALYSIS: STUDIES OF

ACIDS, BASES, AND SALTS BY EMF, CONDUCTANCE,

OPTICAL AND KINETIC METHODS

JULY 1964 TO JUNE 1965

Edited by Roger G. Bates

I. S. DEPARTMENT OF COMMERCE

NATIONAL BUREAU OF STANDARDS

NATIONAL BUREAU OF STANDARDS

Technical Note 271

ISSUED SEPTEMBER 6, 1965

ELECTROCHEMICAL ANALYSIS: STUDIES OF

ACIDS, BASES, AND SALTS BY EMF, CONDUCTANCE

,

OPTICAL AND KINETIC METHODS

JULY 1964 TO JUNE 1965

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 60 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 seven sections and about 85 technical personnel encompassing some 50 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 highly

important. In addition, however, it has been my experience

that informal, annual summaries of progress describing

efforts of the past year can be very valuable in disseminating information. At the National Bureau of Standards such

publications fit logically into the category of Technical

Note. In 1965 we plan to issue these summaries for all of

our sections. The following is the first annual report on

progress of the Electrochemical Analysis Section.

W. Wayne Meinke, Chief

Analytical Chemistry Division

ii
PREFACE

This is the first in a series of annual progress

reports of the Electrochemical Analysis Section of the

Analytical Chemistry Division. The report covers the fiscal

year 1965, which began on July 1, 1964, and ended June JO,

1965.

As a discrete entity, the Electrochemical Analysis

Section dates back to the late 1930 's when a group was

organized by the late S. F. Acree to investigate the electrochemistry of solutions, in particular the interactions of

acids and bases. When the establishment of a national

standard of acidity later became a major concern of the

group, the name "pH Standards Section" was adopted. This

title gave way to "Physical Chemistry Section", and, as the

study of solvent effects on electrolyte processes became

the primary objective about i960, the name "Solution

Chemistry Section" came into being. With the establishment

of the Analytical Chemistry Division some three years later,

the present name was authorized, in order to emphasize the

place of the Section's activities within the frame-work of

the Division's responsibilities.

The mission of the Section has been summarized as

follows:

iii
"The Electrochemical Analysis Section conducts

experimental and theoretical investigations of

a J the properties of pure materials, especiallyweak and strong electrolytes, in aqueous and

nonaqueous media; b) reversible and transport

behavior at electrodes; and c) the application

of these properties and phenomena to the development and improvement of analytical processes.

"Specific programs include the establishment of

standard reference materials and reference data

for pure materials (dissociation constants, activity coefficients, conductivities, resistivities,

electrode potentials, dielectric constants, rate

constants, pH standards), measurement scales (pH

and other acidity functions for the precise estimation of hydrogen ion), and analytical methods

of improved sensitivity and accuracy."

A large share of the Section's efforts is directed

currently toward an understanding of the behavior of acids,

oases, and electrolytes in solution; to the measurement of

the thermodynamic, transport, and kinetic properties of

pure materials in solution; and to the electrical properties of pure liquids and solids. These fundamental investigations are designed to lead to a) the solution of measurement problems through the development of new and improved

standard reference materials, b) the collection of useful

data for pure materials, and c) the development and maintenance of a high competence in special areas of analysis within

the purview of the Section.

iv
It is the purpose of this report not only to review

the individual projects of the Section but to convey as well

an impression of the interrelationships of the separate

activities as they fuse into a single Section program.

The first goal could be achieved in a most satisfactory way

by collecting together the published -- or soon to be published — work of the staff as listed at the end of this

document. The second aim, however, is more elusive. It

can only be met by an integrated summary of the total

Section effort, where accomplishment can be viewed against

the backdrop of the mission, facilities, and personnel of

the organizational unit.

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 and Mrs.

Marguerite Raudenbush In the preparation of this report is

gratefully -acknowledged.

Roger G. Bates, Chief

Electrochemical Analysis Section

Washington, D.C.

June 28, 1965

v
TABLE OF CONTENTS

PAGE

1. FACILITIES AND EQUIPMENT . 1

Emf Methods . 1

Spectrophotometric Methods. ... 2

Isopiestic Vapor Pressure Method 4

Resistance and Impedance Methods. ..... 4

Conductance of Solutions 4

Dielectric Constants of Liquids

Resistivity of Metals 7

Miscellaneous Special Methods 7

2. DESIGN OF AUTOMATED INSTRUMENTATION FOR EMF

MEASUREMENTS 8

5. MEASUREMENT OF ACIDITY 11

pH Standards 11

Redetermination of the pH of Potassium

Hydrogen Phthalate, Primary pH Standard 12

Objectives 12

Procedures 14

Results 14

Standards for pH* in Methanol-Water Solvents 16

Interpretation of pH(x) in Nonaqueous Media 17

Reference Solutions for 50 Wt . Percent

Methanol 19

The Utility of Correction Factors ... 19

Procedures 20

vi
TABLE OF CONTENTS (Con't)

PAGE

Results o 22

Consistency 'of the pH* Scale In 50-Percent

Methanol 23

Standards for pD in Deuterium Oxide 25

Establishment of a pD Scale 26

Phosphate Buffer as a Standard Reference

Solution 27

Buffer System DAc (0,05 m) , NaAc (0.05 m)

in D

2 30

Future Programs 31

Standard Emf of the Cell: Pt;H

p (g.), HI,

Aglj Ag from to 50 °C . . . 33

4. INDICATORS AS REFERENCE BASES FOR ACID-BASE

STUDIES IN INERT SOLVENTS 36

The Need for New Reference Bases 36

Indicators of the Nile Blue Group 38

Structures and Terminology 38

Preparation of the Indicators 39

Determining Relative Strengths of Acids in

Inert Solvents by Spectrophotometry .... 40

Relative Acidic Strengths in Acetonitrile

Using Phenyl Diethyl Nile Blue Base ... 4l

Relative Acidic Strengths in Benzene Using

Tolyl Dipropyl Nile Blue Base

^

Nile Blue Oxazone 45

Solvatochromism vs. Halochromism 45

Nonaqueous Titrations Using Indicators of the

Nile. Blue Group 47

vii
TABLE OF CONTENTS (Con't)

PAGE

Comparative Acidities of Indicators in

Water and 50 Wt . Percent Methanol .... 49

Comparative Acidities of the Indicators in

4:1 Benzene-Methanol 49

5. SOLVENT EFFECTS ON ACID-BASE PROCESSES OF

ANALYTICAL INTEREST 50

Behavior of Cation Acids in Alcohol-Water

Solvents 50

Dissociation of Tris(hydroxymethyl)-

aminomethane in 50 Wt. Percent Methanol. . 51

Dissociation of Ammonium Ion in Methanol-Water

Mixtures 54

Nature of the Solvent Effect on Cation Acids 55

Isotope Effect in the Dissociation of Acetic Acid

and Dihydrogen Phosphate Ion in Heavy Water. 57

Improved Spectrophotometric Techniques; Design

and Construction of a Controlled-Temperature

Cell Block „ . 61

6. AQUEOUS SOLUTIONS OF MIXED SALTS 64

Methods and Procedures ; 64

Results 67

7. BEHAVIOR OF SODIUM-RESPONSIVE GLASS ELECTRODES 70

Method of Application 70

Procedures _ 1

Results

8. CONDUCTOMETRIC DETERMINATION OF TRACES OF WATER

Apparatus 76

Results .0 ~f

viii
TABLE OF CONTENTS (Con't)

PAGE

Prototype Instrument 78

Water Content of Carrier Gas 78

Calibration 79

Fundamental Data 79

Materials 79

Conductances . . 79

Revision of Instrumentation • . . . . 8l

9. REFERENCE MATERIALS FOR DIELECTRIC MEASUREMENTS 84

Tertiary Butanol ........ 84

Benzene „ . . 85

Dimethylsulfoxide-Water Mixtures 88

10. KINETIC METHODS OF ANALYSIS 89

Experimental . . 90

Preliminary Results . ...... 9°

11. PERSONNEL 93

12. PUBLICATIONS AND MANUSCRIPTS, JULY 1964 to

JUNE 1965 o . 9^

13. TALKS, JULY 1964 to JUNE 1965 97

14. LIST OF REFERENCES . . . . „ . . 99

IX
LIST OP FIGURES

FIGURE NO, PAGE

1. Equipment for the precise electromotiveforce measurements on which pH standards

are "based . . . . ......... 15

2. Equipment for the emf measurements on which

pH* standards in methanol-water solvents

are based .................. 21

5. Emf cell for studies of solutions in

deuterium oxide ............... 28

4. General formulas of indicators of the Nile

Blue group. Unless otherwise noted,

R 1 = C^ 38

5. Spectrophotometric curves obtained with

Phenyl Diethyl Nile Blue base 42

6. Comparative reactivities of acids in benzene,

as measured by their tendencies toward association with Tolyl Dipropyl Nile Blue Base:

Trichloroacetic acid, diphenyl phosphate, and

trinitro-m-cresol 43

7. Comparative reactivities of acids in benzene,

as measured by their tendencies toward association with Tolyl Dipropyl Nile Blue Base:

Phenylpropiolic, salicylic, m-nitrobenzoic,

phenoxyacetic, and o-chlorobenzoic acids ... 4-5

8. Solvent effects on absorbance of Nile Blue

A oxazone .................. 46

9. Equipment for emf studies in deuterium

oxide solutions 58

10. Controlled-temperature cell block for

spectrophotometric measurements ....... 63

11. Isopiestic vapor pressure apparatus,

showing the silver dishes and the copper

block in the desiccator 65

12. Equipment for studies of the behavior of

sodium-responsive glass electrodes ..... 72
LIST OP FIGURES (Con't)

13. Conduc tome trie cell for the determination

of small amounts of water in a gas stream .

14. Circuit designed to compensate for the

non-linear relationship between conductance

and amount of water

15. Three- terminal cell for measurements of

dielectric constants

16. Capacitance-conductance bridge assembly

and constant-temperature baths for precise

measurements of the dielectric constants

of liquids

17. Spectrophotometer used to measure the rate

of the reaction between Alizarin Red and

hydrogen peroxide

PAGE

76

82

87

88

91

LIST OP TABLES

TABLE NO.

1. Standard electromotive force of the

hydrogen-silver chloride cell in methanolwater solvents at 25 °C . . . .

2. Values of pH*(S) in 50-percent methanol

at 25 °C .

3. Values of E* for the deuterium- silver

chloride cell in deuterium oxide from

5 to 50 °C

4. Standard reference values of pa-p for the

solution KDpP04( 0.025 m) Na^DPOj^ 0.025 m)

in heavy water .......

5. Standard reference values of pap, for the

solution CH^COOD (0.05 m), CH^COONa (0.05 m)

in heavy water ..........

6. Standard electromotive force (E°) of the

cell HQ

j Hl(m), Agl; Ag from to 50 °C,

in V o . . xi

PAGE

22

24

27

30

31

34
LIST OF TABLES (Con't)

TABLE NO. PAGE

7. Activity coefficients of HI at 10, 25, and

40 °Cj L2 of HI at 25 °C 35

8. Relative acidities of some indicator dyes in

water and nonaqueous solvents. . . 48

9. pK and related thermodynamic quantities for

the dissociation of protonated tris(hydroxymethyl)amlnomethane in 50-percent methanol

from 10 to 40 °0 .... - 53

10. pK values for the dissociation of ammonium

ion in methanol-water solvents at 25 °C. . . 54

11. The dissociation constant of acetic acid and

the second dissociation constant of deuteriophosphoric acid in deuterium oxide from 5

to 50 °C 60

12. Thermodynamic quantities for the dissociation of acetic acid and phosphate ion in

water and in deuterium oxide ........ 6l

13. Comparison of observed and calculated

responses of a sodium-sensitive glass

electrode at 25 °C 74

14. Effect of water on the conductance of

solutions of sulfuric acid in propionic

acid 80

15. Conductance of water and 0.209 M sulfuric

acid in propionic acid at 25 °C (Cell

constant: O.OOIO75) 8l

16. Dielectric constant of samples of tertiary

butyl alcohol at 25 and 30 °C; conductivity at 30 °C 85

17. Effect of drying benzene on the dielectric

constant at 30 °C

xii
ELECTROCHEMICAL ANALYSIS: STUDIES OF ACIDS, BASES,

AND SALTS BY EMF, CONDUCTANCE, OPTICAL, AND

KINETIC METHODS

July 1964 to June 1965

Edited by Roger G. Bates

ABSTRACT

This survey of the activities of the Electrochemical

Analysis Section, Analytical Chemistry Division, covers

the period July 1964 to June 1965. 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. Consequently, a description of facilities and equipment Is

presented first and is followed by brief summaries of the

several lines of work currently under way. Emphasis is

given to the reasons why each study was undertaken. The

main areas of investigation include the study of acidity

and solvent effects in water, deuterium oxide, and watermethanol solvents by emf methods; the development of reference indicator bases for nonaqueous media and reference

materials for dielectric measurements; measurement of the

thermodynamic properties of mixed salt solutions; and inxiii
vestigation of special problems In trace analysis by conductometrie and kinetic methods. The survey concludes

with lists of the members of the Section staff, publications

and manuscripts produced during the year, and talks given

by the staff.

xiv
1. FACILITIES AND EQUIPMENT

The Electrochemical Analysis Section occupies nine

laboratories located in the east wing, fourth floor, of the

Industrial Building. With the exception of two rooms maintained at constant temperature to assure the proper functioning of special equipment, all of the laboratories are

of the "general-purpose" type.

The specialized equipment maintained and developed by

the Electrochemical Analysis Section to support its mission

within the Institute for Materials Research and the Standard

Reference Materials program may be suitably classified in

terms of the basic measurement methods involved. This is

particularly so since considerably different methods are

frequently used to obtain data from which comparable values

of physicochemical constants or parameters for systems of

analytical interest can be derived.

A. Emf Methods

Precise measurement of the electromotive force of

chemical systems using hydrogen electrodes or electrodes

responsive to other ions provide data, for standard reference

potentials, for standard reference values of the acidity or

basicity of buffer solutions, for dissociation constants,

and for the derivation of the thermodynamic quantities

associated with the processes studied.
The Section has three assemblies with which emf measurements with an uncertainty of 0.1 mV or less can be made

in aqueous and semi-aqueous media over the temperature range

to 60 °C. Correspondingly, the pH of reference solutions

can be determined to 0.003 unit with a precision of 0.001

unit or better. Considerable improvement has been made in

these assemblies during the past year, particularly in

regard to simplification of arrangement and operating procedures with improved convenience in bath-temperature

control. The latter objective was accomplished by the

installation of commercial temperature controllers utilizing

the proportional mode of control and thermistor sensors.

High-impedance electrometers are available for potential measurements to 0.05 mV and pH to 0.02 unit with the

glass electrode. A high-impedance electrometric apparatus

utilizing a recording potentiometer with useful sensitivity

of 0.02 mV is in the assembly and development stage. It is

intended for studies of the sodium-ion response of special

glass membranes.

B. Spectrophotometric Methods

Three assemblies for precise measurement of changes in

spectral absorption taking place in the near ultraviolet,

visible, and near infrared are available for study of the

properties of solutions of acids, bases, and salts in both
aqueous and nonaqueous media. When sufficient colorimetric

sensitivity exists, these instruments .are useful for both

quantitative and qualitative analysis for specific constituents or impurities.

These assemblies are built around the Beckman Model DU

and Optica spectrophotometers and the Cary Model 14 recording spectrophotometer. They permit measurement of transmission and absorption in the wavelength range 200 to 3,000 nm,

Under optimum conditions, the results may be accurate to

0.3 to 0.5 percent. A precision of 0.1 to 0.2 percent may

be attained in the range 220 to 800 nm, with considerably

less definition outside this interval. Sample-temperature

control to 0.1 °C can be achieved near room temperature.

The cell temperature in one instrument can be controlled in

the range 15 to 35 °C by the circulation of water from a

constant-temperature bath through a water- Jacketed cell

compartment designed in the Section. Another instrument

is being provided with an electrically-heated cell compartment the temperature of which is controlled with the aid of

a thermistor sensor. A considerable range of temperature

above room temperature should be obtainable.

The spectrophotometric assemblies are in use for the

determination of dissociation constants, the comparison

of relative acidic strengths of acids and bases in various
solvents, studies of rates of reaction, and for the development of methods for the detection of trace quantities of

specific ions.

C

Isoplestic Vapor Pressure Method

An independent approach to the determination of solvent

activity and solute activity coefficients has been provided

through construction of an isoplestic apparatus for the

measurement of the vapor pressure of various simple and

complex aqueous solutions at 25 °C. The method is simple

and elegant both in principle and application, and the

results obtained are in excellent agreement with those of

other thermodynamic methods, where a comparison can be made.

One assembly, consisting of constant-temperature bath,

vacuum vessels, copper blocks, and silver dishes, is available.

D. Resistance and Impedance Methods

1. Conductance of Solutions

The conductivity of solutions of electrolytes is a definite function of the mobilities and charges of the ionic

species present. It is therefore modified by changes in

the temperature, viscosity, and dielectric constant of the

medium. The factors affecting conductivity are not yet

completely understood. Nevertheless, useful information on

equivalent conductances of salts, conductances of individual

ions, dissociation constants of electrolytes, and other
quantities of both theoretical and practical importance

can be Obtained from the measurement of conductance as a function of concentration of electrolyte, temperature, and

solvent properties.

Conductance phenomena are extremely sensitive to changes

in ion concentrations and, hence, to the dissociation of the

electrolyte present and to the extent of such reactions as

neutralization and complexation. Alteration of the solvent

has a profound effect as well. For these reasons, conductance measurements can often serve as a useful tool in

chemical analysis.

A modified audio-frequency Jones bridge assembly with

suitable conductivity cells and a precisely controlled

constant-temperature oil bath are available for measurement

of the conductance of electrolytes in aqueous and nonaqueous

media. In addition, a precision capacitance-conductance

bridge assembly and oil bath is available for general measurement of electrolytic conductance. Under optimum conditions, conductivities can be obtained with either of these

two assemblies with a relative accuracy approaching 0.01

percent with a precision of about 0.001. percent over a temperature range of to 80 °C.

2. Dielectric Constants of Liquids

A knowledge of the dielectric properties of the solvent

is important in appraisal of the results of electrochemical
measurements and to provide useful information regarding the

structure of molecules and the condensed states of matter.

Precise dielectric and conductance data of sufficient definition can serve as standard reference data and, where suitable,

the materials may serve as accurate reference standards for

other properties of concern.

The precision capacitance-conductance bridge assembled

in the Section can be used to measure the static dielectric

constants of semiconducting and nonconducting liquids at

frequencies up to 100 kHz. Absolute measurements of the

dielectric constant can be made with a precision of 0.02

percent to 0.1 percent over a temperature range from -^0

to 1^0 °C.

A commercial transformer ratio-arm capacitance bridge

with a range of 10" to 10" ' farad is also available for

measuring the dielectric properties of low-loss dielectrics

at frequencies up to 100 kHz. In combination with suitable

three-terminal cells, it should permit the characteristic

dielectric constants of well-defined reference and research

materials to be determined with an accuracy of about 0.01

percent and dielectric loss to be determined to about 1

percent.
J>. Resistivity of Metals

The resistivity of pure metals and alloys, or its

reciprocal, the conductivity, is characteristic of a given

metal composition, "both in respect to its magnitude and to

its dependence on temperature. As part of the Standard

Reference Materials program, a. Kelvin double bridge and

accessory items are being acquired and assembled for use in

characterizing metals and alloys through the resistivity

and its temperature dependency.

4. Miscellaneous Special Methods

A sensitive dilatometer and associated apparatus exists

for kinetic studies by an incremental volumetric method.

Determinations of the volumetric change with a sensitivity

7 approaching 2 parts in 10' has been achieved and the method

has been used with considerable success over the temperature

range 15 to 40 °C

.

A magnetic densimeter has been constructed for measurements of the density of pure liquids and solutions. It

utilizes a closed system, and as little as 15 ml of sample

is required. Operation of the instrument consists in finding the point at which a magnetic float is in buoyant equilibrium. The adjustment is accomplished by the calibrated

regulation of an electromagnetic field. When suitable

liquids are used, the densities obtained are estimated to

be accurate to 0.0001 with a precision of 0.00005.

(C. G. Malmberg)

7
2. DESIGN OF AUTOMATED INSTRUMENTATION EOR

EMF MEASUREMENTS

The design of an automated system for control, measurement, and data-handling In emf studies has "been completed.

This assemblage of equipment will make all adjustments and

record all data concerned In a series of emf measurements,

Including time, barometric pressure, temperature, and emf.

In addition, It will control accurately the temperature of

the bath In which the cells are Immersed, will decide when

equilibrium has been reached in all of six cells Cor any

specified fraction of them), and will automatically change

the bath temperature according to a flexible prearranged

program. At the end of each run, the system will make

available all of the pertinent data on a suitable print-out.

Thus the new instrumentation will perform all of the functions that a human operator must now perform after the

cells have been prepared, filled, and placed in the bath.

Moreover, such a system will have several advantages

over human operation. Because it can operate unattended

around the clock and over weekends, It will reduce the time

to complete a run to about one-third of that now required,

even when allowance is made for a 20-percent down-time

factor. Because the system will make decisions on the

basis of preprogrammed criteria, it will do away with errors

8
of human judgment in arriving at a reliable conclusion as

to when equilibrium has been reached. Each run will be

completed in the shortest possible time, and consequently

the system will make the most economical use of costly

reagents (such as deuterium gas, for example) and will enhance the accuracy of data obtained for materials whose

stability in solution is limited. Finally, because it will

free the operator from the tedium of routine measurements,

it will give him more time for the analysis of data and for

the design of experiments which promise to yield the greatest possible amount of useful information.

The proposed instrumentation has considerable built-in

flexibility and is readily adaptable to the performance of

tasks other than routine emf measurements. At the heart

of the proposed system is a computer and teleprinter-punch

unit which will not only perform all of the data-logging

and computation required for precise emf measurements but

will be available for use on other routine computation jobs

originating in the Section and Division, once programs for

such tasks have been developed.

The sensors consist of a digital thermometer, a digital

barometer (uncertainty of 0.01 mm in the range to 800 mm),

a digital voltmeter (accurate to 0.001 percent of full scale

from -1.1 to +1.1 V), and a digital clock. These sensors

are independent units and can be used individually or in
combination in any type of physical experiment which can

utilize them. The data can be handled according to any

program which is within the capability of the computer.

Accordingly, if the emphasis of the work within the

Section should change, or if it should prove desirable to

use the system for more than one type of routine physical

measurement, the probability is high that the system could

easily be adapted to the change.

All of the major elements of this assembly are commercially available and have a. service history which is

considered acceptable. The total cost of the system would

be approximately $55,000. This includes a 20-percent allowance for contingencies and 6 man-months of labor at $10 per

hour for assembly, interfacing, and initial programming.

The basic design of the automated equipment was worked

out by E. W. Hogue of the Measurement Engineering Division;

Mr. Hogue also selected the components for the assembly.

The design was later modified by D. R. Boyle of the

Information Technology Division so as to make use of

computer control.

(R. Gary)

10
3. MEASUREMENT OF ACIDITY

The Section has had primary responsibility over a period of 25 years for the development of a national standard scale of pH and for the selection and certification of

suitable reference materials to provide experimental realization of this scale. The NBS scale utilizes the unit

-log a-rr, where a^ is the conventional hydrogen ion activity

in each of the selected reference solutions. It has been

endorsed, along with the closely-related British Standard

scale, by the International Union of Pure and Applied

Chemistry. Data provided by NBS research projects provide

the experimental basis for the standard pH measurement

procedures approved for use not only in the United Kingdom

but also in Japan and on a less formal basis in most other

industrialized countries.

The major research effort of the past year has been

directed toward the definition of scales suitable for

measurements of acidity in partially aqueous solvents and

amphiprotic media in general. Special attention has been

directed to 50 wt. percent methanol and to deuterium oxide,

and one or more standard reference solutions for each of

these media have been established.

A. pH Standards

The search for new and better standard reference materials for pH measurements is a continuing activity. In

11
addition, there are continual "maintenance tasks". These

include l) the examination and certification of new lots

of materials "before issuance, 2) the recheck of experimental data and recalculation of the reference data as

"best" values of the natural constants are altered, and

J>) the intercomparison of reference points to obtain a realistic evaluation of the accuracy with which the "practical"

experimental pH scale can be fixed.

1. Redetermination of the pH of Potassium Hydrogen

Phthalate, Primary pH Standard ,

a. Method . Standard values for pH measurements,

termed pH(s), are assigned to the NBS primary standard buffers by determination of the acidity function p(aHYr

,-

| ) for

a particular concentration of the buffer with different

small concentrations of added soluble chloride. This function is derived from measurements of the emf of hydrogensilver, silver chloride cells without liquid junction. The

calculation utilizes the measured electromotive force, E,

standard potential of the cell, E°, and other known quantities:

( \ (E - E°) F n fl x p(Wri ) = — "— + log mri (1) h 01 RT In 10 U1

in which y is the individual ionic activity coefficient on

the molal scale, F is the Faraday, and m is molality. The

function p ( a,/^ -, ) ° is evaluated by extrapolation of P( afjYr<-i)

12
to zero concentration of chloride, and pa„ is calculated

by means of the equation

paH

- p(aHYcl

)° + log yc1 (2)

The Bates-Guggenheim convention Is used to evaluate

the Individual ionic activity coefficient needed to obtain

•pa-rr. This convention is based on the formula

AI1/2

log y

= Y72 ^) 01 1 + 1.5 I

'

where A is a constant of the Debye-Huckel theory and I Is

the ionic strength. The final step in the assignment of a standard reference value to the buffer solution is to Identify pa^ with pH(S) In the operational definition of the

experimental pH value.

The partially completed study described in this section

is part of a continuing project for the maintenance of the

NBS standard pH scale, since the 0.05 m solution of potassium

hydrogen o-phthalate Is widely used for the standardization

of pH assemblies. The main objective of the study is the

re- examination of the pH of phthalate solutions, and the

secondary objective or "by-product" is a redetermination

of K-jKp, the product of the first and second dissociation

constants of o-phthalic acid. Values of p(a-HYp-) ) are being

determined at intervals of 5 degrees from to 60 °C for

13
five concentrations of potassium acid phthalate (NBS Standard

l85d) in the range 0.01 to 0.1 molal with three different

concentrations of added potassium chloride.

b. Procedures . The usual procedures used in this

laboratory for precise electromotive force measurements have

been followed, with a few exceptions. It is our usual practice to use palladium- coated hydrogen electrodes in phthalate solutions to avoid reduction of phthalate by platinum

black. Better agreement between the pair of hydrogen electrodes in a cell was obtained at temperatures above 35 °C

when the electrodes were conditioned, for a day before use,

in oxygen-free 0.01 M hydrochloric acid. The cells were

measured first at 25 °C, then from to 35 °C in five-degree

intervals in one continuous series. A different set of

cells was prepared and measured at 25 °C, then from 40 to

60 °C. Final checks were always made at 25 °C. The general

arrangement of equipment -- precision potentiometer, cell

rack, and constant-temperature bath is shown in figure l.

c. Results . At least four values of p(aTjYp-) ) have been obtained for each of the five concentrations of

potassium hydrogen phthalate with three concentrations of

added potassium chloride and at 13 temperatures from to

60 °C. The final least-squares treatment and extrapolation

of the data are still incomplete. However, provisional

results indicate that the pH(S) of 0.05 m potassium hydrogen

14
it

-I'feflfe!

*

Figure 1. Equipment for the precise electromotive-force

measurements on which pH standards are based.

phthalate at 25 °C is within 0.001 pH unit of *the value of

4.008 assigned to this important reference solution on the

basis of experimental data obtained nearly 20 years ago.

The provisional data and calculations for 25 °C are

as follows:

0.05 m KH Phthalate, 0.015 m KC1 p(aHYcl ) = 4.0867

0.05 m KH Phthalate, 0.01 4.0896

0.05 m KH Phthalate, 0.005 4.09^4

0.05 m KH Phthalate, p(aHYcl

)°= 4.0966

-log yc1 (eq.2) = O.O876 paH

(eq.2) = 4.009

(H. B. Hetzer)

15
B. Standards for pH* In Methanol"Water Solvents . The major purpose of this phase of the Section's activity has been the development and application of an acidityscale in amphiprotic solvents (media capable both of accepting and donating protons) containing water mixed with an

organic liquid. Solvent mixtures of this sort are often

used in research and industry to assure the solubility of

slightly-soluble reactants or products. Without an acidity

scale for these media, precise control of analytical and

synthetic processes is impossible. A mixture of equal parts

by weight of methanol and water was chosen as the first such

medium for careful study, in view of its general usefulness

in chemical and biochemical studies.

Operational pH numbers for solutions in many nonaqueous

media can be measured with a pH meter using hydrogen (or

glass) and calomel electrodes. This operational pH scale

is defined as

PH(X) " PH(S) + (RTlnlO)/F

' (1)

where pH(S) is the reference value for suitable aqueous

reference solutions used to standardize the pH meter. In

eq. 1, E

x

and E„ are values of the emf of the pH cell with

the electrodes immersed either in the standard (S) or in the

unknown solution (x). The procedure for assigning values of

pH(S) to aqueous standards has been described in section 3.

A

of this report.

16
1. Interpretation of pH(X) in Nonaqueous Media

The pH(X) determined by eq. 1 with aqueous standards

will often be a reproducible number even when the solvent is

not wholly aqueous. Thus it fulfills the primary requirement of an acidity scale for regulatory and control purposes.

Abandonment of aqueous reference solutions in favor of reference solutions in the mixed solvent can be justified only if

the interpretation of the data furnished and their correla -

tion with acid-base phenomena are significantly improved.

It was shown recently [1] that the potential across the

liquid junction: solution X in alcohol-HpO| concentrated KC1

(aq. ) appears to be substantially constant even though the

acidity of solution X changes over wide ranges , provided that

the composition of the alcohol-water solvent remains unchanged,

Furthermore, glass electrodes have been found to give nearly

the theoretical response to hydrogen ions in alcohol-water

mixtures when the water content is not too low. A consequence of this discovery is that the difference of EL - E^

obtained under optimum experimental conditions may be a valid measure of the relative hydrogen ion activity of the

two media, or of pH(X) - pH(S) as expressed in logarithmic

form in eq. 1. These optimum conditions for mixed solvents

are, as in water, that the acidity of the solution be neither too high nor too low, that the ionic strength be moderate

(<0.2), and that the solute consist of simple buffer acids,

17
bases, and salts. The assignment of a valid and useiu.

value to one or more reference solutions is then sufficient

to define a new pH scale for that particular medium.

The unit designated pa*, and defined by

pa* = -log i^.^ (2)

can be obtained by procedures similar to those on which

aqueous pH standards have been based, utilizing a convention entirely analogous to eq. 3 of section 5.A.I. In

eq. 2, the asterisk and the subscript s indicate that the

quantity a— or y^ is referred to the standard state in the

mixed solvent (s) rather than to that in water (w) . It is now

possible to set up an operational pH* scale for an amphiprotic or mixed solvent as has been done for pH in aqueous

media:

E — E

PH*(X) = PH*(S) + " (3)

(RT In 10)/F

For the reference solutions selected, the pH*(S) is

taken to be equal to pa£ as determined by cells without a

liquid junction. Under the optimum or "ideal" conditions,

pH*(x) will then also have the dimensions of a conventional

hydrogen ion activity or pa.*- in the particular solvent

n

medium of concern.

18
2. Reference Solutions for 50 Wt. Percent Methanol

a. The Utility of Correction Factors , In work

performed prior to the reporting period [1], it was shown

that the difference between the pH(x) determined with aqueous

.standards and the pH*(x) determined with respect to reference

solutions in an alcohol-water medium is nearly a constant

under optimum experimental conditions. This constant was

shown to be a combination of liquid- junction potentials and

the medium effect (or transfer energy) for the hydrogen ion

between the two solvents , water and the alcohol-water mixture.

The value of the constant could be determined either from

suitable emf measurements [1] or from pH- titration curves [2],

The derivation of pH* by the application of correction

terms to pH values found with an assembly standardized in

the usual way with aqueous reference solutions is undeniably

attractive. It may be dangerous, however, for the asymmetry

potential of the glass electrode may vary when the electrode

is removed from one solvent and placed in another of widely

different water activity. The safest procedure is to use

reference solutions of the same solvent composition as the

"unknown" or test solutions.

For this reason, pa£ was determined for four buffer

systems in 50 wt. percent methanol from 10 to 40 °C and over

a range of buffer concentrations by the same methods used to

obtain pa.^ for reference solutions in water. The procedure

involved three steps, l) the determination of E°, the

19
standard emf of the hydrogen- silver chloride cell for the

solvent 50-percent methanol, 2) the measurement of the

acidity function p i^wYn-i) for the reference solutions, and

j) the calculation of pa* from P s( axjYr;i)

^ application of

a chosen convention.

b. Procedures . The four buffer systems selected

were equimolal mixtures of

1) acetic acid, NaAc, and NaCl

2) sodium hydrogen succinate and NaCl

J>) tri s(hydroxymethyl)aminome thane and its hydrochloride

4) KH

2

P04

, NagHPO^, and NaCl.

The chemicals used for the experiments were NBS

Standard Reference Materials or reagent-grade materials,

and some were further purified. "Spectro grade" methanol

was used. Glacial acetic acid was purified by partial freezing and its purity checked by gas chromatography. The succinic acid and sodium succinate from which the sodium hydrogen succinate was prepared were assayed, and the purity of

the tri s(hydroxymethyl)aminomethane was also established by

titration with a standard solution of distilled hydrochloric

acid.

Appropriate stock solutions of the buffers were prepared in water. These were diluted with water and methanol to

yield the desired concentration of the buffer in a 50-percent

methanol-water medium. The ionic strength of each buffer

system was varied from 0.01 to 0.1.

20
The emf of the cell

PtjH

2 (g., 1 atm), buffer in 50-percent MeOH, AgCl; Ag

was measured over the range of temperatures from 10 to 40 °C

at five-degree intervals. In order to obtain the -standard

emf ( E°) of this cell, a series of measurements was also

made from 10 to 40 °C in which the buffer was replaced bysolutions of hydrochloric acid in 50-percent methanol.

Further data at 25 °C in 10, 20, 45, and 70 wt . percent

methanol were also obtained. The potentiometer and associated measuring equipment are shown in figure 2.

Figure 2. Equipment for the emf measurements on

which pH* standards in methanol-water

solvents are based.

21
c. Results . The standard emf ( E°) of the cell

was derived by accepted thermodynamic methods from the

data for cells containing solutions of hydrochloric acid.

The methods and the results have already been described

[3,4]. The standard emf at 25 °C is summarized in table 1,

Table 1. Standard electromotive force of the hydrogensilver chloride cell in methanol-water

solvents at 25 °C.

wt. Percent Methanol E°, V

s '

0.222^4

10 0.21549

20 0o 20901

45 0.19414

50 0.19067

70 O.I6835

A knowledge of the standard emf permits the acidity

function p (aTTY~,-, ) to be calculated for each of the reference

•"s v H'Cl'

buffer solutions for which emf data in methanol-water media

were obtained. The calculation is similar in every way to

the derivation of p(aFr Yr

, 1) in aqueous solutions:

E - E°

P

s (a H

Tcl ) = "log gtV^cl) = log m^

(RT In 10)/F 01

22
By adopting a formula (convention) analogous to that

on which aqueous pH standards are based (eq„ 3 of section

3.A.1), yri could be evaluated and, hence, pag values

obtained from the experimental acidity functions:

pag = Psfa^Cl 5 + l0S s^Cl (5)

The convention was modified to allow for the difference

between the dielectric constants of the two solvents (water

and methanol-water) and for the differences in the densities. The values of pag obtained by eq. 5 were identified

with pH*(s) in the operational definition of a practical

scale of pH* (eq. j5)»

A detailed description of the work, except for the

measurements on buffers of tris(hydroxymethyl)aminomethane

("tris"), has been published [5]- The standard reference

values in 50-percent methanol at 25 °C are collected in

table 2. In addition to their application to the standardization of pH assemblies, these values are useful in the

spectrophotometric determination of the dissociation constants of weak acids and bases in 50-percent methanol.

d. Consistency of the pH* Scale in 50-Percent

Methanol . The internal consistency of the pH* scale defined

by the four buffers in 50-percent methanol was examined by

measurements of the cell

Pt;H

2 (g., 1 atm), buffer (1) in 50^ MeOH | S |^'| buffer (2)

in 50^ MeOH, H

2 (g., 1 atm); Pt

23
Table 2. Values of pH*(S) in 50$ methanol at 25 °C.

Buffer

Ionic

Strength Acetate Succinate Tris Phosphate

0.01 5.591 5.860 7.897 8.232

0.02 5.568 5.818 7.929 8.157

0.03 5.552 5.785 7.952 8.102

0.04 5.5^0 5.757 7.970 8.058

0.05 5.529 5.734 7.985 8.021

0.06 5.520 5.714 7.998 7.988

0.07 5.512 5.698 8.010 7.959

0.08 5.505 5.684 8.020 7.932

0.09 5.499 5.674 8.029 7.907

0.10 5-493 5.666 8.038 7.884

The assigned pa-g values for the acetate and succinate buffers were found to be completely self-consistent, whereas

the phosphate and tris buffers showed anomalies. With tris

buffers, the discrepancy was sometimes as large as 0.05 pH

unit. The residual liquid- junction error which causes the

discrepancy seems to depend on the amount of chloride present in the buffer solution and also on the valence type of

the buffer. More experimental work will, it is hoped,

reveal a practical solution to this problem and lead to

more accurate pH* measurements than are now possible.

(M. Paabo and R. G. Bates)

24
C . Standards for pD in Deuterium Oxide

In an earlier section of this report, it was shown

that standard reference solutions for measuring the practical pH can be provided for some solvents other than water.

Since the assignment of pH values to a reference standard

requires a considerable amount of careful work, the choice

of solvents to receive first attention should be a judicious

one.

There are a variety of reasons why it was felt that a series of reference standards for deuterium oxide would be

valuable to the scientific community. From the practical

point of view, deuterium oxide has moved rapidly from the

status of a laboratory curiosity in 1930 almost to the status

of a commercial solvent. It is found in ton quantities in

many atomic energy installations where it is used as a coolant and a pile moderator, and where, incidentally, corrosion and the variables which control it are a constant

concern. It is an indispensable tool of the biochemist who

uses it in tracer studies of complex biochemical reactions

and processes. Heavy water will no doubt find increasing

application as it continues to become still more readily

available.

, The "acidity" measured in deuterium oxide is not the

activity of hydrogen ion, aH

, but the activity of deuterium

ion, a^ . And, similarly, while the practical pH scale for

HpO is designed to reflect aH

, an analogous scale in D

p

25
should reflect a^. The experimental realization of this

scale is called a practical pD scale.

1. Establishment of a pD Scale

The procedure for establishing such a scale parallels

that used in constructing the practical pH scale. Values of

the quantity p(a-DYcl ) are computed from emf measurements of

the cell

Pt;D

2 (g.), Reference solutions , MCI in D

2

0, AgCl; Ag

in which the Pt;D

2 (g. ) electrode replaces the Pt;H

2 (g.) electrode used for studies in ordinary water. The emf of the

cell is given by the equation

E = E° - k log(a

D

m

clYcl ) (l)

where E° (the "standard emf") is a constant characteristic

of the cell reaction at a given temperature and at unit

activity of the reactants and products, and k is written

for (RT In 10)/F.

The standard emf is, of course, different from the E°

which pertains to the case of hydrogen. It therefore had

to be evaluated before meaningful measurements could be made

on the reference buffer systems. The evaluation was made

from emf measurements of the cell

Pt;D

2 (g.); DCl(m) in D

2 0, AgCl; Ag

at 5- or 10-degree intervals from 5 to 50 a C.

26
Details of the work have been published [6]. The values of

E° are summarized in table 3.

Table 3. Values of E° for the deuterium-silver chloride

cell in deuterium oxide from 5 to 50 C.

t, °c E% V.

5 0.22528

15 0.21931

25 0.21266

35 0.20532

45 0.19733

50 0.19310

These measurements, as well as those taken subsequently on

the reference solutions _, were made in a cell of new design

(figure 3)

.

2. Phosphate Buffer as a. Standard Reference Solution

In choosing solutions to serve as possible reference

standards for pD, the standard materials already used for

the pH scale were considered to be prime candidates. Standard pD(S) values would be determined for them according to

the usual NBS procedures.

The first system to be studied was the equimolal phosphate buffer system [7]. Measurements of the emf of the

27
Figure J. Emf cell for studies of solutions in deuterium

oxide.

28
cell Pt;D

2 (g.), KD2 P0

Z| (0.025 m) , Na^DPO^O. 025 m) ,

NaCl(m'), AgClj Ag were made at 5-degree intervals from 5

to 50 °C and at several levels of m'. The emf (E) and the

standard emf (E° ) were used to compute values of the quantity

p(aDYG1 ) which were extrapolated to the limit of mri = 0,

where the quantity may be designated p(anYf

,-,) :

Pl^Td)^ (E - E°)/k + log m

ci

To obtain a value for -log aD

= pa^, the quantity yci

was evaluated from the expression

" l0

^

cl = ~^

1 + B' a I

'

This expression differs from that used in the case of ordinary water in that the constants A' and B' are different.

This is because the solvents differ in dielectric constant

and density. The ion-size parameter a was taken to be the

same in deuterium oxide as in ordinary water, namely 4.56 A.

Values of pa^ for the phosphate buffer system could

then be computed, and these are shown in table 4. The

practical significance of these values is that they can

serve as reference values pD(s), for the calibration of

equipment and procedures for measuring pD in other laboratories on an operational scale similar to that defined by

eq. 1 of section 3.B. The accuracy of such measurements is

29
Table 4. Standard reference values of pa^ for the

solution KDgPO^ (0.025 m) , Na

2

DP0^ (0.025 m)

in heavy water.

t, °c paD

t, °C paD

5 7.537 30 7.411

10 7.504 35 7.397.

15 7.475 4o 7.386

20 7.450 45 7.380

25 7.429 50 7.377

probably subject to the same restrictions and precautions

that pertain to pH measurements. Details of procedures for

the use of commercial equipment such as the glass electrode

to measure pD in heavy water will be the subject of further

study.

3. Buffer System DAc (0.05 m), NaAc (0.05 m) in D

2

This buffer system is not one of the five current NBS

reference standards for pH. It was chosen for study,

however, because it not only provides another point on the

new pD scale but also because the data obtained might

stimulate further work of theoretical importance by other

workers. Acetic acid has often been used as a reference

acid in the study of isotope effects.

30
The protium form of acetic acid, HAc, is a wellcharacterized material. The ionization of acetic acid in

water has been the subject of several classic studies by

well-known physical chemists.

In the determination of paD

for the solution CH COOD

(0.05 m), CKUCOONa (0.05 m) in heavy water, the same procedure was followed as for the phosphate buffer. The pa^

values obtained are shown in table 5 [8]. These values

can be identified with pD(S) in an operational scale for

pD measurement.

Table 5. Standard reference values of pa^ for the

solution CH,COOD (0.05 m) , CEUCOONa (0.05 m)

in heavy water.

t 9 c paD

t, c paD

5 5.265 30 5.227

10 5.252 35 5.225

15 5.245 40 5.226

20 5.235 45 5.230

25 5.250 50 5.235

4. Future Programs

One of the criteria for a useful acidity scale is that

it be internally consistent. This means that the difference

between the practical pD values for any two standards measured

31
individually in cells without liquid junction should

agree with the difference observed when they are compared

directly in the cell which is used for the overwhelming

proportion of practical acidity measurements. The two

systems already investigated, as well as others that will

be studied from time to time, must be checked against each

other in cells with liquid junction similar to the type

used commonly for practical pH measurements (see section

3>.B.2.d). This comparison has been initiated.

Since the glass electrode Is the most convenient and

widely used indicator electrode for hydrogen ion, the best

way of adapting it to the measurement of deuterium ion requires study. The response of the glass electrode to deuterium ion in D

?

has not yet been fully affirmed. The

ultimate availability of several pD standards will make possible this series of experiments, essential in order to

demonstrate the utility of the glass electrode in deuterium

oxide.

In developing the pH scale, it was important to determine the limits within which the assumptions made in the

formulation of the scale could reasonably be expected to be

valid. The pD range within which reasonable constancy of

the liquid- junction potential is maintained is not yet

known and should be determined.

(R. Gary)

32
D. Standard Emf of the Cell : Pt;H2 (g.)> HI, Agl; Ag from to 50 °C

The standard electromotive forces of the analogous

cells with hydrochloric acid and a silver-silver chloride

electrode [9] and with hydrobromic acid and the silver- silver

bromide electrode [10] have been determined previously. The

silver- silver iodide electrode has some advantages over the

other two, particularly its low solubility in ammonia and

amine solutions. This makes it a very useful electrode with

which to measure the dissociation constants of bases and

the pH values of buffer solutions of high alkalinity.

Measurements of the emf of the cell

Pt;H

2 (g.), HI (m), Agl; Ag

have been made [11] at 11 temperatures from to 50 °C, and

the standard emf (E°) evaluated at each temperature. The

values of E° found can be expressed as a function of temperature by the equation:

E° = - 0.15242 - ^.19 x 10~ 4 (t-25) - 2.84 x 10" 6 (t-25) 2

,

where t is the temperature in degrees Celsius. A comparison

with the earlier data of Owen [12], who used borax-buffered

potassium iodide solutions, is given in table 6.

53
Table 6. Standard electromotive force (E°) of the

cell H

2

j HI (m), Agl; Ag from to 50 °C

In V.

t, °c E°, this work E°, Owen [12]

-0.146^7

5 -0.14719 -0.14717

10 -0.14822 -0.14810

15 -0.14942 -0.14925

20 -O.15081 -O.15067

25 -0.15244 -O.1523O

30 -0.15405 -0.15401

35 -0.15590 -0.15591

40 -0,15788 -0.15792

45 -0.15998

50 -O.16219

The agreement is excellent at 5, 10, 30, 35, and 40 °C,

but there are differences of 0.14 to 0.17 mV at Intermediate temperatures.

Activity coefficients have been calculated up to 1 m

at 10, 25, and 40 °C and those at 25 °C compared with those

derived from isopiestic measurements [13]. The partial

molal enthalpies (L

p

) of these solutions at 25 C

C have also

been derived. The results are given in table 7* The thermochemical calorie, (equivalent to 4.1840 J) is used here.

34
Table 7. Activity coefficients of HI at 10, 25,

and 40 °C; L

2

of HI at 25 °C.

y± y± y± T± L

2

m (10 °c) (25 °c) (25 °c) a (40 °C) 1 -i-l cal raol

0.005 0.931 0.929 - O.928 41

0.01 0.909 0.907 - 0.904 60

0.02 0.882 0.897 - 0.876 77

0.05 0.843 0.839 - 0.835 106

0.1 O.815 0.811 0.798 0.806 130

0.2 0.799 0.793 0.788 O.788 170

0.3 0.803 0.795 0.792 O.789 202

0.4 0.814 0.805 0.803 0.798 226

0.5 0.831 0.820 0.819 0.813 244

0.6 O.851 0.839 0.839 O.832 263

0.7 0.874 0.861 0.862 0.853 281

0.8 0.900 0.885 0.886 0.877 297

0.9 0.927 0.911 0.913 0.903 315

1.0 - - 0.940 - -

a From isopiestic data

.

(H. B. Hetzer, R. A. Robinson)

35
4. INDICATORS AS REFERENCE BASES FOR ACID-BASE STUDIES

IN INERT SOLVENTS

A. The Need for New Reference Bases

Acid-base studies, such as the determination of relative acidic strengths and the determination of total acid

content in nonaqueous solvents, have been impeded by short-

.comings of presently obtainable reference bases. Ideally, a

series of pure crystalline bases of graded strengths should

be easily obtainable for chemical work in varied solvents,

and the series would include strong bases which could be

used without introducing water or alcohols. Two examples

may be given of the limitations of reference bases now in

common use:

1. Quaternary ammonium hydroxides or alkoxides,

used when a very strong base is needed, are not available in

a convenient form. Some are supplied as aqueous or alcoholic

solutions, and in other cases the base has to be prepared

from a halide. An additional disadvantage in their use is

the inevitable introduction of water or alcohol into otherwise anhydrous systems

2» Aliphatic amines, commonly employed when a

weaker reference base will serve, are liquids when of low

molecular weight. Solid aliphatic amines have been prepared

by introducing heavy alkyl groups, but the additional bulk

of these groups makes the nitrogen atom less accessible to

acidic molecules.

36
With concerted efforts, it should be possible to

develop better reference bases. For example, quinuclidine

(C_H_ N) appears to be superior In two respects to trl7 15

ethylamine (one of the most frequently used aliphatic

amines): It is a solid, and it is a little more reactive

(probably because the nitrogen is more exposed). However,

it is hard to obtain because an economical procedure for

its synthesis has not yet been worked out.

An obvious place to look for strong reference bases

is among compounds in which cation formation is favored by

resonance. The relatively high basicity of 1, j5-diphenyl~

guanidlne can be attributed to stabilization of the cation

form by resonance. Other useful solid bases should be found

among guanidines containing aralkyl and alkyl groups. Development work in this area is badly needed.

Indicator bases , another group of compounds in which

cation formation is favored by resonance, offer well-known

advantages. One is the high molar absorptivity which is

characteristic of these compounds, making them very sensitive reagents. Another advantage is that the principal

absorption bands ordinarily do not overlap absorption bands

of solvents and dissolved acids. Little developmental work

has been done in this area. This report summarizes exploratory studies using several indicators of the Nile Blue group,

37
B. Indicators of the Nile Blue Group

1. Structure and Terminology

When using acid-base indicators in non-ionizing solvents it is necessary to specify whether the indicator is

employed as the salt or as the corresponding base or acid.

Figure 4 shows the general structural formulas of the indicators under discussion, as the base (i) and as a salt (il)

formed by adding an acid HX to the base. A related indicator, known as an "oxazone", is also shown (ill).

rr* R^N'^^Nr^'^NR R ,

(I) BASE (II) SALT (ill) OXAZONE

Figure 4. General formulas of indicators of the

Nile Blue group. Unless otherwise noted,

R ' = C

2

H

5

.

In all but one of the indicators under discussion,

R' = C HC . The identities of R in the salts used and the

d o corresponding common names (in parentheses) were as follows:

(1) H (Nile Blue A); (2) C^ (Nile Blue BX); (?) CgH-CHg

(Nile Blue 2B); (4) p-ClCgH^CR^ (4-Chlorobenzyl Diethyl

38
Nile Blue); (5) CgH^ (Phenyl Diethyl Nile Blue). The common

names applied to the salts are used for the corresponding

bases, hut with the word "base" or "anhydro-base" appended.

On comparing the three structural formulas it can be seen

that the salts or bases just mentioned .are related to the

same oxazone ("Nile Blue Oxazone").

One of the anhydro -bases used differed from these

oompounds with respect to the nature of both R and R':

R = p_-CH„C/rHj. and R' = n-C^JHL. Its common name is

"p-Tolyl Dipropyl Nile Blue Base"

.

The "common names" mentioned have been adopted as

everyday substitutes for the very cumbersome systematic

terminology. As an example, the systematic chemical name

for Nile Blue A Base is "9-diethylamino-5-iminobenzo/a7-

phenoxazine"

.

2. Preparation of the Indicators

Ordinarily indicators of the Nile Blue group are prepared as the salts (II ) by condensation reactions (described

in literature on dyes), then converted to the bases (i) by

treatment with alkali. The oxazone (ill) can be prepared

either by a condensation reaction or by heating an aqueous

solution of the salt with an excess of acid. Purification

of the products is accomplished by techniques familiar to

the organic chemist.

39
Nile Blue A and Nile Blue 2B have long been supplied

commercially in impure grades. Pure materials were needed

for our investigations. Most of them were obtained by gift

and were samples of materials that had been synthesized for

possible medical applications [14],

C . Determining Relative Strengths of Acids in Inert

Solvents by Spectrophotometry

In comparatively inert solvents like benzene and

acetonitrile (which were the solvents of major interest),

the relative strengths of acids are measured in terms of

their comparative affinities for the reference base which

has been selected; see, for example, reference [15]. The

best numerical criterion is the value of the equilibrium

constant for 1:1 association of the base (symbolized by B)

and the acid (symbolized by HA). If the tendency toward

formation of BHA is small, overlapping association reactions

will probably occur -- for example, formation of (HA) p

or of B(HA)„. When the tendencies for such competing associations are unknown, a flexible program of measurements

must be adopted. In this work, a fixed concentration of

the indicator base (usually of the order 10_:) M) was mixed

with systematically varying amounts of the acid.

It is desirable, whenever possible, to continue the

experiment to the stage where complete conversion of base

to salt may be considered to have occurred. Entire absorp40
t'ion curves are obtained, so that any spectral shifts indicative of competing associations will not be overlooked.

When favorable experimental conditions have been employed,

association constants can be calculated from spectral data

and the stoichiometry. Only materials of high purity were

used, and precautions were taken to exclude moisture (even

though traces were not likely to affect the results of most

experiments)

.

1. Relative Acidic Strengths in Acetonitrile Using

Phenyl Diethyl Nile Blue Base

The type of spectrophotometric data obtained is illustrated in figure 5. The absorption curve marked "0" was

obtained for Phenyl Diethyl Nile Blue Base in acetonitrile

(8 x 10 -^ M). The other curves were obtained using solutions with the same concentration of base but differing

amounts of two monobasic acids -- hydrogen chloride and

anhydrous diphenyl phosphate. The curve obtained for the

solution which contained 4.0 equivalents of diphenyl phosphate (3.2 x 10 M) plus additional solid acid probably

corresponds closely to the "limiting" curve (the curve

obtained when all of the base has been converted to a salt).

In the experiments testing the reactivity of diphenyl

phosphate, known amounts of the solid acid were mixed with

the base. In the experiments testing the reactivity of

hydrogen chloride, known amounts of the base and of the

corresponding chloride were used.

41
25

Wave number, cm

-

' x IO

- 3

20

80

6 x I0"3

-i 1 1-

Phenyl Diethyl Nile Blue Base,

8xlO"^M in Acetonitrile, plus

400

HCI

Ph2HP04

15

Addnl. solid

-4.0 equivs._

/ 1.0 \

500 600

Wavelength, m^.

700

Figure 5. Spectrophotometric curves obtained with

Phenyl Diethyl Nile Blue Base.

These preliminary results justify the following

conclusions:

(1) The base and acid combined in 1:1 ratio. This conclusion is supported by the spectral evidence that salt formation was nearly complete on adding one equivalent of the

acid. Also, the sharp isosbestic point near 570 mu. (570 nm)

indicates the probable absence of competing associations,

such as formation of B(HA) .

(2) Hydrogen chloride and diphenyl phosphate do not differ

greatly in proton-donor ability.

42
2. Relative Acidic Strengths in Benzene Using

Tolyl Dipropyl Nile Blue Base

Figures 6 and 7 present results obtained in studying

the comparative reactivities in benzene of three strong

acids and five moderately strong acids. Trichloroacetic

acid,, diphenyl phosphate , and trinitro-m-cresol (figure 6)

associate fairly readily with Tolyl Dipropyl Nile Blue Base.

Trichloroacetic acid is more reactive than trinitro-m-cresol

in benzene ;. in keeping with its lower pK value in water.

AA PER CM,

AT 635 my

0.5 -

0.4-

0.3 -

0.2-

0.1 -

1 1 1 1 1 111

- ^^**^*^ _ CI3C.COOH

(pK=0.6KT'

^^

Ph 2 HP04 (pK = 0.3?)-

/ JT ^" TRINITRO-m-CRESOL " / / s^ (pK = 0.8)

<fr*/ -

- (l/f -

r 1 1 1 -1 1 2 4 6 8 10 12 CONCN. OF ACID, EQUIVS.

Figure 6. Comparative reactivities of acids in benzene _, as measured by their tendencies toward

association with Tolyl Dipropyl Nile Blue Base

Trichloroacetic acid,, diphenyl phosphate,, and

trinitro-m-cresol.

^
Diphenyl phosphate seems to be intermediate in strength.

The very low aqueous pK value given in parentheses,, taken

from the literature, was determined under unfavorable

conditions and is probably only a rough approximation to

the true value. The shapes of the curves in figure 6,

and also the well-defined isosbestic points which were

observed in the complete absorption curves, indicate that

the base is associated with these acids in 1:1 ratio in

the concentration range shown. Equilibrium constants were

calculated for the associations involving diphenyl phosphate

and trinitro-m-cresol: — the respective log K values ° assocn.

were 4.4 and 4.13.

The curves shown in figure 7 which were obtained

with the five moderately strong acids, have an S-shape that

is more evident when the data are plotted with an expanded

horizontal axis. With the possible exception of the reaction involving o_-chlorobenzoic acid, overlapping reactions

(probably including formation of both BHA and B(HA)

2 but

chiefly the latter) seem to have occurred under the experimental conditions used. The order of apparent strengths of

these five acids in benzene does not parallel exactly the

relative order in water (see aqueous pK values given in

parentheses). The presence of internal hydrogen bonding

in compounds such as o-chlorobenzoic acid is one explanation

for this.

44
n 1 1 1 1 r

CCI3COOH

0.6p( P K = 0.6)

AA PER CM,

AT 630 imp

0.4

0.2

PhC = CC00H (2.2) -

2000 4000 6000

CONCN. OF ACID, EQUIVS.

8000

Figure 7. Comparative reactivities of acids in benzene,

as measured by their tendencies toward

association with Tolyl Dipropyl Nile Blue Base:

Phenylpropiolic_, salicylic, m-nitrobenzoic , phenoxyacetic, and o-chlorobenzoic acids.

3. Nile Blue Oxazone

The oxazone differs from the two anhydro-bases just

discussed in (a) being a weaker base, and (b) possessing

a marked fluorescence in solution. This fluorescence is

quenched by adding a strong acid.

4. Solvatochromism vs. Halochromism

Hantzsch [16] coined the two terms "solvatochromism" and

"halochromism" to distinguish between two kinds of color

changes of indicator dyes. The first term refers to the

bathochromic or hypsochromic shifts of spectral absorption

curves caused by changing the solvent, and the second term

applies to color changes which accompany salt formation.

45
Halochromic shifts are illustrated in figure 5. Solvate--

chromlsm is illustrated in figure 8. The solvation shifts

Wove number, crrH x I0"3 25 20

exIO"3

40

30

20

10

300 400 500 600

Wavelength, m/i

Figure 8. Solvent effects on absorbance of

Nile Blue A oxazone.

were produced by adding water, in progressively larger

amounts, to dioxane solutions of Nile Blue Oxazone. Analogous shifts were observed on changing from a single solvent

such as benzene to more polar solvents, for example acetonitrile and methyl sulfoxide. The anhydro-bases referred

to above also exhibit solvatochromism, but the phenomenon

is more noticeable when the fluorescent oxazone is used.

46
Solvatochromism occurs very widely. So lv at o chromic

shifts have "been employed as "empirical parameters" of the

relative polarities of solvents; for example, see reference

[17]. Most compounds used in deriving such parameters are

much more solvato chromic than the Nile Blue group of indicator bases. However, it is apparent that well-marked

isosbestic points like those shown in figure 5 cannot be

expected if the solvent composition is allowed to change

in the process of salt formation.

This last point seems worth stressing, because analytical titrations and various non-spectrophotometric procedures

are often carried out with changing solvent composition.

Changes in the medium could be responsible for the puzzling

results sometimes obtained.

D. Nonaqueous Titrations Using Indicators of the

Nile Blue Group

Nile Blue Oxazone can be used as an indicator in

certain titrations with perchloric acid in acetic acid solution, but in general it is less satisfactory than Crystal

Violet. However, the fluorescence might be an advantageous

property in some applications.

Interesting results have been obtained in using some

of the Nile Blue type indicators (as the salts) in nonaqueous titrations. A representative sampling of results

is presented in table 8.

47
ri£

COPp

CO£P

•H

CO

CD!»

-O

Pho

-p

COoH

rt

P! «H

CD6O

CO

<Ho • CO CO p

CD p! •H CD P >

•HH

T3 o

•H CO o

co CO2

CD o > CD •H ?i P O*

cd cd H Pi

cd o P C

CO

CD rHp

cd EH PPP

pq o

rH P O P CD P H H P

CO P CD ft CD

H

p Pi

U P CD P rH pi rH N rH

CD >H P P!h P P ft

fflp CD P l Eh rH CM rH <u H P P

CD P o O

O

H Pm P H p P CO p O

CD P CD P CD

p K. P rH

N P P P p P P O P P a P O p P P CO

CD bH o o o ft CD pq |z; b a p CO P P o o p o P o H P u H p H p 1 1 -^

PM

p P p P p ftl 01 PPPPPH p P p

£2 P

rH

p >H o

Pip Eh

cd CD

p P P rH H H p CD ft CD P P

O) ?i P P S rHP P

Pm

HP P P

>< ^ IS]

p

M

rH rH

H rH P 0J < P • O o O P P p ti CO a H P p P P £ CD CD >> P o p P P P u p p p P P o P. o p P o p P P

inO o o k""1 p 3 p

p P P o o o P o p P P P Pi p P 1 H H H P P p PL| -=f P P P

^-^ O ^—

.-

—> • -—

v

rH CD VD 00 '—s

ft • v_^ • ^t C^

KA

CD H

V£> ft CO

p

CD

a CD • H t- P p. P b- CD CD • H P H —• H N CT>

cd P PU P ft P U P P CD

CD rH rH rH OJ P P < P O O o Pu H

cd Pi CO

p

O P ^ CD CD >s p rH P P P P p p O P P a o p p CO P P o o O CD ft £ p P cd P o o o p O P P p p p H 1 i H p p p P el ftj P P • P 1—1

«H oo

ft

O rH

CD

a i i W i P T3 P

cd P! P P CD O P P ft

ft

O 1 H T3

<H P P O COP

CD P P P P •H O

«H P r\ O P

ft P cd 1 CO P

T3

p

•H CD •H 6 P

rH P

CO CO P H > CO

CO * <H 1 P

CO ft CD

CD P CD CD P P P ctf cd

CD

Sh •H «H

EC

XX

ft

OO P u P ft ft

-H ft ft

cd CO

CDP

rH o o

CO

p P > p P c V

48
1. Comparative Acidities of Indicators In Water

and 50 Wt. Percent Methanol

Two broad classes of indicators are represented in

Table 8 -- "uncharged" acids ( sulfonephthaleins, _p-naptholbenzein) and "cation" acids (the Nile Blue group). Not all

of the indicators listed have been tested in all three of

the solvents represented. So far as comparative behavior

in water and 50 wt. percent methanol is concerned, the

following tentative conclusions may be drawn: (l) The

acidities follow the same relative order in these two solvents. (2) In 50 wt. percent methanol the Nile Blue group

of five salts fall in the expected order of decreasing

acidity. This order is based on the known effects of substituents R on the basicity of nitrogen.

2. Comparative Acidities of the Indicators in

4:1 Benzene-Methanol

The relative positions of Nile Blue 2B and Nile Blue A

with respect to uncharged acids are very different from the

positions observed for solutions in water or in aqueous

methanol. The shifted order makes possible differentiating

titrations of strong and weak acids in benzene-methanol [18]

(M. M. Davis)

49
5. SOLVENT EFFECTS ON ACID-BASE PROCESSES

OF ANALYTICAL INTEREST

Many chemical processes that occur in solution are

essentially interactions between acids and bases. These

processes include the synthesis of many organic chemicals

and drugs, for example, as well as enzymatic and related

physiological operations of importance in clinical studies

and in medical research.

It is often desirable to conduct these operations in

solvents that are not wholly aqueous, either because the

reactants are not sufficiently soluble in water or because

a particular nonaqueous solvent has a favorable effect upon

the extent of the reaction or the rate at which the reaction proceeds. Although some of these changes are quite

profound, they are very little understood.

A. Behavior of Cation Acids in Alcohol-Water Solvents

It seems clear that the effects produced by a change

in the character of the solvent medium are the result of

increases or decreases in the strength of the acids and

bases which take part in the reactions. The decrease in

the strength of weak uncharged acids when the medium is

altered by the addition of an organic solvent, producing a decrease of the dielectric constant, is well recognized.

This effect can be explained readily, in qualitative fashion,

by the electrostatic influence of the lowered dielectric

constant on the work of separating the ions.

50
The behavior of positively charged weak acids such as

ammonium ion and the protonated forms of the common amine

bases is less easy to explain. When these acids (designated

BH ; dissociate in the medium SH as shown by the equation

BH+

+ SH = SH* + B

the number of ions remains unchanged. Hence, electrostatic

terms should be without substantial influence on the position of equilibrium. Nevertheless, the general pattern of

the solvent effect of methyl alcohol on weak acids of this

type is an initial increase of strength when the alcohol is

added to a water-rich solvent. This increase continues up

to a methanol concentration between 60 and 80 wt. percent,

where the ionization constant (K) goes through a maximum

and then falls rapidly as the amount of water in the solvent

becomes small.

1. Dissociation of Tris(hydroxymethyl)aminome thane

in 50 Wt. Percent Methanol

In an earlier study performed in this laboratory [19],

the solvent effect of methanol on the dissociation constants

of anilinium-type indicators was examined by a spectrophotometric method. In order to clarify the factors responsible for the rather peculiar form of the pK curve as

methanol is added to aqueous solutions of cationic acids,

a study of the pK value of protonated tris(hydroxymethyl)-

aminomethane ("tris") in 50 wt . percent methanol was undertaken. These measurements were made by an accurate emf

51
method and extended over a temperature range (10 to 40 °C)

sufficiently wide to permit the enthalpy and entropy changes

on dissociation to be derived as well as the pK values.

In essence, the acidity function p(a„Y^-. ) for "buffer . solutions composed of tris and its hydrochloride in 50-

percent methanol was determined by measuring the emf of

hydrogen-silver chloride cells, for which the standard emf

was already known (section J.B). An "apparent" pK value

was then obtained by substitution of the acidity function

into the expression for the dissociation constant of the

acid tris*H ; the true pK was found by extrapolation of the

apparent values to zero ionic strength. The change of pK

with temperature enables the thermodynamic quantities associated with the dissociation process to be derived readily.

These procedures are set forth in somewhat more detail in

section 5.B, where isotope effects in heavy water are

discussed.

The results obtained for the pK of protonated tris [20]

in 50-percent methanol are compared in table 9 with those

earlier found in water [21]. The corresponding values of

the enthalpy (AH°) and entropy (AS°) changes are also given.

For comparison with the latter, AH° and AS° for the dissociation of ammonium ion and methylammonium ion in 60 wt

percent methanol [22] are given in parentheses. These are

the only thermodynamic data available for the dissociation

of a cationic acid in methanol-water solvents.

52
Table 9. pK and related thermodynamic quantities

for the dissociation of protonated

trls(hydroxymethyl)amlnomethane In

50-percent methanol from 10 to 40 °C.

t. °c

pK

(water)

pK

(50-percent MeOH)

10 8.516 8.273

15 8.361 8.115

20 8.213 7.962

25 8.072 7.818

30 7.934 7.681

35 7.802 7.550

40 7.676 7.426

AH°(25 °C) 47.7 kJ mol

-1

47.9 kJ mol" 1

52.6 (NH4

+

) 53.5 (MeNH_,

+

) 3

AS°(25 °c) 5.4 J deg"

-1 _-l mol 10.9 J deg" -Sno

l" 1

11.8 (NH4

+

)

- 6.4 (MeNH +

)

53
2. Dissociation of Ammonium Ion in Methanol-Water

Mixtures

Standard, emf data for the cell at other methanol-water

solvent compositions exist only at 25 °C, so thermodynamic

data (apart from pK) cannot be obtained. Nevertheless, it

was of interest to determine the course of the pK curve for

ammonium ion at 25 °C from pure water (pK = 9-245) through

60 percent methanol (pK = 8. 591). The results obtained

thus far are summarized in table 10.

Table 10. pK values for the dissociation of ammonium

ion in methanol-water solvents at 25 °C

.

Wt . Percent

Methanol

PK

9.245

20 9.044

50 8.687

60 8.591

70 8.571

54
5. Nature of the Solvent Effect on Cation Acids

As a first step toward explaining these results, the

electrostatic approach was examined. If the ions BH and

SH

p

were identical in size and contour, the work of charging a mole of each in the same medium would be equal. A

change of the solvent would therefore be without influence

on the dissociation constant and the amount of energy absorbed or released when a mole of BH ions dissociates. If

the ions are of different sizes, different amounts of energyare required to charge the two ions, and, consequently, the

dissociation energy will vary with the solvent. A further

refinement can be introduced by assuming a reasonable variation of the dielectric constant from the region of dielectric saturation at the surface of the ion outward to a distance of 4 A where the effective dielectric constant of

the solvent is believed to approach the macroscopic dielectric constant of the solvent.

Although this electrostatic calculation is able to

explain some increase in the strength of cationic acids

when methyl alcohol is added to the aqueous solvent, the

predicted increase is always far less than that actually

observed. Furthermore, the deviations are too large to be

explained by the recognized oversimplifications in the

method of calculation used. Similar attempts to explain

55
the decrease in strength for uncharged acids likewise produce predicted strengths that are lower than is actuallyobserved, yet it is significant that the deviations are

approximately the. same for acids of the two charge types

in a solvent of fixed composition. This observation makes

it appear that a "basicity parameter" characteristic of the

solvent should be superimposed upon the electrostatic parameters in order to explain fully the solvent effect [19].

The unexplained enhancement in strength suggests that

the basicity of the medium is increased when methyl alcohol

is added to water, passes through a maximum, and then decreases as the composition approaches pure methyl alcohol. At

first thought, this observation is surprising, for water molecules are believed to be more basic than methanol molecules.

At room temperature, however, a considerable part of the

water exists in a clustered structure. If methanol were

to cause a breakdown in the clustered structure of the

water, an increase in the total basicity of the medium could

result. It is hoped that further work now underway, as well

as studies elsewhere of the structure-breaking properties

of the alcohols, will shed some light on the validity of

this proposed mechanism.

(R. G. Bates)

56
B. Isotope Effect on the Dissociation of Acetic Acid

and. Dihydrogen Phosphate in Heavy Water

In section J>.C of this report, the electrometric method

for determining P( auTr--i ) of acetate and phosphate buffer

solutions in deuterium oxide as solvent was described.

This was part of work designed to establish a pD scale in

deuterium oxide analogous to the pH scale in ordinary water.

The same experimental measurements yielded the dissociation

constant of acetic acid and the second dissociation of

phosphoric acid in deuterium oxide at temperatures from 5

to 50 °C, together with the associated thermodynamic quantities, namely the changes in entropy, enthalpy, and heat

capacity, on ionization of the acid [7,8], A general view

of the constant-temperature bath and measuring equipment

used for emf studies with the deuterium gas electrode in

deuterium oxide solutions is shown in figure 9.

The dissociation constant of acetic acid in deuterium

oxide is defined by the equation:

K

_ "D-V , TD ?Ae

(1)

"tjAc ^DAc

The acidity function P( ar)Yc

-| ) = ~ log (^dTd^Ci) has been

obtained from the cell measurements described earlier.

Substitution of m-y^ from eq. 1 gives

p(a ycl ) = PK + log

^

+ log ^_ (2)

"W ^Cl^DAc

57
Figure 9. Equipment for emf studies in deuterium

oxide solutions.

In these experiments the concentrations of acetic acid and

sodium acetate were equal, so that

'VAc p(aD7cl ) = PK + log (3)

^Cl^DAc

(in dilute solutions, the ionization of acetic acid destroys

the equality m_. = m» and requires the application of a small correction.

)

Inasmuch as acetate ions and chloride ions have the

same charge, y. and y~^ will have almost the same value;

DAc, an undissociated molecule, will have an activity coefficient close to unity. Hence p(any(-, 1

) hardly varies with

58
concentration and pK can be found as the limiting value of

a plot of p(aLycl ) against the total ionic strength.

The calculation is not so easy for the second constant

of phosphoric acid,

D

2

P04

" ±> H

+

+ DP04=

, K

2

even if, as in the present case,

mD

2

P04

_ ~ "bPO^

Now,

"*2 " ^ vctH

r Cl

ToPO^-

pKQ : pCa^Yp-,) - log ~ W

YC1- YD

2

P04

The activity coefficient term can, however, be approximated

by a Debye-Huckel expression and the right-hand side extrapolated to zero concentration in order to obtain pK

2

for

phosphoric acid in the deuterium oxide solvent.

Dissociation constants of both of these acids have been

measured from 5 to 50 °C. The results are given in table 11,

where the data in deuterium oxide as solvent are compared

with those for water as solvent.

The results for both acids can be fitted to the

equation

A

i pK = — - A, + A,T

m c- J

59
Table 11. The dissociation constant of acetic acid

and the second dissociation constant of

deuteriophosphoric acid in deuterium

o oxide from 5 to 50 C.

Acetic acid Phosphoric acid

t, °c pK in D

2 pK in H2 pKp in DpO pK

2

in H

2

5 5.346^ 4.7701 7.8846 7.2810

10 5.334^ 4.7628 7.849

9

7.2545

15 5.3236 4.7581 7.823^ 7.2324

20 5.3168 4.7558 7.7986 7.2145

25 5.313 4.7558 7.7796 7.2005

50 5.310Q 4.7581 7.766

?

7.1902

35 5.309^ 4.7624 7.754

?

7.1834

40 5.3118 4.7688 7.7484 7.1800

45 5.3l6

?

4.7770 7.743

3

7.1799

50 5.324

5

4.7871 7.743

5

7.1828

Thus, by standard thermodynamic methods, the changes in

enthalpy, entropy, and heat capacity on dissociation can

be derived. Values at 25 °C in heavy water and ordinary

water are given in table 12.

60
Table 12. Thermodynamic quantities for the dissociation

of acetic acid and phosphate ion in water and

in deuterium oxide*.

DAc

in

D

2

HAc

in

H

2

w

in

D

2

H

2

P04

"

in

H

2

AH°

(cal mol ) AS°

(cal deg" mol ) AC°

(cal deg mol )

275

-23.4

-38.6

- 98

-22.1

-36.6

1376

-31.0

-58.4

987

-29.6

-54.1

* 1 cal = 4.1840 J

(R. A. Robinson)

C . Improved Spectrophotometric Techniques: Design and

Construction of a Controlled-Temperature Cell Block

In the spectrophotometric method for determining the

dissociation constant of a weak acid or base, measurements

are made, at a given wavelength, of the absorption of light

by a solution containing the acid in its undissociated state,

another solution in which the acid is completely dissociated

into its ions, and a third solution containing both forms..

The light absorption measurements give information about

61
the relative proportions of each species present in this

last solution. The resulting dissociation constant is known

to be dependent on the temperature, and it is important to

have some provision for temperature control. This may be

accomplished by making the measurements in a temperaturecontrolled room or by thermostating the cells containing

the solutions. Improvement in the latter operation has been

achieved.

A thermostated cell block was constructed using the

metal cell holder supplied with the spectrophotometer. The

holder has provisions for four cells, and a heater was placed

in each of the outer cell positions. The heaters were formed

by winding 7 ft. of 26 B and S nichrome wire, insulated with

glass spaghetti, onto each of two aluminum blocks. The complete heaters fitted snugly into the cell holder. The heater

windings were connected in series.

A thermistor cemented to the cell holder served as a

sensor in a temperature- control circuit. A second thermistor, used for temperature measurement, was mounted on a

movable metal bracket so that it could be placed in the

solution in one of the cells or moved aside when solutions

were being changed. Leads from the heaters and thermistors

terminate in a six-pin socket mounted on the cell holder;

the leads from a corresponding six-pin plug were passed

through a hole in the lid of the cell compartment.

62
The resistance of the measuring thermistor is determined using a D.C. Wheatstone bridge. The leads from the

regulating thermistor are connected to the input of a commercial electronic proportional temperature controller.

The heaters constitute the load for this controller. Provision is made for connecting a 100- or 250-ohm ballast load

in series with the heaters.

Using this apparatus, the temperature of the solutions

can be regulated between room temperature and 50 °C with a constancy of 0.01 °C . The prototype of this cell block is

shown in figure io.

Figure 10. Controlled-temperature cell block for

spectrophotometric measurements.

(B. J. Steel)

63
6. AQUEOUS SOLUTIONS OP MIXED SALTS

The Section's program of research on the properties

of aqueous mixtures of pure salts has been assisted by a grant from the Office of Saline Water of the U.S» Department

of the Interior.

A. Methods and Procedures

In general, this research is carried out by the isopiestic method of vapor pressure measurement. This method

consists essentially in comparing the vapor pressure of

solutions containing two or more salts with the vapor pressure of a solution of a single salt (the "reference" salt).

The apparatus is shown in figure 11. Small silver cups

with close-fitting hinged lids are placed, with lids raised,

on a heavy copper block. The block and cups are placed in

a. vacuum desiccator which has a spherical joint in the center

of the top. Through the spherical joint passes a glass rod

with a "T" on the end. Above the spherical joint Is the

vacuum control stopcock. The spherical joints, when greased,

permit the rotation of the glass rod which in turn rotates

the light wire frame which holds open the lids of the cups.

A gap in the wire frame permits the lids of the cups to fall

one by one as the glass rod in the top of the desiccator is

turned. Thus the lids of the cups can be closed before the

vacuum of the desiccator is broken. During the experiment

the desiccator is immersed in a thermostat.

64
s

5 >

f#-

Figure 11. Isoplestlc vapor pressure apparatus,

showing the silver dishes and the

copper block in the desiccator.

65
A typical determination is carried out in the following way. Two stock solutions, each containing one of the

two salts to be mixed, are prepared. Mixtures of the stock

solutions in different proportions are then prepared by

weight. The solutions so prepared are weighed into the

silver cups which are placed on the copper block (precooled

to reduce vapor loss) inside the desiccator. The desiccator

is evacuated and placed in the thermostat. It is gently

rocked by a. mechanical arrangement in the thermostat in

order to prevent the formation of concentration gradients

in the solutions as they equilibrate with one another.

When the solutions have equilibrated, the lids on the

cups are closed, the vacuum in the desiccator is broken,

and the cups are reweighed. As the cups lose only about

0.0001 g of vapor per minute, it is easy to correct the

weights to the time the vacuum was broken. In each determination, three of the 11 cups in the desiccator contain a solution of an electrolyte such as potassium chloride whose

thermodynamic properties are well known. This solution is

the standard or "reference" solution.

Many of the simple salt systems which are of the greatest general interest can be studied easily and precisely by

the isopiestic method. This method provides the same kind

of data as the measurement of the emf of cells. The determination of the thermodynamic properties of mixtures of

66
electrolytes In solution by the measurement of emf of cells

Is, however, restricted by the availability of suitable

electrodes. It would be difficult, for instance, to devise

an emf method for the determination of the properties of the

mixture water-sodium chloride-potassium chloride. However,

this simple system can be studied with relative ease by the

isopiestic method.

B. Results

The first system studied within the reporting period

was the system water-glycine-potassium chloride. For this

system there were indeed some electrometric data in the

literature covering a narrowly restricted range of concentrations. A broader range of data was desired, and furthermore

it was desirable to test, with experimentally determined

values of the solubility of each component, both the underlying theory and the accuracy of the isopiestic method.

In a paper published early this year [23], data were reported

for the system water-glycine-potassium chloride which agreed

excellently with prior electrometric data for a restricted

range of concentrations. The new data on the activity

coefficients of the constituents of the system also predicted in a most satisfactory way the solubility of each of the

components in the presence of the other.

67
Two other papers [24,25] explored the properties of

mixtures of sodium chloride with barium chloride and of

potassium chloride mixed with the barium chloride. These

two systems behave quite differently. Thus the isopiestic

ratio in the system sodium chloride-barium chloride is a

linear function of the ionic concentration, whereas in the

system potassium chloride-barium chloride the same two

quantities bear a non-linear relationship to one another.

When the total ionic strength is unity, addition of

sodium chloride to barium chloride solutions is found to

increase the activity coefficient of barium chloride, and

the addition of barium chloride to sodium chloride solutions

reduces the activity coefficient of sodium chloride. Thus

the activity coefficients tend to come closer to one another,

In the system potassium chloride-barium chloride, the activity coefficients of both salts are depressed.

At a total ionic strength of about 2.5, the activity

coefficients of sodium and barium chlorides in mixtures of

these salts display a slight reversal of the tendency for

the activity coefficients to come together. The separation

effect on the activity coefficients of potassium and barium

chlorides at this ionic strength is still more marked.

The type of behavior exhibited in the potassium chloridebarium chloride system is atypical and must indicate unusual

ionic interactions in this system.

68
A fourth paper [26] deals with an addltivlty rule for

the vapor pressure lowering of aqueous solutions containing

an alkali metal chloride and an alkaline earth metal chloride. It Is found that the vapor pressure lowering of an

aqueous solution containing two salts A and B can be compounded addltlvely from the vapor pressure lowerlngs of a solution containing A alone and a solution containing B

alone. The addltivlty rule disclosed in this paper is that

the vapor pressure lowering of the mixture of alkaline metal

chlorides and alkaline earth metal chlorides is the sum of

the respective vapor pressure lowerlngs for the salts, each

taken at the total ionic strength of the mixed solution.

In some instances the addltivlty of vapor pressure is

valid to within 0.5 percent. Six systems were examined in

this paper, and the greatest deviation from addltivlty

amounted to 2.0 percent.

A fifth paper [27] deals with the osmotic and activity

coefficients of tris(hydroxymethyl)aminomethane, a compound

of interest both as an acidimetric standard and as a biological buffer substance. This base behaves as an almost

ideal solute even at high concentrations. The activity

coefficients of its hydrochloride are similar to those of

ammonium chloride up to 1 molal, and even in concentrated

solutions the differences are small. It is concluded that

ion-pair formation is not marked.

(V. E. Bower)

69
7. BEHAVIOR OF SODIUM-RESPONSIVE GLASS ELECTRODES

The purpose of this investigation is to determine some

of the advantages and limitations of sodium-responsive glass

electrodes for the analysis of sodium ions in solution and

for the determination of thermodynamic data for solutions

of sodium salts in mixed solvents.

A. Method of Application

If a sodium-responsive glass electrode and a silversilver chloride electrode are placed in a solution of sodium

chloride to form the cell:

Glass || NaCl (m), AgCl; Ag

the emf of the cell should be given by:

2 RT

E = E» lna + (l)

P

In eq. 1, a + is the mean ionic activity of sodium chloride

in the solution and E' includes contributions from the

standard potential of the silver-silver chloride electrode,

the potential of the internal reference electrode, and the

asymmetry potential of the glass electrode.

If the emf of the cell is measured for different concentrations of NaCl, then, provided E

1 is constant ,

2 RT (a + ) E - E_ = AE = 2.303 x log —^-i (2)

1 F (aj±2

70
Thus, provided a + is known for one solution which can

be used as a reference, a + can be determined in other solutions. As a reference solution it is convenient to use an

aqueous solution of sodium chloride, since values of mean

ionic activity coefficients y+ for these solutions are already tabulated in the literature. Hence a + = m y+ is known

for the reference. Since the factor 2.303 RT/P is 0.05916

at 25 °C,

a + AE = 0.11832 log (3)

(a ± )ref

Thus, in order to determine log a+ with a precision of

0.001, AE must be measured with a precision of better than

0.1 mV. It is also essential that E' remain constant during

the determination. The initial problem, therefore, is to

measure the emf of the cell with a. precision of 0.05 mV and

to determine whether the emf of the cell is constant for a reasonable period of time (E 1 constant).

B. Procedures

The cell potentials were measured using a precision

potentiometer. Because of the high impedance of the cell

containing a glass electrode, an expanded-scale pH meter

was used as a null detector. The feedback circuit of this

meter had been modified to further increase its sensitivity.

71
The output of the meter was fed into a 1 mV recorder in

order to study easily the time-dependence of the measured

emf . The sensitivity of the circuitry was adjusted so

that 0.05 mV out of balance corresponded to a recorder

deflection of 0.1 in. Initially, electrical noise was very

troublesome, but with close attention to shielding and the

use of an isolating transformer in the a.c. supply circuit,

the noise could be reduced to less than 0.02 mV. The

thermostat and measuring equipment are shown in figure 12.

Figure 12. Equipment for studies of the behavior

of sodium-responsive glass electrodes,

72
The measuring circuit was tested by measuring the emf

of a standard cell in series with a 100-megohm resistor.

The emf of the cell agreed to ± 0.02 mV with that determined

using a low-resistance potentiometer circuit. Further, the

emf was constant for up to 1 hour, showing that there is

negligible drift in the pH meter.

C. Results

This project is new, and only preliminary results have

thus far been obtained. When measurements were made on the

glass electrode cells, thermostated at 25° ± 0.01 °C, the

emf showed a marked time dependence. In some cases, the

cell emf drifted at a rate of 1 mV per minute and, although

this rate decreased with time, the cells did not attain a steady emf within several hours. It was eventually shown

that this drift could be eliminated by rapidly stirring the

solution in the cell. If the stirring was stopped, the emf

immediately began drifting again. This effect is present

to a smaller extent in new electrodes and is being investigated further.

As a test of the response of the electrodes, the emf

of the cells was measured as a. concentrated stock solution

of sodium chloride was added in weighed aliquots to a stirred solution of sodium chloride in the cell. The results

of a typical run are shown in table 13 . It can be seen

that errors up to 0.25 mV are present in these measurements.

73
Table 13. Comparison of observed and calculated

responses of a sodium- sensitive glass

electrode at 25 °C

Molality

of NaCl

log a + O.II832 A log a + AE

0.06355 -1.2897 (Referenc*

0.12005 -1.0359 0.0300^ 0.03006

0.15043 -0.9461 o.oio6

2

0.0106c5

0.17380 -0.8886 0.0068

Q

0.0066

0.19539 -0.8421 0.0055 0.0052c

The best type of silver-silver chloride electrode to

use in these measurements is being investigated. The

thermal-electrolytic type of electrode is rather sluggish

in response, probably because of its porous nature. Electrodes formed by electrodepositing silver chloride onto

silver wire or chemically deposited silver are being tested

(B. J. Steel)

74
8. CONDUCTOMETRIC DETERMINATION OF TRACES OP WATER

During the past year, work has been initiated on the

application of conductance techniques to the determination

of traces of water in gases. The objectives have been

oriented toward the development of a simple, inexpensive

instrument for the determination of hydrogen in the microelemental analysis of organic compounds. To be useful in

this context, it is necessary to determine 0.5 to 5 nig of

water with an error not greater than 0.01 mg.

The detection scheme was based on the great sensitivity

of the conductance of solutions of sulfuric acid in acetic

acid to small amounts of water [28], It has been found that

the fractional increase of the specific conductance of solutions of 0.02 to 0.04 M H

2S0^ is about 85O times the weight

fraction of water added [29]. Qualitatively, the effect

depends on the relative acidities of the constituents,

namely H

2

S0|. > HOAc > H

2 0. Sulfuric acid is a weak acid in

this solvent, contributing relatively few protons to the

acetic acid. Water is sufficiently basic to accept protons

readily from sulfuric acid but not from the solvent . It

was predicted that propionic acid would provide an even

more sensitive system.

75
A. Apparatus

The prototype model of the instrument consists basically of three' elements: an absorption-circulation system for

transferring water' from the gas stream to the electrolyte,

the conductance cells, and the electrical measuring system.

The absorption system and conductance cell are shown in

figure 13,

Figure 13. Conduct ometric cell for the determination

of small .amounts of water in a gas stream,

76
The absorber consists of a vertical helix of 2-mm glass

capillary tubing. Gas bubbles entering the bottom of the

helix push liquid ahead as they rise and contact the wetted

wall. At the top, the spent gas is vented through a drying

tube. It was originally intended that the absorber would

also provide circulation of the electrolyte through the

conductance cells by gravity flow down the tube at the center

of the helix. The flow rate proved to be too slow, however,

requiring 2 to 3 minutes for one complete circuit of the

apparatus. Consequently, an enclosed glass magneticallyactuated pump is now being constructed to provide rapid

circulation and mixing of the cell contents. The four-way

stopcock provides either for flow through both conductance

cells in series or for closing the reference cell and continuing circulation through the indicating cell. The instrument holds 15 cc. of electrolyte.

The two conductance cells are formed by three parallel

stainless steel disks separated by l/l6-in. polytetrafluoroethylene gaskets. The middle disk is a common electrode

for both cells, and the liquid flow connections are made

to both outside disks. This construction provides for a good thermal contact between the cells, for low cell constants (about 0.005), and for small volumes.

The electrical measuring system is a simple Wheatstone

bridge. The cells constitute adjacent arms so that it is

the ratio of the conductances which is measured. Since the

77
effects of polarization and temperature changes are nearly

the same in each cell, these are largely cancelled. One of

the two low-potential corners of the bridge is connected to

the common electrode and the other to the slider of a 1000-

ohm, 10-turn precision potentiometer which, with 5000-ohm

end coils, is connected across the outside electrodes.

B. Results

1. Prototype Instrument

a. Water Content of Carrier Gas . In trials with

the gas-lift absorption-circulation system, the cell resistance did not reach a constant value after additions of water.

After 10 to 15 minutes of circulation, the conductance continued to increase at a uniform rate. This drift rate was

equivalent to approximately 400 ppm by volume of water in

the "super dry" tank nitrogen used as the carrier gas. Prior

passage of the gas through a bed composed of a molecular

sieve (size 4A) reduced the drift to a value corresponding to

about 150 ppm of water, and this value was also characteristic

of tank argon. The drift in conductance also depended on the

water content of the electrolyte. Argon or pre-dried nitrogen

apparently removed water from solutions of sulfuric acid in

propionic acid when more than 300 ppm (by weight) of water was

present. In any combustion procedure, no more than one liter

of carrier gas should be required to sweep the products into

the analyzer, and hence the correction for moisture in the

g,a^ can probably be made negligible.

78
b. Calibration . Of several methods of calibration which were investigated, the most satisfactory was the

direct addition of liquid water by weight from a micropipet.

Pipets have been made of 5 l~il capacity and a 40-mg tare.

The calibration results can be expressed in terms of the

response factor, F, which is the ratio of the fractional

change of specific conductance to the weight fraction of

water added to the electrolyte system. The expected value

of F was 85O, but calibrations of O.OO96 M and 0.0190 M

sulfuric acid solutions gave values in the range 415-490.

Accordingly, conductances of the systems water- sulfuric

acid-organic acid were examined in more detail.

2. Fundamental Data

a. Materials . Reagent-grade acetic and propionic

acids were distilled through a fractionating column. Approximately 0.4 M sulfuric acid stock solutions were made from

concentrated HpSOj, and were standardized potent iometrically

against potassium acid phthalate. An acetic acid solution,

0.0386 M in HpSO^ and propionic acid solutions 0.052, 0.086,

and 0.209 M in HpSOK were prepared by dilution. Additions

of water were made by weight.

b. Conductances . The conductances of the solutions were measured with conventional conductance equipment

at 25 °C. Cells having constants of O.O616 and 0.001075

were used. The molar conductance of the acetic acid solu-

-1 2 -1 tion was O.O637 ohm cm mole , about 50 percent higher

79
than the literature value [28]. Water was added in amounts

up to 0.18 percent and the resistance (R) measured. A plot

of log R vs. ppm HpO showed appreciable curvature. The

initial slope corresponded to P = 524, while the slope at

the highest water content corresponded to F = 331

The specific conductance of the solvent propionic acid

Molar conductances of the three

sulfuric acid solutions increased with concentration from

nrrll v, -1 -1

was 5 x 10 ohm cm

-6 -S •1 1 8.8 x 10 to 5.4 x 10 ohm """ cm """, reflecting the extremely slight ionization of the electrolyte. The effect of

water on the conductance of each of these solutions was

qualitatively similar to that in acetic acid. Values of F

for the initial additions of water are shown in table 14.

Table 14. Effect of water on the conductance

of solutions of sulfuric acid in

propionic acid.

As expected, F is appreciably larger in propionic acid than

in acetic acid. The conductance data for one of these solutions are shown in table 15.

80
Table 15. Conductance of Water and 0.209 M

sulfuric acid in propionic acid at 25 °C

(Cell constant: 0.001075).

Cell Resistance

Added Water, ppm Measured Calculated

95,490 (95,490)

45.0 91,711

66.4 88,400 88,653

129.2 83,940

240.9 77,960 74,606

421.7 68,740 64,093

1095.0 44,801 42,025

2018.8 28,809 28,542

3. Revision of Instrumentation

The unexpected non-linearity of the water response is

inconvenient in an instrument intended to provide a direct

analytical result. The circuit shown in figure 14 compensates to a considerable extent for the observed behavior.

In this diagram, D represents the potentiometer setting, in

ohms, and R

? is the resistance of the reference cell.

When initially balanced, the resistance of the indicating cell, R, is equal to R and D = 0. As water is added,

Rp remains constant, and the decrease in R is compensated

81
\

DET.

A "1 D B

AAAA/W

c

Figure 14. Circuit designed to compensate for the nonlinear relationship between conductance

and amount of water.

by increasing D. It is readily shown that, for this circuit

(

d In l0\ = s = ( B + C )

d D / d=0 B(A + C)

(1)

where A, B, and C represent the resistances of the elements

identified on the figure. The initial slope S is related

to the initial value of F by

S

Q = -FQ/(dD/dW) (2)

82
where dD/dW is a scale factor relating the dial reading to

water content.

It will be remembered that F was defined in terms of

the change in specific conductance and therefore is opposite

in sign to S. The working volume of the absorber plus indicating cell is 10 cc, so that 100 mg of water provides a concentration of 1 percent. For the system described in

Table 15, F is 1162 and, hence, -S = 0.01162 in order to

represent 1 percent water by 1000 ohms in D. At the initial

balance condition, B = AC/(A+C), and this relation requires

that -AS = (2A-C)/(A+C) . Based on an arbitrary choice of

1.5 for -AS , A = C = 129.1 and B = 64.5. These values

o were used to calculate the third column of Table 15. The

agreement with the observed data could be improved by a small change in F . o

(T. 3. Hoover)

83
9. REFERENCE MATERIALS FOR DIELECTRIC MEASUREMENTS

A. Tertiary Butanol

Tertiary butyl alcohol is an amphiprotic solvent of

low dielectric constant which is completely miscible with

water and most other common solvents. This substance is

easily acquired in reasonable purity and appears to be

stable up to its boiling point of 82.5 °C . It has a sharp

melting point at about 25.6 °C and so it lends itself conveniently to 'purification by fractional distillation and

stepwise recrystallization. The purified compound has a conductivity of less than 10 ohm cm" at 30 °C

.

Samples are easily produced with conductivities well below

_0 _-i _-]

10 * ohm cm , and these can be used with minimal conductivity error in dielectric constant measurements.

In view of these characteristics, accurate measurement

of the dielectric constant of this alcohol can serve to

establish it as a convenient reference standard for comparative measurements, provided consistent values can be obtained

on independent samples. Along with the measurement of the

dielectric constant (e) of the pure alcohol and its aqueous

mixtures now in progress, comparative data have been obtained

on samples from three commercial sources. Values at 25 and

30 °C obtained by a bridge method are shown in table 16

,

along with the measured melting points and respective conductivities (y). Included are determinations made both by

comparative and absolute methods.

84
The results show considerable consistency, and the

usefulness of this compound as a standard reference material

appears promising. Furthermore, the melting point of tertiary butanol has been proposed as a fixed reference point

for temperature measurements and could conceivably be

established as such.

Table 16. Dielectric constant of samples of

tertiary butyl alcohol at 25 and JO °C;

conductivity at JO °C

.

Sample

No.

m.p., °C € 25 e jo Y30

|jmho cm

1A1 11.444 0.0050

1C1 11.436 0.0016

2A 25.

6

Q

12.46

2

11.446 0.0021

2AB 12.46

Q

11.445 0.0008^

2AA1 25.

6

Q

12.458

11.444 . 00006-.

3

3AB 25.

6

6

11.447 0.0001^

5

B. Benzene

Benzene has long been used as a standard reference

liquid in investigations of dielectric properties. It is

not well suited for tnis purpose, inasmuch as water is

appreciably soluble in benzene and influences its dielec-

* trie constant. The results of recommended procedures for

85
drying benzene are in effect not entirely consistent.

Furthermore, in many cases they are extremely time-consuming,

extending to a period of months or more.

Some comparisons of the effectiveness of four drying

agents were made using a common commercial lot of benzene.

This material was purified by distillation and recrystallizatlon and dried with non-indicating Drierite until a melting point of 5.52 °C was attained. Representative

results for the dielectric constant of benzene, shown in

table 17, indicate that Drierite and molecular sieves have

satisfactory drying rates.

Table 17. Effect of drying benzene on the dielectric

constant at 30 C.

Drying agent or condition

Drying

time

(days)

Undried Wet at

22 °C

Drierite Molecular

sieves

CaH

2 Silica

gel

2.2670

2.2665

2.2740

1 2.26^1

3 2.2626

4 2.2625 2.2637 2.2648

5 2.2629 2.2634 2.2646

7 2.2627

10 2.2627 2.2626 2 . 2634

Dr led stock + CaHp

2 2.2626

86
The differences evident in table 17 are sufficientlylarge to explain the scatter of the dielectric constants of

benzene reported In the literature. If the literature value

of 0.54 g/1 is taken for the solubility of water in benzene,

-4

the data suggest a sensitivity in € of 2 x 10 units per

0.01 g of water.

These measurements are made with the three- terminal

cell shown in figure 15. It has a working capacity of 21 pf

with a liquid volume of about 5 ml and is suitable for use

with most low-loss liquids. The capacitance-conductance

bridge assembly and constant temperature baths employed are

shown in figure 16.

Figure 15 . Three-terminal cell for measurements of

dielectric constants.

87
n

Figure 16. Capacitance-conductance bridge assembly

and constant-temperature baths for

precise measurements of the dielectric

constants of liquids.

C . Dimethylsulfoxide-Water Mixtures

Analysis of the data obtained on dimethylsulfoxide

and its aqueous mixtures over the temperature range -5

to 75 °C is in progress. A satisfactory interpretation

of tne results will require further time.

(C. G. Malmberg)

88
10. KINETIC METHODS OF ANALYSIS

The Section's program dealing with the effect of a change in medium (or environment) on the kinetics of reactions has been redirected in part. In conjunction with

the considerable Division effort in trace analytical procedures., emphasis is now being placed on the application of

kinetic measurements to analytical problems.

Reaction-rate methods are important in analytical

chemistry because many quantitative determinations can be

based on measurements of rate. For example, in the analysis

of a mixture containing two or more closely related compounds (such as isomers ), use can be made of the fact that

these compounds undergo the same type of reaction with an

added reagent, but at different rates. Likewise, catalyst

concentrations can be determined from the rates of catalyzed

reactions.

Trace concentrations of many ions can be determined

from the catalytic effects of the ions on certain reactions.

The sensitivity of these methods is, generally, quite high;

however, the reaction chosen is frequently subject to catalysis by other ions. This lack of specificity is a major

problem, and experiments are designed to remedy this defect

by choosing conditions such that the importance of these

interferences is minimized. An understanding of the reaction mecnanism can aid in achieving this goal.

89
Trace amounts of cobalt can be determined from studies

of the rate of reaction between Alizarin Red (a 1,2-dihydroxyanthraquinone dye) and hydrogen peroxide. The usefulness of

this method is being examined, and the conditions necessary

for high sensitivity and accuracy are being studied.

A. Experimental

The rate of reaction between Alizarin Red and hydrogen

peroxide was determined by observing the decrease in absorption of the dye with time. All spectra were recorded, and

all of the absorption measurements were made with the aid of

the recording spectrophotometer shown in figure 17.

The water used was obtained by passing the general

laboratory distilled supply through a mixed bed ion-exchange

column. The disodium hydrogen phosphate solution was prepared from salt available as an NBS Standard Reference

Material. An aqueous solution containing 10 percent sodium

hydroxide was subjected to electrolysis employing a mercury

pool cathode and platinum gauze anode.

B. Preliminary Results

The rate of the cobalt-catalyzed reaction passes through

a maximum value at a pH of about 11.1 in buffer solutions

composed of disodium hydrogen phosphate and sodium hydroxide.

The rate of the uncatalyzed reaction was studied in the

presence of 0.5 x 10~^ M to 1.5 x 10~^ M disodium ethylenediaminetetra-acetate and was found also to pass through a maximum value at a pH of 11.1 in the same buffer solutions.

90
Figure 17, Spectrophotometer used to measure the

rate of the reaction between Alizarin

Red and hydrogen peroxide.

The reaction is first-order both with respect to the

dye and hydrogen peroxide at the pH where the maximum

occurs. The magnitude of the difference in the rates between the catalyzed and uncatalyzed reactions is such that

-Q

a sensitivity of 0.1 x 10 ' M cobalt can be attained.

However, erratic results and a lack of satisfactory reproducibility when cobalt is present poses a serious problem.

All experiments that have been performed up to the

present time suggest that most of the trouble is due to

adsorption of cobalt by the surface of the container in

91
which the reaction is taking place. This process "begins

when cobalt is added to the reaction vessel and continues

while tne reaction is occurring; hence the cobalt concentration and the rate of the reaction change during a run.

A polytetraf luoroethylene (Teflon) surface appears to

present less of a problem in this respect than do either

borosilicate (Pyrex) glass or polyethylene (Nalgene)

surfaces. The suitability of organo- silicon coated surfaces

is being investigated. Residual impurities in the water and

reagents and contamination from containers do not noticeably affect the reproducibility as long as the stock solutions are prepared in the same containers and the number of

such containers is held to a minimum. Residual impurities

will, however, limit the sensitivity of the method.

Future experiments will be designed to study the

tolerance of the reaction to tne effects of certain interferences. Attempts will be made to improve the precision

of the method at high sensitivities.

(R. K. Wolford)

92
11. PERSONNEL

Electrochemical Analysis Section, Roger G. Bates, Chief

R. A. Robinson, Assistant Chief

Group I. Electromotive Force and

Optical Measurements of Acidity

V. E. Bower

R

.

Gary

H. B. Hetzer

M. Paabo

R. A. Robinson

B. J. Steel, On leave from the University of

Adelaide, South Australia.

Group II. Conductance Techniques

T. B. Hoover

Group III. Dielectric Measurements

C. G. Malmberg

Group IV. Spectroscopic Techniques for Nonaqueous Media

M. M. Davis

Group V. Isopiestic Vapor Pressure Measurements

V. E. Bower

R. A. Robinson

Group VI. Kinetic Techniques

R. K. Wolford

Group VII. Solubility Techniques

P. W. Schindler, Guest worker of the Swiss

National Foundation, on leave

from the University of Berne.

Section Secretary

M. Raudenbush

93
12. PUBLICATIONS AND MANUSCRIPTS, JULY 1964 TO JUNE 1965

1. Paabo, M., Robinson, R.A., Bates, R.G.

Standard Electromotive Force of the Hydrogen-Silver

Chloride Cell and the Thermodynamics of Solutions

of Hydrochloric Acid in 50 Weight fo Methanol from

10 to 40 °C.

J. Chem. and Eng. Data 9, 574-376 (1964).

2. Gary, R., Robinson, R.A.

Standard Potential of the Ag-AgCl Electrode in 5$

Aqueous Mannitol.

J. Chem. and Eng. Data 9, 376-378 (1964).

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

Thermodynamics of Aqueous Solutions of Hydriodic

Acid from Electromotive Force Measurements of

Hydrogen-Silver Iodide Cells.

J. Phys. Chem. 68, 1929-1933 (1964).

4. Vosburgh, W.C., Bates, R.G.

Constancy, of a Modified Weston Standard Cell over

Long Periods.

J. Electrochem. Soc. Ill , 997-998 (1964).

5. Hoover, T.B.

Conductance of Copper m-Benzenedisulfonate

Hexahydrate in N-Methylpropionamide from 20 to 40°.

J. Phys. Chem. 68, 3OO3-3OOR (1964).

6. Ong, K.C., Robinson, R.A., Bates, R.G.

Interpretation of Potentiometric Titrations of Weak

Acids in Methanol-Water Solvents.

Anal. Chem. 36, 1971-1972 (1964).

7. Wolford, R.K.

Kinetics of the Acid-Catalyzed Hydrolysis of Acetal

in Dimethyl Sulfoxide-Water Solvents at 15, 25,

and 35°.

J. Phys. Chem. 68, 3392-3398 (1964).

8. Gary, R., Editor

English Translation, Nesmeyanov, A.N.,

"Vapor Pressure of the Chemical Elements", pp. 1-462,

American Elsevier Publishing Co., New York, 1964.

94
9. Gary, R., Bates, R.G., Robinson, R.A.

Second Dissociation Constant of Deuteriophosphoric

Acid in Deuterium Oxide from 5 to 50°.

Standardization of a pD Scale.

J. Phys. Chem. 68, 3806-3809 (1964).

10. Paabo, M., Robinson, R.A., Bates, R.G.

Reference Buffer Solutions for pH Measurements

in 50$ Methanol. Dissociation Constants of Acetic

Acid and Dihydrogen Phosphate Ion from 10 to 40°.

J. Am. Chem. Soc. 87, 4l5-4l8 (1965).

11. Robinson, R.A., Bower, V.E.

Thermodynamics of the Ternary System: Water-Sodium

Chloride-Barium, Chloride at 25 °C

.

J. Research NBS 69A, 19-27 (1965).

12. Malmberg, C.G.

Electrical Conductivity of Dilute Solutions of

"Sea Water" from 5 to 120 °C„

J. Research NBS 69

A

, 39-43 (1965).

15. Kolthoff, I.M., Bruckenstein, S., Bates, R.G.

Acid-Bases in Analytical Chemistry

Interscience Reprint (from Kolthoff-Elving "Treatise

on Analytical Chemistry" ) . Interscience Publishers,

(1965).

14. Bates, R.G.

Acids and Bases in Alcohol-Water Solvents.

Chem. Chem. Ind. (Japan), 18, 68O-685 (1965).

15. Woodhead, M., Paabo, M., Robinson, R.A., Bates, R.G.

Acid-Base Behavior in 50 Percent Aqueous Methanol:

Thermodynamics of the Dissociation of Protonated

Tris ( hydroxymethyl)aminome thane and Nature of the

Solvent Effect.

J. Research NBS 69A, 265 (1965).

16. Bates, R.G.

Acids, Bases and Buffers.

Proceedings Toronto Symposium, Electrochemical

Society. Ann. N.Y. Acad. Sci. (in Press).

17. Bates, R.G.

Acidity Functions for Amphiprotic Media.

Chapter for Chemistry of Non-Aqueous Solutions.

Academic Press, New York. (in Press).

95
18. Bates, R.G.

pH Determination.

Encyclopedia of Industrial Analysis

Interscience, New York. (in Press).

19. Robinson, R.A., Bower, V.E.

Osmotic and Activity Coefficients of

Tri s

( hydroxymethyl)aminomethane and its

Hydrochloride in Aqueous Solutions at 25 °C

J. Chem. and Eng. Data. (in Press).

20. Gary, R., Bates, R.G., Robinson, R.A.

Dissociation Constant of Acetic Acid in Deuterium

Oxide from 5 to 50°. Reference Points for a pD

Scale.

J. Phys. Chem. (in Press).

21. Robinson, R.A., Bower, V.E.

An Additivity Rule for the Vapor Pressure Lowering

of Aqueous Solutions.

J. Research NBS . (in Press).

22. Robinson, R.A., Bower, V.E.

Properties of Aqueous Mixtures of Pure Salts:

Thermodynamics of the Ternary System Water-Potassium

Chloride-Barium Chloride at 25 °C.

J. Research NBS. (in Press).

23. Davis, M.M., Hetzer, H.B.

Titrimetric and Equilibrium Studies Using Indicators

Related to Nile Blue A.

Anal. Chem. (in Press).

24. Woodhead, M., Paabo, M., Robinson, R.A., Bates, R.G.

Buffer Solutions of Tris( hydroxymethyl)aminomethane

for pH Control in 50 Wt. % Methanol from 10 to 40 °C.

Anal. Chem. (In Press).

25. Bates, R.G.

Proposal for the Practical Measurement of pH in

Amphiprotic and Mixed Solvents.

Pure Appl. Chem. (in Press).

26. Bates, R.G.

Book Review - "Acid-Base Equilibria" by E. J. King.

Anal. Chem. (in Press).

96
13. TALKS, JULY 1964 TO JUNE 1965

1. Bates, R.G., "Acids and Bases In Alcohol-Water

Solvents", Chemical Society of Japan, University of

Tokyo, Tokyo, Japan. July 27, 1964.

2. Bates, R.G., "Acids and Bases in Alcohol-Water

Solvents", American Chemical Society, Hawaii Section,

University of Hawaii, Honolulu, Hawaii.

August 3, 1964.

3. Bates, R.G., "pH Measurements in Amphiprotic Solvents",

Gordon Research Conference on Analytical Chemistry,

New Hampton, New Hampshire. August 17, 1964.

4. Bates, R.G., "Acid-Base Behavior in Amphiprotic

Solvents", Chemistry Seminar, Rutgers University,

New Brunswick, New Jersey. October 29, 1964.

5. Robinson, R.A., "Thermodynamic Properties of Multicomponent Solutions", Office of Saline Water,

Department of Interior, Mellon Institute,

Pittsburgh, Pennsylvania. November 12, 1964.

6. Bates, R.G., "Acids, Bases, and Buffers", Symposiumon Current Concepts of Acid-Base Measurements,

New York Academy of Sciences, New York, N.Y.

November 23, 1964.

7. Bates, R.G., "Acidity and Basicity", Joint Board on

Science Education, Randall Junior High School,

Washington, D.C. December 16, 1964.

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

Solvents", Chemistry Seminar, University of

Pennsylvania, Philadelphia, Pennsylvania.

February 24, 1965.

9. Bates, R.G., "pH Standards", Course on "Advances in

Instrumental Analysis", Walter Reed Army Institute

of Research, Walter Reed Army Medical Center,

Washington, D.C. March 11, 1965.

10. Robinson, R.A., "Thermodynamics of Multi component

Systems", Department of Chemistry, University of

South Carolina, Columbia, South Carolina.

March 12, 1965.

97
11. Bates, R.G., "Acids and Bases In Alcohol-Water

Solvents", Research Seminar, Rohm and Haas Company,

Philadelphia, Pennsylvania. March 24, 1965.

12. Bates, R.G ., Paabo, M., Woodhead, M., Robinson, R.A.

"Behavior of Weak Bases in Methanol-Water Solvents",

American Chemical Society, National Meeting,

Detroit, Michigan. April 7, 1965.

13. Davis, M.M., "Titrimetric and Equilibrium Studies Using

Indicators Related to Nile Blue A", American Chemical

Society, National Meeting, Detroit, Michigan.

April 8, 1965.

14. Robinson, R.A., "Properties of Mixed Salt Solutions",

Rensselaer Polytechnic Institute, Troy, N.Y.

April 12, 1965

15. Bates, R.G., "Behavior of Acids and Bases in MethanolWater Solvents", Chemistry Seminar, University of

Kentucky, Lexington, Kentucky. April J>0, 1965.

16. Gary, R., "Reference Solutions for a pD Scale",

Chemical Society of Washington, University of Maryland

May 7, 1965.

17. Bates, R.G., "The Conceptual Basis of Practical pH

Measurements in Water and Other Amphiprotic Solvents",

Chemical Institute of Canada, McGill University,

Montreal, P.Q., Canada. June 1, 1965.

98
14. LIST OF REFERENCES

Bates, R. G., Paabo, M., and Robinson, R. A.,

J. Phys. Chem. 67, 1833 (1963).

Ong, K. C, Robinson, R. A., and Bates, R. G.,

Anal. Chem. 36, 1971 (1964).

Paabo, M., Robinson, R. A., and Bates, R. G.,

J. Chem. Eng. Data, % 374 (1964).

Paabo, M., Bates, R. G., and Robinson, R. A.,

Anal. Chem. 37, 462 (1965).

Paabo, M., Robinson, R. A., and Bates, R. G.,

J. Am. Chem. Soc, 87, 415 (1965).

Gary, R., Bates, R. G., and Robinson, R. A.,

J. Phys. Chem., 68, 1186 (1964).

Gary, R., Bates, R. G., and Robinson, R. A.,

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

Gary, R., Bates, R. G., and Robinson, R. A.,

"Dissociation constant of acetic acid in deuterium

oxide from 5 to 50°. Reference points for a pD scale",

J. Phys. Chem. (in press).

Bates, R. G. and Bower, V. E.,

J. Res. Natl. Bur. Std., 53, 283 (1954).

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

J. Phys. Chem., 66, 1423 (1962).

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

J. Phys. Chem., 68, 1929 (1964).

Owen, B. B., J. Am. Chem. Soc, 57, 1526 (1935).

Harned, H. S. and Robinson, R. A.,

Trans. Faraday Soc, 37, 302 (l94l).

Crossley, M. L., et al., J. Am. Chem. Soc,

14, 573, 578 (1952).

Davis, M. M. and Hetzer, H. B.,

J. Res. Natl. Bur. Std., 60, 569 (1958).

99
[16] Hantzsch, A., Chem. Ber., 55, 953 (1922).

[17] Reichardt, C, Angew. Chem., Intern. Ed. Engl., 4,

29 (1965).

[18] Davis, M. M. and Hetzer, H. B., J. Res. Natl. Bur. Std.,

54, 309 (1955).

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

J. Res. Natl. Bur. Std., 68A, 305 (1964).

[20] Woodhead, M., Paabo, M., Robinson, R. A., and

Bates, R. G., J. Res. Natl. Bur. Std., 69A , 263 (1965).

[21] Bates, R. G., and Hetzer, H. B., J. Phys . Chem., 65 , 667 (1961); Datta, S. P., Grzybowski. A. K., and

Wilson, B. A., J. Chem. Soc. (London), 792 (1963).

[22] Everett, D. H., and Wynne- Jones, W. P. K., Trans.

Faraday Soc. 48, 531 (1952).

[23] Bower, V. E., and Robinson, R. A.,

J. Res. Natl. Bur. Std., 69A, 131 (1965).

[24] Robinson, R. A. and Bower, V. E.,

J. Res. Natl. Bur. Std., 69A, 19 (1965).

[25] Robinson, R. A. and Bower, V. E., "Properties of

aqueous mixtures of pure salts: Thermodynamics of

the ternary system water-potassium chloride-barium

chloride at 25 °C", J. Res. Natl. Bur. Std.,

(in press)

[26] Robinson, R. A. and Bower, V. E., "An additivity rule

for the vapor pressure lowering of aqueous solutions,

J. Res. Natl. Bur. Std., (in press).

[27] Robinson, R. A. and Bower, V. E., "Osmotic and activity

coefficients of tris(hydroxymethyl)aminome thane and

its hydrochloride in aqueous solution at 25 °C",

J. Chem. Eng. Data, (in press).

[28] Kolthoff, I. M., and Willman, A., J. Am. Chem. Soc,

56, 1007 (1934).

[29] Hutchison, A. W., and Hoover, T. B.,

"Research and Development on Water Content of

Concentrated Nitric Acid", Pinal report,

Contract DA-36-O6I-ORD-326 (1955).

100


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