onal Bureau 01 avauuo*"

jibrary, N.W. Bldg

NBS

400

ELECTROCHEMICAL ANALYSIS:

STUDIES OF ACIDS, BASES, AND

SALTS BY EMF, CONDUCTANCE,

OPTICAL, AND KINETIC METHODS

JULY 1965 TO JUNE 1966

: 1

I

t +** ° F *oA

Edited by Roger G. Bates

*<">cau of

'

U.S. DEPARTMENT OF COMMERCE

National Bureau of Standards

IBS I

Hi'.:::.:
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UNITED STATES DEPARTMENT OF COMMERCE • John T. Connor, Secretary

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

NBS TECHNICAL NOTE 400

Issued September 6, 1966

ELECTROCHEMICAL ANALYSIS: STUDIES OF

ACIDS, BASES, AND SALTS BY EMF, CONDUCTANCE,

OPTICAL, AND KINETIC METHODS

JULY 1965 TO JUNE 1966

Edited by Roger G. Bates

Electrochemical Analysis Section

Analytical Chemistry Division

Institute for Materials Research

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

listed or referred to in the open literature.

For sale by the Superintendent of Documents, U.S. Government Printing Office

Washington, D.C., 20402 - Price 55 cents
FOREWORD

The Analytical Chemistry Division was established as

a separate division at the National Bureau of Standards on

September 1, 196j>, and became part of the Institute for

Materials Research in the February 1, 1964, reorganization.

It consists at present of nine sections and about 100 technical personnel encompassing some 45 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 our 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 a Technical

Note. In 1966 we plan to issue these summaries for all of

our sections. The following is the second annual report on

progress of the Electrochemical Analysis Section.

W. Wayne Meinke, Chief

Analytical Chemistry Division

ii
PREFACE

This is the second in a series of annual progress

reports of the Electrochemical Analysis Section of the

Analytical Chemistry Division. The report covers the

fiscal year 1966, which "began on July 1, 1965, and ended

on June 30, 1966.

Many of the processes and reactions of analytical

interest take place in solutions, and a large fraction of

these involve ionized solutes. If the research programs

of the Electrochemical Analysis Section were to "be placed

in a single broad category, undoubtedly "Solution Electrochemistry" would "be a. fair choice, with primary emphasis

on acid-base phenomena, solvent effects on the behavior of

electrolytes, and potentiometry with reversible electrodes.

Competence in polarography and coulometry exists elsewhere

in the Analytical Chemistry Division] hence, these areas

are not a part of the research activity of the Electrochemical Analysis Section. In line with a uniform policy

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

and sample aspects. During the fiscal year Just ending,

about 70 percent of the total effort was devoted to research, while 20 percent was devoted to programs on Standard

Reference Materials and 10 percent to Other-Agency programs.

The outstanding event of the present year was the longawaited move to the excellent new facility at Gaithersburg,

Md. The move and the attendant loss of time during resettlement have inevitably left their mark on the Section's

activity* More serious, however, has been a shortage of

personnel. Two project leaders, Dr. Robert Gary and Dr.

Richard K. Wolford, were chosen as Science and Technology

Fellows and were assigned elsewhere in the Department of

Commerce for 10 months of the reporting period.

iii
A third, Dr. Marion M. Davis, retired from the Section on

December J51, 1965. On the other hand, Dr. Paul W. Schindler

spent nine months in the Section as a guest worker supported

by the Swiss National Foundation.

The purpose of this report is to summarize the broad

program of the Electrochemical Analysis Section and to

convey also the manner in which the individual projects

contribute to the whole. An attempt is made to set forth

in a rather complete way the entire year's activity of the

Section and to reveal the ways in which this specialized

group contributes to the missions of the Division and

Institute of which it is a part.

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 Rauderibush in the preparation of this report

is gratefully acknowledged.

Roger G. Bates, Chief

Electrochemical Analysis Section

Washington, D.C.

June , 1966

IV
TABLE OF CONTENTS

PAGE

1. FACILITIES AND EQUIPMENT 1

The Gaithersburg Laboratories ........ 1

Plans and Needs 5

2. MEASUREMENT OF ACIDITY 5

Concepts and Program 5

pH Standards 10

Redetermination of the pH of Potassium

Hydrogen Phthalate, Primary pH Standard 10

Method ..... 10

Procedures 11

Results 11

Certification of a New Lot of Potassium

Dihydrogen Phosphate ..... 1^

Sodium Hydrogen Succinate, a New Standard

Reference Material for pH Measurements . 15

Standard pH Values for 4-Aminopyridine Buffers

in 50 Percent Methanol 17

Standards for a pD Scale in Deuterium Oxide. . 19

A Scale of Acid Strengths in Benzene 21

Experimental 21

Discussion 22

3. SOLVENT EFFECTS ON ACID-BASE PROCESSES .... 25

Dissociation of Positively-Charged Acids

in Methanol-Water Solvents 25

Ammonium Ion 28

Tris (hydroxymethyl )methylammonium

Ion 30

Solubility Studies of the Medium

Effect ^2

Parameters for the Solvent Effect. . 35

Isotope Effects on the Dissociation of

Acetic Acid in Water and Deuterium Oxide . . 59

v
PAGE

Acid-Base Behavior in Organic Solvents with

Low Dielectric Constants . . 45

Spectrophotometry Studies of the Dissociation

of Phenols 46

4. EQUILIBRIUM STUDIES OP SYSTEMS OF ANALYTICAL

INTEREST 48

Thermodynamics of the Dissociation of

Morpholinium Ion in Water ......... 48

Ionization Constant of Deuterium Oxide

from 5 to 50 °C 51

Response of the Glass Electrode in

Heavy Water 56

5. THERMODYNAMIC STUDIES OF AQUEOUS SOLUTIONS

OF MIXED SALTS . . . 60

6. CONDUCTOMETRIC DETERMINATION OF TRACES OF

WATER 63

Objectives and Principle of Method ...... 63

Modification of Apparatus 63

Calibration with Water 68

Analysis of Errors ..... 71

Status of the Project ...... 73

7. KINETIC TECHNIQUES IN TRACE ANALYSIS 74

8. STANDARD REFERENCE MATERIALS FOR ACIDIMETRY. . 76

Homogeneity Tests of Potassium Hydrogen

Phthalate 84h 76

Certification of Potassium Hydrogen Phthalate

84h as an Acidimetric Standard 78

Status of Certification Procedures for

Standard Reference Materials . 82

9. FOREIGN- CURRENCY PROGRAMS 84

10. PERSONNEL 85

VI
PAGE

11. COMMITTEE MEMBERSHIPS 86

12. PUBLICATIONS AND MANUSCRIPTS,

JULY 1965 TO JUNE 1966 88

13. TALKS, JULY I965 TO JUNE 1966 91

14. LIST OF REFERENCES 94

vll
LIST OF FIGURES

FIGURE NO. PAGE

1. Equipment for the emf measurements on

which pH standards in aqueous solution

are "based . 16

2. Relative acidic strengths in benzene and

water at 25 °C . The strength in

benzene is expressed as the value of

log K for 1:1 association of the acid

(HA) with 1, 3-dlphenylguanidine (DPG) . 24

3. pK values for acids of two charge types

in methanol-water solvents at 25 °C . . 27

4. Medium effects ( my) for ammonia, tris,

hydrochloric acid, and alkali halides

in methanol-water solvents at 25 °C . . 36

5. Observed solvent effects on the pK of

protonated tris and ammonia (solid

lines) compared with the solvent effect

calculated "by eq. ~LJ> (broken lines) . . 37

6. pK 1 in morpholine-morpholinium chloride

"BtLffer solutions as a function of ionic

strength (m

1

) at 0, 25, and 50 °C . . . 49

7. Apparatus for the electrolytic production

of the sodium amalgam used to prepare

the standard solution of NaOD 53

8. Equipment for a direct comparison of* the

deuterium gas electrode and the glass

electrode in deuterium oxide. ..... 58

9. Variation of activity coefficients («y)

with composition in an aqueous mixture

of calcium chloride (B) and magnesium

chloride (o) of total molality 4 . . . 62

10. Conductance cell and circulation system

for the determination of traces of

water 65

11. Schematic wiring diagram for conductance

bridge .... ....... 66

12. Recorder trace of calibrations by

resistance steps and by additions

of water 70

vili
LIST OF FIGURES (con't)

FIGURE NO* PAGE

l^. Comparison of calibration methods .... 72

14. End point determination in a typical

coulometric titration . 80

LIST OF TABLES

TABLE NO. PAGE

1. Standard pH values of 0.05 m potassium

hydrogen phthalate 12

2. Comparison of standard pH values, pH(s) , for present phosphates with those

for the newly certified materials ... 1^

3. Values of pH*(s) for buffer solutions

composed of 4-aminopyridinium

chloride (m) and 4- aminopyridine (m)

in 50 percent methanol . • 19

4. pap values of the buffer solutions 0.05 m

acetic acid, 0.05m sodium acetate

in heavy water 20

5. Acidic dissociation constant of ammonium

ion in methanol-water solvents at

25 °C 29

6. pa^ for buffer solutions of ammonia and

ammonium chloride (each at molality m)

in methanol-water solvents at 25 °C . . 29

7. pK for the dissociation of protonated

tris(hydroxymethyl)aminomethane in

methanol-water solvents at 25 °C.

The solvent effect 32

8. Medium effects for tris(hydroxymethyl)-

aminomethane in methanol-water solvents,

obtained from solubility measurements . jk

9. Thermodynamic quantities for the dissociation of acetic acid and deuterioacetlc

acid in water and deuterium oxide

at 25 °C . 43

ix
LIST OF TABLES (con't)

TABLE NO. PAGE

10. Increase in pK for six substituted

phenols resulting from a change of

solvent composition from pure water (w)

to 50 wt . percent methanol at 25 °C , . 47

11. Thermodynamic quantities for the dissociation of morphollnium, piperidinium,

and pyrrolidinium ions in water

at 25 °C 51

12. Comparison of the ionization constants

of deuterium oxide and water ..... 55

13. Comparison of the changes in enthalpy,

entropy, and heat capacity for the

ionization of deuterium oxide and water

at 25 °C (molal scale) ........ 56

14. Assay of potassium hydrogen phthalate

84h for homogeneity ........... 77

15. Assay of potassium hydrogen phthalate

84h for certification as a standard

for acidimetry ............ 8l

x
ELECTROCHEMICAL ANALYSIS: STUDIES OF ACIDS, BASES,

AND SALTS BY EMF, CONDUCTANCE, OPTICAL, AND

KINETIC METHODS. JULY I965 TO JUNE 1966

Edited by Roger G. Bates

ABSTRACT

This survey of the activities of the Electrochemical

Analysis Section, Analytical Chemistry Division, covers the

period July 1965 to June 1966. 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. A description

of facilities and equipment is presented and the directions

of new programs are indicated. Brief summaries of the

several lines of work now under way are given. The main

areas include the study of acidity measurement in water,

deuterium oxide, and methanol-water solvents, along with

the development and certification of standard reference

materials with which acidity scales in these media can be

defined. Solvent effects and isotope effects are being

studied, and equilibrium data for certain systems of analytical interest have been, obtained. Progress on a conductometric method for determining traces of water is reported,

and the thermodynamic properties of some salt mixtures have

been measured by a vapor pressure method. Kinetic methods

for trace analysis have been examined, and a new lot of

potassium hydrogen phthalate (Standard Reference Material

84h) has been certified for use as an acidimetric standard.

The survey concludes with lists of the members of the

Section staff, publications and manuscripts produced during

the year, talks given by the staff, and committee assignments.

xi

1. FACILITIES AND EQUIPMENT

A. The Gaithersburg Laboratories

On March 23, 1966, the Electrochemical Analysis Section

began the move to Its assigned quarters at the Bureau's

Gaithersburg site. By the end of the same month most of the

equipment had been moved and the job of reconstructing

apparatus and of adapting to the new surroundings had begun.

The Section was assigned space on a single corridor of

the Chemistry Building, one of the seven general-purpose

laboratory buildings. Twelve modules of general laboratory

space and seven smaller modules (with north light) devoted

to offices and limited laboratories fell to the Section.

The total area provided for the Section was a little more

than 3,600 square feet of which slightly less than 300

square feet were given over to the activities of the Section

Office.

Pour of the larger laboratory modules are temperaturecontrolled, nominally to ± 1 °F. In these laboratories

will be carried out those measurements where precise temperature control (to ± 0.001 °C) will be required. All

other offices and laboratories are controlled nominally

to ±2 °F.

The Section has more hood space at the new site than

at the old. Against three hoods at the old quarters in

Washington, the Section now has seven hoods, with provision

for three elephant -trunk hoods in the larger general-purpose

modules. The elephant-trunk hoods are expected to be suspended over oil thermostats which may be operated at high

temperatures

.

The equipment available for use in conductometric,

spectrophotometric, isopiestic, and potentiometric measurements remains essentially the same as that described in
last year's progress report [1]. The equipment for the precise measurement of emf has, however, been augmented by the

addition of a Radiometer pH M4 high- impedance electrometer.

Furthermore, the constant-temperature baths used on the

three emf assemblies are being replaced by new ones. These

are of stainless-steel construction with built-in refrigeration. Temperature control is provided by external electronic controllers of the proportional type, provided with

thermistor sensors.

It is anticipated that a fourth emf assembly will be

added to the inventory of the Section. The new instruments

will be devoted principally to the study of the behavior of

electrolytes dissolved in mixed media consisting of water

and organic solvents,

(V. E. Bower)
B. Plans and Needs

The measurement of acidity has been a major activity

of the Section for many years, and it is here that the

Section's competence has been developed to the highest level,

However, the concept of a practical acidity scale for partially aqueous and nonaqueous solvents has only recently

been developed, despite the fact that techniques for acidity

measurements are critically needed for many media of fixed

composition. It is hoped therefore to continue and accelerate this aspect of the program. With the acquisition

of new constant-temperature baths, the facilities for

studies with hydrogen and deuterium electrodes are excellent,

Research with the glass electrode, to which little attention

has been given in the past, should be expanded, and for this

purpose a precision high-impedance electrometer of the

vibrating-reed or dynamic-condenser types is needed. Measurements in sea water have taken on a new significance with

the recent emphasis on oceanographic research, and studies

of acid-base behavior in this medium will shortly be initiated.

The establishment of a new acidity scale and the selection of reference materials needed to define it require an

extensive series of emf measurements over a range of solute

concentrations, solvent compositions, and temperatures.

A completely automated system which would program the temperature changes, monitor the approach of the cells to

equilibrium, and record all of the essential data has been

designed. It was described in the annual progress report

for 1964-65 [1]. As the broad program of acidity studies

takes form, this type of facility will prove indispensable

and should be constructed soon.

3
The need for accurate pH measurements at high temperatures and. pressures is increasingly being felt. No standards

of reference are at present available for temperatures above

95 °C . The construction of equipment for precise emf measurements of solutions under pressure in the range 100 to 200 °C

should be undertaken in the Section in the near future.

The Section now has competence in thermodynamic studies

of solution equilibria, in the conductance of electrolytic

solutions, and in the study of the rates of homogeneous

reactions in solution. There is a need to supplement these

studies with direct calorimetric measurements of the enthalpies of solution reactions. Furthermore, calorimetric

measurements are assuming a role of increased importance in

analytical chemistry. The Division would, for example,

benefit from the development of competence in enthalpimetric

titrimetry. Furthermore, the new technique of microcalorlmetry is revealing exciting new possibilities in rapid

analysis and in the investigation of reactions in biological

systems for which conventional methods of analysis are completely unsuitable. To begin work in this new area, we

should acquire equipment for microcalorimetric studies in

the near future.

(R. G. Bates)

4
2. MEASUREMENT OP ACIDITY

A. Concepts and Program

Studies performed in the Section over a period of more

than two decades have led to the establishment of the NBS

conventional activity scale of pH. This scale has five

fixed points represented by aqueous solutions of the standard

reference materials issued by the National Bureau of Standards

The development of this pH scale has been described in detail

elsewhere [ 2]

.

In establishing a useful pH scale, one comes face to

face with the dilemma that the experimental methods for

acidity measurement do not yield quantities of exact physical

definition, whereas the hydrogen ion concentration and other

exactly defined quantities cannot be measured accurately In

solutions of unknown composition. The emf cell with hydrogen

electrode and a calomel reference electrode, or the modern

counterpart of this cell, the glass electrode pH meter,

develops an emf that is a function of the acidity of the

solution. This hydrogen ion response depends on the activities of the substances In the solution but is not clearly

related to the activity of the hydrogen ion. Indeed, the

activity of a single ionic species is not uniquely defined

by thermodynamics. It is still possible, however, to establish a conventional scale of hydrogen ion activity through

the arbitrary choice of a numerical scale for the activity

coefficient of a single species of ions.

The NBS pH scale, which is gaining favor gradually

throughout the world, has its experimental base in accurate

emf measurements of cells without liquid junction of the

type

H

2 (g. 1 atm), Buffer soln., Cl", AgCljAg
The standard emf of this cell is known accurately from

to 60 °C and with reduced accuracy to 95 °C; other measurements have extended as high as 270 °C . By the strict application of thermodynamics, therefore, it is possible to

determine the acidity function p(cUjYrin ) in buffer-chloride

solutions and, by an extrapolation procedure, in the

chloride-free buffer solution, if need be. The standard

values of the hydrogen ion activity (pa^) are easily derived

by the relationship

paH

= p(aHYcl ) + log yc1 (l)

provided an estimate of yri can be made. This is accomplished through the adoption of a non-thermodynamic convention. For this purpose, representatives of the British and

American groups engaged in pH standardization have agreed

to accept a particular form of the Debye-Huckel equation

relating y-,, to the ionic strength, I [J>]:

AT 1/2

"log YC1 = 1/2 (2)

where A is a constant of the Debye-Huckel theory and p has

the value 1.5 in water at 25 °C.

It must be recognized that, in theory, there is no

single scale of hydrogen ion activity. An arbitrary or

conventional scale defined in this way will, however, meet

all of the practical requirements. Coupled with the scale

of single-ion activity coefficients fixed by the convention

adopted above (eq. 2), experimental pH values determined

under ideal conditions are completely consistent with ionization constants and other thermodynamic data for the

acid-base systems which fix the pH. It is evidently of

primary importance that the same conventional scale be
adopted by everyone who measures and records pH values.

For this reason, it is imperative that procedures for pH

standardization be recommended by a national standardizing

body such as the National Bureau of Standards or by an

authoritative international standards group such as the

International Standards Organization (ISO) or the International Union of Pure and Applied Chemistry (IUPAC). The

NBS approach to a standard scale of pH has received the

endorsement of the IUPAC and of the American Society for

Testing Materials (ASTM)

.

There is a persistent demand on the part of science

and industry for practical scales of acidity in nonaqueous

solvents and in mixtures of water and organic solvents.

During the past few years, some effort has been directed by

the Section's staff to providing a useful experimental

scale for acidity in alcohol-water media and in heavy water.

The glass electrode has been found to respond satisfactorily

to deuterium ions in heavy water and to hydrogen Ions in

aqueous methanol and aqueous ethanol, at least in solvents

containing no more than 90 wt . percent of the organic

component. It remained to show that the liquid- junction

potential between an aqueous calomel reference electrode

and a nonaqueous or partly aqueous buffer solution was not

a function of the acidity of the latter. Satisfactory

evidence of the constancy of the liquid Junction was obtained for methanol-water mixtures containing as much as 70 wt

.

percent methanol [4].

These studies made it clear that the glass electrode

pH meter is capable of yielding values for the relative

hydrogen ion activity of two solutions both of which have

the same solvent composition. To set up a scale, therefore,

one or more reference points are needed. For practical

reasons, the pH unit pag was chosen. This quantity is

defined as follows:

7
pa* = -log m^^ D)

where m is molality and y„ is the molal activity coefficient based on the standard state in the nonaqueous or

partly aqueous solution (s); in other words, yv becomes

S XI equal to unity at infinite dilution in each particular

solvent or solvent mixture in question. The procedures

used for the establishment of standard values of pH can be

easily extended to yield standard values of pH*. To do so,

the convention which furnishes numerical values of yni is

entirely similar to that used for y„-. in aqueous solutions,

namely eq. 2. Appropriate changes in the constants A and p

are made to reflect the different dielectric constants and

densities of the solvents.

Thus far the solvents selected for study have been

chosen because of their practical importance and freedom

from undesirable side reactions with the hydrogen electrode

or the silver-silver chloride electrode which would impair

the accuracy of the measurements. It is planned to extend

the work to other alcohol-water mixtures and to aqueous

ketones and aqueous amides in the order of the importance

of these solvents in acid-base studies. Attention will

then be directed to certain acidic or basic solvents such

as acetic acid and ethylenediamine. The procedures will

become more laborious when an attempt is made to study

solvents with dielectric constants below the range 30 to 40.

Under these circumstances, supplementary measurements,

possibly by conductance methods, of ion-pair association

constants will be necessary.

8
These procedures are most likely to be successful

when the solvents are amphlprotlc . At the present time

It seems unlikely that methods based on emf measurements

will succeed with aprotic (inert) media, where the dielectric constant is usually low and ionization is slight.

Nevertheless, the measurement of acidity in media of this

sort is of very real practical importance. At the present

time it appears that a different approach to the establishment of acidity scales in aprotic media will have to be

found

.

(R. G. Bates)

9
B. pH Standards

1. Redetermination of the pH of Potassium Hydrogen

Phthalate, Primary pH Standard . a. Method . The assignment of standard values of

pH to the NBS primary standard buffers consists essentially

of four steps. The first of these is the determination of

the acidity function p^^y-,-, ) by measurement of the emf

of hydrogen-silver, silver chloride cells without liquid

junction, containing the buffer solution and three or more

different concentrations of soluble chloride. This function

is calculated by

t \ ( E - E °) Fn ,uy Pla^Tm ) = — + log nu-. (4) n U1 RT In 10 U1

in which a is activity, y the molal ionic activity coefficient, E the measured emf, E° standard emf of the cell,

F the faraday, and m molality. These values of p(a„Yp-, ) are then extrapolated to zero chloride concentration to

obtain the quantity p(a~yYrfi ) . The pa^ is then computed

by the relationship

paH

= p(aHYcl

)° + log ycl (5)

The Bates- Guggenheim convention is used to evaluate log Ynn

(see eq. 2, section 2. A). Finally, the standard reference

value pH(s) is identified with the conventional pa^ value

of the buffer.

10
b. Procedures . Except for a few changes in

technique, the usual procedures for emf measurements were

followed. As is the usual practice, hydrogen electrodes

coated with palladium "black were used to avoid reduction

of the phthalate by hydrogen in the presence of platinum

black. Several different palladium solutions were used,

including one prepared from palladium of exceptionally

high purity. No differences in behavior were observed

among electrodes prepared from the different solutions.

The potassium hydrogen phthalate used was NBS Standard

Reference Material l85d.

The cells were measured first at 25 °C, then from

to 35 °C at five-degree intervals in one series. A different set of cells was then measured at 25 °C and from 40

to 60 °C. Final checks were always made at 25 °C . All of

the cells contained 0.05 ni potassium acid phthalate and

either 0.005, 0.010, or 0.015 m potassium chloride.

c. Results . The linear extrapolation of values

of p ( a-rrYp -, ) to obtain the quantity p(a,/Yp-| ) at each temperature was carried out by the method of least squares.

The ionic strengths, I, were computed from table 2 of the

paper of Hamer, Pinching, and Acree [5]. The values of

pH(S) obtained as described were fitted by the method of

least squares to an equation of the form

pH(S) = ~ + B + CT + DT2

(6)

with A = 2042.45, B = -1^.0228, C = 0.044210, D = -3.3761

-5

x 10 . The standard deviation o at the 13 temperatures is 0.0010.

-5

x 10 . The standard deviation of the calculated values

11
The results are shown In table 1. Values of pH(S)

now In use, based on the experimental data of Hamer and

Acree [5,6] are given in the last column for comparison.

Table 1. Standard pH values of 0.05 m potassium hydrogen

phthalate.

t °c p(aHYcl

)° -log yc1 paH

a pH(S) V(s)

4.096

Q

0.084 4.012 4.012 4.003

5 4.089

9

.085 4.005 4.005 3.999

10 4.087

5

.086 4.002 4.002 3.998

15 4.0876 .086 4.001 4.001 3.999

20 4.089

9

.087 4.003 4.003 4.002

25 4.096 .088 4.009 4.008 4.008

30 4.102

4.1094

.088 4.014 4.014 4.015

35 .089 4.020 4.023 4.024

40 4.124 .090 4.034 4.033 4.035

45 4.1356 .091 4.045 4.045 4.047

50 4.149

9

.091 4.059 4.058 4.060

55 4.165

5

.092 4.073 4.073 4.075

60 4.l8l

9

.093 4.089 4.089 4.091

a Calculated by eq. 6

b Present NBS values.

(H. B. Hetzer)

12
2. Certification of a New Lot of Potassium

Dihydrogen Phosphate

For acid-base studies, the equimolal aqueous mixture

of potassium dihydrogen phosphate and disodium hydrogen

phosphate provides a most useful reference solution. These

two salts, packaged separately, are distributed as Standard

Reference Materials by the National Bureau of Standards.

A standard buffer solution prepared according to the directions on the- certificate accompanying these materials provides an accurately defined reference point near the middle

of the pH scale.

The chemicals which are used for these reference

materials are analytical reagents of high grade, certified

only for the pH value they yield when mixed and prepared

according to directions. Because these salts are bought

in rather large lots by the Bureau, recertification is only

occasionally required. The high quality of these salts is

an assurance that the pH value will change very little with

change of lot. Nonetheless, because the pH standards are

certified to the highest precision that present techniques

and theory permit, minor variations in pH between lots

occur, and a detailed restudy is required for certification

of each new lot.

During the last year such a study was undertaken because a new lot of potassium dihydrogen phosphate had been

acquired. The study shows that standard buffers prepared

with the new lot will have a lower pH, by a few thousandths

of a unit, than buffers made with the old material. The

tentative results listed in table 2 show the change of pH

that has been observed. The values given relate to the

solution containing the two phosphates, each at molality

of 0.025.

13
Confirmatory measurements are now under way, and the final

results will be listed in the certificate issued with the

new lot of potassium dihydrogen phosphate.

Table 2. Comparison of standard pH values, pH(S), for

present phosphates with those for the newly

certified materials.

t °c PH(S), present materials Tentative pH(S), new lot

6.984 6.982

5 6.951 6.949

10 6.923 6.922

15 6.900 6.897

20 6.881 6.876

25 6.865 6.862

30 6.853 6.849

35 6.844 6.840

40 6.838 6.834

45 6.834 6.831

50 6.833 6.831

(V. E. Bower)

14
5. Sodium Hydrogen Succinate, a New Standard Reference

Material for pH Measurements

Sodium hydrogen succinate in aqueous solution provides

useful buffer solutions with pH near 4.7. Mixed with either

succinic acid or dlsodlum succinate, it can be used to make

buffers whose pH varies, according to the mixing proportions, from about 3.5 to about 6.0. The relative ease of

purification of these three compounds and the pH range which

they cover has led to their widespread use as reference

buffers for the determination of the dissociation constants

of weak acids.

For the standardization of the glass electrode in the

slightly acid pH range, potassium hydrogen phthalate solutions (pH about 4) are very often used and are indeed among

the most reliable standard buffers available. However, as

a buffer medium in the spectrophotometric determination of

dissociation constants or in spectrophotometric analysis,

potassium hydrogen phthalate is of limited usefulness because it absorbs strongly in the ultraviolet region of the

spectrum. Furthermore, potassium hydrogen phthalate is

unstable in contact with the platinum-hydrogen electrode.

Solutions of sodium hydrogen succinate have pH values

not greatly different from those of potassium hydrogen

phthalate, yet they have the great advantage of stability

in contact with hydrogen and finely-divided platinum.

Inasmuch as succinate buffers are free of the difficulties

outlined above, it was considered desirable to add this

new pH standard to those which define the NBS conventional

pH standard scale.

Samples of sodium hydrogen succinate were prepared and

measurements begun this year. Preliminary results point to

a stable well-behaved system capable of precise measurement.

15
The measurements of the kind required for the establishment of a pH standard are carried out on a potentlometric

assembly such as that shown in figure 1. This arrangement

is temporary, but the appearance of the permanent station

will not differ greatly from that shown in the figure.

(V. E. Bower)

Figure 1. Equipment for the emf measurements on which

pH standards in aqueous solution are based.

16
C . Standard pH* Values for 4-Aminopyridine Buffers in

30 Percent Methanol

As a continuation of the development and standardization of an acidity scale for amphlprotic media, pa^t values

In 50 percent methanol were determined for a fifth buffer

In this series, namely for a mixture of 4-aminopyridine

and Its hydrochloride. This "buffer has found use In the

pH range 8.5 - 9.5 In place of the borax buffer when the

latter has undesirable side reactions. 4-Aminopyridine Is

a solid base which can easily be purified. It is moderately

soluble in a variety of solvents such as water, alcohol,

benzene, toluene, and chloroform. It has also been proposed

as a primary acidimetric standard.

A method for determining standard values for pH measurements in aqueous buffer solutions has been developed in

this laboratory, and the same general procedure was followed

in 50 percent methanol medium for the five buffer systems

that have been studied.

The 4-aminopyridine was the same material used for

determining the dissociation constant of this base in water

in this laboratory [7]. A stock solution of the 4-aminopyridine and standard hydrochloric acid was prepared in

50 percent methanol. This was diluted further with water

and methanol to yield buffer mixtures with a total ionic

strength varying from 0.1 to 0.01. The acidity function

p (attT^-, ) was measured over a temperature range from 10

to 40 °C with the aid of the cell

Pt;H

? (g., 1 atm.), Buffer in 50 percent methanol,

AgClj Ag

17
It was calculated from the emf of this cell and the

standard emf of the hydrogen-silver chloride cell in

50 percent methanol "by the use of the eq. 4 of section

2.B.I. The standard emf, E°, in 50 wt. percent methanol

has already been determined and the results published

[8], By using a convention for evaluation of the singleion activity coefficient, Ypv analogous to that used in

water (eq. 2 of section 2. A), the pa* values were obtained

from the experimental acidity functions by the relationship

paH

= P s(

aH^Cl) + l0g

B

(*VC1 ) (?)

These pa|t values are given in table ~5 t where they

are identified with pH*(S), the standard reference values

in the operational definition of the pH* for alcohol-water

solvents. These values can be used for practical measurements of pH* in 50 percent methanol by standardizing the

pH meter with a reference solution of the same solvent

composition as that of the test solution. This procedure

avoids to some degree the variation of the asymmetry potential of the glass electrode due to changes in solvent

composition when the electrode is transferred from one

solution to another of different water content.

(M. Paabo)

18
Table 3. Values of pH*(S) for buffer solutions

composed of 4-aminopyridinium chloride (m)

and 4-aminopyridine (m) in 50 percent methanol.

m 10 °c 25 °C 40 °C

0.02 9.048 8.629 8.262

0.04 9.088 8.668 8.305

o 06 9.116 8.695 8.332

0.08 9.157 8.715 8.353

0.10 9.155 8.732 8.368

D. Standards for a pD Scale In Deuterium Oxide

The procedures described in earlier sections of this

report, which have led to standard reference data for a pH

scale in water and a pH* scale in alcohol-water solvents,

can readily be adapted to the definition of a pD scale for

deuterium oxide (heavy water). The cell used for measurements of this type is

Pfc;D 2 (g, 1 atm.), Buffer soln. containing 01 in DpO,

AgGlj Ag

The acidity function p(a

T) Yr

, ]

) in the buffer solution can be

determined, inasmuch as E°, the standard emf of the cell,

is known [93- The convention for Yp-i, needed to derive pa~

from p(a

T) Yp-| )t is "that proposed by Bates and Guggenheim [3]

see eq. 2 of section 2. A), in which A and p are modified

to reflect the difference between the dielectric constants

and densities of ordinary water and heavy water.

19
It is thus possible to define a pa— scale for solutions in deuterium oxide. For this purpose, pa„ has

already been established over a range of temperatures for

the equimolal mixture of potassium dideuterium phosphate

(0.025 tn) and disodium deuterium phosphate (0.025 rn ) in

deuterium oxide as solvent [10].

A further pa^ standard has now been established -- the

equimolal mixture of acetic acid (0.05 m) and sodium acetat

(0.05 m) in deuterium oxide [11]. Measurements of the

above cell with this buffer solution and a deuterium gas

electrode gave values of p(anyr

, 1 )

, and extrapolations to

zero chloride concentration gave p(anYr

,-

) )

. Using values

of A and P in the Bates -Guggenheim convention for log

Ypi appropriate to the density and dielectric constant of

heavy water and to the temperature, values of pa~ were

calculated. These are given in table 4, They may be regarded, along with those for the phosphate and acetate

buffers, as standard values, pD(s), for an operational

scale of pD in deuterium oxide.

Table 4, pa- values of the buffer solution 0.05m acetic

acid, 0.05 m sodium acetate in heavy water.

t °c pa^ t °C paD

5 5.265 30 5.227

10 5.254 35 5.226

15 5.243 40 5.227

20 5.236 45 5.230

25 5.230 50 5.236

(R. A. Robinson)

20
E. A Scale of Acid Strengths in Benzene

Benzene is the most important of the hydrocarbon

solvents. Work on this project has been directed to

developing methods and standard reference materials for

determining relative strengths of acids and bases in

benzene, as well as for determining the total acid and

base content of benzene solutions. The results of these

studies have considerable generality; for example,, they

are applicable to other hydrocarbon media as well as to

halogenated hydrocarbons. Such media are often referred

to as "inert" or "aprotic" (the latter term signifying

that they have little or no ability to accept or donate

protons); another adjective often applied to them is

"differentiating"

.

A paper was prepared, based on unpublished experimental work by M. M. Davis and M. Paabo and bearing the

title "Comparative Strengths of Aliphatic Acids and Some

Other Carboxylic Acids in Benzene at 25°"; it has been accepted for publication by the Journal of Organic Chemistry . Excerpts from the abstract and discussion follow.

1. Experimental

In continuation of earlier measurements by Davis and

Hetzer [12] the strengths of the following acids (HA) in

benzene were determined: all of the straight-chain aliphatic acids containing from 2 to 20 carbon atoms except

the 15- and 19- carbon acids; isobutyric and trimethylacetic

(pivalic) acids; and o_-benzoylbenzoic, trans-cinnamic,

2,4- and 2, 5-dimethyl-2-furoic, and 1- and 2-naphthoic acids

Like most other procedures for comparing strengths of

Brizfasted (hydrogen) acids in aprotic solvents, the method

requires the use of a standard reference base and a standard indicator acid; it differs from other methods in

21
the choice of reference base and reference acid and also

in the treatment of data.

2. Discussion

In the method developed, the comparative strengths of

acids are denoted by values of equilibrium constants (KRTTA ) for the 1:1 association of acid (symbolized by HA) and

reference base (symbolized by B). The association product

is believed to consist of hydrogen-bonded ion pairs. The

logarithm of &„,, is regarded as being the analog of the

corresponding pK value for an aqueous solution of the acid.

(in the past, a more usual practice was to express comparative acidities in aprotic solvents as ratios, taking acetic

acid or benzoic acid as the basis for comparison.

)

The reference base adopted, 1, 2-diphenylguanidine, is

stronger than the reference bases ordinarily used. For

example, it associates readily with carboxylic acids, com4 7 monly giving equilibrium constants in the range 10 to 10.

By comparison, the extent of other hydrogen bonding reactions under the experimental conditions used is slight

and can be neglected. However, slightly more accurate

association constants can be obtained by correcting for

the carboxylic acid present as the dimer (double molecules),

and in the process of correction, a monomer-dimer equilibrium constant for the carboxylic acid can be deduced;

dimerization constants thus deduced compare favorably with

results by other methods.

The reference indicator acid employed, ~$ l , j5" , 5

' , 5

n - tetrabromophenolphthalein ethyl ester ( "bromophthalein

magenta E" ) is well suited for determining the strengths

of aliphatic and aromatic carboxylic acids in aprotic

solvents, because it matches them well in strength.

22
The results of this Investigation are In harmony with

the previous results, whose validity was Indicated by

clear-cut linear relationships between log K^tta values

(B = diphenylguanidine, solvent = benzene) and aqueous pK

values (and also acidic strengths in alcohols and partly

aqueous solvents). Figure 2 shows the separate linear cor- relations obtained in the latest work for meta- substituted

and ortho-substituted benzoic acids. o_-Benzoylbenzoic

acid is an interesting exception to the general behavior

in that it conforms to the "normal" (meta-) relationship

instead of to the special relationship found for the other

ortho-substituted acids; this exceptional behavior of

jo-benzoylbenzoic acid has not yet been explained.

This method of determining acidities in benzene Is

applicable to varied chemical types of acids and is also

useful for acids that are difficultly soluble, scarce In

amount, or unstable toward moisture.

(M. M. Davis)

23
Log K assocn (DPG, Benzene)

5

-

m-su

o-Benzoylbe

2-Naphthoi<

trans-Cinnamic-*

^CJCPj

4 L |Aliphotic Acids —i i__

substd

Acids

enzoic

benzoic

5 4 3

pK a (Water)

Figure 2. Relative acidic strengths in benzene and

water at 25 °C. The strength in "benzene

is expressed as the value of log K for

1:1 association of the acid (HA) with

1, >-diphenylguanidine (DPG).

Open circles, results from previous work

Half-filled circles, values for three

aliphatic acids (acetic, isobutyric,

trimethylacetic), this work

Filled circles, this work

24
3. SOLVENT EFFECTS ON ACID-BASE PROCESSES

A. Dissociation of Positively-Charged Acids in MethanolWater Solvents

The process by which a cationic weak acid dissociates

is one of particular interest. A positively charged acid

such as ammonium ion dissociates into an uncharged conjugate "base and a solvated proton by a reaction that can be

written in general terms as follows

BH+

+ SH = SH

2

+

+ B (8)

where SH represents an amphiprotic solvent. From the

electrostatic point of view, this reaction is a simple

one, as there is no net change in the number of ions. The

electrostatic contribution to the energy of dissociation

should be rather small, and this type of process is therefore well suited to a study of the parameters necessary to

account for the observed solvent effects on the dissociation. To do this, the electrostatic effect must be estimated in order to demonstrate whether it alone is sufficient

to explain the changes in dissociation constant actually

found.

For the past few years the solvent effect of methanol

on the dissociation of cation acids has been a subject of

interest in this laboratory. A detailed study of the dissociation of the protonated form of tris(hydroxymethyl)-

aminomethane (tris, THAM) in 50 wt. percent methanol from

10 to 40 °C was recently published [13]. Further work on

this acid in solvents containing 30, 70, and 90 wt. percent

methanol at temperatures of 15, 25, and 35 °C has now been

completed. In addition, the dissociation of ammonium ion

in five methanol-water solvents at 25 °C has been studied.

25
The typical pattern for the solvent effect of methanol

on the dissociation of a positively charged weak acid is

represented by the curves at the bottom of figure 3.

Methanol causes a decrease in the pK until a. minimum is

reached at a concentration of 60 to 80 percent methanol,

after which the pK rises sharply and may exceed that in

pure water. This behavior is in marked contrast to that

shown by uncharged weak acids such as acetic acid and the

phenols, represented by the curves at the top of the figure.

It has long been recognized that the contrasting behavior shown by the two types of curves of figure 3 can be

explained qualitatively on electrostatic grounds „ Thus

the large increase in pK produced by adding methanol to an

aqueous solution of an uncharged weak acid reflects the

increased difficulty in separating the oppositely charged

ions (hydrogen and the anion base) in the medium of lowered

dielectric constant produced by addition of the organic

constituent. The minimum in the pK curves for cationic

acids cannot be explained, however, on simple electrostatic

grounds alone. Recent efforts in this laboratory have been

directed toward an identification of the other parameters

which, together with the dielectric constant, can explain

in a fairly satisfactory way the existence of a minimum

and the course of the pK curve. Analyses of data for both

uncharged and positively charged acids have led to the

suggestion that the solvent effect on both types reflects

a "basicity effect", namely, an increase in the total basicity of the solvent molecules represented by SH in eq. 8.

26
J:

Q.

40 60

WEIGHT PERCENT METHANOL

Figure 3. pK values for acids of two charge types

in methanol-water solvents at 25 °C.

27
Reasonably concordant values for the magnitude of this

"basicity effect have "been obtained from the analysis of

data for the dissociation of weak acids of the two types.

The nature of this basicity effect and the identification

of the other parameters on which the solvent effect depends are still being investigated.

(R. G. Bates)

1. Ammonium Ion

Ammonium ion is the simplest and most important of

the positively charged acids. Furthermore, the approximately spherical structure of the ammonium ion adds a certain amount of reliability to the calculations of the

electrostatic work of charging the ion in a medium of known

dielectric constant. It therefore seemed desirable to

study the solvent effect of methanol on the dissociation of

ammonium ion.

Inasmuch as the standard electromotive force of the

hydrogen- silver chloride cell in methanol-water solvents

covering a wide range of solvent compositions is now known

with some accuracy, the conventional emf method was chosen.

Contributing to the success of this experimental method is

the fact that the solubility of silver chloride in ammonia

solutions decreases as the solvent is enriched with methanol.

The cell is represented as follows;

Pt;H

2 (g, 1 atm), NH,(m), NHi|_Cl(m) in MeOH-H2

0, AgCl; Ag

The measurements were made at 25 °C only, and the solvent

compositions were 10, 20, y^.h, 50, and 70 wt. percent

methanol. The results are given in table 5 as values of

28
Table 5. Acidic dissociation constant of ammonium ion

in methanol-water solvents at 25 °C.

wt. percent Me OH P( S

K)

9.245

10 9.146

20 9.044

33.4 8.893

50 8.687

60 8.591

70 8.571

Table 6. pa* for buffer solutions of ammonia and

ammonium chloride (each at molality m) in

methanol-water solvents at 25 °C

m

pa* in ~

20 percent

MeOH

50 percent

Me OH

70 percent

Me OH

0.02 9.126 8.798 8.711

0.04 9.158 8.838 8.762

0.06 9.180 8.866 8.797

0.08 9.198 8.888 8.823

0.10 9.212 8.905 8.843

29
p( K)j the subscript s signifies that the activity coefficients in the equilibrium expression are based on a standard state in the methanol-water medium, becoming unity

at infinite dilution in that medium. The pa.* values given

in table 6 were calculated as described in section 2.C.

These acidity functions may be regarded as reference data

useful in the study of the dissociation of other acid-base

systems in methanol-water solvents. In view of the difficulty of preparing the buffer solutions, however, these

values are not recommended as reference standards for pH*

measurements. This work has already been published [14],

(R. G. Bates)

2. Tris (hydroxymethyl ) methyl ammonium Ion

Tris(hydroxymethyl)aminomethane is a solid, crystalline base obtainable in a state of high purity. It is

moderately soluble both in water and in methanol and is

widely used as a standard reference material in acidimetry

as well as for pH control. Like ammonium ion, protonated

tris is presumably a roughly spherical ion whose electrostatic energy can be estimated by the Born calculation.

All of these considerations have led us to choose this base

for the study of the methanol solvent effect on acidic

dissociation.

The conventional emf method, which has been used extensively in this laboratory, was inapplicable to a determination of the pK values at temperatures other than 25

°C, for the standard emf of the hydrogen- silver chloride

cell is known only at 25 °C in several of the methanol-water

solvents included in the study, It was, however, desired

to include other temperatures in order to obtain information

30
on the enthalpies and entropies of dissociation. A titration method was therefore devised, utilizing a cell without

liquid junction consisting of a glass electrode and a silversilver chloride electrode. As in earlier work [15], concentrations were adjusted so that the titration was performed in a medium of essentially unchanging ionic strength,

constant chloride ion concentration, and, of course, fixed

solvent composition. In essence, data for the following

two cells were obtained:

1) Glass IjHClO^m), NaCl(m), AgCl; Ag

and

2) Glass II HClOj, (m), Tris, NaCl(m), AgCl; Ag

where the composition of the methanol-water solvent was

uniform throughout and the ionic strength was 2m.

The difference in the emf of these two cells may be

expressed in terms of the pK of protonated tris (TH ) as

follows:

PK - ^

+ log

^

+log ^H+^C1>2 (9)

2.3026RT/P m-nvp (VlM^Cl)!

The last term of this expression is a combination of activity coefficients which, under the conditions of the experiments, should be small. Experience showed that this

was indeed the case and that this term could be ignored

for the low concentrations of electrolytes used. The values

of pK found in 30, 70, and 90 wt . percent methanol at 25 °C

are compared in table 7 with the pK in water [16,17] and in

50 percent methanol [13]. The last column gives the solvent

effect, defined as p( K)-p( K)

.

(P. W. Schindler and R. G. Bates)

31
Table J, pK for the dissociation of protonated

tris(hydroxymethyl)aminomethane in

methanol-water solvents at 25 °C. The

solvent effect.

Wt. percent

Me OH P( SK) P( SK)~P(W

K)

8.072

30 7.940 -0.1^2

50 7.819 -0.253

70 7.855 -0.217

90 8.462 +0.390

3. Solubility Studies of the Medium Effect

The pK values determined as described in the preceding

section are in actuality p( K) values; K is the dissociation constant expressed in activities, a, defined by

a = m( y), where m is molality and the activity coefficient

s s y becomes unity at infinite dilution in the particular

smethanol-water solvent in question. Activities related to

the conventional aqueous standard state are designated a

w

when it is necessary to distinguish the two types. The

fundamental relationship between the two numerical scales

for the activity coefficient is

y = y. y (lO)

w' s

r m r v

'

which may be regarded as a definition of y, the "medium

effect activity coefficient" [4]. The magnitude of this

medium effect is determined solely by the Gibbs energy of

transfer of the species from one standard state to the other.

32
For the process

Species (in H? 0) = Species (in solvent s)

the Gibbs energy of transfer AG° is given "by

AG° = RT In y (ll)

A consequence of this relationship is that the solvent

effect on pK is a combination of the medium effects for the

ions and molecular species involved in the dissociation

process:

f

( K)-p(K) = log f-^lS-j

m\ TBH+

This expression can be written in another form which involves only mean activity coefficients:

p( sK)-p(

wK) = log myB

+ 2 [log mY± (HCl)-log

my± (BHCl)] (l?)

In order to elucidate as thoroughly as possible the

factors on which the solvent effect depends, we have undertaken to evaluate the separate medium effects occurring in

eq. l^. For this purpose, we have studied the Gibbs energy

and corresponding medium effect for the transfer of tris

from water to methanol-water solvents. This was done by

means of solubility measurements in the media at 15, 25,

and 35 °C. Inasmuch as saturated solutions of tris in

water and in a methanol-water solvent are both in equilibrium with solid crystalline base, we may write

%^1 = ms(w^2 = Vs^m^ < 14 )

33
where 1 represents the saturated aqueous solution and 2

represents the saturated solution in the methanol-water

solvent. From this relationship,

m

mTt

=

w

(v y) w ' ' 1

m

w m ( SY)

(15)

m

Thus the ratio of solubilities yields an approximate value

for the medium effect of tris base (T)j a more accurate

value could be obtained by correcting for the difference in

the "salt effect activity coefficients" (,/yh W _L and ( ay) S i— derived from independent measurements. The approximate

values of ym for tris are adequate for the present studies

m 'T

of the solvent effect. They are summarized in table 8.

(P. W. Schindler and R. G. Bates)

Table 8. Medium effects for tris (hydroxymethyl)aminomethane

in methanol-water solvents, obtained from

solubility measurements.

Wt. percent

Me OH

log

rn^T

(15 °C)

log& Ym m

rrji (25 °c)

log m

YT

(35 °C)

10 0.067 0.055 0.045

20 0.135 0.114 O.095

30 0.208 0.180 0.154

40 0.285 0.254 0.221

50 0.379 0.343 0.303

60 0.499 0.459 0.412

70 0.653 0.604 0.549

80 0.8^4 0.791 0.729

90 1.094 1.032 0.935

100 1.322 1.297 1.231

34
4. Parameters for the Solvent Effect

With a. knowledge of the medium effect for trls base,

it is possible to make some interesting comparisons between

the predicted solvent effect and that actually observed

„

The medium effect for the transfer of hydrochloric acid can

be derived readily from the standard electromotive force

for the hydrogen-silver chloride cell [18,19] containing

hydrochloric acid in methanol-water solvents. The medium

effect for tris hydrochloride has unfortunately not been

determined. Data for sodium chloride and lithium chloride

are, however, available from the work of Akerlof [20].

The values for these three electrolytes are compared with

those for undissociated tris and ammonia in figure 4. It

is evident that the medium effects for the two alkali

halides do not differ greatly and show a. continuous, nearly

linear, rise with the weight percentage of methanol. On

the other hand, the medium effect for hydrochloric acid

remains low at all water-rich solvent compositions and does

not rise rapidly until the solvent composition has reached

60 percent methanol.

Figure 5 shows that data such as these can account for

the minimum in the plot of pK for a cation acid as a function of the composition of the methanol-water solvent.

The broken lines of the figure were calculated by eq. 13

using the data for lithium chloride as an approximation to

the unknown medium effects for ammonium chloride and tris

hydrochloride. The pK values actually observed are shown

as solid lines. The minimum in the pK curve can thus be

attributed to the fact that the free energy of transfer of

alkali metal salts and ammonium salts increases much more

rapidly when methanol is added to their aqueous solutions

than does the free energy of hydrochloric acid at the same

35
o

El.O20 40 W

wt. % methanol

Figure 4. Medium effects ( y) for ammonia, tris,

hydrochloric acid, and alkali halides

in methanol-water solvents at 25 °C

36
20 40 60 80 100

wt. % methanol

Figure 5. Observed solvent effects on the pK of

protonated tris and ammonia (solid lines)

compared with the solvent effect

calculated by eq. 13 (broken lines).

37
solvent compositions. One can only speculate as to the

cause of this behavior. If normal behavior is represented

by the course of the curves for the alkali metal salts

shown in figure h, it must be concluded that at low methanol

concentrations the escaping tendency for hydrochloric acid

from the methanol-water media is unexpectedly low. This is

just another way of expressing the "basicity effect" earlier

postulated [21]. An increase in basicity as methanol is

added could indeed be attributed to a breakdown in the

structure of the water. It would appear that the proton

is solvated primarily by water molecules in solvents containing less than 60 wt . percent methanol and that the rapid

rise in methanol-rich solvents reflects the increased escaping tendency of protons solvated by the relatively weakly

basic methanol molecules.

(R. G. Bates)

38
B. Isotope Effects on the Dissociation of Acetic Acid

In Water and Deuterium Oxide

The dissociation of "ordinary" acetic acid, CH^COOH,

in "ordinary" water, H

p

0:

CH^COOH ?= H+

+ CH^COO" (l6)

3 3

is a process about which much quantitative information has

been available for many years. Thus, in 1932 Maclnnes and

Shedlovsky [22] measured the dissociation constant,

K =

"^H4" m

CH^COO- TH+ YCH^COOH / 17n m

CH^COOH ^CH-.COOH

3 3

at 25 °C by a conductance method. At the same time,

Harned and Ehlers [23], using an emf method, determined

the dissociation constant from to 60 °C . The concordance

between the two results at 25 °C is excellent, and for this

temperature we can write pK-. = 4.756 with confidence.

Acetic acid, however, can be dissolved in heavy water

(deuterium oxide, D? 0), and there is good reason to believe

that there is exchange between deuterium and the hydrogen

of the carboxyl group in acetic acid but no exchange with

the hydrogen atoms in the methyl radical. Thus "ordinary"

acetic acid in deuterium oxide forms the molecule CH^COOD,

3

and this in turn dissociates according to the process

CH COOD £=^D+

+ CH^COO" ( l8

)

39
with a dissociation constant

"%+ mCH^COO- YD

+ 7CH,C00-

K"2Q = 2 ^—

(19)

mCH^COOD YCH^COOD

The thermodynamics of this dissociation process has

"been studied recently in the Section and a report has been

published [11]. In addition to measurements of the pKp

value between 5 and 50 °C, pD (= -log a^) values were

evaluated for buffer solutions containing 0.05 ni CH-,C00D

and 0.05 m CH-^COONa in deuterium oxide, thus continuing the

work towards establishment of a pD scale in deuterium oxide

which had commenced with measurements of the pD values of

solutions of 0.025 m KD

2

P0^ and 0.025 m Na

2

DP0^ in deuterium oxide (see section 2.D).

Besides the "ordinary" form of acetic acid, CH^COOH,

a deuterated form is available. This acid, which has the

formula CD-^COOD, is sometimes called acetic acid-dh. When

dissolved in ordinary water, the following exchange would

be expected:

CD^COOD + Ho0£r==.CD^C00H + HDO (20)

with the production of acetic acid-d^ which in turn dissociates,

CD^C00H^=yH++ CD^COO" (2l)

Corresponding to this equilibrium, there is a dissociation

constant K^t

m

H

+ "^CD^COO

- TH

+ ^UD^COO"

K^ = 2 2 ( 22 ) m

CD,C00H YCD^C00H

40
The thermodynamics of this dissociation process has "been

studied during the current year [24], and details of the

results are given below.

In addition, there is a fourth dissociation process

which completes the series. If acetic acid-d^ is dissolved

in deuterium oxide, the dissociation process is

CD

3

C00D±zz^D+ + CD^COO" {2J>)

with

mD

+ m

CD^COO~ YD+ YCD,COCT

K, = -^ ^—

(24)

m

CD^COOD ^CD^COOD

This dissociation reaction has also "been the subject of

study during the current year. The results are now in the

process of publication.

Measurements were made of the emf of the cell

i Pt;H2

(g, 1 atm), CD

3C00H(m

1

), CD,C00Na(m

2

)., NaCl(nO, AgCljAg

The pK-, value of acetic acid-d^ in ordinary water is given

in terms of the emf (E) of this cell by the expression

j pK^ ; = PK

3

- log (Yd-THAc^^Ac-

= (E-E°)/k + log nu + log (m

1

-mH+ )/(m

2

+mH+ ) (25)

where Ac" = CD^COO" HAc = CD^COOD, and E°, the standard emf

of the cell in ordinary water as solvent, has already been

measured [25]. As usual, k represents (RT In 10)/F.

41
It would be anticipated that the term log

(Ypi-YwAc^/Vac- would he small, and this proved to be true.

Almost horizontal, straight-line plots of pK' against the

ionic strength gave extrapolated values of pK., at 10 temperatures from 5 to 50 °C.

In order to obtain the pKj, value of acetic acid-dr in

deuterium oxide, the cell

Pt;D2 (g, 1 atm), CD^COOD^), CD

3C00Na(m

2

), NaCl(nw), AgCljAg

was employed, using deuterium oxide as solvent. The standard

emf, E°, of this cell in deuterium oxide has already been

measured [9K The pK^ value of acetic acid-dh is given by

the expression:

PK4 a PK

4

- log (Ycl -YDAc )/YAc -

= (E-E°)/k + log rru + log (m

1

-mD+ )/(m

2

+mD+ ) (26)

where Ac" = CD^COO" and DAc = CD^COOD. The intercept pK^

is obtained by extrapolation of a plot of pK/, against

ionic strength.

A knowledge of the variation of the dissociation

constant of an acid with temperature leads to values of

the changes in enthalpy, AH°, entropy, AS°, and heat

capacity, AC , for the dissociation process. These values

at 25 °C for the four dissociation processes now under

consideration are collected in table 9- It will be observed

that the entropy changes for CH^COOH and CD^COOH in ordinary

water are almost the same.

This means that the exchange reaction:

CD^COOH + CHUC00- i=^ CH^COOH + CD^COCT (27)

42
proceeds with virtually no change in entropy. This is

also true for the exchange reaction

CD,COOD + CH^COO"

3 3

CH^COOD + CD^COO3 3 (28)

in deuterium oxide. However, this simple relationship does

not hold for the enthalpy changes. Here small, but perhaps

significant, differences are found.

Table 9« Thermodynamic quantities for the dissociation

of acetic acid and deuterioacetic acid in water

and deuterium oxide at 25 °C.

(1 cal = 4.1840 J)

pK AH°

cal mol"

AS°

cal mol deg

ac;

cal mol , -1deg

CH^COOH in Ho 3 2

4.756 -98 -22.1 -37

CD^COOH in Ho 3 2

4.772 -69 -22.1 -34

CH^COOD in Do 3 2

5.313 +275 -23.4 -39

CD^COOD in Do 3 2

5.325 +279 -23.4 -37

43
Calculations can also be made of the temperature at

which the dissociation constant will have its maximum

value and what this maximum value will be. In all four

instances, this temperature is found to lie close to room

temperature, being about 8 °C higher for deuterium oxide

than for ordinary water as solvent.

It is a common practice to list dissociation constants

on the molality scale of concentration, moles per kg of

solvent. For ordinary water, this is equivalent to moles

per 55»51 moles of solvent, but in heavy water it corresponds to moles per 49,94 moles of solvent. A more valid

comparison of dissociation constants is obtained by expressing them in terms of the mole fraction scale. When this

is done, it is seen that the pK values of CH_,COOH in

ordinary water and CH^COOD in deuterium oxide differ by

about 0.5 unit. This is about the difference found between

CTuCOOH in ordinary water and CTUC00D in deuterium oxide.

This relationship can be expressed in another way: the

pK values of CH,C00H and CD,C00H in ordinary water differ

by only about 0.01, and this is true also for the pK values

of CH^COOD and CD-,C00D in deuterium oxide. Thus, the

main isotope effect can be attributed to the change of

solvent. Deuteration of the methyl radical of acetic acid

itself produces a very small isotope effect.

(R. A. Robinson)

44
C. Acid-Base Behavior in Organic Solvents with Low

Dielectric Constants

A definitive monograph with the above title is being

prepared for publication by the Analytical Chemistry

Division; it is also to be published by Academic Press

(possibly in a condensed version) as a chapter in one of

a series of volumes on "The Chemistry of Non-Aqueous

Solvents", edited by Dr. J. J. Lagowski of the University

of Texas. The main topics to be covered (in the sequence

indicated) are as follows: I. Introduction. II. AcidBase Concepts. III. Classification of Organic Solvents.

IV. The Role of Hydrogen Bonding in Aprotic Media.

V. Quantitative Aspects of Acid-Base Behavior in Aprotic

Solvents.

The longest section will be Section IV, which will

deal with the following sub-topics: A. Self-association

of nitrogen- containing bases. B. Self-association of

acids (for example, phenols, carboxylic acids, nitric

acid, phosphorus-containing organic acids). C. Hydrogenbonded ion pairs. D. Conjugate Ions. E. Solute-solvent

interactions. F. Theoretical aspects of hydrogen bonding,

The discussion In the various sub-sections is based on

varied kinds of experimental evidence, principally from

measurements of conductance, dielectric polarization,

colligative properties, and absorption spectroscopy

(electronic, infrared, NMR).

As of May 1966, the literature research for the monograph was at least 95 percent complete. The writing of

Section IV was about two-thirds completed. There were

about 140 typed pages of manuscript, 18 figures, and

41 tables.

(M. M. Davis)

45
D. Spectrophotometric Studies of the Dissociation of

Phenols

The Section has devoted considerable attention to

acid-base effects in 50 wt. percent water-methanol as

solvent. The acids studied have been mainly of the A+B

charge type such as ammonium ion, the protonated form of

tris ( hydroxymethyl)aminomethane, and 4-aminopyridinium ion.

The pK values of acids of this type are lower in 50 percent

methanol than they are in water, and this difference is a measure of the "solvent effect," which is, in turn, a combination of the medium effects for the individual species

involved in the dissociation process (see section 3. A. J>)

.

It is known that the solvent effect is of the opposite

sign for acids of the charge type A°B~. This is true for

aliphatic acids such as acetic acid and for aromatic acids

like benzoic acid. Here the pK value is higher in 50 percent

methanol than it is in water.

Measurements have now been made with a number of acids

of the A°B~ type derived from phenol. A' spectrophotometric

method was used, together with buffer solutions whose pa*

values in 50 percent methanol have been determined in

earlier emf studies. In all cases, the pK value in 50

percent methanol was found to be greater than in water.

The differences observed between the pK value in 50 percent

methanol and in water at 25 °C are listed in table 10. The

solvent effect is seen to amount to about 0.6 in pK, but it

is clear that there is a considerable dependence on the

nature of the substituent groups in the phenol. The work

is now being prepared for publication.

(B. J. Steel and R. A. Robinson)

46
Table 10. Increase In pK for six substituted phenols

resulting from a change of solvent composition

from pure water (w) to 50 wt. percent methanol

at 25 °C.

Acid ApK = p( sK) P(wK )

o-Nitrophenol

m-Nitrophenol

p-Nitrophenol

4-Methyl-2-nitrophenol

2, 6-Dichlorophenol

Pentachlorophenol

0.687

0.615

0.536

0.687

O.889

0.458

47
4. EQUILIBRIUM STUDIES OF SYSTEMS OF ANALYTICAL INTEREST

A. Thermodynamics of the Dissociation of Morpholinlum

Ion in Water

The aqueous dissociation constants of a number of

nitrogenous bases over a temperature range have been measured previously in this laboratory by the emf method. Earlier

studies of the three ethanolamines [see reference 26] and

of the series of aminoalcohols derived from t-butylamine [27]

made possible comparisons of the basicities and other thermodynamic properties of -these structurally related compounds.

It was therefore of Interest to compare the heterocyclic

amine, morpholine, with the cyclic aliphatic amines piperidine [28] and pyrrolidine [29] studied earlier. The three

compounds have the following structures:

Piperidine Pyrrolidine Morpholine

Morpholine is of industrial Importance in organic chemical

reactions, as a solvent, and as a corrosion inhibitor In

steam boilers.

The dissociation constant was determined from to

50 °C by the measurement of the emf of hydrogen-silver,

silver chloride cells without liquid junction containing

buffer solutions composed of morpholine and Its hydrochloride. Details of the work are contained in a report

now in the process of publication.

Electromotive force measurements of cells of the type

PtjH2

(g, 1 atm), |(CH2

) 2 0(CH

2

) 2

NH

2Cl( mi ), |(CH

2

) 2 0(CH

2

) 2

NH (m2

).

AgCl; Ag

48
9.12

* 8.46 -

7.92

Figure 6. pK In morpholine-morpholinium chloride buffer

solutions as a function of ionic strength (m-, ) at 0, 25, and 50 °C

.

49
where m-, and m

? represent the molalities of the salt and

of the free amine, respectively, were made at 11 temperatures from to 50 °C. An expression for calculating the

"apparent" dissociation constant (K' ) from the emf data

was obtained by substitution of the acidity function

p( axjYri-| ) into the expression for the dissociation constant

of the protonated base. These values of the apparent dissociation constant were calculated from the measured emf

E (corrected to 1 atm of dry hydrogen) by the equation

pK^ = (E-E°)/k + log (m^/m

2

) - 2Am];/2 (29)

in which E° is the standard emf, k is (RT In 10)/F, and

A is the Debye-Hiickel constant on the molal scale. Straightline plots of pK* against m, , positioned by least squares,

gave limiting values of pK at 1, =0 with a standard deviaa. _l tion of 0.001 pK unit. Representative extrapolation lines

are shown in figure 6. The pK^ values were expressed as a function of temperature (T) in °K by an equation of the

Harned-Robinson form

1663.29

pK = + 4.1724 - 0.0042239T (30)

Values of the enthalpy, entropy, and heat capacity

changes for the dissociation process were calculated by

the application of the standard thermodynamic formulas to

eq. 30. The thermodynamic data for the dissociation of

morpholinium ion and the two related substances at 25 °C

are summarized in table 11. It may be noted that the

introduction of an oxygen atom into the pyrrolidine ring

decreases the basicity more than 2.8 pK units at 25 °C,

(H. B. Hetzer)

50
Table 11. Thermodynamic quantities for the dissociation

of morpholiniurn, piperidinium, and pyrrolidinium

ions in water at 25 °C

(1 cal = 4.1840 J)

Acid pK^ a AH° AS° AC°

P

cal mol cal deg

mol

cal deg

-,-1 mol

Morpholiniurn ion 8.492 9,330 -7.6 11.5

Piperidinium ion 11.123 12,760 -8.1 21.0

Pyrrolidinium ion 11.305 13,020 -8.1 16.2

B. Ionization Constant of Deuterium Oxide from 5 to 50 °C

In earlier sections of this report, a number of

thermodynamic measurements using deuterium oxide as a

solvent have been described. For acid-base studies in

heavy water, a knowledge of the ionization constant of

deuterium oxide is indispensable. In view of the considerable uncertainty in the literature values, this constant

has now been redetermined.

The cell chosen was that of Abel, Bratu, and Redlich

[30],

Pt;D

2 (g, 1 atm), NaOD, NaCl (in D

2 0), AgCl; Ag

and the emf was measured over the temperature range 5 to

50 °C. A constant ratio of the molalities of NaOD and NaCl

was maintained, but the total ionic strength was varied

over a range sufficient to permit an accurate extrapolation

to zero concentration of the ionized solutes.

51
A solution of sodium hydroxide (NaOD) in heavy water

was prepared by an adaptation of the method of Marsh and

Stokes [31], using the apparatus shown in figure J. The

separatory funnel of 1000 ml capacity contained 500 ml of

triply-distilled mercury and 400 ml of a 30 wt. percent

solution of sodium hydroxide in deuterium oxide. Electrolysis was carried out for 3 hr at a current of 4 A using a

2 platinum anode 2 cm in area. Stirring of the mercury was

necessary to prevent solidification of the amalgam at the

mercury-solution interface. At the conclusion of the

electrolysis the amalgam was allowed to run into the flask

containing 600 ml of deuterium oxide through which C0

p

- free, dry nitrogen had been bubbled for some hours , The

reaction between sodium amalgam and water is very slow at

room temperature. However, when the flask was maintained

at a temperature of 80 to 90 °C on a hot plate for eight

hr, sufficient reaction took place to yield a solution

about 0.5M in NaOD.

The ionization constant of deuterium oxide on the

molal (m) scale is defined by the equation

K

m

= m

D+'

m0D-^D+^0D- / aD

2

^l)

Values of pK were calculated from the equation: ^ m

pK

m

S pKm

" l0S ^C1-^0D- " lo%

aD

2 = (E-E°)/k + log mcl--/m

aD

- , (32)

where k is written for (RT In 10)/F. Values of E* have been

determined earlier [9]. The pK 1 varies little with ionic

strength (i), even with changes in m^., ../m^- ratio; this is

to be expected since both the activity coefficient term

YCl'TOD" an<^ aDo0 are c l° se to unity.

52
C0 2 -FREE

N 2 GAS

Figure J. Apparatus for the electrolytic production of

the sodium amalgam used to prepare the standard

solution of NaOD.

53
The values of pK' were fitted, to equations linear in

I by the method of least squares, and the intercepts (pK ) at zero ionic strength were determined in this way. The

values of pK on the molal scale (pK ), the molarity scale

(pK ), and the mole fraction scale (pK^) are summarized in

table 12 and compared with the corresponding quantities for

the self-dissociation of ordinary water [32]. Values of

the ratio K(HpO)/K(D

? 0) on the molal scale are also given.

The pK from 5 to 50 °C is given as a function of T, the

temperature in °K, by the equation

4913.14

pK = - 7.5H7 + 0.0200854T (33)

The difference between the experimental value and that

calculated by eq. 33 did not exceed 0.001 in pK at any

temperature. The thermodynamic quantities AH°_, AS , and AC° for the ionization of deuterium oxide have been

calculated with the use of eq. 33. They are compared with

similar data at 25 °C for ordinary water in table 13.

The results for H

p

have been recomputed to the basis

°C = 273.15 °K.

The value of pK = 14.869 at 25 °C is in good agreement with 14.86 found by Gold and Lowe [33]. The work of

Wynne-Jones [34] is not inconsistent with these two values

if one of his points is discarded. Other discordant results recorded in the literature are based on extrapolations to pure DpO of data for DpO-HpO mixtures and involve

considerable uncertainty.

(A. K. Covington)

54
Table 12. Comparison of the Ionization constants of deuterium

oxide and water from 10 to 50 °C.

D

2

H

2

K

m

(H

2 0)

t, °c pK pK pK,, ^ m c ^ N pK^ pK pK, T m ^ c ^ N

K

m

(D

20)

10

20

25

30

40

50

15.526 15.439 17.224

15.136 15.049 16.384

14.955 14.869 16.653

14.784 14.699 16.482

14.468 14.385 16.166

14.182 14.103 15.880

14.535 14.535 16.279

14.167 14.169 15.9H

13.997 14.000 15.741

13.833 13.837 15.577

13.535 13.542 15.279

13.262 13.272 15.006

9.79

9.31

9.08

8.93

8.57

8.32

55
Table 13. Comparison of the changes In enthalpy, entropy,

and heat capacity for the ionization of

deuterium oxide and water at 25 °C (molal scale).

t, °G D

2

H 2°

AH°, cal mol" 5

25

50

15,371

14,311

12,884

14, 427

13,526

12,312

-1 -1

AS°, cal mol deg 5

25

50

-16.76

-20.43

-25.03

-15.55

-18.68

-22.59

AC°, cal mol" deg" 25 -54.8 -46.6

t (°C)

max v ' 221 239

-log K ° max

12.356 11.382

C. Response of the Glass Electrode in Heavy Water

Although there are indications that the glass electrode

functions as well in heavy water as it does in ordinary water,

no confirmation of the response appears to have "been made by

direct comparison with the deuterium gas electrode. An

empirical correction of 0.40 unit in order to obtain pD from

the operational pH (based on aqueous standards) has, however,

been recommended [35,36]:

56
pD = pH + 0.40 (24)

This correction, determined with the use of solutions of

strong acid and strong alkali, converts the operationallydefined pH into values on a scale of pD.

The esta.blishment of a pa^ scale "based on the emf

of cells without liquid junction and the deuterium gas

electrode now makes it possible to confirm the deuterium

ion response of the glass electrode in buffer solutions

as well as in strong acid and strong alkali. Collateral

measurements now under way will also furnish a check of

the validity of the correction factor in eq. J>k . It is

hoped they will shed some light also on the origin of this

correction factor, that is, on whether the difference of

0.40 unit results in large part from a difference in the

standard potentials of the glass ||D and the glass IIH

couples or whether differences in the potentials at the

liquid junctions DpO KC1 (aq.) and H

p

| KC1 (aq.) are of

significance also.

The following solutions were prepared*

1. 0.05417 m DC1

2. Potassium tetroxalate, about 0.05m

5. Deuterioacetic acid, NaAc, and NaCl, each at a molality of 0.104^

4. KD

2 ?0^ and Na^DPO^, each at a molality of 0.02502

5. Borax and NaCl, each at a molality near 0.1

6. NaOD and NaCl, each at a molality near 0.01

The glass electrodes, four in number, were of four different

commercial types. Three were new, and the fourth had been

used for several months in aqueous solutions. The new

electrodes were conditioned in heavy water for five days

before use. The emf measurements were made on a commercial

pH electrometer of the chopper type. The voltage scale was

57
Figure 8. Equipment for a direct comparison of

the deuterium gas electrode and the

glass electrode in deuterium oxide*

58
calibrated against a. precise potentiometer and was considered reliable to 0.2 raV.

Measurements of the cell

Ptj D

2 (g), Soln in D2 j|glass

were made at 25 °C . The deuterium gas electrodes were allowed two hours to reach equilibrium. Each glass electrode

was placed successively in each of the six cells containing

the six test solutions, with intermediate rinsing by the

next solution before transfer. As a check, each electrode

was returned to the first solution at the end of the series.

The arrangement of cells and electrodes is shown in figure 8.

All of the glass electrodes followed the potential of

the deuterium gas electrode within 0.5 mV in the first four

solutions. In solutions 5 and 6 (pH 9.2 and 12.4, respectively), there was evidence that the alkaline error is larger

than would be expected for electrodes at the same pH in

ordinary water. Drifts of the electrodes on transfer from

one solution to the next were slight, but those electrodes

which had been conditioned in heavy water and used only in

this medium performed best.

(A. K. Covington and R. G. Bates)

59
5. THERMODYNAMIC STUDIES OF AQUEOUS SOLUTIONS

OF MIXED SALTS

The program for investigating the thermodynamic

properties of solutions containing two dissolved salts is

being carried out in cooperation with the Office of Saline

Water of the U. S. Department of the Interior.

The research utilizes the isopiestic method of measuring the aqueous vapor pressure of a solution. The immediate

result of an Isopiestic experiment is, therefore, a knowledge of the vapor pressure of the solutions, which can also

be expressed in terms of the water activity, a , of the

solution or the osmotic coefficient, (L . Further calculations lead to a knowledge of the activity coefficients and

chemical potentials of each of the salts in the solution.

Work on the system water- sodium chloride-barium

chloride was described in last year's report [1] and has

been published [37]. In the current year, work on the

three systems:

water-potassium chloride-barium chloride

water-calcium chloride-magnesium chloride

a water-sodium chloride-calcium chloride

has been completed. A paper on the first of these systems

has been published [38]* an(^ manuscripts describing work

on the other two are in press.

Data are also available for the system water-potassium

chloride-calcium chloride, and the analysis of these results is under way. A paper has also been published on

the thermodynamics of the system water-potassium chlorldeglycine [39]. Another paper, already published, describes

the extent to which the vapor pressure lowering of a solution of two salts can be calculated from the vapor pressure

lowerings of the separate, single salt solutions [40].

60
It has been found that the activity coefficients of

two salts B and C in a mixed solution of constant total

ionic strength can be represented by the equations

log yn

= log y£ - oc m (35) 'B ° r B B"T5

and

log yc = log y° - ^

c mB

(36)

Here y° is the activity coefficient of a salt in its own

solution in the absence of the other salt. For example,

in the system water-calcium chloride (B)-magnesium chloride

(C) at (mB+m

c

) = 4, -K^ = -O.O373 and oC

Q = O.O371. The

activity coefficient of calcium chloride in its own solution

at 4m is 2.93 and, hence, its activity coefficient in a solution containing 2m calcium chloride and 2m magnesium

chloride is 3*48. Thus the activity coefficient of calcium

chloride is considerably higher in the mixed salt than in

a solution of the same total molality consisting only of

calcium chloride and water.

Similarly, the activity coefficient of magnesium

chloride in its own solution at 4m is 5.53, and it is 4.67

in a solution containing 2m calcium chloride and 2m magnesium chloride. Thus the activity coefficient of magnesium

chloride is considerably lower in the mixed salt solution.

tr t r

Trace activity coefficients, yband yc

, can also

be calculated. These are the activity coefficients of one

salt when its concentration is reduced to trace amounts,

m-p. —> or rru —> 0. Thus the trace activity coefficient

of calcium chloride as rru —> 0, m~ —> 4 is 4.13, and that

of magnesium chloride as m-, —'? 4, m^ —> is 3*93*

61
These relations are illustrated in figure 9. It will "be

observed that the admixture of one salt with another tends

to equalize the activity coefficients, so that, in the

limit, the trace activity coefficients are almost, "but not.

completely, identical.

(R. A. Robinson)

0.8

0.7 -

0.6

0.5 -

0.4

iog/2

log 7* r

Figure 9„ Variation of activity coefficients (y) with

composition in an aqueous mixture of calcium

chloride (b) and magnesium chloride (C) of

total molality 4.

62
6. CONDUCTOMETRIC DETERMINATION OP TRACES OP WATER

A. Objectives and Principle of Method

Work has continued on the development of an Instrument

for the determination of water In amounts of 0.5 to 5 nig.

The water Is to be determined in a few hundred cc of a gas

stream, such as the combustion product obtained in the

microanalysis of organic compounds for hydrogen. The principle of the instrument is based on the large effect of water

upon the conductance of 0.2M solution of sulfuric acid in a propionic acid solvent. The specific conductance, K , of

this system at 25 °C is well represented by the empirical

equation:

105 X = (0.012705 + 9.785W)/(1.12860 - 130. 85W) (57)

where W is the weight fraction of water in the solution.

Thus, the addition of 0.1 percent of water nearly doubles

the conductance.

B. Modification of Apparatus

New conductance cells, circulation system, and

Wheatstone bridge were constructed to reduce the working

volume of solution, to eliminate several sources of leakage

and contamination, and to improve the electrical measurement s

.

Electrodes were made of platinum foil silver-soldered

to nickel plates for rigidity and were separated by an

annular gasket, 1/^2 in. thick, of poly-(tetrafluoroethylene),

designated PTPE. The electrolyte cavity of the cell was

38 mm in diameter and 0.8 mm wide. This provided about the

same cell constant as the previous cell but only one-fourth

the volume. External tubing was replaced by l/l6 in. PTPE

tubing with Luer fittings, and the 4-way stopcock of the

previous apparatus was replaced by a. PTPE plug valve

.

63
A serious problem has been the design of a suitable

circulating pump which would be chemically inert, have a

small volume, and be able to circulate the entire electrolyte through the conductance cell in about one minute or

less. The previous report [l] described a proposed glassenclosed magnetic plunger operated by external solenoids.

This model, with its associated check valves and tubing,

required a liquid volume of about 20 cc. The major disadvantage, however, was serious overheating of the solution

by the solenoids. The centrifugal design shown in figure 10

was adapted from Petriconi et al . [4l]. These authors did

not furnish dimensions, but the present model appears to be

close to the smallest practicable size.

The rotor is made from quarter-inch round PTFE rod

2.5 cm. long and contains embedded permanent magnets. This

length barely provides a useful pumping action in light of

the unavoidable clearances required by the glassblowing.

Under the conditions of use, the rotor can be run at about

1,000 rpm before it gets out of phase with the driving

magnet. At this rate it circulates 6 to 8 cc of electrolyte

per minute through the cell. It was found necessary to

install a spindle on the shaft in order to keep the rotor

centered in the housing. A particular advantage of the

design is the provision for gas scrubbing. The sample

stream of gas is introduced near the periphery of the housing and is thoroughly dispersed and mixed with the electrolyte before it leaves through the vertical axial tube. An

inner axial tube, 1 mm in inside diameter, provides both an

upper bearing for the spindle and a channel for the downward

flow of returning electrolyte while the spent gas bubbles

are rising through the outer channel.

64
Figure 10. Conductance cell and circulation system

for the determination of traces of water.

For test purposes, a "breadboard" bridge was wired

according to the schematic diagram shown in figure 11.

The major changes were the introduction of a Wagner earth

branch and selector switches to permit more versatility in

range, precision, and linearizing circuits. The balancing

section of the ratio arms was replaced by four 100-ohm steps

in series, across each of which the 10- turn precision potentiometer could be shunted successively by a. selector switch,

thus increasing the range of dial settings four-fold.

65
R

5K

r\ 8K <

Lioo-i-iooJLlOO^lOO^

t 1 1200

I5K L 2260

STEP

3 OK 3 2 10 4520

AAM/WWNMAAAAAA/WW

rh G

Figure 11. Schematic wiring diagram for conductance bridge.

66
It was originally planned that the fixed reference arm,

in parallel with the variable shunt C, would be a closely

similar conductance cell, as indicated in figure 11. This

cell, filled with the starting electrolyte solution, would

provide automatic compensation for the initial conductance

of the solution, polarization, and any temperature changes.

Because of various erratic effects and drifting of the

initial balance, described further below, this reference

cell was replaced by a fixed 5000-ohm resistor during most

of the tests. Provision was made for plug-in capacitors

in parallel with either ratio arm or with the Wagner earth,

as required for balance.

The previous report [1] described a bridge circuit

which approximately compensated for the non-linear conductance function (eq. 37). Further analysis has shown

that a. bridge with unequal arms can be designed which will

provide either an unbalance signal, or a dial reading at

null balance, which is directly proportional to the water

concentration of the electrolyte.

Equation 37 can be rearranged to the equivalent form

W = a(l - r)/(c + br) (38)

where a = 0.143^ b = 110.43, c = 16.624, and r is the

specific resistance of the solution relative to that of the

anhydrous solvent. For a simple Wheatstone bridge, it can

be shown that the ratio (e) of the output signal to the

input signal is given by

e = (SC-R)/[(1 + B)(C + R)] (39)

where S is the ratio of ratio arms, A/B; C is the equivalent

resistance of the reference arm, and R is the resistance of

the unknown arm. If it is required that e = kW, where k is

a proportionality constant, then a comparison of the equa67
tions reveals the conditions:

SC = 1 (40a)

(1 - S)S = kc/a (40b)

and

(1 +S) = kb/a (40c)

with the result that S = b/c, C = c/b, and k = bc/[a(b+c)],

in terms of the parameters of the conductance formula set

forth in eq. ;>8; S is 6.64 and k has about half the maximum

value of de/dW for an equal-arm bridge. This loss of

sensitivity is the price of a linear response.

A corresponding treatment in terms of the change in

resistance ratio required to bring the bridge to null

balance gave a corresponding value of k and, further, defined absolute values for the fixed resistors in A and B

such that the 400 ohms available for dial settings should

correspond to any desired concentration range of water.

Three such ranges were provided in the bridge.

C. Calibration with Water

Calibration of the system by means of weighed additions

of water from a micropipet have, in general, confirmed the

response of the electrolyte system and the circuitry as

derived from the conductance equation. The precision of

these measurements, however, was very disappointing. When

fresh electrolyte was put in the clean, dry instrument,

the initial balance condition fluctuated erratically, but

somewhat cyclically, for an hour or two, eventually settling down to a steady drift toward increasing conductance.

When the apparatus stood idle for several days at a time,

there was little change, but as soon as circulation of the

solution was resumed the same drift was observed; hence the

calibrations had to be made against the background of changing reference and conductance. It seemed improbable that

68
contamination was responsible for the effect, because leakage or corrosion would be expected to produce a localized

concentration of contaminant during static conditions.

Resumption of circulation would then produce a sharp peak

as the concentration gradient entered the cell, but such a peak was never observed. After five or six hours of circulation, the solution became noticeably cloudy, and eventually a fine precipitate settled in the pump chamber. Recently

it has been established that the precipitate is a nickel

compound. Thus, corrosion of the short nickel connecting

tubes between the cell and external PTFE tubing must be a factor in the poor reproducibility.

In an automatic analytical instrument it is desirable

that the output be displayed directly on a chart or dial.

This requires that, after an initial adjustment of the

bridge, the unbalance signal be related to the water content

of the electrolyte. In order to obtain a direct graphical

indication of performance, the output from the bridge was

fed through a. full-wave rectifier to a chart recorder with

a full-scale range of 1 mV. In order to calibrate the

chart divisions against the bridge dial, the successive

step positions were first recorded. The bridge was then

rebalanced on the first position and three additions of

water were made. The resulting trace is shown in figure 12.

It is immediately apparent that the successive steps did

not produce equal pen displacements. This result is still

unexplained, but it may be related to the fact that a true

null balance could not be obtained on the chart, perhaps

because of a lack of shielding on the bridge or a faulty

rectifier. For this measurement, the bridge and recorder

were connected to the same ground.

69
0.1 0.2 0.3 0.4 0.5 0.6 0.7 0.8 0.9 1.0 mv.

Figure 12. Recorder trace of calibrations by resistance

steps and by addition of water,

70
Although It is less obvious, the chart record for

water additions Is also nonlinear. In curve I of figure 1J>,

"equivalent" step positions, as derived by an admittedly

crude Interpolation of the chart calibration, are plotted

against the water additions. The system contained 13.224 g

of electrolyte, so the final point represents nearly 0.1

percent water. The chart shows a very rapid response after

the addition, of water, but there seems to be a small cyclic

change with a period of 5 min as complete mixing takes

place. Despite the rather noisy record, it is evident that

0.2 mg of water is detectable in the more sensitive region

of the chart

.

For a more careful calibration, the output of the

bridge was connected to a tuned amplifier-null detector

combination, and the bridge was rebalanced after each addition of water. The result of one such calibration is

shown in curve II of figure 13 . The difference in slopes

of curves I and II Is due to different' positions of the

range switch. The expected slope of curve II is 0.0573

while the observed value is 0.05^5.

D. Analysis of Errors

Consideration has been given to the ultimate precision

which might be obtainable. For observation of the bridge

unbalance signal, the relation between W and signal, as

-2

outlined in section 6.B, is dW/de = 1 x 10 . It is doubtful that a digital a.c. voltmeter costing under $3,000

could measure the signal with an error much less than 0.1

percent. Thus, the probable error in W would be 1 x 10 , corresponding to 0.1 mg of water in 10 g of electrolyte.

As related to the microanalysis of organic compounds, this

uncertainty corresponds to about 0.01 mg of hydrogen, or

1 percent absolute on a sample weighing 1 mg. The method

71
6 8

WATER, mg

10

Figure 13. Comparison of calibration methods.

Curve I, derived from figure 12

as a function of time

Curve II, obtained by the

null-balance method

72
would thus be suitable only for routine or uncritical

applications. The null-detector balance method of operation places less stringent requirements on the detector

but is limited by the accuracy of the potentiometer (0.1

percent). The corresponding analysis indicates that the

probable error in W would be an order of magnitude less,

permitting the determination of hydrogen in a 1 mg sample

with an absolute error of 0.1 percent. To extend the

accuracy further would require precision conductance techniques and several decades of precision resistors in the

bridge. This entire analysis is based on the application

of eq. 37 for the conductance of the electrolyte system.

It is unlikely that this empirical relationship is accurate

to better than 0.5 percent over the working range, and therefore a careful calibration will be required.

E. Status of the Project

Steps which are called for immediately are the replacement of nickel in the cell and the provision of better

shielding and placement of the bridge components. There is

room for further improvement in the absorption- circulation

system in the direction of reducing the volume and the time

required to reach a steady value of the conductance.

(T. B. Hoover)

73
7. KINETIC TECHNIQUES IN TRACE ANALYSIS

The principal investigator was on assignment to the

Coast and Geodetic Survey as a Science and Technology

Fellow during most of the year. For this reason, there is

very little progress on this project to report

During the early months of the reporting year, however,

a summary describing the use of reaction rate methods in

analytical chemistry was written. This summary was prepared

for internal use as an aid to the Division and Section in

formulating future plans for work in this area. Reaction

rate methods were subdivided Into three general classes,

namely catalytic methods, enzyme-catalyzed reactions, and

differential rate methods.

Catalytic methods take advantage of the susceptibility

of certain reactions to catalysis by substances present in

trace amounts, and are, therefore, directly applicable to

problems in trace analysis. Nearly 90 different reactions

which have been successfully employed for this purpose were

identified In the literature. The summary contains a table

which lists 32 elements that have been determined by the

catalytic method, summarizes the concentration range or

sensitivity, and notes any serious interferences that have

been studied. The transition metals constitute most of

the group of 32 elements.

The current status of enzyme-catalyzed reaction methods

was discussed, although the literature search was not as

exhaustive as It was for the catalytic methods.

74
Quantitative methods "based on the use of differential

reaction rates depend on the fact that two or more related

compounds in a mixture may undergo the same type of reaction

at different rates- These methods are generally applicable

to binary and often to ternary mixtures of organic compounds

such as homologs or isomers containing the same functional

group. The summary lists all of the binary and ternary

mixtures for which special procedures have been established

and reported in the literature.

(R. K. Wolford)

75
8. STANDARD REFERENCE MATERIALS FOR ACIDIMETRY

A. Homogeneity Tests of Potassium Hydrogen Phthalate 84h

The purpose of this work was to determine if the

commercial "candidate" lot of potassium hydrogen phthalate

was of sufficient uniformity and purity to warrant analysis

and certification as an NBS acidlmetric standard.

The material was analyzed "by weight titration with

aqueous carbonate-free alkali standardized against NBS

potassium hydrogen phthalate 84d. The value: of 99.98o

percent purity, obtained by Bates and Wichers [92] by

differential titration with the hydrogen electrode, was

used for sample 84d. The procedure was essentially that

suggested on the NBS certificate for the use of this

material as an acidimetric standard, except that larger

samples (approximately 2 g) were titrated. The base was

added from an enlarged weight buret delivering approximately ] 120 g of solution.

The end point at pH 8.7, the inflection point obtained

experimentally by Bates and Wichers, was detected by means

of a glass electrode. The change in pH per drop of base

was recorded for several drops below and the one immediately

above pH 8.7. The buret was then weighed. From the known

weight of one drop of base, the weight which would have

changed the pH to 8.7 was interpolated. Results of analyses

of samples taken from the top, middle, and bottom of each

of the seven containers are given in table 14. The assay

values are given in percent.

With the three sets of duplicates averaged, the

standard deviation for the variability of the results is

0.0073. The conclusion from the analysis of variance was

that, since the differences from container to container

76
and from top to bottom in a particular container were not

sufficiently larger than the residual error, one may regard the entire lot as being homogeneous.

Thanks are due to Mr. J. M. Cameron for the statistical analysis of these data and to Dr D W. J. Youden for

assistance in the design of the experiments.

(H. B. Hetzer)

Table 14. Assay of potassium hydrogen phthalate 84h

for homogeneity.

Location

Container

I II III IV V VI VII

Top

Middle

Bottom

99-972

99.982

99.985

99.985

99.984

99.975

99.972

99.989

99.980

99.973

99.986

99.970

99.986

99.991

99.971

99.990

99.990

99.992

99.973

99.975

99.970

99.978

99.973

99.990

77
B. Certification of Potassium Hydrogen Phthalate 84h

as an Acidimetric Standard

The constant-current coulometric method was chosen

for the determination of the assay of a new lot of potassium hydrogen phthalate, to be issued as a Standard

Reference Material for acidimetry. This method has the

advantage over other conventional methods of analysis that

it involves the titration of the sample with a measured

quantity of electricity. The results are therefore referred to fundamental physical constants such as the volt,

ohm, and second, all of which can be easily and accurately

measured.

By combining Faraday's law of electrolysis with Ohm's

law, one can calculate the amount of electrochemical reaction during electrolysis by means of the equation

iAt

(no. of equivalents) = (4l)

where i is the current in amperes, At' is the elapsed time

in seconds, and F is the faraday. The latest value for

the faraday is 96,487.0 C equiv" . Thus in a hydrogen

coulometer the quantity of hydroxyl ions produced as the

titrating reagent can be determined with a high degree of

accuracy and precision from the known magnitude of the

current passed and the time elapsed.

The technique for precise coulometric titration of

acids and bases has been worked out previously by Taylor

and Smith [43]. The facilities used in this work were part

of the equipment of the Micro chemical Analysis Section and

have been described and illustrated in the annual report

of that section for the year 1965 [44]. Acknowledgment is

made to George Marinenko for instruction in the use of

the equipment

,

78
In general, the experimental procedures developed by

Taylor and Smith were followed. From the seven containers

of potassium hydrogen phthalate, the homogeneity of which

had been studied by H. B. Hetzer (see section 8. A), six

samples were taken at random for coulometric study. Twogram samples were ground gently, dried at 120 °C for two hr,

and titrated by passing a known amount of current through

a cell consisting of a platinum cathode and a silver anode,

with a 1 M solution of potassium chloride as the supporting

electrolyte. The titration was performed under a nitrogen

atmosphere, and the end point of the titration was determined by noting the pH changes, indicated by a glass-calomel

combination electrode, as the last 0.05 percent of the total

current was passed in small increments through the solution.

A typical end point determination is shown in figure 14.

Curve a represents the pH indicated by the glass electrode

as a function of time. Curve b is the differential plot

from which the inflection point of the titration curve was

determined. An error of 0.005 percent in the total assay

is represented by 25 seconds on the plot.

The results obtained for 17 titrations of the new lot

are given in table 15. The values in curly brackets indicate duplicate titrations of the same grind. The data were

analyzed by Mr. J. M. Cameron, and the purity of the sample

was found to be 99.9921 percent, with a standard deviation

of O.OO65 for a. single determination. The analysis of

variance supports the statement, with approximately 95

percent confidence, that 99 percent of all samples have

assay values lying in the interval 99.9921 ± 0.0288.

(M. Paabo)

79
3.8

1 r i rH

0.005%

50 100 150 200 250 300 350 400

sec

Figure 14. End point determination in a typical

coulometric titration.

Curve a, pH change as a function of time

Curve b, differential plot used to

locate the end point

80
! Table 15. Assay of potassium hydrogen phthalate 84h for

certification as a. standard for acldlmetry.

(values In percent)

Location

Container

I II rv V VI

Top /99. 99751

(99.9952)

100.0024

(99.98221

1 99. 9869J

100.0018

Middle 199.9907)

[99.9967]

(99.9863)

199.9856]

99.9964

Bottom i99. 99261

l99.99!50j

(99.9856)

!99.9825j

(99.9959)

(99.99443

81
C . Status of Certification Procedures for Standard

Reference Materials

Precise acidimetry has always presented an unrelenting

challenge. Coulometric titrations and differential potentioraetric titrations with the hydrogen electrode have shown

themselves capable of sensitivities of a few thousandths of

one percent in the assay of standard materials for acidimetry.

As conducted here, these methods are providing accuracy which

we believe is not exceeded elsewhere. The statistical examination of the data, for potassium hydrogen phthalate 84h

suggests strongly that the accuracy of the analysis surpasses the degree of homogeneity now attainable on a large lot

of material. The sensitivities of these refined methods

are likewise so great that distinctions between the inflection point of the titration curve and the true equivalence

point may have to be made. These differences are not clearly understood, and studies of end points may have to be

undertaken.

Somewhat different problems are encountered in the

certification of reference standards for pH measurements.

The experimental emf measurements which lead unambiguously

to values of the acidity function p(aHYr

,-,) have been developed to a high level of precision in our laboratory.

An accuracy of 0.002 pH unit is attainable in water and in

the other media we have studied, over most of the experimental temperature range. However, the assigned pH(s) is

derived from this well-defined acidity function by introduction of an arbitrary convention whose "accuracy" cannot

be assessed. Furthermore, the instruments used almost

universally in the measurement of pH utilize a cell with

liquid junction. The vagaries of the liquid- junction

potential are such that the experimental pH values can

82
rarely be considered reproducible to better than 0.01 unit

regardless of the refinements in the standard values for

the reference materials with which the assembly is standardized. There remains always the challenge to reduce the

variance in the assigned pH(s)j an even greater challenge

is presented by the prospect of designing an improved

instrument for practical pH measurements.

(R. G. Bates)

83
9. FOREIGN-CURRENCY PROGRAMS

Two projects being conducted at universities in India

and financed through the provisions of Public Law 480 are

sponsored by the Electrochemical Analysis Section. A brief

description of these programs follows.

A project entitled "Properties of Electrolytes in

Nonaqueous Media" is being conducted at Jadavpur University

in Calcutta under the supervision of Dr. M. N. Das and

Dr. K. K. Kundu. At the present time Dr. Das and Dr. Kundu

are studying the properties of weak electrolytes in glycolwater solvents over a. range of temperatures. Work has been

done on the standard emf of the hydrogen-silver chloride

cell in glycol-water media, and similar studies of the

hydrogen -silver bromide cell have been initiated. Data.

for the pK values of monoethanolammonium ion and diethanola

ammonium ion have also been obtained in these solvents.

A project entitled "Physicochemical Studies in Nonaqueous Solvents" is being pursued at Panjab University,

Chandigarh, under the supervision of Professor Ram Chand

Paul. Professor Paul's current interests include the

behavior of Lewis acids in molten acetamide solvents.

Conductometric methods have been used, and several solvates

have been isolated and identified. The work is being

extended to studies of the solvation of electrolytes in

general, including protonic acids, in amide solvents.

Dissociation constants of weak electrolytes will be determined, and an investigation of ion-pair formation in

these media is planned.

(R. G. Bates)

84
10. PERSONNEL

Roger G. Bates, Section Chief

R. A. Robinson, Assistant Section Chief

Marguerite Raudenbush, Section Secretary

Electromotive Force and

Optical Measurements of Acidity

V. E. Bower

A. K. Covington, On leave from the University

of Newcastle upon Tyne, England.

R. Gary

Hannah B. Hetzer

Maya Paabo

R. A. Robinson

Barry J. Steel, On leave from the University of

Adelaide, South Australia.

Conductance Measurements

T. B. Hoover

Dielectric Measurements

C. G. Malmberg

Acid-Base Studies in Inert Solvents

Marion M. Davis

Isopiestic Vapor Pressure Measurements

V. E. Bower

R. A. Robinson

Kinetic Studies

R. K. Wolford

Solubility Studies

P. W. Schindler, Guest worker of the Swiss

National Foundation, on leave

from the University of Bern.

85
11. COMMITTEE MEMBERSHIPS

Roger G. Bates

Member, Subcommittee on Hydrogen Ion Determinations,

ASTM Committee E-l, Sub. 22.

Councilor, American Chemical Society, Washington Section.

Member, NRC Committee on Symbols, Units, and Terminology

for the Physical Sciences (U.S. National Committee for

ISO/TC-12).

Member, U.S. National Committee for IUPAC

.

Member, Commission on Electrochemistry, IUPAC.

Member, Commission on Symbols, Terminology, and Units,

IUPAC

.

Associate Member, Commission on Electroanalytical

Chemistry, IUPAC.

Member, Subcommittee on Physiocochemical Symbols and

Terminology, Committee on Physical Chemistry, NRC.

Member, Advisory Board, Analytical Chemistry . Member, Board of Editorial Advisors, Analytica Chimica

Acta . Member, Board of U.S. Civil Service Examiners

(Chemistry Panel).

Member, Washington Editorial Review Board (NBS).

Member, Committee on Units and Usage (NBS).

Member, Museum Committee (NBS).

Marion M. Davis

Member, American Association of University Women

Educational Foundation, Committee on Fellowships

to American Women.

Member, Committee on Long-range Planning,

Chemical Society of Washington.

Member, Board of U.S. Civil Service Examiners

(Chemistry Panel).

36
Robert Gary

Treasurer, Society of The Sigma Xi, D.C. Chapter.

R. A. Robinson

Secretary, Commission on Electroanalytical Chemistry,

IUPAC

.

8T
12. PUBLICATIONS AND MANUSCRIPTS, JULY I965 TO JUNE I966

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

Standardization of Analytical Data Obtained with SilverSilver Chloride Electrodes in Methanol-Water Solvents.

Anal. Chem. 37, 462-464 (1965).

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

Thermodynamics of the Ternary System: Water-GlycinePotassium Chloride at 25 °C from Vapor Pressure

Measurements

.

J. Res. Natl. Bur. Std. 69A, 131-135 (1965).

3- Robinson, R.A., Bower, V.E.

An Additivity Rule for the Vapor Pressure Lowering of

Aqueous Solutions.

J. Res. Natl. Bur. Std. 69A, 365-367 (1965).

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

Osmotic and Activity Coefficients of

Tris(hydroxymethyl)aminomethane and its Hydrochloride

in Aqueous Solutions at 25 °C

J. Chem. and Eng. Data 10, 246-247 (1965).

5. 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. 69, 2750-53 (1965).

6. 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. 37, 1291 (1965).

7. 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. Res. Natl. Bur. Std. 69A, 439-448 (1965).

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

Electrochemical Analysis: Studies of Acids, Bases,

and Salts by EMP, Conductance, Optical, and Kinetic

Methods.

NBS Technical Note 271 (September 6, 1965).

9. Davies, J.M., McQue, B., Hoover, T.B.

Heat Transferred by Decomposition Products from

Cotton Fabrics Exposed to Intense Thermal Radiation.

Textile Res. J. 35, 757-769 (1965).

88
10. Bates, R.G.

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

Anal. Chem. 37, 57A (1965).

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

Dissociation of Ammonium Ion in Methanol-Water Solvents.

J. Phys. Chem. 70, 247-251 (1966).

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

Dissociation of Acetic Acid-d^ in Aqueous Solution

and Related Isotope Effects from to 50°.

J. Phys. Chem. 70, 540-543 (1966).

13. Davis, M.M., Hetzer, H.B.

Titrimetric and Equilibrium Studies Using Indicators

Related to Nile Blue A.

Anal. Chem. 38, 451-461 (1966).

14. Bates, R.G., Robinson, R.A.

Acid-Base Behavior in Methanol-Water Solvents.

Chemical Physics of Ionic Solutions, Chapter 12,

pp. 211-235, John Wiley and Sons, Inc., New York, N.Y.

1966.

15. Bates, R.G.

Electrometric Methods of pH Determination.

Chapter 26 in "Standard Methods of Chemical Analysis",

6th ed., IIIA, pp. 521-532, P. J. Welcher, Ed.

D. Van Nostrand Co., Inc., Princeton, N.J., 1966.

16. Bates, R.G.

Acids, Bases, and Buffers, in "Current Concepts of AcidBase Measurements", Ann. N.Y. Acad. Sci. 133, 25-33 (1966)

17. Davis, M.M., Paabo, M.

Comparative Strengths of Aliphatic Acids and Some

Other Carboxylic Acids in Benzene at 25°.

J. Organic Chemistry 31, 1804-1810 (1966).

18. Bates, R.G.

pK Determination.

Encyclopedia of Industrial Chemical Analysis, Vol. III.

Interscience Publishers, New York. (in Press).

19. Bates, R.G.

Acidity Functions for Amphiprotic Media

Chapter for Chemistry of Non-Aqueous Solutions.

Academic Press, New York. (in Press).

20. Bates, R.G.

Proposal for the Practical Measurement of pH in

Amphiprotic and Mixed Solvents.

Pure and Appl. Chem. (in Press),

89
21. Stokes, R.H., Robinson, R.A.

Interactions In Aqueous Nonelectrolyte Solutions.

I. Solute-Solvent Equilibria.

J. Phys. Chem. (in Press).

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

Dissociation Constant of Morpholinium Ion and

Related Thermodynamic Quantities from to 50°.

J. Phys. Chem. (in Press).

23>. Robinson, R.A., Bower, V.E.

Properties of Aqueous Mixtures of Pure Salts.

Thermodynamics of the Ternary System: Water-Calcium

Chloride-Magnesium Chloride at 25 °C

.

J. Res. Natl. Bur. Std. (in Press).

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

Dissociation of Acetic Acid-d4 in Deuterium Oxide

from 5 to 50° and Related Isotope Effects.

J. Phys. Chem. (in Press).

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

Properties of Aqueous Mixtures of Pure Salts:

Thermodynamics of the Ternary System Water-Sodium

Chloride-Calcium Chloride at 25 °C

.

J. Res. Natl. Bur. Std. (in Press).

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

Behavior of the Glass Electrode and Other

pH-Responsive Electrodes in Biological Media.

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

27. Covington, A.K., Robinson, R.A., Bates, R.G.

The Ionization Constant of Deuterium Oxide

from 5 to 50°.

J. Phys. Chem. (in Press).

28. Schindler, P.W.

Heterogeneous Equilibria Involving Oxides, Hydroxides,.

Carbonates and Hydroxide Carbonates.

Advances in Chemistry. (in Press).

29. Robinson, R.A., Bates, R.G.

The Dissociation of Some Substituted Phenols in

50-Percent Aqueous Methanol as Solvent.

J. Res. Natl. Bur. Std. (in Press).

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

Dissociation of 4-Aminopyridinium Ion in 50 per cent

Methanol-Water and Related Acidity Functions from

10 to 40°C.

Anal. Chem. (in Press).

90
13. TALKS, JULY 1965 TO JUNE 1966

1. Bates, R.G., "Nonaqueous pH Measurements", Analytical

Chemistry Group, American Chemical Society,

Pittsburgh Section, Mellon Institute,

Pittsburgh, Pennsylvania. October 8, 1965.

2. Bates, R.G., "Solvent Effects on Acidity in Nonaqueous

Media", The Electrochemical Society, Buffalo, N.Y.

October 11, 1965.

3. Bates, R.G., "Solvent Role in Acidity", American Chemical

Society, Susquehanna Valley Section, King's College,

Wilkes-Barre, Pennsylvania. October 20, 1965.

4. Bates, R.G., "Acid-Base Processes in Alcohol-Water Solvents",

Chemistry Seminar, Michigan State University,

East Lansing, Michigan. November 16, 1965.

5. Bates, R.G., "Acid-Base Behavior in Water-Like Solvents",

Wabash Valley Section, American Chemical Society,

Terre Haute, Indiana. January 11, 1966

.

6. Bates, R.G., "pH Numbers -- their Measurement and Meaning",

Combined Student Group, Rose Polytechnic Institute and

Indiana State University, Terre Haute, Indiana.

January 12, 1966.

7. Bates, R.G., "Acid-Base Behavior in Alcohol-Water Solvents",

Chemistry Seminar, Indiana University, Bloomington,

Indiana. January 12, 1966.

8. Bates, R.G., "Dissociation of Ammonium Ion in Methanol-Water

Solvents", American Chemical Society, National Meeting,

Phoenix, Arizona. January 20, 1966.

9. Covington, A.K., "Dissociation of Moderately Strong Acids",

Department of Chemistry, Carnegie Institute of Technology,

Pittsburgh, Pennsylvania. February 21, 1966.

-0. Covington, A.K., "Dissociation of Moderately Strong Acids",

Reactor Chemistry and Chemistry Division, Oak Ridge

National Laboratory, Oak Ridge. Tennessee.

March J, 1966.

LI. Covington, A.K., "Dissociation of Moderately Strong Acids",

Brookhaven National Laboratory, Upton, Long Island,

New York. March 14, 1966.

91
12. Covington, A.K., "The Hydrogen-ion Response of the

Glass Electrode", Department of Chemistry, Brown

University, Providence, Rhode Island. March 16, 1966.

13. Covington, A.K., "Dissociation of Moderately Strong Acids",

Department of Chemistry, Rensselaer Polytechnic Institute,

Troy, New York. March 17, 1966

.

14. Bates, R.C., "Solvent Effects on Acid-Base Behavior in

Amphiprotic Media", Chemistry Colloquium, Pennsylvania.

State University, University Park, Pennsylvania.

March 17, 1966

.

15. Covington, A.K., "The Hydrogen-ion Response of the Glass

Electrode", Exploratory Chemistry Research Department,

Corning Glass Works, Corning, New York. March 18, 1966.

16. Schindler, P.W., "Heterogenous Equilibria Involving Oxides,

Hydroxides, Carbonates and Hydroxide Carbonates",

American Chemical Society, National Meeting,

Pittsburgh, Pennsylvania. March 24, 1966.

17. Bates, R.G., "Potentiometric Measurements and pH Standards",

Course in Advanced Laboratory Instrument at ion,

Walter Reed Army Institute of Research, Washington, D.C.

March 25, 1966.

18. Covington, A.K., "The Hydrogen-ion Response of the Glass

Electrode", The Electrochemical Society, National

Capital Section, Reiss Science Center, Georgetown

University, Washington, D.C. April 7, 1966

.

19. Covington, A.K., "Dissociation of Moderately Strong Acids",

Clarkson College - St. Lawrence University Sigma Xi Club,

Clarkson College, Potsdam, New York. April 11, 1966.

20. Bates, R.G., "Role of the Solvent in Acid-Base Behavior",

Chemistry Seminar, Bucknell University,

Lewisburg, Pennsylvania. April 11, 1966.

21. Bates, R.G., "pH: Its Meaning and Measurement",

Physical Chemistry Class, Bucknell University,

Lewisburg, Pennsylvania. April 12, 1966.

22. Covington, A.K., "Dissociation of Moderately Strong Acids",

Department of Chemistry, McGill University,

Montreal, Canada. April 13, 1966,

92
Bates, R.G., "Solvent Effects on Acid-Base Behavior",

Chemistry Seminar, University of the State of New York

at Buffalo, Buffalo, New York. April 20, 1966.

Bates, R.G., "Acid-Base Chemistry of Partially Aqueous

Media", Chemistry Seminar, George Washington University,

Washington, D.C. April 29, 1966.

Covington, A.K., "Dissociation of Moderately Strong Acids",

Rockefeller University, New York, New York. May 5, 1966

.

Bates, R.G., "Behavior of the Glass Electrode and

Other pH-Responsive Electrodes in Biological Media",

Conference on Bioelectrodes, New York Academy of

Sciences, Woldorf -Astoria Hotel, New York, New York.

June 6, 1966.

93
14. LIST OF REFERENCES

[1] Electrochemical Analysis: Studies of Acids, Bases, and

Salts by Emf, Conductance, Optical, and Kinetic Methods . July 1964 to June 1965 . (R.G. Bates, Ed. ),

NBS Technical Note 271 (September 6, 1965).

[2] Bates, R.G.

Determination of pH , Chap. 3 and 4.

John Wiley and Sons, Inc., New York, 1964.

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

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

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

J. Phys. Chem., 67, 1833 (196^).

[5] Hamer, W.J., Pinching, G.D., and Acree, S.F.

J. Res. Natl. Bur. Std., 36, 47 (1946).

[6] Hamer, W.J. and Acree, S.F.

J. Res. Natl. Bur. Std., ^2, 215 (1944).

[7] Bates, R.G., and Hetzer, H.B.

J. Res. Natl. Bur. Std., 64A, 427 (i960).

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

J. Chem. Eng. Data, 9, Jjk (1964).

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

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

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

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

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

J. Phys. Chem., 69, 2750 (1965).

[12] Davis, M.M. and Hetzer, H.B.

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

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

J. Res. Natl. Bur. Std., 69A , 26^ (1965).

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

J. Phys. Chem., 70, 247 (1966).

94
[15] Bates, R.G., Siegel, G.L., and Acree, S.F.

J. Res. Natl. Bur. Std., 30, 347 (1943).

[16] Bates , R.G. and Hetzer, H.B.

J. Phys. Chem., 65, 667 (I96l).

[17] Datta, S.P., Grzybowski, A.K., and Wilson, B.A.

J. Chem. Soc. (London), 792 (1963).

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

Anal. Chem., 37, 462 (1965).

[19] Olwa, I.T.

Scl. Rpts. Tohoku Univ., Ser. 1, 41, 47 (1957).

[20] Akerlof, G.

J. Am. Chem. Soc, 52, 2353 (1930).

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

Chap. 12 in "Chemical Physics of Ionic Solutions"

(B. E. Conway and R. G. Barradas, Eds.), pp. 211-235

John Wiley and Sons, Inc., New York, 1966.

[22] Shedlovsky, T. and Maclnnes, D.A.

J. Am. Chem. Soc, 54, 1429 (1932).

[23] Harned, H.S. and Ehlers, R.W.

J. Am. Chem. Soc, 54, 1350 (1932); 55, 652 (1933).

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

J. Phys. Chem., 70, 540 (1966).

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

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

[26] Bates, R.G. and Allen, G.F.

J. Res. Natl. Bur. Std., 64A , 343 (i960).

[27] Hetzer, H.B., Robinson, R.A., and Bates, R.G.

J. Phys. Chem., 66, 2696 (1962).

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

J. Res. Natl. Bur. Std., 57, 153 (1956).

[29] Hetzer, H.B., Bates, R.G., and Robinson, R.A.

J. Phys. Chem., 67, 1124 (1963).

95
30] Abel, E., Bratu, E., and Redlich, 0.

Z. Physik. Chem., A173, 353 (1935).

31] Marsh, K.N. and Stokes, R.H.

Aust. J. Chem., 17, 740 (1964).

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

Trans. Faraday Soc, 36, 973 (1940).

33] Gold, V. and Lowe, B.M.

Proc. Chem. Soc. (London), 140 (1963).

34] Wynne-Jones, W.F.K.

Trans. Faraday Soc, 32, 1397 (1936).

35] Glasoe, P.K. and Long, F.A.

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

36] Sa.lomaa, P., Schaleger, L.L., and Long, F.A.

J. Am. Chem. Soc, 86, 1 (1964).

37] Robinson, R.A. and Bower, V.E.

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

38] Robinson, R.A. and Bower, V.E.

J. Res. Natl. Bur. Std., 69A , 4^9 ( 1965 )

.

39] Bower, V.E. and Robinson, R.A.

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

40] Robinson, R.A. and Bower, V.E.

J. Res. Natl. Bur. Std., 69A, 365 (1965).

4l] Petriconi, G.L., Montefinale, A.C., and Papee, H.M.

J. Sci. Instruments, 42, 662 (1965).

42] Bates, R.G. and Wichers, E.

J. Res. Natl. Bur. Std., 59, 9 (1957).

43] Taylor, J.K. and Smith, S.W.

J. Res. Natl. Bur. Std., 63A, 153 (1959).

44] Analysis and Purification Section: Summary of Activities

July 1964 to June 1965 (J. K. Taylor, Ed. J.

NBS Technical Note 273 (October 1, 1965), PP. 14-16.

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