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Electrochemical

Analysis Section

Summary of Activities

July 1968 to June 1969

503

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U.S. DEPARTMENT OF COMMERCE

National Bureau of Standards
NATIONAL BUREAU OF STANDARDS

National Bureau of Standards ' was established by an act of Congress March 3, 1901 . Today,

in addition to serving as the Nation's central measurement laboratory, the Bureau is a principal

focal point in the Federal Government for assuring maximum application of the physical and

engineering sciences to the advancement of technology in industry and commerce. To this end

the Bureau conducts research and provides central national services in four broad program

areas. These are: (1) basic measurements and standards, (2) materials measurements and

standards, (3) technological measurements and standards, and (4) transfer of technology.

The Bureau comprises the Institute for Basic Standards, the Institute for Materials Research, the

Institute for Applied Technology, the Center for Radiation Research, the Center for Computer

Sciences and Technology, and the Office for Information Programs.

THE INSTITUTE FOR BASIC STANDARDS provides the central basis within the United

States of a complete and consistent system of physical measurement; coordinates that system with measurement systems of other nations; and furnishes essential services leading to accurate and

uniform physical measurements throughout the Nation's scientific community, industry, and commerce. The Institute consists of an Office of Measurement Services and the following technical

divisions:Applied Mathematics—Electricity—Metrology—Mechanics—Heat—Atomic and Molecular Physics—Radio Physics -—Radio Engineering - —Time and Frequency -—Astrophysics '-' —Cryogenics.-

THE INSTITUTE FOR MATERIALS RESEARCH conducts materials research leading to im- proved methods of measurement standards, and data on the properties of well-characterized

materials needed by industry, commerce, educational institutions, and Government; develops,

produces, and distributes standard reference materials; relates the physical and chemical properties of materials to their behavior and their interaction with their environments; and provides

advisory and research services to other Government agencies. The Institute consists of an Office of Standard Reference Materials and the following divisions: Analytical Chemistry—Polymers—Metallurgy—Inorganic Materials—Physical Chemistry.

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provides technical services designed to aid Government agencies in the selection, acquisition,

and effective use of automatic data processing equipment; and serves as the principal focus

for the development of Federal standards for automatic data processing equipment, techniques,

and computer languages. The Center consists of the following offices and divisions: Information Processing Standards—Computer Information — Computer Services — Sys- tems Development—Information Processing Technology.

THE OFFICE FOR INFORMATION PROGRAMS promotes optimum dissemination and

accessibility of scientific information generated within NBS and other agencies of the Federal government; promotes the development of the National Standard Reference Data System and a system of information analysis centers dealing with the broader aspects of the National Measure- ment System, and provides appropriate services to ensure that the NBS staff has optimum accessibility to the scientific information of the world. The Office consists of the following

organizational units:

Office of Standard Reference Data—Clearinghouse for Federal Scientific and Technical Information : ' —Office of Technical Information and Publications—Library—Office of Public Information—Office of International Relations.

1 Headquarters and Laboratories at Gaithersburg, Maryland, unless otherwise noted: mailing address Washington, D.C 20234.

- Located at Boulder. Colorado 80302.

3 Located at 5285 Port Royal Road. Springfield. Virginia 22151.
UNITED STATES DEPARTMENT OF COMMERCE

Maurice H. Stans, Secretary

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

NBS TECHNICAL NOTE 503

ISSUED AUGUST 1969

Nat. Bur. Stand. (U.S.), Tech. Note 503, 99 pages (Aug. 1969)

NBTNA

Electrochemical Analysis Section

Summary of Activities

July 1968 to June 1969

Edited by Roger G. Bates

Electrochemical Analysis Section

Analytical Chemistry Division

Institute for Materials Research

National Bureau of Standards

Washington, D.C. 20234

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 $1
FOREWORD

The Analytical Chemistry Division was established as

a separate division at the National Bureau of Standards on

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

Materials Research in the February 1, 1964, reorganization.

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

competences from activation analysis and atomic absorption

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

and in turn the sections which they comprise, are charged

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

material or service analysis. In addition, it is their

responsibility to inform others of their efforts.

Formal publication in scientific periodicals is a highly

important output of our laboratories. In addition, however,

it has been our experience that informal, annual summaries

of progress describing efforts of the past year can be very

valuable in disseminating information about our programs.

A word is perhaps in order about the philosophy of these

yearly progress reports. In any research program a large

amount of information is obtained and techniques developed

which never find their way into the literature. This

includes the "negative results" which are so disappointing

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

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

important enough to warrant publication--yet can be of great

interest and use to specialists in a given area. Finally

there are the experimental techniques and procedures, the

designs and modifications of equipment, etc., which often

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

11
Thus our progress reports endeavor to present this

information which we have struggled to obtain and which we

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

it appears will not be treated fully in regular publications

are considered in some detail here. Other results which are

being written up for publication in the journal literature

are covered in a much more abbreviated form.

At the National Bureau of Standards publications such

as these fit logically into the category of a Technical Note

In 1969 we plan to issue these summaries for all of our sections. The following is the fifth annual report on progress

of the Electrochemical Analysis Section.

W. Wayne Meinke, Chief

Analytical Chemistry Division

111
PREFACE

This is the fifth annual progress report of the

Electrochemical Analysis Section of the Analytical Chemistry

Division. It covers the fiscal year 1969, which began on

July 1, 1968, and ended on June J>0, 1969

The Section concerns itself particularly with ionic

processes occurring in solution, with areas of analytical

measurement where ionic equilibria play a part, and with

the explanation of solution behavior in terms of the

interactions of ionic solutes with solvent molecules. The

special interests of the members of the Section lie for the

most part in solution electrochemistry, electrochemical

thermodynamics, and ionic transport phenomena. In the

context of modern analytical chemistry, their competences

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

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

research and sample aspects. During the fiscal year just

ending, about 40 percent of the total effort of the Section

was devoted to research and about 60 percent to the development of standard reference materials.

The permanent staff of the Section remained at four

during the year. In addition, the Section enjoyed the

valuable services of Dr. Marinus Alfenaar of the University

of Utrecht from July until late September, when his temporary appointment was concluded. In October Dr. Harry P.

Thun of the University of Ghent arrived to spend a year in

the Section as a guest worker. Furthermore, the Section

was augmented during the summer months of 1968 by the

return of Edgar S. Etz and James T. Clark, Jr., and by the

presence of Robert S. Carter, a student at Landon School.

iv
Both Mr. Etz and Mr. Clark are graduate students at Clarkson

College of Technology, where they are working under the

direction of Professor Donald Rosenthal. Professor Rosenthal

remains a consultant to the Section.

With a permanent staff of only four, it has obviously

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

the past year has been limited to emf studies and conductance

measurements. Unlike the formal publications that appear in

the scientific journals, this report summarizes the total

program of the Section, incomplete studies as well as finished work, and stresses its relationship to the missions

of the Division and Institute.

In order to specify adequately the procedures, it has

been necessary occasionally to identify commercial materials

and equipment in this report. In no case does such identification imply recommendation or endorsement by the National

Bureau of Standards, nor does it imply that the material or

equipment identified is necessarily the best available for

the purpose.

The assistance of Mrs. Rosemary Maddock, Mrs. Marguerite

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

report is gratefully acknowledged.

Roger G. Bates, Chief

Electrochemical Analysis Section

Washington, D.C.

June 30, 1969

v
TABLE OF CONTENTS

1. FACILITIES AND EQUIPMENT

2. ACIDITY MEASUREMENTS

Concepts and Program

Standards for Clinical pH Measurements. . Tri s ( hydroxymethy1 ) aminomethane in

Isotonic Saline

Dissociation of Protonated ,r Bis-Tris"

in Water

Measurement of Acidity in Sea Water . . . Acidity Standards for Mixtures of

2-Methoxyethanol (80 Weight Percent)

and Water (20 Weight Percent

Media of Low Ionic Strength ....

Media of I = 1.0

3. STANDARDIZATION OF ION-SELECTIVE ELECTRODES

The pH Convention

Proposed Conventional Activities for

Calibrating Ion-Selective Electrodes. . Measurement of Electrode Response ....

Equipment

Experimental Results

Sodium Electrode

Fluoride Electrode

Calcium Electrode

Bromide Electrode

Liquid-Junction Errors

PAGE

1

6

8

15

16

17

17

19

20

21

25

25

25

25

25

27

29

30

vi
PAGE

4. EQUILIBRIUM DATA FOR IONIC SYSTEMS 32

Dissociation of Phosphoric Acid in

Deuterium Oxide 32

Measurement of the First

Dissociation Constant 32

Thermodynamic Quantities for

the Dissociation J>6

Deuterium Isotope Effects 37

First Dissociation Constant of Piperazinium

Ion in 50 Weight Percent Methanol-Water

Solvent 4l

Dissociation of Protonated Tris (hydroxy

methyl)aminomethane in Aqueous

Alcoholic Solvents 43

Thermodynamics of Hydrochloric Acid in

N-Methylpropionamide Solvents 46

Thermodynamics of Hydrochloric Acid in

Mixtures of 2-Methoxyethanol

80 Weight Percent) and Water

'20 Weight Percent) 49

Vapor Pressures of the Solvent .... 49

Dielectric Constants of the Solvent. . 52

Determination of Acidity Functions . . 52

5. OCEANOGRAPHIC CONDUCTANCE STANDARDS 55

Apparatus 55

Calibration of Cells 57

Conductance of Natural and Synthetic

Sea Water 62

6. PREPARATION AND PROPERTIES OF SOLVENTS ... 63

Dimethylformamide 63

2-Methoxyethanol 63

Sulfolane 64

Vll
PAGE

7. STANDARD REFERENCE MATERIALS ACTIVITY .... 67

Summary f . 67

Standard Materials for Acidity

Measurements 68

Standard Materials for Conductivity

Measurements 69

Acidimetric Assay of Bilirubin 70

General Considerations 70

Conductance Measurements 71

Emf Titrations 71

A Standard for Ion-Selective Electrodes . . 75

8. FOREIGN-CURRENCY PROGRAMS 76

9. IMPACT OF SECTION'S ACTIVITIES 77

10. PERSONNEL 78

11. COMMITTEE MEMBERSHIPS 79

12. PUBLICATIONS AND MANUSCRIPTS, JULY 1968

TO JUNE 1969 80

13. TALKS, JULY 1968 TO JUNE I969 82

14. LIST OF REFERENCES 85

viii
LIST OP FIGURES

FIGURE NO. PAGE

1. Emf laboratory equipped, for the study

of ion-selective electrodes 1

2. Equipment for standardization of

clinical pH measurements 7

3. Pathways leading to the activity

coefficients of single ionic

species 23

4. Mean activity coefficients of sodium

fluoride from a cell without

liquid junction 26

5. Mean activity coefficients of calcium

chloride from a cell without

liquid junction . 27

6. Comparison of "experimental" values

of the activity coefficient of

calcium ion with values furnished

by four conventions 29

7. Liquid- junction errors in the

determination of pCl 30

8. Removal of activity coefficient effects

by extrapolation in the determination

of the dissociation constant of

phosphoric acid in deuterium oxide . . 34

9. Deuterium isotope effects (ApK)

for weak acids at 25 °C 38

10. Cell vessel for determining pK by

potentiometric titration without

a liquid junction 44

11. Comparison of the activity coefficient

of hydrochloric acid in NMP with

that in water 47

12. Apparatus for measuring the vapor

pressures of solvents 50

IX
LIST OP FIGURES (con't)

FIGURE NO.

13.

14.

15.

16.

17.

18.

Janz-Mclntyre bridge for conductivity

measurements

Two-electrode conductivity cells . . Oil thermostat for conductivity

measurements

Four-electrode conductivity cell . . Still for the purification of

organic solvents

Cell for the potent iometric titration

of bilirubin

PAGE

56

57

59

60

65

73

LIST OF TABLES

TABLE NO.

1. Emf of the cell: Pt;Ho, Tris-HCl(m),

Tris(m), NaCl(0.l6-m), AgCljAg

at 25 °C. Values of P(aHYcl ) . .

2. pKa of protonated Tf bis-tris" from

to 50 °C

3. Thermodynamic quantities for the

dissociation of protonated

rr bis-tris u at 25 °C

4. pajj values for buffer solutions

composed of bis-tris(m) and

bis-tris-HCl(m) at 25 °C

5. Residual liquid- junction errors for

bis-tris buffer solutions at 25 °C

6. Standards for pA measurements . . .

7. Dissociation constant of phosphoric

acid in deuterium oxide

from 5 to 50 °C

PAGE

8

11

11

13

13

24

35

x
LIST OP TABLES (con't)

TABLE NO. PAGE

8. Thermodynamic constants for the

dissociation of phosphoric acid

in deuterium oxide and water

at 25 °C 36

9. Deuterium isotope effect on the

thermodynamic constants for the

dissociation of weak acids at 25 °C . 40

10. Values of p( K) for the dissociation

of protonatied weak bases in 50 wt %

methanol-water at 25 °C and related

solvent effects 42

11. pK values for tris«H+ in 50 wt %

aqueous alcohols 45

12. Vapor pressures of a mixture of

2-methoxyethanol (80 wt %) and

water (20 wt %) from 10 to 50 °C . . 51

13. Values of Pw (ajjYQ]_) and WY± for

hydrochloric acid in a mixture

of 2-methoxyethanol (80 wt %)

and water (20 wt %) at 25 °C .... 53

14. Standard reference materials for

acidity measurements 68

15. Potentiometric titration of bilirubin . 74

xl
ELECTROCHEMICAL ANALYSIS SECTION: SUMMARY OF ACTIVITIES,

JULY 1968 TO JUNE 1969

Edited by Roger G. Bates

ABSTRACT

This survey of the activities of the Electrochemical

Analysis Section, Analytical Chemistry Division, covers

the period July 1968 to June 1969 . An attempt is made to

summarize a year's progress on the technical projects of

the Section, to indicate the composition and capabilities

of the unit as a whole, and to stress the Section's role

in the mission of the Institute. Brief summaries of several

lines of work under way are given. The concept of a broad

program on acidity measurements is outlined, and new efforts

to provide assistance in clinical acid-base measurements are

summarized. Progress toward reference standards for the

calibration of ion-selective electrodes is reported.

Equilibrium data for phosphoric acid in heavy water, for

piperazinium ion in 50 wt % methanol-water, and for hydrochloric acid in a solvent composed of 2-methoxyethanol

(80 wt %) and water (20 wt %) have been obtained. Considerable effort has been devoted to the development of new

standard reference materials, including standards for conductivity measurements. Lists of publications, talks, and

committee assignments of staff members are given.

Key Words

Acidity, Analysis, Conductivity, Electrochemical analysis,

Electrochemistry of solutions, Ionic activity, Ion-selective

electrodes, pD measurements, pH measurements, Medium effects,

Solvent effects, Standard reference materials

xi:
1. FACILITIES AND EQUIPMENT

Excellent facilities for the Section's programs are

provided on the second floor of the Chemistry Building at

the Gaithersburg site of the National Bureau of Standards

The space devoted to research and development of standard

reference materials consists at present of a conductance

laboratory and four well-equipped emf laboratories. The

instrumentation room has been outfitted to accommodate

the instrument for measuring dielectric constants of

solvents, an automatic titrator, and several pH meters.

One of the emf laboratories, specially equipped for the

study of ion-selective electrodes, is shown in figure 1.

Figure 1. Emf laboratory equipped for the study of

ion-selective electrodes.
During the year, one of the four emf laboratories was

devoted to studies of acid-base behavior in nonaqueous and

mixed solvents. A second was devoted to work in deuterium

oxide with deuterium gas electrodes and to the study of

aqueous buffer systems of clinical interest. Investigations

of ion-selective electrodes were conducted in a third, while

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

over to studies of acid-base processes in sea water. The

highly precise equipment for the measurement of conductivity

was used throughout the year in establishing conductivity

standards for oceanographic purposes.

This emf installation is sufficiently large to justify

an automatic data read-out center capable of recording

automatically and simultaneously the emf of all the cells

under study in each of the laboratories. A multichannel

digital voltmeter with an accuracy of 10 |j,V is needed. It

would be coupled to a print-out unit from which a record

of all the emf data could be obtained and, when desired,

a tape could also be produced for computer analysis.

Measurements of all cells at desired intervals could be

programmed in advance.

Because of staff limitations, it has not yet been

possible to initiate work on the accurate measurement of

pH at high temperatures and pressures or to supplement our

present capability with calorimetric studies of ionic

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

additional facilities and equipment will be needed.

(R. G. Bates)
2. ACIDITY MEASUREMENTS

A. Concepts and Program

The measurement of acidity continues to be one of the

major concerns of the Section. The concepts and approaches

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

progress reports for 1966 [1] and 1967 [2],

The NBS pH scale is fixed by a series of standard

reference solutions chosen to cover uniformly that region

of the scale in which the emf of the pH cell is a linear

function of the pH of the solution in which the electrodes

are immersed. Corrections are made, if necessary, for

imperfections in the response of the glass electrode. Unlike the pH scale of the British Standards Institution [J>]

,

which recognizes only a single primary standard, the NBS

approach permits a series of primary reference solutions.

The search for new and better standards is going on continually. A new standard may be selected because it extends

the useful range of the standard scale or because it shows

a stability superior to those of existing standards in its

pH range

.

The pD scale in heavy water bears the same fundamental

relationship to the deuterium gas electrode that the pH

scale bears to the hydrogen gas electrode. Basically,

the phenomena which are regarded as a manifestation of

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

the hydrogen ion. For this reason, the pD scale is not

related to the pH scale. The same methods can nonetheless

be utilized for establishing acidity standards in deuterium

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

water.
During the past year, two new standards for the pH

scale and three standards which fix a practical pD scale in

deuterium oxide were described in published papers [4,5].

The new pH standard solutions contain potassium dihydrogen

citrate (0.05 molal) and a mixture of sodium bicarbonate

and sodium carbonate (each 0.025 molal). Solutions of the

same salts in the same molalities, prepared in deuterium

oxide, serve as pD standards, together with a mixture of

potassium dideuterium phosphate and disodium deuterium

phosphate (each 0.025 molal). Sodium bicarbonate and sodium

carbonate were added to the list of Standard Reference

Materials. The citrate, phosphate, and carbonate salts

serve equally well for preparing pD and pH standard solutions; exchange of hydrogen for deuterium when the protium

salts are dissolved in heavy water alters the acidity so

slightly that it causes no concern. A source of pure potassium dihydrogen citrate in quantity is still being sought.

The extension of the pH concept to media other than

pure water has received considerable attention in the

Section during the past few years. Solvent effects on

acidity are not only of great theoretical interest but have

an important bearing on the behavior of equilibrium systems

of analytical interest. Practical acidity scales in nonaqueous and partly aqueous solvents would, for example,

greatly facilitate the study of metal complexation with a variety of ligands that are themselves proton donors or

acceptors. As a guest worker in the Section, Dr. Harry P.

Thun is turning his attention to the establishment of a useful acidity scale for a mixed solvent system with which

a range of dielectric constants can be achieved. For this

study, mixtures of 2-methoxyethanol (methylcellosolve) and

water were chosen. This work is aimed at setting up standard scales for the measurement of acidity at low ionic

4
strengths in these mixed solvents and also in media of

ionic strength 1.0 (l molal with respect to sodium

perchlorate) . Complexation studies are often conducted

in media of a constant, rather high, ionic strength in

order to minimize changes in activity coefficients. It

is therefore useful to have available secondary acidity

standards in special solvent media of this sort.

With the increased attention being given to clinical

measurements, secondary standards for other media of

special significance to biomedicine are assuming importance in the Section's future plans. Of most immediate

concern is isotonic saline, or 0.16M sodium chloride.

Preliminary work in this medium has already begun. These

studies parallel closely the work in sea water which has

been conducted at a relatively low level during the past

two years

(R. G. Bates)
B. Standards for Clinical pH Measurements

1 . Tris(hydroxymethyl)aminomethane in Isotonic Saline

Acidity measurements in biological media with glass

electrodes, standardized in the usual aqueous reference

buffer solutions, often involve a considerable error due

to the change in liquid- junction potential when the glass

electrode is transferred from the relatively dilute standard

buffer to a medium of higher ionic strength. This source of

error can be minimized or eliminated by developing a new set

of reference buffers in a medium of a composition close to

that of the biologic fluid. Furthermore, buffer solutions

of this sort are very useful for pH control in biomedical

research

.

A sodium chloride solution with a constant total ionic

strength of . 16 (isotonic saline) approximates closely the

colligative properties of blood plasma. Tris(hydroxymethyl)-

aminomethane (tris, THAM) has been selected tentatively as

the first pH standard in this new medium. Emf measurements

of the cell

PtjH

2 (g), Tris-HCl(m), Tris(m), NaCl (0.l6-m), AgCljAg

were made from to 50 °C . The concentration of free base

was kept equal to that of the hydrochloride, and enough

sodium chloride was added to make the total ionic strength

equal to 0.16. The molality of the base was varied from

0.005 to 0.05 rnol kg . The acidity function p(aHYp. ) was derived from the measured emf E of the cell, and

its variation with buffer composition was studied. The

following equation was used:
E-E

p(aH

Ycl ) = + log m

cl (1)

where E° is the standard emf of the cell and k is

(RT In 10)/F. The results at 25 °C are given in table 1.

Cell vessels and constant-temperature bath can be seen

in figure 2.

Figure 2. Equipment for standardization of clinical

pH measurements.

7
Table 1. Emf of the cell: Pt;H" 2

, Tris-HCl(m), Tris(m),

NaCl(0.l6-m), AgCl;Ag at 25 °C. Values of

p(aH

Ycl ).

m E PCagYd

0.005004 0.76292 8.342

.01001 .76322 8.3^7

.01999 .763^0 8.350

.0^001 .76^46 8.351

.05001 .76354 8.353

To date, only one series of measurements has been made

When the work Is resumed, sufficient data to permit the

assignment of primary standard pH values to a tris buffer

in isotonic saline will be obtained.

2. Dissociation of Protonated "Bis-Tris" in Water

A new buffer material suitable for biochemical work

in the pH range from 5. 5 to 7-5 has recently been made

available commercially. This compound, 2, 2-bis (hydroxy

methyl ) -2, 2 ' 2" -nitrilotriethanol [ (H0CH2

CH

2

) 2

NC ( CH 20H ), ]

,

or "bis-tris", is a solid, easily purified, water-soluble

compound with pK value of 6.46 at 25 °C [6].

a

The structure of bis-tris is closely related to that

of tris. Since it does not contain primary amine groups

which can react with carboxyl groups, bis-tris is preferred

over tris for pH control in studies of enzyme-catalyzed

reactions. An accurate knowledge of the acidic dissociation constant of bis-tris over a range of temperature will

be useful in applying the new buffer system to pH measurement and regulation in biological work.

8
Bis-tris was purified by recrystallization from absolute ethanol (assay 99*98 percent as determined by potentiometric titration). An attempt to prepare the bis-tris

hydrochloride by passing anhydrous hydrogen chloride gas

into a solution of bis-tris dissolved in methanol yielded

a product whose chloride content was only 99.80 percent of

the theoretical. For this reason, the solutions used for

the emf measurements of the cell

Pt;Hp(g), Bis-Tris-HCl(m), Bis-Tris(m), AgCljAg

was prepared from bis-tris and a standard solution of

hydrochloric acid. The molality (m) of the free base was

kept equal to that of the hydrochloride salt; m was varied

from 0.01 to 0.1 mol kg" . The temperatures at which the

measurements were made ranged from to 50 °C.

The emf of the cell is related to the dissociation

constant of protonated bis-tris by the equation

(E-E°)F m^ + 7BT+ YC1- pK = + log mri - + log -2±- + log J2± ±±- (2)

2.3026 RT L' ± m

BT YBT

where E° is the standard emf of the cell. The "apparent"

dissociation constant pK' is defined as follows: * a

2AI

pK

a = p^ci) T7 (5)

i+Bai 1/cf

where A and B are the Debye-Huckel constants, I is the

total ionic strength, and a is the ion-size parameter.

For these measurements, the buffer ratio was unity,

and hydrolysis of the buffer substances was negligible.
The true value of pK is equal to pK* at 1=0; it was obtaina a

ed by extrapolating values of pK

1 to zero ionic strength,

a

using the method of least squares. Table 2 lists the experimental pK values together with the a values which gave

a

the smallest standard deviation (a) for pK . The method of least squares was also used to express

)K as a a the equation

the pK as a function of absolute temperature according to

a

pK = - + B + CT (4)

a <J>

Between and 50 °C, pK is given by

a

pK = 12&7.855/T + 2.7905 - 0.00210396T (5) a

Standard changes of enthalpy, entropy, and heat capacity

in the dissociation of protonated bis-tris were derived

from the constants of equation 5, and the values at 25 °C

are summarized in table ~$

.

10
Table 2. pK of protonated "bis-tris" from to 50 °C

.

a

t, °c ^K

a

Ion-size

parameter (A) a

6.9314 1.9 0.0002

5 6.8^44 2.0 .0002

10 6.7428 2.2 .0001

15 6.6535 2.4 .0001

20 6.5665 2.6 .0001

25 6.4835 2.8 .0001

30 6.4012 2.7 .0001

35 6.3212 2.9 .0002

40 6.2441 3.2 .0002

45 6.1689 3.2 .0002

50 6.0977 4.4 .0005

Table 3. Thermodynamic quantities for the dissociation

of protonated "bis-tris" at 25 °C.

AH°: 28,238 J mol

-1

AS°: -29.4 J K" 1 mol" 1

AC : 24 J K" 1 mol-1

11
From the emf data one can also derive acidity functions which may be regarded as reference data useful in

the study of biological systems at these particular total

ionic strengths. The pa^ values for bis-tris buffer

solutions in water were evaluated with the aid of two relations. The first was used to interpolate p(aRyr

,-, ) at

even values of m:

1/2 2AI x p (¥ci' = pK

a

" BI +

1+Bil l/2 (6)

where -B is the slope of the plot of pK ' as a function

a

of the ionic strength. The second equation was

pajj = -log ajj = p(aH

7cl ) + log yc1- (7)

The single-ionic activity coefficient y^-was defined

by the Bates-Guggenheim convention [7]

Ml/2

- los ^cr = 7JZ

(8)

1+1.5 I

The pa^ values for equimolal buffer solutions are given

in table 4.

12
Table 4. pa„ values for buffer solutions composed of

bls-tris(m) and bis-trls .HCl(m) at 25 °C.

m paH

0.01 6.532

.02 6.551

.03 6.565

.04 6.576

.05 6.586

.06 6.595

.07 6.603

.08 6.610

.09 6.617

.10 6.623

Table 5. Residual liquid- junction errors

for bis-trls buffer solutions at 25 °C.

m=I paH

0.01 6.532

0.02 6.551

0.05 6.586

0.1 6.623

pH - pa^

pH A

3

6.517 -0.015

6.529 -0.022

6.547 -0.039

6.566 -0.057

13
The internal agreement of these pa^ values with the

pH scale fixed "by the NBS primary standards was examined

by the intercomparison of bis-tris buffer solutions with

the 0.025 m phosphate buffer [pH(S) = 6.865 at 25 °C ] in

the following cell

Standard

Pt;H , solution 2

(s)

KC1 Hg, KC1

satd. calomel satd.

Bis-Tris

solution, H j Ft

(x)

where each vertical line represents a liquid junction.

The operational pH of bis-tris solutions is given by the

equation

pH = 6.865 + (9)

(RT In 10)/P

where E^ and E~ are the potentials of the hydrogen electrodes in the compartments containing bis-tris and standard

solution, respectively, measured with respect to the

calomel reference electrode. If there is no residual

liquid- junction potential, the operational pH values should

be in close agreement with the corresponding assigned pa^

values. The degree of internal consistency can be judged

by the results given in table 5.

The deviations (a) noted for bis-tris may be compared

with deviations for the following nitrogen-base buffer

systems at I = 0.02: ammonia A = -0.014, 4- amino pyridine

A = -0.008, and tris A = -0.020 [8], The results suggest

that the more dilute bis-tris buffer solutions should be

used for pH standardization purposes.

(M. Paabo)

14
C . Measurement of Acidity In Sea Water

This work has as its objective the development of

standards for pH measurements in media approximating sea

water in composition. These reference standards would be

used to adjust pH equipment for the accurate determination

of hydrogen ion concentration or activity in oceanographic

studies. Materials for the preparation of the standards

would be made available through the Standard Reference

Materials program.

During the reporting year, this project was active in

the Section only during the summer months which James T.

Clark, Jr., spent at NBS as a guest worker. During the

academic year, Mr. Clark continued the work as part of his

graduate program at Clarkson College of Technology.

Electromotive force data for the cell

PtjH

2

, HC1 in SSW, AgCljAg

at several concentrations of synthetic sea water (SSW) have

now been obtained at 5* 15* 25* and 35 °C. The molality of

hydrochloric acid was varied over a sufficiently wide range

to allow the standard potentials for the cell to be obtained by the usual extrapolation procedure. The data are now

being analyzed with the aid of a computer.

Amine buffer systems have been chosen as secondary

standards for sea water, in order to avoid undesirable

interactions occurring with phosphates, borates, and other

organic buffer materials. The emf measurements which will

lead to a certification of the pH of a buffer solution

composed of tris(hydroxymethyl)aminomethane and its hydrochloride in SSW have been started.

(R. G. Bates)

15
D. Acidity Standards for Mixtures of 2-Methoxyethanol

(80 Weight Percent) and Water (20 Weight Percent)

Nonaqueous solvents and mixtures of organic solvents

with water are becoming of increasing interest to complexation chemists. Indeed, many promising organic ligands

cannot be investigated in aqueous solutions as they are

too insoluble in water. Of the many organic solvents available, few meet the chief requirements of stability with

time and acidity, ability to dissolve many organic and

inorganic compounds, sufficiently high dielectric constant,

and a high boiling point. In addition, ease of purification is of the utmost importance. As the overwhelming

majority of stability data for complexes have been obtained

through emf measurements, mainly pH measurements, it becomes urgent to establish reliable standard buffer solutions

in suitable solvents.

After unsuccessful experiments with dimethylformamide

were concluded (see Section 6. A), attention was given to

2-methyoxyethanol, known commercially as methylcellosolve.

This liquid is an extremely good solvent with a high boiling point (123.4 °C). It has low hygroscopicity and is

readily miscible with water. As it has a rather low dielectric constant (16.0 at 25 °C), it seemed preferable to

start the investigation with mixtures of the solvent with

water. The mixture containing 80 wt percent of the organic

solvent seemed particularly attractive as the dissociation

constants of a number of organic compounds had already been

determined in this mixture by Simon, Heilbronner, and their

co-workers [9,10].

16
1. Media of Low Ionic Strength

Insofar as the ligands are acids or salts of organic

acids, most measurements intended for the determination

of stability constants are carried out in the region pH 2

to 6. Moreover, most complexation chemists work in media

of constant ionic strength, most frequently using sodium

perchlorate or lithium perchlorate as inert electrolytes.

Even in media of low ionic strength, therefore, the usual

NBS standard buffers cannot be used, as most of them contain potassium ions which cause the precipitation of insoluble potassium perchlorate. Hence, two new buffer

solutions will be examined, namely O.O^M oxalic acid +

0.016M sodium oxalate, and 0.025M sodium dihydrogen

phosphate + 0.025M disodium hydrogen phosphate, both in an

80 wt percent 2-methoxyethanol - 20 wt percent water solvent.

For both buffers, the pH(S) will be determined in the same

way as for other NBS buffers, namely through emf measurements in cells without liquid junction, together with the

Bates -Guggenheim convention (equation 8) for the activity

coefficient of the chloride ion.

2. Media of I = 1.0

Inasmuch as most stability constants are determined

in media of high ionic strength, very often I = 1.0, it

will be of importance to assign pH(s) values to buffers of

ionic strength I = 1.0. The same two buffer solutions

selected for low ionic strengths will be used, and the

ionic strengths will be adjusted to 1.0 by the addition of

suitable quantities of sodium perchlorate. The pH(s)

values can be obtained in three different ways:

a. through comparison with the primary NBS buffers

in cells with a liquid junction,

17
b. through comparison with the buffers described in

the previous section, also in cells with a liquid junction,

and

c. through measurements in cells without liquid

junction after choice of a reasonable convention for the

activity of the chloride ion in 1M solutions of sodium

perchlorate. In essence, procedures a and b will lead

to secondary pH standards, while c will yield primary

standards. In work under way, these three approaches will

be examined in detail.

(H. P. Thun)

18
5. STANDARDIZATION OF ION-SEIECTIVE ELECTRODES

Recent progress in the development of a variety of

commercial ion-selective electrodes has given the chemist

access to the reasonably selective determination, by

electrochemical methods, of an ever-expanding number of

individual ionic species. The application, seemingly

unlimited, that these electrodes may have in the fields

of chemistry and biology, in industry and in the clinical

laboratory emphasizes the pressing need for standardization

methods and standard reference materials for use with ionselective electrodes.

A proposed method of standardization, based on the same

concepts used in the standardization of pH (hydrogen ion

activity), was outlined in the last progress report [11]

and in a lecture presented at the Symposium on Ion-Selective

Electrodes held at NBS Gaithersburg, January J>0 and Jl,

1969, and soon to be published [12]. This procedure is

being recommended because of the wide acceptance of the

convention used to arrive at the practical activity of the

hydrogen ion in fixing NBS standard pH scale. The convention entailed in the definition of the NBS standard pH scale

is reasonable as well as consistent with the response of the

glass electrode and other practical electrochemical sensors.

It is the purpose of the present line of investigation

to demonstrate how the pH convention can be extended to a

variety of activity standards for ion-selective electrodes.

There are, however, several sources of error which influence

this measurement of ionic activity. These errors, which

include imperfect response of the electrodes, variation of

potential with time or other influences (stirring, etc.),

and errors due to liquid- junction potentials, are the

subject of study in the Section.

(B. R. Staples)

19
A. The pH Convention

The convention on which a practical pH scale has been

based is discussed in detail in a host of previous publications. The proposed method of adapting the convention to

provide activity scales for ion-selective electrodes will

be described in the sections to follow.

For pH assignment, the emf of a cell without liquid

junction is measured and an acidity function p(aTjYf

, 1) is

evaluated from it (see equation l):

P(aH

Ycl ) = -log(mH+YH+Ycl -) (10)

The activity of the hydrogen ion in a chloride-free

buffer solution is formally related to the acidity function

in the limit of rru-^O by

paH

= p(aHYcl

)° + log yc1 _ (ll)

The paH

is then calculated by introducing a convention [7]

f° r ?C1- : _ AI l/2

log Yci- = 77i

(12)

1+1.5 i

'

The operational pH value for an "unknown" solution X

is defined in terms of a reference standard pH(S), and the

values of the emf of a pH cell containing the standard (E„)

and the unknown (EY

) by

Ey-Eq

pH(X) = pH(S) + (13)

(RT In 10)/P

In the NBS standardization procedure, the conventional pa„

values obtained as described above for selected buffer

solutions are identified with pH(s) in the operational

definition (equation 13).

(B. R. Staples)

20
B. Proposed Conventional Activities for Calibrating

Ion-Selective Electrodes

The experimental arrangement by which "pA" measurements with an electrode selective for the ion A are made

is entirely analogous to that for pH measurements. The

electrode is usually combined with a suitable reference

electrode such as a calomel electrode in a cell with liquid

junction:

Electrode (sel. for A), Soln. x| Bridge, Reference Electrode

where the liquid junction is indicated by a vertical line.

It is reasonable, therefore, to define operational pA

values in the same manner that operational pH values are

defined. If the ion A is a cation with n unit charges

(EY

-E )nP

PA(X) - PA(S) +

R/ln

S

10 (14)

and if A is an anion with n unit charges

(EY-E

q )nF

PA(X) = PA(S) - X S

(15)

RT In 10

The only difference between equations 14 and 15 is the

sign of the last term.

21
The convention proposed for the assignment of single

ionic activities to standard solutions of the ions A is the

same one used in the calculation of hydrogen ion activities,

namely equation 12. This convention yields values of the

activity coefficient of chloride ions and therefore permits

the sodium ion activity in a sodium chloride solution to

be obtained from the known mean activity coefficients

of NaCl:

Y+ (NaCl)

Wh--= ( l6 ) Na

YdIf the activity coefficients of sodium and chloride ions

are taken to be a function only of temperature and ionic

strength, it becomes possible to use these values as

pathways to other single ion activity coefficients.

Using the values of y„,_, for example, one can obtain,

through solutions of chlorides, the activity coefficients

of cations other than sodium as shown schematically in

figure 5- Similarly, one can use Y-vr a + and proceed through

sodium salt solutions to other single anion activity

coefficients. As in the standardization of pH, these

conventional values of pa„ for selected standard reference

solutions are identified with pA(s) in the operational

definitions set forth in equations 14 and 15 . In order to

demonstrate the utility of the activity scales set up in

this way, however, it remains to be shown that the responses of the electrodes are consistent with the defined

activities of the respective ions in solutions of a standard electrolyte.

22
Convention

*•>,NO,"

'KCI *\* 'KBr

"7Br- ^KNCh -*7

NO,

'CoCI, rCa2 + VCaBr2~T*"7Br- 7Ca(N03) 2

"' *

7N03-

Figure j5. Pathways leading to the activity coefficients

of single ionic species.

Table 6 lists the standard pA values for solutions

which have thus far been proposed. Other ions will pose

more difficult problems, especially when no completely

ionized reference salt can be found. Cadmium, for example,

forms highly complexed species with chloride ion in solution, yet a cadmium-selective electrode senses only the

free ion and must be standardized in a solution in which

the activity of that ion is known. One possibility in

cases of this sort is to proceed through the perchlorate

23
salts, most of which are highly dissociated. Further

study is being devoted to elaborating procedures for ions

which present special difficulties.

Table 6. Standards for pA measurements.

Materi al Molality pCl pNa pCa pP

NaCl 0.001 3.015 3.015

0.01 2.044 2.044

0.1 1.110 1.108

1.0 0.204 0.162

CaCl

2 0.000333

0.00333

0.0333

3.537

2.653

I.887

0.333 1.104

NaF 0.001 3.015.

0.01 2.048

0.1 1.126

(B. R. Staples)

24
C. Measurement of Electrode Response

1. Equipment

The apparatus for measuring the emf of cells containing

ion-selective electrodes was shown in figure 1 (Section l).

It consists of a vibrating-reed electrometer and a precision

potentiometer which opposes the emf from the electrometer.

The small difference in emf, usually 1 mV or less, is then

amplified and displayed on a strip-chart recorder. It is

hoped that a digital voltmeter with a degree of computer

control may be obtained in the future.

2. Experimental Results

To date, all of the ion-selective electrodes have shown

a nearly ideal response, i.e., a Nernstian slope when the

emf of the cell is plotted as a function of the single ion

activity.

a. Sodium Electrode

Measurements with the sodium glass electrode, Corning

NAS 11-18, were described previously [11]. The electrode

was found to respond ideally in sodium chloride solutions

-4 of concentrations from 10 M to 1 M. Its response was

measured in a cell without liquid junction Na(gl); NaCl(m),

AgCl;Ag, where m is molality. The measurements in sodium

fluoride solutions, described below, depend for their interpretation on the demonstrated Nernstian behavior of the

sodium glass electrode.

b. Fluoride Electrode

The potential of the fluoride electrode, Orion Model

94-09 (lanthanum fluoride type), was measured against each

of the sodium glass electrodes and a saturated calomel

electrode of the fiber type at temperatures of 10, 25, and

40 °C and in a concentration range from 10" M to 1.0 M.

25
The cell without liquid junction is represented by

Na(gl); NaP(m), P(sel)

It was standardized in a solution of sodium fluoride with

a molality of 0.01 mol kg

Mean activity coefficients of sodium fluoride were

calculated from the emf data. The results at 25 °C are

shown in figure 4, where the curve is drawn through the

known activity coefficients of solutions of NaP. The

experimental values are shown as circles.

Figure 4. Mean activity coefficients of sodium fluoride

from a cell without liquid junction.

26
c . Calcium Electrode

The response of a calcium-selective electrode of the

liquid ion-exchanger type (Orion Model 92-20) was studied

in a cell without liquid junction represented by

Ca(sel); CaCl

2 (i), AgCljAg

where I is ionic strength (three times the molality of

calcium chloride). The cell was standardized at I = 0.01,

using the known activity coefficient of calcium chloride

[15] • The mean activity coefficients of calcium chloride

at 25 °C obtained experimentally are plotted in figure 5

as a function of ionic strength. The curve is again drawn

through the mean activity coefficients obtained by other

thermodynamic methods and recorded in the literature [Y3, 14],

The agreement is quite satisfactory, confirming the accuracy

of the response of the calcium electrode in solutions of

calcium chloride.

Figure 5. Mean activity coefficients of calcium chloride

from a cell without liquid junction.

27
In order to examine the response of the calciumselective electrode in mixtures of sodium chloride and

calcium chloride, emf measurements of this same cell with

aqueous salt mixtures of a total molality of 1.0 mol kg"

were made. Poor reproducibility was observed, especially

in solutions of high sodium content. In calcium chloride

solutions at 25 °C, however, the calcium-selective electrode was found to be reproducible to 0.2 to 0.3 mV (0.01

pCa unit) and to reach equilibrium in about 1 min after

immersion in the solution.

The response of the Beckman calcium electrode,

No. 39608, was found to be Nernstian at 25 °C. Further

determinations are under way.

The response of the calcium-selective electrode in

calcium chloride solutions was also explored by measurements of a pCa cell with a saturated calomel reference

electrode and a liquid junction [12]. The cell was standardized at I = 0.01. The results are shown in figure 6,

where -log yr p. , that is pCa + log mn 2+} i s plotted as

a ua

a function of the logarithm of the ionic strength. Curve 1

was calculated by the convention proposed here, curve 2 by

2 the Maclnnes convention, and curve 3 by assuming the z relationship embodied in the Debye-Huckel formula, which

leads to y~ p+ = Y+(CaClp). Curve 4 resulted from setting

yn 2+ equal to the mean activity coefficient of calcium

a chloride, as suggested by Shatkay [15].

28
0.8

+ 0.4 -

CVJoO

8* o

-0.4

Figure 6. Comparison of "experimental" values of the

activity coefficient of calcium ion with

values furnished by four conventions.

d. Bromide Electrode

The Beckman bromide electrode, no. J59602, (solid

membrane type) has been found to respond nearly ideally at

25 °C. Further studies at temperatures of 10 and 40 °C

are now going on.

(B. R. Staples)

29
D . Liquid-Junction Errors

The influence of the composition and concentration of

the solution on the liquid- junction potential error (AE .

)

j

is shown in figure 7 [12] . A value of AE . of 3 mV corresponds to 0.05 pCl unit. The cell consisted of a silversilver chloride electrode used in conjunction with the saturated calomel electrode. It was standardized in a sodium

chloride solution of molality 0.01 mol kg , and all of the

measurements were made at 25 °C. The "true" activity coefficients of chloride ion in each of the four salt solutions

or aqueous salt mixtures were calculated by equation 12.

6mV

3mV

6mV

T 1 1 INFLUENCE OF AE; ON pCl MEASUREMENTS ^

NaCI

NaCI/KN0,(M)

CaCI 2

3 NaCI/KCI (hi)

-2

log I

Figure J. Liquid- junction errors in the determination

of pCl.

30
It is seen that the variation of the liquid- junction

potential error with ionic strength is about the same for

all four types of solutions, but for the solutions of

calcium chloride and the two mixtures it is 1 to 2 mV different from that for pure solutions of sodium chloride.

The corresponding error of 0.02 to 0.04 pCl unit could be

eliminated only by standardizing the cell in a solution

containing the same salt or salt mixture as is present in

the "unknown" solutions.

On the other hand, an error of about 4 mV (0.07 pCl

unit) is incurred when the cell is standardized in 0.01 m

sodium chloride and used for measurements in solutions of

the same salt at a molality of 1.0 mol kg" . These estimates are consistent with calculated liquid- junction

potentials [16], To eliminate this concentration effect

on the liquid- junction potential, it would be necessary

to have a series of standard solutions of concentrations

spaced over the entire range in which the electrode gives

useful results. For example, if the electrode were standardized in 0.5 M sodium chloride, the error in 1.0 M sodium

chloride would be reduced to about 1 mV (0.017 pCl unit).

(M. Alfenaar)

51
4. EQUILIBRIUM DATA FOR IONIC SYSTEMS

A . Dissociation of Phosphoric Acid in Deuterium Oxide

1 . Measurement of the First Dissociation Constant

As a continuation of the study of acid-base equilibria

in heavy water, it was of interest to determine the dissociation constant of a moderately strong acid in order to

obtain more information on the relationship between the

d-isotope effect and acidic strength. The first dissociation constant of phosphoric acid ( pK-, = 2.148 in H

p

at

25 °C) has been determined previously in heavy water only

at 25 °C [17,18]

.

In view of the difficulties in measuring pK for an

acid as strong as phosphoric, it seemed worthwhile to redetermine the dissociation constant from 5 to 50 °C by

accurate emf measurements of the cell

Pt;D

2 (g), DCl(m), KD

2

P04

(4m), AgCljAg

The buffer mixture had a composition ratio of 3 moles

KD

2

P0^ to 1 mole D P0^ to 1 mole KC1. The molality of

D^PO^ was varied from C.003 to 0.05.

An expression relating the dissociation constant K-, to the emf E of the cell was obtained by combining the

Nernst equation for the emf with the mass law:

(E "E

° )F Vpo, ^pcv^cipK = + log m + log —2 + log I : *

(17)

2.3026 RT 0± D

2

P0JJ

yD

2

P0

4

Since the last term of equation 17 is negligible in dilute

solutions (and varies linearly with ionic strength at higher

concentrations), the approximate pK value (pK' ) is given by

m-nL +

pK* = P(ain Yri ) + log ^—

(18) 1 D 01 3m + m

J)+

32
The acidity function p(cLyp,) is equivalent to the first

two terms on the right of equation 17- A small correction

for the ionization of phosphoric acid had to be made. This

was done by estimating the deuterium ion concentration from

the acidity function p(&n

Y-,-) ):

2AI 1/2 d 1/2

-log m

D+ = p(aDTci ) - ° (19)

1+BaI 17 ^ d

L/

d

o

where A and B are the Debye-Huckel constants for heavy water,

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

d is the density of heavy water.

The value of pK, , that is, the limit of pK^ at 1=0,

was obtained by fitting the data, to an equation linear in

I by the method of least squares. The ion-size parameter

was varied between J5-5 and 7.0 A, and the average value

which gave the best fit at all temperatures was 5-0 A.

A plot of the data at 25 °C is shown in figure 8. The

pK-, values found at each temperature were also fitted by

the method least squares to the following equation

pK^ = A/T + b + CT (20)

where T is in kelvins. The values of the constants are

given in table 7> together with the calculated pK-, values

(molal scale) in heavy water and the corresponding pK-,

values obtained in ordinary water by Bates [19].

33
The values in water were checked by repeating the experimental procedures described above with ordinary water as

solvent

.

2.420 No

i i i t=25°C PV2.420

CT = 0.0005

0*5fi

2.415

2.410 o \ -

>K|

2.405 -

2.400 . O

-

2.395 i 1 1 .05 .10 IONIC STRENGTH

.15 .20

Figure 8. Removal of activity coefficient effects by

extrapolation in the determination of the

dissociation constant of phosphoric acid

in deuterium oxide.

54
Table 7« Dissociation constant of phosphoric acid in

deuterium oxide from 5 to 50 °C.

t°c pK-, in DpO pK., in H^O pK

1

in H

2

C ApKd

5 2.345 2.071 2.065 0.274

10 2.361 2.088 2.084 .273

15 2.379 2.106 2.104 .273

20 2.399 2.127 2.125 .272

25 2.420 2.148 2.148 .272

30 2.443 2.171 2.171 .272

35 2.468 2.196 2.196 .272

40 2.494 2.222 2.222 .272

45 2.521 2.249 2.248 .272

50 2.550 2.278 2.276 .272

a

b

c

d

Calculated from the equation pK-, = 843.979/T

- 4.5714 + O.OI39555T

Reference [19]

This work

(Column 2 ) - (Column 3)

(M. Paabo)

35
2. Thermodynamic Quantities for the Dissociation

The thermodynamic constants for the dissociation of

phosphoric acid in heavy water at 25 °C have been determined from the temperature coefficient of pK-, and are given

in table 8. They can be compared with those for the dissociation of phosphoric acid in ordinary water. From these

values, one can calculate the standard thermodynamic constants for the transfer process from ordinary water to

heavy water, with results discussed in the next section

of this report.

Table 8. Thermodynamic constants for the dissociation of

phosphoric acid in deuterium oxide and water

at 25 °C.

in D

2

a in H

2

b

AH , J mol" 1 -7595 (12) -7650

AS°, J K"

1 mol" 1 -71.8 (o.i) -66.8

AC°, J K" 1 mol" 1

P

-159 (2) -154

a

b

Values in parentheses refer to the uncertainties

in thermodynamic constants as estimated by the

method of variance analysis

Reference [19]

(M. Paabo)

36
3 . Deuterium Isotope Effects

Soon after the discovery of heavy water it was found

that weak acids are weaker in this solvent than in ordinary

water. It was soon apparent that the increase in pK over

that in ordinary water, known as the "deuterium isotope

effect", is not the same for all acids but varies in some

manner with the strength of the acid. The exact relationship between ApK, that is, pK (in D

p 0) - pK (in Hp0), and

the pK of the acid in water is not yet completely clear.

It has long been agreed that the isotope effect for moderately strong acids is smaller than that for very weak

acids, and it has been suggested that the isotope effect

varies linearly with pK (in H

p 0) [20].

Exact tests of the linear relationship are rendered

difficult by the scarcity of precise data for ionization

constants in deuterium oxide. This situation is slowly

being remedied by the appearance of new data, in large

part furnished by careful measurements made in this laboratory. For this reason, Dr. R. A. Robinson, formerly

a member of the Section, was led to re-examine the data

on isotope effects during the past year. The results have

recently been published, together with data for the isotope

effect on two protonated amines, a substituted phenol, and

the mean of pK., and pKpfor a polybasic carboxylic acid

(citric acid) [21]. The cation acids for which pK values

were reported are o-nitroanilinium and m-nitroanilinium

ions; very few acids of this charge type have been studied

in deuterium oxide.

37
A summary of the results for the deuterium isotope

effect (ApK) on the dissociation of 15 acids is shown in

graphical form in figure 9- It is clearly evident that

the data do not fall on a single straight line.

5 10 pK(in H20)

15

Figure 9- Deuterium isotope effects (ApK) for weak

acids at 25 °C.

1 o-Nitroanilinium ion

2 Bisulfate ion

3 Phosphoric acid

4 m-Nitroanilinium ion

5 4-Chloro-2, 6-dinitrophenol

6 Citric acid (pK +pKQ )/l/2 14 Bicarbonate ion

9 Deuterioacetic acid

10 Hydrogen succinate ion

11 ~5, 5-Dinitrophenol

12 Dihydrogen phosphate ion

13 Boric acid

'2

7 Succinic acid

8 Acetic acid

15 Self-dissociation of

solvent

38
It is more difficult, however, to classify the deviations

from linearity in terms of structure and charge type. One

notes that the smallest values of ApK are found for two

mineral acids, sulfuric (2nd stage) and phosphoric acid

(1st stage). It should be noted that the value of ApK

for the latter (point 3) has been redetermined very recently with the deuterium gas electrode in cells without liquid

junction [22] and found to lie very close to the straight

line drawn in the figure (see Section 4.A.1). Likewise,

those acids, numbered 12 through 15, which are the weakest

studied, lie on or close to the straight line; they may

also be classified as inorganic.

Of the acids remaining, 1 and 4 are positively- charged

substituted anilinium ions; 5 and 11 are neutral phenols;

7, 8, and 9 are neutral carboxylic acids; and 6 and 10

involve carboxylic acids with a negative charge. All of

these acids are organic, but both aliphatic and aromatic

types are represented. Quite evidently the new data just

reported (points 1, 4, 5, and 6) alter quite profoundly the

concept of an isotope effect that decreases in a regular

manner as the strength of the acid increases. Taken as a

whole, the data tend to minimize the importance of charge

type as a parameter determining the magnitude of ApK.

They do suggest, however, that organic acids may behave

differently from inorganic acids. Further work may be

necessary in order to demonstrate whether aromaticity plays

a part in the solvent effect.

The pK values of a sufficient number of weak acids

have now been determined over a range of temperature

in both water and deuterium oxide to make worthwhile a

comparison of the enthalpy, entropy, and heat capacity effects. Such a comparison is made in table 9; the acids are

arranged in order of increasing pK (decreasing strength).

39
The figures given in each column represent the excess of

the indicated thermodynamic quantity (AH°, AS% or AC°) in

deuterium oxide over the value in water. The differences

in enthalpy and heat capacity are seen to increase as the

acid becomes weaker, but the isotope effect on the entropy

change appears to be nearly independent of the acidic

strength.

Table 9« Deuterium isotope effect on the thermodynamic

constants for the dissociation of weak acids

at 25 °C

a

Acid AH° (d)-AH°(h) AS°(d)-AS°(h) AC°(d)-AC°(h)

Phosphoric 57

b -5.0 C

- 4

C

Acetic 1561 -5.4 - 8

Acetic-d-2

3

1456 -5.4 -17

Primary phosphate ion 1628 -5.9 -18

Bicarbonate

ion 2564 -5.9 +15

Solvent 3284 -7.3 -34

d = in D

2

0; h = in H" 2

J mol

C

J K"

1 mol" 1

(R. G. Bates)

4o
B. First Dissociation Constant of Piperazinium Ion

in 50 Weight Percent Methanol-Water Solvent

As part of a study of the process by which cationic

weak acids dissociate,, the acidic ionization constants of

t ris(hydroxymethyl)aminomethane [23], ammonia [24], and

4- amino pyridine [25] have been determined in the 50 weight

percent methanol-water solvent.

Work is now in progress on diprotonated piperazine

.

This acid has a double positive charge, and a positively

charged conjugate base is formed in the first dissociation

step.

PzH*4" + SH <=± PzH

+

+ SH^

where Pz is written for the neutral base piperazine and SH

represents the amphiprotic solvent. The two dissociation

constants of piperazinium ion in water were determined

recently in this laboratory [26] thus making it possible to

estimate the medium effect p( K) - p( K). The subscripts

s and w indicate the standard state to which the thermodynamic quantity is referred: 50 percent methanol (s)

or water (w)

.

Emf measurements were made at seven temperatures from

10 to 40 °C for the cell

Pt,H

2 (g),Pz-2HCl(m),Pz.HCl(m) in 50 wt % MeOH-H

2

0, AgCljAg

The concentrations of the two hydrochlorides were kept

equal and the molality of each varied from 0.0025 to

0.05 mol kg . A slight correction was made for the ioniz-

++ ation of PzH?

. The molality of hydrogen ion was estimated

with the aid of the acidity function p (a^Yp-,).

The method of calculating p( K-, ) followed closely that

outlined above in Section 4.A.I. In this instance, however,

the activity coefficient term was, as expected, far from

41
negligible. For the calculation of p( K-,)' it was estiS J- mated by a Debye-Huckel equation having the same form as

the last term in equation 19 (the coefficient 2 in the

numerator being replaced by 4, as required by the charges

of the species participating in the dissociation process).

The value of p( K1 ), that is, the limit of p( K-, )

' at

zero ionic strength, was calculated by the method of least

squares for several values of the ion-size parameter (a)

at 25 °C. The smallest standard deviation for p( K-, ) was

obtained with a=2 A. The resulting medium effects for

piperazinium ion and other weak bases studied are given

in table 10.

It is evident that the initial decrease in pK for

a

positively charged weak acids now generally recognized to

occur when methanol is added to the aqueous solvent, is

accentuated by doubling the charge on the cation acid.

Further work on this base in other solvent compositions

at 25 °C is planned.

Table 10. Values of p( K) for the dissociation of

protonated weak bases in 50 wt % methanolwater at 25 °C and related solvent effects.

Acid

Piperazinium (first step)

Protonated tris

4-Aminopyridinium

Ammonium

P(sK

a) PW A

3

4.583 5.333 -0.750

7.818 8.075 -0.254

8.518 9.114 -0.596

8.687 9.245 -0.558

a A=p(

sKa

) - p( w

K

a

)

(M. Paabo)

42
C. Dissociation of Protonated Tris(hydroxymethyl)-

aminomethane in Aqueous Alcohol Solvents

Acid-base behavior in mixed solvents of which water

is one component has been a subject of interest in the

Section for a number of years. These studies have the

aim of shedding light on ion-solvent interactions and

other medium effects. Investigations of this sort, it is

hoped, may eventually point the way to a single acidity

scale for different mixed solvent systems. The nature of

solvent effects on the equilibrium behavior of electrolytes in mixed solvents was discussed in earlier progress

reports [2,11] and in a lengthy review published during

the past year [27].

Schindler, Robinson, and Bates [28] determined

the dissociation constant of the protonated form of

tris(hydroxymethyl)aminomethane in methanol-water solvents

containing 30 , 50, 70, and 90 wt percent methanol by a

potentiometric titration method. This procedure is relatively simple, utilizing glass electrodes in a cell without

liquid junction, yet capable of yielding results accurate

to better than 0.01 in pK.

We have now extended these measurements to demonstrate the influence of the nature of the alcohol, studying the pK of protonated tris in aqueous solvent mixtures

containing 50 wt percent ethanol, n-propanol, or tert - butanol. All of the measurements were made at 25 °C.

^
The cell

Glass electrode; HC1 (0.01m) + tris in 50 wt %

alcohol-water, AgCljAg

was used. The cell vessel is shown in figure 10. The

initial emf (E-, ) was measured at equilibrium before any

tris had been added.

Figure 10. Cell vessel for determining pK by

potentiometric titration without a

liquid junction.

44
Successive additions of a solution 0.01 m in HC1 and 0.1 m

in tris (m=molal) in the alcohol-water solvent under study

led to emf values designated Ep. In this way, both E-, and

Ep were determined at a constant chloride molality and a constant total ionic strength, while the buffer ratio

(tris/tris .HCl) was changed. A similar procedure, using

the hydrogen gas electrode instead of the glass electrode,

was developed in the Section many years ago [29].

The activity coefficient term in the expression for pK

is of such form as to be negligible in dilute solutions [28];

hence, a close approximation to the pK is obtained from the

equation

(E p-E-,) nv, +

pK =

2 1

+ log -£2 (21)

(RT In 10)/F 0.01mB

where rrLRH+ and rru are respectively the molalities of protonated tris and the free tris base in the solution for

which Ep was measured.

Preliminary results for the pK of protonated tris in

50 wt percent aqueous methanol, ethanol, n-propanol, and

tert-butanol at 25 °C are given in table 11. On the basis

of these results, the pK seems to be lower in ethanol than

in the other alcohols of the series.

Table 11. pK values for tris«H+ in 50 wt % aqueous alcohols.

Alcohol

(50 wt %) PK

Solvent

density

Methanol 7.799 0.9125

Ethanol 7-654 0.9138

n-Propanol 7.780 O.9095

tert-Butanol 7.758 O.8985

45
Further work is under way. In a qualitative manner,

the trend in pK suggests a possible influence of solvent

structure. The differences in pK are not large, and it

is possible that the comparison may have to be made at

equal mole fractions of alcohol, where the numbers of solvent molecules are the same. The pK of tris«H in water

at 25 °C is 8.075 [30]

(B. R. Staples, R. S. Carter, Jr.)

D. Thermodynamics of Hydrochloric Acid in

N-Methylpropionamide Solvents

Electromotive force studies in N-methylpropionamide

(NMP) and in mixtures of NMP with cyclohexane or decalin

have been conducted in the Section by Edgar S. Etz as part

of his graduate program at Clarkson College of Technology.

The experimental work was described in the two preceding

annual reports [2,11]. Measurements have been made from

5 to 55 °C

Mr. Etz returned to the Section for the summer months

of 1968, devoting his time to a detailed computer analysis

of the data. His efforts resulted in l) values of the

standard emf of the hydrogen-silver chloride cell when the

solvent is NMP or a mixture of NMP with one of the two

hydrocarbons studied, 2) the activity coefficient 3 partial

molal heat content, and partial molal heat capacity of

hydrochloric acid in these nonaqueous solvents, and 3) the

changes in Gibbs energy, enthalpy, and entropy on the

transfer of one mole of HC1 from the standard state in

A

7 ater to the standard state in these three nonaqueous

solvent systems.

46
Accurate thermodynamic data for electrolytes in

anhydrous systems have rarely been obtained, and consequently these results are of considerable scientific

interest. Pure NMP has a very high dielectric constant

(175.6 at 25 °C). Interionic attractive forces might

therefore be expected to be small, and indeed the activity

coefficient of hydrochloric acid in this solvent exceeds

O.96 for all concentrations (C) less than 0.1 mole per

liter, as shown in figure 11.

Figure 11. Comparison of the activity coefficient of

hydrochloric acid in NMP with that in water

47
Through the study of NMP-hydro carbon mixtures

"isodielectric" with water it was hoped to compare solvent

basicities. When electrostatic contributions are approximately equalized by this device, it is reasonable to expect

that the energy change in the transfer of hydrochloric acid

from water to a nonaqueous solvent will reflect primarily

the relative magnitudes of the solvation interactions in

the two solvents. The interaction between the proton and

solvent proton acceptors (bases) is likely to be important

here. Because of the large difference between the dielectric constants of the pure solvents water and NMP, such a

comparison of transfer energies is not likely to be very

useful as a measure of relative basicities.

Thus far, it has not been possible to interpret the

results in terms of a difference in intrinsic basicity

of the two solvents water and NMP. This conclusion is

consistent with the growing recognition [31] that "intrinsic basicity" is not a meaningful concept in hydrogenbonded mixed solvent systems.

A portion of this work was presented by Professor

Donald Rosenthal in the Fisher Award Symposium held on

April 17 , 1969, as part of the National Meeting of the

American Chemical Society at Minneapolis, Minnesota.

(R. G. Bates)

48
E. Thermodynamics of Hydrochloric Acid in Mixtures of

2-Methoxyethanol (80 Weight Percent) and Water

(20 Weight Percent)

In order to base standards for pH in 80 wt percent

2-methoxyethanol on emf measurements, it is first necessary to determine the standard emf of the hydrogen- silver

chloride cell for this solvent system. The requisite

measurements, made over a range of temperatures, will yield

not only the standard emf "but also information on the

thermodynamic behavior of hydrochloric acid in this mixed

solvent

.

1. Vapor Pressures of the Solvent

In order to make the hydrogen pressure corrections to

the measured emf values, the vapor pressures of the solvent

at different temperatures have to be known. The apparatus

for the vapor pressure determination is shown in figure 12.

A flask containing the solvent is submerged in a water bath

with temperature control to ±0.01 °C . The flask is connected with a U-tube mercury manometer and the pumping

system. To obtain a good vacuum, three pumps were used in

series; these were a mercury diffusion pump, an oil diffusion pump, and a mechanical pump. Because 2-methoxyethanol

attacks the high vacuum grease, Teflon stopcocks had to be

used.

Before connecting the flask to the vacuum line, the

temperature of the solvent was lowered to -72 °C, so that

the evacuation of the system could not cause any change in

the composition of the liquid. To avoid condensation of

vapors in the manometer, the latter was placed in a controlled air bath with a transparent plastic panel, so that the

manometer could be kept at 5^ °C throughout the measurements

49
The rest of the apparatus, except for the flask, was wrapped with heating tape and kept at a temperature well above

50 °C . Before starting the measurements, the pressure in

the apparatus was reduced below 1 |jm Hg. The manometer was

read with a cathetometer giving a precision of 0.05 mm Hg.

Figure 12. Apparatus for measuring the vapor pressure

of solvents.

50
The measured, values of the vapor pressures of the

binary solvent mixture 2-methoxyethanol (80 wt percent) - water (20 wt percent) In the temperature range 10 to 50 °C

are given in table 12. Least-squares treatment of the

results yielded the following equation:

2,419.46

log p = 11.502 (22)

T

_2 where p is the vapor pressure in Nm and T is the temperature in kelvins (l normal atmosphere = 1.01325xl0J Nm ).

Table 12. Vapor pressures of a mixture of 2-methoxyethanol

(80 wt fo) and water (20 wt %) from 10 to 50 °C

.

_2

t, °C p, Nm p, mm Hg

10 907 6.8

15 1266 9.5

20 1733 13.0

25 2480 18.6

30 3333 25.0

35 4486 33.6

40 5993 44.9

45 7846 58.9

50 10159 76.2

51
2. Dielectric Constants of the Solvent

The dielectric constants of the mixed solvent were

measured at 10, 25, and *50 °C, using freshly distilled

water and spectro-grade acetone as reference materials.

The measurements were carried out with the Multidekameter

instrument. The temperature of the cell was adjusted to

within 0.1 °C of the desired temperature by circulating

water from a constant-temperature bath. The following

values of the dielectric constant have been obtained:

34.7 at 10 °C, 31.5 at 25 °C, and 27.8 at 50 °C.

3. Determination of Acidity Functions

Electromotive force measurements of cells without

liquid junction of the type

Pt;H2

(g)j HCl(m) in 2-methoxyethanol-H2

0, AgCljAg

were made with 12 different molalities (m) of hydrochloric

acid in the solvent 2-methoxyethanol (80 wt percent) - water (20 wt percent) over the temperature range 10 to 50 °C.

The temperature was controlled to 0.01 °C, and the emf

was measured with a precision potentiometer. The observed

potentials were very stable, and the reproducibility of

duplicate cells was nearly always better than 0.1 mV.

These data will serve for a calculation of the standard

emf of the cell for this mixed solvent system.

Values of the acidity function Pw (arrYpi )> referred

to the aqueous standard state, have been derived from

the emf, with the use of the standard emf obtained from

measurements in water solvent [32]. Acidity functions

at 25 °C are given in table 13. The mean activity coefficient of hydrochloric acid ( y+ ), referred to the standard

state in pure water, is also of interest. It was calculated from the acidity functions by the equation

52
-log W

Y ± = [Pw^hW + log m]/2 (23)

and is listed in the last column of the table.

Table 13 . Values of Pw

( aTj

rVni^ and w^+ for hydrochloric

acid in a mixture of 2-methoxyethanol (80 wt %)

and water (20 wt %) at 25 °C.

m x 10-5 emf, mV pwl¥ci) w^±

4.05 401.38 0.634 7.59

6.01 382.32 0.483 7.40

9.31 364.28 0.368 6.78

10.90 356.14 0.299 6.79

17.07 338.17 0.190 6.15

21.83 326.50 0.100 6.04

32.38 311.58 0.019 5.45

42.13 299 . 18 -0.077 5.32

58.58 287 . 19 -0.136 4.83

78.55 273.70 -0.237 4.69

94.24 267.77 -0.258 4.39

106.20 261.16 -0.318 4.43

The magnitude of the activity coefficient (4 to 8)

demonstrates that hydrochloric acid has a higher free

energy or "escaping tendency" in the mixed solvent

than in water. This result is consistent with data for

53
other mixed solvents such as dioxane-water [33]. It may

be attributed largely to the lowered dielectric constant

produced by the addition of 2-methoxyethanol to water, with

the consequent increase in the energy of charged species

in the medium. It is likely that a contrary influence,

obscured by this larger electrostatic effect, is exerted

by ion pairing, some of which is to be expected in a solvent

of dielectric constant near J>0

.

(H. P. Thun)

54
5. OCEANOGRAPHIC CONDUCTANCE STANDARDS

A. Apparatus

The program of conductivity research has been directed toward the attainment of the highest accuracy in the

measurement of the conductance of aqueous salt solutions.

The direct objective is the certification of sodium chloride

solutions suitable for the calibration of oceanographic

salinometers . The needs for that project were outlined in

the annual progress report for 1968 [11].

Since last year the major modification of equipment

was the incorporation of the Janz-Mclntyre bridge in a

single cabinet, shown in figure 1^. At the same time, the

following changes were made to the bridge:

(1) The output from the impedance unbalance meter

was brought to an external d-c amplifier, permitting interpolation of the lowest decade of the resistance arm to

0.01 ohm.

(2) Shields of the cell leads were connected to the

reversing switch so that the shields of the low-potential

leads are in the guard circuit of the comparator, while

the shields of the high-potential leads are grounded.

(3) Provision was made for connecting a fixed 10,000-

ohm coil in parallel with the unknown in order to extend

the range of measurements to higher impedances. In this

configuration, the four-terminal method of measurement

does not strictly eliminate the effect of lead impedance,

but the small correction can be calculated readily.

55
Figure lj5. Janz-Mclntyre bridge for conductivity

measurements.

Five cells, having constants of approximately 10, 40,

80, 120, and 160, were constructed. Three of these are

shown in figure 14. The others are similar, differing

only in the length and diameter of the capillary section.

All have 2.5 cm diameter bright platinum electrodes and

four leads terminating in shielded connectors.

(T. B. Hoover)

56
Figure 14. Two-electrode conductivity cells.

B. Calibration of Cells

Three of the cells were calibrated at 18 and 25 °C

with 0.1 demal potassium chloride solutions prepared with

salt from two different sources. Statistical analysis of

the results showed a standard deviation of the cell constants of 100 ppm. About half of this is accounted for by

a significant trend of the deviations in the direction of

higher constants at the end of the week of measurements.

57
Another series of calibrations of all five cells with

0.01, 0.1, and 1 denial solutions from a single supply of

potassium chloride has been completed experimentally, also

at 18 and 25 °C, but the results have not been fully

analyzed

.

In addition to the calibration corrections for the

resistances of the bridge coils, two other significant

corrections need to be applied to the raw data. These are

for variations of temperature and for frequency effects.

Temperatures of the oil thermostat, which is shown

in figure 15, were measured with two platinum resistance

thermometers. In 28 measurements, the maximum difference

between the readings of the two thermometers was 3.6xl0~^K,

while the mean difference and standard deviation were

-4 r -4

2x10 K and 16x10 K, respectively. A continuous record

of the temperature for several hours indicated cyclic

fluctuations of ± 0.001 K with a period of 2-3 minutes.

The heat capacity of the cell probably smooths out some of

this variation, but the absolute temperature must be considered uncertain to at least 0.002 K. This uncertainty

is equivalent to about 40 ppm in the conductance of the

solution.

Frequency effects resulting from polarization of the

electrodes are an inherent source of error in electrodetype conductivity cells. It was considered desirable to

avoid platinization of the electrodes because of the uncertainty associated with adsorption effects on platinum

black. Consequently, other experimental means of reducing

the effects of polarization were investigated. These

included: (l) differential measurement between cells with

different cell constants in the reference and unknown arms

of the bridge [j54], (2) constant impedance, in which the

reference arm was a fixed resistor and the variable resist58
Figure 15 . Oil thermostat for conductivity measurements

ance was in series with the cell, and (3) four-electrode

cells, such as that shown in figure 16. In the latter

method the electrodes are part of the lead impedances,

and polarization effects were expected to cancel by the

four-terminal method of measurement. The cylindrical

potential electrodes shown in the photograph were placed

in a region of minimum potential gradient.

In each of these procedures very large frequency

effects were observed, which probably were a consequence

of introducing an impedance in series with the cell.

One current lead then was neither at the detector potential nor at the output potential of the bridge.

59
Figure 16. Four-electrode conductivity cell.

Admittance from this lead to its shield would act as a

frequency-dependent shunt across other portions of the

circuit. When the impedance "between current and potential

electrodes of the four-electrode cells was made small, as

shown, the effect was reduced, but not to the le/el found

for the conventional two-electrode cells.

The mathematical correction for polarization effects

by extrapolation is strictly empirical, although Feates,

Ives, and Pryor [35] solved the balance conditions for a

simple model of electrode polarization. The Janz-Mclntyre

bridge is not very suitable for investigating frequency

effects because it provides only four output frequencies.

60
Nevertheless, a computer program to fit the observations

to the Feates-Ives-Pryor model showed that the highfrequency limit of extrapolation was always very close

to the value measured at 10 kHz. On the other hand,

measurement of some of the same cells with the unshielded

Jones-Dike bridge at nine or ten frequencies showed a

precisely linear relation between resistance and reciprocal

frequency between 1 and 20 kHz, provided the resistance

was less than 1500 ohms.

As a further check of results obtained with the JanzMclntyre bridge, the same solution was measured in a cell

before and after the electrodes had been coated with

palladium black, and again after stripping off the palladium electrolytically. As expected, the calculated

double-layer capacity and leakage resistance varied considerably with electrode treatment. The high-frequency

resistance of the electrolyte also varied significantly

according to the Feates-Ives-Pryor calculation but was

quite consistent according to the simpler inverse frequency

extrapolation. It was concluded that the theoretical

model [35] is not applicable to the systems under investigation.

(T. B. Hoover)

61
C . Conductance of Natural and Synthetic Sea Water

Specific conductances of samples of synthetic sea

water used for acidity measurements (Section 2.C) were

measured at 15 and 25 °C. In general, the conductances

were about 7 percent lower than the values corresponding

to natural sea water of the same chlorinity [36]. Two

ampoules of Standard Sea Water, Lot P50, were also measured at 15 and 25 °C in each of two cells. These results

lay within a range of 50 ppm at each temperature and were

in close agreement with the published values of Reeburgh

[36]. Twenty-two samples of natural sea water varying in

salinity from 0.5 to 40^o have been measured at 20 °C for

use by the Naval Oceanographic Instrumentation Center in

calibrating salinometers

.

(T. B. Hoover)

62
6. PREPARATION AND PROPERTIES OF SOLVENTS

A. Dimethylformamide

Dimethylformamide and water appeared to form a suitable

solvent system for the study of the stability of complexes

dependent on acidity and dielectric constant. Dimethylformamide boils at 153 °C, has a dielectric constant of 36.

7

at 25 °C, and dissolves readily most organic as well as many

inorganic compounds. It can be easily purified, but it has

the disadvantage of high hygroscopicity and is subject to

hydrolysis catalyzed by acids and bases. Dimethylformamide

of the highest grade commercially obtainable was stated to

contain less than 0.01 percent of water. The solvent was

distilled to remove possible organic impurities. After the

bottles had been opened several times during normal use,

the water content was found to be 0.015 percent by Karl

Fischer titration.

Unfortunately, dimethylformamide had to be abandoned

after preliminary emf measurements were made. The emf was

unstable, the marked drift presumably being caused by

hydrolysis and the formation of complexes with silver chloride [37,38].

(H. P. Thun)

B. 2-Methoxyethanol

Reagent-grade 2-methoxyethanol (methylcellosolve) was

acquired from commercial sources. It was stated to contain

less than 0.03 percent water. Gas chromatographic analysis,

performed by D. Enagonio of the Separation and Purification

Section, confirmed this statement; no other impurities

could be detected.

(H. P. Thun)

63
C . Sulfolane

A new dipolar aprotic solvent, tetramethylene sulfone

(sulfolane), with a moderately high dielectric constant

(43 at 30 °C), has attracted recent attention. It has

been used as a solvent for potentiometric titrations, for

kinetic and spectrophotometric measurements, and for

polarography. As a continuation of our program of development and application of acidity scales in nonaqueous media,

it was of interest to use this new solvent in emf studies.

Commercial sulfolane was purified by vacuum distillation from sodium hydroxide pellets [39] through a 1-meter

column packed with fine glass rings. This particular still

has been used previously in the Section for the purification of various organic solvents. It is shown in figure 17

The original sample had a conductivity of 1.7x10 ft"

cm . A gas chromatographic analysis performed by Dr. R. T.

Leslie of the Separation and Purification Section indicated

that the sample contained about 0.5 percent water and about

1.8 percent of an unidentified compound.

The first distillation at 4 mm Hg reduced the conductivity to 3 -7-7 .8x10 ' ft cm . The product of a

second distillation at 4 mm Hg and 126 °C had a conductivity of 6.1x10 ft cm , and its appearance was clear

and colorless as compared to the original amber sample.

Emf measurements at 25 °C were made of the cell

Pt;H

2 (g), HC1 in sulfolane, AgCl;Ag

The cell solution was prepared by passing anhydrous

hydrogen chloride through sulfolane. The color of the

solution soon changed to pale yellow, and in a few days

it resembled the amber shade of the original sample.

64
Figure 17 . Still for the purification of organic solvents,

The molality of hydrogen chloride was about 0.0175 as

determined by titration with aqueous sodium hydroxide solution using bromothymol blue indicator to detect the end

point. The emf of the cell never became constant; it increased from -0.02 to +0.28V in the course of 2-1/2 days.

The concentration of hydrochloric acid was the same at

the end of the run as at the beginning.

65
A second cell containing 0.026 molal HC1 in a mixture

of 50 moles percent sulfolane and water was studied. No

coloration of the aqueous solution of sulfolane was observed. The emf of the cell came to an equilibrium value of

about 0.255 V, but the emf readings were very sensitive to

the rate of hydrogen bubbling. An increase in the rate

reduced the emf by several millivolts, and a decrease of

the rate increased the emf. At a constant rate of flow

the emf was stable to about 0.5 mV for a period of 12 hours

The cause of the instability was not determined. Coetzee

and Bertozzi [40] found the hydrogen electrode to function

normally in sulfolane. It is likely that the source of

instability is the silver-silver chloride electrode.

(M. Paabo)

66
7. STANDARD REFERENCE MATERIALS ACTIVITY

A. Summary

Material SRM No. Determined Analyst

Sodium Bicarbonate (191)

Sodium Carbonate (192)

Potassium Dihydrogen Phosphate (2l86l)

Disodium Hydrogen Phosphate (2l86ll)

Sodium Bicarbonate (2191)

Sodium Carbonate (2192)

pH BRS

pH BRS

pD MP

pD MP

pD MP

pD MP

Ion-selective electrodes

(in process)

Sea Water Conductivity

(in process)

Bilirubin

(in process)

pH in Sea Water

(in process)

(3024767) PA BRS, MA

(3024759) Conductivity TBH

(3024761) Assay

(3022753) pH

TBH

ESE

Analyst

BRS - B. R. Staples

MP - M. Paabo

MA - M. Alfenaar

TBH - T. B. Hoover

ESE - E. S. Etz

67
B. Standard Materials for Acidity Measurements

With the completion of the analysis and certification

of potassium dihydrogen phosphate, disodium hydrogen

phosphate, sodium bicarbonate, and sodium carbonate [11],

these materials were added to the list of NBS Standard

Reference Materials for the measurement of pH and pD.

Potassium dihydrogen citrate promises to be a most useful

reference material for acidity measurements both in HpO

and DpO, and it is regrettable that no commercial source

of this salt in pure form has yet been found. In anticipation of the issuance of this material, however, SRM numbers

have already been assigned. Table 14 is a summary of the

Standard Reference Materials for acidity measurements;

although the potassium dihydrogen citrate is not yet available, it is listed in the table.

Table 14. Standard reference materials for acidity

measurements

SRM No

l85d

1861 c 18611b

187a

188

189

190

191

192

pH Standards

Acid potassium phthalate

Potassium dihydrogen phosphate

Disodium hydrogen phosphate

Borax

Potassium hydrogen tartrate

Potassium tetroxalate

Potassium dihydrogen citrate

Sodium bicarbonate

Sodium carbonate

SRM No

21861

2186II

2190

2191

2192

pD Standards

Potassium dihydrogen phosphate

Disodium hydrogen phosphate

Potassium dihydrogen citrate

Sodium bicarbonate

Sodium carbonate

(R. G. Bates)

68
C . Standard Materials for Conductivity Measurements

The oceanographic conductance standards discussed In

Section 5 should find application as well in the calibration of conventional conductivity cells used in research.

Although it is fundamentally desirable that such standards

should be determined absolutely, e.g., in terms of the

physical dimensions of a suitable cell and the absolute

ohm, the difficulties of such a determination dictate that

all sources of error should be systematically investigated.

Several efforts in this direction have been described above.

On the other hand, the general adoption of the Jones and

Bradshaw [4l] aqueous potassium chloride conductance standards for research on electrolyte solutions emphasizes the

practical advantage of retaining these chemical standards

even though they may be shown to be in error. We have

adopted as a provisional reference the specific conductance

datum of Jones and Bradshaw for 0.1 demal potassium chloride

at 25 °C. As part of the development of accurate conductivity techniques, inconsistencies among the other eight

standards of Jones and Bradshaw will be sought.

The adoption of a reference material (potassium chloride) for conductivity permits some specifications for purity

to be set up. The mobilities of most inorganic ions are sufficiently similar that ionic impurities at the ppm level are

not a serious source of error. Exceptions are the hydrogen

and hydroxyl ions which have high mobilities in water. It

can be shown that if the pH of 5-5 M potassium chloride or

of 5.0 M sodium chloride is within 2 units of neutrality

the contribution to the specific conductance is less than

10 ppm. This requirement is not particularly stringent but

cannot be ignored. Nonelectrolyte impurities interfere only

as a weighing error and are not serious at the ppm level.

69
Water, however, may be present at significant levels, especially in hygroscopic salts.

A comparison of 0.1 demal solutions was made between

a commercial "ultrapure" sample of potassium chloride and

a laboratory preparation that had been purified by the

procedure of Pinching and Bates [42] and in addition had

been treated by ion exchange to reduce other alkali metals

to a negligible concentration. The conductance of the

solution of the commercial sample was 0.07 percent lower

than that prepared from the fused laboratory preparation of

potassium chloride. When the commercial material was dried

for three hours at 300 °C it lost 0.045 percent of its

weight. The conductance of a 0.1 demal solution of the

dried material was again 0.035 percent lower than that of

the reference solution. Coulometric assay of the solutions

showed that the former was also 0.0^6 percent lower in

chloride. The discrepancy is believed to be due largely

to mother liquor entrained in the crystals of the commercial specimen.

(T. B. Hoover)

D. Acidimetric Assay of Bilirubin

1. General Considerations

Effort has been directed toward developing a stoichiometric assay procedure for bilirubin, which is urgently

needed as a biological standard reference material. Bilirubin, which has a molecular weight of 584.7, has two

carboxyl groups, suggesting that these functional groups

might be titratable in a suitable system. Bilirubin was

known to be very insoluble in water but quite readily

soluble in dimethylsulfoxide (DMS0) and in strong alkali.

70

-
In alkaline aqueous solution it is readily degraded byoxygen and light. Consequently, titrations were carried

out in closed vessels, under argon atmosphere, and shielded

from direct light. Reagents were added through a septum by

means of a hypodermic syringe. Efforts were made to avoid

a pH greater than 10 and to keep exposure to stronger alkali as brief as possible consistent with complete solution

of the sample.

2

Conductance Measurements

Exploratory measurements were made in dimethylformamide

— T — 1 — 1 (solvent conductance 4x10 ' « cm ) and in DMSO (solvent

_o _-] _ -i conductance 4x10 ' ^ cm ) . In both solvents bilirubin

appeared to be a very weak electrolyte and the salt form

differed little in equivalent conductance from the titrant

(tetrabutylammonium hydroxide). Consequently, the end

points of the titrations were barely detectable.

In the aqueous system, bilirubin was dissolved in an

excess of 0.05 M sodium hydroxide and back-titrated with

0.05 M hydrochloric acid. Significant differences in the

slopes of the titration curve were observed for the three

major regions of the titration (free alkali, titration of

bilirubin, and excess acid), but the transitions were highly

rounded, preventing precise location of the end points.

The large excess of strong electrolyte needed to dissolve

the sample initially was a serious drawback in the conductometric technique.

3. Emf Titrations

One object of potentiometric titrations in water was

to confirm published observations on the solubility, ionization, and pH at half neutralization of bilirubin [4^].

An alkaline solution of bilirubin was back-titrated with

0.05 M hydrochloric acid while the pH was measured with a

71
combination glass-calomel electrode that had been standardized in phosphate buffer at pH 6.86. In confirmation of

the above reference, the curves showed a single, sharp

break at pH 8.3, when the acid form of bilirubin began to

precipitate; the half -neutralization pH at 22.5 °C was 8.20.

No separation of the two carboxyl groups was detectable.

In the titration range in which bilirubin was precipitating,

equilibrium was reached very slowly. After an addition of

acid, the pH immediately decreased and required as long as

half an hour to return to the equilibrium value, which was

reproducible to ± 0.05 pH unit in different runs. In contrast to the published report [4j5], back titration with

standard alkali did not retrace the equilibrium curve but

always showed hysteresis in the direction of higher pH.

The freshly precipitated bilirubin, however, redissolved

more readily than the original sample. These results indicate that bilirubin, because of Its extreme insolubility

at pH less than 8, precipitates in a colloidal form that

traps some of the ionic or salt form and only very slowly

approaches a stoichiometric composition. Consequently,

an indirect approach was adopted of dissolving the sample

In a minimum of alkali and titrating only the excess base.

The most satisfactory procedure was to weigh the

sample into the clean, dry cell under argon and add DMSO

equivalent to 5 to 10 percent of the final volume. This

was not enough to dissolve the sample completely but it

wet and dispersed the bilirubin. Next, 0.05 M sodium

hydroxide was added, equivalent to not more than 100 percent

excess, and the solution was immediately diluted with sufficient water to cover -the pH electrode. The standardized

electrode was inserted and the cell contents were titrated

by gravimetric additions of 0.02 M hydrochloric acid to pH 8

or until a definite break in the pH curve was reached.

72
The cell used for 2-3 mg samples of bilirubin is shown

in figure lo . A similar cell having a capacity of 25 cm

was used for larger samples. In use, the cells were wrapped in aluminum foil as a light shield. The results of

some titrations are summarized in table 15. Runs 10, 11,

and 12 show that repetitive titrations of the same sample

were reproducible to about 0.5 percent, even over a 24-hour

period, indicating that instability of the bilirubin was

not serious. The agreement between different samples is

much poorer, and the causes of this lack of reproducibility

are still under investigation.

Figure 18 Cell for the potentiometric titration

of bilirubin.

(T. B. Hoover)

73
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74
E. A Standard for Ion-Selective Electrodes

It is anticipated that a Standard Reference Material,

sodium chloride, will be issued in the near future, together

with a provisional certification of values for the activities of sodium and chloride Ions. When this first standard

is made available, a detailed publication will be prepared

describing methods for the standardization and use of ionselective electrodes.

(B. R. Staples)

75
8. FOREIGN- CURRENCY PROGRAMS

The two projects sponsored by the Electrochemical

Analysis Section at universities in India and financed

through the provisions of Public Law 480 continued active

during the reporting year. A brief description of these

programs follows.

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

Dr. K. K. Kundu

(a) Standard potentials of the Ag-AgI electrode in

ethylene glycol-water mixtures.

(b) Dissociation constants of tris(hydroxymethyl)-

aminomethane in ethylene glycol.

(c) Studies in isodielectric media.

1. Standard potentials of the Ag-AgBr electrode

in propylene glycol-methanol solvent systems.

2. Autoprotolysis constants of propylene glycolmethanol solvent systems.

5. Standard potentials of the Ag-AgI electrode

in propylene glycol-methanol solvent systems.

(d) Standard potentials of the Na-Na+ electrode in

ethylene glycol-water systems.

Physicochemical Studies in Nonaqueous Solvents , Panjab University, Chandigarh, Professor R. C. Paul

Recent work is described in the following publication

and manuscript:

Paul, R.C., Singla, J. P., Narula, S.P.,

Transference Numbers and Ionic Solvation of Lithium

Chloride in Dimethylformamide,

J. Phys. Chem., £3, 7^1 (1969).

Paul, R.C., Singla, J. P., Narula, S.P.

Transference Numbers and Ionic Solvation of Some

Alkali Salts in Dimethylformamide (in manuscript form),

76
9. IMPACT OF SECTION'S ACTIVITIES

The impact of the Section's activities on the technical and scientific public can best be judged in terms

of the distribution of Section publications, number of

invited talks, and requests for advice and consultation

In this connection, the following statistics, compiled

for the calendar year 1968 unless otherwise indicated,

are pertinent

.

Reprints requested 68l

Manuscripts reviewed for

outside journals 22

Project proposals reviewed 2

Significant consultations 27

federal and state

agencies 8

industries 13

universities 6

Committee memberships 16

Publications (7/68 to 6/69) 11

Talks (7/68 to 6/69) 21

scientific societies 12

universities 3

industries 2

government agencies 4

77
10. PERSONNEL

Roger G. Bates, Section Chief

Marguerite Raudenbush, Section SecretaryConductance Measurements T. B. Hoover

Electromotive Force and

Acidity Measurements

pH, Ion-selective

electrodes

Heavy water,

mixed solvents

Nonaqueous solvents

Consultant

B. R. Staples

M. Alfenaar (July to Sept.)

from University of

Utrecht, Holland

J. T. Clark, Jr.

(July to Sept .

)

guest worker from

Clarkson College of

Technology

R. S. Carter, Jr.

(June to Aug.

)

guest worker from

Landon School

M. Paabo

H. P. Thun (Oct. to July)

guest worker from

University of Ghent,

Belgium

E. S. Etz (July to Sept.)

from Clarkson College

of Technology

D. Rosenthal, Clarkson

College of Technology

78
11. COMMITTEE MEMBERSHIPS

R. G. Bates

Member, Subcommittee on Hydrogen Ion Determinations,

ASTM Committee E-l, Sub. 22.

Councilor, American Chemical Society,

representing the Division of Analytical Chemistry.

Member, ACS Council Committee on Publications.

Chairman, Undergraduate Awards Committee

ACS Division of Analytical Chemistry.

Member, NRC Committee on Symbols, Units, and

Terminology for the Physical Sciences

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

Member, Commission on Electrochemistry, IUPAC

Member, Commission on Symbols, Terminology, and

Units, IUPAC.

Associate Member, Commission on Electroanalytical

Chemistry, IUPAC.

Chairman, Subcommittee on Physico chemical Symbols and

Terminology, Committee on Physical Chemistry, NRC.

Member, Committee Yl, Abbreviations for Science and

Engineering, U.S.A. Standards Institute.

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

Standards Institute.

Member, Board of Editorial Advisers,

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

(Chemistry Panel).

Member, Museum Committee (NBS).

T. B. Hoover

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

B. R. Staples

Member, Ad Hoc Division Committee on Electronic Data

Systems, Analytical Chemistry Division, NBS.

79
12. PUBLICATIONS AND MANUSCRIPTS, JULY 1968 TO JUNE 1969

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

Electrochemical Analysis Section: Summary of

Activities, July 1967 to June 1968,

NBS Technical Note 453 (July 1968).

2. Davis, M.M.

Acid-Base Behavior in Aprotic Organic Solvents,

NBS Monograph 105 (August 1968).

3. Bates, R.G.

Equilibrium Properties of Acids and Bases in

Amphiprotic Mixed Solvents, in Hydrogen-Bonded

Solvent Systems , A. K. Covington and P. Jones, Eds.),

Taylor and Francis Ltd., London, 1968, pp. 49-86.

4. Pool, K.H., Bates, R.G.

Thermodynamics of Hydrochloric Acid in 95 Volume

Per Cent (92.41 Mass Per Cent) Ethanol from

E.M.F. Measurements at 5 to 50 °C,

J. Chem. Thermodynamics, 1_, 21-29 ( 1969 )

.

5. Paabo, M., Bates, R.G.

Standards for a Practical Scale of pD in Heavy Water,

Anal. Chem., 4l, 283 (1969).

6. Staples, B.R., Bates, R.G.

Two New Standards for the pH Scale,

J. Res. Nat. Bur. Stand., 73A, 37 (1969).

7. Hoover, T.B.

The N-Methylpropionamide-Water System. Densities

and Dielectric Constants at 20-40°,

J. Phys. Chem., 73, 57 (1969).

8. Staples, B.R., Atkinson, G.

Structure and Electrolyte Properties in Bolaform

Electrolytes. III. The Hydrodynamics of Potassium

Salts of Several Rigid Bolaform Disulfonic Acids in

Dioxane-Water Mixtures at 25°,

J. Phys. Chem., 73, 520 (1969).

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

Intracellular pH,"

Physiol. Revs., 49, 285 (1969).

80
10. Robinson, R.A., Paabo, M., Bates, R.G.

Deuterium Isotope Effect on the Dissociation

of Weak Acids in Water and Deuterium Oxide,

J. Res. Nat. Bur. Stand., JJA, 289 ( 1969 )

11. Bates, R.G.

Inner Reference Electrodes and Their Characteristics,

in Glass Microelectrodes (M. Lavallee, 0. P. Schanne,

and N. C. Hebert, Eds . )_, John Wiley and Sons,

New York, 1969, pp. 1-24.

12. Bates, R.G.

Medium Effects and pH in Nonaqueous Solvents, in

Solute-Solvent Interactions (C. D. Ritchie and

J. P. Coetzee, Eds. ) Marcel Dekker, New York, 1969 (in press).

15. Bates, R.G.

Hydrogen Ions

Encyclopaedia Brittanica (in press).

14. Bates, R.G., Alfenaar, M.

Activity Standards for Ion-Selective Electrodes, in

Ion-Selective Electrodes (R. A. Durst, Ed.)

NBS Special Publication 314, 1969 (in press).

15. Paabo, M., Bates, R.G.

Dissociation of Deuteriocarbonate Ion in

Deuterium Oxide from 5 to 50 °C,

J. Phys . Chem., 73 (in press).

16. Bates, R.G.

Measurement of Reversible Electrode Potentials, in

Techniques in Electrochemistry

(E. Yeager and A. Salkind, Eds.

)

Wiley-Interscience, New York (in press).

17. Staples, B.R., Turner, D.J., Atkinson, G.

An Apparatus for the Determination of Rates of

Proton Transfer and Other "Very Past Reactions,

Chem. Instrumentation (in press).

81
1J>. TALKS, JULY 1968 TO JUNE I969

1. Bates, R.G., fr Acid-Base Interactions in HydrogenBonded Mixed Solvents", Chemistry Seminar,

University of Florida, Gainesville, Florida,

July 15, 1968.

2. Bates, R.G., "Problems in Monitoring Ion Concentrations

in Dilute Aqueous Systems", Environmental Engineering

Seminar, University of Florida, Gainesville, Florida,

July 15, 1968.

3. Hoover, T.B., "Dielectric Friction and Entropy of

Conductance of Ions in Solution", 156th National ACS

Meeting, Atlantic City, N.J., September 9, 1968

.

4. Paabo, M., "Thermodynamics of the Dissociation of

Deuterocarbonate Ion in Heavy Water from 5 to 50 °C",

156th National ACS Meeting, Atlantic City, N.J.,

September 9, 1968.

5. Alfenaar, M., "Ionic Activity Scales and Standards

for Ion-Specific Electrodes"', 156th National ACS

Meeting, Atlantic City, N.J., September 12, 1968

.

6. Bates, R.G., "Practical Approaches to Acidity

Measurement in Amphiprotic Solvents",

Physical Chemistry Seminar, Tyco Laboratories, Inc.,

Waltham, Mass., October 30, 1968

.

7. Bates, R.G., "Measurement of Acidity in Alcohol-Water

Mixtures", Connecticut Valley ACS Section

University of Connecticut, Storrs, Conn.,

November 4, 1968

.

8. Bates, R.G., "Solvent Effects on Protolytic Equilibria

in Hydrogen-Bonded Mixed Solvents", Western Vermont

ACS Section, St. Michael's College, Winooski, Vt .

,

November 6, 1968

.

9. Bates, R.G., "Solvent Effects on Protolytic Equilibria

in Hydrogen-Bonded Mixed Solvents",

Central Massachusetts ACS Section,

Assumption College, Worcester, Mass.,

November 7, 1968.

10. Bates, R.G., "Solvent Effects on Protolytic Equilibria

in Hydrogen-Bonded Mixed Solvents",

Maine ACS Section, Portland, Maine,

November 8, 1968

.

82
11. Bates, R.G., "Effect of the Solvent on Acid-Base

Reactions In Solution", Chemistry Visitors Day,

University of Kentucky, Lexington, Kentucky

November 22, 1968

.

12. Bates, R.G., "Effect of Solvent Properties on the

Strengths of Carboxylic and Cationic Acids in

Methanol-Water Solvents", National Institutes of

Health — National Bureau of Standards Joint

Colloquium, NIH, Bethesda, Maryland,

January 7, 1969

.

13. Bates, R.G., "Requirements for pH Analysis"

Course on Advances in Instrumental Analysis,

Walter Reed Army Institute of Research,

Washington, D.C., January 16, 1969

.

14. Bates, R.G., "Activity Standards for Ion-Selective

Electrodes", Symposium on Ion-Selective Electrodes,

NBS, Gaithersburg, Maryland, January J>0, 1969.

15. Bates, R.G., "Measurement of Acidity in Alcohol-Water

Solvents", Northern New York ACS Section,

Clarkson College of Technology, Potsdam, N.Y.,

February 6, 1969

.

16. Staples, B.R., "Activity Scales for Ion-Selective

Electrodes", Middle Atlantic Regional Meeting of ACS,

Washington, D.C., February 15, 1969

.

17. Bates, R.G., "Solvent Effects on Acid-Base Equilibria

in Hydrogen-Bonded Mixed Solvents", Analytical Group

North Jersey ACS Section, Seton Hall University,

South Orange, N.J., February 24, 1969

.

18. Paabo, M., "Deuterium Isotope Effect on the First

Dissociation of Phosphoric Acid and Related

Thermodynamic Quantities", 157th National ACS Meeting,

Minneapolis, Minnesota, April 17, 1969

.

19- Bates, R.G., "Solute-Solvent Interactions and

Acid-Base Behavior", Fisher Award Address,

157th National ACS Meeting, Minneapolis, Minnesota,

April 17, 1969.

83
20. Staples, B.R., "Ion-Selective Electrodes and Activity

Scales", Central Electricity Research Laboratories,

Analytical Chemistry Section, Leatherhead, Surrey,

England, April 30, 1969.

21. Bates, R.G., "pH", Symposium on an International

Standard Reference Materials Program,

National Bureau of Standards, Gaithersburg, Maryland

May 22, 1969.

84
14. LIST OF REFERENCES

Electrochemical Analysis; Studies of Acids, Bases , and Salts by Emf , Conductance, Optical, and Kinetic

Methods. July 1965 to June 1966"

R. G. Bates, Ed.), NBS Technical Note 400

September 6, i960).

Electrochemical Analysis Section: Summary of

Activities, July 1966 to June 1967 . R. G. Bates, Ed.), NBS Technical Note 423

October 1967)

•

British Standard 1647: Specification for pH Scale , British Standards Institution, London (1961).

Staples, B.R., Bates, R.G.

J. Res. Nat. Bur. Stand., 73A, 37 (1969)

.

Paabo, M., Bates, R.G.

Anal. Chem., 41, 283 (1969) .

Lewis, J.C.

Anal. Biochem., 14, 495 (1966).

Bates, R.G., Guggenheim, E.A.

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

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

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

Simon, W., Lyssy, G.H., Mbrikofer, A., Heilbronner, E

Zusammenstellung von Scheinbaren Dissoziations - konstanten im Lbsungsmittelsystem Methylcellosolve/

Wasser \ Juris-Verlag Zurich. Band I, 1959.

Sommer, P.F., Simon, W.

Ibid . Band II, 1961. Band III, 1963.

Electrochemical Analysis Section: Summary of

Activities, July 1967 to June 1958"! R. G. Bates, Ed.), NBS Technical Note 453

'July 1968).

Bates, R.G., Alfenaar, M.

Chapter 6 in Ion-Selective Electrodes . (R. A. Durst, Ed.) NBS Special Publication 314,

1969.

85
13] Shedlovsky, T. Madnnes, D.A.

J. Amer. Chem. Soc, 59, 503 (1937).

14] Robinson, R.A., Stokes, R.H.

Electrolyte Solutions , 2nd ed . , appendix 8.10

Butterworths, London, 1959.

15] Shatkay, A.

Anal. Chem., 39, 1056 ( 1967 )

.

16] Shatkay, A., Lerman, A.

Anal. Chem., 4l, 514 (1969).

17] McDougall, A.O., Long, F.A.

J. Phys. Chem., 66, 429 (1962).

18] Salomaa, P., Schaleger, L.L., Long, F.A.

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

19] Bates, R.G.

J. Res. Nat. Bur. Stand., 47, 127 (1951).

20] Bell, R.P.

The Proton in Chemistry , Chapter 11

Cornell University Press, Ithaca, N.Y., 1959.

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

J. Res. Nat. Bur. Stand., 73A, 289 ( 1969 )

.

22] Paabo, M., Bates, R.G.

Unpublished work.

23] Woodhead, M., Paabo, M., Robinson, R.A., Bates, R.G.

J. Res. Nat. Bur. Stand., 69A , 263 (1965).

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

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

25] Paabo, M., Robinson, R.A., Bates, R.G.

Anal. Chem., 38, 1573 (1966).

26] Hetzer, H.B., Robinson, R.A., Bates, R.G.

J. Phys. Chem., 72, 208l (1968).

27] Bates, R.G.

In Hydrogen-Bonded Solvent Systems (A. K. Covington

and P. Jones, Eds.), p. 49.

Taylor and Francis, London, 1968.

86
Schindler, P., Robinson, R.A., Bates, R.G.

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

Bates, R.G., Siegel, G.L., Acree, S.F.

J. Res. Nat. Bur. Stand., 30, 347 (1943).

Bates, R.G., Hetzer, H.B.

J. Phys. Chem., 65, 667 (1961).

Pranks, F., Ives, D.J.G.

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

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

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

Bates, R.G.

Determination of pH, Theory and Practice , Chapter 7

John Wiley and Sons, New York, 1964.

Ives, D.J.G. , Pryor, J.H.

J. Chem. Soc, 2104 (1955).

Feates, F.S., Ives, D.J.G., Pryor, J.H.

J. Electrochem. Soc, 103, 580 (1956).

Reeburgh, W.S.

J. Marine Res., 23, 187 (1965).

Mandel, M., Decroly, P.

Nature, 182, 794 (1958).

Butler, J.N.

J. Phys. Chem., 72, 3288 (1968).

Garnsey, R., Prue, J.E.

Trans. Faraday Soc, 64, 1206 (1968).

Coetzee, J.F., Bertozzi, R.J.

Anal. Chem., 4l, 860 (1969).

Jones, G., Bradshaw, B.C.

J. Amer. Chem. Soc, 55, 1780 (1933).

Pinching, G.D., Bates, R.G.

J. Res. Nat. Bur. Stand., 37, 311 (1946).

Overbeek, J.T.G., Vink, C.L.J., Deenstra, H.

Rec Trav. Chim., 74, 8l (1955).

87

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