Bureau of

ry,

fc- ,!i 11961 NBS TECHNICAL NOTE

Electrochemical

Analysis Section:

SUMMARY OF ACTIVITIES

JULY 1966 TO JUNE 1967

423

***" ° F *o„ IS /

U.S. DEPARTMENT OF COMMERCE

National Bureau of Standards
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UNITED STATES DEPARTMENT OF COMMERCE

Alexander B. Trowbridge, Secretary NATIONAL BUREAU OF STANDARDS • A. V. Astin, Director ^J^TECHNICAL NOTE 423

ISSUED OCTOBER 1967

Electrochemical

Analysis Section: SUMMARY OF ACTIVITIES JULY 1966 TO JUNE 1967

Edited by Roger G. Bates

Electrochemical Analysis Section

Analytical Chemistry Division

Institute for Materials Research

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

listed or referred to in the open literature.

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

Price 50 cents.
FOREWORD

The Analytical Chemistry Division was established as a separate division at the National Bureau of Standards on September 1, 1963, and became part of the Institute for Materials Research in the February 1, 1964, reorganization.

It consists at present of nine sections and about 100 tech- nical personnel encompassing some 45 different analytical

competences from activation analysis and atomic absorption t< 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 o 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 highlj

important output of our laboratories. In addition, however,

it has been our experience that informal, annual summaries o progress describing efforts of the past year can be very valu- able 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 require months to perfect and

yet all too often must be covered in only a line or two of a journal article.

ii
Thus our progress reports endeavor to present this infor- mation 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 1967 we plan to issue these summaries for all of our sections. The following is the third annual report on pro- gress of the Electrochemical Analysis Section.

W. Wayne Meinke, Chief Analytical Chemistry Division
PREFACE

This is the third annual progress report of the Electrochemical Analysis Section of the Analytical Chemistry

Division. The report covers the fiscal year 196"J , which began on July 1, 1966, and ended on June 30, 1967

The Section concerns itself particularly with ionic processes occuring in solution, with areas of analytical measurement where ionic equilibria play a part, and with

the explanation of solution behavior in terms of the inter- actions of ionic solutes with solvent molecules. We do not attempt to cover the entire field of electrochemistry or even of electroanalytical chemistry. The special interests

of members of the Section lie for the most part in solution

electrochemistry, electrochemical thermodynamics, and ionic transport phenomena. In the context of modern analytical

chemistry, their competences bear most directly on potentiometry, electrometric titrations, and conductometric analysis. In line with a uniform policy of the Division, the Section's programs have both research and sample aspects. During the fiscal year just ending, about 70 percent of the total effort was devoted to research while 30 percent was devoted to Standard Reference Materials.

The permanent staff of the Section was reduced by four during the year, and under staffing Inevitably left its im- print on the progress of our research programs. At the conclusion of their appointments as Commerce Science and Technology fellows, Dr. Robert Gary and Dr. Richard K. Wolford accepted positions in scientific administration

outside the Division. Later In the year Vincent E. Bower

and Cyrus G. Malmberg left the Section by transfer to other

MBS Divisions. However, Dr. Bert R. Staples joined the
Section in May. In September Dr.' Donald Rosenthal came

to us in the capacity of Visiting Expert and remained until

February. During the same period, James T. Clark, Jr., a graduate student of Dr. Rosenthal, was a Guest Worker in

the Section. Philip K. Hall of the University of Wisconsin

at Marshfield joined us for the summer of 1966

The purpose of this report is to summarize the total activity of the Electrochemical Analysis Section during

the reporting period. Unlike the formal journal publications that emanate from the Section, this report offers an opportunity to present the program as an integrated whole,

to set forth the objectives of this specialized group, and

to underscore the relationships between the Section's

activities and 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 prepara- tion of this report is gratefully acknowledged.

Roger G. Bates, Chief Electrochemical Analysis Section

Washington, D.C. June 30 , 1967
TABLE OF CONTENTS

PAGE

1. FACILITIES AND EQUIPMENT 1 Space and Equipment 1 Plans and Needs 4

2. MEASUREMENT OF ACIDITY 6 Concepts and Program 5 pH Standards: Sodium Hydrogen Succinate .... 8 pH* Standardization: Status of Acidity

Standards for Alcohol-Water Media 10 pD Standards 13 Present Status 13 Potassium DIhydrogen Citrate 16 Measurement of Acidity In Biological Media ... 18 Measurement of Acidity in Sea Water 20 Behavior of Glass Electrodes In Deuterium

Oxide 22 Method 23 Experimental 25 Results 25

J. SOLVENT EFFECTS ON IONIC PROCESSES 29 Ion-Solvent Interactions 29 Medium Effects for Chlorides in Methanol- Water Solvents 31 Medium Effects for Individual Ionic Species . . 35 Potentiometric Studies in N-Methylproplonamide . 38 Acid-Base Behavior in Aprotic Organic Solvents . 39
PAGE EQUILIBRIUM DATA FOR IONIC SYSTEMS 40 Dissociation of PIperazinium Ions In Water ... 40 Dissociation Constants of Substituted

Anilines and Phenols 45 Isopiestic Vapor Pressure Measurements 47 Spectrophotometry Studies in N-Methylpropionamide 48 CONDUCTANCE AND TRANSPORT BEHAVIOR OP ELECTROLYTES 49 Objectives and Approach 49 Analysis of Errors of the Jones-Dike

Conductance Bridge 50 PREPARATION AND PROPERTIES OF SOLVENTS 54 Ethanol (95 Volume Percent) 54 N-Methylpropionamide 56 NONAQUEOUS TITRATIONS: DETERMINATION OF HYDROGEN FLUORIDE 62 FOREIGN- CURRENCY PROGRAMS 66 IMPACT OF SECTION'S ACTIVITIES 67 PERSONNEL 68 COMMITTEE MEMBERSHIPS 69 PUBLICATIONS AND MANUSCRIPTS, JULY 1966 TO JUNE 1967 70

TALKS, JULY 1966 TO JUNE I967 73 LIST OF REFERENCES .......... 75
LIST OF FIGURES FIGURE NO. PAGE

1. One of the emf laboratories 1

2. Laboratory for conductance studies 2 3. Equipment for precise emf measurements

with glass electrodes 2

4. Assembly of equipment for electrometric titrations 9

5. Assembly of apparatus for studies of

the behavior of glass electrodes In heavy water 26

6. Emf cell used for the direct and indirect comparison of glass electrodes with the deuterium gas electrode in heavy water . . 27

7. Emf cell with liquid junction, containing

a calomel electrode prepared with heavy water 27

8. Direct comparison of two glass electrodes with deuterium gas electrode over a range of pD 28

9. Differences between the medium effects

for individual ionic species in methanol-water solvents 34

10. Extrapolation curve for obtaining pK-j_ from the values of pK-[ 42

11. Variation of pK^ for piperazlnium ion

with temperature 44

12. Relationship between the pK of substituted protonated anilines (pK^) and that of the corresponding substituted phenols

(pK

p

) 46

13. Switch panel used with conductance bridge

to provide a resistor in series or parallel with the cell and to reverse

the cell leads 53

14. Cell used for measurements of relative dielectric constants in NMP and NMP-cyclohexane mixtures 57
LIST OP FIGURES (con't) FIGURE NO. PAGE

15. Excess molar volume (V 1 p-N1V1 -NpV2 ) of NMP- water mixtures at 30 C as a function

of the mole fraction of water (N2 ) .... 59

16. Dielectric constant deviations

(D]_ 2 -N]_D]_-N2D 2 ) of NMP-water mixtures

at 20, 30, and 40 °C as a function of the mole fraction of water (Np) 60 LIST OF TABLES TABLE NO. PAGE

1. Stability of cells of the type Pt;H2 , NaHSuc (0.02m), KC1 (m), AgCljAg 10

2. Standard values of pH*(S) in 50 wt percent methanol 11

3. Standard reference values of pa.p for an acetate buffer solution and a phosphate buffer solution in deuterium oxide from 5 to 50 °C 14

4. Correction factors for operational pD measurements in heavy water solutions . . 26

5. Medium effects for chloride salts in methanol-water solvents at 25 °C 33

6. Medium effects (log mYu) for the transfer of hydrogen ion from water

to methanol-water solvents 37

7. Comparison of thermodynamic constants

for the dissociation of piperazinium,

morpholinium, piperidinium, and pyrrolidinium ions at 25 °C 43

8. Dielectric constants of N-methylpropionamide-cyclohexane

mixtures at 25 °C 6l

9. Analysis of variance for fluoride titrations 64

10. Analysis of variance for standardization

of base 64
ELECTROCHEMICAL ANALYSIS SECTION: SUMMARY OF ACTIVITIES,

JULY 1966 TO JUNE 1967 Edited by Roger G. Bates ABSTRACT

This survey of the activities of the Electrochemical

Analysis Section, Analytical Chemistry Division, covers

the period July 1966 to June 1967 . An attempt is made to summarize a year's progress on the technical projects of the Section In such a way as to stress the program and capabilities of the organizational unit as a whole. A description of facilities and equipment is presented, and

the directions of the programs are indicated. Brief sum- maries of several lines of work now under way are given. The main areas include the measurement of acidity, with

progress toward the establishment of standards for pH, pH*

and pD being reported. Studies of the behavior of glass electrodes in deuterium oxide are described and solvent effects on Ionic processes are discussed. Equilibrium

data for certain Ionic systems have been obtained, modifications of a precision conductance bridge are described,

and a method for titrating hydrogen fluoride In dimethylformamide is outlined. The survey concludes with lists

of the members of the Section staff, publications and manuscripts produced during the year, talks given by the

staff, and committee assignments. Key Words Electrochemical analysis, Analysis, Acidity, Glass electrodes, pH measurements, Medium effects, Conductance

measurements, Electrochemistry of solutions
1. FACILITIES AND EQUIPMENT

A. Space and Equipment

As presently arranged, the Section's facilities and equipment provide five well-equipped laboratories for highly accurate emf measurements, a conductivity laboratory,

an instrumentation room, a solvents preparation laboratory,

a balance room, and space for spectrophotometric and iso- piestic vapor pressure studies „ One of the emf laboratories

is shown in figure 1 and a corner of the conductance lab- oratory is seen in figure 2.

>

Figure 1. One of the emf laboratories

•

Four new constant-temperature baths of stainless-steel construction were received during the year. One of them

can be seen in figure 3. These thermostats are provided

with 1/2-hp built-in compressors which, in spite of a rather
Figure 2. Laboratory for conductance studies.

Figure j5. Equipment for precise emf measurements with

glass electrodes.
high noise level, appear to function satisfactorily when properly adjusted. Considerable delay was involved, however, before the manufacturer succeeded in putting the controls into a condition which warranted acceptance. Further modifications were then necessary in order to provide each bath with a brass framework to support the cells and a manifold to distribute the purified hydrogen

gas. Improved temperature control was also provided by external commercial electronic controllers of the propor- tional type utilizing thermistor sensors. Although these have proved generally satisfactory, they are not entirely

free from drifts which become evident when a constant setting is maintained for more than 5 hr

Other modifications, desirable but not essential, are now under way. They will improve the stirring by redirecting the impulse current and by doubling the stirring capacity. The cooling coil will be removed from the corner near the stirrers and extended around the inner wall of the working area. Finally, a drain of larger diameter will

be provided, in order to permit the temperature of the bath to be changed with minimum loss of time. New equipment acquired during the year included two potentiometer facilities, each consisting of a new design

of highly accurate potentiometer with built-in working

cell, null instrument, and standard cell of the unsaturated

variety. One of these new instruments is seen in figure j5 • The potentiometer can be read directly and easily to 10 u.V or better. The unsaturated standard cells will not be used as primary standards, however, and for this purpose

two temperature-controlled laboratory standard units were

acquired. Each contains four saturated standard cells mounted in a box and maintained at a constant temperature

by electrical means.
A commercial conductivity bridge of intermediate accuracy but high convenience was purchased for use in performing conductivity titrations. For the expanding

studies in nonaqueous solvents, the routine rapid measurement of dielectric constants will be necessary. For this

purpose, a new instrument has been purchased, together

with cells suitable for the dielectric constant range

2 to 90. It is planned to acquire another cell which will

extend the range of this instrument to 200 provided suitable standards can be found. The new laboratories at Gaithersburg are proving, as expected, to be fundamentally well designed and well adapted

to the conduct of the Section's operations. Some modifications were inevitable, and most of these have been completed during the current reporting year. Wall cabinets, badly

needed In the laboratories but not installed when the building was new, were acquired gradually and put in place. Drain lines were run to connect with the constant-temperatun

thermostats which are an integral part of the emf equipment

B. Plans and Needs

In general, one of the five emf laboratories in the Section is devoted to studies of aqueous solutions and

to pH standardization, while the others will be devoted

to the Investigation of nonaqueous and mixed solvents, when additional trained personnel to conduct these projects

can be found. Needed to complete this unit is an automatic data read-out center which would record automatically and simultaneously the emf of all cells under study in any or

all of the five laboratory rooms. Multichannel digital voltmeters of the required accuracy (10 (iV up to 1.1 V)

are now available commercially. The data produced would

be programmed for computer analysis, and a permanent record

of the emf values would be filed. At the present time, a
more elaborate automated system capable of programming

temperature changes, monitoring the approach of the cells

to equilibrium and recording the barometric pressure,

does not appear practical.

During the past year, the need for developing com- petence in the area of ion-specific electrodes has been

increasingly felt. This project should include glass

electrodes, the potentialities of which are far from being

exhausted. For studies of these high-impedance systems,

a precision high-impedance electrometer of the vibratingreed type will be needed. As the work progresses, it is likely that time constants of the electrode response will

need to be measured, and for this purpose an oscilloscope

and electrometer of rapid response will have to be acquired.

As noted in our last report [1], there is a continuing

need for methods for the accurate measurement of pH at high

temperatures and pressures. No standards of reference for temperatures above 95 °C are available at the present time. Construction of equipment for precise emf measurements of solutions under pressure in the range 100 to 200 °C can only be undertaken when a scientist to undertake this project has been found.

(R. G. Bates)
2. MEASUREMENT OF ACIDITY

A. Concepts and Program

The concept of the Section's programs on the measurement of acidity was set forth in detail in the annual progress report for 1966 [1] and will not be repeated here.

A few remarks concerning the essentially dynamic

nature of the national pH scale are perhaps in order. According to the NBS approach, the pH is defined in terms

of five primary standards [2]. In the long-term view, however, these standard reference values of pH may he improved and altered as better experimental measurements

are obtained. Furthermore, a search is continually under

way for new standard reference materials of improved

properties, greater stability, and enhanced usefulness. For example, solutions of potassium hydrogen tartrate,

one of the present standards, are subject to molding and may be useless for standardization purposes 24 hr after

preparation. Potassium hydrogen phthalate, in most respects an excellent standard and widely used, is never- theless unstable in contact with platinum-hydrogen elec- trodes. For these reasons, two new promising standard reference materials for the pH range 3.5 to 4.7 are under

consideration. These are sodium hydrogen succinate and potassium dihydrogen citrate. In establishing acidity scales for nonaqueous and mixed solvents, the same concepts are applicable. In addition, management of the programs for acidity in nonaqueous media requires the selection of the most important solvents for study. At the present time pri- orities are assigned after consideration of a) the extent

to which a particular solvent or solvent mixture is used

in science and industry. This decision Is based on observations of published work and on personal inquiries,
b), the compatibility of a solvent system with the electro- metric pH techniques used almost universally for acidity

measurement, and c) the utility of the precise tools for standardization in each particular medium. The NBS stand- ardization procedure relates acidity as directly as possible

to the hydrogen electrode potential. If this electrode

cannot be used in a particular solvent, the validity of any standard scale is impaired to some degree. By far the largest number of pH measurements are made

in media that are nominally aqueous. Some of these media

are of special importance and, due to their fixed composi- tion with respect to predominant solutes, may be considered

to advantage as discrete solvent systems. In this category

fall sea water, blood plasma, and certain concentrated salt media used widely in the study of the stabilities of metal

complexes. When a measurement problem of sufficient magnitude exists in a medium of this sort, it is the NBS policy

to establish standards for these particular media of special

importance. At the present time a need appears to exist for reference standards in blood plasma (or isotonic saline)

and in synthetic or natural sea water. Current procedures for nonaqueous and partially aqueous solvents lead to the establishment of an acidity scale based on the unit pH* [3] . This is a conventional activity

unit referred to a different standard state in each solvent. Hence, it cannot provide the intercomparison of acidities

in two or more solvents that has long been sought. This general scale of pH will only become a reality when an acceptable scale of ionic medium effects has been set up [4]. Medium effects for individual ionic species cannot

be derived by thermodynamic means. Nevertheless, there is

a strong belief that the medium effect (or the related

ionic transfer energy) is a real and useful physical concept.
There is some hope that a valid procedure may be found to separate medium effects for electrolytes into their cation

and anion constituent parts or to calculate from theory the ionic transfer energy of a single reference ion. Success

in this venture would bring to reality the generalized

acidity scale envisioned by Br^nsted [5]. Consequently, medium effects continue to claim the interest and attention

of members of the Section.

(R. G. Bates)

B. pH Standards: Sodium Hydrogen Succinate

In spite of its extremely wide use as a pH standard, potassium hydrogen phthalate is not an ideal material for

this purpose. The most suitable standard for acidity measurements should be stable in contact with the electrode formed by hydrogen gas adsorbed on platinum, as phthalate buffers are not. Furthermore, reference solutions for pH are useful not only in the standardization of electrometric assemblies but also In the preparation of indicator stan- dards for optical ( colorimetric) measurements of pH. Although potassium hydrogen phthalate can be used without difficulty in the visible region of the spectrum, it absorbs strongly in the ultraviolet region and consequently

is unsuitable for many spectrophotometric studies of acid- base behavior including the optical method for the deter- mination of dissociation constants.

As a useful alternative to the phthalate standard, our attention has fixed on sodium hydrogen succinate and on potassium dihydrogen citrate, both of which are free from

the disadvantages encountered in the use of phthalate

solutions,, No concentrated effort has yet been devoted to either of these materials. However, some time was spent during the past year (before the loss of V. E. Bower by
transfer to the Electricity Division) in attempts to pre- pare a sample of sodium hydrogen succinate with the proper

stoichiometric composition. For reasons not clearly

apparent, these were not very successful. When pure suc- cinic acid and disodium succinate were combined in the proper proportions and the mixture crystallized, the salt obtained did not have the expected composition, within the desired close tolerances. Equipment used for electrometric

titrations of this sort is shown in figure 4.

Figure 4. Assembly of equipment for electrometric

titrations

Nevertheless, one series of emf measurements was made

with an equimolal mixture of succinic acid and disodium

succinate from to 25 °C at intervals of five degrees C. Four cells were included in the series. In each the molality of (synthetic) sodium hydrogen succinate was 0.02 mol kg while the molality of added potassium chloride
was 0.015 or 0.01. The emf displayed the expected high

stability over a period of 36 hr, with good agreement between the Initial and final emf at 25 °C (see table l). The study of succinate mixtures was, however, discontinued

at this point until further attempts to prepare a suitable

lot of the acid salt should prove successful. Table 1. Stability of cells of the type Pt;H

2

, NaHSuc (0.02m), KC1 (m), AgCl;Ag.

m(KCl) - 0.015 m (KC1) = 0.01 Initial emf (V, 25 °C) 0.60090 0.60113 O.5832I 0. 58327 Final emf (V, 25 °C) 0.60100 0.60115 O.58316 O.5833O

(R. G. Bates)

C. pH* Standardization; Status of Acidity Standards

for Alcohol-Water Media

In the past several years, considerable attention has been given to acid-base behavior in 50 wt percent aqueous

methanol. In a paper published early in 1965 [6], reference buffer solutions for pH* measurements in 50 percent methanol from 10 to 40 °C were proposed. These solutions consisted of acetate buffers, solutions of sodium hydrogen

succinate, and phosphate buffers, all of which contained

added sodium chloride, necessary for the proper functioning

of the silver-silver chloride electrode. The pa^r* values

of three useful reference standards for acidity measure- ments in 50 percent methanol are given in table 2.
Table 2. Standard values of pH*(s) In 50 wt percent

methanol.

HAc (0.02m) KH

°c NaAc (0.02m) NaCl (0.02m)

10 5-56

15 5.55

20 5.54

25 5.54

50 5-54

55 5.54

40 5.55

NaHSuc (0.02m) Na^PO^ ( 0.02m) NaCl (0.02m) NaCl ( 0.02m)

5.81 7-94

5.79 7.92

5.77 7.90

5.76 7.88

5.75 7.87

5.74 7-86

5.74 7-86

Ac = acetate

Sue = succinate These solutions are useful for standardizing an operational

scale of pH* in this solvent medium. For this purpose, operational definition

PH*(X) = PH*(S) + ^-^ (1) (RT In 10)/P
In subsequent work, pa-rr* values were also assigned to two buffer solutions of a different charge type. These

solutions consisted of a weak base, tris(hydroxymethyl)-

amlnomethane ("trls") or 4-aminopyridine, and Its correspond

Ing hydrochloride. When pH cells of the glass electrode

type were standardized with the first series of buffer

standards, measurements of the tris buffers showed that internal inconsistencies existed among the two sets of standards. Even in the aqueous medium there is some evi- dence that the liquid- junction potential between a con- centrated solution of potassium chloride and buffer solu- tions of weak base-salt type may differ appreciably from

that existing in the cell containing standard buffers of

the weak acid-salt type [7]. It is evident that the 50 percent methanol solvent system requires further study before the accuracy of practical pH measurements in this medium can approach that obtainable in aqueous solutions. The nature of the ab- normality in liquid- junction potentials needs further

investigation. In addition, the acetate, succinate, and phosphate standards should be studied with varying concen- trations of added chloride, in such a way that the pa^* values in the total absence of chloride can be obtained. For standardization purposes, the addition of chloride to the buffer solution is an inconvenience that should be avoided if possible. When these studies have been success- fully completed, suitable standard reference materials will

be acquired and certified for issuance.
In view of its wide usefulness as a solvent in acid- base studies, 95 percent ethanol has been chosen as the next solvent system for investigation. The preparation of highly pure and well-characterized solvent is under study

and the measurement of ion-pair constants in this medium

by conductance methods will be initiated shortly.

(R. G. Bates)

D. pD Standards

1. Present Status

The annual progress reports for the past two years

[1,8] have summarized all of the experimental work, completed to this time, that has been directed specifically

toward the development of standard reference solutions for

the measurement of acidity in deuterium oxide. During the past year, the major thrust in the heavy water studies has been an examination of the response of the glass electrode

to deuterium ion. This work is described in some detail

in section 2. a of this report. A knowledge of the behavior

of the glass electrode in heavy water is an essential pre- requisite to the drafting of procedures suitable for accurate measurements of pD in practice. Thus far, conventional pa^ values have been assigned

to three series of solutions in heavy water, namely deuterium chloride [9], a buffer mixture of potassium

dideuterium phosphate and disodium deuterium phosphate [10], and an acetate buffer composed of acetic acid (CH^COOD)

and sodium acetate [11]. The pa- values for one selected solution in two of these three series are listed in table 3, and the solution compositions appear at the foot

of the table. The details of the assignment procedures

and the nature of the convention used are set forth in the original publications cited above.

13
Table 3. Standard reference values of pa^ for an acetate buffer solution and a phosphate buffer solution

In deuterium oxide from 5 to 50 °C. Acetate Phosphate

t, °C solution solution

5 5 . 27 7-54

10 5.25 7.50

15 5.24 7.47

20 5.24 7.45

25 5.23 7.43

30 5.23 7.41

35 5.22 7- 40

40 5.23 7.39

45 5.23 7.38

50 5.24 7.38

Compositions (in mol kg ): Acetate solution—CH\C00D (0.05), CH^COON (0.05) Phosphate solution—KDPO^ (0.025), Na^PO^ (0.025)

It is well known that a solution of a strong acid is not a good reference for standardizing pH cells with liquid

junction. Similarly, the solution of deuterium chloride

in deuterium oxide, when brought into contact with a con- centrated aqueous solution of potassium chloride, displays

an abnormal liquid- junction potential relative to that

found for buffer mixtures in the same solvent medium.

14
in table 3 are quite satisfactory as reference solutions.

The phosphate mixture is readily prepared from potassium

dihydrogen phosphate and disodium hydrogen phosphate, both

of which are already available as NBS Standard Reference

Materials. The contamination of the heavy water solvent by the exchanged hydrogen when this solution is prepared is slight and usually of little consequence. For example, the deuterium isotopic composition is reduced by 0.075 percent when the two phosphates are each present at a molality

of 0.025 mol kg

-1

. Unfortunately, the acetate buffer solution is not easy

to prepare. The best procedure is probably to add sodium

carbonate in the proper amount to a solution prepared by dissolving a weighed quantity of pure acetic acid in heavy

water. Both materials are hygroscopic and difficult to weigh accurately in the small quantities usually required. The issuance of acetic acid as a standard reference material

does not appear feasible at the present time. For 'the most useful standardization of a practical pD scale, therefore, other buffer systems should be studied.

To replace acetate at the lower end of the pD scale, suc- cinate and citrate buffers are likely possibilities. For

the upper end, a carbonate buffer solution shows promise

of usefulness. Contamination of the solvent with hydrogen

would be minimal with sodium hydrogen succinate and a buffer

solution composed of sodium bicarbonate and sodium carbonate,

as only one atom of hydrogen is contained in molal amounts

of these buffer systems. Potassium dihydrogen citrate would be slightly less suitable in this respect, but its other favorable characteristics warrant its further study. For example, the buffer capacity of solutions of primary

citrate salts exceeds that of acid succinates at the same
molal concentration. Furthermore, potassium dihydrogen

citrate has a considerably lower pD than does sodium hydrogen succinate when these salts are dissolved at the

same concentration in deuterium oxide. Both salts appear

suitable for issuance as well-characterized standard refer- ence materials. When the pa„ values of two or more suitable reference

solutions have been obtained by means of emf measurements

of cells without liquid junction, they can be identified with pD(S) in an operational definition of the practical

pD scale:

PD(X) = PD(S) + Z—£ (2) (RT In 10)/P

The usefulness of this practical scale will, however, depend on the purity of the response of the glass electrode

to deuterium ion and on the regularity of the liquid

junction potential between solutions (X) and (s) in heavy

water and the concentrated aqueous solution of potassium

chloride with which they make contact at the liquid junction. These are factors which will require further study.

(R. G. Bates)

2 . Potassium Dihydrogen Citrate

In a continuation of the program of selection and certification of standard reference materials for pD measurements, potassium dihydrogen citrate was selected as the next material to be studied. At the present time two buffer mixtures, one containing potassium dideuterium

phosphate and disodium deuterium phosphate and the other containing acetic acid and sodium acetate, have already been proposed as primary buffer .standards in heavy water.

16
However, the second buffer mixture is rather inconvenient

to prepare. It would therefore be preferable to use potassium dihydrogen citrate, which is a solid material and can be prepared and purified with relative ease. In solution

this salt is converted, by exchange with the heavy water, into potassium dideuterium citrate, and its solutions in heavy water exhibit a pD value near 4.5.

a. Method . The standard values of pa_ for the newly selected buffer material will be derived by a method which follows closely the usual NBS procedure for establishing primary pH standards in ordinary water. The acidity function p(a

DYcl ) for a particular buffer concentration

with different amounts of added soluble chloride will be determined by measuring the emf of the cell

Pt ;D 2 (g), KH

2Cit(m

1

), KCl(m

2

) In D

2

0, AgCljAg (3) over a range of temperatures. The measured emf (E), the standard emf (E°), and natural constants are then used in

the computation of acidity functions p(an 7r

, 1 ). The function

p(a

T.7ri

)° is evaluated by extrapolating the values of

p(aT.Yni ) to zero concentration of the added chloride. The pa-p. values are derived by means of the equation

paD

= p(aD

Ycl

)° + log yc1 (4) where the ionic activity coefficient of chloride ion is estimated by the use of the Bates-Guggenheim convention [12]

AT 1/2 -log y = M (5)

17
In eq 5, A is a constant of the Debye-Hlickel theory, the magnitude of which reflects the dielectric constant and density of the heavy water solvent, and I is the ionic

strength.

b. Experimental . The potassium dihydrogen

citrate was prepared according to the directions of Kolthoff

[13]. Citric acid was neutralized with a calculated amount

of potassium carbonate to yield the desired primary salt. Two lots of potassium dihydrogen citrate were prepared by

this method. They were assayed by weight titration with a standard solution of sodium hydroxide which itself had been

standardized against a- standard solution of distilled hydrochloric acid. The endpoint of the titration was determined potentiometrically . The first lot assayed 100.01 percent

(standard deviation 0.04 percent) and the second lot gave 100.03 percent (standard deviation 0.03 percent), Indicating

that the material was sufficiently pure for further study.

A few emf measurements of cell 3 have been made in order to select the most suitable concentration for the proposed

reference solution.

(Maya Paabo)

E. Measurement of Acidity in Biological Media

There is an increasing awareness of the role that the Analytical Chemistry Division can, and should, play in improving the measurements on which the success of bio- medical research and the accuracy of clinical diagnoses

depend. Likewise, the Section is gradually becoming more deeply involved in the special problems encountered in the measurement of the acidity of biological media.

Some years ago, a standard reference phosphate buffer

for pH measurements of blood was established [14], This work was undertaken to meet the needs of clinical chemists

18
and research workers for a. standard solution of pH about 7.4 at 37 °C which would facilitate the discrimination of

pH differences in blood with an accuracy greater than 0.01

unit. This standard is widely used. Three conferences of the New York Academy of Sciences

have served to strengthen the lines of communication be- tween the NBS specialists in pH and research workers in biomedicine. The first took place in December i960 and was entitled " In Vitro and In Vivo Effects of Amine Buffers"

(see reference [15]), while the second, entitled "Current Concepts of Acid-Base Measurement", was held in November

1964. The third, devoted to the subject "Bioelectrodes", took place in June 1966. The chief of the Section was an invited speaker at all three of these meetings. Reprint

requests for the second of these lectures [16] have exceeded

350 and are still being received. As an outgrowth of this conference, R. G. Bates was asked to serve as a member of two ad hoc committees, one dealing with acid-base terminology and the other with methodology. The committee reports

have been published [17]. The proceedings of the third conference have not yet appeared . A comparatively recent development in the electrochemical investigation of biological systems is the glass microelectrode . These electrodes are of both the openand closed-tip designs, and the tip diameters may be as small as 0.1 |j,m. They have been used successfully to measure pH and cation concentrations as well as membrane potentials. An illuminating conference devoted to intracellular

glass electrodes was held at the University of Montreal in May 1967, at which the chief of the Section presented an invited lecture on factors determining the choice of inner reference systems for these electrodes. He has likewise been invited to be a co-author of a review article on micro19
electrodes, the major portion of which is being compiled

by Drs. Butler, Waddell, and Poole of the Medical School, University of North Carolina. These conferences, together with a visit to the glass-electrode facility of the Corning Glass Works in the spring of 1967^ have produced suggestions for new standard reference materials capable of enhancing the reliability

of clinical and research data. One proposal is for the establishment of a standard with the same pH, ionic strength, sodium ion concentration, and protein concen- tration as blood plasma. Another proposal is for a pH standard in isotonic saline. These suggestions are both meritorious and feasible, and it is hoped that they can

be incorporated into the Section's programs in the near future

.

(R. G. Bates)

F. Measurement of Acidity in Sea Water Growing interest in the study of the oceans demands

that techniques be developed for accurately determining

the hydrogen ion concentration or activity in the seas. In the interest of serving the research community, the NBS has set out to develop pH standards for media approx- imating sea water in composition. These standards would

be made available through the Standard Reference Materials

program, as are the present materials for standardizing

pH equipment for general use. To achieve this goal, data must be accumulated on the effect of sea water on the activity of hydrogen ion and on the response of pH electrodes

.
The initial phase of this work is being performed in collaboration with James T. Clark and Dr. Donald Rosenthal

of Clarkson College of Technology. A chloride solution has been prepared which simulates the composition of natural

sea water. The composition of this simulated synthetic

sea water (SSW) is as follows: NaCl 0.4158 mol kg

-1 MgCl

2 0.05417

CaCl

2 0.01034

KC1 0.00923

SrCl p 0.000158

The ionic strength of the SSW is varied by dilution. This composition is similar to that of Lyman and Fleming's

formula for 19$ chlorinity sea water [18]. The emf is being determined over the temperature range

5 to 45 °C for dilute solutions of hydrochloric acid in SSW, and in diluted concentrated solutions of similar relative

composition, using the cell with a hydrogen electrode

and a silver-silver chloride electrode. The following

molal concentrations of hydrochloric acid are being studied:

0.005, 0.01, 0.033, 0.066, and 0.10. The acidity function

pa-rYp., [that is, -log(aTT+Y„ 1 - ) ] is derived from the measured

emf and the known concentrations, and its variation with

the chloride concentration is being studied. Thus far information has been obtained for most of the chloride and HC1 concentrations. Repeat determinations

must be run in several cases to check on the reproducibility

of the emf measurements. In addition, some of the solutions

have yet to be measured. In all, approximately 50 percent

of the experimental work has been completed.
In addition to the above study with the hydrogensilver chloride cell without liquid junction, a similar

study with the glass-calomel system Is to he undertaken.

This will involve measuring the glass electrode pH in these same solutions of HC1-SSW in order to measure the effect of salt concentration on the potentials of the glass electrode.

(J. T. Clark, Jr.)

G. Behavior of Glass Electrodes In Deuterium Oxide

It is a common practice to use commercial glass electrodes for measurements of acidity in water, in many

nonaqueous solvents, and also in deuterium oxide. Various

workers have found that in order to obtain pD from the operational pH value, referred to the usual pH standards prepared In ordinary water, an empirical correction of 0.40

pH unit should be added to the pH meter reading [19,20]. Recently, however, a pD scale based on conventional values

of pa^ in certain buffer solutions has been established through measurements of the emf of the cell [10,11]

Phosphate and acetate buffer solutions were used In the

work. To examine the experimental basis for a practical

scale of pD, a study designed to affirm the deuterium

response of the glass electrode and the validity of the empirical correction factor was undertaken [1]. Glass electrodes were compared both directly and indirectly

against deuterium gas electrodes in buffered solutions covering a range of 12 pD units.
1. Method

The experimental procedures consisted of two major

parts. First, the behavior of several glass electrodes from different manufacturers was observed in buffer solu- tions in heavy water after the glass electrodes had been

standardized in a buffer solution made up in ordinary- water . A saturated aqueous calomel electrode was used as the reference electrode. The cell can be represented by

calomel

KC1 (satd.) standard buffer ; glass (7) in H o The apparent pD can be calculated from the emf (E) of this cell by the following equation:

pD (apparent) = pH(s) +

(RT In 10)/F

(8)

glass electrode is immersed in the standard or in the "unknown" buffer solution in heavy water, respectively,

each half cell being measured against the calomel reference

electrode. The pH(s) is the assigned value of the standard solution on a conventional activity scale in water. Measurements were also made of the same type of cell in which the glass electrodes were replaced by deuterium

or hydrogen gas electrodes:calomel KC1 (satd.) standard buffer ,

The apparent pD values obtained from measurements of these

two cells with liquid junction were then compared with the

23
corresponding pa^ values obtained from measurements of cell 6 containing the same buffer solutions in heavy water

as did cells 7 and 9. The empirical correction factor

5 results from a comparison of pa^ with the apparent pD derived from the emf of the cells with liquid junction:

= pa_ - pD (apparent) (10)

In the second procedure, the behavior of glass elec- trodes in heavy water was observed after these electrodes

had been standardized in a reference buffer solution pre- pared in heavy water. The cell with liquid junction used contained a saturated calomel electrode also made up In heavy water. This cell avoids the differences in liquid

junction potential between two different solvents. It can be represented by Pt;D2 or calomel KC1 (satd.) standard buffer : Pt;D 2 or glass (11)

The operational pD furnished by this cell will be designated pD to distinguish it from pD (apparent) found when

standards in ordinary water were used and the bridge solu- tion consisted of saturated KC1 in ordinary water:

pD = pD(S) + • (RT In 10)/F

(12)

Again E„ and E„ are the emf values of the left and right half cells measured against the calomel reference electrode.

24
By combining the values of pD and the pa^ derived from independent measurements of cells without liquid junction, correction factors were obtained for both glass and gas electrodes

5 = paD

- pD (13) Finally, in addition to the indirect- comparisons of the glass electrodes with deuterium gas electrodes as outlined

above, a direct comparison was made by means of the following cell: Pt;D (g), buffer in D

2

0; glass (l4)

2. Experimental

The heavy water used was of high isotopic purity

(about 99-75 percent). Measurements of the emf of cells with glass electrodes were made on a commercial pH electrometer of the battery type. Systems of low electrical resistance were measured with a precision potentiometer. The electrometer and other associated measuring

equipment are shown in figure 5- All measurements were performed at 25 °C . The buffer systems selected for study covered the pD range from 1 to 13 • Preparation of the solutions, electrodes, and cells followed the usual pro- cedures of the Section. Two of the cell vessels used for

the study of glass electrodes are illustrated in figures 6 and 7.

3. Results

The operational correction factors for six of the buffer systems studied have been calculated and are given

in table 4. It is evident that the correction factors which need to be applied to pD measurements based on pH standards in ordinary water are constant but differ in magnitude for glass and gas electrodes. The correction

appears to be nearly zero when the glass electrode is 25
Figure 5. Assembly of apparatus for studies of the behavior of glass electrodes in heavy water. Table 4. Correction

measurement

fa

s ctors for operational pD in heavy water solutions Buffer Standard buffer

and reference elec- trode in HpO

Standard buffer

and reference elec- trode in D

2

8 (gas) app v& ; *app (glass) 5 (gas) 5 (glass Citric acid +O.O67 +0.470 +0.003 -0.001

D

2Suc, NaDSuc +0.068 +0 . 466 -0.001 -0.018 DAc NaAc +0 . 077 +0.460 +0.009 -0.002 Borax +O.O78 +0.472 +0.016 +0 . 034

NaDCO,, Na o C0 +0.079 +0.459 +0 . 007 +0.043

Sue = succinate

Ac = acetate

26
Figure 6. Emf cell used for the direct and indirect comparison of glass electrodes with the deuterium gas electrode in heavy water.

Figure 7. Emf cell with liquid junction, containing a calomel electrode prepared with heavy water.

27
standardized in a reference buffer solution prepared in heavy water. The direct comparison of two of the glass electrodes with the deuterium gas electrode is illustrated

in figure 8. These glass electrodes follow the potential

of the deuterium gas electrode within 0.5 mV over most of

the pD range studied. There is evidence, however, that

some glass electrodes show greater alkaline errors than

would be expected in the corresponding aqueous solutions.

Figure 8. Direct comparison of two glass electrodes with the deuterium gas electrode over a range of pD

.

(Maya Paabo)

28
3. SOLVENT EFFECTS ON IONIC PROCESSES

A. Ion-Solvent Interactions

As solvent effects on the equilibrium behavior of strong and weak electrolytes are examined in ever greater

detail, it becomes increasingly evident that electrostatic

effects related to the changing dielectric constant of the medium cannot account for the observed changes. The electrostatic treatment alone is moderately successful in amphiprotic hydrogen-bonding media of relatively high

dielectric constant. At the other extreme, in aprotic

solvents of low dielectric constant, other factors play

a dominant role. The properties of solvents--dielectric

constant, hydrogen-bonding capacity, basicity, and solvating power, for example--vary widely, and consequently all gradations in behavior may be expected and are indeed

found in practice.

A convenient way of examining the solvent effect on

an Ionic process involving two or more species is through

the transfer free energy AG°, namely the free energy change

for the transfer

I (in water, w) = i (in solvent s) between the standard states In the solvents w and s: AG° =

s G°(i) - w G°(i) (15) Another useful quantity for comparison purposes It the "medium effect" y. or "medium activity coefficient", which

is simply related to the free energy of transfer [21]: log m

71 = AG°/(RT In 10) (l6)

29
Transfer free energies and medium effects for uncharged

species, for neutral combinations of ions, and for certain other ionic combinations can be obtained with full thermo- dynamic rigor by experimental means. If the transfer energy

for any single ionic species could be measured or calculated, these combinations would permit a scale of medium effects

for all ions to be set up. So far this is not possible. Nevertheless, differences in the medium effect for two ionic species are often revealing, and information of this sort can often be derived from measurements of solvent effects

on equilibrium constants such as the dissociation constants

of weak acids. In this way, it has been possible to identify several influences of qualitative significance. The increased free energy of an ion as the dielectric constant decreases has

long been recognized and is given semi-quantitative expression through the Born equation [22]. The relationship between dielectric constant and the formation of ion pairs

has also been elucidated [23]. When the hydrogen-bonding

capacity of the solvent is markedly reduced, an entirely

different group of interactions may become important [24]. The efforts of the conjugate bases of protonated species HA and BH to achieve stability may lead to the formation

of the homoconjugates AHA and BHB , to analogous heteroconjugates, and to higher aggregates (see, for example, references [25] and [26]. Anions such as picrate and other phenolates which are "delocalized oscillators" may interact particularly strongly with solvent molecules through en- hanced dispersion forces [27]. It is quite evident that an understanding of Ion-solvent interactions would provide the key to an explanation of the many diverse aspects of what are known as solvent effects.

30
Furthermore, scales of ionic medium effects, whether based

on experiment or theory, would provide the basis not only

for a single scale of electrode potentials but for a single "universal" acidity scale. For these reasons, ion-solvent interactions provide a unifying theme for a large part of the research activity

of the Section. This subject is of interest here not only from the equilibrium point of view but also in the study

of transport processes such as conductance and diffusion.

It is believed that progress in this area is an essential prerequisite to advancements in the electrochemistry of nonaqueous media

.

(R. G. Bates)

B. Medium Effects for Chlorides in Methanol-Water Solvents

Solvent effects on the acidic dissociation of positivelycharged weak acids were discussed in some detail in the annual progress report for 1966 [1]. This process can be written in general terms as follows

BH

+

+ SH = SHg + B ( 17 ) where B represents an uncharged base such as ammonia or tris(hydroxymethyl)aminomethane ("tris", THAM) . A Guest Worker in the Section, Professor Paul W. Schindler, studied

the dissociation of tris in methanol-water solvents with

compositions ranging from 30 to 90 wt percent methanol and determined the solubility of the base over the entire composition range from to 100 percent methanol at 15, 25, and 35 °C. These data, together with results in the literature for the energies of transfer of hydrochloric acid and

the alkali chlorides, from water to water-methanol solvents,

31
have made it possible to account for the minimum in the pK

of protonated tris and ammonium ion as methanol is added

to the aqueous solvent. A manuscript on this subject has been prepared for publication.

The difference between p K in the methanol-water

mixture and p K in the aqueous medium depends on a combina- tion of the medium effects y. for the species participating in the dissociation process. For the dissociation of tris-H+ (designated TH+

), p

s K - PwK = logrn^T + l0g raV* log m%i+ < l8 ) which may also be written in terms of the mean medium

effects for the two chlorides HC1 and THC1:

P*K " K,K = A + 2 log y - 2 log *____ (19) Inasmuch as medium effects for individual ions in a given

solvent medium are, by definition, completely independent

of other counter ions, an equation analogous to eq 19 could be written for bromides, iodides, perchlorates, and

so on.The solubility measurements permit reasonably accurate

values for

mYT

to be obtained, and YHC1 is known with con- siderable precision from extensive measurements of the standard emf of cells with hydrogen electrodes and silver- silver chloride electrodes containing hydrochloric acid in methanol-water solvents (see reference [28]). Hence, the measurements of p g K - pwK permit the medium effect for the transfer of tris hydrochloride from water to methanol-water

solvents to be calculated by eq 19 . The values obtained

are compared in table 5 with similar data for lithium

chloride and sodium chloride derived from solubility mea- surements [29] . 32
ble 5. Medium effects for chloride salts in methanol-water solvents at 25 °C. Wt 'MeOT Cent l0g mYTMC1 l0S mYLiC1 l0g mYNaC1

30 0.319 O.36I 0.492

50 O.563 0.620 0.853

70 O.867 0.936 1.282

90 1.240 1.369 1 . 87C

Although medium effects and ionic transfer energies

for individual ionic species are not experimentally obtainable (see next seetion), eq 18 shows that differences between the logarithms of the medium effects of ions of

like charge can be derived unambiguously from measurements

of solvent effects on the pK for certain dissociation

equilibria. Furthermore, it is evident that similar information can be obtained from medium effects for electrolytes measured by emf, solubility, and other methods.

Thus, log mTM+

- log m

yH+ = log m

ym

- log mTHX (2Q) where M+ may be, for example, an alkali ion and X~ a halide

ion. Figure 9 is a plot of some of these differences

calculated from pK measurements and from data given in the literature. The form of the curves shown in Figure 9 is suffic- iently unusual as to require comment. The sharp rise, independent of whether M

+ (or BH+) is tris-H+, ammonium

ion, or lithium ion, shows that the transfer energy of the

33
I.V 1

' T —Z^+i /XNH4

+ \

0.8 E jF/*~~^ ^ Li+- f

- ft6 JV/ (CH 2

0H)3 CNH 3

+ +

o> 0.4 o ///

02 - ///

^ 1 1 1 1 20 40

WT.

'

60 Me OH

Figure 9. Differences between the medium effects for individual ionic species in methanol-water

solvents

.

proton is much smaller than those of these three other

cations at methanol concentrations from to 70 wt percent. Above 70 wt percent, however, the curves begin to fall off, suggesting that the free energy of the proton in methanolrich solvents is rising more rapidly than that of the other

cations. It is tempting to attribute this observation to

a "basicity effect" [30] . In these terms, the basicity of 34
methanol-water solvents would be greater than water Itself

in the low range of methanol concentrations but would decrease rapidly as the amount of water in the mixture

becomes low. There is no general agreement as to the ef- fective basicity of alcohols and alcohol-water mixtures [31].

(R. G. Bates)

C . Medium Effects for Individual Ionic Species

The realization of several long-term goals of analytical and physical chemists intimately involved in measurement problems depends on the establishment of a valid and useful scale of medium effects for individual ionic species. Scales of electrode potentials, for example, have been set up in several different solvents, but these values must necessarily always be referred to an arbitrary zero, usually

the standard potential of the hydrogen, hydrogen ion couple

in each particular solvent under study. Fusion of scales

of relative electrode potential into a single unified

system would demand a knowledge of the standard hydrogen

electrode potential in each of these solvents relative to

a single, solvent-independent, reference point. The latter

would most conveniently be the standard hydrogen electrode

in the aqueous medium. The difference between the potential

of the standard hydrogen electrode in a given solvent and

that in water is directly related to the transfer free energy and the medium effect for the hydrogen ion (proton): RT

wE°(H) - s E°(H) = — In m

YH

(21)

35
Similarly, the standard scales of pH* being established in the Section for mixed solvents and other amphiprotic

media are based on conventional values of paH

* defined in terms of the activity of hydrogen ion related to the standard state in each solvent or solvent mixture:

paH

* = -log (r^-

s 7H

) (22)

Although these scales are capable of furnishing relative

values of the acidity of different solutions in the same solvent, they cannot relate the acidities of solutions in different solvent media to a single reference point. The situation is thus analogous to that encountered with electrode potentials. Again it is the medium effect for hydrogen ion or the proton that is needed in order to realize a general paH scale which would be based on the aqueous standard state regardless of the solvent. Only

such a general scale could interrelate the acidity of different solvent media in a valid and useful quantitative

fashion:

(lVsVmV = paH* - l0g

rrJH ^

There have been many suggestions for interesting and novel approaches by which the medium effect of a single ionic species might conceivably be evaluated. As we have already indicated, the derivation of the medium effect of one ionic species in each medium is a necessary and suf- ficient condition for setting up a scale of medium effects

for all ions in that medium through the use of experimentally accessible medium effects for ion combinations. Although

36
many of these proposals are eminently reasonable as well

as highly ingenious, the results furnished thus far are in serious disagreement. The variation even extends to the sign of the logarithm of the medium effect, as shown in table 6, where results for the medium effect of hydrogen

ion in methanol-water solvents are given. Table 6. Medium effects (log y„ ) for the transfer of hydrogen ion from water to methanol-water

solvents

2] Feakins, deLigny and Tomkins [33] Alfenaar [34] -0.37 -0.19

-2.05 -1.02

-2.37 -3-91

It is clear that the Section is intimately concerned

with all efforts to establish scales of ionic medium ef- fects and ionic transfer energies. During the past year, considerable effort was devoted to reviewing and evaluating

recent progress in this interesting and controversial area. The results of this examination were incorporated in a chapter entitled "Medium Effects and pH in Nonaqueous

Solvents" which was prepared for a book Solvation Inter- actions to be published by Marcel Dekker (J. F. Coetzee

and C. D. Ritchie, Eds.). This material also formed the basis for the keynote address to the Symposium on Electrochemistry in Nonaqueous Media at the Dallas meeting of the Electrochemical Society in May.

(R. G. Bates)

Wt percent I zmailc Me OH

10 _ 43.1 0.35

50 0.43

90 1.45

100 3-3

37
D. Potentiometrlc Studies in N-Methylpropionamide Dawson and his associates [35] have studied N- methylamides, which have high dielectric constants. One

of these compounds, N-methylpropionamide, is a liquid

stable over a. wide temperature range, freezing at —4o °C and giving reproducible dielectric constant and density measurements at temperatures in excess of'150 °C. The dielectric constant of the liquid is 176 at 25 °C . Furthermore, the liquid amide has remarkable solvent properties,

being completely miscible with water and several of the hydrocarbons of low molecular weight. Many organic and inorganic compounds are likewise quite soluble in N- methylpropionamide

.

A study of acid-base equilibria and pH measurements

in this solvent has been initiated. Accurate and repro- ducible potentiometric measurements were made of dilute hydrochloric acid solutions at 25 °C using a cell consisting

of a hydrogen electrode and a silver-silver chloride e- lectrode. Collateral measurements were also made at 10, 40, 55, 85, and II3.6 °C. The hydrochloric acid solutions

were somewhat unstable at the higher temperatures, and hence the emf data at these temperatures are less accurate than those obtained at 25 °C. The acidity function p(a

H

fyfH

-, ) can be calculated from

the emf values, and the standard emf can be derived by an extrapolation procedure. The measurements have been

completed, but the analysis of the results is still under

way. (Donald Rosenthal)
E. Acid-Base Behavior In Aprotlc Organic Solvents Under special arrangement, Dr. Marion Maclean Davis,

who left the Section at the end of 1965, devoted the calendar year 1966 to the completion of a comprehensive

and important manuscript reviewing the behavior of acids

and bases in aprotic solvents. This work will he published

as an NBS Monograph in the near future, and a briefer

version will appear as a chapter in one of the volumes of The Chemistry of Non-Aqueous Solvents (j. J. Lagowski, Ed.). The scope of the manuscript, which is entitled "Acid-Base Behavior in Aprotic Organic Solvents", is indicated by the following abstract:

,r A unified picture of acid-base behavior in aprotic organic solvents is presented, based on an extensive survey of the literature and experimental

results of the author and associates. Evidence

given to support this picture includes data pertaining to colligative properties of acids, bases, and salts and also conductance, dielectric constants, distribution between immiscible solvents, and spectral absorption in the infrared, visible, and ultraviolet.

The solvents emphasized are hydrocarbons and halo- hydrocarbons, but data for dipolar aprotic solvents

(for example, acetone, acetonitrile, and nitrobenzene)

are included. Contrasts in acid-base behavior and in acidity and basicity scales in aprotic and water-like

solvents are discussed.

"The role of hydrogen bonding in aprotic solvents

is discussed at length. Important types of hydrogen- bonded structures include chelate rings; self- associated acids, bases, and salts; hydrogen-bonded

Ion pairs; and homo- and heteroconjugate cations and anions. Examples are given in which hydrogen bonding

of these types affects such properties as the absorp- tion spectrum of a salt, the catalytic effect of an

acid, and the accurate location of a titration end- point .

(R. G. Bates)

39
4. EQUILIBRIUM DATA FOR IONIC SYSTEMS

A. Dissociation of Piperazinlum Ions in Water

As part of a study of heterocyclic nitrogenous bases, the dissociation constant of piperidine, C,-H. =NH, was measured several years ago in this section [36]. The twofold object of the work was to study the effect of structure on the basicity and other thermodynamic properties

of this substance and to look for further useful buffer

solutions for pH standardization and control. In a con- tinuation of this work, a study was later made of the dissociation constant of pyrrolidine, ChHg=NH, a compound

that differs from piperidine in having a f ive-membered ring

instead of a six-membered ring [37]. More recently, similar studies have been concerned with morpholine,

0=ChHo=NH, a compound in which an additional hetero atom

is introduced into the piperidine ring [38]. During the past year, work has been completed on piperazine, HN=C2,Ho=NH, a base that can, on protonation,

form two acidic species, namely H 2 N=Ci|Hg=NH and H2N=C2jHg=NH2 • The dissociation of these two acids can be represented

schematically by

PzH

2++

+ H

2 = H

+ + PzH

+

(24) and

PzH+ + H

2 = H

+

+ Pz (25) where Pz is written for the neutral base piperazine.

Corresponding to these two stages of dissociation, there

are two dissociation constants, K-, and Kp respectively,

and it is therefore possible to study the effect of both

NH group on the dissociation of Both the base and its salts are of

40
importance as pharmaceuticals, the- salts especially being

used in veterinary medicine as anthelmintics. Electromotive force measurements were made, using two cells which can "be represented

PtjH

2

, Pz.2HCl(m

1

), Pz-HCl(m

2

), AgCl;Ag (26) and Pt;H

2

, Pz-HCl(m

2

), Pz(m ), AgCljAg (27) The emf E of cell 26 is related to the first dissociation

constant (K-, ) of diprotonated piperazine by the equation

(E-E°)/k + log m

cl

_ + log (m

1 /m

2

) = pKl - log (7Pzh 2++ .7c1 -/yPzH+ ) (28) where k stands for (RT In 10)/F and E° is the standard emf of the hydrogen-silver chloride cell in water.

A quantity pK^ can be defined as follows:

l /p

pK^_ = (E-E°)/k + log m

cl _ + log (m1 /m2

) - 4AI ±/c: (29) where A is the slope parameter of the Debye-Huckel equation. Then the relationship

pK^ = pK

1 + bl (30) holds to a good approximation and pK-, can be obtained by extrapolating values of pK* to zero total ionic strength

(i), using the method least squares. An extrapolation of the data at 25 °C is shown in figure 10.
pK

Figure 10. Extrapolation curve for obtaining pK-, from

the values of pK-I . Cell 27 was used for the determination of the second dissociation constant, the equilibrium constant for re- action 25. The emf of this cell is related to pKp by the equation

PK

2

- log ( YpzH+ . Toi_/Vpz ) (31 42
A small correction was made to the buffer- ratio term, log (rru/rruj, because of the slight hydrolysis of the base. The function used to obtain the true value of pKp by extrapolation of pKp values to zero ionic strength was, therefore

m?+mf)H 1 /P pK

2 = (E-E°)/k + log m

cl - + log ——^ - 2AI X/ ^ (32)

Measurements were made over a range of temperature

from to 50 °C. The variation of both pK-, and pK

p with temperature is considerable; that for pK-, is shown in figure 11. Thermodynamic quantities for the dissociation

of piperazinium ions and certain of their analogues at

25 °C, derived from emf measurements made in this section, are given in table 7-

Table 7. Comparison of thermodynamic constants for the dissociation of piperazinium, morpholinium,

piperidinium, and pyrrolidinium ions at 25 °C. pK AH° AS° ac;

J mol

-1

J deg

-1 ,-1 mol

J deg

-1 mol" 1 Piperazinium ion First stage 5.333 31,080 + 2.2 87 Second stage 9.73x 42,870 -42.5 73 Morpholinium ion 8.492 39,030 -31.7 48 Piperidinium ion 11.123 53,390 -33.9 88 Pyrrolidinium 11.305 54,470 -33.7 68

43
25 TEMPERATURE, Figure 11. Variation of pKn for piperazinium ion with

temperature. (Circles are experimental

points; the line represents the function

found by least squares treatment of the

data.

)

The values of pK-, and pK

? for piperazinium -ions, namely 5.333 and 9.73-, at 25 °C, are to he compared with

8.492 for morpholinium ion at the same temperature. Thus the charged NH p group in monoprotonated piperazine

This is to be expected, for one positively

group should facilitate the dissociation of morpholine.

charged NH 2 the other NHp"

1" group. It may be observed, however, that

an oxygen atom in the heterocyclic ring is more powerful

in its effect on the dissociation of the NHp group than

is an NH group in the ring, pK for morpholinium ion (8.492) being smaller than pK„ for piperazinium Ion (9.73-1 )• (H. B. Hetzer)

44
Dissociation Constants of Substituted Anilines

and Phenols

The dissociation constants of substituted anilines

and phenols, and particularly of the nitro-substituted

compounds, are of interest to the Section, partly in its search for new indicators and partly in the study of the effect of structure on the dissociation equilibria. Data

for 4-chloro-2-methylaniline, 2, 4-dibromoaniline, 2,6- dibromoaniline, 2-methyl-5-nitroaniline, 4-methyl-2-

nitroaniline, and 4-methyl-j5-nitroaniline and the corresponding disubstituted phenols have been obtained [39]

The results are shown in figure 12, along with some data for the six dichloroanilines and the six dichlorophenols and a number of monosubstituted anilines and phenols. It is evident that, with some exceptions, there

is a linear relation between the pK value of a substituted

aniline (pit.) and that of the corresponding phenol (pKp). This linearity might be useful in providing at least an approximation to the unknown value of a substituted aniline

if that of the corresponding phenol were known, and vice versa .It is, however, the exceptional cases in the figure which are of particular Interest. Linearity fails for

j?-methoxy, o_-methyl, and o_-nitro monosubstitution. Among

the disubstituted compounds, it fails for the 4-chloro-2-

methyl, the 2-methyl-5-nitro, and the 4-methyl-2-nitro

compounds. In two of these there is an o_-methyl group and in the third an o_-nitro substuent. Thus any anomalies in

the monosubstituted compounds carry over, more or less unchanged, to the disubstituted compounds.

45
2-CH3 -5-N0o

D2,6-DiCI • 4-CH 3-2-N02 •2,6-DiBr

pK p Figure 12. Relationship between the pK of substituted protonated anilines (pKA ) and that of the corresponding substituted phenols (pKp).

46
As part of a study of the medium effect, the pK

values of a number of phenols have been determined in 50 wt percent methanol [40,4l]. The pK value is higher in every

case in 50 percent methanol than it is in water, as the following results for ApK, that is pK(MeOH-H

20)-pK(H

20), show: _o-Nitrophenol ApK = = 0.687 m-Nitrophenol 0.615 p-Nitrophenol 0.536 4-Methyl-2-nitrophenol O.687

2, 6-Dichlorophenol O.889 Pentachlorophenol 0.458 The pK value is, on the average, 0.645 unit higher in

50 percent methanol than in water, but the difference is specific to each substituted phenol.

(R. A. Robinson)

C . Isopiestic Vapor Pressure Measurements

As part of a program carried out in collaboration with

the Office of Saline Water, thermodynamic properties of the following aqueous systems have been determined and the results published: sodium chloride - barium chloride potassium chloride - barium chloride sodium chloride - calcium chloride magnesium chloride - calcium chloride Data for the system: potassium chloride - calcium chloride

are now being calculated. The support of the Office of Saline Water was terminated November 15, 1966.

(R. A. Robinson)

47
D. Spectrophotometrlc Studies in N-Methylpropionamlde Acid-base equilibria have been studied spectrophotometrically in N-methylpropionamide at 25 °C for a number

of indicator substances of different charge types. Accurate measurements have been made on p-bromoaniline,

N,N-dimethyl-p-bromoaniline, m-nitroaniline, 4-chloro2, 6-dinitrophenol, dimethylpicric acid, and metacresol

purple. The equilibrium constant at infinite dilution in N-methylpropionamide can be obtained from these measurements. The detailed analysis of results has not yet been

completed. Comparison of results with those for aqueous solutions should give useful insights into the nature of solvation in N-methylpropionamide, medium effects, and the

effect of dielectric constant.

(Donald Rosenthal)

48
5. CONDUCTANCE AND TRANSPORT BEHAVIOR OP ELECTROLYTES

A. Objectives and Approach

Transport properties of electrolytes in solution, and electrical conductance in particular, provide information

on the behavior of Ions that is not obtainable from equi- librium thermodynamics. Practical electrochemical operations, including many analytical methods, must necessarily

be carried out under non-equilibrium conditions and may

be limited by the rates of the fundamental chemical and physical processes. Conductance measurements can contribute to an under- standing of non-equilibrium properties of electrolyte

systems in two ways. First, from a phenomenological

approach, such measurements help to define transport, or friction, coefficients. Although the formalism and interpretation of such coefficients is not standardized

In practice, it is recognized that they provide a unique,

quantitative measure of the pair-wise interactions among

the species in a solution as a function of concentration. These coefficients are calculated from an appropriate

combination of conductance, transference, and diffusion

data. Second, from a mechanistic approach, conductance measurements and current theories concerning them yield estimates of the hydrodynamic properties of ions. An appreciable discrepancy between the Stokes radius of an Ion and other estimates of its effective size is usually

an indication of specific solvent effects. Conductance data can also provide thermodynamic data, particularly with respect to ion-pair equilibria. In many nonaqueous solvents these results are not readily

obtainable by other means, and in any case the conductance

values provide confirmation of thermodynamic data.

49
A study of conductance cell design has been initiated, using parallel-plate electrodes separated by cylindrical

quartz or glass spacers which form the body of the cell.

It is anticipated that a wide range of cell constants

can be provided by a small selection of interchangeable

components. The use of very thin spacers has already

proved useful in making cells with low constants and very

small volume for monitoring distillations and conducto- metric titrations. If the constants of such demountable

cells prove to be sufficiently reproducible, it may be possible to construct absolute cells in which the con- ductance of a solution is measured in absolute units on

the basis of linear dimensions or capacitance of the empty cell.

(T. B. Hoover)

B. Analysis of Errors of the Jones-Dike Conductance Bridge

A commercial Wheatstone bridge based on the design

principles of Grinnell Jones [42] is widely used for precise conductance measurements because of its convenience

and reliability. Within the operating limits set by the manufacturer [4;?], namely frequencies below 2 kHz and resistances less than 60,000 ohms, the claimed limit of error of less than 0.02 percent appears reasonable. In

order to reduce the effects of electrode polarization,

however, it is frequently desirable to extend measurements

to higher frequencies. This is especially true when platin- ization of the electrodes may lead to undesirable adsorption

or catalytic reactions.

The Jones-Dike bridge has also been used in this Section for audio-frequency dielectric constant measurements when the cell capacitance is of the order of 1000 pP.

50
In this application also it is desirable to carry the measurements to higher frequencies in order to eliminate

electrode polarization effects. In earlier conductance

studies anomalous effects were observed at frequencies

greater than 8 kHz with our present bridge. Accordingly,

a systematic study was undertaken to identify the sources

of error and determine to what extent they could be reduced. The experimental approach was based on an unpublished

theoretical analysis due to C . G. Malmberg and outlined

below.In this treatment of the simple Wheatstone bridge, the ratio arms are represented by subscripts 1 and 2, and the unknown and measurement arms by J> and 4, respectively. Junctions 1-4 and 2-3 correspond to the high potential

corners of the bridge. Each arm is represented by a resistance and capacitance in parallel, so that the admit- tance of arm j is given by Y . = l/R . + i«C.. The balance condition Y-,Y^ = Y^Y^ is solved with the following approxiR

1C2 ^ R

4

C 3 The difference between distributed capacitances of the ratio arms is given to a very close approximation by

R

1R 4

c

51

(33)
By placing sufficiently large resistors and capacitors in arm 3 the difference between FU and R^ can be made signi- ficant and a very small difference between the ratio arms will be revealed. Measurements were made at 5 to 50 kHz, with R^ = 50,000

and 100,000 ohms and C^ from 600 to 1400 pF. Allowance

was made for the losses of the capacitors in calculating

the effective parallel resistance in arms 3 and 4. The

2 results showed (R, - R^) nearly proportional to C (rather than to CU), and the external capacitors had much greater apparent loss in arm 4 than in arm 3. The internal

circuit of the bridge was examined and found to have been

changed from the design specified in reference [45] . The slide wire and end coil (Rp and R^ of figure 3 in reference

[43]) were shunted by an additional 15-ohm coil. The values

of R

2 and R^ were adjusted so that the effective range of

the R

2 dial was -0.05 to 0.15 ohm, but with R

2 set at zero

the additional resistance in the measurement arm was nearly

13 ohms. This amount was exactly compensated by an ad- ditional coil (Rh

) in the adjacent (unknown) arm in series with X.,-Xp. When the ratio arms are closely matched this modification has no effect on the accuracy of ordinary

conductance measurements. At high frequencies, however,

and with a large capacitance in the unknown arm, the 13-ohm

series resistance acts as a frequency-dependent shunt. Leads from the ends of the new coil, R, , were brought

to binding posts on the bridge panel so that the coil could

be shorted out when desired. Repetition of the tests de- scribed above indicated (Cp - C-, ) was about -2 pF. The introduction of 5-, 5-, and 10-pF silvered mica capacitors

in series with each other and in parallel with the ratio

coil of arm 2 reduced the indicated unbalance to less than

0.5 PP. With 1000 pP and 100,000 ohms in arm 3 the dif- ference between R-, and ~Ru was 0.05 percent at 20 kHz.

52
In conductance studies it is frequently desirable to work outside the normal resistance range of the bridge by connecting a resistor of high quality either in series or

in parallel with the cell. The external switch panel shown

in figure 13 was designed to facilitate the operation and avoid changing lead connections. A reversing switch was included to permit Interchanging the cell connections as

a check on asymmetry of stray capacitance or inductance.

The switches were of an anti-capacity telephone type. A further provision for using either the 2-terminal or 4- terminal mode of cell connection is relatively unimportant

with the Jones-Dike bridge because of the slide-wire shunt device used for fine balance. The fixed resistance of this circuit exceeds the usual cell lead resistance which the 4-terminal operation is intended to eliminate.

Figure 13. Switch panel used with conductance bridge to provide a resistor in series or parallel with

the cell and to reverse the cell leads.

(T. B. Hoover)

53
6. PREPARATION AND PROPERTIES OF SOLVENTS

A. Ethanol (95 Volume Percent) Because of its excellent solvent properties for many

organic compounds, 95 percent ethanol (190 proof) finds widespread use in medical, pharmaceutical, and organic

chemical fields. At the same time, its water content is sufficient to permit reproducible pH measurements with the glass electrode. Consequently, it is an appropriate medium

for investigating solvent effects on electrolytes, both by

potentiometric and conduct ometric methods, and for evaluating pH* standards. Published methods for the purification of ethanol usually apply to the anhydrous solvent. They often start with the water azeotrope in order to avoid benzene which

may be present in industrial absolute alcohol. The most frequently recomme'nded steps, which may be used in various

combinations, are: removal of most of the water with

calcium oxide, prolonged refluxing under nitrogen, treatment with silver sulfate or trinitrobenzoic acid, fractional

distillation, and removal of the last traces of water by treatment with aluminum amalgam or with magnesium ethylate, followed by a final distillation. Ethanol with by far the lowest conductance that has been reported was prepared by mixed-bed ion exchange treatment of otherwise carefully

purified and dried solvent [44]. With a view to simplifying and standardizing the pre- paration of good quality ethanol, several of the above

steps were examined in some detail. Most of them could be eliminated if the deionization treatment were equally

effective for stock 190-proof alcohol. A commercial mixed anion-cation exchange resin was equilibrated, with respect

to water content of the azeotrope, by continuous extraction

54
for seven days. The solvent was distilled from a boiler, percolated through the resin and returned by siphon to the boiler. In this way It was expected that water would con- tinuously accumulate in the boiling flask, together with

any free electrolyte or soluble organics that might be extracted from the resin. By circulating a limited amount

of solvent, it was hoped that the exchange capacity of the resin would not be reduced appreciably. About 40 g of resin (dry weight) containing 60 g of water was treated with 500 cc of 95 percent ethanol. Because of difficulty

in controlling the distribution of liquid between the extractor and the boiler, only 150 cc remained in the boiling flask at the end, and the alcohol in contact with the resin contained 15 percent water. U.S. P. 190-proof alcohol having an initial specific conductance of 6.67 x 10 ' ohm cm

- was passed through a column containing the treated

resin. The conductance of the product was 1.43 x 10 ohm cm . Because the latter value was an order of magnitude greater than desired, attention was turned to the preparation of absolute alcohol. Molecular Sieve 4A was used for the removal of water

from U.S. P. 190-proof alcohol. The capacity was about 20 g of water per 100 g of sieves, so that 3.5 1 (l gal) of alcohol required at least 1100 g of the sieves. After

standing five days until no more bubbles were released from

the drying agent, 2.5 1 of alcohol containing 0.35 percent water was decanted. An additional 350 cc was recovered by room-temperature distillation, but this material apparently

underwent some dehydration and polymerization. In order

to regenerate the sieves, it was necessary to boil them

with repeated changes of distilled water until no more odor

of ethanol was evident before drying at 250 °C

.

55
The decanted alcohol, which contained considerable

fine particulate matter from the sieves, was held at reflux temperature for 20 hr while argon was "bubbled through

the liquid continuously. Distillation through an efficient, glass-packed column gave a product (A) having a specific

-7 -1 -1 conductance of 1.0 x 10 ohm cm . Distillation from

0.05 percent trinitrobenzoic acid gave a conductance of

1.6 x

10~ 8

. Refluxing product A (above) over magnesium ethylate

and distillation under argon atmosphere gave material with

a water content less than 0.01 percent and specific con- ductance of 3.08 x 10 ohm cm . Amalgamated aluminum

foil seemed relatively ineffective as a drying agent because of the difficulty in keeping an active surface during

the reflux period.

B. N-Methylpropionamide

N-Methylpropionamide (MP), a solvent with a high dielectric constant, has been used in the Section for con- ductance, viscosity, and kinetic measurements. During the

past year the material has been characterized more fully

in terms of its physical properties. Densities and dielectric constants have been measured from the freezing

point (-30 °C) to 150 °C. All samples used for the recent measurements were reclaimed by fractional distillation

from old stock, but several different distillations showed

no difference in properties. The water content by Karl Fischer titration was 0.002 percent. The dielectric con- stant was remeasured on one sample that had been prepared

and measured six years ago. The agreement with both the original value and the recent series was excellent. Except

for the range 20 to 40 °C, all measurements were made by

C. G. Malmberg.

56
Densities were determined in two dilatometers and checked between 20 and 40 °C with two different pycnometers.

All results are represented within . 005 percent by the equation

0.95069 - 8.136 x 10 H

t + 2.921 x 10

- 1.79 x 10" 9 X? (34) where t is the temperature in °C. Dielectric constants were determined in three cells, one absolute and the other two relative. One of the relative cells, shown in figure 14, had parallel plate

electrodes; the other had concentric cylindrical electrodes.

Figure 14. Cell used for measurements of relative dielectric constants in NMP and NMP-cyclohexane mixtures.

57
Both relative cells were calibrated with air, nitrobenzene,

and water. Measurements were made both with a Jones-Dike

bridge and with a General Radio capacitance bridge. The results are represented within 0.2 percent by

log1Q D = 2.36486 - 5.043 x 10~^ t + 9.846 x 10 t

2

- 1.915 x 10" 8

t 3 (35) where D is the dielectric constant at t °C. The binary system NMP-water is of interest because

both components are highly polar hydrogen-bonding liquids

of different structures. Water appears to be characterized

by ice-like pseudo-crystalline aggregates, while the pro- perties of NMP are best explained as due to long chain-wise associations of the molecules. Densities and dielectric

constants were measured over the whole composition range

and from 20 to 40 °C. The excess molar volumes shown

in figure 15 are typical of mixtures of hydrogen-bonding

liquids with water, but the minimum is about 10 percent

lower than has been reported for any other aqueous system. The dielectric constants show similar large negative deviations. In figure 16 deviations from a linear mole fraction

mole fracThe polarizations, P, calculated for this system are characteristic

of strong associations. The P vs. l/T curve has a large negative slope, but the slope is not quite as great as has been found for methanol or N-methylacetamide . These results

suggest that NMP breaks up the ice-like structure of water

in favor of the closer-packed monomer structure. On the other hand, water forms hydrogen-bonded crosslinks between

NMP chains, resulting in a closer-packed structure with a smaller resultant moment.

58
EXCESS MOLAR VOLUME

MOLE FRACTION WATER

Figure 15. Excess molar volume ( V

1 2~N

1V1 _N2V2^ of NMP-water mixtures at 30 °C as a function

N-Methylpropionamide is miscible with many organic

liquids, raising the possibility of a series of binary

mixtures isodielectric with water at some one temperature

but quite different in other properties. Such a series should help to separate electrostatic from chemical effects

in ion-solvent interactions. In order to identify such a series, dielectric constants were measured for mixtures of NMP with cyclohexane, carbon tetrachloride, decalin, and p-xylene. Normal hydrocarbons appear to be incompletely

59
miscible with NMP. Partial results for one system, NMP- cyclohexane ., are given in table 8. Deviations from a linear mole-fraction relationship are slight. Densities

are not yet available to show whether the system is more nearly ideal on a volume mixture basis.

NMP-WATER - DIELECTRIC CONSTANT DEVIATIONS

MOLE FRACTION WATER

Figure 16 . Dielectric constant deviations (D-, p-ISLD-, _N

pDp) of NMP-water mixtures at 20, 30 and 40 °C as a function of the mole fraction of water (Np

).

60
Table 8. Dielectric constants of N-methylpropionamidecyclohexane mixtures at 25 °C. Dielectric

constant

0.1872

.3197 .4996

.7786

.9522 1.0000

175.65

133.77

109.16 79.34

33.30

5.29 2.015

- 9.38

-10.99

- 9.56

- 7.16

- 5.02

(T. B. Hoover)

61
7. NONAQUEOUS TITRATIONS: DETERMINATION OF HYDROGEN FLUORIDE

An improved procedure for the determination of 5- to 10-mg quantities of fluoride was needed for the assay of calcium fluoride. A pyrohydro lytic procedure was available

for decomposing the sample and converting fluoride quantitatively to hydrogen fluoride. The applicable methods in

the literature fall into two classes: formation of metal- fluoride complexes, e.g . with aluminum or thorium, and conversion to an alkali salt which can be titrated as a base by various techniques. The direct acidimetric titra- tion of HF in a basic solvent offered the possibility of eliminating several steps.

N, N-DImethylformamide (DMF) was selected as a promising

solvent, and sodium methylate in methanol was chosen as the titrant. Potentiometric, conductometric, and indicator

methods for endpoint detection were evaluated. Since water

will accompany HF in the distillation from the sample, it

is important that moderate amounts of water should not Interfere with the accuracy or precision of the method. Conductometric titrations were made in a circulating

cell having a capacity of 20 cc and provided with bright

platinum dipping electrodes. The results of four titrations

of 5-mg quantities of HF had a standard deviation of 2 percent. The titrant was 0.2 M sodium methylate in methanol. The slope of the conductance curve changed sign at the equi valence point, but the fore-curve was nonlinear and the principal uncertainty lay In defining the intersection of the fore- and post-curves. Up to 10 percent water in the DMF solvent had no effect on the shape of the titration plot

but at the higher water content a fine precipitate formed

as the equivalence point was approached. In 50 percent DMF-water no break in the titration plot was discernible.

62
Potentiometric titrations were made with glass vs . saturated calomel electrodes using an automatic titrator.

The same titrant as above was added volumetrically from a calibrated 1-cc buret, and 20 cc of solvent was used in each case. The standard deviation of four titrations of 3-mg quantities of HP was j5 percent.

All of these titrations were made in the presence of two drops of a 0.05 percent solution of phenol red (phenolsulfonephthalein) in DMF. Since the indicator color

change seemed to coincide with the instrumental- endpoints, further attention was given to optimizing conditions for

the indicator titration. Sample and titrant were added by weight from hypodermic syringes and the volume of solvent

was reduced to 10 cc. Experience with pyrohydrolysis of fluorspar samples had showed that some tungsten was included in the distillate. To check for possible interference, sodium tungstate was added to some of the samples to give

a ratio of tungsten to fluorine of 0.048.

A statistical design was set up with two solvent compo- sitions (l and 5 percent water) and three different treatments. These were standardization against benzoic acid, titration of HF, and titration of HF plus tungstate. Each combination was run once in a randomized order in five replications of the experiment. Blank titrations were made

separately. The analysis of variance of the results is summarized

in tables 9 and 10. The standard deviation for the de- termination of hydrogen fluoride is consistent with that for the standardizations and corresponds to 1 percent for

a single determination. There was no significant effect

due to the differences in solvent composition. The highly

significant effect caused by tungstate appeared to be

65
Table 9- Analysis of variance of fluoride titrations. Degrees of freedom

Sums of squares

of deviations f Solvent 1 0.00002738 1.44 Tungsten 1 .0023588 124.2 Blocks 4 .00001095 O.58 Error 13 .00001899

Varlate: g NaOMe solution/ g HF solution

Mean: I.7386 Standard Deviation: 0.0044

Table 10. Analysis of variance for standardization of base. Degrees of Sums of squares freedom of deviations

Solvent 1 0.08490 2.44 Error 8 .03487

Varlate: g NaOMe solution/ g benzoic acid

Mean: 48.097 Standard Deviation: O.187

64
directly equivalent to an alkali correction; that is, the titer of the HF was reduced by the amount of sodium added with the tungstate.

Some preliminary trials were made with solvent con- taining 10 and 25 percent water. In both cases a precipitate formed during the titration. The blank correction

for the 10 percent solvent was remarkably large and variable,

and the precision of analysis in the 25 percent solvent was

poor.

No trials were made with actual fluorspar samples, but it is anticipated that the HF distillate could be collected directly in a small volume of DMF rather than

in aqueous sodium hydroxide as has been the practice. If the tungsten is volatilized as tungstic acid, it may be

so insoluble in DMF that It will cause no interference.

(T. B. Hoover)

65
8. FOREIGN-CURRENCY PROGRAMS

The two projects sponsored by the Electrochemical

Analysis Section at universities in India and financed through the provisions of the Public Law 480 continued

active during the reporting year. A brief description

of these programs follows.

A project entitled "Properties of Electrolytes in Nonaqueous Media" is being conducted at Jadavpur University

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

Dr. K. K. Kundu. Considerable work is being done on the properties of weak electrolytes in glycol-water solvents over a range of temperatures. During the past year, a manuscript by U. Sen, K. K. Kundu, and M. N. Das entitled "Standard potentials of the silver-silver chloride electrode

in ethylene glycol and its aqueous mixtures at different temperatures and related thermodynamic quantities" was accepted for publication in the Journal of Physical

Chemistry.

A project entitled "Physicochemical Studies in Non- aqueous Solvents" is being pursued at Panjab University,

Chandigarh, under the supervision of Professor Ram Chand

Paul. Professor Paul's continuing interests Include the behavior of Lewis acids in solvents quite unlike water, ion-pair formation, and the general subject of the solva- tion of electrolytes. Some studies of weak electrolytes

by conductivity methods are also a part of his program.

(R. G. Bates)

66
9. IMPACT OF SECTION'S ACTIVITIES

The impact of the Section's activities can best he judged in terms of the distribution of Section publications, invited talks, and requests for advice and consultation. In this connection, the following statistics, compiled for

the calendar year 1966, are pertinent.

Reprints requested « 1026 Publications distributed

without a request 1105 Manuscripts reviewed for outside journals 55 Project proposals reviewed 5 Significant consultations 2j5 universities 8 industries 8 government 4 research

institutes 3 Committee memberships 16 Publications in 1966 22 Talks given in 1966 18 scientific societies 8 universities 9 government 1

67
10 . PERSONNEL

Roger G. Bates, Section Chief

R. A. Robinson, Assistant Section Chief Marguerite Raudenbush, Section Secretary Electromotive Force and Acidity Measurements

V. E. Bower (left Dec. 1966)

J. T. Clark, Jr. (Sept. to Feb.)

Guest worker, from Clarkson College

of Technology

R. Gary (left July 1966) Hannah B. Hetzer

Maya Paabo

R. A. Robinson

D. Rosenthal (Sept. to Feb.) Visiting Expert, on leave from Clarkson College of Technology

B. R. Staples (reported in May 1967) Conductance Measurements

T. B. Hoover

P. K. Hall (July to Sept.) from the University of Wisconsin, Marshfield Dielectric Measurements

C. G. Malmberg (left Jan. 1967

)

Isopiestic Vapor Pressure Measurements

V. E. Bower (left Dec. 1966 ) R. A. Robinson

Kinetic Studies

R. K. Wolford (left July 1966)
COMMITTEE MEMBERSHIPS

Member, Subcommittee on Hydrogen Ion Determinations

,

ASTM Committee E-l, Sub. 22

Councilor, American Chemical Society,

representing the Division of Analytical Chemistry.

Member, NEC Committee on Symbols, Units, and Terminology for the Physical Sciences

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

Member, U.S. National Committee for IUPAC

Member, Commission on Electrochemistry, IUPAC.

Member, Commission on Symbols, Terminology, and Units, IUPAC. Associate Member, Commission on Electroanalytical

Chemistry, IUPAC. Member, Subcommittee on Physicochemical Symbols

and Terminology, Committee on Physical Chemistry,

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

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

Chimica Acta . Member, Washington Editorial Review Board (NBS). Member, Committee on Units and Usage (NBS). Member, Museum Committee (NBS). Member, Advisory Committee on the Senior Lunch Club

(NBS).

A. Robinson

Secretary, Commission on Electroanalytical Chemistry,

IUPAC

69
12. PUBLICATIONS AND MANUSCRIPTS, JULY 1966 TO JUNE I967

1. Paabo, M., Bates, R.G., Robinson, R.A. Dissociation of Acetic Acid-dj| In Deuterium Oxide from 5 to 50° and Related Isotope Effects.

J. Phys. Chem. JO, 2075-77 (1966).

2. Stokes, R.H., Robinson, R.A. Interactions in Aqueous Nonelectrolyte Solutions.

I. Solute-Solvent Equilibria.

J. Phys. Chem. £0, 2126-31 (1966).

3. Robinson, R.A., Bower, V.E. Properties of Aqueous Mixtures of Pure Salts. Thermodynamics of the Ternary System: Water-Calcium

Chloride-Magnesium Chloride at 25 °C.

J. Res. Natl. Bur. Std. 70A, 305-311 (1966).

4. Robinson, R.A., Bower, V.E. Properties of Aqueous Mixtures of Pure Salts: Thermodynamics of the Ternary System Water-Sodium

Chloride-Calcium Chloride at 25 °C

J. Res. Natl. Bur. Std. J_0A, 313-318 (1966).

5. Hetzer, H.B., Bates, R.G., Robinson, R.A. Dissociation Constant of Morpholinlum Ion and Related Thermodynamic Quantities from to to°.

J. Phys. Chem. JO, 2869-2872 (1966).

6. Paabo, M., Robinson, R.A., Bates, R.G. Dissociation of 4-Aminopyridinium Ion In 50 per cent Methanol-Water and Related Acidity Functions from

10 to 40 °C. Anal. Chem. 38, 1573-1575 (1966).

7. Bates, R.G. Acidity Functions for Amphiprotic Media

Chapter for Chemistry of Non-Aqueous Solutions. Academic Press, New York 1, 97-128 (1966).

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

The Ionization Constant of Deuterium Oxide from 5 to 50°.

J. Phys. Chem. J_0, 3820-24 (1966).

9. Robinson, R.A., Bates, R.G. The Dissociation of Some Substituted Phenols in 50-Percent Aqueous Methanol as Solvent.

J. Res. Natl. Bur. Std. 7_0A, 553-56 (1966) 70
10. Bates, R.G. Proposal for the Practical Measurement of pH In Amphlprotlc and mixed Solvents.

IUPAC Information Bulletin No. 26, pp. 50-54 (1966).

11. Bates, R.G. The Acheson Medalist, Warren C. Vosburgh, His Life

and Career.

J. Electrochem. Soc. 113 , 309C (1966).

12. Schindler, P., Robinson, R.A., Bates, R.G. Mediumeffekte bei der Dissoziatlon von protoniertem

Tris(hydroxymethyl)-amino-methan ("Tris") in Wasser-Methanol Mischungen.

Chimia 20, 447-448 (1966).

13. Bates, R.G.

pH Determination. Encyclopedia of Industrial Chemical Analysis. John Wiley and Sons, Vol. 3, 146-161 (1966)

(P. D. Snell and C. E. Hilton, eds.).

14. Bates, R.G. Hydrogen-Ion Concentration. Encyclopedia of Chemical Technology, 2nd edition. John Wiley and Sons, Vol. 11, 380-390 (1966) (Kirk and Othmer, eds.)

15. Steel, B.J., Robinson, R.A., Bates, R.G, Dissociation Constant of m-Nitrophenol in 50 Wt. Percent Methanol-Water Solvent from 25 to 40 °C and Related Medium Effects.

J. Res. Natl. Bur. Std. 71A, 9-12 (1967).

16. Robinson, R.A. The Dissociation Constants of Some Disubstituted

Anilines and Phenols in Aqueous Solution at 25 °C.

J. Res. Natl. Bur. Std. 71A, 213-218 (1967).

17. Bates, R.G., Covington, A.K. Behavior of the Glass Electrode and Other pH-Responsive Electrodes in Biological Media.

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

18. Bates, R.G. Measurement of pH. Handbook of Biochemistry and Handbook of Chemistry and Physics. (in Press).

71
19. Bates, R.G. Medium Effects and pH in Nonaqueous Solvents Chapter for Solvation Interactions

Marcel Dekker publisher.

20. Robinson, R.A. The Dissociation Constants of Some Substituted Nitrophenols in Aqueous Solution at 25 °C

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

21. Schindler, P., Robinson, R.A., Bates, R.G. Solubility of Tris(hydroxymethyl)aminomethane in Water-Methanol Solvent Mixtures and Medium Effects in

the Dissociation of the Protonated Base.

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

22. Bates, R.G. Hydrogen Ions. Encyclopaedia Brittanica. (in Press).

72
13. TALKS, JULY 1966 TO JUNE 1967

Bates, R.G., "Solvent Effects on the Dissociation of Protonated Bases", Gordon Research Conference on Analytical Chemistry, New Hampton, N.H. August 25, 1966. Bates, R.G., "Nature of the Solvent Effect of Methanol

on the Dissociation of Protonated Tris(hydroxylmethyl)-

aminomethane" (paper by Schindler, Robinson, and Bates), American Chemical Society meeting, New York, N.Y. September 13, i960". Robinson, R.A., "The Ionization Constant of Deuterium

Oxide from 5 to 50°" (paper by Covington, Bates, and Robinson) American Chemical Society meeting, New York, N.Y. September 14, 1966. Bates, R.G., "The Acheson Medalist, Warren C. Vosburgh, His Life and Career", The Electrochemical Society

national meeting, Philadelphia, Pa. October 11, 1966. Bates, R.G., "Acid-Base Behavior in Alcohol-Water

Solvents", Chemistry Lecture Series, University of Michigan, Ann Arbor, Michigan. November 10, 1966. Bates, R.G., "Thermodynamic Aspects of Acid-Base

Interactions in Nonaqueous Media", Analytical Chemistry

Colloquium, University of Illinois, Urbana, 111. November 15, 1966. Bates, R.G., "Solvent Effects on Acid-Base Processes", Analytical Chemistry Colloquium, Purdue University, West Lafayette, Indiana. November 16, 1966. Bates, R.G., "pH and Other Electrochemical Acidity

Functions for Nonaqueous Media", National Capital

Section, The Electrochemical Society, Georgetown

University, Washington, D.C. December 1, 1966. Bates, R.G., "Acids and Bases in Alcohol-Water Solvents", Passaic Valley Section, American Chemical Society, Montclair State College, Montclair, N.J. January 12, 19 67 Bates, R.G.. "Potentlometric Measurements and pH Standards , Course In Advanced Laboratory Instrumentation, Walter Reed Army Institute of Research, Washington, D.C. March 30, 19 6j

73
11. Bates, R.G., "Medium Effects on Ionic Processes In Nonaqueous Solvents", Symposium on Electrochemistry I Nonaqueous Media, The Electrochemical Society, Dallas, Texas. May 10, 19 6j

.

12. Bates, R.G., "inner Reference Electrodes and Their

Characteristics", Conference on Intracellular Glass Microelectrodes, University of Montreal, Montreal,

Canada. May 24, 1967.

13. Bates, R.G., "ionic Processes and Energies of Transfer from Water to Nonaqueous Media", Chemistry Seminar, McGill University, Montreal, Canada. May 26, 1967.

74
14. LIST OF REFERENCES Electrochemical Analysis: Studies of Acids, Bases , and Salts by Emf , Conductance, Optical, and Kinetic Methods. July 1965 to June 19 bb

.

(R. G. Bates, Ed . ), NBS Technical Note 400 (September 6, 1966)

.

Bates, R.G.

J. Res. Natl. Bur. Std . , 66A, 179 (1962). Bates, R.G. and Robinson, R.A. "Chemical Physics of Ionic Solutions"

(B. E. Conway and R. G. Barradas, Eds.), p. 211. John Wiley and Sons, Inc., New York, 1966. Bates, R.G. "Medium effects and pH in nonaqueous solvents". Chapter in Solvation Interactions (J. F. Coetzee

and CD. Ritchie, Eds.). Marcel Dekker.

In press. Br^nsted, J.N.

Rec. Trav.. Chim., 42, 718 (1923).

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

J. Am. Chem. Soc, 87, 415 (1965). Bates, R.G., Paabo, M., and Robinson, R.A.

J. Phys. Chem., 67, 1833 (1963). Electrochemical Analysis: Studies of Acids, Bases , a'nd Salts by Emf, Conductance, Optical, and Kinetic Methods. July 1964 to June 1965 - NBS Technical Note 271 (September 6, 1965)

.

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

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

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

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

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

J. Phys. Chem., 69, 2750 (1965). Bates, R.G. and Guggenheim, E.A. Pure Appl. Chem., 1_, 163 (i960). Kolthoff, I.M. "Acid-Base Indicators", p. 253. The Macmillan Co., New York, 1937.

[14] Bower, V.E., Paabo, M., and Bates, R.G.

J. Res. Natl. Bur. Std., 65A, 267 (l96l).

75
[15] Bates, R.G.

Ann. N.Y. Acad. Sci., 92, 34l (1961).

[16] Bates, R.G.

Ann. N.Y. Acad. Sci., ±_33, 25 (1966).

[17] Ann. N.Y. Acad. Sci., 133, 251, 259 (1966).

[18] Lyman, J. and Fleming, R.H.

J. Marine Res., 3, 13^ (1940).

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

J. Phys. Chem., _64, 188 (i960).

[20] Salomaa, P., Schaleger, L.L., and Long, P. A.

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

[21] Bates, R.G. "Determination of pH" , Chap. 7. John Wiley and Sons, Inc., New York, 1964.

[22] Born, M.

Z. Physik, 1, 45, 221 (1920).

[23] Bjerrum, N. Kgl. Danske Vidensk. Selskab., 7, No. 9 (1926).

[24] Parker, A.J. Quart. Rev. (London), 163 (1962).

[25] Coetzee, J. P., Padmanabhan, G.R., and Cunningham, G.P. Talanta, 11, 93 (1964).

[26] Bruckenstein, S. and Saito, A.

J. Am. Chem. Soc, 87, 698 (1965).

[27] Grunwald, E. and Price, E.

J. Am. Chem. Soc, 86, 4517 (1964).

[28] Paabo, M., Bates, R.G., and Robinson, R.A. Anal. Chem., 37, 462 (1965).

[29] Akerlof, G.

J. Am. Chem. Soc, _52, 2353 (1930).

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

Bates, R.G.

J. Res. Natl. Bur. Std . , 60A, 263 (1965).

[31] Pranks, F. and Ives, D.J.G.

Quart. Rev. (London), 20, 1 (1966).

[32] Izmailov, N.A. Dokl. Akad. Nauk SSSR, 12J_, 104 (1959); 1^9, 884, 1103, 1364 (1963).

76
Bennetto, H.P., Feakins, D., and Turner, D.J.

J. Chem. Soc, 1211 (1966). Tomkins, R.P.T., Thesis, Birkbeck College, University of London,

1966. deLigny, C.L. and Alfenaar, M. Rec. Trav. Chim., 84, 8l (1965). Dawson, L.R., Sheridan, R.C., and Eckstrora, H.C.

J. Am. Chem. Soc, _6_5, 1829 (l96l). Bates, R.G. and Bower, V.E.

J. Res. Natl. Bur. Std., 57, 153 (1956). Hetzer, H.B., Bates, R.G., and Robinson, R.A.

J. Phys. Chem., 6j_, 1124 (1963). Hetzer, H.B., Bates, R.G., and Robinson, R.A.

J. Phys. Chem., 70, 2869 (1966). Robinson, R.A.

J. Res. Natl. Bur. Std., 71A , 217 (1967). Robinson, R.A. and Bates, R.G.

J. Res. Natl. Bur. Std., 70A, 553 (1966). Steel, B.J., Robinson, R.A., and Bates, R.G.

J. Res. Natl. Bur. Std., 71A , 9 (1967). Jones, G. and Josephs, R.C.

J. Am. Chem. Soc, 50, 1049 (1928). Dike, P.H. Rev. Sci. Instr., 2, 379 (1931). Felici, N.J. Brit. J. Appl. Phys., 15, 801 (1964).

77

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