UNITED STATES DEPARTMENT OF COMMERCE • C. R. Smith, Secretary

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

Acid-Base Behavior in

Aprotic Organic Solvents

Marion Maclean Davis

Institute for Materials Research

National Bureau of Standards Washington, D.C. 20234

U^S^ , National Bureau of Standards.Monograph 105 .

r » »

Issued August 1968

For sale by the Superinlendent <if Documents. U.S. Government Printing Office Washington, D.C. 20402 - Price $2.25
NOV 2 9 1368

1^4 2 5 '46

Library of Congress Catalog Card No. 67-62078
Foreword

During the past 50 years the American chemical industry has made available for common uses

a great variety of organic solvents, in some of which acids and bases behave very differently than

in water. For example, the order of relative strengths of a series of acids or bases may be altered

by a change of solvent. This is especially evident when acid-base behavior in water is compared

with that in hydrocarbons and halogenated hydrocarbons. To the latter two groups, which are often called "inert" or "aprotic" solvents, belong important liquids like benzene, toluene, cyclohexane, and carbon tetrachloride. Industrially important materials such as drycleaning solvents,

lubricants, motor fuels, refrigerants, and transformer oils are additional examples.

From 1941 to the end of 1965 the National Bureau of Standards maintained, in response to requests from industry and other Government agencies, a research program designed to ascertain

and explain acid-base behavior in aprotic organic solvents, as well as to develop methods and

reference materials for determining total acid and base content and relative strengths of acids and

bases in such media. Supplementing the experimental program, a record was maintained of rele- vant work being performed in other laboratories. The Office of Naval Research and the Air Force

Office of Scientific Research contributed financial support.

This Monograph consists of a comprehensive, critical survey of present knowledge in this area, based on numerous kinds of chemical and physical measurements. More than 90 tabulations of

data are included. The conclusions reached have broad scientific applicability, and the Monograph

will be a useful guide both for those concerned with practical chemical programs and those en- gaged in fundamental researches. Its detailed discussion of the varied types of hydrogen bonds

which form in aprotic media will be valuable to scientists attempting to develop an improved theory

of hydrogen bonding.

A. V. ASTIN, Director.

iii
Contents

Page

1. Introduction 1

2. Acid-base concepts 2

2.1. Present-day concepts 2

2.2. Formulation of acid-base reactions in ionizing solvents 2

2.3. Formulation of acid-base reactions in nonionizing

solvents 3

2.4. Hantzsch's views 4

2.5. The role of the solvent in ionization 4

3. Classification of organic solvents 5

3.1. Dielectric constants of pure solvents 5

3.2. Dielectric constants of mixed solvents 6

3.3. Br0nsted's eight classes 8

3.4. Comparisons of aprotic solvents 9

3.4.1. Aromatic solvents as bases 9

3.4.2. Unsaturated aliphatic hydrocarbons as bases 11

3.4.3. Halogen compounds as bases 12

a. Alkyl halides 12

b. Carbon tetrachloride 13

3.4.4. Carbon tetrachloride as an electron acceptor 13

3.4.5. C-H containing compounds as proton donors 13

a. Chloroform and related solvents 13

b. Acetylenic hydrocarbons 14

c. Other studies of C-H containing solvents 16

3.4.6. Chemical reactivity of halogenated hydro- carbons 16

3.4.7. "Dipolar aprotic" solvents 17

3.4.8. Concluding discussion of solvents 18

4. The role of hydrogen bonding in aprotic media 19

4.1. Self-association of nitrogen-containing bases 19

4.1.1. Aliphatic amines : 20

4.1.2. Aromatic amines 21

4.1.3. Amidines, guanidines, and miscellaneous hetercyclic bases 22

4.2. Self-association of acids 24

4.2.1. Phenols 25

a. Self-association of phenol in aprotic solvents 25

b. Hindered phenols 25

c. Intramolecular hydrogen bonding in phenols 28

d. The deuterium isotope effect in phenols 30

4.2.2. Carboxylic acids 31

a. General discussion 31

b. Dimerization constants 31

c. Hetero-association 34

d. Solvent effects on dimerization 35

e. Chelated acids 35

f. Thermodynamic constants 36

g. Possible equilibrium between open-chain and

cyclic dimers 37

h. Thiolobenzoic acid 37

4.2.3. Nitric acid 37

4.2.4. Phosphorus-containing organic acids 38

a. Nature of aggregates formed 38

b. Solvent effects 38

c. Monomer-dimer equilibrium constants 39

4.2.5. Sulfinic and sulfonic acids 40

a. Sulfinic acids, R-S(= 0)0H 40

b. Sulfonic acids, K-SO2OH 40

4.2.6. tert. -Butyl hydroperoxide and deuteroperoxide.

Deuterium isotope effects on hydrogen bonding by Br0nsted acids 40

4.3. Hydrogen-bonded ion pairs 41

4.3.1. Conductance 41

a. Walden's work 42

b. Wynne-Jones' postulate 42

c. Investigations of Kraus, Fuoss, and associates... 43

d. Summary 45

4.3.2. Dielectric constants 46

a. Dielectric polarization of substituted ammonium

picrates and halides in benzene 46

b. Dielectric polarization of other acid-base complexes in benzene 47

4.3.3. CoUigative properties 49

a. Cryoscopy 50

Page

b. The differential vapor pressure (DVP) method... 51

4.3.4. Spectral absorbance in the visible and ultraviolet regions 52

a. Bromophthalein magenta 53

b. Halogen derivatives of phenolsulfonephthalein... 58

c. Nitrophenols 58

d. a-Naphthol, jS-naphthol, and phenol 61

4.3.5. Spectral absorbance in the infrared 62

a. Measurements of crystals....- 62

b. Results using dry aprotic solvents 64

c. Results with wet aprotic solvents 73

d. Further discussion 75

4.3.6. Nuclear magnetic resonance spectroscopy 76

a. Measurements of crystals 76

b. Studies of solutions 76

4.3.7. Concluding discussion of hydrogen-bonded ion pairs 77

4.4. Hydrogert bonding of cations to neutral proton acceptors 79

4.4.1. Introduction 79

4.4.2. Solvated cations 79

a. Solvated hydronium ions 79

b. Miscellaneous solvated cations 81

4.4.3. Homoconjugate cations 82

a. Studies using aprotic or dipolar aprotic solvents... 82

b. Studies of solid salts of N-oxides 85

4.4.4. Miscellaneous heteroconjugate cations 87

4.5. Hydrogen bonding of anions to neutral proton donors 88

4.5.1. Introduction 88

4.5.2. Homo- and heteroconjugation of OH ' ions 89

4.5.3. Homoconjugate aliphatic and aromatic car- boxylate ions 89

a. Solid acid carboxylates 89

b. Homoconjugation of carboxylates in solutions 91

4.5.4. Homoconjugate phenolate ions 97

a. Examples of homoconjugation 97

b. Examples of non-homoconjugation 100

4.5.5. Homoconjugation of halide ions 102

a. Homoconjugation in solutions 102

b. Homoconjugation of X~ in solids 104

4.5.6. Homoconjugation of other organic and inorganic

acids

a. Perchloric acid 105

b. Other acids 105

4.5.7. Heteroconjugate anions 105

a. Chemical behavior 105

b. Physical measurements 108

4.6. Concluding discussion of hydrogen bonding 112

4.6.1. Definition of a hydrogen bond. Systems which

form hydrogen bonds 112

4.6.2! Some recent concepts of hydrogen bonding 113

5. Acidity and basicity scales in aprotic organic solvents... 114

5.1. Early attempts to determine relative acidities 114

5.2. Log Kbha scales of acidity and basicity 115

5.3. Methods used in determining values of log A^bha 116

5.4. Tables of log ^bha, and AS obtained using aprotic solvents 118

5.4.1. Aromatic and substituted aromatic hydrocar- bons 118

a. Benzene 118

b. Anisole and chlorobenzene 123

c. Miscellaneous aromatic solvents 125

5.4.2. Aliphatic hydrocarbons 126

5.4.3. Carbon tetrachloride 126

5.4.4. Chloroform 128

5.4.5. Miscellaneous aprotic solvents 129

6. Acid-base titrations in completely aprotic solvents 133

6.1. Introduction 133

6.2. Titrations with indicator dyes in completely aprotic solvents 133

6.3. Instrumental procedures for detecting end-points in completely aprotic solvents 133

6.3.1. Conductance titrations 133

iv
Contents — Continued

Page

6.3.2. Dielectrometric titrations 133

6.3.3. Photometric titrations 134

6.3.4. Thermometric titrations 135

6.3.5. Titrations by the DVP method 135

6.3.6. Cryoscopic titrations 135

6.3.7. Other possible instrumental procedures 135

a. Refractive index 135

b. Density 135

c. Optical rotation 135

6.4. Reference acids and bases for aprotic organic

solvents 135

6.4.1. Acids 135

a. Hydrogen chloride 135

Page

b. p-Toluenesulfonic acid 135

c. Picric acid 135

d. Trichloroacetic acid 135

e. Diphenyl phosphate 135

6.4.2. Bases 135

a. Aliphatic amines 135

b. Amidines 136

c. Derivatives of guanidine 136

7. List of references 137

8. Author index 145

V

Acid-Base Behavior in Aprotic Organic Solvents

Marion Maclean Davis

A unified picture of acid-base behavior in aprotic organic solvents is presented, based on an

extensive survey of the Hterature 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 acids upon which attention is centered

are proton-donor compounds that are measurably ionized in water, such as aliphatic and aromatic

carboxylic acids, substituted phenols, and mineral acids. The bases of principal interest are likewise compounds capable of forming ions in water, for example, aliphatic and aromatic amines and derivatives

of guanidine or pyridine. The solvents emphasized are hydrocarbons and halohydrocarbons, 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 hydro- gen Donding of these types affects such properties as the absorption spectrum of a salt, the catalytic

effect of an acid, and the accurate location of a titration endpoint.

Key Words: Acid-base behavior, acidity and basicity scales, aprotic organic solvents, hydrogen bonding, titrations

1. Introduction

At the present time increasing attention is being

given to physicochemical investigations of acidbase behavior in organic solvents of comparatively

inert character. Such solvents include aliphatic and

aromatic hydrocarbons and halogenated hydrocarbons, which commonly have dielectric constants

(e) in the range 2 to 10 and are almost devoid of

acidic or basic properties. A second group of

solvents of somewhat higher dielectric constant

(21 to 36) is also comparatively inert in character.

Examples of the latter group are acetone, acetonitrile, and nitrobenzene. (Di)methyl-sulfoxide

(e = 46) is sometimes included in the second group. Having only feeble acid-base properties, these

two groups of solvents do not interact strongly

with acidic solutes such as carboxylic acids,

phenols, and mineral acids, or with basic solutes

like amines and derivatives of guanidine or pyridine.

Consequently, they are not "masking" or "leveling"

solvents like water and low molecular weight

alcohols, but instead are "differentiating" solvents. (Solvents of the first group have also been referred

to as "inert" or "indifferent.") The differentiating

behavior of the two groups of solvents has proved

to be very valuable in chemical analyses and

separations, as well as in fundamental investigations of chemical equilibria and kinetics. The current activity in studying acid-base

systems in these solvents follows a long period of

neglect, which was partly caused by overemphasis

on the role of ionization in chemical reactions.

One consequence of this overemphasis on ionization has been a widespread belief, slow to expire, that acid-base activity is practically nonexistent

in nonionizing solvents like benzene. Later, after

this belief was recognized as being incorrect,

attempts were made to apply the quantitative

acid-base formulations found to be highly satisfac- tory for aqueous solutions to solutions in benzene,

chloroform, etc. The failure of such attempts

discouraged further investigations in this area. An example of beliefs which are now recognized

as erroneous can be found in a noted textbook on

qualitative analysis which was widely used into the 1920's [ir:

"A solution of hydrogen chloride in a poorly

ionizing medium, like benzene or toluene, is an

extremely poor conductor . . . indicating only a

trace of ionization. Such a solution behaves chemically, also, quite differently from the aqueous

solutions of hydrogen chloride with which we are

familiar: dry steel nails, dropped into it, remain

almost unchanged — there is no marked evolution

of hydrogen . . . We find thus ... in the an- hydrous benzene solution an absence of ionization, as indicated by lack of conductivity, and along with

this, a lack of the familiar action of hydrogen

chloride as an acid . . . ."

' All references are listed at the end of the review.

1
Later it was recognized that such tests, using

heterogeneous systems, were not valid criteria of

chemical reactivity, and it was readily demonstrated—for example, by tests using indicator dyes —

that hydrogen chloride and other acids function as acids in benzene and similar solvents. In justice to the noted author of the above cita- tion, it should be pointed out that this textbook

predated the chemical revolution initiated by the newer acid-base concepts of Br0nsted, Lowry,

and Lewis and also predated awareness of the important role of hydrogen bonding in acid-base

behavior. Moreover, in later pages of the same book

the author acknowledged that certain experiments

by others suggested that hydrogen chloride in benzene can react as nonionized molecules, and

he advocated reserving decision as to whether the

observed reactions were caused by nonionized

molecules or by traces of hydrogen ion until after more extensive investigations.

As is well known, the Br0nsted-Lowry concept of

acid-base relationships paved the way for systematic investigations of chemical behavior in all types of nonaqueous solvents. However, Br0nsted's

equation representing an acid-base reaction,

Ai + B2^A2 + B,, (1)

cannot be appUed without modification to benzene

and similar media. This topic will be discussed

further in section 2. The role of hydrogen bonding

will be reviewed in section 4. Through the combined efforts of many investi- gators, applying modern concepts and varied tech- niques, a great deal has been learned in recent decades about the interactions of hydrogen ("Br0nsted") acids with bases in inert organic solvents. The principal objective of this review is to present

a unified picture of our current knowledge in this area, with examples of the kinds of evidence ob- tained by different types of physicochemical investi- gations. Particular emphasis will be placed on

acid-base behavior in hydrocarbons and halogenated

hydrocarbons. Attention will be centered on sys- tems consisting of strong or moderately strong

protogenic and protophilic compounds — that is, acids and bases for which pKa and pKb values in water are obtainable.

2. Acid-Base Concepts

2.1. Present-Day Concepts

Irrespective of the particular acid-base defini- tions adopted, there is general agreement at the

present time that basic properties in compounds are associated with (a) the presence of lone pair ("n-")

electron donor atoms, such as nitrogen or oxygen,

and also with (b) the presence of delocaUzed

("tt-") electron systems, as in aromatic and

unsaturated aliphatic hydrocarbons. Classes (a) and (b) differ very greatly in basic strength, however,

so that the basic properties of n-electron donors can

be studied very satisfactorily in benzene and other

aromatic solvents. There is also general recognition that the be- havior of hydrogen acids with bases often parallels

closely that of nonprotogenic "acids" like boron

trichloride or sulfur trioxide [2]. This parallehsm can be illustrated by the reaction of tri-ra-butylamine

with hydrogen chloride and with sulfur trioxide in benzene:

BusN + SOs^BusN-SOa, (2) BusN + HCl^BugNHCl. (3) The products of reactions 2 and 3 have almost

identical dipole moments in benzene — 7.2 and 7.17,

respectively [3, 4, 5]. However, tributylammonium

chloride differs from tributylamine-sulfur trioxide

in having the capacity to undergo ionic dissociation

in suitable solvents:

BU3NHCI ^ BU3NH+ + CI- (4)

(ignoring solvation of ions).

Equation 4 indicates a major reason for treating

protogenic and nonprotogenic acids separately.

Additional reasons have been pointed out by various

authors [6, 7, 8, 9, 25, 26]. Appropriately, nonprotogenic acidUke compounds

have come to be known as "Lewis acids," while the

term "Br0nsted acids" has become a synonym for "hydrogen" or "protogenic" acids. There is no

dualism in the case of bases — Lewis emphasized [2] his agreement with Br0nsted's views [10] in regarding organic bases, as well as such ions as halide and

acetate, as being true bases.

2.2. Formulation of Acid-Base Reactions

in Ionizing Solvents

In applying the Br0nsted-Lowry acid-base

concept, Br0nsted [10, 11, 12, 13] formulated acidbase reactions in ionic terms. Doing so has proved

to be satisfactory for reactions in ionizing media,

but a different formulation is needed for acid-base

reactions in inert solvents. To bring out the contrast, let us first recall the essence of how Br0nsted

stated and applied his acid-base concept:

1. The definition of acids and bases is given by

the scheme

2
A (acid) B (base) + H+ . (5) A and B are called a "conjugate pair."

2. Acids and bases can be neutral or ionic, as

illustrated below.

Type of acid

or base Examples of acids Examples of bases

Cationic Neutral Anionic

NH.+ CH3COOH

HOOC.COOH2NC6H4C6H4NH^CH3COO3. Acids and their conjugate (corresponding)

bases (see eq 5) cannot have the same electrical charge.

4. Equation 5 illustrates a "simple" acid-base

equilibrium. A mixture of two simple systems gives

a "double" acid-base system and acid-base equilibrium. Thus,

A,^Bi + H+, (6)

B2 + H+^A2, and (7)

adding eqs 6 and 7, Bj + Ai^Bi + Aa. (8) The reaction between triethylamine and hydrogen

chloride may serve as a practical illustration of the

general eqs 6-8:

HC1^C1- + H+, (9)

Et3N + H+^Et3NH+, (10)

and Et3N + HCl^Et3NH+ + Cl-. (11)

(basea) (acidi) (acid2) (basei)

The equilibrium constant corresponding to eq 11 is K= (Et3NH+)(Cl-)/(Et3N)(HCl), (12)

where the symbols enclosed in parentheses signify the stoichiometric concentration of the species

concerned multiplied by the activity coefficient. {Brackets will be used in place of parentheses

whenever activity coefficients are assumed to be

approximately 1, or to be cancelled.)

We should bear in mind that all such formulations

ignore any acidic or basic role played by the solvent.

2.3. Formulation of Acid-Base Reactions in

Nonionizing Solvents

When acid-base reactions occur in relatively

inert solvents like benzene, the data are found not

to fit an equation patterned after eq 12. To illus- trate, the association of triethylamine with the

phenohc indicator bromophthalein magenta E

(3',5',3",5"-tetrabromophenolphthalein ethyl ester)

in benzene, studied by spectrophotometry in the

visible region, was found to conform instead to eqs 13-14:

EtsN + HIn (the indicator) ^ Et:,NHIn, (13)

K^soen. = [EtaNHIn] /[HInJ[Et3N]. (14)

The equihbrium constant, A^assocn ^ for this reaction

has the value 2.74 X 10^ (25°, molar units) [14, 15]. In general terms, eqs 13-14 can be formulated,

respectively, as B (base) + HA or A (the acid) S (sah), (15)

/^assocn.-[S]/[A][B]. (16)

The reaction between triethylamine and bromophthalein magenta E in benzene is but one of many

that have been studied in hydrocarbon and halo- genated hydrocarbon solvents and found to conform

to eqs 15-16. In such solvents any tendency to react in the sense of eq 11 is usually of neghgible

importance.

A simple way of demonstrating that acid-base

reactions conform to the general pattern of eqs

15-16 is to plot log [S]/[A] against —log [B], as in figure 1. The six sets of points were obtained for

I

' I' ' ' '

I

' ' ' ' I' ' '

I ' ' ' 'I

' ' ' ' I 1.0 -

20 2.5 3.0 3.5 4.0 4.5 50 5.5 - LOG [B] Figure 1. Conformance of acid-base equilibria in mixtures of triethylamine with the isomeric dinitrophenols in benzene,

to the assumed pattern of reaction: B (base) + A {acid) ^ S

(salt). See M. M. Davis, J. Am. Chem. Soc. 84, 3623 (1%2). ,Rgf No. 16.) acid-base equihbria in benzene in which triethylamine was the base and the acid was one of the six isomeric dinitrophenols [16]. It is easily seen that

if the six reactions conform to eqs 15—16 the experimentally obtained points should fall — as they do —

on or near straight Hues with the slope —1. Then,

where log [S]/[A] = 0, - log [B] = log Xassocn.- To measure the relative strengths of acids (bases)

in a comparatively inert solvent hke benzene, a suitable reference base (acid) must be added.

From figure 1 it is apparent that, with triethylamine

as the reference base, 2,6-dinitrophenol is the most

acidic, and 2,5-dinitrophenol the least acidic, of the

isomers. The same relative order of strengths in benzene has been obtained with diphenylguanidine

as the reference base [17].

3
2.4. Hantzsch's Views [18-21]

Hantzsch maintained that a chemical method,

such as the tendency toward salt formation with

indicator bases in an inert solvent, gives the most

logical measure of the relative strengths of acids,

and he estimated the relative strengths of several

acids by semiquantitative measurements of the ex- tent of formation of colored salts with basic indicator

dyes in solvents like benzene and chloroform [18,

19]. Two of his associates, Weissberger and Fasold

[22], hoped to demonstrate the correctness of

Hantzsch's views as applied to salt formation

between trichloroacetic acid and methyl yellow

(p-dimethylaminoazobenzene, /VJV-dimethyl-p-phenylazoaniline) in benzene and chloroform. However, the values obtained for A^assocn. (eq 16) were not constant and showed a trend. With the benefit

of our much more complete knowledge about such

systems, several possible explanations can now be

seen for the failure of the particular system studied

to conform to eq 16. These include the monomerdimer equilibrium of trichloroacetic acid, complex

anion formation, and the cis-trans isomerism of

methyl yellow. A much more favorable system is the

equiliDrium mixture of tribenzylamine, picric acid,

and tribenzylammonium picrate, for which A!^assocn. values were obtained by dielectric polarization

measurements [23] and by spectrophotometry

in the visible region [20]. The results by the .two methods were in excellent agreement. Slightly less accurate association constants had been obtained

earlier for mixtures of amines with bromophthalein

magenta E [14]. In nitrobenzene, ionic dissociation is not negligible. Witschonke and Kraus [24] studied the equilibria involving anilinium, dimethylanihnium, and

pyridinium picrates by the conductance method,

and deduced that two equilibria coexist for these

salts in nitrobenzene:

BHA^BH+ + A-, (17) andBHA^B + HA. (18)

The equilibrium constant corresponding to eq 18 is, of course, the reciprocal of /^assocn. (eq 16). The

ionic dissociation constant and the equilibrium

constant for dissociation of the salt into the component base and acid were both evaluated from the

conductance data. To illustrate, for dimethylanilinium picrate the following values were obtained

(molar units, 25° 2): K^^^ (eq 17)=-4.1 X lO-^-

/^dissocn. (eq 18) = 2.5xl0-4.

There is now an abundance of evidence that strong acids (for example, picric acid) and bases (for example, triethylamine or diphenylguanidine) — and

not merely the weak acids and bases previously

known to associate through hydrogen bonding —

interact according to the pattern of eqs 15-16 when

dissolved in inert solvents, and the stoichiometry

^ Throughout this review temperatures are in °C.

indicated is widely accepted as being correct. However, sometimes the schematic formula of the

salt formed by association of a base B with an

acid HA is written as BH+A", the process of association is referred to as "ionization," and the salt is said to consist of "ion pairs." Other investigators

(including the writer) believe this practice oversimplifies the situation, considering the evidence

(see sec. 4.3) that many, at least, of such salts consist of hydrogen-bonded ion pairs. To indicate

the role of hydrogen bonding, the structure of the

salt may be written as BH+...A-, the process

may be referred to as "partial ionization" or

"partial proton transfer," and the salt may be said

to consist of "hydrogen-bonded ion pairs." Shatenshtein [25], the late N. A. Izmailov [26], and N. D. Sokolov [27] are among proponents of such views.

2.5. The Role of the Solvent in Ionization

Br0nsted wanted to formulate acid-base reactions

independently of any specific solvent, as in eq 11,

and in doing so he deliberately ignored the role of

the solvent in promoting ionization. He was well aware, of course, that ionization of solutes depends

on the dielectric constant of the solvent and on its acidic and/or basic character, also. The oversimplification was sometimes criticized, notably by

Izmailov, who wrote [26]: "Our investigations have

shown that Br0nsted's theory is too schematic and

provides for only one type of chemical interaction

between dissolved acids and the solvent (only the

exchange of protons), with the result that this theory

does not allow the explanation of all the peculiarities

in the effect of solvents, and in particular, does not explain the differentiating action on the strengths

of acids, which is the result of more complex in- dividual interactions in solvents of different types."

Izmailov developed a "theory of the dissociation of

acids and bases" [26], which he formulated in considerable detail. The theory was based on exhaustive

studies by Izmailov and co-workers of the relative strengths of acids in different types of solvents,

using such varied physical methods as cryoscopy,

Raman spectroscopy, and emf measurements. The

essence of the theory is that acids themselves are

not ionized, but instead, the ionizing species consist

of molecular compounds of the acids with the sol- vent. The structures of these molecular compounds

may vary according to the chemical structures of

the solute and solvent, thus accounting for differ- ences in the acid strength of a given solute in different solvents. A number of investigators,

notably KendaU (see ref. [28] and later papers),

have also argued that the ionization process in- volves association of solutes with the solvent. Shatenshtein has presented a complete, concise

discussion of Br0nsted's formulation of acid-base processes, together with his own comments and

criticisms [25]. Davis has pointed out [5, 14-17, 20-21, 29-33]

that results obtained in studying the behavior of

hydrogen acids with bases in organic solvents

4
(mainly benzene) by means of spectrophotometry

in the visible region are in harmony with the evi- dence from other kinds of physicochemical investigations reported in the literature. The accumulated

evidence points very convincingly to an important

and varied role of hydrogen bonding in acid-base

interactions involving hydrogen acids. When the

role of the solvent in acid-base interactions was

ignored, it was difficult to reconcile the Br0nsted

concept with the Lewis concept of acid-base reactions. This can be illustrated by comparing the

reaction of triethylamine with hydrogen chloride as formulated by the Br0nsted scheme with the for- mulation of the reaction of triethylamine with a

typical Lewis acid, boron trichloride (eqs 11, 19).

Et3N + HCl^Et3NH+ + Cl-, (ii)

EtsN + BCI3 ^ EtaN • BCI3 (19)

However, applying what has been learned about the

stoichiometry of acid-base reactions in nonionizing

solvents (eqs 15-16) and taking into account the

part played by an ionizing solvent such as water led

Davis to formulate a scheme [5, 21] by which the

two concepts may be harmonized, as illustrated by

eqs 20-25.

1. The Ionization of an Acid, HA, in Water. H2O + HA ^H20H+...A-, (20)

(base) (acid)

H2OH+...A-+ H2O ^ H2OH+...OH2 + A-,

(base) 1(21)

Combining eqs 20 and 21 gives:

2H20+HA^H20H+...OH2 + A- (22)

2. The Ionization of a Base, R3N, in Water.

R3N +H0H^R3NH+...0H-, (23)

(base) (acid)

R3NH+...OH-+ H2O

(base) ^ R3NH+...OH2 + OH-, (24)

Combining eqs 23 and 24 gives:

R3N+2H20^ R3NH+...OH2 + OH-. (25) In these two sets of equations water represents

the solvents which are capable of functioning both

as a base and as an acid. The first equations (eqs

20 and 23) represent associations which are anal- ogous to Lewis acid-base reactions, while the sec- ond equations (eqs 21 and 24) show separation into solvated ions. The third equation in each set is the sum of the first two. Those who accept these

formulations as valid look upon the primary acidbase reaction as being an association, and formation of separate, solvated ions as constituting a secondary reaction.

Equation 21 is based on the assumption [5] that

in water solutions protons are likely to be associated with at least two water molecules, rather than

just the one molecule implied by the formula H3O+.

This assumption has found added support in work

of the past decade which has led to the conclusion

that H3O+, the hydronium ion, associates with

further water molecules, forming the fairly stable species H3O+ • 3H2O when sufficient water is pres- ent [34]. Likewise, there is strong experimental

evidence in support of eq 24 [5]. In proposing concrete pictures of how water can bring about ionization of acids and bases

through specific hydrogen bonding reactions it is not essential that solvation of anions be considered.

It is obvious, however, that the two sets of equations

would have been still more complete if attention

had also been paid to solvation of A" and OH".

Structures such as HOH ... A" are highly probable.

Important effects of hydrogen bonding on acidbase behavior will be discussed in greater detail

in section 4.

3. Classification of Organic Solvents

In the preceding discussion water and chemically

related solvents like the alcohols have been re- ferred to as masking, leveling, or ionizing solvents. In comparison, benzene and chemically related

solvents have been referred to as indifferent, inert,

differentiating, or nonionizing. In an earlier day

benzene was also characterized as an "associating" solvent.

3.1. Dielectric Constants of Pure Solvents

A more definite basis for classifying solvents is according to the magnitude of the dielectric constant. Table 1 contains available values of the

dielectric constant, e, of eight homologous series of

organic solvents, and table 2 contains the dielectric constants of additional currently important

organic solvents. Water, formerly considered to have an extremely high dielectric constant, has

been overshadowed in this respect by some of the newer solvents, for example, A'^-methylpropionamide (e=176) [39a]. As would be expected, ali- phatic hydrocarbons have the lowest dielectric constants (e =S 2). The most familiar organic solvents — usually

compounds of low molecular weight — may differ 5
Table 1. Values of dielectric constant (°C) for homologous solvents having the general formula normal-CnH2n + iX

n =

H Br CN COCH3 COOH COOC2H5 NH2 OH

0 3.8 (25) 115 (20) 21.1 (21) 58 (16) 7.2 (25) 16.9 (25) 78.54 (25)

1 1.7 (- 173) 9.82 (0) 37.5 (20) 20.7 (25) 6.15 (20) 6.02 (25) 9.5 (25) 32.63 (25) 2 9.39 (20) 27 (20) 18.5 (20) 3.30 (10) 5.65 (19) 6.94 (10) 24.30 (25)

oo 1.61 (0) 8.09 (25) 20 (21) 1 S 4. (901 9 97 (901 S 10 (181 90 1 (9'il 4 7.07 (20) 17 (21) "13.1 (20) 2.7 (20) 4.71 (18) 5.3 (21) " 17.5 (25) 5 1.84 (20) 6.32 (26) ion /oA\ 12.U (ZU) 0 AO ill \

6 1.89 (20) 5.82 (25) 10.4 (20) 2.59 (71) ' 13.0 (25) 7 1.92 (20) 5.33 (25) 2.4 (20) 12.1 (22) 8 1.95 (20) 5.00 (25) ^9.96 (25) 9 1.97 (20) 4.74 (25) 10 1.99 (20) 4.44 (25) 3.55 (20) 8.1 (20)

11 2.00 (20) 4.73 (-9) 3.44 (20) 12 2.01 (20) 4.07 (25) 6.5 (25) 14 3.84 (25) 4.72 (38) 15 3.89 (20) 16 3.71 (25) 3.82 (50) 18 3.53 (30) 3.42 (58)

" Unless indicated otherwise the data were taken from reference [35].

^ 4-Methyl-2-pentanone (methyl isobutyl ketone). From reference [36].

References:

[35] A. A. Maryott and E. R. Smith, Table of Dielectric Constants of Pure Liquids, NBS Circ. 514, Aug. 10, 1951.

[36] C. G. Malmberg, NBS, unpublished data.

appreciably in the value of the dielectric constant depending on the polar groups present (for ex- ample, compare the values of e for CH3COOH,

CH3NH2, CH3OH, CH3CN, and CH3NO2). However, in aliphatic homologous series there is generally a trend toward very low values with increasing carbon content (see tables 1 and 2). The hydrocarbons show a slight trend in the opposite direction

(table 1, second column). From consideration of

the examples given in these tables it is apparent

that if solvents with e ^ 20 were arbitrarily regarded

as having high constants, and those with e < 20 as having low constants, most of the common organic

solvents would belong to the second class. That

is, besides the hydrocarbons and brominated hydrocarbons, the second class would include valeronitrile and higher nitriles, all of the ketones except

acetone and methyl ethyl ketone, all of the ali- phatic carboxyUc acids except formic acid, all of

the ethyl esters, and all of the primary aliphatic

amines, as well as n-butanol and higher alcohols. The dielectric constants of secondary and tertiary

aliphatic amines are still lower than those of primary ahphatic amines.

Schlundt [42] seems to have been the first person

to report that a decrease in the value of the dielectric constant with increasing molecular weight is the general rule for homologous series, though

similar observations had been made still earlier

for a few series. It is not surprising, of course, that the value of € drops off as polar groups become in- creasingly screened by aliphatic groups. The value of e also depends on the location of

polar groups in the molecule. To illustrate, the fol- lowing values of e (22°) were obtained for 1-, 2-, 3-, and 4-heptanol, respectively: 12.1, 9.21, 6.86, and

6.17 [35]. Analogously, values of e (20°) found for twenty of the isomeric octanols [35] range from 10.3

for 1-octanol or isooctanol to 2.87 for 4-methyl-4-

heptanol.

The relation between dielectric constant and

structure has been treated exhaustively by

Smyth [43]. Enough has been said to indicate that, by itself, the magnitude of the dielectric constant does not provide a very helpful basis for classifying solvents.

3.2. Dielectric Constants of Mixed Solvents

Mixed solvents are used frequently in acid-base

titrations and in studying solute-solvent interactions. For instance, "ASTM Medium," which is

6
Table 2. Values of dielectric constants {I °C) of selected organic solvents''

Alcohols Halogenated Hydrocarbons Nitro Compounds

Benzyl alcohol

(io-Butyl alcohol

rer/. -Butyl alcohol (36a|

Cyclohexanol

Ethylene glycol

i'so-Propyl alcohol [36].

Amines

Aniline Benzylamine

Dibenzylamine

Ethylenediamine.

2-Methylpryidine.

Morpholine

Piperidine

Pyridine

Pyrrole

Triethylamine

Ether

Anisole

1.4-Dioxane

Ethyl ether Tetrahydrofuran [37].

13.1 17.7 12.5

15.0 38

19.2

(20)

(25)

(25) (25)

(25)

(25)

6.89 (20) 4.6 (21) 3.6 (20) 14.2 (22) 9.8 (20) 7.33 (25) 5.8 (22) 12.3 (25) 7.48 (18) 2.42 (25)

4.33 (25) 2.21 (25) 4.34 (20) 7.39 (25)

Benzotrifluoride Carbon tetrachloride Chlorobenzene

Chloroform

o-Dichlorobenzene

1.1- Dichloroethane [38].... 1.2- Dichloroethane [38].... Dichloromethane

Hexachloro-l,3-butadiene.

Tetrachloroethylene

9.18 (30) 2.23 (25) 5.62 (25) 4.81 (20) 9.93 (25) 9.90 (25) 10.23 (25) 9.08 (20) 2.55 (25) 2.30 (25)Nitrobenzene... Nitroethane

Nitromethane...

1- Nitropropane.

2- Nitropropane.

Phenols

Hydrocarbons

Phenol....

o-Cresol.. m-Cresol.

p-Cresol..

Benzene

Cyclohexane

Decahydronaphthalene .

Mesitylene

Phenylacetylene

Tetrahydronaphthalene. Toluene

2,2,4-Trimethylpentane.

o-Xylene

m-Xylene

p-Xylene

2.27 (25) 2.02 (25) 2.2 (20) 2.28 (20) 2.98 (25) 2.76 (20) 2.38 (25) 1.94 (20) 2.57 (20) 2.37 (20) 2.27 (20)Miscellaneous

Benzonitrile Carbon disulfide Cyclohexanone

Dimethylformamide [39].... Formamide

A'-Methylformamide [39].... A'-Methylpropionamide [39a] Methyl sulfoxide [40] Tetramethylguanidine [39b] Tetramethylurea [41] Tributyl phosphate

^ Unless otherwise indicated data were taken from reference [35].

References:

[35] A. A. Maryott and E. R. Smith, Table of dielectric constants of pure Hquids, NBS Circ. 514, Aug. 10, 1951. Supt. of Documents,

U.S. Govt. Printing Office, Washington, D.C.

[36] C. G. Malmberg, Natl. Bur. Std.. unpublished data.

[36a] W. Dannhauser and L. W. Bahe, J. Chem. Phys. 40, 3058 (1964).

[37] F. E. Critchfield. J. A. Gibson, Jr., and J. L. Hall, J. Am. Chem. Soc. 75, 6044 (1953).

[38] J. T. Denison and J. Ramsey, J. Am. Chem. Soc. 77, 2615 (1955).

[39] G. R. Leader and J. F. Gormley, J. Am. Chem. Soc. 73, 5731 (1951).

[39a] T. B. Hoover, J. Phys. Chem. 68, 876 (1964). [39b] A. Popov. M. L. Anderson, and R. N. Hammer, Abstrs. Papers 145th Meeting Am. Chem. Soc, New York, N.Y., Sept. 1963,

p. 21N.

[40] J. J. Lindberg and J. Kenttamaa, Suomen Kemistilehti B33, 104 (1960).

[41] A. Liittringhaus and H. W. Dirksen, Angew. Chem. Intern. Ed. Engl. 3 , 260 (1964).

used in standard methods for determining the acid

and base numbers of petroleum products [44], as

well as in determining the acid numbers of marine

and vegetable oils [45], consists of 50.0 percent of toluene, 49.5 percent of isopropanol, and 0.5 per- cent of water, by volume. The dielectric constant

is about 8 [46]. In the area of fundamental investi- gations, Fuoss and co-workers have been investi- gating electrolyte-solvent interactions systematically

since about 1950, using various mixed solvents so

as to gain insight into interactions between sol- vents as well as between solvent and solute [47-49].

Figures 2 to 5 contain curves showing the di1 electric constants of some important binary solvent

I mixtures as functions of the composition. It is clear that the magnitude of e does not vary with

changing composition in a quantitatively predictable way. This is not surprising, considering that the structure of a mixture at any given temperature

will depend on the size, the shape, and the relative orientations of component molecules, as well as the strengths of intermolecular forces of attraction. Additional structural changes will occur, of course, when solutes are added.

By interpolation in the curves, a rough idea of

the dielectric constants of some of the solvent mixtures currently popular can be obtained. For ex- ample, the value so obtained for 4:1 benzenemethanol by volume, a solvent mixture which has

been recommended for titrating carboxylic acids

[50], is about 6.7, assuming the absence of impurities

and negligible volume change on mixing.

7
—-^...Dimeihyl sulfoxide

''s Methonol "--...^^^ AQUEOUS SOLVENTS (25°)

\ ^ \ ^ V Ethanol

Tetrahydrofuron

^^'-^--^j^^^ Acetone

^""*^^--fc.*;.,_^^^^_tef1- Butanol-^

1 1 1 02 0-3 0-4 0-5 06 0-7 01 MOLE FRACTION OF ORGANIC SOLVENT

, Acetonitrlle MIXTURES WITH CARBON TETRACHLORIDE

(25°)

e 20

Ethanol

0! 0 2 0-3 0-4 0-5 0-6 0-7 0-8 0-9 10 MOLE FRACTION OF CARBON TETRACHLORIDE

Figure 2. Dielectric constants of aqueous organic solvents as a function of composition (25°).

Sources of data: Water-methyl sulfoxide, ref. 40: water mixed with methanol,

etlianol, n-propyl aleohol. or acetone, ref. 51; water-fert. butyl alcohol, ref. 52; water- dioxane, ref. 53; water-tetrahydrofuran, ref. 37.

Figure 4. Dielectric constants of binary solvent mixtures as a

function of composition (25°). Sources of data: Carbon tetrachloride-acetonitrile, ref. 61; carbon tetrachloride- nitrobenzene, refs. 62H>3; carbon tetrachloride-methanol, ref. 64; carbon tetrachloride- ethanol, ref. 65.

Figure 3. Dielectric constants of binary solvent mixtures as a

function of composition (25°). Sources of data: Benzene-methyl sulfoxide, ref. 54; benzene-nitromethane, ref. 55: benzene-acetonitrile, ref. 56; benzene-methanol, ref. 57; benzene-ethanol, ref. 58; benzene-n-butyl alcohol, ref. 59; benzene-l,2-dichloroethane, ref. 60.

Figure 5, Dielectric constants of binary solvent mixtures as a

function of composition (25°),

Sources of data: Dioxane-acetonitrile, ref. 66; dioxane-ethanol, ref. 67; dioxane- acetone, ref. 68: dioxane-chloroform (20°), ref. 69.

3.3. Br0nsted's Eight Classes

Br0nsted proposed [12] subdividing solvents into eight types, taking into consideration the acidic

and/or basic properties and the values of the di- electric constant. This scheme is shown in, table 3. Plus and minus signs indicate a "higher or lower

value" of the property referred to. Suggested examples of the eight types have been added.

For solvents of type 1 Br0nsted [13] coined the

designation "amphiprotic," to signify that they are both protogenic and protophilic and are therefore

"completely active protolytically." Solvents of type

8 represent the opposite extreme in properties. To

these solvents Br0nsted applied the adjective

"aprotic," to signify that they are "completely

indifferent protolytically."

Benzene, cited as an example of type 8, is now

well known to be a 7r-electron donor, and thus a base according to the most modern concept. However, its basicity is very low relative to that of n-electron donors [5]. Br0nsted recognized, of course, that all sorts of

transitions between the eight main classes of sol8
Table 3. Brflnsted's Scheme for Classifying Solvents [121'^

Type Dielec- Rela- Rela- Examples

No. tric Con- live tive (e, approx.)

tant, e Acidity Basicity

1 + + + Water (78), methanol (33). 2 + + Hydrogen cyanide (115), hy- drogen fluoride (84, 0°). 3 + + yV-Methylpropionamide (176), tetramethylurea (23). 4 + — — Acetone (21), acetonitrile

(36) , dimethylformamide

(37) , methyl sulfoxide (46), nitrobenzene (35). 5 + + feri. -Butyl alcohol (11), cyclohexanol (15), tert.- pentyl alcohol (6). 6 - + Acetic acid (6), propionic

acid (3). 7 + n-Butylamine (5), pyridine

(12), tetramethylguanidine

(11), dioxane (2), ethyl ether (4), tetrahydrofuran

(7). 8 Benzene (2.3), cyclohexane

(Z.U), z,Z,4-trimelnyipen- tane (1.9), carbon tetra- chloride (2.2), 1,2- dichloroethane (10).

^ Plus and minus signs indicate a comparatively high or low

value of the property referred to. In the selection of examples,

solvents of e S= 20 were arbitrarily regarded as having a comparatively high dielectric constant, that is, to have a "plus" value

for this property. Correspondingly, solvents of e < 20 were re- garded as having a low ("minus") value. Compounds commonly

considered to be good proton donors (acceptors) are assigned

plus values for relative acidity (basicity). Reference:

[12] J. N. Br0nsted, Chem. Ber. 61, 2049 (1928).

vent are conceivable. For the study of acid and

base properties, he said [12], a medium of type 4

would be of special interest, and he emphasized the

importance of solvents of type 8 in these words [13]: "The properties of acids and bases may then by no means, as is generally believed, be investigated in

I aqueous solutions, but must on the contrary be

1 investigated in solutions of benzene-like, aprotic

character if we would gain the most comprehensive

and deepest insight into the nature of these sub- stances." Moreover, he said [12]: "A thorough study

of the molecular state [in benzene and similar sol1 vents] must precede an exact quantitative investigation of acid and base strength [in such solvents]."

j This rough classification is useful, even though

I few would be in complete agreement about the most

i suitable categories for particular solvents. It will be noted that in the selection of examples for table

3, the value € = 20 was arbitrarily taken as separating

solvents of relatively high and relatively low di- electric constants. The dielectric constant of a

solvent bears a relationship, of course, to its acid- base character.

As previously indicated, this review is mainly

concerned with acid-base behavior in solvents of

type 8. Solvents of type 4, on account of their comparatively slight acidic and basic character,

are the most nearly similar in their influence on

solute behavior, and useful parallels can be drawn

between the two types.

3.4. Comparisons of Aprotic Solvents

The behavior of solutes dissolved in acyclic

saturated hydrocarbons is generally found to re- semble closely that of the gaseous solutes, but in other hydrocarbons there is evidence of more or

less solvent-solute interaction. For this reason, the

qualifying adverb "comparatively" or "relatively"

should always be understood when cycloparaffins

and aromatic or other unsaturated hydrocarbons

are referred to as "inert" solvents. At the same

time, any acid-base properties of these solvents

should not be overstressed, for all have far less acid-base character than such familiar acids and

bases as carboxylic acids, phenols, amines, ethers,

and the hke.

Halogenated hydrocarbons, likewise, are inert only in a comparative sense. Proton donor capacity

in halogenated compounds containing C-H, such as chloroform, has been amply demonstrated, and more recently, evidence of hydrogen bonding to halogen has been presented. Furthermore, in the

last several years there has been increased realiza- tion of the tendency of halogenated hydrocarbons

to react with nitrogen bases and with oxygen. Such

reactions are accelerated by irradiation.

3.4.1. Aromatic Solvents as Bases While benzene is the most commonly used hydrocarbon solvent, occasionally it is replaced by a related aromatic solvent like toluene or chloro- benzene, for such reasons as a more desirable boil- ing point or dielectric constant. Recognition that these solvents differ slightly in basic character is helpful in accounting for solvent effects sometimes

observed. The comparative basicities of aromatic

solvents will be discussed only briefly here. This

topic has been reviewed recently by Shatenshtein

[25], Briegleb [70], Arnett [71], and many others. Basic tendencies in benzene have been recog- nized for at least 25 years (see Davis's summary

of the evidence [5]). Considerable knowledge of the

basic character of aromatic hydrocarbons has

resulted from many investigations of the past 15

years, stimulated by Benesi and Hildebrand's

experimental studies of the association of aromatic

hydrocarbons with the Lewis acid iodine [72] and

the theoretical studies of "charge-transfer complexes" by Mulliken and co-workers [73-76]. In these investigations two of the most widely used

criteria are (a) the stabiUty constants of complexes

9
with Lewis or Br0nsted acids, determined by measuring spectral changes in the ultraviolet or visible regions, and (b) the comparative magnitudes of

shifts in infrared stretching frequencies. (The

interacting matericds are generally diluted with

solvents like heptane, cyclohexane, or carbon

tetrachloride.)

The examples of comparative stability constants

given in table 4^ are representative of many data

Table 4. Kc (literjmoles) values for associations in CCU (25°)^

TT-Electron Donor

Electron Acceptor

I2 ICl

0.15

0.16

0.54 Toluene 0.87 o-Xylene 0.27 1.24 0.31 1.39 0.31 1.51 Mesitylene 0.82 "3.70

0.88 6.43 Hexamethylbenzene 1.35 "13.2 Bromobenzene 0.13 0.32 m-Chlorotoluene bO.32

'>0.24

^Data are from Andrews and Keefer [77] except as noted

otherwise. ••From Ogimachi, Andrews, and Keefer [78].

References:

[77] L. J. Andrews and R. M. Keefer, J. Am. Chem. Soc. 74,

4500 (1952).

[78] N. Ogimachi, L. J. Andrews, and R. M. Keefer, J. Am.

Chem. Soc. 77, 4202 (1955).

which show that the basic character of the benzene

ring increases with increasing number of methyl

substituents. For example, analogous results were

obtained by selective extraction from heptane with

the reagent HF-BF3 [79], and by determining the

capacities of methylbenzenes to dissolve hydrogen chloride at —78.5° [80]. Such results would

be expected because of the well-known electron- releasing capability of the methyl group [81]. Halogen atoms, on the contrary, tend to attract electrons from the benzene ring, and it is not sur- prising to find (table 4, last column) that bromobenzene and m- and p-chlorotoluenes are poorer

TT-electron donors toward iodine monochloride than

benzene and methylbenzenes.

Olah and co-workers [82], from their own and

published data on shifts of infrared fundamentcd

H-X stretching frequencies, concluded that the

relative order of decreasing basicity of halobenzenes

is as follows:

CeHe > CeHsI > CeHsBr > CeHsCl > CeHsF.

Table 5 indicates there to be a consistency of infra- red results with various proton donors. Olah has

pointed out [82] that occasional inconsistencies — as, for example, the relatively high stability of the iodobenzene-Ag+ complex — probably can occur because

of the capacity of halobenzenes to act both as

TT-electron and ra -electron donors. Part of the results cited in table 5 cire from investigations by Josien

and co-workers. Additional results of interest may

be found in the original papers [83-84].

Differences in the TT-electron donor (hydrogen

bond acceptor) capacities of benzene, toluene, and

chlorobenzene provide a highly reasonable explanation for Popovych's results [85] in determining the extent of association of several pyridine-type bases

with the reference acid tetrabromophenolsulfoneTable 5. Shift (cm ') of infrared fundamental H-X stretching frequencies of proton donors in halobenzenes'

Electron Donor

Proton Donor

HCl HBr HI PhOH MeOH MeOD n-BuOH Pyrrole PhC^CH

CCl, -55 -40 -26 -44 -42 -31 -40 -35 -25

QHsF -76 -40 -40 -48 -44 -33 -50 -42 -33

CeHsCl -107 -77 -48 -76 -63 -43 -64 -55 -36

CeHsBr -117 -82 -54 -86 -69 -51 -70 -60 -38

CeHsI -125 -90 -63 -95 -77 -53 -71 -63 -40

-130 -106 -66 -97 -80 -60 - 74 -70 -41

^ Adapted from the compilation by Olah and co-workers [82]. Values given are for the shift in stretching frequency upon transfer from the gaseous state to solution.

Reference:

[82] G. A. Olah, S. J. Kuhn, and S. H. Flood, J. Am. Chem. Soc. 83, 4581 (1961).

' Unless stated otherwise, constants cited in the tables and text are either values

calculated by the original authors or logarithmic equivalents of these values. 10
phthalein (bromophenol blue). The association constants (see eq 16) are given in table 6. Values for

Table 6. Effect of solvent on some acid-base association constants (25°)

'^assocn. X 10"^ (liter/mole)

Base

Pyridine

Isoquinoline.

2-Picoline

3-Picoline

4- Picoline

2,4-Luti(line.

Chlorobenzene

(5.62)

2.50

14.6 28.6 438

Benzene

(2.27)

1.88 1.40 4.75 3.64

19.8 12.9 9.51 6.67

18.0 11.8 179 113

Toluene

(2.37)

^ From O. Popovych, J. Phys. Chem. 66, 915 (1962). Dielectric constants are given in parentheses. The reference acid was

bromophenol blue (tetrabromophenolsulfonephthalein), first color change (colorless to yellow).

Reference:

[85] O. Popovych, J. Phys. Chem. 66, 915 (1962).

log Kb (log ^assocn.) in the three solvents are linearly

related, but the Kas&om. values are greatest in chloro- benzene and least in toluene. Thus, in this particular

reaction chlorobenzene is the most nearly inert of

the three solvents, as it competes the least with the

pyridine-type bases.

There has been considerable interest in geometrical aspects of the interaction of benzene and other

aromatic hydrocarbons with electron acceptor

molecules. Proton magnetic resonance techniques

have been found to be particularly suitable for such

studies [86-87]. The vr-electron system in aromatic

molecules is commonly regarded as being a rela- tively exposed region of negative charge, effective

in a direction perpendicular to the aromatic ring. In the interaction of benzene with chloroform,

acetonitrile, and acrylonitrile, Schneider [86] has

interpreted the observed proton shifts as being

indicative of the structures shown in figure 6. Hatton and Richards [87] have proposed structures representing the directed interactions between

benzene and three amides (A'^,A^-dimethyl- and

diethylformamide, A^,A^-dimethylacet amide).

Recently, Creswell and Allred [88] studied the

association of benzene with chloroform in cyclohexane solution by means of nuclear magnetic

resonance spectroscopy. The equihbrium constant

(25°) is about 1.1 (m.f.)-i or about 0.11 Uter/ moles.

3.4.2. Unsaturated Aliphatic Hydrocarbons as Bases Using phenol and p-fluorophenol as reference

proton donors. West [89] obtained infrared spectroscopic evidence that weak hydrogen bonding to unsaturated aliphatic hydrocarbons occurs in car- bon tetrachloride solution. Part of West's results

H-^CFigure 6. Examples of directed solvent-solute interactions involving benzene. Top, left and right: Interaction with chloro- form and acetonitrile, respectively. Bottom: Interaction with

acrylonitrile.

Reproduced from W. G. Schneider, J. Phys. Chem. 66, 2653 (1962), with the permission of the author and the publisher. (Ref. No. 86.)

is given in table 7. From the relative shifts of the

0-H band the following conclusions seem justified:

Table 7. Shifts (cm~') of 0-H band for associations in CCU ^

Hydrocarbon

Proton Donor

CeHsOH P-F-C6H4OH

1-Me-cyclohexene 113 2-Me-2-butene 108 116 2-Me-l-butene 104 106 Cyclohexene 95 99

2-Hexene 85 1-Hexene 69 74 1-Octene 62 69

52 59

1-Hexyne 92 96

Mesitylene 73 78 p-Xylene 66 69 Toluene 58 61 Benzene 47 49

^Data of West [89], abridged and rearranged.

Reference:

[89] R. West, J. Am. Chem. See. 81, 1614 '1959).

(a) p-fluorophenol is a slightly better proton donor

than phenol; (b) olefins are better proton acceptors

than aromatic hydrocarbons; (c) basicity increases

with increasing alkyl substitutions at the double

bond; (d) conjugation (as in isoprene) would appear

to reduce basicity; (e) the triple bond in 1-hexyne

confers greater basic character than the double

bond in 1-hexene. Conclusions (a)-(d) fulfill expectations. Conclusion (c) is the same as the earlier conclusion of Brown and Brady [80], based on the

comparative solubihties of hydrogen chloride in varicrus hydrocarbons at — 78.5°. Basic properties of acetylene groups will be

referred to again in section 3.4.5.

280-613 OL-68—

11
3.4.3. Halogen Compounds as Bases Haloaromatic compounds as tt -electron donors were discussed briefly in section 3.4.1.

a. Alkyl halides. Recently several investigators

have determined values of At- (shifts in infrared

0-H stretching frequencies) and A// for hydrogen

bonding to halogen which is covalently bonded to carbon — that is, for hydrogen bonding to halogen

in alkyl halides, RX. A fuU discussion of such comparatively weak interactions is outside the intended

scope of this review, but examples of results ob- tained are given in tables 8 to 10. Table 8 gives values of t^v and thermodynamic

constants, obtained by West and co-workers for

Table 8. Thermodynamic properties and spectral shifts in hydrogen bonding of phenol to alkyl halides (CCI4, 25°) ^

Electron Donor, RX -AS°

caljdeg

(cm-') kcaljmole kcaljmole mole

Fluorocyclohexane 53 3.1 1.31 6.1 Chlorocyclohexane 66 2.2 0.87 4.5 Bromocyclohexane 82 2.0 0.85 4.0 lodocyclohexane 86 1.7 0.82 3.0

ra-Butyl ether 278 6.0 2.45 11.8

^ From West and co-workers [90].

Reference:

[90] R. West, D. L. Powell, L. S. Whatley, M. K. T. Lee, and

P. von R. Schleyer, J. Am. Chem. See. 84, 3221 (1962).

hydrogen bonding of phenol to four cycloalkyl

halides, and also to n-butyl ether, in carbon tetra- chloride solution [90]. As the authors pointed out, the spectral shifts indicate that the relative basicities

of the monohalocyclohexanes follow the order

RI > RBr > RCl > RF,

in agreement with earlier infrared spectral results

for n -butyl halides [91], while the enthalpy changes

indicate the opposite order of basic strengths. The

explanation suggested (ascribed to H. J. Bernstein)

is that At- measures only the weakening of the bond

PhO-H, whereas A//° measures the overall energy

change accompanying the formation of the H bonded

complex PhOH...XR and the simultaneous

weakening of the PhO-H and X-R bonds. The

data of table 8 show the alkyl halides to be decidedly

weaker electron donors than n -butyl ether. Infrared results of Jones and Watkinson for hydrogen bonding of phenol to n-heptyl halides

(table 9) [92] show the same trends in Ai^ and

Ah as in the work of West and co-workers, although

the spectral shifts and enthalpy changes are

smaller. The authors call attention to the probability

that "free" phenol molecules are hydrogen bonded

Table 9. Values of AH and Av for hydrogen bonding of 0.01 M

phenol to n-heptyl halides (CCl2 = CCI2, 25°) "

Halide (1 M) -AH

cm"' kcaljmole

CjHisF " 28.4 2.54

C7H,.5F 39.7 2.13 CtH.sCI 58.7 1.65 CrH.sBr 69.2 1.57 C7H,5l 71 1.25

"From Jones and Watkinson [92]. For bonding to PhOH with

the exception noted.

" Bonded to PhOD.

Reference:

[92] D. A. K. Jones and J. G. Watkinson, J. Chem. Soc, p. 2366

(1964).

to the solvent (tetrachloroethylene), and they suggest that greater inductive electron release from

cyclohexyl to halogen than from ra-heptyl may

account for the larger enthalpy changes observed

by West and co-workers (compare tables 8 and 9). Table 10 shows results obtained by Jones and

Watkinson [92] using a constant electron donor

Table 10. Values o/AH and Av for hydrogen bonding of m- and

p-substituted phenols to n-heptyl fluoride (CCl2 = CCl2, 25°)"

Substituent in PhOH Ap -AH

cm~^ kcaljmole

p-OMe 36.6 1.98 m-OMe 40.4 2.19 p-F 41.9 2.15 p-Cl 44.6 2.26

p-Br 45.4 2.31

p-I 46.4 2.43 m-N02 53.5 2.89 P-NO2 59.4 3.31

" From Jones and Watkinson [92].

Reference:

[92] D. A. K. Jones and J. G. Watkinson, J. Chem. Soc, p. 2366

(1964).

(n-heptyl fluoride) and several substituted phenols

as proton donors. This experiment yielded a fairly

good linear correlation of Ai^ and AH values. The comparative AH values for formation of

n-C7H,.,F...D0Ph and n-C7Hi.5F...HOPh (see table 9) indicate that the deuterium bond is stronger

than the hydrogen bond, in agreement with previous

evidence [92].

12
Krueger and Mettee [93] have measured fre- quency shifts and formation constants for numerous

methanol-base adducts in carbon tetrachloride.

Like the other workers cited [90-92] they concluded

that shifts in 0-H stretching frequencies upon

hydrogen bond formation are not a measure of the

overall bond strength.

b. Carbon tetrachloride. Infrared data are thought

to indicate that proton donors such as methanol,

pyrrole, and phenol [94, 95] form weak bonds to chlorine in carbon tetrachloride as a result of the

tendency of X-H dipoles to "[seek] out an appropriately charged polar group . . . with which to associate." Although the carbon tetrachloride molecule has zero dipole moment, the Cl~-C+ bond

has polar character [95]. In other words, carbon

tetrachloride is not the completely inert medium it was formerly assumed to be.

3.4.4. Carbon Tetrachloride as an Electron Acceptor

About 30 years ago Earp and Glasstone, follow- ing up previous indications of molecular compound

formation between carbon tetrachloride and ethers,

made dielectric polarization measurements of the

following systems [96-97]: Carbon tetrachloride

mixed with ethyl ether, isopropyl ether, dioxane,

acetone, or quinoline; carbon tetrabromide mixed

with ethyl or isopropyl ether; hexachloroethane

mixed with ethyl ether. The results in all cases

indicated a weak tendency for 1 : 1 association, and

for some systems approximate association constants were obtained. The authors considered that the union was a chemical one, with oxygen or

nitrogen functioning as an electron donor.

Recent work has led to analogous conclusions.

From calorimetric and cryoscopic data, Goates and

co-workers [98] deduced that benzene and carbon

tetrachloride form a 1 : 1 addition compound, with

the approximate heat of association — 3.5 kcal/mole.

They suggested that formation of this complex

may arise from 7r-bonding between the aromatic

ring and the empty 3-d level of chlorine. For the

corresponding p-xylene-carbon tetrachloride complex the heat of association is about — 6.3 kcal/mole.

Subsequently, Anderson and Prausnitz [99], using

ultraviolet spectrophotometry, determined association constants for the 1 : 1 complex of carbon tetra- chloride with benzene and also with mesitylene

(in n-hexane). The association constant calculated

for the complex formed with benzene is about 0.01

(liter/mole, 25°), and for the complex with mesityl- ene it is about 0.1. Anderson and Prausnitz attribute

this complex formation to charge transfer, with

the aromatic compound probably acting as the

electron donor. However, Person [99a] has questioned the vaHdity of the low K values obtained.

Sharpe and Walker observed that the dipole

moments of most amines are higher in carbon tetra- chloride than in benzene, and concluded that this indicated interaction of the amines with carbon

tetrachloride [100]. On the assumption that inter- action can be inferred when the increment exceeds

0.04 Debye unit, they concluded that ethyl ether

and butyl ether also interact with carbon tetra- chloride. In their view, also, the ether is the electron

donor. On the basis of their own work and a review

of work by others, they believe that tertiary amine

oxides, esters, phosphines, and sulfides are additional substances which interact with carbon

tetrachloride.

McKinnon and Williamson's recent heat of mixing studies [100a] indicate strong interaction of

carbon tetrachloride with dioxane, while Whetsel

and Kagarise's infrared data [101] are indicative

of a weak specific interaction between carbon

tetrachloride and cyclohexanone.

3.4.5. C-H Containing Compounds as Proton Donors

The ability of hydrogen in C-H containing

solvents like chloroform and phenylacetylene to become bonded to electron donor groups was re- ferred to in the beginning part of this section (also note table 5 and fig. 6). Organic chemists have long known that C-H

containing compounds with "negative" groups ad- jacent to the carbon possess exceptionally "active"

or "labile" hydrogen. It has also been realized for more than 35 years [96, 97] that acetone and ether

form crystalline compounds with chloroform, and

that anomahes in the behavior of their liquid mixtures with chloroform indicate intermolecular

association. It now seems highly logical to attribute

such phenomena to hydrogen bond formation.

a. Chloroform and related solvents. Nevertheless,

when Glasstone and Earp first interpreted the dielectric polarization behavior of mixtures of haloforms

(or related C-H containing compounds) with acetone, ethers, or quinoline on the basis of hydrogen

bonding, some skepticism was expressed about the

ability of hydrogen in a C-H bond to play the

part of a proton donor [96, 97]. Soon, however, a vast assortment of confirmatory evidence was obtained. Part of this was reviewed briefly by Pimentel and

McClellan [102], and some of the recent work has

been mentioned in further studies using nuclear

magnetic resonance spectroscopy [88] and infrared

spectroscopy [104]. The following is a partial Ust

of techniques which, in addition to the two tech- niques just mentioned, have been used in studying

interactions of chloroform with various bases:

Dielectric polarization, freezing point diagrams,

heats of mixing, the parachor, solubilities, and vapor

pressure measurements [88, 102-104]. Pimentel and

McClellan concluded their review by remarking: ". . . The evidence in favor of association of

chloroform with bases is conclusive. The evidence

that this association is of the H bonding type is substantial . . . ." One of the strong indications that hydrogen bonding must occur in the association of bases with

chloroform or other C-H containing molecules comes from the comprehensive solubihty studies

of Zellhoefer, Copley, and Marvel. For example,

13
they found that the solvent capacities of halo- methanes for basic molecules follow the order

CH3X < CH2X2 < CHX3 »CX4,

where X = C1 or F. The many solutes tested in- cluded ethers, thioethers, esters, ketones, and primary and tertiary amines. The failure of alcohols

to associate with chloroform was attributed to their

strong tendency toward intermolecular self-asso- ciation [105, 106]. The order given above is exactly

that predicted on postulating that C-H hydrogen

is involved in the bonding and that its activity is progressively enhanced on continued replacement

of methane hydrogens by halogen. Heats of mixing

data [107] show bromine to be a little less effective

than chlorine in promoting activity of C-H hydrogen. The phenyl group also appears to have less activating effect than chlorine.

Infrared spectroscopy has been a particularly

favored method for studying hydrogen bonding

by chloroform and related C-H containing com- pounds [101, 102-104, 108-112]. Besides adding

strong support to earher deductions from other

kinds of data (dielectric polarization, solubilities,

heats of mixing, etc.), it has contributed supplementary conclusions.

For example, in a comparative study of hydrogen bonding of chloroform, sym.-tetrachloroethane,

or pentachloroethane (diluted with n-heptane) to phosphorus oxychloride, triethyl phosphate, and

tri-n-butylphosphine oxide [111], the order of

hydrogen bonding found for the three phosphoryl

compounds tested was

n-BuaPO > (EtO)3PO > POCI3,

in harmony with Kosolapoff and McCullough's

previous conclusions from heats of mixing data [113]. The comparative hydrogen bonding ability found

for the three halohydrocarbons was CHCI3 > CHCI2CCI3 > CHCI2CHCI2.

A few years ago Kearns suggested [114] that, besides the 1 : 1 complexes like Me2C=O...HCCl3

previously assumed to exist, 1 : 2 association can

also occur, giving a complex of the general type

indicated by (I). Me2C=0 . . .HCCljCl. . .HCCl^

(I) By assuming occurrence of 1 : 2 association, certain

"excess" free energy and heats of mixing data can

be accounted for. Whetsel and Kagarise, from

infrared studies of the association of acetone or cyclohexanone with chloroform (diluted with

cyclohexane) [101], agree that both 1 : 1 and 1 : 2 complexes are probably formed, but prefer Las- combe's formula (H) [115] for the 1 : 2

.•HCCl, Me2C=o: 3

'• HCCl^

(II)

complex. In Lascombe's formula both molecules of

chloroform are hydrogen bonded to the carbonyl

group. Lowering the temperature from +27° to — 55° produced spectral effects which were consist- ent with the inferences drawn from changes in concentration. For the 1 : 1 association of cyclohexanone with chloroform Whetsel and Kagarise

calculated K= 1.2 ±0.3 (liter/mole, 30°). The association constant for acetone-chloroform complexing was the same within the experimental uncertainty.

This value is about 10 times greater than earUer

estimates. For the 1 : 2 association they estimated

K = 0.2. Whetsel and Lady [112] recently studied the

complexing of chloroform with cyclohexylamine,

aniline, and several aniline derivatives (in cyclohexane). In the case of most ring-substituted anilines, log K values for amine-chloroform association were found to be linearly related to pKa values for the amines in water. Examples of equilibrium constants and heats of reaction obtained by Whetsel

and Lady are given in table 11, which also contains

results for other associations involving chloroform.

Results in which A^assocn. is given in mole fraction"^

units are presented separately in table 12. (As the

data in these tables are shown mainly to permit

general conclusions and comparisons, detaUs about

conditions of measurement and estimates of precision and accuracy given in the original papers have

been omitted.)

As more is learned about intermolecular reactions, some of these values may need revision.

During the past decade the self-association (di- merization) of chloroform has been discussed and

studied considerably. On the basis of our preceding

discussion self-association would certainly be ex- pected, and there are numerous experimental

reasons [121, 122] for accepting its existence as demonstrated. The dimerization constant 0.013

(liter/mole, 25°) has been obtained [116] from NMR

data (see tables 11 and 12). Base-haloform association constants that have been calculated without

taking chloroform dimerization into account, and

also possible 1:2 association, are very likely some- what in error. As mentioned earlier in section 3.4, chloroform

is thought to bond weakly to benzene, and still more weakly to carbon tetrachloride (see table 11). Comparative K values for association of acetone

with haloforms, like heats of mixing data [107],

indicate that chlorine is more effective than bro- mine in activating the C-H hydrogen (see tables

11 and 12). Bromine in turn has a greater activating

effect than iodine. The activating effects of fluorine

and bromine appear to be about equally great

(table 12). Huggins and Carpenter [123] have deduced from

NMR data that silicochloroform, SiHCls, is a

slightly weaker proton donor than chloroform.

b. Acetylenic hydrocarbons. Hydrogen bonding of

acetylene and related hydrocarbons to electron

donors has been a topic of growing interest in the

past 25 years. Such bonding was predictable from

14
Table 11. Equilibrium constants and heats of reaction for

base-haloform associations "

Proton Electron Temp. Method -A// Helcrence Donor Donor

°C I iter/ mole kcal/mole CHCI3

CHCI3

CCI4 NMR 0.01 [88] CHCI3 25 NMR 0.013 [116] CHCI3 C(iH(i 25 NMR 0.11 2.0 [88] CHCI3 PhNHa 25 IR 0.51 1.7 [112]

CHCI3 EtsN 25 NMR 0.51 [88] CDCI3 EtaN IR 0.36 [118] CHCI3 MejCO 30 IR 1.2 [101] CDCI3 MejCO 31 IR '•O.QO 3.5 [117]

CDClzBr MejCO 31 IR ^0.80 3.3 [117] CDBr3 MeaCO 31 IR 0.45 2.7 [117] CHCI3 CeHnNH^ 25 IR 1.10 3.6 [112]

^ Except as noted otherwise, the mixtures were diluted with cyclohexane.

" No diluent. n-Hexane was used as diluent.

References:

[88] C. J. Creswell and A. L. AUred, J. Phys. Chem. 66, 1469 (1962). [101] K. B. Whetsel and R. E. Kagarise, Spectrochim. Acta 18, 329 (1962). [112] K. B. Whetsel and J. H. Lady, J. Phys. Chem. 68, 1010 (1964). [116] C. F. Jumper, M. T. Emerson, and B. B. Howard, J. Chem. Phys. 35, 1911 (1961). [117] R. E. Kagarise, Spectrochim. Acta 19, 629 (1963). [118] G. M. Barrow and E. A. Yerger, J. Am. Chem. Soc. 76, 5247 (1954).

Table 12. Equilibrium constants and thermodynamic constants for base-haloform associations (cyclohexane, 25°) ^

Proton Donor

CHCI3

CHCI3

CHCI3

CHCI3

CHF3

CHF3

CDF3

CHCI3 CHBra

CHI3

Electron Donor

ecu

CHCI3

CeHe

EtjN

THE"

THE

THE

THE

THE

THE

^assocn. (m.f.)-'

0.1 0.16

1.06

4.2 4.57

4.77 4.95

5.78

4.81 4.64

-^H

(kcall mole)

1.97 4.15

2.4

2.59

2.73

3.6

2.6

1.6

-^s

(caljmole

deg)

6.5

11.0

5.2

5.59

5.97

8.6

5.6

2.3

Reference

[116]

[116]

[88]

[88] [119]

[120]

^ [120]

[119]

[119]

[119]

^ By proton NMR spectroscopy except as noted otherwise. Tetrahydrofuran.

' By fluorine NMR spectroscopy.

References:

[88] C. J. Creswell and A. L. Allred, J. Phys. Chem. 66, 1469

(1962).

[116] C. F. Jumper, M. T. Emerson, and B. B. Howard, J. Chem.

Phys. 35, 1911 (1961). [119] C. J. CresweU and A. L. Allred, J. Am. Chem. Soc. 85,

1723 (1963). [120] C. J. Creswell and A. L. Allred, J. Am. Chem. Soc. 84,

3966 (1962).

the labile nature of C-H hydrogen in these compounds. Copley and HoUey [124] measured the solubility of acetylene in numerous organic solvents, and,

in harmony with expectations, found its solubilities

in ethers, esters, and ketones to be greatly in excess

of those predicted by Raoult's law. They also found,

from heats of mixing data, that phenylacetylene

(ethynylbenzene) forms 1:1 complexes with N,Ndimethylacetamide, ethyl ether, acetone, cyclohexylamine, and methyl acetate. It wiU be recalled

that acetylene is very soluble in acetone and has

customarily been transported as an acetone solution, for safe handhng [125]. Shortly after Copley and HoUey's work, Stanford and Gordy [126] obtained supporting evidence

of hydrogen bonding from infrared spectra of

phenylacetylene mixed with oxygen- or nitrogen-containing solvents (acetone, dioxane, triethyl phosphate, A'^jA^-dimethylacetamide, pyridine, and re- lated bases). Recent investigations, principally by

infrared spectroscopy and/or NMR spectroscopy,

have yielded much additional evidence of C-H...0

15
and C-H...N formation in acetylene and related

compounds, as well as evidence of C-H bonding

to other n- or 7r-electron donors. Most of the work,

though of general interest, is not closely pertinent

to our topic, and only a few examples will be mentioned. Some of the additional work is reviewed in references cited [104, 127-30].

Brand and co-workers [127] found that in general the comparative infrared spectral behavior

of acetylene derivatives correlates well with their chemical reactivities. They noted also that gasliquid chromatography supplies evidence for hydrogen bonding — whenever hydrogen bonding

to the stationary phase can occur, the retention

time of an acetylene derivative is substantially

increased. For 1:1 association of phenylacetylene

with ethyl ether they computed ^assocn =1-1 (mole

fraction units, CCI4, 29°), AH = -1A kcal/mole,

and z\5 = — 4.5 cal/mole deg.

Kreevoy and co-workers [128], applying both IR

and NMR spectroscopy, estimated for the phenylacetylene-pyridine 1:1 complex an association constant in the range 0.1-0.2 (hter/mole, CCI4).

In view of the 7r-electron donor character of

acetylene derivatives, mentioned earlier (sec. 3.4.2), a tendency toward self-association is to be expected

in C-H containing acetylenes. Brand and co- workers found such evidence [127] in the infrared

spectra of all compounds which they studied in the

liquid state — for example, n-hexylacetylene (1-octyne), phenylacetylene, and ethyl propiolate (propynoic acid ethyl ester). Josien and co-workers [129]

demonstrated self-association of 1-heptyne and

1-deutero-l-heptyne (in CCI4) by infrared spectros- copy, as well as association of these two hydrocarbons with carbon disulfide, ethers, ketones,

and aromatic hydrocarbons. Like other proton

donors, acetylenes associate more strongly with

aromatic hydrocarbons as the latter increase in basic character.

West and Kraihanzel's more recently published

infrared investigations [130] included a fuller study

of the basic properties of acetylenes and the self- association of "terminal" acetylenes. They pointed

out that acetylene derivatives such as phenylacetylene and propargyl halides have more than one

electron donor center — the acetylene bond and also

TT-electrons of the aromatic ring and ^-electrons on

the halogen atoms. West and Kraihanzel found that

self-association of 1-alkynes is very incomplete, even in the pure liquids, and so has only a sUght

effect on the macroscopic physical properties.

c. Other studies of C-H containing solvents. Shatenshtein (ref. [25], sec. Ill) has discussed the

behavior of hydrocarbons as acids in types of reactions with which this review is not concerned, such

as the formation of organometaUic compounds and

hydrogen isotope exchange reactions with bases.

Recently Allerhand and Schleyer published a comprehensive survey, based on infrared spectros- copy, of the ability of C-H groups to serve as proton donors in hydrogen bonding [104]. Their results were obtained using two strong proton acceptors — pyridine-d'5 and methylsulfoxide-<i6 (in carbon

tetrachloride). Among their conclusions are the

following:

(1) A marked spectral shift is definite evidence

for hydrogen bonding; the absence of a detectable

band shift or alteration in band intensity means

that hydrogen bonding, although not necessarily

absent, is unimportant.

(2) The capacity of a C-H group to act as a pro- ton donor in hydrogen bonding varies with the

carbon hybridization,

C(sp)-H > C(5p2)-H > C(sp3)-H,

and increases with the number of adjacent elec- tron-withdrawing groups.

(3) More than one electron-withdrawing group

must be attached to an sp^-hybridized carbon for any C-H hydrogens to participate in hydrogen

bonding.

(4) When three or more chlorines are attached to the benzene ring the aromatic protons become

labile enough for bonding to strong proton acceptors. As an example, sym.-trichlorobenzene bonds to pyridine.

(5) The relative proton donor capacities of the

haloforms is

CHBrs > CHI3 > CHCI3.

(This conclusion is in disagreement with results cited in sec. 3.4.5a, in which the reference base was acetone or tetrahydrofuran.)

(6) Methyl groups do not hydrogen bond, and

neither do C-H groups in alkyl chains, aldehydes,

formates, or formamides.

(7) C-H groups can have a very wide range of

proton donor abilities. Spectral shifts range to well over 100 cm"^

3.4.6. Chemical Reactivity of Halogenated Hydrocarbons

In the use of halogenated hydrocarbons as "inert"

solvents, their chemical reactivity has sometimes

been overlooked. Chloroform is well known to react with atmospheric oxygen when kept in color- less glass containers [131]. Carbonyl chloride (phosgene) and hydrogen chloride are products. Before

its employment as an aprotic solvent, the ethanol

(0.5-1%) added to chloroform as a stabilizer must,

of course, be removed [131]. Dichloromethane be- haves similarly to chloroform [131]. In the writer's

laboratory, 1 ,2-dichloroethane [ethylene chloride)

of high commercial grade has been found to ac- quire a pronounced acidic reaction within a few

days when stored in a colorless glass container

in diffuse daylight, but no acidic reaction is ob- served after prolonged storage in an amber glass

bottle. Formation of acid is detected conveniently

by adding a benzene solution of an indicator base

such as methyl yeUow.

16
Miiller and Ehrmann [132] exposed several

halogenated hydrocarbons to radiation from a mercury vapor lamp in the presence of dry nitro- gen or oxygen. With nitrogen, no change was

observed, but with oxygen, reaction occurred, one

product being hydrogen chloride. The other principal products identified are listed in table 13.

Table 13. Products formed on irradiating halogenated hydro- carbons in the presence of dry oxygen "

Compound Irradiated Main Product

1 ,2-Dichloroethane Monochloroacetic acid sym.-Tetrachloroethane Dichloroacetyl chloride Pentachloroethane Trichloroacetyl chloride Trichloroethylene Dichloroacetyl chloride

Tetrachloroethylene Trichloroacetyl chloride Chlorobenzene No change

Benzyl chloride Benzoic acid 1,1-Dichlorotoluene Benzoyl chloride

^ From Miiller and Ehrmann [132].

Because of the considerable use of carbon tetra- chloride as a solvent in infrared spectroscopy,

evidence of its reactivity with nitrogen bases warrants special mention. Several workers have observed its reaction with triethylamine at room

temperature [133-136]. Protection from daylight

retards the reaction. Joesten and Drago [137],

when studying the association of triethylamine

with phenol in carbon tetrachloride, minimized

the concentration of free amine in contact with

carbon tetrachloride by first mixing the amine

with the phenol and then adding the solvent.

Reference:

[132] E. MuUer and K. Ehrmann, Chem. Ber. 69B, 2207 (1936). Stevenson and Coppinger [136] studied the reaction of triethylamine with several tetrahalogen

derivatives of methane, and concluded that the

order of reactivity of the halo-derivatives is

CBr4 > CBrCls > CCU > CFCI3.

They beheve that 1:1 amine-halide association occurs first, with subsequent formation of amine

hydrohalide and chloroform. Reaction also occurs

in mixtures diluted with isooctane. Di-n-propylamine and tri-n-propylamine were found to react

similarly with carbon tetrachloride.

Reaction of piperidine with carbon tetrachloride

has also been observed [138, 139]. The reaction was at first thought to be catalyzed by copper in a chromium-plated brass hypodermic needle [138],

but later observers of the reaction [139] claimed that copper is dissolved, and therefore is not a catalyst

for the reaction. Still others [155, 157] have com- mented on the reactivity of amines with carbon

tetrachloride.

In checking claims that chloroform likewise re- acts with triethylamine, Williams [140] concluded

that purified chloroform is "almost unreactive,"

the seeming reactivity of chloroform B.P. being

due to impurities (mainly to CH2BrC]). These conclusions were supported by Foster [134J. Sadek and Fuoss reported unexpected reactivity

of carbon tetrachloride with ethanol [65J. Mixtures

of the two solvents showed much higher conduct- ance than either of the separate liquids, and after several hours the conductance had more than

doubled. Methanol is a great deal less reactive

with carbon tetrachloride. Mass-spectrometric ex- amination of products formed by irradiating a mixture of carbon tetrachloride and ethanol indicated

that hydrogen chloride and carbon monoxide are among the compounds formed. Phosgene may be

an intermediate product.

3.4.7. "Dipolar Aprotic" Solvents Parker [141] has devoted considerable study to the behavior of a group of solvents which he characterizes as "dipolar aprotic" solvents. These are defined as "solvents with dielectric constants > 15,

which . . . cannot donate suitably labile hydrogen atoms to form strong hydrogen bonds with an

appropriate species." The following are examples:

Acetone, acetonitrile, benzonitrile, dimethylacetamide, dimethylformamide, methyl sulfoxide, nitro- benzene, nitromethane, and sulfolane (tetramethylenesulfone). These solvents have dipole moments

in the range 2.7-4.7 [141, 142] and dielectric constants in the approximate range 21-46.5. On refer- ring to table 3 (sec. 3.3) it will be noted that solvents

such as these were assigned to Br^nsted's fourth

class. "Dipolar aprotic" solvents, like the aprotic

solvents just discussed, are not completely lacking

in acid-base properties even though they do not

display strong tendencies toward hydrogen bond

formation. As will be seen in section 4, the limited

hydrogen bonding capacities of solvents of Br0nsted's fourth and eighth classes causes hydrogen

bonding of solutes to become much more apparent

than in aqueous and alcoholic solutions.

Current investigations, in which dipolar aprotic

solvents are being much used, have led to a revival

of interest in varied kinds of evidence that some

solvents of this group may form dimers, as well as weak 1:1 complexes with other dipolar aprotic

solvents. Such complex formation is believed to be due predominantly to antiparallel, electrostatic,

multipolar interactions involving localized polar

groups rather than entire molecules [143—145].

Some years ago H0jendahl [146] postulated an inter- action of the sort to account for the variation of the

polarization of nitrobenzene in benzene solutions

as a function of its concentration. The variation

could be accounted for fairly well in the dilute range of concentrations on assuming that part of

the nitrobenzene molecules associate in pairs with

opposed dipoles, the resultant moment being zero.

17
The postulate of specific interactions between

neighboring molecules is not always needed in accounting for "static" dielectric constants of polar

liquids. Buckley and Maryott [147] modified the

Onsager-Kirkwood theory of the static dielectric constant of polar liquids [148, 149] by considering

the influence of molecular shape, and found that on treating some molecules as prolate spheroids

and others as oblate spheroids the "deviation

factor," G, became unity or very nearly so, over a wide range of temperature, signifying agreement

of theory and experiment. Such cases included

benzonitrile, propionitrile, chlorobenzene, ethyl

bromide, ethyl ether, methyl ethyl ketone, and

nitromethane. It was not necessary to postulate

specific molecular interactions or dipole associations. In the case of nitrobenzene, however, the

assumption of association seems needed. (Hydrogen

bonded liquids were excluded from consideration.)

Nitromethane, another dipolar aprotic solvent,

is widely thought to be self-associated. The de

Maines and co-workers [150] have itemized many

types of experiments which indicate self-associa- tion of nitromethane in the pure liquid and when

diluted with inert solvents. From ultraviolet ab- sorption data they calculated closely similar di- merization constants for carbon tetrachloride and

ra-heptane solutions (105.1 and 104.2, respectively,

liter/moles, 20°). The respective values found for A// (kcal/mole) and AS (cal/mole deg) were —0.34

and -12.8.

3.4.8. Concluding Discussion of Solvents The investigations summarized in the preceding

paragraphs show convincingly that practically edl so-called "inert" solvents display measurable

acid-base character when examined (in suitable

environments) by very sensitive modern techniques,

thus extending prior chemical evidence and conclusions from older physical measuring techniques.

Although hydrogen bonding is by no means the

only factor involved in solute-solvent interactions

(for example, see WiUiams' and HeJlam's reviews

of infrared studies of molecular interactions [151,

152]), there seems clear evidence that very com- monly it is a major factor. Bellamy and co-workers

[95] (see table 14) showed that the frequency shifts of the N-H stretching mode of pyrrole (serving as reference proton donor) in 20 solvents of varied . character increase progressively from n-hexane, one of the most inert solvents, to pyridine, a well- recognized hydrogen bond former with proton

donors. It can be seen by inspection that the fre- quency shifts (Av/v X 10^) are not linearly related

to the dielectric constants of the solvents. Extending earlier work by Gordy with numerous

basic solvents and different proton donors (HCl,

MeOD, P2O) [153], Bellamy and co-workers compared the frequency shifts of the N-H stretching

mode of pyrrole with corresponding frequency

shifts for other proton donors in solvents ranging

from inert to recognized proton acceptors. Part of

the results obtained by (or cited by) Bellamy and

Table 14. Relative frequency shifts (^v/f X 70') in hydrogen bonding of

pyrrole to organic solvents ^

Solvent (e) AkIv X 103 (N-H)

Solvent (e) Avlu X 10^ (N-H)

n-Hexane (1.9).... Carbon tetrachloride (2.2). Chloroform (4.8) Carbon disulfide (2.6) Chlorobenzene (5.6)

syro.-Tetrachloroethane (8.2 "). Bromobenzene (5.4 '') 5ym.-Tetrabromoethane (T.O*"). 1,2-Dichloroethane (10.2) Nitromethane (35.9)

6.8

8.5

12.5 13.9

14.7

15.6

16.4 17.6

18.4

19.8

Benzene (2.3) Nitrobenzene (34.8)... Mesitylene (2.3) Acetonitrile (37.5) Ethyl acetate (6.0)....

Acetone (20.7) Cyclohexanone (18.3) Dioxane (2.2) Ethyl ether (4.3) Pyridine (12.3)

20.4

22.9

25.8 30.6

34.6

39.6

47.3

50.2

50.4

88.1

^ Referred to the vapor state. Excerpts from table 1 of Bellamy and co-workers [95].

" From ref. [35].

References:

[95] L. J. Bellamy, H. E. HaUam, and R. L. WUliams, Trans. Faraday Soc. 54, 1120

(1958).

[35] A. A. Maryott and E. R. Smith, Table of Dielectric Constants of Pure Liquids,

NBS Circ. 514, Aug. 10, 1951.

18
co-workers are shown in figure 7. Clearly, the

shifts for the four different proton donors indicated

100 1 1 1 1 1 1 1 r

C^vlv X 10^

t^vlv X io3, Pyrrole

Figure 7. Frequency shifts of the HBr, MeOD, and BioHu

stretching modes in a variety of solvents, compared with the corresponding frequency shifts of the NH stretching mode of

pyrrole.

Based on data obtained or cited by L. J. Bellamy and co-workers [95].

(pyrrole, HBr, MeOD, BioHh) are linearly related,

to a good approximation, even though the proton

donors vary considerably in donor capacity. This

result led the authors to observe [95|: "It would seem that the nature of any complex formed due

to the association of, say, pyrrole with benzene is essentially similar to that formed by any of the other

solutes with the same solvent. This suggests that hydrogen bonding ... is the major factor in- volved in all cases." One of the main purposes of section 3 has been

to show that most of the aprotic solvents are not

truly inert, but show proton donor-acceptor behavior,

thus affecting quantitative relationships observed

in studying interactions between stronger acids and

bases. When this is borne in mind, the aprotic

solvents are admirably suitable media for ascer- taining differentiating kinds of acid-base behavior

that are masked in amphiprotic and other active

solvents. In years to come they should find much

more extended use in analysis and other chemical

applications.

For the most part, the remainder of the discussion wiU be concerned with much stronger proton

donor-acceptor reactions, involving compounds

generally thought of as acids and bases because

they are measurably ionized in water.

4. The Role of Hydrogen Bonding in Aprotic Media

We have come to the central theme of this dis- formation of conjugate and heteroconjugate ions,

cussion, namely, the varied types of hydrogen bond- The objective of this section is, first, to assemble

ing which determine overall acid-base behavior experimental evidence about various kinds of hydroin aprotic solvents. gen bonding interactions, obtained by many pro- Hydrogen bonding plays a major role in all cedures, such as cryoscopy and ebuUioscopy, con- types of solvents. It is responsible for the most ductance, electric polarization measurements,

characteristic properties of water as a solvent — for and absorption spectroscopy. The section will example, its high dielectric constant, its "levehng then be concluded with a brief discussion of theoeffects" on dissolved acids and bases, and the for- retical aspects of hydrogen bonding in aprotic

mation of "iceberg" structures about inert solutes— solvents,

and a widening recognition of its importance is responsible for recent intensified interest and new 4.1. Self-Association of Nitrogen-Containing

approaches in the study of aqueous behavior. Bases

Hydrogen bonding is equally important in aprotic

solvents. In section 3.4 it was shown that even these Self-association through formation of NH...N

solvents, although frequently referred to as "in- bonds is now known to occur rather widely among

ert," display significant and differing hydrogen heterocyclic bases, aromatic amines, and also ali- bonding tendencies. However, in these solvents phatic amines, both in the liquid and solid states

the interactions of principal interest are the much and also as concentrated solutions in inert solvents,

stronger interactions between dissolved- species. Compounds not having a hydrogen atom attached

To comprehend acid-base behavior in aprotic sol- to nitrogen — for example, triethylamine and pyri- vents one must become acquainted with different dine — do not, of course, undergo such self-as- types of intermolecular hydrogen bonding reac- sociation.

tions that are likely to be occurring simultaneously. Like types of hydrogen bonding that were dis- as well as with the relative strengths of such cussed earlier, NH...N formation was inferred

interactions. some years ago from such evidence as solubility

Intramolecular hydrogen bonding (chelation) as- behavior and cryoscopic or ebuUioscopic data, with sumes greater importance in aprotic solvents than later substantiation by newer techniques like di- in leveling solvents. Among the intermolecular electric polarization, infrared spectroscopy, and

hydrogen bonding reactions of particular impor- NMR spectroscopy. The NH...N bond seems to tance are self-association of acids, bases, and salts, be comparatively weak — weaker, for example,

formation of hydrogen-bonded ion pairs, and the than the 0H...0 bonds formed by aggregation of

19
alcohol molecules. Occasional failures to detect

its formation have been ascribed to use of too

dilute solutions, too limited a concentration range,

or instruments of insufficient sensitivity. In some

cases, NH is considered to become bonded to a

TT-electron system.

Studies of NH...N formation have followed

the common pattern of first obtaining varied qualitative evidence, v^ith succeeding attempts to determine association and thermodynamic constants. Investigations are now in the second stage.

The small association constants are experimentally

difficult to determine, for such reasons as sensitivity

of the bases towards atmospheric moisture (which

may affect the extent of aggregation), susceptibility

to oxidation, and reaction with solvents like carbon

tetrachloride. The magnitude of the association constant will be somewhat affected by the degree

of inertness of the aprotic solvent selected. The

association constants obtained also depend on

assumptions made with regard to the composition of the hydrogen-bonded aggregates (that is, whether they are dimers, trimers, tetramers, mixed

species, etc.) and their structure (linear or cyclic).

Consequently, the association constants deter- mined in different laboratories are often not in close agreement. Comparison of results is impeded

somewhat because the choice of concentration

units varies with the areas of interest of investi- gators, some of whom prefer mole fraction units,

while others prefer molar or molal units. The symbols Kc, Km, and Kx (or Kn) are generally used to designate equilibrium constants in molar, molal,

and mole-fractional units, respectively. In dilute

solutions, the following relations are approximately

correct for association reactions [78, 170]:

Kc = Kmld^ of solvent, (26)

Kc = Kx (mol. wt. of solvent)/(1000 solvent).

(27)

(It will be realized, of course, that the reciprocals

of these conversion factors must be applied to dissociation constants [32].) Most authors now refer to the various self-aggregates as "multimers" [153a] or as "n-mers,"

avoiding the term "polymers" because of its dif- ferent significance. "Ohgomer" is another new

term, appHed when n has a small value [169-171].

4.1.1. Aliphatic Amines

Not surprisingly, tertiary aliphatic amines, which

cannot form NH...N bonds, often show strik- ingly different behavior from that of primary and

secondary amines. Marvel and coworkers [107]

noted, for example, that the heat of rriixing of

dimethylcyclohexylamine with chloroform is much

higher (approx. 1.2 kcal/mole, 3°) than that of

cyclohexylamine (0.8 kcal) or n-butylamine (0.7 kcal), and suggested that the low values for the

primary aliphatic amines may be due to N-II...N

association. In later studies, Lambert and

Strong [154] measured the second virial coefficients

of ammonia and primary, secondary, and tertiary

methyl- and ethylamines (as vapors, 20 to 130 °C),

and interpreted the results as showing dimerization to occur with all of the compounds except

tertiary methyl- and ethylamine. From Raman and

infrared spectral data for pyrrolidine over a wide

temperature range, Evans and Wahr [155] concluded

that gaseous pyrrolidine is monomeric, but most,

if not all, molecules of liquid pyrrolidine (25°) are associated by N-H...N bonds, while there is partial association in carbon tetrachloride solutions. (Pyrrolidine was observed to react slowly

with carbon tetrachloride; compare sec. 3.4.6.) Linnell and co-workers [156] determined the

dimerization constant of pyrrolidine in carbon

tetrachloride from infrared data, obtaining the

approximate value 0.89 (Uter/mole). Feeney and

Sutcliflfe [157] studied n-mer formation of monoand diethylamine, isobutylamine, and ammonia

(CCI4) by NMR spectroscopy, while Bystrov and

Lezina [158], also applying NMR, investigated

N-H...N formation in the following bases

(CCI4): Methyl- and dimethylamine, ethyleneimine,

trimethyleneimine, pyrrolidine, piperidine, and

hexamethyleneimine.

Wolff and co-workers have performed very detailed studies of hydrogen bonding by primary

aliphatic amines, employing Raman spectroscopy

[159], vapor pressure measurements [160, 161],

and infrared spectroscopy [162, 163]. The follow- ing primary amines, dissolved in aliphatic hydrocarbon solvents, were included in the investigations: Methyl, ethyl, n-propyl, n-hexyl, and /i-decyl.

From Raman spectroscopy they concluded that on the assumption that the amino nitrogen is 5p^-hybridized the most probable structures are ones in which only one hydrogen of the amino group

forms a hydrogen bond, the other remaining free. Infrared data (—110 to +190°) supported this conclusion for amines in the liquid state, but in crystal- lized methylamine the results seemed to indicate

that both hydrogens of the amino group take part

in hydrogen bonding. For carbon tetrachloride

solutions the existence of the monomer and of cyclic dimers and trimers was deduced. The mean

association energy of methylamine (CCI4) was cal- culated to be about 2 kcal/mole (within the limits of error), in agreement with the result obtained

from vapor pressure measurements (mostly, —55

to + 20°) of the system methylamine-n-hexane. The same mean energy of association was found for the other amines studied. At low temperatures and

high amine concentrations, besides NH...N bonds,

C-H...N bonds are also thought to form.

Vapor pressure data led Wolff and associates to the conclusion that at room temperature and a methylamine mole fraction of < 0.15, the associated particles consist of monomers, dimers, trimers,

and tetramers. Representative data which they ob- tained are shown in tables 15 and 16. Table 15 compares association constants obtained for methyl-,

ethyl-, and /i-propylamine (ra-hexane, —20 to +20°,

20
Table 15. Comparative association constants of primary ali- phatic amines from vapor pressure data (n-hexane, mole

fraction unitsY

Amine t, °c

MeNH2 + 20 " 1.52 2.56 3.85 0 1.99 3.46 5.17 — OAzU 2.67 4.72 7.04

EtNHs + 20 1.17 2.06 3.13 0 1.55 2.74 4.12 -20 2.02 3.62 5.41

n-PrNH2 + 20 1.00 1.89 2.86 0 1.31 2.37 3.56 -20 1.92 3.38 5.04

^ Selected from results of Wolff and co-workers [161].

"/i:,, 2 (liter/mole) = 0.198.

Reference:

[161] H. Wolff, A. Hiipfner, and H.-M. Hiipfner, Ber. Bunsenges.

Physik. Chem. 68 , 410 (1964).

mole fraction units). The constants were computed

for stepwise conversion of monomer to the dimer,

trimer, and tetramer. The following conclusions were drawn: (1) Under similar conditions, primary

amines are less associated than alcohols; (2) for each amine, the equilibrium constants for dimer,

trimer, and tetramer formation are different, and

increase in the order named. (This means that the

associations cannot be described as "associations

derived by repeatable steps," an assumption often

made in studying n-mer formation.); (3) with in- creasing carbon content, the amines become less highly associated; (4) with decreasing temperatures, all association constants show the expected

[ increases; (5) the heats of formation for the different

I reactions are not measurably different (all are ~ — 2 kcal/mole).

I Table 16 indicates how the chain length of the

I solvent affects the K values for w-mer formation

from methylamine. The results shown are repre-

: sentative of those obtained with ethyl- and propylamine. It is noteworthy that when ^assocn. values are expressed in mole fraction units, they show a de- crease with increasing chain length of the solvent,

whereas /^assocn. values in molar concentration units

show the opposite trend (see footnotes to table 16).

I' 4.1.2. Aromatic Amines

Solubility determinations [105] and infrared data

[164-1661 furnished part of the early evidence for

i NH...N formation by aniline. Recent results,

^ obtained mainly by spectroscopy (IR, UV, and

NMR), have yielded conflicting conclusions, but

again, those favoring NH...N formation as the

explanation for associative tendencies offer the more convincing arguments.

Table 16. Effect of diluent and temperature on self-association of methylamine (mole fraction units)

Diluent t, °C Kl, 2 K3,

ra-Butane + 15 ''1.93 3.33 4.96 0 2.33 4.08 6.08 — 20 3.06 5.45 8.13 -40 3.88 7.04 10.5

n-Hexane + 20 ^' 1.52 2.56 3.85 0 1.99 3.46 5.17 — 20 2.67 4.72 7.04 -40 3.68 6.62 9.86

re-Nonane + 20 "1.28 1.92 2.88 0 1.68 2.75 4.12 -20 2.32 3.98 5.93 -40 3.28 5.80 8.63

" Table 4 of Wolff" and co-workers (abridged) [161]; from vapor

pressure data. "K (liter/mole) = 0.190.

"^ K (liter/mole) = 0.198. "K (liter/mole) = 0.230.

Reference:

[161] H. Wolf, A. Hiipfner, and H.-M. Hiipfner, Ber. Bunsenges.

Physik. Chem. 68, 410 (1964).

Examples of other aromatic amines for which

NH...N formation has been postulated are:

o- and p-Toluidine, o- and m-chloroaniline, diphenylamine, a-naphthylamine, and A'^-methyl- and A^- ethylaniline [165-166]. Intermolecular NH...N

formation is not expected, of course, nor observed

[165] in the case of orfAo-substituted aromatic

amines like o-nitroaniline and methyl anthranilate,

for which there is very abundant, well-known evi- dence of intramolecular hydrogen bonding (chelation). Moreover, steric factors are important in NH...N formation, as in all other knoAvn types of

hydrogen bonding.

Whetsel and co-workers [167] measured the

near-infrared N-H absorption bands of aniline and

a few other primary aromatic amines in several

solvents (CeHia, CCL,, CHCI3, CS2, CeHe, MeCN) at various concentrations, and observed an appreciable effect of concentration and solvent on the

positions and intensities of the bands. The bearing

of these results on qualitative and quantitative

analyses of aromatic amines by near-infrared spec- troscopy was pointed out. More detailed studies

have since been undertaken to elucidate the molecular interactions which cause the spectral shifts. A recent paper by Lady and Whetsel [168] reports

intensive studies of the first overtone N-H symmetric stretching band of aniline (C6H12. 10 to

70°). The data were analyzed by different methods,

including graphical methods of Rossotti and co- workers [169-171]. For concentrations up to 4 M, the results are consistent with the postulate

21
that monomer coexists with hydrogen bonded

dimer and (cyclic) tetramer. A similar model is thought to be valid for m-toluidine and m-chloroanihne, also. Up to 0.2 M, a simple model involving only dimerization holds fairly well for anihne

The heats of formation found for aniline dimei

and tetramer are about —1.6 and —7.7 kcal/mole,

respectively. The corresponding association constants, which are somewhat more uncertain, are about 0.2 and 3-4 liter/mole [168]. Dimerization

constants of about 0.3 and 0.15 (cyclohexane, liter/ mole) have been estimated for /V-methyl- and A'^- ethylaniline, respectively.

Dearden [172] investigated the self-association

of aniline (C6H12) by ultraviolet spectroscopy. In analyzing the data, he assumed that all associated

species have the same molar absorptivities (which

he considers to be not strictly true), and also assumed that either dimer, trimer, or stepwise

multimer formation takes place; the respective constants obtained are about 2 liter/mole, 80-

(liter/mole)^, and 2.7 liter/mole.

It should be borne in mind that bases like ani- hne, which are sufficiently good proton donors for self-association, are readily hydrogen bonded to other basic species. As an example. Few and Smith

found (in agreement with earlier Russian work)

[173] that the apparent molar polarization of ani- line and methylaniline at infinite dilution is much

greater in dioxane than in benzene, while with

dimethylaniline, the difference is much smaller;

these results were attributed to bonding of amino

hydrogen to oxygen of dioxane, with resulting modification of the charge distribution. In more recent work, Dodd and Stephenson [174] have studied

N-H infrared frequency shifts of amines in sol- vents hke acetone and pyridine, attributing the

shifts to hydrogen bonding, with the aromatic amine

serving as the proton donor. Many other possibilities for hydrogen bonding between primary or secondary

aromatic amino groups and other basic groups are

readily imagined. The interacting groups may, of

course, be present in the same or different molecules.

4.1.3. Amidines, Guanidines, and Miscellaneous

Heterocyclic Bases

Self-association through N-H...N formation

has been studied most extensively for bases of

the types listed in table 17. The references cited

in the table are mostly to recent papers; some of

these review earlier work.

Pyrrole (III) as a proton donor in hydrogen bonding was mentioned earlier (see tables 5 and 14,

fig. 7, and related discussion). Happe concluded

(NMR data, cyclohexane) [178] that pyrrole forms

a dimer in which the NH of one ring is bonded to TT-electrons of a second ring. He estimated Ki,2—4'.S

(m.f.)~' (which is equivalent to about 0.46 liter/ mole), assuming either an open dimer (with one

free NH) or a closed ("cyclic") dimer (no free NH).

This value of Ki, > corresponds to about 1 percent

dimerization in 0.01 M solution. The solutions examined were considered dilute enough to exclude

formation of significant amounts of higher multimers. Lumbroso reported (IR data, CCL4, 21°) [180] a /Ci,2 value of about 0.2 Uter/mole, assuming a

cyclic structure. Dielectric polarization measure- ments yielded the following approximate values

(beUeved to be underestimates): 0.14 (C6H12); 0.09

(CCI4). In benzene or dioxane, the dimerization constant was estimated to be about 0.06, a value

considered unrehable because both solvents are basic enough for hydrogen bonding to pyrrole.

Lumbroso shares Happe's view that in the dimer

NH is bonded to 7r-electrons of a second ring, but

favors the closed structure because it accounts

best for the dielectric polarization data, even though

hydrogen bonds would be Unear in open dimers,

bent in closed dimers.

Several workers have studied the 1:1 association

of pyrrole with other bases (for example, pyridine,

triethylamine, dioxane) [176-180]. The most probable value of /^assocn. (Utcr/mole) for the pyrrolepyridine complex seems to be in the range 2.5 to 2.7 [176-179]. A recently obtained value for AH

of association is —3.2 kcal/mole (25°, CCI4, IR data)

[179]. The hydrogen bond in this complex is probably of the N-H...N type, though some bonding

to the 77-electrons of pyridine is considered possible [178].

Formation of NH...N bonds by pyrazole (IV)

and some of its alkyl or aryl derivatives in aprotic

solvents (CeHe, CCI4, CHCI3, and others) was first studied cryoscopically [181], then by infrared [171,

182] and NMR [183] spectroscopy. Details of the

carefully analyzed results will be found in the papers

cited. It would appear that in solutions of pyrazole

in aprotic solvents (CCI4, CHCI3, CH2CI2, CS2,

CeHe), at stoichiometric concentrations =s 1 A/, the

monomer, cychc dimer, and cycUc trimer coexist. The cyclic trimer seems to have a practically strain- less structure. Derivatives such as 3,5-dialkylpyra- zoles show similar behavior. The extent of association is much greater than in the case of pyrrole.

Saturated solutions appear to contain significant

amounts of multimers up to the decamer. Vinogradov and Kilpatrick [182] found that 3,5-dimethylpyrazole is much less self-associated in benzene

than in carbon tetrachloride; this result is pre- dictable, of course, from the ability of benzene to compete as a proton acceptor (see sec. 3.4.1). Hunter and Marriott, on determining the molecular weights of substituted amidines (V) and imidazoles (VI) (which can be regarded as cyclic amidines) in naphthalene over a range of concentration,

concluded that members of these two classes have

pronounced tendencies toward self-association by

NH...N formation [184]. In accord with their expectations, compounds without an NH group (as, for example, 1,2-dimethylbenzimidazole) are not

associated. Moreover, they observed that imidazoles

have much lower melting points if the 1-position is substituted, indicating weaker intermolecular

forces. From similar measurements, Roberts

[185] concluded that A'^,7V'-diphenylformamidine

22
Table 17. Structures of some self-associating bases

Class of base General structure Formula number

References

Pyrrole

1 2 3 4 5 NH—CH=CH—CH=CH

1 1

(III) [165, 166. 175-180]

Pyrazole

1 2 3 4 5

1 1

(IV) riTi loi looi

[171, 181-183]

Amidine

1 2 R,N—C(R)=NR (V) [181, 184, 1851

Imidazole

1 2 3 4 5 NH—CH=N—CH=CH

1 1

(VI) [169, 171, 181, 184, 186-188]

Guanidine

1 2 3 R,N—C(=NR)—NR. (VII) [181, 184]

References:

165] W. Gordy and S. C. Stanford, J. Am. Chem. Soc. 62, 497 (1940). 166] N. Fuson. M.-L. Josien, R. L. Powell, and E. Utterback, J. Chem. Phys. 20,

145 (1952). 169) F. J. C. Rossotti and H. Rossotti, J. Phys. Chem. 65, 926, 930, 1376 (1961). 171] D. M. W. Anderson, J. L. Duncan, and F. J. C. Rossotti, J. Chem. Soc, p. 140, 4201. 2165 (1961). 175] N. Fuson and M.-L. Josien, J. Chem. Phys. 20, 1043 (1952). 176] S. N. Vinogradov and R. H. Linnell, J. Chem. Phys. 23, 93 (1955). 177] N. Fuson, P. Pineau, and M.-L. Josien, J. Chim. Phys. 55, 454 (1958). 178] J. A. Happe, J. Phys. Chem. 65, 72 (1961). 179] H. J. Wimette and R. H. LinneU, J. Phys. Chem. 66, 546 (1962). 180] H. Lumbroso, J. Chim. Phys. 61, 132 (1964). 181] N. E. White and M. Kilpatrick, J. Phys. Chem. 59, 1044 (1955). 182] S. N. Vinogradov and M. Kilpatrick, J. Phys. Chem. 68, 181 (1964). 183] V. F. Bystrov, I. I. Grandberg, and G. I. Sharova, Opt. Spectry. (USSR) (English Transl.) 17, 31 (1964). 184] L. Hunter and J. A. Marriott, J. Chem. Soc, p. 777 (1941). 185] R. M. Roberts, J. Am. Chem. Soc 72, 3608 (1950). 186] D. G. O'SuUivan, J. Chem. Soc, p. 3278 (1960). 187] H. Zimmermann, Z. Elektrochem. 65, 821 (1961). 188] N. Joop and H. Zimmermann, Z. Elektrochem. 66 , 440, 541 (1962).

1 2 3 -4. 5 NH-CH=CH-CH=CH

1 2 3 A 5 NH-CH=N-CH=CH

I I

(III) (VI)

12 3-15

NH-N=CH-CH=CH

R2N-C(=NR)-Nll2

(IV) (VII)

1 2 R2N-C(R)=NR

(V)

(PhNH-CH=NPh) is extensively associated in benzene at 6° and appreciably associated in naphthalene at 80°. White and Kilpatrick [181] also studied the association of this compound in benzene

cryoscopically, deduced that a cycUc dimer is

23
formed, and calculated the dimerization constant

Ki,2, to be about 69 (kg of solvent/mole).

Self-association of amidines, though initially

Hunter and Marriott's main interest, might have

escaped recognition if the much more highly asso- ciated cyclic analogs, the imidazoles, had not been

included in the investigations. As these authors

pointed out, the pronounced associative tendencies

of the imidazoles are probably due to their more

acid character. Amidines are among the strongest

organic bases. As their structure (V) indicates, the

cation formed on proton addition (at the N-2 position) will be stabilized by resonance. However,

simple amidines like diphenylformamidine and

-acetamidine, though obtainable commercially and

of potential interest as reference bases, have not been very widely studied. Tri-A'^-substituted amidines, not now commerically offered, should not be

measurably associated, and thus might prove to be

especially valuable reference bases for nonaqueous

solutions.

Self-association of imidazoles, and also of the

closely related benzimidazoles, by N-H...N

formation has received considerable attention.

Hunter and Marriott studied more than 30 bases

of this type, finding that 1-alkyl- or 1-aryl-substituted

compounds have normal molecular weights in naphthalene, while the results obtained for all compounds unsubstituted in the 1-position indicate

extensive self-association. A later, more complete

analysis of cryoscopic data (CeHe) for 2-n-butylbenzimidazole [181] indicated very extensive linear

association; this conclusion has been supported by

infrared data for various imidazole derivatives

[169, 171, 186-188].

Anderson and associates [171] investigated the nature of oligomers of imidazole (CCI4). They found

that significant concentrations of all oligomers up

to the dodecamer are present in saturated solutions,

but pointed out that imidazole, in contrast to pyrazole (which, as previously mentioned, forms

cyclic dimers and trimers), cannot form cyclic

oligomers smaller than the decamer, for steric reasons. The linear association of imidazoles is presumed to be stabilized by resonance. It was

found [171] that the hydrogen bonding equilibria

of imidazole (CCI4, 18 ± 1°) can be represented

quantitatively by dimerization, with /Ci, 2 = 234 ±15

hter/mole, and subsequent stepwise N-H...N

formation, with /Cassocn. (for each step) = 760 ±20

liter/mole.

Zimmermann and Joop [187, 188], who have been

interested especially in theoretical aspects of

NH...N bonding in imidazole and methylimidazole, have carefully reviewed and extended

the available information on this topic. Their results,

from varied experimental data (x-ray, IR, Raman,

NMR, isotope effects, dipole moments), support

the other evidence of extensive linear association.

Self-association of 1 ,3-diphenylguanidine (VHI)

has been observed cryoscopically, first in naphthalene (-80°) [184] and later in benzene (-5.5°,

0.005-0.1 molal) [181]. 1,2,3-Triphenylguanidine

1 2 3 PhNH-C(=NH)-NHPh

(VIII)

and 1,1,2,3-tetraphenylguanidine seem to be unassociated in Uquid naphthalene [184]. White and

Kilpatrick concluded from their cryoscopic data for diphenylguanidine in benzene [181] that a linear

dimer and a cycUc trimer form successively, with

the association constants (molal units) Ki,2 — ~53

and A^i,3=~ 290 or 350 (according to the method

used for calculation). The structures which they

suggested for the dimer and trimer would be formed

by hydrogen bonding of the type

PhNH-C (

) =N ( Ph ) , . .H-NH-C (-NHPh ) =NHPh

This type of N-H...N formation .^n be thought

of as arising from a hypothetical tautomeric form

of diphenylguanidine (IX), in which -NH2 is the

1 3 2 PhNH-C(=NPh)-NH2

(IX)

proton donor group, and =NPh the proton acceptor.

It seems more reasonable for PhNH- to serve as the proton donor, and =NH as the proton acceptor.

Considering the scientific and industrial usefulness

of this compound and some other substituted

guanidines, further studies of their associative

behavior are very desirable. An additional need is for compounds such as (X), for example, 1,3-diphen1 2 3 PhNH-C(=m)-NHPh

(X)

yl-2-methylguanidine, to be made easily available

for determinations of their behavior and properties.

The main purpose of the foregoing brief re- view has been to present evidence that N-H...N

bonding, once believed to be non-existent, is of

rather common occurrence. Though comparatively

weak, such bonding must be taken into consideration in interpreting data and planning the most

favorable conditions for experimental studies. Its occurrence seems especially likely when steric factors do not interfere and the structures formed

are stabilized by resonance.

4.2. Self-Association of Acids

Self-association of acids through formation of hydrogen bonds of the type 0-H...0 — sometimes forming linear aggregates and sometimes,

cyclic structures — is one of the most extensively

studied and best known examples of hydrogen

bonding, though much of the work is widely scat- tered through the literature. The first evidence of

self-association of acids in aprotic solvents came

from molecular weight determinations and distribu- tion ratio data [189]; later this was confirmed and

24
augmented by numerous other kinds of physical

and chemical data. The most familiar examples of self-associating

acids are carboxylic acids, which have been studied

in the vapor, liquid, and solid states, as well as in aprotic and more active solvents. In a limited number of cases, monomer dimer equilibrium constants in aprotic solvents have been obtained, and

also related thermodynamic constants. Phenols,

nitric acid, sulfonic acids, and phosphoric acid

esters of the general formulas (R0)P(0)(0H)2 and

(R0)2P(0)0H, where R may be an alkyl or an aryl group, have been found to display analogous

associative tendencies, as would be expected on

structural grounds.

4.2.1. Phenols

The associative behavior of phenols in aprotic

solvents is a subject of importance, for some of

the most important reference materials for aprotic

solvents are derived from phenol. For example,

picric acid (2,4,6-trinitrophenol) is strong enough

to serve as a standard acid in titrations [14]. Methylpicric acid (trinitro-m-cresol) [20], dinitrophenols

[16], mononitrophenols, and various related compounds, which are weaker than picric acid, are valuable (or potentially valuable) acid-base indi- cators and reference acids for measuring the relative strengths of bases. Another very important group

of phenolic compounds is the sulfonephthalein

family of indicators — for example, bromophenol

blue (3' ,3",5' ,5"-tetrahTomophenolsulfon€phthalein)

[29]. StiU another phenoUc indicator of considerable

interest is bromophthalein magenta E (3',3",5',5"-

tetrahromophenolphthalein ethyl ester) [14]. Phenol is a much stronger acid {pKa = 10) than

aliphatic alcohols. Its proton donor capacity, as

well as other chemical and physical properties, are weU accounted for by the resonance hypothesis

[190-191], according to which phenol is a resonance

hybrid of Kekule and quinoid structures. As a result

of this resonance, the C-0 bond has partial

double-bond character, and the benzene ring and

OH group tend to lie in the same plane. This planar

structure leads to geometrical isomerism in ortho- substituted phenols (see formulas XI and XII).

H

(cis) (trans)

j (XI) (XII)

IHowever, in phenol itself and symmetrical 2,6-

y disubstituted phenols, the cis and trans isomers

'iare equivalent. Self-associating tendencies are in- Ijfluenced by the nature of any substituents ortho to the OH group. The relative donor capacities of phenols in hydro-

^igen bond formation are expected, of course, to

I parallel their capacities for proton transfer to a

base, and thus to increase as pKa (—log Ka) values

diminish.

a. Self-association of phenol in aprotic solvents. Table 18 is a compilation of equilibrium constants

for self-association of phenol in several aprotic

solvents, obtained by different methods over a

period of 30 years. It is evident that most of the

investigators represented have favored the assumption of a monomer-dimer equilibrium at the lowest

concentrations studied. However, Saunders and

Hyne [196, 199] consider a six-membered cyclic

trimer to be a Ukely structure, and think that their data (for 0.01-5 M solutions) are best represented

by a monomer-trimer equihbrium. Furthermore,

they believe that the NMR data of Huggins and

coworkers [195] were not obtained at low enough

concentrations for application of the "limiting

slope" method used in calculating their Ki,2 value. Saunders and Hyne do not dispute the existence of

a monomer-dimer equilibrium in solutions too dilute

for measurement by present NMR techniques. In agreement with most workers, they think it likely that after the first few steps in multimerization the

successive constants approach a limit or vary

uniformly.

CoggeshaU and Saier [194], who based their cal- culations (for solutions about 0.06-0.95 M) on the

assumption of different values for Ki, 2 and Kn, „ + i, found that various alcohols and alkylphenols have

constants of the same general magnitude.

Davison [200] determined Ki,2 values for o- and

p-cresol and for m- and p-nitrophenol in ben- zene cryoscopically (~ 5 °C). The relative values

foUow the same order as the acidities, that is, o-Me = p-Me < m-NOz ^ P-NO2. Other phenols for which monomer-dimer constants have been cal- culated include m-cresol (C6H12) [172], p-cresol

(CsHia, CCI4) [172, 198] and 4-chlorophenol [198].

Interactions of phenol with various aprotic sol- vents which were referred to earlier (sec. 3.4) offer a ready explanation for any small solvent effects on

its self-association. Philbrick [192] attributed the

comparatively low self-association of phenol in nitrobenzene (table 18) to its association with the

solvent in the manner PhOH...ON(0)Ph. The

distribution method which he used for obtaining

dimerization constants is now known to lead to erroneous results in many cases because of the

water carried into the organic layer. In some in- stances, there is increased association, and in others, decreased association.

StiU another factor which affects association constants is temperature control. As would be

expected, the extent of association decreases

with rising temperature, and vice versa. Considering

the various factors mentioned, the values in table 18 seem in reasonably good agreement when expressed

in the same concentration units, but additional

systematic study is clearly desirable.

b. Hindered phenols. The discovery of hindered

phenols by Stillson and co-workers [201] has in25
Table 18. Equilibrium constants for self-association of phenol in various aprotic solvents "

Solvent Temp. Method " Kn. Reference

CeHsCHs

CeHsCl C6H5NO2

CsHe

CCL, CCI4

CCL,

CCI4

CCI4

ra-CeHn CCI4

CeHia

(°C) 25

2525

25

25

Room

28

21 Room

Room

30 Room

D

D

D

D

VPL

IR

IR NMR

NMR

UV

UV

IR UV

0.575

0.843 0.648

0.196

'0.570

1.39 M3±7

"2-5

"4.5

'7.22

1.28

4.78

15.4

2.26

2.94

1.45

[192]

[1921 [192]

[192]

[193]

[193a]

[194] [195]

[196]

[197]

[197]

[198]

[172]

Experiments are listed in chronological order. They did not all cover the same concentration

range.

" D, distribution between aq. and org. layers; VPL, vapor pressure lowering; other symbols refer

to famiUar methods of absorption spectroscopy. For the reaction 2 PhOH ^ (PhOH)2; Uter/mole unless otherwise indicated.

" For the reaction 3 PhOH ^ (PhOHjs; literVmole-'. •'For the reaction (PhOH)„ + PhOH ^ (PhOH)„+,; molar units,

f AW ==-2.4 kcal/mole. ^(Mole fraction)"'; 5 percent by vol. of cyclohexane was present as reference solute.

*" Provisional values. '(Mole fraction)"'; AH = — 5A kcal/mole. The dimer is the predominant multimer up to about

0.2 M.

References

[192] F. A. Philbrick, J. Am. Chem. Soc. 56, 2581 (1934). [193] E. N. Lassettre and R. G. Dickinson, J. Am. Chem. Soc. 61, 54 (1939). [193a] H. Kempter and R. Mecke, Z. Physik. Chem. B46, 229

(1940). [194] N. D. Coggeshall and E. L. Saier, J. Am. Chem. Soc. 73,

5414 (1951). [195] C. M. Huggins. G. C. Pimentel, and J. N. Shoolery, J. Phys.

Chem. 60, 1311 (1956). [196] M. Saunders and J. B. Hyne, J. Chem. Phys. 29, 1319

(1958); 31, 270 (1959). [197] M. Ito, J. Mol. Spectry. 4, 125 (1960). [198] M. M. Maguire and R. West, Spectrochim. Acta 17, 369

(1961). [172] J. C. Dearden, Can. J. Chem. 41, 2683 (1963).

spired numerous studies of their spectroscopic

(IR, UV, NMR) behavior [194, 202-213]. These

phenols, which contain a bulky group (most commonly, t-Bu) in the 2- and 6-positions, do not give

the conventional chemical tests and reactions of

phenols. In pursuing their study by IR spectroscopy,

Coggeshall devised (see table 19) [202] a useful

classification of phenols as "unhindered," "partially hindered," or "hindered" on the basis of the

shifts of the phenohc OH band on self-association

(CCI4). Since it is generally agreed that shifts of

the 0-H stretching frequency measure the extent

of hydrogen bonding interactions, the very small'

shifts observed (table 19) for the chemically inactive

(hindered) phenols furnish a direct indication! that bulky ortho-groups obstruct hydrogen bonding.i

It has been found, as anticipated, that in some partially hindered phenols self-association does not seem to proceed beyond the dimer stage, while in hindered phenols, even dimerization does not occur. Coggeshall's infrared data have been confirmed t and extended by other workers [204, 205, 207-11]

Although 3- and 5-alkyl groups have no direct sterit effect on reactions of phenolic OH, they reinforce

26
Table 19. Classification of orxho-alkylphenols according to infrared spectral shifts observed on selfassociation (CCLi)

Nature of Alkyl Substituent in AX, /X Classification of the phenol

2-Position 6-Position

None or small None or small >0.15 Unhindered. Large None or small 0.04-0.15 Partially hindered. Large Large 0.04 Hindered.

^ From Coggeshall [202].

Reference:

[202] N. D. CoggeshaU, J. Am. Chem. Soc. 69, 1620 (1947).

the hindering effect of alltyl groups in the 2- and

6-positions. As an example, in the classification scheme of Sears and Kitchen [204], the relative "hydrogen bonding indexes" range from 0 (for complete hindrance, as in 2,6-if-octyl-4-methylphenoi) to 1 (for no hindrance, as in phenol); on this scale, the indexes for 2,6-dimethyl-, 2,4,6-trimethyl-,

2,3,6-trimethyl-, and 2,3,5,6-tetramethylphenols fall off in the sequence 0.33, 0.29, 0.20, and 0.14. This

is an example of what has been called the "buttressing effect" [214]. The 3- and 5-methyl groups are presumed to force the 2- and 6-methyl groups closer

to OH, thereby enhancing their hindering effects on OH.

Although the buttressing effects seem to be well

established, it should not be forgotten that methyl

(and other alkyl) substitutions tend to inhibit association because they reduce slightly the proton donor

ability of phenol, as shown, for example, by the

aqueous pK values for the cresols [215]. Hunter has

emphasized an additional reason for not attributing

chemical unreactivity and abnormal spectral be- havior in 2,6-dialkyl-phenols entirely to physical

obstruction of the approach of phenolic OH to other reactive groups [206]. In his view, bulky ortho

substituents can prevent the OH group from achieving planarity with the benzene ring, thus inhibiting

the resonance which makes phenols so much more

strongly acidic than alcohols; in consequence,

hindered phenols are "reduced to the status of

tertiary alcohols." His opinion is not shared by

Puttnam [209] or by BeUamy and WiUiams [210],

who consider the OH group in hindered phenols to be largely coplanar with the ring. From studying

effects of numerous proton acceptor solvents on the

infrared spectra of hindered and unhindered

phenols, BeUamy and WiUiams concluded that association of OH with other reactive groups, which

occurs very rapidly in the case of phenol, is retarded

when bulky groups are placed adjacent to OH,

because coUisions suitably oriented for association are then less frequent.

Puttnam's recent reexamination and extension of some of the earher IR work on phenols, using a

higher resolution spectrometer [209], yielded conclusions which agree very substantiaUy with those

from earlier studies. His principal conclusions with

respect to alkyl-substituted phenols may be summed

up as foUows:

(1) In all phenols, the OH group is coplanar with

the aromatic ring.

(2) A single alkyl (R-) substitution — other than

f-Bu — adjacent to the OH group has only a slight effect on hydrogen bond formation.

(3) R groups in both the 2- and 6-positions cause

a reduction in hydrogen bond formation. However,

for R's other than t-Bu, a small amount of higher

multimers tends to coexist with the monomer and

dimer.

(4) The hindering effect of ortho R's increases

with increasing branching. Phenols containing ^-Bu

in both ortho positions exist practically exclusively

in the monomeric form.

These conclusions agree very weU with the association constants for some hindered and partiaUy

hindered phenols (CCLj) which are coUected in table 20.

Table 20. Association constants of some hindered and partially hindered phenols (CCLi, room temp.) "

Substituents present

5-Bu

t-Bu

t-Bu

f-Bu

i-Pr

i-Pr t-Bu

t-Bu

t-Bu

4-

t-Bu

Me

t-Bu

i-Pr Me

t-Bu

Method

IRIR

IR

IR

IR

IRNMR

NMR

NMR

NMR

NMR

1.39

1.72

1.00

1.54 0.96

1.04

1.7

1.3 1.0 « 0.05 0.05

K.n,n+l

2.94

2.94

0.72

"No substituents were in the 3- or 5-positions. IR results are I from "CoggeshaU and Saier [194]; in the original article, the numerical values given corresponded to dissociation constants. NMR results are from Somers and Gutowsky [212].

" For the reaction 2 PhOH (PhOH)2, where PhOH signifies the substituted phenol. Concn. units: IR results, liter/mole; NMR, (mole fraction)*".

' For the reaction (PhOH)„ -I- PhOH ^ (PhOH)„+i.

" Bellamy and Williams [210] concluded (IR data) that this compound is a monomer even when undiluted.

To explain why phenols containing halogen or methoxy in the 2-position have two infrared absorption bands in place of the single band observed

for monomeric phenol and symmetrical 2,6-disubstituted phenols [216], Pauling postulated the

existence of cis- and trans-isomers (XI and XII)

[190]. Double bands were not observed for thymol

280-613 OL-68—

27
and o-cresol [216]. However, Puttnam [209] detected

two absorption bands for 2-f-butylphenol, and con- cluded that the configuration is mainly trans-; he

also detected two bands for 2-methyl-6-t-butylphenol.

Ingold [211] has concluded (IR frequencies and

intensities) that 2-alkylphenols exist in the cis- and

?rares-configurations both in solution and in the

vapor state.

c. Intramolecular hydrogen bonding in phenols

[190. 206, 209, 216-24]. Chelate ring formation by

hydrogen bonding is too well known a phenomenon

to require detailed treatment. Ortho-nitro- and ortho- halophenols are among the best known examples of

chelated structures, the tendency toward chelation

being especially strong in o-nitrophenols because of resonance as well as favorable geometry. The

hydroxy and methoxy groups are other protonacceptor groups with recognized capacities to form

stable rings by bonding to phenohc hydrogen. In

general, 6-membered rings are the most stable,

though 5- and 7-membered rings also occur. Obviously, miramolecular hydrogen bonding must occur

at the expense of mfermolecular hydrogen bonding.

Chelate ring formation stabilizes a phenol molecule in the cis-configuration (XI). Jaffe has pointed

out that the entropy change on intramolecular

hydrogen bond formation in molecules such as o-chlorophenol is practically negligible, because of

the favorable geometry, and A// and AG are essentially equal. Apparently this is one reason why very

weak intramolecular hydrogen bonds can form in preference to much stronger intermolecular bonds, even in highly polar solvents [225]. Infrared spectroscopy has contributed the largest

volume of systematic knowledge about inter- and

intramolecular hydrogen bonding by phenols. Commonly, though not exclusively, the studies of

bonding involve observations and interpretations of

changes in frequency and/or intensity of the

0-H stretching vibration in the 3-micron region of the spectrum, which occur on changes of solvent or on various substitutions in the phenol molecule. It has already been mentioned that the more acidic

the phenol, the more likely it is to become hydrogen

bonded to electron donors. The stronger the hydrogen bond formed, the greater the change in fre- quency and intensity. This is illustrated (for changes

in frequency) by the data in table 21, which are from

work by Ingraham, Corse, and co-workers [226]. These authors found that the spectral shifts are

linearly related to the Hammett cr values [227] by

the following equation: v=v» + p„(T. (28)

In eq 28 vo and v are the respective frequencies of

the unsubstituted and the m- or p-substituted phenol;

(J and p are, respectively, the Hammett substituent

and reaction constants. Puttnam [209] found that the frequency shifts for phenols with one ortho- substituent are similarly linearly related to Hammett

cr values, provided the frequencies used are those

for the irans-isomers.

Table 21. Hydroxyl frequencies {cm-') of some 3- and 4-substituted phenols in CCU and C2CI4 ^

Substituent CCI4 C2CI4

4-MeO 3615.6 3616.0 4-OH 3615.6 U-Bu 3612.1 3609.7 4-Me 3612.8 3611.7 3,5-Me2 3611.7 3609.7

3-Me 3611.3 3610.1 None 3610.1 3610.1 4-Ph 3609.3 3607.4 3-OH ,3610.1 3608.9

4-Cl 3608.5 3607.0

4-Br 3607.4 3606.2 3-NO2 3599.9 3597.6 4-NO2 3594.5 3591.1

" Excerpts from Table I of Ingraham, Corse, and co-workers

[226].

Reference:

[226] i L. L. Ingraham, J. Corse, G. F. Bailey, and F. Stitt, J. Am.

Chem. Soc. 74, 2297 (1952).

In intermolecular hydrogen bonding, the infra- red OH frequencies shift with dilution, as hydrogenbonded species give place to the monomeric form.

Intramolecular hydrogen bonding can be recognized

by the constancy of frequencies on dilution. Intermolecular hydrogen bonding of phenol to 77-electron donors such as aromatic and unsaturated

aliphatic hydrocarbons was discussed earlier (see tables 5 and 7). As would be anticipated, intra- molecular bonding of similar types has been

detected — for example, in o-phenylphenol [216] and

o-allylphenol [223]. o-Hydroxybenzonitrile (o-cyanophenol) is in- capable of forming a chelate ring because of

unfavorable geometry — the -OH and -CN groups

are too far apart. Consequently, it resembles its m- and p-isomers in chemical and physical behavior

[228, 206, 220].

Table 22 contains examples of results obtained

by Baker and Shulgin [223] in measuring the OH

frequency shifts produced by various ortho substi- tutions in the phenol molecule. From their data

they concluded that the strengths of intramolecular

hydrogen bonds between OH and ortho groups

increase generally in the order:

F < O < Cl< Br < I< S < =0 < N. The increase in from 346 cm"' for o-nitrophenol

to 437 cm"' for 2-nitro-3,5-dimethylphenol was

explained as a buttressing effect which increases

the strength of the hydrogen bond. The decrease

to 82 cm"' for 2-nitro-3,4,6-trichlorophenol was

attributed to a steric interaction large enough to 28
Table 22. IR frequency shifts ofphenols on hydrogen bonding to oTtho-substituent groups {CCh, 0.7-0.0007 M) "

Substituents and their position

Ai;, 0—H,

2- 3- 4- 5- 6-

F

(cm-') 185256

6163

63

67

78

83

83

85

105 244

329

346

388

437 "82

.OMe.

CI

/"II /- T T / - T T

CI CI

CI LI Cl

15- CI NO2

Br

L.n2C^l NO2

rMeaN

MejN Me Me

NO2

NO2 NO2

NO2 Me

CI Me

NO2 CI CI

_

Excerpts from Table II, Baker and Shulgin [223].

" Attributed to K. Bradley.

Reference:

[223] A. W. Baker and A. T. Shulgin, J. Am. Chem. Soc. 80,

5358 (1958).

displace the nitro group from the plane of the ring,

thereby lengthening the distance between OH and

NO2 and reducing the electron donor capacity of

the NO2 group, with a consequent weakening of the

intramolecular hydrogen bond. Later, Baker and

Kaeding [222] revised the relative order of the four

o-halophenols as follows: I< F < Br < CI.

They based the new order on spectral data obtained for six 2,6-halophenols having different halogen atoms in the 2- and 6-positions. The

authors' revised explanation, which will not be

repeated here, takes into account repulsive energy

due to overlapjpin^ of the halogen atoms' lone pair

orbitals and the 0-H bonding orbital. It is note- worthy that solutions (in hexachloro-l,3-butadiene)

of the unsymmetrical 2,6-dihalophenols exhibited

a pair of absorption bands.

In an early estimate, based on the ratio of the

areas of the IR absorption bands (CCI4), Pauling

[190] concluded that the cis and trans forms of

o-chlorophenol are present in the ratio 10/1. Baker

[221], using exhaustively purified materials, has

arrived at the following cis/trans ratios for o-halophenols:

o-Cl, 56/1; o-Br, 38/1; o-I, 13.5/1.

Allan and Reeves, from the chemical shift parameter

in NIVIR spectroscopy [229], have arrived at the same

cis/trans ratios for o-chloro- and o-bromophenols

(CS2 solutions, 25°), and a slightly different value

(19/1) for o-iodophenol. (In the calculations, solvation of the trans-form was considered to be negligible.) In agreement with Baker [221], the enthalpy

of cis-trans isomerism was found to be 2-3 kcal/mole.

Like others, Allan and Reeves find no evidence of an intramolecular hydrogen bond in o-fluorophenol;

the attractive force maintaining it in the cis configuration is considered to be electrostatic in nature,

because the fluorine radius is too small for bonding

to phenolic hydrogen.

In studying temperature effects on hydrogen

bonding. Finch and Lippincott [230] observed that the higher frequency band for o-cUorophenol (that

is, the band corresponding to the mframolecularly

bonded OH frequency) remained constant both in position and intensity in the temperature range

2-25 °C while, in contrast, the miermolecularly

bonded OH frequency decreased 36 cm~' when the

temperature dropped from 27 to — 3 °C. The absorption curves (for CCI4 solutions, concentration not

stated) are shown in figure 8. The changes in fre0

20

X 40

60

80

"^^"^^^^^2

9 8 ° K r

283° A\

275° AV/Ol

nUainterO-CHLOROPHENOL

1 1 1 I 3000 3200 3400 3600

cm"'

Figure 8. Infrared absorption bands of intermolecularly and

intramolecularly bonded OH in o-chlorophenol (CCLj) at three temperatures; it will be seen that the position and intensity of

the "inter-" band vary with temperature, whereas position and

intensity of the "intra-" band remain constant.

Reproduced from J. N. Finch and E. fi. Lippincott, J. Phys. Chem. 61, 894 (1957), with the permission of the authors and publisher. (Ref. No. 230.) quency and intensity of absorption for the trans form are in the direction expected on lowering the

temperature, because of increased intermolecular

hydrogen bonding.

There have been numerous other IR investigations of o-halophenols, including several very recent

studies [231-233]. Brown and co-workers [231],

29
who are interested in the comparative biological

effects of various ortAohalophenols, have been

investigating systematically electronic, steric, and

solvent effects on the position and intensity of the

0-fI stretching absorptions in numerous ortho- bromophenols possessing additional substituent groups. Such studies have potential value to those

employing IR spectroscopy in analyses, as well as

to those interested in developing additional, or

better, acid-base indicators.

The numerous conclusions of Brown and co- w^orkers include the following:

(1) A methyl group in the 6-position of an o-bromophenol causes the molecule to exist solely in the

chelated conformation.

(2) The phenol group in o-bromophenols shows

a preference for intramolecular OH... Br formation

over intermolecular bonding to ether (0H...0Et2).

(3) Examination of o-bromophenol in CCI4 and

MeCN, singly and mixed, indicates that the di- electric constant of the solvent is of minor

importance.

There is disagreement about the relative abihties

of the four halogens to become bonded to ortho-OW.

Thus, Bourassa-Bataille and co-workers [232].

from measuring the IR spectra of the four o-halo- phenols in CCl4-Et20 mixtures, decided on the

relative order,

F < I< Br < CI,

while Jones and Watkinson [233] decided on the

relative order (CCl2=CCl2 solutions)

I< Cl< Br.

Both of these series disagree in part with the order

deduced by Baker and Kaeding (see above).

Information about associative tendencies in phenols is needed for the proper selection of

phenolic indicators for quantitative spectral studies. In view of the preceding discussion, 2,6-dibromophenols would be expected to be monomeric in solvents like benzene, and their spectral behavior

should not be complicated by cis-trans isomerism.

For such reasons, Davis and Schuhmann [14], when performing spectrophotometric acid-base

studies in benzene with the phenolic indicator,

bromophthalein magenta E (which can be thought

of as a substituted 2,6-dibromophenol), considered

it justifiable to treat it as a monomeric species. In actuality, no deviation from Beer's law was

detected in the concentration range employed

(1.6 X 10-5 to 3 X 10" M).

Effects of additional substituents on the strength

of the intramolecular hydrogen bond in o-nitro- phenol is another subject of continuing interest. The investigations frequently include studies of

the persistence of the chelate structure when

various polar solvents are added to the aprotic

medium. Some of the earlier evidence was reviewed

briefly by Davis [16]. The work cited and additional

results of Davis indicate that in aprotic solvents

strong intramolecular hydrogen bonds form in o-nitrophenol itself and in the dinitrophenols which

contain an o-nitro group (that is, the 2,4-, 2,5-, and

2,6-isomers).

Baker and Shulgin's work [223] on IR spectral

effects of substitutions in o-nitrophenol was mentioned a little earlier (see table 22). In another recent

study, Dearden and Forbes [235] obtained ultra- violet absorbance data for o-nitrophenol and also made comparisons with the m- and p-isomers. The

chelate structure of o-nitrophenol was found to persist on changing the solvent from cyclohexane

to ether or ethanol — only very minor shifts occurred

in the positions of the absorption bands. Ultraviolet shifts (expressed in wave numbers) were found

to parallel approximately IR results previously

reported. Von Keussler and Rossmy [234] and

Cardinaud [236] have studied the comparative

behavior of some phenols and deuterophenols, and

discussed theoretical implications of the results. From comparisons of the IR spectral behavior of

the three nitrophenols (CCI4), Cardinaud deduced

that both 0-NO2C6H4OH and 0-NO2C6H4OD exist only as the cis, chelated form; there was no trace

of a band for free OH or OD. The absorption of the

chelated form was not perturbed significantly by

addition of dioxane. He suggests that p-nitrophenol

tends to form dimers by "head-to-tail" association

in which both nitro groups are bonded to an OH

group. In contrast, he considers that m-nitrophenol

behaves analogously to phenol, forming noncyclic aggregates. The solubility of both m- and

p-nitrophenol in CCI4 is too little for aggregation to proceed beyond the dimer. Dioxane greatly en- hances the solubility of p-nitrophenol in CCI4, and

is thought to associate with it in the manner,

Nyquist [237] has studied the absorption of various

or^/io-substituted phenols in the 7(0-H) (out-ofplane) frequency region. As indicated in other

studies, a large substituent in the 6-position of

o-nitrophenol seems to enhance the strength of the

intramolecular hydrogen bond by bending the OH

group closer to NO2.

The conclusions from spectroscopic studies of

nitrophenols merit wide dissemination. If nitro- phenols are to continue playing an important role as acid-base indicators and reference acids in non- aqueous and aqueous solutions, there will be need

for increased consideration of factors which affect their proton donor ability and spectral behavior.

For example, it seems doubtful that 2,4-dinitrophenol, which has been used as a reference acid in a considerable volume of work in nonaqueous sol- vents, for studying the comparative strengths of

nitrogen-containing bases, is the most suitable

indicator obtainable.

d. The Deuterium Isotope Effect in Phenols.

Effects of replacing hydrogen by deuterium are 1 discussed briefly in section 4.2.6. )

30
4.2.2. Carboxylic Acids

a. General discussion. It has long been known that carboxylic acids, both aliphatic and aromatic, have

a pronounced tendency to dimerize, in the vapor

state as well as in aprotic solvents. In fact, from some of the early studies, using relatively concentrated solutions (about 0.01 m and stronger), it was concluded that only double molecules were

present [247]. Later work with more dilute solu- tions revealed the existence of a monomer-dimer

equilibrium, and attempts were made to evaluate

equilibrium constants. Possible formation of aggregates higher than the

dimeric form has not been excluded, and the for- mation of trimers, in particular, has sometimes

been considered. However, the monomer and dimer

are commonly regarded as being the predominant

species in aprotic solvents. Generally, dimeric

carboxylic acids are assumed to be cyclic, though

occasionally it is suggested that part of the molecules are "open."

A large proportion of the investigations have in- cluded benzoic acid. Formulas indicating the three

most likely conformations of the monomeric acid

are shown below (Xllla, Xlllb, and XIV). Con-

«\ /« bo 0

(XIII:0 (.ailb) (XIV)

figurations (Xllla) and (Xlllb) are identical in benzoic acid itself, but are nonidentical in un- symmetrically substituted benzoic acids. Infrared

evidence supports (XIII) [238-241]. Mansel Davies

[242] indicates a hydrogen bond between the car- bonyl groups; however, the two oxygens are not considered to be equivalent. When the carboxyl

proton can become bonded to an ortho substituent,

as in o-methoxybenzoic acid (XV) or o-hydroxybenzoic (salicylic) acid, configuration (XIV) is favored.

C H

(XV) The accepted structure for most carboxylic

acid dimers is indicated by (XVI); the 0-H...0

bonds are considered to be linear. This has sometimes led to the conclusion that hydrogen bonds

must always be linear, but there are many known

cases of "bent" hydrogen bonds, notably in the case of intramolecular hydrogen bonding. A well- known example is dimeric sahcylic acid (XVII), in which there is simultaneous intra- and inter- molecular hydrogen bonding [190b]. Schneider, in considering the requirements for a stable hydrogen

(XVI) (XVII) bond, X-H...Y, to form between a proton donor

group XH and an electron donor group Y, con- cluded that hydrogen bonding will be strongest

when the direction of the lone pair orbital of Y

and the direction of the hydrogen bond are col- linear [243]. The only requirement for XH is that the charge distribution of its orbital must leave the

proton sufficiently exposed. These requirements

are well met in the dimerization of carboxylic acids.

The groups composing (XVI) are considered to

lie essentially in the same plane, and it is apparent

that conformational isomerism can be looked for

in the dimers of unsymmetrically substituted benzoic acids. For example. Brooks and co-workers

[244] have discussed conformational isomerism in the dimer of o-methoxybenzoic acid. The chelate structure of salicylic acid, stabilized

by resonance, is thought to persist in the salicylate

ion. Branch and Yabroff [245] proposed this as an

explanation of the unexpectedly great strength of

salicylic acid in water, and subsequently Baker

[246] offered an analogous explanation (involving a doubly chelated mesomeric structure) for the still greater strength of 2,6-dihydroxybenzoic acid in aqueous solutions. Such chelate structures should

be at least equally stable in aprotic solvents. In accord with this expectation, Davis and Hetzer [32]

found both salicylic acid and p-orsellinic (2,6- dihydroxy-4-methylbenzoic) acid to behave as strong

acids in benzene by the criterion of their tendencies

to associate with diphenylguanidine.

b. Dimerization constants. Probably no equilibrium process has been studied by a greater variety

of physical methods than the monomer-dimer equihbrium of carboxylic acids (eq 29).

2 RCOOH ^ (RC00H)2. (29)

The following examples, selected mainly from re- cent work which yielded equilibrium constants

as well as qualitative evidence, illustrate the di- versified approaches: Cryoscopy [247, 248], ebul- lioscopy at atmospheric pressure [249] and at re- duced pressures [250], isopiestic data [251, 252]

acid catalysis of the inversion of menthone [253],

spectroscopy in the visible region [254, 254a], dielectric polarization data [255-258, 267b], distribution

between immiscible solvents [259-260], heat of

dilution data [261], ultrasonic absorption [262, 263],

infrared absorption [249, 264-267], and ultraviolet

absorption [268, 269]. NMR spectroscopy [270] has

also been applied in studying carboxylic acid dimerization, but no association constants appear to have

been reported so far. Useful summaries of attempts to measure the

self-association of carboxylic acids were included

in papers by LeFevre and Vine [255] and Allen and

31
Caldin [250b], covering literature up to about 1953. The equilibrium constant for dimerization (eq

29) is most commonly symbolized by Ku 2 (or Kvz). The rather confusing practice in some of the earlier literature of referring to the reciprocal of the dimerization constant, K2, 1 (which corresponds to the

process (RC00H)2 ^ 2 RCOOH), as the "dimerization constant" has been largely discontinued.

An alternative symbol, Ka, has been used at times

for the dimerization constant, but is increasingly

appHed, instead, to distribution of acids between

aqueous and aprotic solvents. The distribution method for determining Ki,2

values was widely used in the past but has been

criticized increasingly on the ground that dimerization tendencies cannot be assessed accurately

in wet solvents. Some 30 years ago, Fredenhagen

[272] observed that cryoscopic results for solutions

of organic acids dissolved in benzene which had

not been carefully dried did not indicate the ex- pected double molecule formation until the molar

concentration of the acid exceeded that of the water

present in the benzene. In similar experiments

Bell and Arnold [247] studied the cryoscopic behavior of trichloroacetic acid in dry benzene and

in an equimolar mixture with water in benzene.

The results indicated that a hydrate of the composition CCI3COOH • H2O forms in wet benzene in preference to the acid dimer. Wenograd and Spurr

[260] determined the dimerization of several ali- phatic acids in both dry and wet carbon tetrachloride, by infrared spectroscopy and distribution data,

respectively. Their results indicated that the values

obtained by the distribution method do not apply

to acid association in the dry solvent. Christian

and coworkers [273] have strongly criticized the

distribution method, adding the following statement: "We also wish to emphasize that the presence of

tiny amounts of dissolved water can lead to erron- eous results in the case of other methods for de- termining association constants for carboxylic

acids. In fact, we suspect that a major reason for the lack of agreement between various literature values of association constants of carboxylic acids

in solution is that small, variable amounts of dis- solved water have been present in most of the sys- tems investigated." Their work indicated that acetic acid, like trichloroacetic acid [247], forms

the monomer monohydrate in preference to the

anhydrous dimer when dissolved in wet benzene.

Of the many dimerization constants which have

been determined, relatively few are on a comparable

basis. Differing aprotic solvents have been used —

benzene, toluene, chlorobenzene, cyclohexane,

heptane, carbon tetrachloride, chloroform, and

others — and temperatures of measurement have

varied from the freezing point to the boiUng point

of the solvent. However, the most recent measure- ments have been made either at (or near) 25° or at a series of temperatures which permit the calculation of values for 25°. Table 23 contains results obtained for benzoic acid and some aUphatic acids

Table 23. Log Ki,2 values for miscellaneous carboxylic acids (CCU, 25°)

Acid pK,.. Water"

Log Ki.2 (liter/mole)" Reference

Dry sol- vent (IR) Wet sol- vent (D)

Benzoic 4.20

0.64

1.26 2.86

4.31 4.76

4.84

4.87

3.37

2.43 2.72 3.04

3.20

3.29

3.38

3.30

3.24

3.44

3.39

3.27 3.29

[267a]

[267]

[264]

[264]

[264]

[260]

[264]

[260]

[267, 259]

[266]

[260]

[260. 259]

[260]

[260]

Trichloroacetic

Dichloroacetic Chloroacetic Phenylacetic

Acetic

Valeric Propionic

Caproic

2.51

3.20

2.85; 2.65

2.99 Laurie

" pK„ values are from the compilation by Robinson and

Stokes [271]. IR, infrared spectroscopy; D, distribution between immiscible solvents.

References:

[259] M. Davies, P. Jones, D. Patnaik, and E. A. MoelwynHughes, J. Chem. Soc, p. 1249 (1951). [260] (a) J. Wenograd and R. A. Spurr, J. Am. Chem. Soc. 79,

5844 (1957); (b) J. Wenograd, Dissertation Abstr. 16, 38

(1956). [264] J. T. Harris, Jr. and M. E. Hobbs, J. Am. Chem. Soc. 76,

1419 (1954). [266] J. J. Lindberg and C. Majani, S uomen Kemistilehti B 37,

21 (1964). [267] H. E. Affsprung. S. D. Christian, and A. M. Melnick,

Spectrochim, Acta 20, 285 (1964). [267a) Y. I'Haya and T. Shibuya, BuU. Chem. Soc. Japan 38,

1144 (1965). [271] R. A. Robinson and R. H. Stokes, Electrolyte Solutions, 2d ed. Academic Press, New York, 1959.

in carbon tetrachloride solutions (25°) during the

past 10 to 15 years. The comparative tendencies

for dimerization are expressed as log Ki,2 values

for presumably water-free solutions (see 3d column).

In some cases log Ki,2 values were also obtained

for water-containing solutions (column 4). The lat- ter values are all distinctly lower than those for "dry" carbon tetrachloride solutions, supporting

the view that hydrogen bonding to water reduces

the extent of self-association through hydrogen

bond formation. In the two cases (trichloroacetic

and acetic acids) where two or more log Ki, 2 values

are available, the agreement is not good. However,

two generalizations appear justifiable: (1) In comparing the log Ki,2 values for carbon tetrachloride

solutions with the corresponding tendencies for ionic dissociation in water — see pKa (~log Ka)

32
values in column 2 -it seems evident that the A greater range of log K\,2 values is available

weaker the acid in water, the greater the tendency for benzene solutions of benzoic and substituted

for self-association in carbon tetrachloride. (2) The benzoic acids; see table 24, column 4. As mentioned

pK„ values cover a greater numerical range than in the footnotes to this table, calculations were

the log K\,2 values. made where needed to convert the log K\^2 values

Table 24. Comparative values o/ pK;,, K", and log K|, 2 /or some carboxylic acids {25°}

Acid Log K , Log Ki, 2, Method Keterence Water D h benzene Benzene

{literj mote) (literjmote) X Ul 1.00 p jZOOJ (^A 1 |zooJ 9 17 1 09 I1 Phenylpropiolic e 0 0 1 OCX |250] A 1 7 9 £^9 Z.OZ r |ZOOj

9/11 T Z.oO 9 fKf\ Ao.zz99 9 c^l Z.Ol p

T

1 [zozj 0-Lhlorobenzoic J 9 0/1 o.Uo [250] 2.37 P [258] 2.41 I [252]

o-r luorobenzoic ^ 3.27 5.77 2.48 P [258] 2.68 I [252]

Q 0

o.ol O.UO 2.62 T1 [252] 0.00 O.UO 9 A9 z.Oz T

1 roc 01 [ZdZ\ O.oD O.UO 9 Al Z.Ol TTH, [zoUJ 9Z.Oo T

1 [zozJ

tn-r luorobenzoic o.oO 9 AC Z.OO TPIn [zozJ o-Toluic 3.91 4.94 2.63 I [251a] p-Bromobenzoic 6.\fl 0.00 2.82 T

1 [252] p-Iodobenzoic o.yo

c oc D.OO 2.74 T1 [252] p-Chlorobenzoic 3.98 5.82 2.83 I [252] p-Fluorobenzoic 4.14 5.61 2.78 IR [252] 4.20 5.26; 2.79 I [251a]

'5.32 2.75 E [250] 2.84 P [258] 2.82 I [252] 2.80 U [263] m-Toluic 4.27 5.13 2.86 I [251a] p-Toluic 4.37 5.03 3.38 E [250] p-Methoxybenzoic 4.47 4.92 2.75 E [250] o-Methoxybenzoic 4.09 3.7 0.59 I [252]

"From compilations in [32] and [271]. Log K" values are for association of the carboxylic acid with diphenylguanidine; see Davis and Hetzer [32].

' Values are cited in chronological order. Some of the values for 25° were obtained by

interpolation or extrapolation in a plot of log K versus lOOO/r. Log ^^1,2 values originally given

in molal or mole^ fraction units were converted to molar units by applying eq 26 or 27 (see sec.

4.1.1). P. dielectric polarization; CAT, catalytic method; I, isopiestic; E, ebullioscopic; IR,

infrared. R. A. Robinson, provisional value. 'See Davis and Paabo, ref. [15]. References:

[250] G. Allen and E. F. Caldin, (a) Trans. Faraday Soc. 49, 895 (1953); (b) Quart. Rev. (London)

7, 255 (1953). [251a] F. T. Wall and F. W. Banes, J. Am. Chem. Soc. 67, 898 (1945). [252] H. Dunken and G. Jiiger, Z. Chem. 3, 432 (1963). [253] A. Weissberger, J. Am. Chem. Soc. 65, 102 (1943).

[2.55] R. J. W. LeFevre and H. Vine, J. Chem. Soc, p. 1795 (1938). [258] K. Palm and H. Dunken, Z. Physik. Chem. (Leipzig) 217, 248 (1961). [263] W. Maier, J. Chim. Phys. 61, 239 (1964). [271] R. A. Robinson and R. H. Stokes, "Electrolyte Solutions," 2d ed. (Academic Press, Inc., New York, N.Y., 1959). 33
to the same concentration units (liter/mole) and

temperature (25°). The agreement among multiple

values obtained by diflferent methods is better than

in the examples shown in table 23; for example,

see the various log Ki,2 values obtained for o-chlo- robenzoic, m-iodobenzoic, and benzoic acids. The

trend towards increased self-association in the

aprotic solvent with decreasing ionization (increasing pKa value) in water which was noted in the pre- ceding table is more strongly evident.

The relative tendencies for the acids to dimerize

in benzene may also be compared with their acidic

strengths in benzene; these are indicated by the

log K" values (column 3), which measure the rela- tive affinities of the acids for the base diphenylguanidine (see eqs 15-16, section 2.3). The log K"

values were shown earlier [32] to parallel the ten- dencies for ionization of the acids in water — that

is, the greater the ionization of the acid in water,

the greater its capacity for association with di- phenylguanidine to form hydrogen-bonded ion

pairs. The data in table 24 reveal that acids with

the greatest tendencies toward self-association

combine the least readily with diphenylguanidine.

The correlation of strong self-associative ten- dencies in aprotic solvents with reduced manifestations of acidic behavior in water was previously

evident to others, even on the basis of much more

limited evidence [255, 253, 257b, 248, 250]. Barton

and Kraus [248] observed that log Ki,2 values for benzoic, o-bromobenzoic, cinnamic, and phenylpropiohc acids in benzene (at about 5.5°) are ap- proximately hnearly related to the pKa values for water solutions (25°), but 2-phenylpropionic (hydrocinnamic) acid exhibited a puzzling failure to con- form to this relationship. Subsequent work of Allen

and Caldin [250] supplied an explanation for the

deviating behavior of 2-phenylpropionic acid. On

examining a larger body of data they discovered

that carboxylic acids can be subdivided roughly

into two groups: (1) Acids in which the — COOH

group is conjugated with the side chain (benzoic,

substituted benzoic, cinnamic, crotonic, and phenylpropiohc acids); (2) acids in which —COOH is not conjugated with the side chain (aliphatic acids,

including phenylacetic and 2-phenylpropionic acids).

The acids of group (1) seem to be more highly

dimerized than those of group (2). In comparing the isomeric toluic and monohalobenzoic acids (table 24) it wiU be observed

that the extent of dimerization increases in the order

o- < m-< p-. The dimerization constant of p-toluic

acid seems unexpectedly large. No explanation has

been offered for this. Interestingly, a dimerization constant for 2,4-dimethylbenzoic acid in benzene

which was determined indirectly by Davis and

Paabo [254a] proved to be almost identical with

the value reported for o-toluic acid [251a]. The

indirect method used (which was applied earlier

to benzoic acid [254], giving log A^i,2 = 2.72) consists of measuring the absorbance at 540 nanometers

(m/u.) of various mixtures of diphenylguanidine,

bromophthalein magenta E (3',3",5',5"-tetrabromophenolphthalein ethyl ester), and a carboxyhc

acid in benzene, then calculating both the equilibrium constant for 1:1 association of the car- boxylic acid with diphenylguanidine and the

dimerization constant of the carboxylic acid by

fitting various assumed values of Ki,2 to the data. Results of I'Haya and Shibuya [267a] in a very

recent investigation of the dimerization of benzoic

acid in carbon tetrachloride and in chloroform support their claim that measurements of the inte- grated intensities of absorption bands give better

values than peak intensity measurements. The

data used were for the carbonyl stretching-vibration bands of the monomeric and dimeric forms of

benzoic acid. Averaging the results for 14 solutions

in carbon tetrachloride having stoichiometric concentrations in the range 0.0003-0.073 M they

obtained log A:i^(~25°) = 3.37, the R.S.D. (relative

standard deviation) of the values of Ki,2 being 8.3 percent. For chloroform solutions (15 mixtures,

0.00013-0.076 M) they obtained log /^i, 2 = 2.60,

R.S.D. for Ki,2 being 7.3 percent. The dimerization

constants, unlike results from earlier investigations (see ref.. 267a, literature cited), did not exhibit a concentration dependence.

From dielectric polarization data for cyclohexane

solutions Thyrion and Decroocq [267b] have obtained the following values of log Ki 2 (25°): CH3COOH, 4.32; CHCI2COOH, 3.48; CCI3COOH

(10 solutions, 0.0042-0.051 M, R.S.D. for 2 7.6%),

3.18; CF3COOH (10 solutions, 0.0045-0.06 M,

R.S.D. of Ki,2 values 11.2%), 2.51. Values of pKa

for the first three acids are given in table 23, and

work of Henne and Fox [267c] indicates that pK

for trifluoroacetic acid is ~ 0.19. Thus, it is evident

that the same kind of parallelism noted earher

holds for this series of acids — that is, the stronger

the acid in water the less the tendency toward dimerization. It will be noted that the dimerization constants for the first three acids in cyclohexane

are substantially greater than corresponding values

for carbon tetrachloride solutions (table 23) — much

closer agreement would have seemed reasonable.

c. Hetero-association. As exemplified in recent work by Affsprung and associates [267], mixtures

of carboxylic acids exhibit the expected tendency

to form "hetero"-dimers as well as "homo"-dimers.

From IR data these authors determined Ki,2 for \ both acetic and trichloroacetic acids, and also

Xassocn. fof th^ /le^ero-associatiou, using solutions

in which the two acids were present in equimolal

concentrations. When reduced to molar units (CClj,

25°) the log K values found for the three associative

processes, in the order mentioned, are 3.30, 2.43,

and 3.65, indicating that the hetero-dimer is formed

in preference to the homo-dimers. Although this result should not be attributed solely to statistical i factors, it is worth noting that in systems containing two monomeric acids, A and A', in equal concentration, random collisions of A with A' are more { probable than collisions of A with A or of A' with A'. «, 34
d. Solvent effects on dimerization. In the light of

present-day knowledge one would expect carboxylic acids to associate with solvents that are capable

of forming hydrogen-bonded complexes, at the

expense of self-association. Employing either their electron donor capacity or proton donor capacity,

or both, carboxylic acids are capable of forming

complexes of the types illustrated below (XVIII-XX).

.0...HCC1, ,0 0...H-O.. R-Cr ^ R-C( R-C( H 0-H 0-H...OEt_ 0-H...O^

2 I Et

(XVIII) (XIX) (XX)

TT-Electron donors may be included among the

solvents expected to form complexes of type (XIX).

The effects of different aprotic solvents on the extent of dimerization seem to be completely consistent with expectations based on our previous

comparisons of these solvents (section 3.4). Thus,

trichloroacetic acid (compare tables 23 and 24) ap- pears to be less dimerized irMrenzene, a 7r-electron donor, than in carbon tetrachloride. Analogously,

other carboxylic acids have'been found to be less dimerized in benzene than iji heptane [256] or in carbon tetrachloride [274]. Brooks and co-workers

[244], who studied the dimerization of benzoic acid

in 7i-hexane, carbon tetrachloride, carbon disulfide,

and chloroform (IR spectra, equimolar solutions),

found that these solvents produced an increased

monomer-dimer ratio in the sequence named.

It is appropriate to say a few words about the

effects of polar solvents, in view of their frequent

admixture with aprotic solvents. Brooks and as- sociates [244] found that benzoic acid is not measurably dimerized in acetonitrile, even at fairly

high concentrations (1 M). Lindberg and Majani

[266], who were interested in the association of

acetic acid with methyl sulfoxide, concluded (IR

spectra) that both 1:1 and 2:1 association occur. The structures postulated are shown below (XXI,

XXII):

0 O...H-^.C^ ^ CH -C'^ CH -C^ 0

^ ^0-H...0=S(CHj)2 ^ ^0-H...0=C(CHj)2

j (XXI) (XXII)

! In acetic acid-methyl sulfoxide mixtures diluted

' with carbon tetrachloride, log K for the formation

of (XXI) appears to be about 2.58 (liter/mole, 25°).

j The importance of such studies, as contributions

ij to^ increased understanding of solvent effects on

j acid-base behavior, as well as for correct interpre-

' tations of absorption spectra, is easily appreciated.

In two very recent investigations of interactions of

1 carboxylic acids with varied solvents, Lascombe

i and associates [275] and CoUings and Morgan [276]

have examined solvent-produced shifts of the car- bonyl stretching frequency of the acids in the region

j 1700-1800 cm~'. The comprehensive investigations by Lascombe et al. dealt with acetic acid

(CH3COOH and CH3COOD) and also propionic,

butyric, caproic, and benzoic acids in some 20 to 30 different solvents of the three types repre- sented in (XVIII-XX). The spectra of acetic

acid-butanol mixtures were interpreted as indicating the presence of three species — a 1:1 aggregate,

RCOOH...O(H)R (compare XIX), and two species of

the composition RCOOH-2BuOH. One of these is assigned a structure analogous to (XX), and the other

is presumed to have the same structure except that the two alcohol molecules are not connected by a hydrogen bond.

Carboxylic acid-solvent interactions have also been studied recently by means of ultraviolet spec- troscopy [268, 269, 277, 278]. In an investigation

of this type, Forbes and Knight [278] concluded that benzoic, p-toluic, salicylic, and p-hydroxybenzoic

acid are more prone to dimerize than to associate

with ethyl ether in cyclohexane solutions. However, the acids are thought to be monomeric in 100 percent ether. In discussing the determination of dimerization constants by the distribution method (sec. 4.2.2b)

it was emphasized that carboxylic acids seem prone

to associate with water molecules in aprotic solvents

at the partial expense of dimerization. However,

the dimeric structure is believed to persist to some

extent in aqueous solutions, and to account for anomalous deviations in the behavior of carboxylic

acids at low concentrations where the Debye-Hiickel

theory should be applicable [279-281]. MacDougall

and Blumer [279] calculated that Ki,2 for acetic acid

in water is 0.185 (liter/mole, 25°), and Katchalsky

and associates [280] deduced a closely similar value

(0.16); they also computed several other dimerization constants, ranging from 0.04 for formic acid

to 0.75 for benzoic acid. Davies and Griffiths [281],

while agreeing that acetic acid is appreciably dimerized in water, arrived at a somewhat different association constant. In this connection it is noteworthy

that the acids found to be the most highly dimerized

in benzene or carbon tetrachloride (tables 23 and

24) are the ones least ionized in water. It is also of interest that benzoic acid seems to be more highly

dimerized than acetic acid both in water and in aprotic solvents.

e. Chelated acids. As noted above (sec. 4.2.2a),

orf/io-hydroxybenzoic (salicylic) acid and related

acids seem to be strengthened by chelate ring for- mation. In contrast, intramolecular hydrogen bonding in o-methoxybenzoic acid (XV) will obviously

reduce the extent of its involvement as a proton

donor in intermolecular reactions. Both infrared

[282, 283, 244, 252] and ultraviolet [269] spectral

data indicate the presence of chelated o-methoxybenzoic acid in solvents like carbon tetrachloride

and cyclohexane. It would also be inferred from the

comparatively low value (0.59) for log Ki,2 found

for o-methoxybenzoic acid in benzene (see table 24)

that the acid is largely in the monomeric, intra- molecularly hydrogen-bonded form. The abnormally

high dipole moments of related naphthoic acids

(for example, l'methoxy-2-naphthoic and 3-methoxy35
2-naphthoic acids) have likewise [284] been ex- plained in terms of chelation. In accord with all these findings, Davis and Hetzer [32] observed that

in benzene solution o-methoxybenzoic acid associ- ates much less readily with diphenylguanidine

than do most other benzoic acids. Chelation of

o-methoxybenzoic acid is not evident in solvents

which can become hydrogen bonded to the -OMe

group, like water, alcohols, and dioxane-water [32].

It is also reduced in extent by the addition of ethyl

ether to carbon tetrachloride solutions [244]. Intramolecular hydrogen bonding has also been

postulated for the orf/io-halobenzoic acids, which seem anomalously weak both in ethanol [285] and

in benzene [32]. However, infrared studies lately

reported [244, 252] were interpreted as signifying

absence of intramolecular hydrogen bonding. The

monomeric acids exhibit two absorption peaks

(as do also the corresponding esters) in carbon tetrachloride and chloroform solutions; these have

been attributed to the presence of the two conformational isomers (compare Xllla and Xlllb) be- cause all four o-halobenzoic acids show only single,

closely similar monomeric hydroxyl bands. We

have noted above, however, that carbon tetrachloride and chloroform are not completely inert sol- vents (sec. 3.4.3b and 3.4.5a), and it does not seem

certain that the conclusions can be generaUzed to all aprotic solvents. The strong evidence of intramolecular hydrogen

bonding found for o-nitrophenol does not appear

in the spectral [252] or chemical [32] behavior of

o-nitrobenzoic acid, in which chelation is not geo- metrically favored,

/. Thermodynamic constants. Comparative ther- modynamic data for the monomer-dimer equilibrium

of carboxylic acids are available in a limited number

of cases [250-252, 263]. Table 25 contains a comTable 25. Thermodynamic constants for the monomer-dimer equilibrium of some carboxylic acids {literjmole, benzene, 25°)^

Acid Method -AG° -AH -A5° A///0.298AS Reference

kcal/mole kcal/mole caljdeg mole

Benzoic I 3.81 8.42 15.5 1.8 [251al Benzoic E 3.76 7.70 13.2 1.95 [250a] Benzoic P 3.84 9.29 18.3 1.7 [252] o-Bromobenzoic P 3.29 8.14 16.3 1.7 [252] o-Chlorobenzoic E 3.31 8.31 16.8 1.7 [250a]

o-Chlorobenzoic P 3.29 8.12 16.2 1.7 [252] P 3.66 9.65 20.1 1.6 [252] o-Iodobenzoic P 3.33 8.77 18.2 1.6 [252] o-Methoxybenzoic P 0.80 2.06 4.25 1.6 [252] o-Nitrobenzoic P 2.62 7.50 16.4 1.5 [252]

o-Toluic I 3.58 8.39 16.1 1.7 [251al m-Bromobenzoic P 3.58 9.06 18.4 1.65 [252] m-Chlorobenzoic P 3.57 8.92 17.9 1.7 [252] m-Fluorobenzoic IR 3.62 8.47 16.3 1.75 [252] m-Iodobenzoic E 3.57 7.65 13.7 1.9 [250a]

m-Iodobenzoic P 3.65 9.32 19.0 1.6 [252] m-Toluic I 3.91 9.26 17.9 1.7 [251a] p-Bromobenzoic P 3.84 9.35 18.5 1.7 [252] p-Chlorobenzoic P 3.86 9.36 18.5 1.7 [252] p-Fluorobenzoic IR 3.79 9.33 18.6 1.7 [252]

p-Iodobenzoic P 3.74 9.23 18.4 1.7 [252] p-Methoxybenzoic E 3.74 6.57 9.5 2.3 [250a] p-Toluic E 4.61 8.72 13.8 2.1 [250a] Phenylpropiolic E 4.03 8.75 15.2 1.9 [250a]

^ Where necessary, values for 25° were obtained by interpolation or extrapolation in a plot of log A 1, 2 versus lOOO/T. Part of the

log Ki, 2 values, originally given in mole fraction units, were converted to molar units by applying eq 27 (see Sec. 4.1.1). Thermodynamic

constants were calculated, when necessary, by conve.ntional methods.

'' I, isopiestic; P, dielectric polarization; E, ebuUioscopic, IR, infrared.

References:

[250a] G. Allen and E. F. Caldin, Trans. Faraday Soc. 49, 895 (1953). [251a] F. T. Wall and F. W. Banes, J. Am. Chem. Soc. 67, 898 (1945). [252] H. Dunken and G. Jager, Z. Chem. 3, 432 (1963). 36
pilation of values for phenylpropiolic acid and a

series of benzoic acid derivatives; several methods

of measurement are represented [250a, 251a, 252].

Table 26 contains thermodynamic constants for the

Table 26. Thermodynamic constants for the monomer-dimer

equilibrium of benzoic acid in different solvents (Uteri mole, 25T

Solvent -AC -AH -AS AH/0.298

AS

caljdeg kcal/mole kcal/mole mple

Cyclohexane 6.07 13.7 25.6 1.8 Carbon tetra- 5.69 13.6 26.4 1.7 chloride Chlorobenzene 4.45 10.1 19.0 1.8 Toluene 4.00 9.6 19.0 1.7 Benzene 3.82 8.6 16.1 1.8 Chloroform 3.60 9.3 19.3 1.6

^ Converted from to Kc units by means of eq 27 (see sec.

4.1.1). Thermodynamic constants were calculated by conventional procedures. From results of Maier and associates [263]. Reference:

[263] W. Maier, J. Chim. Phys. 61, 239 (1964).

monomer-dimer equilibrium of benzoic acid in several solvents, derived from ultrasonic absorption

data [263]. The values for benzoic acid in benzene,

together with the results cited in table 25, provide

four independent measurements by different methods. The constants have all been converted (when

needed) to the same (molar) units of concentration

(see eq 27). For reasons given earlier (sec. 4.2.2b),

data obtained by the distribution method have not been included.

The following points may be noted in table 25:

(1) All of the thermodynamic constants are negative in sign; (2) values of — AH are mostly in the

range 8-9.4 kcal/mole; (3) — AS values are mostly

in the range 13-20 cal/deg mole; (4) the ratio AHimS, where T=298 °K, is 1.7 for about half the cases listed, the range of values being 1.5-2.3. o-Methoxybenzoic acid conforms well to the last generalization. There is likewise an approximately

linear relation between AG and AH. It should be

remembered, of course, that these acids have closely

similar structures. From comparing the four sets

of values for benzoic acid (see also m-iodobenzoic

acid) it is apparent that the reproducibility attained

in such measurements is not yet great enough for significance to be attached to deviating cases. In an earlier paper [33] Davis and Hetzer com- mented on the approximate constancy of the ratio A///0.298 AS for a smaller group of benzoic acids.

Still earlier, signs of a linear correlation of A// and

AS (based on a very limited amount of data) and the

theoretical basis of such a correlation were dis- cussed by Allen and Caldin [250b].

The approximate constancy of the ratio AHI 0.298 AS upon changing the solvent is indicated

by the data of table 26. The numerical values of

AH and AS are noticeably greater for benzoic acid

when in the two solvents cyclohexane and carbon

tetrachloride than in the aromatic solvents or chloroform — they are closer to the values pertaining

to dimerization in the vapor phase [250b, 263].

These results are consistent with the evidence,

previously mentioned, of tendencies for association of carboxylic acids with aromatic solvents and

chloroform.

g. Possible equilibrium between open-chain and

cyclic dimers. Although the possible presence of a dimer containing only one hydrogen bond is rec- ognized, as noted in the introductory discussion of carboxylic acids, dimeric carboxyUc acids are usually assumed to be predominantly in the cyclic

form when dissolved in aprotic solvents, and calculations of dimerization constants have commonly

been made under this assumption. Not long ago

[286] Bellamy and associates raised the question

again, calling attention to the previous suggestion

by Feneant [287] that the breadth of the Vs band

of carboxylic acids may result in part from the pres- ence of a mixture of cyclic and open-chain dimers.

In Bellamy's view, this subject merits further

study.

h. Thiolobenzoic acid. Thiolobenzoic acid,

PhC(=0)SH, studied by infrared (CCU), ultraviolet (C6H12), and NMR spectroscopy (CeHe), as well

as cryoscopicaUy (ChHh), has been found to exhibit

inappreciable tendency toward dimerization [288]. According to evidence summarized by Pimentel

and McClellan [102], the — SH group is a relatively

poor proton donor, although instances of hydrogen

bonding have been observed with strong electron

donors.

4.2.3. Nitric Acid

Self-association of nitric acid in aprotic solvents

has been inferred from data for its distribution

between concentrated aqueous solutions and aprotic

solvents such as benzene, toluene, kerosene, and

chloroform. This work originated in the need for interpreting data obtained in extracting strongly

acidic materials from aqueous solutions into aprotic

solvents by the aid of basic substances like dialkyl

hydrogen phosphates and tri-n-octylamine [289-291]. In the work of Hardy and associates [289] infra- red absorption curves for the extracts in aprotic

solvents indicated strongly the presence of a dimeric

form as the predominant species, possibly in the

form, of a monohydrate. The structure favored by

the authors is (XXIII). The solvents used in the

0...H 0 HO-N ^O...HO-N(' ^o...h/ 0

(XXIII) 37
infrared studies were benzene, carbon tetrachloride,

cyclohexane, n-dodecane, kerosene, and toluene.

Larger aggregates are considered likely to form in very highly concentrated solutions.

Verstegen [290] determined dimerization constants for nitric acid in benzene, o-dichlorobenzene,

and chloroform. His own results and those of other

workers are shown in table 27. It should be borne in

from orthophosphoric acid by partial esterificatioh, as indicated in (XXIV) and (XXV).

\oH),

(XXIV)

^(0R)2

P=0 \

OH

(XXV)

Table 27. Dimerization constants for nitric acid in wet aprotic

solvents^

Solvent Log K,, 2 Reference

(Literjmole) Kerosene 3.60 [289] Benzene 3.09 [290] 2.77 [289] Xylene 2.48 [291] o-Dichlorobenzene 2.00 [290] Chloroform 1.70 [290]

" From Verstegen's compilation of his own and published

results, all obtained by distribution measurements [290].

References:

[289] C. J. Hardy, B. F. Greenfield, and D. Scargill, J. Chem.

Soc, p. 90 (1961). [290] J. M. P. J. Verstegen, J. Inorg. Nucl. Chem. 26, 1085 (1964). [291] U. Bertocci and G. Rolandi, J. Inorg. Nucl. Chem. 23,

323 (1961). mind that all of these values are for wet solvents. Why o-dichlorobenzene occupies its relative position in table 27 is not clear, but the relative order

of the other solvents is in harmony with our preceding discussion of aprotic solvents. Attention should be called to a communication

by Hogfeldt [292], who has investigated the distri- bution of nitric acid between water and several

aromatic hydrocarbon solvents [293]. Hogfeldt

disagrees with deductions of Hardy and associates

[289] and of Verstegen [290] in regard to dimer for- mation in the organic layer. In his opinion the

method used for determining water in the organic

phase is not accurate enough to warrant such defi- nite conclusions. He has suggested an alternative

formulation of the distribution reaction [292, 293].

4.2.4. Phosphorus-Containing Organic Acids

The practical importance which phosphoruscontaining organic acids attained within the past

20 years in solvent-extraction processes created

a need for information about their states of aggregation in various organic solvents. These processes

involve distribution of the acids between aqueous

and organic layers.

The phosphorus-containing acids chiefly used

up to now in the extraction processes are derived

Commonly used examples of (XXIV) are n-butyl

(dihydrogen) phosphate and phenyl (dihydrogen)

phosphate (sometimes referred to, respectively, as "butylphosphoric acid" and "phenylphosphoric

acid"). Examples of (XXV) are di-n-butyl (hydrogen)

phosphate and diphenyl (hydrogen) phosphate

(occasionally called "dibutylphosphoric" and "di- phenylphosphoric" acids).

Acids with alkyl or aryl attached directly to phosphorus are less well known, and their nomenclature has not been firmly established. Examples

are: Benzenephosphonic (phenylphosphonic) acid

(XXVI), butyl benzenephosphonate (XXVII), and

benzenephosphonous (phenylphosphinic) acid

(XXVIII). Phosphonous acids may have two R

groups attached to P, as in (XXIX).

Ph-P=0

(XXVI)

. ^OBu

Ph-P=0

(XXVII)

H

/ Ph-P=0 ^OH

(XXVIII)

Ph

/ Ph-P=0

(XXIX)

Infrared surveys have shown that acids of these

six types, all of which contain the P(=0)OH or P(=0)(0H)2 grouping, have a very strong tendency

when in solvents like carbon tetrachloride and car- bon disulfide to form aggregates containing

0-H...0 bonds [294, 295]. This conclusion has

been reinforced by studies employing additional

measuring techniques (notably distribution and

cryoscopy) [296-311].

a. Nature of aggregates formed. There seems to be agreement on the following points:

(1) Acids with the grouping P(=0)OH usually

form dimers in aprotic solvents, resembling car- boxylic acids. The dimers are thought to have a

cyclic structure analogous to that of carboxylic

acid dimers (XVI). Acids of types (XXV), (XXVII),

and (XXIX) are included in this category. (XXVIII)

is an exception; according to Kosolapoff and

Powell [309] it forms trimeric aggregates (cryoscopy

in benzene and naphthalene); these are thought

likely to be cyclic.

(2) Acids containing the grouping P(=0)(0H)2

form linear aggregates of magnitude varying with

concentration and temperature. Types (XXIV) and

(XXVI) belong to this category.

(3) The tendency of phosphorus-containing or- ganic acids to form aggregates is very great — it exceeds that of carboxylic acids.

b. Solvent effects. Like carboxyhc acids, phosphorus-containing acids can form aggregates with

38
solvents having proton donor capacity (for example,

CHCI3), electron donor capacity (ethers, ketones,

etc.), or both (water, alcohols); compare section

4.2.2b, formulas (XVIII-XX). Any such aggregate

formation naturally reduces the extent of self- association.

While studying the distribution of di-n-butyl

phosphate between water and organic solvents,

Dyrssen [296] observed that it has an appreciable

tendency to associate with tributyl phosphate

(which is well-known for its electron donor capacity—for example, see Davis and Schuhmann [14]). Peppard, Ferraro, and Mason have made several

studies of solvent effects on the self-association of

acids of types (XXIV), (XXV), and (XXVII), by

cryoscopy, infrared spectroscopy, isopiestic measurements, and, most recently, by NMR spectroscopy [299-301]. From isopiestic data [300] they

reached the following conclusions: "In general,

the monobasic acids are dimers in solvents such as

re-hexane, cyclohexane, benzene and carbon tetra- chloride; monomers in methyl alcohol; and inter- mediate in chloroform and acetone." In studying

solutions of 2-ethylhexyl (dihydrogen) phosphate

they deduced the presence of large aggregates (of

molecular weight 8-14.5 times the formula weight)

in n-hexane and cyclohexane, hexamers in benzene

and carbon tetrachloride, dimers in acetone, and monomers in methanol [300]. Acetic acid is another

solvent in which the acids appear to be predominantly monomeric [299]. In this connection, it seems pertinent to note that Davis and Hetzer [313]

found that diphenyl phosphate, which is strongly

acidic in water, acts as a weak acid in glacial

acetic acid; heteroassociation of the solute and sol- vent was suggested as an explanation (see footnote

15 of ref. 313).

Further examples of solvent effects will be noted

in the next section.

c. Monomer-dimer equilibrium constants. A very

limited number of log Ki,2 values, obtained from

extraction experiments, are available for dialkyl

and diaryl phosphates; quantitative evidence does

not seem to have been obtained for the other classes

i of organophosphorus acids.

j Table 28 presents log Ki,2 values for di-n-butyl

phosphate in 10 wet organic solvents and in water;

these were obtained by Dyrssen and coworkers

[296, 297] or by Hardy and associates [303, 304].

Although the conditions which had to be employed

were far from ideal, expected trends in solvent

effects are discernible. For example, dimerization

appears to be less extensive in vr-electron donor

i solvents (toluene, benzene) than in aliphatic hydroj carbons or carbon tetrachloride, and still less ext tensive in n-electron donors (isopropyl ether, methyl

isobutyl ketone). As expected, the dimer/monomer

I ratio is lowest in water and methyl isobutyl carbinol.

1 Ferraro and Peppard's recently reported study

I [302] of the dimerization of some dialkyl and diaryl

J phosphates in carbon tetrachloride (dry and wet,

) 25°), using NMR spectroscopy, shows clearly the

Table 28. Solvent effects on the dimerization of di-n-butyl

phosphate in various (wet) organic solvents (2S°f

Organic solvent Mineral acid

in aqueous layer Log /Ci 2 Refer- ence

Uteri mole n-Hexane 0.1 M HNO:, 6.87 I297J

Kerosene 1 M HNOr) 5.78 [303]

0.1 M HNO:, 6.49 [297J

1 M HNO3 5.33 [303]

Toluene 1 M HNO3 5.09 [304] Benzene 1 M HNO3 4.88 [303]

Chloroform 0.1 M (Na, H)C104 4.48 [296]

0.1 M HNO.3 4.61 [297] 0.1 M HCIO4 4.21 [303]

1 M HCIO4 4.36 [303] Nitrobenzene 1 M HNO3 3.55 [303]

Isopropyl ether 0.1 M HNO3 2.29 [297]

1 IQ

Water " None 1.11 [296]

Methyl isobutyl carbinol... 0.1 MHNOa <-l [297]

" Based on data from distribution experiments. Saturated solution of di-re-butyl phosphate.

References:

[296] D. Dyrssen, Acta Chem. Scand. 11, 1771 (1957). [297] D. Dyrssen and L. D. Hay. Acta Chem. Scand. 14, 1091

(1960). [303] C. J. Hardy and D. Scargill, J. Inorg. Nucl. Chem. 11,

128 (1959). [304] B. F. Greenfield and C. J. Hardy, J. Inorg. Nucl. Chem.

21, 359 (1961).

interference with the dimerization process which

water produces (see their fig. 3). They calculated

for log K\, z (liter/mole) the approximate value 3.18,

which is much lower than the values obtained by

the distribution method (see table 28), but close to the value reported for chloroacetic acid (see table

23). If water acts merely as a competitor in hydrogen

bond formation, a higher value would have been

expected for dry CClj solutions. Further consideration of the role of water is obviously needed.

(Note the following discussion of sulfinic and sul- fonic acids.) The value of —AH (kcal/mole) per

hydrogen bond was estimated as being 5.5-5.8 —

this is smaller than Maier's value (6.8) for benzoic

acid in CCI4 [263] (see table 26). Krasovec and Jan [307] have compared the log

Ki,2 (wet chloroform) and pKa (water) values of

several diaryl phosphates, using the distribution

method; see table 29. As in the case of the carboxyhc acids listed in tables 23 and 24, the acids having

the lowest acidic strength in water appear to possess

39
Table 29. Dimerization constants (wet chloroform) and pKg

values (water) ofsome diaryl and dialkyl phosphates (R2HPO4) "

Identity of R Log Ki, 2 Reference

liter/mole p-Chlorophenyl 3.44 0.20 [307] Phenyl 3.76 0.26 [307] Benzyl 3.90 0.70 [307] 2-Naphthyl 4.02 0.74 [307] 4.24 0.40 [307] Ethyl (25°) 4.46 0.73 [298]

ra-Butyl (25°) 4.48 1.00 [296]

" Both sets of values were calculated from distribution data;

f = 20 °C except as indicated otherwise.

References:

[296] D. Dyrssen. Acta Chem. Scand. 11, 1771 (1957). [298] D. Dyrssen, S. Ekberg, and D. H. Liem, Acta Chem.

Scand. 18, 135 (1964). [307] F. Krasovec and J. Jan, Croat. Chem. Acta 35, 183 (1963).

the greatest capacity for self-association. Di-p-tolyl

phosphate (Hlie p-toluic acid) seems to behave

anomalously.

Ashby and Kosolapoff [310] determined the ap- parent molecular weights of the three isomeric

tolylphosphonic acids — see (XXVI) — cryoscopically and ebuUioscopically, using (when solubilities permitted) the solvents acetone, camphor, naphthalene, and water. The extent of aggregation of the

isomers appears to increase in the order ortho

< meta < para.

Additional findings are discussed in the refer- ences cited.

4.2.5. Sulfinic and Sulfonic Acids

a. Sulfinic acids, R-S(=0)OH. Self-association

of sulfinic acids has been studied in benzene, carbon

tetrachloride, and nitrobenzene.

Wright [314] determined the extent of association

of several carboxylic acids and of several sulfinic acids (benzene-, p-toluene, 1- and 2-naphthalenesulfinic acids) cryoscopically in benzene and nitro- benzene, and found the sulfinic acids to be more

highly associated than the carboxylic acids — ap- parently they form trimers and hexamers. The

association of both carboxylic and sulfinic acids was found to be greater when water was present.

Detoni and Hadzi [315] studied by IR and UV

spectroscopy the association of benzene-, p-toluene-,

2-naphthalene-, and o-methoxybenzenesulfinic acids

in carbon tetrachloride, both dry and moist. Of

various suggested structures for these acids they

favor (XXX). They concluded that benzenesulfinic

R-S

(XXX)

acid forms aggregates greater than the dimer and,

in agreement with Wright, found that this acid is much more highly associated in wet than in dry

solvent. For carboxylic acids the opposite seems to be true (see table 23).

b. Sulfonic acids, R-SO2OH. Kehiaian and Nenitzescu's ebulliometric measurements [316, 317] of

anhydrous benzene solutions of benzenesulfonic

acid led to the conclusion that this acid does not

dimerize, but instead forms higher aggregates. The

association constants {K„^ „ + i) for the different steps

are held to be the same, and values found at the

temperatures 52.6, 65.2, and 80.1 (°C) are, respectively, 10.9, 8.75, and 6.95 (molal units); the extrapolated value for 25° is 18.4. This is equivalent (see

eq 26) to log /^„,,i-,-i (hter/mole)= 1.21. The values

found for A/7 (kcal/mole) and AS (cal/deg mole) are — 3.7 and —6.8, respectively.

4.2.6. tcrt.-Butyl Hydroperoxide and Deuteroperoxide.

Deuterium Isotope Effects on Hydrogen Bonding by

Br^nsted Acids

Recently published investigations of Walling and

Heaton [318] have provided information about the

self-association and other hydrogen-bonding be- havior of f-butyl hydroperoxide and f-butyl deuteroperoxide in carbon tetrachloride. Their results,

which are summarized in table 30, are based on IR

Table 30. Log K, AH, and AS for some complex-forming reac- tions o/«-BuOOH and /-BuOOD (CCh)''

Complex formed Log Kassocn. -^H -AS

30° 70°

caljdeg kcal/mole mole f-BuOOH + PhCH = CH2 -0.57 -0.77 2.39 10.5 nBuOOH + CeHe -0.82 -0.96 1.58 9.0 r-BuOOH + o-CsHiClz -0.82 -0.89 0.74 6.2 f-BuOOH + PhCl -1.00

lt-BuOOH)2 + 0.28 -0.22 5.95 18.4 t-BuOOD + PhCH = CH2 -0.08 -0.30 2.7 9 f-BuOOD + CeHe -0.60

(i-BuOOD)2 +0.22 -0.19 4.9 15

" Results obtained by Walling and Heaton (IR spectroscopy)

[318],

Reference:

[318] C. Walling and L. Heaton, J. Am. Chem. Soc. 87, 48 (1965). and distribution data. These results indicate that

i-butyl /iyc?roperoxide dimerizes more readily than

f-butyl c?ew^eroperoxide; a cyclic structure (XXXI)

R\/-\

H H Vo(

R

(XXXI) 40
is postulated for these dimers. In contrast, the

deuteroperoxide seems to be more reactive than the

hydroperoxide in associating with benzene and

styrene. Walling and Heaton have called attention

to analogous cases. For example, in IR studies of Plourde and West

[319], it was found that the enthalpy of 5e//-association i^ greater for phenol-A than for phenol-c?,

whereas in other associations with bases the enthal- py of reaction is greater for phenol-c? than for phenol-/i; see table 31. (Also see table 9.) These

Table 31. Enthalpy changes for some complex-forming reactions of phenol-h and phenol-d ^

Electron donor — AW, kcal/mole Proton donor

PhOH PhOD

Benzene 1.84 2.18 Ethyl Acetate 4.81 5.07 Tetrahydrofuran 5.55 5.86 n-Butyl Ether 6.00 6.17 Phenol-ft 4.53 Phenol-c? 4.24

"Results of G. R. Plourde, from IR measurements [319]; solvent and concentration units are not specified.

Reference:

[319] G. R. Plourde, Dissertation Abstr. 22, 1400 (1962).

Studies showed, moreover, that the free energy of

association with bases is almost the same for phenol-A and phenol-o?, since a greater enthalpy

change is almost balanced by a greater entropy

change. The intramolecular deuterium bond was

found to be slightly stronger than the hydrogen

bond in o-chlorophenol and o-phenylphenol, and

much stronger in o-iodophenol. In explanation of

the results it was suggested that "relatively weak"

hydrogen bonds (not necessarily all hydrogen bonds)

are stronger in the case of deuterium because the

"lower zero point vibrational energy of OD . . . makes the OD bond shorter than the OH, allowing

closer approach of the base and greater inter- action. . . . The reversal in the isotope effects [in the self-association reaction] may be due to in- creased basicity of the oxygen in phenol-/i."

The relative strengths of hydrogen and deuterium

bonds have also been compared in trifluoroacetic

acid. Taylor and Templeman's investigations of CF3COOH and CF3COOD in the vapor phase [320]

indicated that the hydrogen bond formed in self- association is somewhat stronger than the deuterium

bond. Similarly, data of Maryott and associates

[257a] indicate that acetic acid-/i is more dimerized

than acetic acid-c? in benzene (30°). On the other

hand, CresweU and Allred [120] found in NMR

studies that fluoroform-c? bonds more readily than

fluoroform-/i to tetrahydrofuran.

4.3. Hydrogen-Bonded Ion Pairs In our preliminary discussion of acid-base reactions in aprotic solvents (sec. 2) it was emphasized

that such reactions should be formulated as association reactions, in a manner illustrated by eq 30,

B + HA ^ BRA, (30)

base acid salt

in which the organic base B is generally an aliphatic

amine or derivative of guanidine, or else one of the

weaker bases related to anihne or to pyridine. Salts of the class BHA are often said to consist of "ion

pairs," as expressed by the structure BH+A". However, this is an incomplete designation in a number

of cases where, as we shall presently see, there is substantial evidence of cation-to-anion hydrogen

bonding. It is desirable to recognize the occurrence

of such bonding by using the term ^'hydrogenbonded ion pairs'' symbolically represented by

BH+...A-.

Cation-to-anion hydrogen bonding in incompletely

substituted ammonium salts (R3NHA, R2NH2A,

RNH3A) is a very plausible explanation for certain

pronounced differences in their behavior and that of completely substituted ammonium salts (R4NA)

when dissolved in aprotic solvents. Differentiating

behavior has been discovered, for example, in measurements of such properties as conductance,

dielectric polarization, and spectral absorption.

Varied experimental evidence wiU be presented,

mainly pertaining to solutions in aprotic solvents,

although deductions from studies of hydrogen bonding in crystalline solids will also be included. It will also prove valuable to include examples of

experiments in which dipolar aprotic solvents

(e.g., nitrobenzene) were used. As we saw earher,

dipolar aprotic solvents can be assigned to Br0nsted's "class 4," that is, solvents of comparatively

high dielectric constant but weak proton-donor and

proton-acceptor capabilities. Like the aprotic sol- vents, dipolar aprotic solvents have scarcely any

masking effect on differentiating acid-base behavior.

Contrasts in the behavior of completely substi- tuted ammonium salts as distinguished from in- completely substituted ammonium salts were first discovered more than 50 years ago [322]. Additional

kinds of differentiating behavior have been dis- covered in the past 20 years in comparative studies

of acid-base reactions yielding tertiary ammonium

salts and those yielding secondary or primary am- monium salts. It will be shown that such additional

examples can be explained satisfactorily in terms

of hydrogen bonding reactions that are only possible

in salts having more than one proton attached to the ammonium nitrogen.

4.3.1. Conductance

Extended understanding of acid-base behavior in aprotic solvents has been one of the important byproducts of investigations of the conductance of

41
salts in nonaqueous solvents — in particular, the

classic investigations associated with such names

as Walden, Kraus, and Fuoss.

a. Walden s work. Walden's electrochemical

investigations of binary salts in numerous non- aqueous solvents extended over more than 20

years. In these studies, which were summarized in the Baker lectures for 1927-1928 [322], the anions

represented were mainly perchlorate, picrate, and

halides; and primary, secondary, tertiary, and

quaternary ammonium ions were among the principal cations represented.

Walden was very much struck by his discovery

that some of the binary salts known to be equally

strong in aqueous, methanolic, or ethanolic solu- tions (using conductivity as the measure of strength) can be subdivided into strong and weak electrolytes

in acetone. In further work he found that tetraalkylammonium halides and picrates are strong electrolytes in nonaqueous solvents like acetone or ethylene chloride, as well as in water, methanol,

and ethanol. In contrast, mono-, di-, and trialkylammonium halides and picrates, although strong

electrolytes in water and the alcohols, become weak

or only moderately strong electrolytes in acetone,

benzene, and halogenated hydrocarbons.

Walden concluded that the strength of a binary

salt depends both on the nature of the cation and

the nature of the anion, but the former exerts the

greater effect. In fact, he said: "The differences

between the tetra-alkylated ammonium salts (particularly the chlorides) on the one hand, and the

di- and tri-alkylated salts on the other, are so large

that it would indeed be permissible to speak of

them as entirely different electrolytes. . . ." The

following series was given as expressing the rela- tionships among these salts with respect to conductivity and percentage dissociation at equivalent

dilutions:

tetra- > tri-(mono-) > di-alkylated ammonium salt.

His corresponding series for anions (in salts containing the same cation) is:

Pi- > CIO4- > I- > NO3- > Br- > C1-.

In subsequent work Walden and Birr found that nitromethane is an additional differentiating solvent

[323]. In a different type of investigation Walden and his associates compared the mobilities of isomeric

substituted ammonium ions in different solvents

[322, 324]. For example, Walden and Birr measured

the mobilities of three isomeric cations — tetra- methylammonium, diethylammonium, and mono-«- butylammonium — in water and in acetonitrile [324]. All three ions were found to have approximately

the same mobility in acetonitrile, but not in water. Instead, the mobilities decrease in the following

order:

tetramethylammonium > diethylammonium

> butylammonium.

Walden and Birr inferred that the dissimilar mobilities of the isomeric cations in water result from

different degrees of solvation — the tetramethylammonium ion being probably unhydrated; diethylammonium ion, partly hydrated; and monobutylammonium ion, more completely hydrated. The

similar mobilities in acetonitrile were thought to indicate that the three cations are not solvated to any extent in this solvent.

b. Wynne-Jones' postulate. Wynne-Jones seems to have been the first person to suggest that the marked

reduction in the conductivities of incompletely

substituted ammonium salts in non-hydroxylic sol- vents, which his own experimental work showed to occur in nitrobenzene as well as in the solvents used

by Walden, may be explained on the assumption of

hydrogen bonding of anions to cations [325]. To

indicate such bonding in a trialkylammonium chlo- ride he wrote its structure as R3N:H:C1, applying

the concept of bicovalent hydrogen, used at that time to explain hydrogen bonding, and saying that the assumption of such bonding can account for the

low conductivities of incompletely alkylated am- monium salts in non-hydroxyUc solvents, as well as

for the necessity of adding a hydroxyUc solvent to assist ionization and for the effects of different

solvents on ionic mobilities that had been observed

by Walden and co-workers. He reasoned that when

a salt can exist in a non-ionic (hydrogen-bonded)

form its separation into cations and anions will be

assisted when the solvent can interact with the

partly substituted ammonium ion, thus separating

it from the anion, as indicated by eq 31:

R3N:H:C1 + S (a hydroxylic solvent molecule)

^R3N:H:S+ + C1-. (31)

Additional molecules of solvent can become hydrogen bonded to mono- and dialkylammonium ions. This suggestion, with the structures of hydrated

cations now written more specifically — for example,

R3NH+...OH2 — has become widely accepted,

as it explains other phenomena besides the promotion of ionic conductance by hydroxylic solvents

and differences in the mobilities of isomeric cations

in such solvents. In harmony with Walden's findings, Wynne-Jones

observed that the hydrochlorides of bases like piperidine, dipropylamine, and diphenylguanidine

are extremely weak electrolytes in nitrobenzene.

(In partial disagreement with Walden, he concluded

that the perchlorates of the bases studied are more

highly conducting than the picrates in nitrobenzene.)

Wynne-Jones also noted the differing behavior of incompletely substituted ammonium salts in five solvents of closely similar dielectric constants. In three of these solvents (acetone, acetonitrile, nitroj benzene) the salts are only weak electrolytes, even

42
though in the other two solvents (methanol, ethanol)

they are strong electrolytes.

c. Investigations of Kraus, Fuoss, and associates

[24, 326-336]. Striking differences in the conduct- ance behavior of completely and incompletely

substituted ammonium salts were also observed by

C. A. Kraus and R. M. Fuoss and their co-workers

in the course of extensive investigations of "properties of electrolytic solutions." Kraus concluded

[329] that electrolytes in benzene solution show

three types of behavior, which are illustrated by the

three conductance curves of figure 9. These curves

i-Atn4NSCN ;TYPE2lXr

i-Am4NPl r

yi-Am^NHPI

1,1.1. LOG C

Figure 9. Conductance plots for three different types of salts in benzene.

Reproduced with minor changes from C. A. Kraus, j. Chem. Educ. 35, 324 (1958), with the permission of the author and the publisher. (Ref. No. 329.)

were obtained by plotting the logarithm of the

equivalent conductance against the logarithm of

the (molar) concentration. At concentrations less than about 1 M each curve of figure 9 shows at least one minimum. The minima at the lowest concentrations studied are attributed to overlapping reactions

involving such species as "ion pairs," "simple

ions," and "triple ions." Thus, representing any

uni-univalent salt by the formula MX, the following

equilibria are supposed to coexist:

M+X-^M+ + X- (32)

M+X--f M+;=iM+X-M+ (33a)

X- + M+X-^X-M+X-. (33b)

Equation 32 indicates an equilibrium between ion

pairs and simple ions, the corresponding equilibrium

constant being referred to as the "ion-pair dissociation constant." Equations 33a and 33b indicate the

formation of positively or negatively charged triple

ions by association of ion pairs with simple ions. At high concentrations more complex aggregates

may be formed,

i Here, we are interested only in the equilibrium

I represented by eq 32, which is believed to account

well for conductance behavior at very low concentrations—for example, for the tetraalkylammonium

I salts in the concentration range extending from

I about 10~^ to 10"^ M, where the conductance rises

gradually from its minimum value. Experimental

results which are not shown in figure 9 indicate an

analogous gradual rise in the conductance of the

tertiary ammonium picrate at very low concentrations. Kraus and Fuoss have devised procedures

for calculating ion-pair dissociation constants,

examples of which are given in tables 32 to 35. In table 32 the ion-pair dissociation constants (expressed as log K values) of several quaternary

ammonium picrates in benzene and five other

organic solvents are compared with the constants

for picrates with at least one hydrogen atom at- tached to the ammonium nitrogen. The table illus- trates Kraus' two main reasons for considering, in agreement with Wynne-Jones, that cations and

anions of partly substituted ammonium salts interact

through non-Coulombic as well as through Coulombic forces:

(1) Ion-pair dissociation constants of completely

substituted ammonium picrates are more than

1000-fold greater than those of incompletely substi- tuted picrates, not only in benzene, chlorobenzene,

and ethylene chloride (all with relatively low dielectric constants), but also in nitrobenzene (e ~ 35). The close similarity of the constants for the picrates

having NH+, BuNH+, BU2NH+, or BU3NH+ as the

cation is especially noteworthy [24, 327-329, 335].

(2) The enormous disparity in the ion-pair dissociation constants of ammonium picrate and tetra- alkylammonium picrates almost disappears when

pyridine is the solvent, though its dielectric constant

is only about a third that of nitrobenzene. Such

behavior is to be expected if we accept the view

that the ammonium ion is hydrogen bonded to picrate in nitrobenzene solution, but to pyridine

when the latter is the solvent [332]. The differ- ences can be represented by the two structures H3NH+...Pi- and H,,NH+...Py.

The last two salts listed in table 32 under the

heading "Other Ammonium Salts" are hydroxxammonium picrates. That is, they are derived from

tertiary amine oxides instead of from amines, and

oxygen, rather than hydrogen, is directly attached

to the ammonium nitrogen. The two examples

given -MeaNOHPi and PhN(Me2)0HPi-seem to be

about equally strong as electrolytes, having ion- pair dissociation constants about a tenth as great as those of ammonium picrate and partly alkylated

ammonium picrates. As in the preceding cases, hydrogen bonding of cation to anion is an attractive

and plausible explanation for the very low tendency

toward ion-pair dissociation, as well as for the

marked disparity in the dissociation constants of

mef/io:rytnmethylammonium and /lyfl^rojcytrimethylammonium picrates.

Table 33 contains ion-pair dissociation constants which indicate a dependence of these values on the nature of the anion. Thus, in comparing values for pyridonium perchlorate and pyridonium picrate in ethylene chloride and in comparing values for butylammonium perchlorate and butylammonium picrate

in nitrobenzene we observe the perchlorate to be

280-613 OL-68—

43
Table 32. Log K for ion-pair dissociation of some quaternary and other ammonium picrates in several organic solvents (25°) ^

Solvent (e) Quaternary Ammonium Salts Other Ammonium Salts

Cation LogK Reference Cation Log K Reference

Benzene (2.27) j-Am4N + -17.05 [326] i-Am3NH + -20.60 [326]

Chiorobenzene (5.63) BU4N + -7.73 [330] BU3NH+ - 12.68 [330]

ttnylene cnlonae (lU. zo) R11.N + — ^ fid Rii NH+ — 7 fift

t .Do

Pyridine (12.3) Me4N + BU4N + -3.17

-2.91

[332]

[332] NHJ -3.55 [332]

Acetone (20.7) Me4N+

BU4N+

-1.95

-1.65

[333]

[334]

-2.95 [333]

Nitrobenzene (34.8)

BU4N + Me4N + PhN + Me3

MejN + OMe

'^-0.7 -1.40

-1.39

-1.60

[327]

[335]

[335]

[335]

BuNH+

BU2NH + BU3NH+

Me3NH+

PhN(Me2)H+

Me3N0H + PhN(Me2)0H+

-3.84

-3.83

-3.81 -3.72

-3.82

-4.39

-4.77

-4.72

[24]

[24]

[24]

[24] [335]

[24] [24]

[24]

''Log K corresponds to the reaction R4N+Pi" ^ R4N+ + Pi", where R may be alkyl, aryl, or hydrogen.

Some of the K values have been rounded off from three significant figures. Concentrations are in molar units. Bu signifies n-butyl. Approximate value; the constant is too large for accurate determination.

References:

[24] C. R. Witschonke and C. A. Kraus, J. Am. Chem. Soc. 69, 2472 (1947). [326] R. M. Fuoss and C. A. Kraus, J. Am. Chem. Soc. 55, 3614 (1933). [327] C. A. Kraus, J. Phys. Chem. 43, 231 (1939). [330] R. L. Mcintosh, D. J. Mead, and R. M. Fuoss, J. Am. Chem. Soc. 62, 506 (1940). [331] D. J. Mead, C. A. Kraus, and R. M. Fuoss, J. Am. Chem. Soc. 61, 3257 (1939). [332] D. S. Burgess and C. A. Kraus, J. Am. Chem. Soc. 70 , 706 (1948). [333] M. J. McDowell and C. A. Kraus, J. Am. Chem. Soc. 73, 3293 (1951). [334] M. B. Reynolds and C. A. Kraus, J. Am. Chem. Soc. 70, 1709 (1948). [335] E. G. Taylor and C. A. Kraus, J. Am. Chem. Soc. 69, 1731 (1947).

Table 33. Ion-pair dissociation constants of some incompletely

substituted ammonium salts containing different anions "

Solvent (e) Cation Anion Log K Refer- ence

Benzene (2.27) i-Am3NH+ Pi- -20.60 [326] Toluene (2.38)

BU3NH+ ci- - 18.82 [336] Ethylene chloride (10.23) PyH+' ClOj "-6.3 [327] PyH+ Pi- "-7.4 [327] Nitrobenzene (34.8) BuNHj CIOt -2.60 [24] BuNHj Pi -3.83 [24] BU3NH+ Pi- - 3.72 [24] BU3NH+ I- -4.02 [24]

* Log K corresponds to the reaction BH+A~ ;=i BH+ + A". Concentrations are in molar units, and ; = 25° unless otherwise

indicated.

" At 35°.

' For convenience, pyridinium ion (PyH+) is listed as being

a substituted ammonium ion.

" Approximate value. References:

[24] C. R. Witschonke and C. A. Kraus, J. Am. Chem. Soc. 69, 2472 (1947). [327] C. A. Kraus, J. Phys. Chem. 43 , 231 (1939). [336] V. Deitz and R. M. Fuoss, J. Am. Chem. Soc. 60, 2394

(1938). the more highly dissociated saU. Kraus regarded

these examples as further evidence of non-Coulombic attraction between cation and anion, per- chlorate ion being known to possess a weaker

affinity for the proton than picrate ion [327, 328]. Part of the results obtained by Kraus and Witschonke in studying ion-pair dissociation of picrates

in nitrobenzene [24] have already been shown in table 32. In additional measurements, summarized

44
in table 34 (as log K values), they studied three

picrates (aniUnium, dimethylanilinium, and py- ridinium) which undergo "acid-base dissociation"

(eq 18) as well as ion-pair dissociation (eq 17). Acidbase association constants are given in the table

in place of acid-base dissociation constants. Results

for anihnium picrate in acetonitrile , obtained by

French and Muggleton [337], are included. It should

be noted that the acid-base association constants were inferred from the conductance values and the

stoichiometry, not measured directly. This table indicates that acid-base behavior is likely to seem

more complex in dipolar aprotic solvents than in aprotic solvents in which, as we have seen (table 32),

ion-pair dissociations are an almost negligible

accompaniment of acid-base association.

We noted in discussing table 32 that picrates of

all types undergo ion-pair dissociation to about the same extent in the basic solvent pyridine and

attributed this behavior to formation of cations of

the type BH + ...Py. Any other base would be

expected to perform analogously, to an extent dependent on its electron-donor ability. Kraus and

Witschonke tested this assumption by adding

pyridine to nitrobenzene solutions of four picrates

which exhibit a strong tendency toward formation

of hydrogen-bonded ion pairs, namely, NH4Pi,

Bu;,NHPi, PhN(Me.,)OHPi, and Me.NOHPi (compare table 32) [24]. Water, triethylamine, and piperidine were also added, in varying amounts, to nitro- benzene solutions of MeaNOHPi. Examples of re- sults obtained are shown in table 35. As predicted.

Table 34. Values of log Kbha CL'^d log Kbh+. \~for some picrates undergoing both acid- base dissociation and ion-pair dissociation (25°) "

Solvent (e) Cation Log Kbha Log Kbh+. a- Reference

Nitrobenzene (34.8) PhNH+ 2.10 -4.7 [24] PhN(Me2)H+ 3.60 -4.39

PyH+ 4.79 -4.26

Acetonitrile (36.2) PhNHj 2.63 -2.98 [337]

" Log KsHA corresponds to the equilibrium B + HA ^ BHA, and log Kbh+. A-corresponds

to the equilibrium BHA^BH+ + A". Concentrations are in moles/liter.

References:

[24] C. R. Witschonke and C. A. Kraus, J. Am. Chem. Soc. 69, 2472 (1947). [337] C. M. French and D. F. Muggleton, J. Chem. Soc, p. 2131 (1957).

^ Condensed from results of C. R. Witschonke and C. A. Kraus, J. Am. Chem. Soc. 69, 2472 (1947).

Reference:

[24] C. R. Witschonke and C. A. Kraus, J. Am. Chem. Soc. 69,

2472 (1947).

adding any one of these bases, which except for water have dielectric constants far lower than that of nitrobenzene (see table 2), produced increases in the extent of ion-pair formation. Effects were proportionate to the concentration of added base, and also paralleled the accepted order of electron-donor

abihties of the bases.

d. Summary. In our brief review of conductance

experiments by Walden, Wynne-Jones, Kraus,

Fuoss, and their associates we have seen that all of the experiments indicate a strong tendency for hydrogen bond formation between cations and

anions of salts that are produced by combining

hydrogen acids with ammonia, amines, pyridine, or

tertiary amine oxides. The principal evidence is: (1) Such salts undergo far less ion-pair dissociation than their quaternary ammonium analogs in aprotic solvents and likewise in dipolar aprotic

solvents;

Table 35. Effects of bases on ion-pair dissociation constants of ammonium and partly substituted ammonium picrates in nitrobenzene (25°)

Cation Base Added Molarity of Base LogK

NHJ None -3.84

C5H5N 0.039 -3.63

BU3NH+ None -3.72 C5H5N 0.036 -3.63

PhN(Me2)0H+ None -4.70

C5H5N 0.035 -3.31

Me3N0H+ None -4.77 H2O 0.01 -4.65 C5H5N 0.001 -4.45 C5H5N 0.010 -3.92 C5H5N 0.037 -3.48

EtsN 0.001 -3.30

Et3N 0.038 -3.06 CsH.oNH 0.001 -3.28 CsH.oNH 0.010 -2.97

45
(2) Insignificant differences in ion-pair dissociation are observed when the sahs are transferred to a basic solvent such as pyridine or when a basic com- pound is added to their solutions in aprotic or dipolar aprotic solvents. We are justified in concluding that nonpolar and

dipolar aprotic solvents are differentiating solvents

because they are deficient in proton donor-acceptor

abilities. This similarity in the behavior of the two

groups of solvents enables us to make predictions

about acid-base behavior in nonpolar aprotic sol- vents by analogy with acid-base behavior in dipolar

aprotic solvents — or vice versa — when direct

experimental evidence is lacking.

Very little can be said about the extent of ionic

dissociation of hydrogen acids dissolved in aprotic or dipolar aprotic solvents. In the earliest recorded

experiments, Kablukoff studied the electrical conductivity of hydrogen chloride in carefully dried

benzene, xylene, hexane, and ether [338]. He re- ported that addition of this acid produced very

slight reductions in the resistance of the solvents. Such evidence of conductance is commonly attrib- uted to minute traces of water or other basic impurities [339]. According to Kraus [327], hydrogen

chloride dissolved in nitrobenzene is a very weak

electrolyte, while perchloric acid is a rather strong

electrolyte in the same solvent. La Mer and Downes

[340] stated that addition of trichloroacetic acid to benzene does not change the conductivity, at least

in amounts up to 0.04 M. Maryott [341] commented

that acids or bases have almost negUgibly low con- ductances in dilute benzene solutions.

We have stressed the minuteness of the tendency

for trialkylammonium picrates and related salts to undergo ion-pair dissociation in benzene and similar

solvents. However, it should be pointed out that acid-base titrations have been performed conductimetrically in benzene, by La Mer and Downes [340]

and by Maryott [341]. Such titrations are possible

because the conductance method is so sensitive that minute increases in ion concentrations accompanying salt formation are measurable. Moreover, the

concentrations used in the titrations referred to were much larger than those at which simple ions are most likely to be formed, and triple ions or still larger aggregates were probably more prevalent.

More will be said about triple ions in section 4.4.

4.3.2. Dielectric Constants

In the discussion of figure 9, which contains curves showing the conductance of salts dissolved

in benzene as a function of concentration, it was

stated that the shapes of the curves have been

explained by Kraus and associates [329] in terms of

equihbria between simple ions, ion pairs, and triple

ions (eqs 32, 33a, and 33b). Our attention will continue to be focused on ion pairs. Some of the best

direct evidence for the existence of this species comes from the dipole moments derived from dielectric constant data. Values of dipole moments of salts

obtained by several investigators are given in tables 36 to 38. The variation of the dielectric constant of salt solutions with concentration provides additional

information about association or dissociation of

dissolved molecules. For example, such curves have

demonstrated differences in the aggregation of

primary, secondary, and tertiary ammonium salts. Illustrative curves will be discussed in the following

section (see fig. 10). a. Dielectric polarization of substituted ammonium picrates and halides in benzene. Kraus and

co-workers [342] demonstrated that quaternary and

tertiary ammonium picrates and halides have large

polar (dipole) moments in benzene solution. Subsequently, Maryott [23, 343] confirmed this finding

and determined the dipole moments of additional

salts, including some primary and secondary ammonium picrates. Results of both authors are shown in table 36. It is clear that the dipole moments of the

salts are far larger than the moments of the component bases and acids, which are in the range

0-2 D (Debye unit). According to the more recent measurements (which are thought to be a little more accurate than the values first obtained), the

dipole moments of partly alkylated ammonium

picrates appear to be in the vicinity of 11-12 D, the

dipole moments of completely alkylated ammonium

picrates being several units higher. A similar dif- ference is seen on comparing the moments for Bu4NBr and BusNHBr. These values are consistent

with the assumption of ion pair formation. The

lower moments of the incompletely alkylated salts indicate a closer association of cation and anion,

in harmony with the postulate of hydrogen bonding

[329].

Figure 10 was constructed by Maryott [343], from

dielectric polarization data for several partly alkylated ammonium picrates, to compare effects of

altering the concentration of the solute and also changing the solvent from benzene to dioxane. In

this figure the increment Ae/A^2 is plotted against

A^2. In all cases, Ae/A^2 is quite high at low solute content. When dioxane is the solvent Ae/A^2 changes

very little during a 10-fold increase in A^2. The same

is true for benzene containing tributylammonium

picrate, in agreement with results of Kraus and

associates [328, 342] for triamyl- and tributylammonium picrates and tributylammonium chloride, bromide, and iodide. The interpretation given

[342, 343] is that the molecular state of these tertiary

ammonium salts is not affected by changes of concentration.

A decided drop of with increasing A'^2 is observed, however, for primary and secondary

ammonium picrates in benzene. The occurrence of

this fall-off and its absence in dioxane are strong

indications of association through hydrogen bonding.

Maryott [343] found his data for dipropyl- and

dibutylammonium picrates to be consistent with a monomer-dimer equilibrium, with the respective

log Ki,2 values 2.07 and 2.26 (converted to molar

46
Table 36. Dipole moments of some substituted ammonium salts in benzene "

Quaternary Ammonium Salts (25°)" Partly Alkylated Ammonium Salts (30°)

• Cation Anion fj,, Debye - Cation Anion fjL, Debye units

units

i-Am4N + Pi- 18.3 i-Am3NH + Pi- 11.9 "13.3 j-Am4N+ SCN- 15.4 BU3NH + Pi- 11.9 (12.2) " 13.1 " 11.4 BU4N + Pi- 17.8 " 15.0 Et3NH + Pi- 11.7 BU4N + ClOj 14.1 Et2(PhCH2)NH + Pi 11.8 BU4N + Br- 11.6 (PhCH2)3NH + Pi- 12.0 BU4N + CH3COO- , 11.2 BU3NH + I- "8.1 "7.7 BU.3NH + Br- "7.6 BU3NH + ci- "7.2 Bu.NHj Pi- 11.5 (12.1) Pr2NH + Pi- 11.5 C«H,7NH3+ Pi (12.1) C,aH37NH+ Pi- (12.3)

^ A few of the measurements were made in other solvents, as indicated in footnotes

c and d. All alkyl radicals were straight-chain except i50-amyl. "From work of C. A. Kraus and co-workers as summarized in ref. [342]. Results of A. A. Maryott (30°) [23, 343] except as noted otherwise. Values given in parentheses were determined in dioxane.

" Determined by Davies and Williams (18-20°) [344] in benzene, dioxane, ra-xylene, or mixtures of these solvents.

References:

[23] A. A. Maryott, J. Res. NBS 41, 7 (1948). [342] C. A. Kraus, J. FrankUn Inst. 225, 687 (1938). [343] A. A. Maryott, J. Res. NBS 41,1 (1948). [344] M. Davies and G. WiUiams, Trans. Faraday Soc. 56, 1619 (1960).

units). Aggregation of the primary alkylammonium

picrates apparently is still more extensive.

We have just noted (fig. 10) that in the case of

tributylammonium picrate in benzene the ratio

is not measurably affected by altering the

concentration of the salt. However, Maryott ob- served [23] that decreases with increasing

dilution of tribenzylammonium picrate in benzene,

indicating that this salt undergoes partial dissociation into picric acid and tribenzylamine. He devised

a procedure for calculating both the dipole moment

of the salt and the (1:1) acid-base equilibrium con- stant (at 30° and 40°). The dipole moment (see table

36) was found to be essentially the same as values

obtained for stable picrates. It is noteworthy that

this method of determining an acid-base equilibrium constant yielded values in almost perfect concord- ance with one obtained from spectral absorbance

data (25°) for solutions containing a fixed concentration of picric acid and varying quantities of tri- benzylamine [20] .

b. Dielectric polarization of other acid-base complexes in benzene. Numerous additional measure- ments of dipole moments of salts and other acidbase complexes have been reported in recent years, but it seems hkely that in some instances the values

are too low. These are instances in which experi220

200

OcNH,Pi (Dioxane) Bu,NHPi (Benzene)

OcNH,Pi (Benzene)

10

Figure 10. Variation ofdielectric constant of benzene or dioxane

with concentration of dissolved primary, secondary, or tertiary ammonium picrate. The rate of change of the dielectric con

stant with changing mole fraction of solute (^elN-z) is plotted

against N2.

Reprinted with permission from A. A, Maryott, J. Res. NBS 41. 1 (1948). (Ref, No. 343,) 47
mental conditions, only sketchily described, appear

to have been such that the acid-base complex would

be partly dissociated into the constituent acid and

base, and the description (also sketchy) of the

method used in calculating dipole moments does

not indicate that the investigators were aware of, and made proper allowance for, the acid-base

dissociation.

Table 37 contains results of Sobczyk and Syrkin

[345, 346] for eight 1:1 acid-base complexes in benzene. The bases represented comprise five ahphatic amines (two tertiary, two secondary, and

one primary) in addition to pyridine and two related

bases. The acids represented are picric acid and

five carboxylic acids (benzoic, propionic, and the

three chloro derivatives of acetic acid). It will

participate in hydrogen bonding, as indicated by

formula XXXII.

,0...H ,R

\0-H--- \

(XXXII)

Tertiary aliphatic ammonium carboxylates were

found to have the same dipole moment in benzene

as in dioxane, and are thought to be unassociated

in these two solvents. In contrast, secondary am- monium salts dissolved in dioxane were observed

to show increased polarity with dilution, becoming more like tertiary ammonium salts. Sobczyk and

Syrkin also concluded that piperidinium picrate

is strongly associated in benzene. These findings

parallel Maryott's results (see fig. 10 and related

discussion).

be noted that the dipole moments given for the

picrates of pyridine-type bases are only slightly

lower than the values for picrates of stronger bases

(compare table 36). Sobczyk and Syrkin have concluded that the

dipole moments of carboxylates of tertiary aliphatic

amines and pyridine-type bases parallel in magnitude the proton-donor abilities of the acids

involved. However, the strengths of the acidic

components are not thought to affect measurably

the dipole moments of secondary ammonium car- boxylates, which are found to be much lower than

the moments of either tertiary aliphatic ammonium

or pyridinium-type carboxylates. These extremely

low moments are attributed to a stable monomeric

structure in which both of the ammonium hydrogens

HPi

" 11.7

10.1 10.3

10.9

Table 38 contains results obtained by Gur'yanova

and Beskina [347] in an investigation of the association of benzoic acid in benzene, at a stoichiometric

concentration in the range 0.034-0.123 M, with

successive portions of the following bases: Two

primary amines (n-BuNH), PhNH2), four secondary

amines (EtaNH, i-ProNH, n-BuaNH, C5H10NH), and

two tertiary amines (EtsN, n-AmsN). They per- formed both "dielectrometric titrations" and

"cryoscopic titrations," and made the following

deductions from the combined results:

(1) At the concentrations used, benzoic acid exists as the dimer and it reacts as this species.

(2) All of the amines except aniline reacted with

benzoic acid to form products with large dipole

moments. This was shown by the immediate sharp

rise in dielectric permeability of solutions on com- mencing to add an amine. The rise in permeability

continued until the stoichiometric ratio CbjCa

Table 37. Dipole moments of some 1:1 acid-base complexes in benzene^

Base

p., Debye units (20-25°), where acid was:

CsHsCOOH C2H5COOH CH2CICOOH CHCI2COOH CCI3COOH

CjH.oNMe

Et:,N CjH.oNH

EtaNH

EtNH.

Pyridine

Quinoline

Acridine

3.87

3.72

2.39

2.86

2.88

2.75

5.97 6.37

2.87

2.17

(1.52) 4.67 4.77

4.98

7.87 8.78

8.82 3.15

7.78 8.32

8.60

^ From work of L. Sobczyk and associates [345-346].

" Result of A. A. Maryott (30°); see table 36.

References:

[345] L. Sobczyk and J. K. Syrkin, Roczniki Chem. 30 , 881, 893 (1956): 31, 197 (1957). CA 51, 7786, 14347 (1957). [346] L. Sobczyk, in Hydrogen Bonding (Ljubljana, 1957), (D. Hadzi and H. W. Thompson, Eds.), p. 323 (Pergamon

Press, New York, 1959).

48
Table 38. Dipole moments of benzoic acid-amine complexes in benzene

Amine (/a)

ft (25°), Debye units

B(HA)2

BHA<^ .(BHA)2'^

n-BuNH2 (1.44) EtaNH (1.26) j-PrzNH (1.20) n-BujNH (1.27) CsH.oNH (1.09) EtaN (1.02) n-AmsN (1.03)

(3.92) 4.62

4.80

4.69

4.57 4.88

4.72

3.02

2.88

2.75

3.11 3.28 4.38

3.87

^ Results of E. N. Gur'yanova and I. G. Beskina [347]. Tlie stoichiometry was determined cryoscopically. Intermediate product. B = amine, HA = benzoic acid. Final product.

R ,R \ /

H / •.

0 \0 Ph-C^ ^C-Ph \-H 0^

(XXXIIIa)

\/^

N / •. ,H H

,0'" "^0

Ph-c^ ;c-ph

\

/"

N

H

.0 \

^0—H 0^

(XXXIIIb)

,0 R / /-R Ph-C^ Nr^

(xxxv)

Reference:

[347] E. N. Gur'yanova and I. G. Beskina, J. Gen. Chem. USSR

Eng. Transl. 33, 914 (1963).

equalled 0.5. At this stage the formation of a complex of the composition B(HA)2 was inferred.

(3) In the interval where CbICa was in the range

0.5-1.0 complex formation followed a different pattern for primary and secondary amines than for

tertiary amines. With tertiary amines the dielectric permeability rose slightly, while with the other two

classes of amine there was a decrease. From these

two patterns of behavior and the cryoscopic data

it was deduced that primary and secondary amines

formed benzoates of the composition (BHA)2 whereas tertiary amines formed benzoates of the general

formula BHA.

(4) After attaining the ratio CblC„ = 1 all systems

again behaved similarly, there being no indication

of the formation of any additional species. Structures postulated for the different species —

B(HA)2, (BHA)2, and BHA -are indicated below

(XXXIII-XXXV). Formulas XXXIIIa and XXXIIIb

apply, respectively, to secondary and tertiary ammonium benzoates.

Piperidine-benzoic acid mixtures are the only ones measured both by Gur'yanova and Beskina

and by Sobczyk and Syrkin. As we have seen, they

agree that both hydrogens attached to a secondary

ammonium nitrogen are involved in hydrogen bonding, but their conclusions differ in detail. Whereas

the latter pair of authors, who postulated only 1:1 acid-base association, yielding the structure BHA,

calculated 2.39 as the dipole moment of piperidinium benzoate (in benzene), two higher values —

4.57 and 3.28 — were obtained in the work of

Gur'yanova and Beskina as the moments of the

salts postulated to have the respective compositions B(HA)2 and (BHA)2.

In summary, those who have studied acid-base

associations in benzene by measuring changes in

(miv)

the dielectric polarization of systems similar to the

examples given consider hydrogen bond formation

to be involved, and there is strong evidence that multiple hydrogen bonding occurs in secondary

and primary ammonium salts (at least, those of the

types which have been discussed here). Unfortunately, much of the dipole moment data so far reported for weak acid-base complexes is of questionable validity because the extent of complex

formation does not seem to have been given adequate consideration. IMoreover, the assumption of 1:1 acid-base association as the only form of combination may not have been warranted in some

cases. The value of the method is enhanced when

used in ponjunction with other types of measurement.

4.3.3. Colligative Properties

The study of colligative properties is an important

adjunct to investigations of possible hydrogen bonding in ion pairs. For example, in the preceding section we noted that Gur'yanova and Beskina, when

studying the structures of some carboxyUc acidamine complexes in benzene, performed cryoscopic

measurements in addition to dielectric polarization

measurements, and that the molecular weights

deduced from cryoscopic data led them to postulate

a dimeric hydrogen-bonded structure for a secondary ammonium carboxylate rather than the monomeric hydrogen-bonded structure postulated by

Sobczyk and Syrkin.

An earlier example of conclusions based in part on a study of colligative properties is found in the

extensive investigations of physical properties of

salts in benzene solutions made by Kraus and

associates. The salts investigated, though mainly

quaternary ammonium salts, also included a few

tertiary ammonium salts. Part of their deductions

from conductance and dielectric polarization meas- urements have been presented (sections 4.3.1c and
4.3.2a). Results which they obtained from cryoscopic

determinations will be summarized in this section,

together with overall conclusions [348-351].

A newer method of studying coUigative properties

is the "differential vapor pressure (DVP) method"

used by Bruckenstein and Saito [352] for investigating various acid-base equilibria in benzene.

a. Cryoscopy. The salts studied by Kraus and

coworkers included three tertiary ammonium salts,

namely, tri-t-amylammonium picrate [348, 350],

tri-n-butylammonium picrate [349], and tri-n-butylammonium iodide [349, 351]. Freezing point

determinations were made over a wide range of concentrations: from 0.001 molal to the highest concentration which the solubility made feasible. When

values of n (the apparent molecular weight of a

salt divided by the formula weight) were plotted as a function of the logarithm of the molal concentration

it became apparent that the solutes studied may be

classified in three general types [329]. These types

are illustrated in figure 11 by the following examples: Type 1, i-AmgNHPi; type 2, i-AmjNSCN;

type 3, f-Am4NPi.

30

LOG m

Figure 11. Relation of association numbers (n) to the molal

concentration, as determined cryoscopically for three different types of salts in benzene.

Reproduced with minor changes from C. A. Kraus, J. Chem. Educ. 35, 324 (1958), with the permission of the author and the publisher. (Ref. No. 329.)

Figure 11 shows very strikingly that n, which is a measure of the extent of self-aggregation of the

solutes, is much greater for tetra-tso-amylammonium

thiocyanate (type 2) than for the related tertiary

(type 1) and quaternary (type 3) ammonium picrates

Moreover, the curve for the thiocyanate displays

a pronounced maximum at concentrations near

0.1 molal, but there is no indication of a maximum

in the other two curves. Instead, the degree of

aggregation rises very slowly with increasing

concentration.

Table 39 contains additional evidence regarding

the extent of aggregation of tertiary and quaternary

ammonium salts in benzene. Of the salts studied,

the two tertiary ammonium picrates (which show

similar behavior) have the lowest tendency to form

aggregates. For example, at 0.1 m concentration the

aggregation of i-AmsNHPi is still not extensive

(n=1.54). At the same concentration aggregation

of n-BusNHI is somewhat greater (n — 2.95). Tetra-f- amylammonium picrate, the only salt assigned to type 3, seems to be a little more highly aggregated,

but its low solubiUty in benzene prevented measure- ments at concentrations greater than 0.0038 m. The much more highly associated salts assigned to type 2 (that is, the quaternary ammonium thio- cyanates, iodide, and perchlorate) show individ- ual variations in behavior: the maxima in n versus

log m curves are different in magnitude and occur

at different concentrations.

The examples given in figure 11 are the same as

in figure 9, which displays three types of conduct- ance behavior in benzene as a function of concentration. In previous discussion of figure 9 (sec. 4.3.1c)

it was pointed out that portions of such conductance

curves have been interpreted in terms of overlapping

equilibria which involve simple ions, ion pairs,

triple ions, and more complex (undefined) aggregates. It is noteworthy that the conductance curve

for i-Am4NSCN, which was found cryoscopically

to form very large aggregates, has more inflections

than the curves for t-Am4NPi and i-AmaNHPi,

indicating the existence of a greater number of

competing equilibria.

There is now a general recognition that dissociation of salts into cations and anions in water and

other good ionizing media is well explained in terms of ion-solvent interactions. Obviously, when

ion-solvent interactions are virtually non-existent,

solute-solute interactions will assume greater importance, and they will be expected to vary in strength with such factors as the charge, composition, polarizability, size, structure, and symmetry

of ions, as well as on their concentration. Such

expectations have been borne out in the extensive

investigations of tertiary and quaternary ammonium

salts in benzene by Kraus and associates. As we

have seen, these investigators discovered such salts

to exhibit at least three differing patterns of associative behavior, as well as individual differences. A

partial summary of their overall conclusions [348-

351, 329] follows.

(1) In the most dilute salt solutions studied (below

the first minimum in conductance), ion pairs are

the predominant species, although there is a measurable concentration of simple ions (see 4.3.1a).

Triple ions, (H h) or ( —I—), are not present in measurable quantities.

(2) Triple ions are formed in increasing proportion as the concentration increases. At first, their formation wiU not affect the proportion of simple

ions appreciably, but a concentration will be reached

where increasing conductance caused by triple

ion formation will be offset by a decreasing conduct- ance resulting from association of simple ions to form ion pairs. At this point the conductance will have its minimum value, since further increases in

50
Table 39. Variation with concentration of the association numbers (n) of some salts in benzene"

Moles of salt per 1000 g of benzene

0.001 0.01 0.02 0.05 0.1 0.2 0.3 0.4 0.5

n-BusNHPi

i-AmaNHPi

A!-Bu,NHI AgC104

i-Am4NPi''

n-Bu4NSCN

i-AmjNSCN

n-AmN(n-Bu:,)I «-Bu4NC104 _

1.01

1.01

1.12 1.15

1.3

2.4

3.2

3.1 3.0

1.07

1.07

1.4

1.7

(1.12)

(1.6)

(1.30)

(2.25)

(1.54) (2.95)

(1.91)

(4.4)

(2.18)

(5.75)

(2.40)

(6.65)

(2.52)

(7.15)

4.2

7.8 , 4.4

5.0

(6) J0.9

(5.8)

(12) (19.5)

(9.8)

(20) (25.6) 16.6

(30) (23.5) (21.6)

(31.6)

(15.5)

(21.4)

(20.7)

(12.2)

(19.7)

(29)

17.0

^ From results reported by Copenhafer and Kraus (see their tables I and II) [350] and by Young

and Kraus (Table IV) [351]. Values in parentheses were obtained by interpolation.

" At the solubility limit (0.0038 m) n = 1.4.

References:

[350] D. T. Copenhafer and C. A. Kraus, J. Am. Chem. Soc. 73 , 4557 (1951). [351] H. S. Young and C. A. Kraus, J. Am. Chem. Soc. 73 , 4732 (1951).

concentration will increase the probability of triple

ion formation.

(3) A succeeding stage is quadrupole formation.

To account for dielectric polarization data it is postulated that in quadrupoles the constituent

dipoles are sometimes in parallel and sometimes in anti-parallel alignment. Figure 11 and table 39 indi- cate that f-AmaNHPi is only 6 percent associated to quadrupoles at the lowest concentration studied

cryoscopically, while at the highest concentration,

association has reached a little more than the

quadrupole stage. It will be recalled that no signifi- cant self-association of tertiary ammonium salts was detectable in dielectric polarization measurements, which are made at much lower concentrations (see 4.3.2a). Association of i-AmaNPi is some- what more extensive, about 40 percent being

associated to quadrupoles at the highest concentration studied (0.0038 m).

(4) Except at very low concentrations the experimental results have not provided definite clues as

to the composition of aggregates.

(5) The type of behavior exhibited by a salt seems

to depend on the following constitutional factors: In type 1, the salt is composed of two large ions

which are hydrogen bonded; in type 2, a large

cation is combined with a smaller anion; and in type 3, the cation and anion are both large, as in type 1 (but are not, of course, hydrogen bonded).

(6) The following are some characteristic differ- ences in the physical behavior of the three types:

Type 1. The ion-pair dissociation constant

(table 32) and dipole moment (table 36), as well as

the extent of aggregation, are comparatively small.

(All these properties are consistent with the hypothesis of hydrogen bonding.) The conductance minimum

appears at a relatively high concentration, and the

increase in conductance between the minimum and

maximum is great.

Type 2. The ion-pair dissociation constant and

the dipole moment are both relatively large. The

first conductance minimum lies at much lower

concentrations than for type 1, and the increase

between minimum and maximum is much smaller.

Conductance increases occur in a series of steps.

Type 3. The ion-pair dissociation constant and

dipole moment are both somewhat larger than for type 2. (As previously noted, only one example of

this type has been observed, and the highest con- centration studied was 0.0038 M.)

It can hardly be expected that phenomena observed for the mixtures frequently dealt with can

be readily understood when concentration effects on the behavior of individual solutes in aprotic

solvents are so varied and imperfectly understood.

The need for much further study of aprotic solu- tions is clear. In the present state of knowledge the

most dilute range appears best suited for attempts

to ascertain quantitative relationships, with proper

attention to possible competing interactions of

solutes with glass and other container materials.

b. The differential vapor pressure (DVP) method

[352]. In this method the state of aggregation of

solute particles is determined by measuring the

difference in temperature of a pure solvent and a

solution having the same equifibrium partial pres- sure of solvent. Interest in its use has been stimulated by recent developments in thermistors and

other new instrumentation. It has the great merit

of permitting flexibility in the choice of temperature.

Bruckenstein and Saito [352] explored the applicability of the method to the study of acid-base equihb51
ria in benzene. The reacting base (B) and acid (HX) were contained in a thermostated chamber (apparently held at or near the laboratory temperature).

The reactants included several carboxylic acids of

the acetic or benzoic acid series and representative

primary, secondary, and tertiary aliphatic amines.

After the determination of by measurements of

voltage changes, values of i (eq 34) were plotted as a function of R (eq 35), as indicated in figure 12.

i = (Af for B + HX, Ch = Jc-molar)

/(Ai for CioHs soln., x-molar) (34)

R — ChxICq (35)

Total Acid / Total Base Figure 12. Hypothetical acid-base titration curves in benzene;

see text for explanation.

Reproduced with minor changes from S. Bruckenstein and A. Saito, J. Am. Chem.

Soc. 87, 69b (1965), with the permission of the authors and the pubhsher.

(Ref. No. 352.)

The analytical (stoichiometric) concentration of the

base, symbolized here by Cb, was usually 0.01 M,

but was also 0.05 M and 0.1 M in some experiments.

The several titration "curves" shown in figure 12 are hypothetical, and correspond to various situa- tions in which the products formed have compositions such as BHX, BHX • HX, and (BHX)2. The

shape of a curve will be modified if a reactant is dimeric or partly dimerized. The hypothetical curves correspond to the five situations described

below.

(1) ABC: B and HX are monomers and react

quantitatively, the only reaction product being

BHX.

(2) A'BC: Like (1) except that B is partly

dimerized.

(3) ADE: B and HX, which are monomers, form

BHX and then BHX • HX. There are no additional

products.

(4) ABC: Like (3) except that a little BHX • HX

is formed at the same time as BHX.

(5) AFDE: Monomeric B and HX associate in

1 : 1 ratio to form (BHX)2. (No other product is formed.)

The curves will have rounded portions (see short

dashes) whenever reactions do not proceed

quantitatively.

Possible use of the DVP method for determining

the self-association constants of the bases and acids was tested. The i-values obtained for 0.01-M amines

differed from 1 by no more than 5 percent, but i for 1,3-diphenylguanidine (another base included in the investigations) indicated partial self-aggregation

(compare sec. 4.1.3). That is, the value of i (0.87) cor- responds to 78 percent monomer, the remainder be- ing the dimer (if these are assumed to be the only

species present). The method failed to yield accurate dimerization constants for the carboxylic acids.

However, the constant obtained for benzoic acid was in approximate agreement with results by other

methods (see table 24). An example of the third type of curve (ADE) was

obtained on adding trichloroacetic acid to the

tertiary amine benzyl-7V,A'^-dimethylamine; that is, base and acid reacted quantitatively to form monomeric BHX, which added further HX to give

BHX • HX. In contrast, 0-01-M dibenzylamine,

though associating in 1 : 1 ratio with trichloroacetic

acid, yielded (BHX)2; further acid gave a practically

quantitative yield of the acid salt BHX • HX. (In

both cases nonquantitative formation of BHX(HX)2

was thought to occur on the continued addition of

HX.) These conclusions as to the composition of

the salt species formed parallel those of Gur'yanova

and Beskina using benzoic acid instead of tri- chloroacetic acid (see table 38). As noted above, Gur'yanova and Beskina deduced

hydrogen-bonded structures for secondary and

tertiary ammonium benzoates in benzene from

their dielectric polarization data (compare formulas

XXXIII-XXXV), and related acid-base systems

studied in carbon tetrachloride are believed from

infrared evidence to form hydrogen-bonded ion

pairs. However, Bruckenstein and Saito, who also performed infrared measurements, consider their infrared data to indicate the absence of hydrogen

bonding. Infrared evidence wiU be considered

further in section 4.3.5.

4.3.4. Spectral Absorbance in the Visible and

Ultraviolet Regions

The differentiating character of benzene as a sol- vent has been very strikingly illustrated in systematic studies of acid-base reactions employing an

indicator dye as one of the reactants. The most

extensive investigation of this type so far performed

has been reported in a series of papers dealing with

visual and spectrophotometric observations of the

behavior of bromophthalein magenta (BPM) —

which is a phenolic indicator of the triphenylmethane class — following the addition of different types of organic bases. Solutions of bromophthalein

magenta in benzene and other aprotic solvents are yellow in color. The yellow tone can be converted

52
at will to practically any color of the rainbow by

adding the proper base in a suitable amount. Such

behavior is, of course, a marked departure from

indicator behavior in aqueous solutions, where all bases (or all acids) cause uniformly similar color

changes. The extraordinary range of colors exhibited

in reactions of bromophthalein magenta with bases

in aprotic solvents has received a logical explanation in terms of hydrogen bonding [14, 5, 21, 31, 15]. This section will be devoted largely to describing

and discussing the color reactions of bromophthalein magenta. However, it should be emphasized at the outset that this compound is by no means unique

in signaling the occurrence of hydrogen bonding

through easily observable changes in electronic

absorption spectra. For example, several indicators

which are widely known through their extensive use

in aqueous solutions have been found to show behavior analogous to that of BPM when in aprotic

solutions. Examples are halogen derivatives of

phenolsulfonephthalein (phenol red), such as bromophenol blue and bromocresol green [29]. Further

examples have been found among nitrophenols

such as picric acid [20], trinitro-m-cresol (methylpicric acid) [20], and 2,4-dinitrophenol [17]. Hydrogen bonding has also been inferred from the occurrence of spectral shifts in the near-ultraviolet

region — for example, in studies of the association of

a- and /3-naphthol with triethylamine and other

proton acceptors in n-heptane solution [353]. Two papers on spectrophotometric studies of

complex anions containing actinide elements

[413, 414] will be referred to in section 4.3.5c.

a. Bromophthalein magenta. The accompanying

formulas show that bromophthalein magenta

(XXXVII) bears a close structural relationship to 3',5',3",5"-tetrabromophenolphthalein (XXXVI) and

also to bromophenol blue (3',5',3",5"-tetrabromophenolsulfonephthalein) (XXXVIII). Davis and

Schuhmann, who studied both the ethyl and the

n-butyl ester of tetrabromophenolphthalein (XXXVII,

R=Et or n-Bu), adopted the respective names and

symbols "bromophthalein magenta E" (BPM-E)

and "bromophthalein magenta B" (BPM-B) for these compounds [14]. The two esters differ slightly

in physical properties, such as the lower melting

point and greater solubility in hydrocarbons shown

by BPM-B, but no differences in their chemical

behavior as phenols were observed.

It is readily perceived that (XXXVI) may be re- garded as the isomer of a hypothetical carboxylic

acid in which the proton forming part of the -OH

group attached to ring 1 has migrated to convert the

factoid COO- to a -COOH group. The yellow

color of the esters (XXXVII) is explained by the

quinoid structure of ring 1. Evidently the lactone

ring of (XXXVI) is too stable for ready conversion

of (XXXVI) to its yellow quinoid isomer or for salt formation with added bases when in benzene solution. In contrast, the esters react instantaneously

with alkylamines and other bases of comparable

strength — not, however, with bases as weak as

aniline and pyridine.

OH

(xxxvi) (xxxvn)

(XXXVIII) (XXXIX)

"Sultones" (XXXVIII) are much more reactive

than lactones (XXXVI). Thus, bromophenol blue,

although it gives a colorless solution in dry benzene,

reacts instantaneously in 1:1 ratio with added

bases, strong or weak, to form the yellow salt indi- cated by (XXXIX). The spectral absorbance of

(XXXIX) is very similar to that of the corresponding

form of a closely related indicator, bromocresol

green (3',5',3",5"-tetrabromo-2',2"-dimethylphenolsulfonephthalein). Transmittance curves for the

latter indicator are shown in figure 13. The three

WAVELENGTH, nm

Figure 13. Spectral transmittance of bromocresol green, 5X10~^ M in benzene: As prepared (faintly dashed line); in the presence of five molar equivalents of acetic acid (curve A ); and in the presence of one molar equivalent of mono-, di-, or tri-n-amylamine ( curves I to 3, respectively).

Reproduced from figure 5 of M. M. Davis, P. J. Schuhmann. and M. E. Lovelace,

J. Res. NaU. Bur. Std. 41, 27 (1948). (Ref. 29.) 53
essentially identical curves numbered 1 to 3 were

obtained by adding one equivalent of mono- di-, and

tri-n-amylamine, respectively. The absorbance of

(XXXIX) is very similar to that of (XXXVII) [29]. The "first stage" in the neutralization of bromophenol blue and other sulfonephthaleins is not observed in their aqueous solutions, as the sultone

structure (XXXVIII) is unstable in water and

yields the yellow ionized species analogous to (XXXIX). In solvents like benzene, toluene, chloro- benzene, and anisole it has provided a means for estimating the relative strengths of aniline or pyridine derivatives [85, 354] and of several carboxylic

acids [355]. Quaternary ammonium salts of bromophthalem

magenta E or B (XL) are found to give blue solutions.

(XL)

whatever the class of solvent used (see table 40,

last column). The shift to a deep color is explained

by the resonance hypothesis. One obviously

important valence-bond structure is that in which

the electronic patterns of rings 1 and 2 are quinoid

and phenolate, respectively, and a second important

structure is the one in which these electronic pat- terns are reversed. Figure 14 shows the type of

€ K 10-3

350 450 550 650

Wavelength, nm

Figure 14. Spectral absorption curves of bromophthalein

magenta E (or B) and several of its salts in benzene solution.

See M. M. Davis and P. J. Schuhmann, J. Res. NBS 39, 221 0947), figures 21 and 22. (Ref. Nu. 14.)

spectral absorption curve obtained for quaternary

ammonium salts of bromophthalein magenta dis- solved in benzene (see curve labeled "BU4N salt")

[14]. The most striking feature of the absorption curve is a strong band of symmetrical appearance,

centered near 605 nanometers (millimicrons). As

mentioned already, the solutions have a blue tone by transmitted light. Benzene solutions containing

a quaternary ammonium salt together with un- changed bromophthalein magenta will, of course, have colors ranging from green-blue to yellow- green, according to the relative proportions of the

acid and its quaternary ammonium salt. Solutions containing tertiary ammonium salts of

bromophthalein magenta are strikingly dissimilar in Color to solutions of quaternary ammonium salts of

this acid. The dissimilarity can be appreciated by

comparing the absorption curves given in figure 14

for the two types of salt. It will be noted that the curve labeled "EtaNH salt" (which is typical of

absorption curves obtained for tertiary ammonium

salts) has Amax. near 540 nanometers (nm) instead of near 605 nm. It resembles the curve for the «-Bu4N

salt in being symmetrical, but it is broader and the

intensity of absorption is only about half as great.

It is well to take note here that the usual procedure

for obtaining absorption curves for tertiary and other

incompletely substituted ammonium salts is to add

an amine to a benzene solution of the indicator acid,

at first in minute quantities and then in gradually

increasing amounts until a limit has been reached

beyond which shifts in absorbance are no longer

measurable. (It is assumed in this discussion that the base itself is colorless in the visible region and

that its cation also is colorless.) The need for an

excess of amine is apparent when we recall that

salts of the type BHA when in aprotic solvents are

in equilibrium with the constituent acid and base.

In the reaction of bromophthalein magenta with

triethylamine in benzene, the association constant is ~ 2.7 X 10" (molar units, 25°) [15]. With tertiary amines, the limiting color is magenta.

As would be expected, the intermediate colors ob- served are various orange and scarlet tones. Table

40 shows (see next to last column) that tertiary am- monium salts of bromophthalein magenta display

the magenta color when dissolved in hydrocarbons,

halogenated hydrocarbons, ethers, and carboxylic

acid esters (tributyl phosphate behaves differently

[14]). On the other hand, in ketones, alcohols, bases

like aniline or pyridine, and in nitroparaffins or nitroaromatic solvents, the characteristic color is the same as for quaternary ammonium salts. In interpreting these observations, Davis and co- workers adopted an explanation along the same lines as those offered by Wynne-Jones [325] and by Kraus

[327] (see section 4.3.1 and 4.3.2). That is, they con- cluded that in tertiary ammonium salts of bromophthalein magenta E and B the cation is hydrogenbonded to the anion, as indicated by (XLI). A

more general representation of this structure is BH+...A- (or BH+...OR) [14]. The magenta tone.

54
Table 40. Effect of solvent and added base on the color of Bromofihthalein Magenta (BPM)

Solvent Type of base added'

0 I II III IV

Hydrocarbons:

Benzene (2.3) Carbon tetrachloride (2.2, 20°) Cyciohexane (2.0, 20°) Toluene (2.4)

Yellow

Yellow

Yellow Yellow

Red-purple

Red-purple

Red-purple

ivcu pui pic

Purple-blue

Purple-blue

Purple-blue

Purple-blue

Magenta

Magenta

Magenta

Magenta

Blue."

Blue."

Blue."

Di ue.

Halogenated Hydrocarbons:

Chlorobenzene (5.6) Chloroform (4.8, 20°) 1,2-Dichloroethane (10.2)Yellow

Yellow

Ypllf)w

Red-purple

Red-purple

IVla^C ivl P ( ) ri t Q 1 1 Id.

Red-purple

Red-purple

]VT iVla^Cll M treti-i t ala

Magenta

Magenta

]Vlagenta

Blue."

Blue."

Dl Uc.

Ethers:

Anisole (4.3) 1,4-Dioxane (2.2) Ethyl ether (4.3, 20°)Yellow

Yellow

Yellow

Magenta

Magenta

Magenta

Magenta

Magenta

Magenta

Magenta

Magenta

Magenta

Blue."

Blue." Blue."

Esters:

Benzyl benzoate (4.9, 20°) Butyl stearate (3.1, 30°) Ethyl acetate (6.0) Methyl benzoate (6.6, 20°) Tributyl phosphate (8.0, 30°)Yellow

Yellow

Green-yellow Green-blue

iji uc^l cell

Red- purple

Magenta

Red-purple

Red-purple oiue

Red-purple

Magenta

Red-purple

Red-purple

Rliip <"

iviagenia Magenta

Magenta

Magenta

Blue."

Ketones:

Acetone (20.7) Cyclohexanone (18.3, 20°) Methyl isobutyl ketone (13.1, 20°)Yellow to blue " Yellow to ^reen*^ Yellow to blue*"Blue"

Blue " Blue"

Blue" Blue"

Blue"

Blue " Blue" Blue"Blue."

Alcohols:

Ethanol, 95% (27.9?) ra-Propyl alcohol (20.1) n-Butyl alcohol (17.5) 2-Heptanol (9.2, 22°)

Blue"

Green-blue

Green-yellow

Blue"

Blue " Blue" Blue"

Blue"

Blue " Blue " Blue"

Blue" Blue"

Diue Red-purple

Blue."

Nitrogenous Bases:

Aniline (6.9, 20°) Pyridine (12.3)

Green-blue

Green-blue

Purple-blue Green-bluePurple-blue Blue"Purple-blue Blue"

Blue."

Micellaneous:

Acetonitrile (37.5, 20°) Nitrobenzene (34.8) Nitromethane (35.9, 30°)

Green-blue Yellow

Yellow

Blue" Green ' Purple-blueBlue"

Green-blue

Blue"

Blue" Green

Blue"

Blue."

^ Visual observations using BPM-B; although BPM-B is more soluble in hydrocarbons than BPM-E no differences in chemical

behavior have been observed. This is a condensation of table 7 of M. M. Davis and P. J. Schuhmann, J. Res. NBS 39, 221 (1947) [14].

^ Numbers in parentheses are values of e, at 25° unless otherwise indicated.

' Key to identity of bases: 0, none; I, n-CsH.jNH.; II, n-BuaNH; III, Et;,N; IV, /i-Bu^N salt of BPM-B.

" Blue without purple tinge. Yellow in concentrated solution, bluer on dilution. 'The yellower color tones observed in PhN02 are to be expected because the solvent itself is yellow.

55
intermediate in the color scale between yellow and

blue, is consistent with the hypothesis that the pro- ton exerts a residual effect on the anion, as indicated

in (XLI). It should be emphasized that the cation is regarded as being held by a specific bond to a phenolate-type oxygen. (It should also be noted that such bonds are not as stable as ordinary chemical

bonds and are regarded as being continually broken

and re-formed, like other kinds of hydrogen bonds.)

In summary, the formation of hydrogen-bonded ion

pairs is regarded as being an "incipient ioniza- tion ... an intermediate step in the complete

transfer of a proton . . ." [14, 20, 5, 25]. The difference in the structure of quaternary

ammonium salts and that of incompletely substi- tuted ammonium salts is conveniently pointed Up

by referring to the first group as consisting of "ion

pairs" [326-329] and the latter group as consisting

of "'hydrogen-bonded cation pairs" [356, 31]. Rupture of the hydrogen bond in tertiary am- monium salts (R3NH + ...X") can occur either so

as to yield R3N and HX or so as to form R3NH + and X~. The second type of change appears not to occur — at least with tertiary ammonium salts of

bromophthalein E or B — unless a second base is present to form a new type of hydrogen bond (XLII)

BaseR

I H—N—

I R

(XLII)

with the cation. This reaction may be expressed in general terms by eq 36 [24, 327, 5, 21]:

If the added base is as strong as aniline or pyridine,

such a reaction appears to occur in the case of the

triethylammonium salt of bromophthalein magenta

E or B, despite the low dielectric constants of these

solvents. On adding water or an alcohol reaction

(eq 36) is probably additionally promoted by the

following reaction:

X- + ROH^RO-H+...X- (37)

R.,NH^..X-+base ^ R.,NH+... base + X" (36)

However, for our present discussion the essential

reaction is the one represented by eq 36, which

accounts for the color shift from magenta to blue. (Some weight should probably be ascribed to the

dielectric constant of nitrobenzene or nitroparaffins.)

Ethers (for example, dioxane) and esters of carboxylic acids are not basic enough to bring about reaction (36), judging by the magenta color of solutions.

Secondary ammonium salts of bromophthalein

magenta in benzene show a color behavior which,

though astonishing when first seen, has found a

satisfactory explanation in terms of hydrogen bonding. In referring again to figure 14 it will be noted

that the absorption curve labeled "Et2NH2 salt" (which resembles very closely the curves obtained

for other secondary ammonium salts) is similar to the curve obtained for the BU4N salt. That is, in both

cases the principal absorption band in the visible region is steep, rather narrow, and fairly symmetrical. However, Xmax. for the secondary ammonium

salt is displaced about 30 nm toward the ultraviolet

(that is, to a region near 575 nm). Moreover, the

"wings" of the main absorption band are more

pronounced. (The seemingly lower intensity of

absorption of the quaternary ammonium salt is attributed, at least in part, to difficulties in ascer- taining precisely its molar absorbance because it reacts with glass by ion exchange [14].) The difference in the color of benzene solutions of

quaternary ammonium and secondary ammonium

salts of bromophthalein magenta has been indicated

by describing the former as "blue without purple

tinge," and the latter as "cornflower blue" or "purple-blue" [14]. The range of colors obtained on

adding a secondary amine to bromophthalein

magenta (in benzene) in gradually increasing

amounts until the limiting color has been attained

will be different, obviously, from the sets of colors obtained by mixing the indicator with a tertiary

ammonium or quaternary ammonium salt. The finding that secondary ammonium salts give

cornflower blue solutions in benzene (and in other

hydrocarbon solvents), instead of magenta solutions,

is explained by postulating that two molecules of

salt associate by hydrogen bonding in the manner

indicated in (XLIII). That is, association occurs because the "spare" ammonium hydrogen of each

molecule of salt can become bonded to the quinoid

oxygen (ring 1) of the other molecule [14]. (It seems

probable that the two anions lie in essentially parallel

planes, united by two cations.) The dimeric struc- ture should be stabilized by resonance.

56
R

+ 1 0....H—N—H--.-0

It might be thought that the cornflower blue color

is caused by traces of impurities (for example,

water), rather than by the formation of a dimeric

salt. However, Davis and Schuhmann found [14] that the cornflower blue is obtained only in the absence of impurities. A mere trace of an electron- donor molecule (for example, ethanol) causes a color

shift toward magenta. Subsequent large additions of

alcohol cause a shift from magenta back to blue, the

final blue tone being, however, the "blue without

purple tinge" which is characteristic of quaternary

ammonium salts in all solvents. The shift from cornflower blue toward magenta is satisfactorily explained by postulating that the

"spare" secondary ammonium hydrogen becomes

bonded to the traces of added base (XLIV), this new

R

O....H—N—H----Base R

hydrogen bond being formed at the expense of the

former bond to quinoid oxygen of a second molecule of indicator. That is, the salt is now a monomeric species, consisting— like (XLI) — of hydrogenbonded ion pairs. It is also easily understood why

secondary ammonium salts of bromophthalein

magenta give magenta solutions in dioxane, anisole,

and other ethers (see table 40), in contrast to their

cornflower blue tone in hydrocarbon solvents. The ultimate shift back to blue in a largely alco- holic medium is explained in the same way as for

tertiary ammonium salts — compare (XLII).

To recapitulate our conclusions, the very striking color phenomena just discussed, like the phenomena

observed in studies of conductance, dielectric

polarization, and coUigative properties, indicate

that acid-base behavior in aprotic solvents is significantly affected by two types of hydrogen bonding, namely,

(1) Hydrogen bonding of cations to anions, which

can occur in tertiary, secondary, and primary

ammonium salts; and

(2) Bonding of "spare" cationic protons to elec- tron-donor atoms, which is to be looked for in reactions involving secondary or primary amines.

Obviously, effects on acid-base behavior which

arise from hydrogen bonding of the second type

will be enhanced when a secondary amine is re- placed by a primary amine. Studies of the behavior

of bromophthalein magenta E with primary amines

in benzene have shown that the expected complications exist [14]. Clearly, the simplest overall pat- tern of acid-base behavior will occur in reactions

involving a tertiary amine, making it desirable to select this type of base as a reference compound

for systematic comparisons of acidic strengths in aprotic media.

In the introductory discussion of bromophthalein

magenta (E or B) it was mentioned that this acid

reacts measurably in hydrocarbon solvents with

bases comparable in strength to aliphatic amines.

The most extensive investigations of basic behavior

in benzene made so far have involved tertiary ali- phatic amines or 1,3-diarylguanidines [14, 5, 15, 21,

31]. Both groups of bases yield magenta solutions.

The following bases have also been found (in qualitative tests) to yield magenta solutions: Hexamethylenetetramine [5], most common alkaloids

[14], diphenylformamidine [17], and diphenylacetamidine [17]. In detailed studies minor diff"er- ences in the absorption bands of magenta solutions

may be observed — for example, the main absorption band obtained on adding 1,3-diphenylguanidine

is located a few nanometers nearer the ultraviolet than the corresponding band produced by adding

triethylamine [15].

It has been concluded tentatively that the fol- lowing guanidine derivatives resemble secondary

aliphatic amines in their association with bromophthalein magenta: Monoalkyl-, monoaryl-, 1,1- dialkyl-, and 1,3-dialkylguanidines [17]. The instability of the dimeric structure assigned

to secondary ammonium salts in hydrocarbon sol- vents upon addition of traces of an ether, alcohol, or other type of base suggests that the color behavior

of a secondary amine with bromophthalein magenta

will be modified when ether or hydroxyl groups are present within the molecule. In keeping with ex- pectation, adding a small amount of morpholine to bromophthalein magenta causes the original yellow

color to change to cornflower blue, but on continued

addition of morpholine the cornflower blue color

shifts toward magenta. Similarly, a trace of methylmonoethanolamine produces a blue color shghtly

tinged with red, while large amounts cause a shift toward magenta [14]. Adding a tertiary amine oxide to bromophthalein

magenta in a hydrocarbon solvent produces a still

57
different color tone (rose), as shown in figure 14

(see curve labeled "MesNO salt") [14, 31]. Tri- benzylamine oxide and trimethylamine oxide have

been found to produce almost identical absorption curves [31]. The salts formed in these cases

are considered to have the general formula

R3NOH+...A". The characteristic absorption band

in the visible region is fairly symmetrical and resembles that for tertiary ammonium salts except in being

somewhat flatter and displaced about 20 nm toward

the ultraviolet. It is inferred from this displacement

that the phenolic proton of bromophthalein magenta

remains closer to its original position after bonding

to a tertiary amine oxide than after interaction with

a tertiary amine.

The anion of bromophthalein magenta is exceptionally polarizable, and this accounts for its strongly differentiating behavior with various types

of cations. With acids having less polarizable anions

the differentiating phenomena and their underlying causes are less easily recognized. The examples

given suggest many other potential uses of this indicator for studies of basic behavior.

b. Halogen derivatives ofphenolsulfonephthalein.

The close similarity of the structures of bromophthalein magenta (XXXVII) and bromophenol

blue which has been half-neutralized by a base

(XXXIX) has been pointed out. In the "second-stage

reaction," which involves neutralization of phenolic

OH, bromophenol blue shows a varied behavior

closely similar to that just described for bromophthalein magenta E (or B), and so do related sul- fonephthaleins, for example, bromocresol green, iodophenol blue (the iodine analog of bromophenol

blue), and tetrabromophenol blue (3',5',3",5"-

tetrabromphenoltetrabromosulfonephthalein) [29]. For lack of more suitable indicators the sul- fonephthaleins, though designed for applications

in aqueous solutions, were used in early studies of

acid-base behavior in benzene and other aprotic

solvents. An important example is work by LaMer

and Downes, who described very complex color

phenomena in the following words [357]: "If Brom

Phenol Blue is added to a benzene solution of pure

[carboxylic acid such as dichloroacetic acid,

salicylic acid, or benzoic acid], the solution remains

colorless. As base [piperidine or dimethylamine] is added, a yellow color appears which cannot be dis- charged by addition of acid. It becomes more in- tense the greater the amount of base added, until a reversible color change from yellow to red takes

place, analogously to the color change in water

from yellow to purple. If the acid is very dilute, the

color will change progressively from yellow to a yeUow-purple mixture, to a yellow-red mixture

and finally to red without any yellow present on the

addition of base. On the other hand, if Brom Phenol

Blue is added to a benzene solution of a pure base,

a purple color is observed. If the base is moderately

concentrated (greater than 0.01 M) the addition of

acid changes the purple to a red color and finally

when an excess of acid is added, changes the red

to a yellow. The change from purple to red is an

irreversible change, for no matter how great an

excess of base is added to the red solution, the

purple color wiU not be regenerated. If the base is extremely dilute, the addition of an excess of acid

will change the purple color directly to yellow. . . ." This complex behavior can be understood by ap- plying the observations and conclusions pertaining

to secondary ammonium salts of bromophthalein

magenta in benzene. [The color changes just re- referred to should not be confused with a physical

phenomenon, known as "dichromatism," which

bromophenol blue and certain other sulfonephthalein indicators exhibit under some conditions. For

example, in its alkaUne range in water (pH 4.6 or

greater) bromophenol blue transmits blue light

when in thin films, whereas thicker layers of the same solution transmit light of a red tone. Behavior of this sort is observed when a light-absorbing

material has two or more absorption bands in the

visible which differ greatly in intensity. In such a case the color transmitted also depends on the

quality of the source — for instance, a solution trans- mitting a blue color in a room illuminated by daylight

will appear red under illumination by an artificial source which does not emit blue light [358].] As noted earUer, the first stage of neutralization

of bromophenol blue and related sulfonephthaleins

provides a useful reaction for estimating relative strengths of aniline- or pyridine-type bases in aprotic

solvents — a role that bromophthalein magenta cannot fill. In general, however, the sulfonephthaleins

are less suitable reagents for acid-base studies in aprotic solvents, as the acids and their salts are not very soluble and the first- and second-stage

reactions tend to overlap [29].

c. Nitrophenols. As indicated in our earlier discussion, the compounds that have been most

extensively studied with respect to their con- ductance (sec. 4.3.1) and dielectric polarization

(sec. 4.3.2) in benzene and other aprotic solvents

have been quaternary or incompletely alkylated

ammonium picrates. We have seen that two important conclusions from such studies are: (1) Ion- pair dissociation of imcompletely alkylated ammonium picrates is much less extensive than that of

quaternary ammonium picrates (see table 32); and

(2) the dipole moments of incompletely substituted

ammonium picrates (~ 11-12 Debye units) are lower by several units than the moments of quaternary ammonium picrates (see table 36). Such re- sults led to the deduction that the incompletely

substituted ammonium salts exist largely as hydrogen-bonded ion pairs when in benzene and

like solvents. The comparative electronic absorption spectra of quaternary ammonium picrates and partly substituted ammonium picrates in benzene give additional strength to this conclusion. As representative evidence, absorption curves are given in figure

15 for picric acid and three types of picrates,

all in benzene solution (5 X 10"^ M). It will be seen

that the curves for these picrates, although all

58
WAVE NUMBER, Cm"! X 10-3 35 30 25 20 —

I , , ,

1

. . . ,

I

. , 1 PICRATES IN BENZENE

WAVELENGTH, nm

Figure 15. Comparative spectral absorption curves of picric acid and some of its salts (5 X M) in benzene and in water.

See M. M. Davis Abstrs. Papers 145th Meeting Am. Chem. Soc, New York, N.Y., September 1963, p. 18T. (Ref. No. 359)

having the same basic shape (the envelope of two

overlapping bands), are increasingly shifted

toward the infrared in the order:

EtsNOHPi < EtaNHPi < Et4NPi

[359]. This is the same as the order of spectral

shifts observed for benzene solutions of bromophthalein magenta and several of its salts (fig. 14), and the interpretation is similar. The curves presented in figure 15 are representative for at least a few members of the three

classes of salts indicated. For example, within each

of the following groups the salts (in all cases,

5 X 10"^ M in benzene) were found to have the same absorbance within experimental uncertainty at wavelengths > ~ 300 nm: (1) Et4NPi,

?i-Pr4NPi, and n-Bu4NPi [360]; (2) EtaNHPi, n- BusNHPi, and i-AmsNHPi [361]'. Furthermore, the

absorbance curve determined for 5 X 10~^ M

EtsNOHPi [361] is not measurably different from

that determined for a mixture of HPi (same molarity)

with two equivalents of (PhCH2)3NO [31]. Some other examples of shifted curves for pic- rates in benzene, and also for trinitro-/n-cresol

and two of its salts, were reported earlier [31]. The spectral absorbance of the picrates Usted

above and also that of numerous other picrates was determined in water when solubilities per- mitted [360, 361]. Doing this provided a good test of purity, as all pure picrates yielded the same

absorbance curve in water (see fig. 15) — except, of course, for varying contributions by different

cations. This identity of absorbance in aqueous

solutions obviously occurs because of the essentially

complete ionization of all picrates in water, in- cluding picric acid itself (pKa = 0.33 [362]). Figure

15 provides a very good illustration of the leveling

character of water as a solvent in contrast to the

differentiating character of benzene.

WAVE NUMBER, Cm-I X IQ-^ 35 30 25 20 -T—r 1 . .

1

.

• 1 .

I I . . .

1

WAVELENGTH, nm

Figure 16. Comparative spectral absorption curves of di-o- tolylguanidinium picrate (o-T2GHPi), tribenzylammonium

picrate, and triethylammonium picrate in benzene [20, .367].

Differentiating absorption curves of salts in ben- zene are not limited to the examples given. Additional examples are presented in figure 16. The

three curves were determined carefully using

purified materials (usually 5 X \Q~^M). The curve shown for tribenzylammonium picrate is the cal- culated as well as experimentally determined limiting curve obtained in studying the association of

picric acid with tribenzylamine [20]; the absorbance

of 1,3-di-o-tolylguanidinium picrate was found to be almost identical with that of 1,3-diphenylguanidinium picrate [361]; and, as mentioned earlier,

the absorbance curve for EtaNHPi is practically

identical with the curves obtained for n-BuaNHPi

and i-AmsNHPi. It is apparent, however, that the

electronic absorption of tertiary ammonium pic- rates is not always the same. Comparing the curves

for tribenzylammonium and triethylammonium

picrates one sees that the "primary" (higherfrequency) band, though appearing in the same

(or very nearly the same) region in each case, may

vary in intensity. It is also evident that the lower

absorbance of triethylammonium picrate in the

region of the primary band is compensated by

greater absorbance in the region of the "secondary" band. The secondary band seems strongest

in the case of ditolylguanidinium picrate; its primary

band seems to be shifted a little toward higher

frequencies. (An analogous hypsochromic shift was observed in comparing the absorption curves

of the diphenylguanidinium and triethylammonium

salts of bromophthalein magenta E [15].) Picrates (5 X 10~W) prepared from three

dialkylamines (piperidine, diethylamine, and di-n- propylamine) yielded almost exactly the same

absorption curve in benzene [361]. Although this curve is perceptibly different from the one obtained

for EtsNHPi, etc., the difference is much less

280-613 OL-68—

59
striking than for secondary and tertiary ammonium

salts of bromophthalein magenta at the same concentration, consisting merely of a slightly reduced

intensity of absorption in the region ~ 335 to ~ 405 nm. It will be recalled that Maryott postulated an

appreciable tendency for secondary ammonium

picrates to dimerize in benzene, thus accounting

for a pronounced falling off of the dielectric polarization with increasing concentrations (see fig. 10

and related discussion).

Von Halban and Szigeti noted additional ex- amples of differentiating spectral behavior in a survey of some dioxane, chloroform, and ethylene

chloride solutions of picrates and naphthopicrates

[363]. The two "partial" bands of Pi^ were sharply

separated, as in benzene. The difference between

quaternary ammonium salts and incompletely substituted ammonium salts was particularly note- worthy.

Analogous differentiating effects have been ob- served for 2,4-dinitrophenolates. For example, in determining the equilibrium constants for 1:1 association of 2,4-dinitrophenol with mono-, di-, and

tri-n-butylamine in chlorobenzene, Bayles and

Chetwyn obtained three different limiting absorption curves [364]. Figure 17 presents absorption

e X 10-3

2,4-Dinilrophenol (DNP)

and Some of Its SoltS:

n Benzene

In Water

n-Bu4N Salt

No Solt

Figure 17. Comparative spectral absorption curves of 2,4-di- nitrophenol and some of its salts in benzene and water [77].

curves for 2,4-dinitrophenol and its EtaNH and

n-Bu4N salts in benzene, and also for the phenol

and its sodium salt in water [17]. As with picrates,

two overlapping bands are clearly detectable in the curves for benzene solutions. The intermediate

location of the curve for the EtsNH salt supports

the inclusion of this compound in the category of

hydrogen-bonded ion pairs.

A recent review [365] has divided progress in electronic absorption spectroscopy into two main

periods. During the first period, which extended

into the mid-1940's, absorbance data were inter- preted according to classical vibration theory, as exemplified by Lewis and Calvin's "theory of elec- tronic oscillations" [366, 367]. According to this theory a planar organic molecule having a conjugated system which extends in two directions will

have two optical axes, and when the molecule

absorbs light there will be electronic oscillations along both axes, giving rise to absorption bands of

different frequencies. The one at lowest frequencies,

caUed the x-band, is associated with the most ex- tended conjugated system. The second band is referred to as the y-band. The resonance hypothesis

has provided a good explanation why ortho-mivophenol (XLV) and para-nitrophenol (XLVII) both

have a rather strong electronic absorption band, es- pecially when in the anionic state [368], and it is reasonable to identify these bands with the x- and

y-bands, respectively, of Lewis and Calvin. The

principal absorption band of 2,6-dinitrophenol

(XLVI), like that of o-nitrophenol, should be an

x-band. Only one resonance structure (XLVIII)

is given here for 2, 4-dinitrophenol, but obviously

this acid and its salts should have an x-band as well as a y-band; the same holds for picric acid

(XLIX) [359]. OH

l=N02

(XLV) OH

II

OH

II 02N=|'^\-N02

(XLVI)

OH—NO2 NO2

(XLVII)

O2NOH

II i^^\-N02

(XLVIII)

A02

(XLIX)

During the second main period in the development of electronic absorption spectroscopy, which

has seen an extensive application of quantum mechanics in the interpretation of spectra, the view

has been adopted that the strongest absorption

bands in the spectra of benzene derivatives are displaced bands of benzene itself [365, 369-371].

The approximate locations and intensities of the

benzene bands are as follows: (1) A very strong

band near 180 nm (e ~ 10^); (2) a less intense band

near 200 nm (e ~ lO^ to lO'*); (3) a still weaker

band near 255 nm (e ~ 300). Various names or code designations have been advocated for the

three bands; for this discussion it is convenient to adopt the adjectives ""^second primary,'"' ""primary,"

and "secondary," which were appUed to the ben- zene bands in the order Usted above [369]. We

shall be concerned only with the primary and sec- ondary bands, as the displaced "second primary"

band does not appear in the spectral range under

discussion.

Relationships between the primary and sec- ondary bands of benzene and the displaced primary

and secondary bands appearing in substituted

benzenes were especially well demonstrated in

60
Doub and Vandenbelt's measurements of the

spectra of an extensive series of mono- and polysubstituted derivatives, dissolved in water [369-

371]. An observation which is especially pertinent

to this discussion is their finding that the absorption curve of )8-resorcylic (2,4-dihydroxybenzoic)

acid approximates a summation of the curves for

salicylic (o-hydroxybenzoic) and p-hydroxybenzoic acids: p-Hydroxybenzoic acid has a strong,

symmetrical band at X 255 (e 14,000); salicylic

acid has a weaker band at k 303 (e 3,600); and

corresponding bands appear for /3-resorcylic acid at k 256 (e 13,600) and k 295 (e 5,500). The structural

similarity of )8-resorcylic acid and 2,4-dinitrophenol is quickly recognized; however, the two

like substituents in the former compound are electron-donor groups and the third substituent is an electron-withdrawing group, while the reverse is true for 2,4-dinitrophenol. It is evident that Doub

and Vandenbelt's primary and secondary bands

correspond, respectively, to Lewis and Calvin's

y- and x-bands.

Although the primary and secondary bands of

picrates and 2,4-dinitrophenolates overlap strongly

when water is the solvent (see figs. 15 and 17),

it can be recognized that their shapes and relative

intensities are about the same as for benzene

solutions. (The resolution of strongly overlapping

band systems is aided by referring to figures which

Vandenbelt and Henrich constructed by making

summations of two symmetrical absorption bands;

the widths and relative intensities of the bands were

varied, as well as the separation of band maxima

[372, 371].) The pronounced overlapping observed

for aqueous solutions seems to result from a "red-shift" of the primary band, and this shift can

be attributed to solvation of the nitro group, since

hydrogen-donor solvents, especially water, are known to cause marked bathochromic shifts in the

electronic absorption of para-substituted nitro- benzenes [371]. Another example of a strong

bathochromic effect of solvation of a para-quinoid

isomer can be seen in comparing the positions

and intensities of absorption bands of o- and p- nitrophenols as given in a tabulation by Wheland [368]. The values obtained, (a) in 0.1 -A'^ aqueous perchloric acid and (b) as the anions in water,

are as follows: (a) o-Nitrophenol, k max 350 nm,

e max 2,500; p-nitrophenol, k max 320 nm, e max

9,000. (b) o-Nitrophenolate, A max 420 nm, e max

5,000; p-nitrophenolate, k max 400 nm, e max

15,000. It is evident that the band locations are not

greatly different for the ortho- and para-isomers,

and again presumably because hydration of the

latter has a strong bathochromic effect. The intensity of absorption seems always to be

substantially greater for primary than for sec- ondary bands.

Comparisons have been made of the absorption

curves of tetra-n-butylammonium and triethylammonium mono- and polynitrophenolates when

dissolved in benzene [359, 17]. In this solvent the

primary and secondary bands are always much more

widely separated than in water. This is illustrated -very strikingly in figure 18, which compares the

WAVE NUMBER, cm"' X

10"''

€ X 10-3 20

15

10

5

1 ——. . ,

TRIETHYLT 1 . 1 ' ' 1

AMMONIUM SALTS

OF NITROPHENOLS

IN BENZENE /-vHPl

p-NP,--/

f / / / / /

* //

- y

1

V \ 2,6-DNP

* \ •

300 350 400 450 500 WAVELENGTH, nm

Figure 18. Comparative spectral absorption curves of tri- ethylammonium salts of p-nitrophenol, 2,6-dinitrophenol,

and picric acid in benzene [359].

absorption curves of the EtaNH salts of p-nitrophenol, 2,6-dinitrophenol, and picric acid in benzene. (Note: The curve for EtsNHPi was deter- mined with a 5 X 10~^ M solution; the other two

absorption curves are limiting curves, obtained on

adding EtsN in excess to lO""* M 2,6-dinitrophenol

or p-nitrophenol.)

The absorption spectra of mono- and polynitroanilines, determined both in benzene [17] and in methanol [373], show the expected parallelism

with the spectra of corresponding nitrophenols. In molecules possessing primary and secondary

bands the primary band has sometimes been found

to gain in intensity at the expense of the secondary

band, or vice versa. For instance, Kamlet and as- sociates recently discovered, while studying effects of A-alkyl and A,A-dialkyl substitutions on the ultra- violet absorption of 2,4-dinitroaniline in methanol,

that bulky substitutents on the amino nitrogen may

lead to inhibition of (+R2N=Ci—Ca^NOj) reso- nance and enhancement of (+R2N=Ci —C4=N02)

resonance [373]. It is conceivable that steric in- fluences are at least partly responsible for the

band intensity shifts indicated in figure 16. This brief review has shown the exceptional

sensitivity of electronic absorption spectroscopy

as a tool for studying the properties of acidic in- dicators like bromophenol blue, bromophthalein

magenta, 2,4-dinitrophenol, and picric acid in aprotic solvents. Obviously, considerably more

labor is required than in studying acid-base equilibria in a leveling solvent, but far more information

is attained.

d. a-Naphthol, ft-naphthol, and phenol [353]. Nagakura and Gouterman have evaluated acid-base

association constants from shifts in the near-ultraviolet absorption bands of phenol and a- and )8- naphthols caused by adding triethylamine, ethyl

ether, or nitromethane. It is well known that except

61
for EtsN these are all weak proton donor or ac- ceptor compounds. The spectral shifts were as- cribed to hydrogen bond formation for reasons such

as the following: (1) The absorption of each phenolic

acid (HA) shifts to longer wavelengths when a proton acceptor (B) is added, and curves for solu- tions containing varying proportions of base and

acid intersect at one or more sharply defined

isosbestic points. (2) Satisfactory equilibrium constants were calculated on the assumption of 1:1 acid-base association.

The calculated association constants (Xbha) ^re

listed in table 41. Judging relative reactivities by

Table 41. Comparative equilibrium constants (Kbha) far some

acid-base associations in n-heptane (25°) as evaluated from

near-ultraviolet absorbance data ^

B (proton

Kbha, where HA is:

acceptor) a-Naphthol )3-Naphthol Phenol

EtsN 121 103 83.8 Et20 15.7 14.5 10.8 MeN02 3.3 2.6

''See S. Nagakura and M. Gouterman, J. Chem. Phys. 26,

881 (1957), Table IV. A^bha is the equilibrium constant corre- sponding to B -I- HA «=* BHA. Approximate molar concentrations were: Cha, (2.2 to 4.3) X 10-»; Cb, 0.01 to 0.5.

Reference:

[353] S. Nagakura and M. Gouterman, J. Chem. Phys. 26,

881 (1957).

the magnitudes of ^bha values, it is clear that a-naphthol is the strongest of the three acids and

phenol, the weakest. It is also clear that, as would

be expected, triethylamine and nitromethane are, respectively, the strongest and weakest of the proton

acceptors.

In infrared spectroscopy the relative strengths of

hydrogen bonds have often been judged from the

comparative shifts of 0-H stretching frequencies on hydrogen bond formation. It is noteworthy that corresponding shifts of ultraviolet absorption bands

of a- and ^-naphthol are different for the La (primary) and Lb (secondary) bands: when the primary

bands are being compared the shift is greater for a-naphthol than for /3-naphthol, but when sec- ondary bands are being compared, the reverse is true [353].

4.3.5. Spectral Absorbance in the Infrared The extensive use made of infrared spectroscopy

during the past quarter century for studying hydrogen bond formation has been illustrated nu- merous times in preceding discussion. Only a small fraction of the IR studies have been con- cerned with hydrogen bonding between cations

and anions. Since 1950 there have been a few in- vestigations of this type of bonding in dry aprotic

solvents, and since 1960, additional investigations

have been made using wet aprotic solvents. The

second group of studies are a result of the widespread interest in processes for extracting acids

from aqueous solutions into amine-containing

aprotic solvents. A third group of investigations has

dealt with hydrogen bonding in crystalline salts. The association of cation-anion pairs by hydrogen

bonds seems to be closely similar in crystals and

in aprotic solvents, and results from all three

types of investigation wiU be summarized.

It is well known that the formation of a hydrogen

bond X—H...Y is very commonly recognized

by the following changes in the X-H stretching

band: a shift to lower wave numbers; an increase

in intensity: a considerable broadening; and, in many cases, a splitting to two or more component

bands. Hydrogen bonding in amine salts, N^-H...X^

is often inferred from changes involving the

N-H stretching band [374-376], but other infrared

absorption bands also contribute supplemental or substitute information.

a. Measurements of crystals. It is appropriate to include mention of some deductions made by studying compilations of x-ray diffraction data. Early in 1951, in a symposium on "Hydrogen Bonds,"

Donohue reported ". . . the rather remarkable

fact that only in very exceptional cases does a hydrogen atom bonded to nitrogen or oxygen occupy

a position such that hydrogen bond formation is impossible. Thus in the case of the hydrochlorides

. . . all three hydrogen atoms of the —NH3+ groups

form ^-H...C1 bonds . . ." [377]. In a subsequent study of more extensive x-ray diffraction data, FuUer reached complete agreement with

Donohue's main conclusions, stating that "almost

without exception, molecules in a crystal arrange

themselves so that the maximum number of hydrogen bonds is formed . . ." [378].

The infrared measurements of crystals were

usually made with mulls, a liquid hydrocarbon and/

or a halogenated hydrocarbon being employed as the suspending medium. In some cases a film of

solid was prepared by sublimation or by evaporation of solvent. Early studies of salts were made by

Lord and Merrifield [374], who reported that the

IR spectra of tertiary amine hydrohalides indicate

the presence of strong N+-H...X~ bonds. For

example, they observed that the spectrum of triethylamine hydrochloride shows no band in the

usual N+-H stretching region; instead, there is a very strong band (a well-defined doublet with two

weaker satellites) centered at 2540 cm"^ The same

band appears with slight modifications in the spec- trum of a chloroform solution. In the view of these

authors the characteristics of the displaced bands

"are incompatible with a simple electrostatic model

of the hydrogen bond . . . (instead) a quantummechanical model is undoubtedly required."

Waldron observed independently [379], in a detailed study of M-eNH:,+Cl-, that the N-H stretch- ing frequencies are lower in the cation than in the

neutral molecule, and he concluded that "part of

62
the reduction in frequency is undoubtedly due to hydrogen bonding to the anion."

The sohd hydrohalides of 19 different amines

were studied in detail by Sandorfy and associates,

and in many cases measurements were also made

with solutions in water [375, 376]. The bases rep- resented were primary, secondary, and tertiary

amines, both aliphatic and aromatic, and also pyridine; HCl, HBr, and HI were the acids repre- sented. Omitting all details, the main observations

and conclusions may be summarized as follows:

(1) Hydrogen bond formation was detected in all of the hydrohalides. Its existence was shown by a displaced NH+ stretching band and by the other

spectral changes considered characteristic of hydrogen bond formation (that is, splitting, increased

intensity and breadth).

(2) In the solid state, the hydrogen bonds are of

the type N+-H...X~, whereas in aqueous solutions,

they are of the N+-H...0 type.

(3) In aliphatic amine hydrohalides the hydrogen

bond is very largely electrostatic in character.

Delocalization of charge is more extensive in aromatic amine hydrohalides.

(4) Tertiary amines form stronger N+-H...X~

bonds than do primary and secondary amines. This

is consistent with the greater concentration of positive charge in tertiary ammonium ions.

(5) A tertiary ammonium ion forms a stronger

hydrogen bond with CI" than with Br" or I". Marzocchi, Fryer, and Bambagiotti [380] studied

the N-H stretching band in the salts formed by

1:1 association of hexamethylenetetramine with

HCl, HBr, HI, and HCIO4, and attributed the struc- ture of the f (N-H) bands of the hydrohalides to N+-H...X~ bonding. They concluded that there

is little, if any, hydrogen bonding in the perchlorate.

They believe that a quantum-mechanical treatment

of the hydrogen bond is required to account for hydrogen bonding in salts like those studied, but

consider that no theory so far proposed accounts

adequately for the spectral changes observed. As

in the IR measurements by Sandorfy and associates

the comparative spectral shifts indicated that the

strengths of cation-to-anion hydrogen bonds de- crease in the order CI" > Br~ > I". In analogous investigations of Lipovskii and

Nikitina a shifted t'(N-H) band of crystalline

trialkylammonium fluorides was attributed to the

structure R3N+-H...F" (R = Et or C10H21). As

expected, the hydrogen bond to F" was found to be

even stronger than the bond to CI" [381]. Hydrogen bonding in crystalline ammonium

halides has received careful study. An example is an investigation of NH4F and ND4F (25 and - 195 °C)

in which the expected evidence of strong N+-H...F"

bonds was found. The authors emphasized that

the stretching N-H band of NH4F is not broad,

showing that band broadening is not an invariable

consequence of strong hydrogen bonding [382]. {Neutron diffraction data for ND4Br (- 195 to 200 °C)

should also be noted: they indicated that the formation of N+-H...X" is an important stabilizing

influence in ammonium halides [383).)

Table 42 is a compendium of conclusions from

Table 42. Occurrence of hydrogen bonds N+-H...X" in crystalline ammonium and substituted ammonium salts as deduced by IR spectroscopy "

N+ —

IN n

..A Present N+-H...X- Weak or Absent

Anion ixeierenccs Anion References

Br" [388] CIO4- [386, 388] Ch [388] BFj [384-386, 388]

I- [388] BClj [388] HFj [386, 388] BPhf [386, 388] N3 [386] PFe [386, 388] NCO- [386] BeF4= [385]

NCS- [386] SiFf [385] NO2- [386] TiF,T [385] NO3- [386. 388] [385]

" S04= [389] PbCl,v [385, 388]

' S03F- [386] SnCliT [385]

^ CH3S05 [386, 388] PtCl= [385] PtBrg= [385]

[385] CrFe^ [385]

" At room temperature. The cation was NH+, MeNH+, MeaNHj,

Me3NH+, Me3ND+, Et3NH+, Et3ND+, PyH+, or PyD+. Additional references to hydrogen bonding in haUdes are given in the ac- companying text. Similar results were obtained with (CioH2i)3NH+ salts. The

samples were viscous liquids. A weaker hydrogen bond is thought

to form to an oxygen of HSOj. Hydrogen bonding to SO4 causes

a decrease in the symmetry of this ion [389]. This salt may resemble NH4I, where the cation is thought to form one hydrogen bond to I" and to be free to rotate about this bond [386].

References:

[385] B. Cox and A. G. Sharpe, J. Chem. Soc, p. 1798 (1954). [386] T. C. Waddington, J. Chem. Soc, p. 4340 (1958). [388] R. H. Nuttall, D. W. A. Sharp, and T. C. Waddington, J. Chem. Soc, p. 4965 (1960). [389] A. A. Lipovskii and M. G. Kuzina, Russ. J. Inorg. Chem.

(English Transl.) 10, 740 (1965). [384] G. L. Cote and H. W. Thompson, Proc Roy. Soc. (London)

A210, 217 (1951).

some additional investigations. In drawing general

conclusions from these and the other studies just

reviewed it is necessary to bear in mind that the

data and interpretations of diff^erent investigators are sometimes in partial disagreement. (For ex- ample, Mathieu and Poulet [387] believe that the

free stretching N-H frequency of NH4+ should be

assigned a value higher than values used by Waddington [386] and others.) However, the combined

results indicate complete agreement that strong

N+-H...X" bonds exist in sohd ammonium and

63
incompletely substituted ammonium halides.

The polyatomic anions also listed in column I are

of types considered from other kinds of physical

evidence to be highly polarizable, so that their in- clusion in this group is not in conflict with expectation. Neither is it surprising for IR spectroscopic

measurements to indicate that hydrogen bonds fail to form between NH^ or related cations and the

complex anions listed in column 3 — or, at most, to be very weak. Obviously, in such anions the

negative charge on a single oxygen or halogen will be much smaller than in the anions of column 1, and they will behave as weaker bases.

However, in some cases hydrogen bonding to complex anions is believed to occur. For example,

Lipovskii and Kuzina [389] consider that Et3NH+

and (CioH2i)3NH+ form hydrogen bridges to the

dioxotrisulfatouranate(VI) ion, [U02(S04)3]~^- In

this anion bidentate coordination of each SO4

to uranium is postulated. The evidence for hydrogen bonding is a decrease in the symmetry of

SOf as well as a displaced i'(N—H) frequency.

From similar evidence they concluded that SO^

forms strong bonds to tertiary ammonium ions. Moreover, as indicated in table 42, Waddington and associates inferred that two related anions,

SOsF^ and CH3SO3, can also become hydrogen

bonded to cations. In contrast with their findings

for SOf, Lipovskii and Kuzina observed no dis- tortion of CIO4 . Infrared data of Hadzi and associates indicate

strong hydrogen bonding in adducts of pyridine

oxide and 2-picoline oxide with HCl and HBr and

also in the adduct of trimethylamine oxide with

p-nitrophenol [390, 391]. (Compare figs. 14 and 15 and related discussion.)

b. Results using dry aprotic solvents. The principal infrared evidence for the existence of hydrogenbonded ion pairs in dry aprotic solvents has come

from investigations by Barrow and associates

[392-395]. In a series of experiments resembling

photometric titrations. Barrow and Yerger studied

spectral changes accompanying the gradual addition of triethylamine [392], diethylamine [393], or «-butylamine [394] to carbon tetrachloride or chloroform containing varying fixed concentrations of acetic acid. The infrared absorbance data were used to ascertain the compositions and probable structures of acid-base adducts formed at various relative and total concentrations of acid

and base, and also to determine the relative stabilities of different products as measured by equilibrium constants. In conjunction with the experiments they also determined equilibrium constants

for the association of EtsN with CDCI3 {118] and

for the dimerization of acetic acid [265] in the same

two solvents. Subsequently, Barrow investigated

the reaction of pyridine with some aliphatic car- boxylic acids in chloroform [395] and Bell and Bar- row studied the association of dimethyLformamide,

pyridine, and triethylamine with phenol-rf, pnitrophenol, and p-nitrophenol-c? in this solvent

[396]. Some of the experiments and their principal conclusions will be summarized briefly.

Reactions between acetic acid and triethylamine

in carbon tetrachloride. In separate experiments

triethylamine was added gradually to acetic acid

of five different initial concentrations in the range

0.001 M to 1.0 M. Absorbance measurements were made at several wavelengths, each wavelength

being near the peak of an absorption band. The

principal bands measured were ones attributed to carboxylate, carbonyl in dimeric acetic acid and

also in the monomeric acid, and OH in the dimer.

As expected, adding EtsN caused some bands to diminish in intensity and others to become stronger.

Some important conclusions from this investigation were:

(1) Half-neutrahzing 0.1-M AcOH in CCI4 with

EtsN produces an acid salt, Et3N(HOAc)2, which

has considerable stability. The assigned structure

(L) is shown below.

-1/2,0...H—

\ CH,— C ,C—CH, ^O...H 0

(L)

In (L) the anion is shown as having carboxylate

character. This term does not imply complete ioniza- tion of the salt, but merely that the anionic charge

is about equally distributed on the two oxygens,

since one oxygen is hydrogen bonded to the cation

and the other is hydrogen bonded to an extra molecule of acetic acid.

(2) Continued addition of triethylamine causes

gradual conversion of the acid salt to the 1 : 1 adduct,

EtaNHOAc (LI).

CH,—

3 \ - +

0 ...H NEtj

(LI)

It is important to note that in (LI) the attachment of the cation to the anion does not involve merely

general electrostatic forces. Instead, the cation is considered to form a specific attachment to one of the anionic oxygens, "undoubtedly by a hydrogen

bond." The negative charge resides essentially on

the bonded oxygen instead of being shared by both

oxygens. The analogy between (LI) and structures previously presented (XLI, XLIV) for certain hydrogen bonded salts of bromophthalein magenta is easily recognized.

(3) The acid salt (L) is not formed in measurable

amounts when Et3N is added to acetic acid solu- tions as dilute as 0.001 M. Instead, adding Et3N

appears to lead directly to the production of the

1 : 1 salt (LI).

64
Reactions between acetic acid and triethylamine

in chloroform. In experiments closely paralleling

those performed in carbon tetrachloride the results

for chloroform solutions were as follows:

(1) As before, the first product obtained on adding

triethylamine to 0.3-A/ acetic acid was an acid

salt (L).

(2) Continuing the addition of EtsN led gradually

to the formation of a 1:1 salt. However, its IR spec- trum differs in some respects from that of (LI), and it was concluded that the salt can be solvated

by CHCI3, as indicated in (LII). This interpretation

-1/2 ^O...HCCLj

cH,—c;\0...H NEt. -1/2 ^

(LII)

is supported by the observation that on adding a

very large excess of EtsN the spectrum shifts toward

that of (LI); such a shift is predictable from the

ability of EtaN to associate with HCCI.3 (see section

3.4.5a) in competition with the C=0 group. An

additional conceivable structure is one resembling

(LII) except in having HCCI3 hydrogen-bonded to both oxygens. The over-all reaction leading to its formation is indicated by eq 38: CH3COOH + 2Et3N + 2HCCI3 ^ CH3C(O...HCCl3)2-+ (Et3N)2H+. (38)

However, (LII) is considered a more likely product.

Log K values for reactions involving acetic acid

and triethylamine. Table 43 contains comparative

equilibrium constants (expressed as the logarithms)

for reactions involving acetic acid and aliphatic

Table 43. Values of log K (25-30°, molar units) for some acid-base reactions, from infrared absorbance data ^' '''

Stoichiometric Concns. Reacting Species Product Solvent Log/^

HOAc Base (1) (2)

Base == Et3N [392]

0.001

0.10, 0.30

0.001

0.3

0.001-0.025

0.05-0.50 0.0005-0.005

0.15-0.50

HA

B(HA)2

HA

B(HA)2

B

B

B

B

BHA

2BHA

BHA

2BHA

CCI4

CCl,

CHCI3

CHCI3

2.90

0.20

3.48

1.34

Base = EtzNH [393]

0.0008

0.0008-0.01

0.10, 0.30 0.0008

0.10, 0.30

0.0002-0.004

0.0016-0.020

0.075-0.40 0.0004-0.004

0.10-0.60

HA

B(HA)2

HA

B(HA)2

B 2BHA

BB

B

BHA

(BHA)2

(BHA).

BHA

2BHA

CCI4

CCl, CCI4 CHCI3

CHCI3

3.45 2.38

1.20 3.48

2.00

Base = n -BuNHs [394]

0.0008

0.0008

0.0008-0.008 0.0008-0.0064 HA

HA

B

B.HCCl.,

BHA

BHA.HCCI3

CCI4 CHCI32.78

3.45

" From average values of K given in tables of Barrow and Yerger [392-394].

" HA signifies acetic acid; B signifies EtsN, Et2NH, or n-BuNH2.

' See accompanying discussion.

References:

[392] G. M. Barrow and E. A. Yerger, J. Am. Chem. Soc. 76,5211 (1954). [393] E. A. Yerger and G. M. Barrow, J. Am. Chem. Soc. 77,4474 (1955). [394] E. A. Yerger and G. M. Barrow, J. Am. Chem. Soc. 77, 6206 (1955).

65
amines in carbon tetrachloride and chloroform.

The first section of this table contains data for reactions where EtsN was a participant. In the

calculations of equiUbrium constants allowance was made for the estimated extent of two competing reactions, namely, the self-association of acetic

acid [265] and the association of EtsN with CDCI3

(which was assumed to behave Uke CHCI3) [118].

The values of log /^i,2 adopted for acetic acid in carbon tetrachloride and chloroform, respectively, were 3.30 and 2.40; log K for 1:1 association of

EtgN with CHCI3 was assumed to be -0.44 [392].

Assuming all of the equilibrium constants to be

substantially correct, one may deduce the following:

(1) In CCI4 (at room temperature) AcOH has a greater tendency to dimerize than to associate in

1:1 ratio with EtsN.

(2) In CHCI3, the reverse is true. A twofold explanation which can be offered for the solvent effect

is that (a) acetic acid is less completely dimerized

in CHCI3 than in CCI4 (see section 4.2. 2d) and (b) solvation of the 1:1 salt by CHCI3 as in (LII) has a stabihzing effect. (Although the 1:1 association of

CHCI3 with EtsN is a competing reaction, its extent

is negligibly small.)

(3) The conversion of Et3N(HOAc)2 to EtsNHOAc

by adding an excess of EtsN occurs much more

extensively in CHCls than in CCI4. As in (2), this can be attributed largely to solvation of the 1:1 salt by CHCI3: see (LII) and eq 39.

CHsC(0. . .H0C0CH3)0-. . .H+NEts

+ Et3N + 2CHCl3 ^ 2CHsC(O...HCCl3)0-.. .H+NEts (39)

Reactions in carbon tetrachloride and chloroform

involving acetic acid and diethylamine. Products

formed and log K values. Yerger and Barrow's

experiments with diethylamine were similar to the ones using triethylamine. Some parallelism of

results was expected, of course, but the ability of

the diethylammonium ion to form two bonds to proton acceptors was a complicating factor. In

brief, the conclusions and explanations were as

follows:

(1) Just as with EtsN, the half-neutralization of

relatively concentrated (0.1 M) acetic acid in either

solvent yields an acid salt (LIII). (However, the

are to the same oxygen. The formation of (LIII) is thought to occur by the addition of Et2NH to dimeric

acetic acid (eq 40). (CHsCOOH)2 + Et2NH ^ (CH3COOH)2HNEt2. (40)

Log K for this reaction (CCI4) was estimated to be

approximately 4.18.

(2) In CCI4 the complete neutralization of acetic

acid is beUeved to yield two products — a monomeric

NBt, / \

-1/2^0...H H... 0^-1/2 CH.-C '9-TMi - -^y.

NEt„ +

CH,—C'

' V..H

••• / H —NEt-

(LIV) (LV)

salt (LIV), which is formed in dilute solutions, and

a dimeric salt (LV), which predominates in concentrated solutions. Log K for the dimerization (eq 41) was estimated to be 2.38. 2CH3COOH2NEt2 ^ (CHsCOOH2NEt2)2. (41) The dimeric salt was not detected in chloroform

solutions, but only a monomeric salt thought to be

solvated by chloroform (LVI), an analog of (LII).

-1/2^p...HCClj CH3-C:

-1/2 0,..H

••• / H NBt_

(LVI) Structure (LIV) is probably not completely planar.

It is more likely that the hydrogen bonds from

Et2NHj to O" are directed toward the nonbonded.

orbitals, thus permitting maximum interaction.

Similar considerations apply to the other diethylammonium salts. Two additional structures — see (LVII) and

(LVIII) — were considered as possible alternatives

to (LIV) [393], but were rejected for reasons that

will not be detailed here.

-1/2

-1/2^0...H 0

H NBt„

(LIII)

acid salt was not formed in detectable amounts in 0.0008 M acetic acid solution.) Structure (LIII) is the same as (L) except in having two hydrogen

bonds from the cation to the anion; the two bonds

-1/2

CH-—

' V.

-1/2

NBt„

(LVII)

CH,—

3 \

0

0...H—N—Bt.

I 2

H

(LVIII)

(3) The equiUbrium constants for reactions in- volving Et2NH are either about the same as for the

corresponding reactions involving EtsN or some- what larger. For example, log K for the 1:1 association of B with HA in chloroform has the same value

66
(3.48) when B is Et2NH as when B is Et.-jN. A reasonable explanation for the similar association constants is that in both EtiNHaOCOCHa and

EtsNHOCOCHs the C=0 group is hydrogen bonded

to CHCI3, as shown in (LII) and (LVI); such solvation would help to stabihze the carboxylate structure. A corresponding stabilization results from the

hydrogen bonding of excess HA to BHA, giving

B(HA)2-see (L) and (LIII). The dimerization of diethylammonium acetate,

yielding (LV), is obviously not possible for a salt with only one proton attached to the ammonium

nitrogen. The three examples of hydrogen bonding

just mentioned all bring about increased association of B with HA.

An explanation along similar lines was proposed

by Davis and Paabo [15] to account for the following

observations: In benzene solution, benzoic acid is almost as strong as bromophthalein magenta E when

the criterion of strength is the value of log K (25°,

molar units) for 1:1 association with Ph2G (1,3- diphenylguanidine), the values of log K for benzoic

acid and bromophthalein magenta being 5.32 and

5.41, respectively. However, when EtsN is used as

the reference base instead of Ph2G, the corresponding values of log K are 3.58 and 4.44, signifying that benzoic acid is an appreciably weaker acid than

bromophthalein magenta. The structure postulated

for triethylamrponium benzoate is (LIX) — a close

analog of Barrow and Yerger's proposed structure

for triethylammonium acetate (LI). Davis and

Paabo suggested that diphenylguanidium benzoate

can assume a cyclic, more stable structure (LX)

by dual hydrogen bonding of the cation to the anion.

Bt

IPh

(LIX)

H H

\ / Ph-N N-Ph \/

II

H H

IPh

(One of the resonating

structures)

(LX)

The cycHc structure postulated for diphenylguanidinium benzoate is unusual (perhaps almost unique),

being constructed by hydrogen bonding of a reso- nant cation to a resonant anion. Phenolates of tri- ethylamine and diphenylguanidine (for example,

their salts with bromophthalein magenta) would be

less dissimilar than the two carboxylates shown

above [15]. Reactions in carbon tetrachloride and chloroform

involving acetic acid and n-butylamine. Products

postulated and log K values. «-Butylamine, like diethylamine and triethylamine, is convertible into an acid salt (LXI) in concentrated CCI4 and CHCI3

solutions. The 1:1 salts, which are close analogs of

diethylammonium salts, are the monomeric acetate

(LXII) and its dimer (LXIII), formed in CCI4, and

the solvated monomeric acetate (LXIV) which is formed in CHCI3. According to table 43 the 1:1 monomer, although less stable in CCU than the

corresponding diethylammonium and triethylammonium acetates, is equally stable in CHCI3, thus

providing additional support for the postulate that the stability of alkylammonium carboxylates is enhanced by hydrogen bonding to proton-donor

solvents.

-V2, CH-— c;

:o . . .H-

-1/2^

\

• H-N 0 Bu

(LXI)

0 . : .H-N—Bu

(LXII)

-1/2° •i:";''-^^

(LXIV)

-<..

N—H.. Bu

(LXIII)

In summing up their overall conclusions from

these sets of studies Yerger and Barrow stated that

in all systems containing acetic acid and a primary,

secondary, or tertiary aliphatic amine a hydrogen

bond forms between the alkylammonium ion and one oxygen atom of the acetate ion. Multiple hydrogen

bonding to a single oxygen atom occurs when more

that one cationic proton is available for bonding,

and appears to lead to the formation of bridged

dimers in relatively concentrated solutions of

primary and secondary ammonium acetates [394].

We have already seen (sec. 4.3.2b) that dielectric polarization data for related acid-base systems in benzene have received similar interpretations.

The Barrow and Yerger studies have also illus- trated how the apparent strengths of bases, as measured by ability to form 1:1 adducts with acids,

may be altered by a change of solvent [393]. According to table 43, by the criterion of log K values for

1:1 adduct formation the three amines vary in strength in the following order, when dissolved in CCU:

Et2NH > EtsN > n-BuNH2.

By the same criterion the strengths of EtsN and

n-BuNH2 become equal to that of Et2NH in CHCI3.

67
It seems reasonable to attribute this "leveling

effect" in large measure to the ability of CHCI3 to form hydrogen bonds to carboxylate atoms of alkylammonium acetates, thereby enhancing the stability

of these salts. Other studies of tertiary amine-carboxylic acid

interactions in aprotic solvents. Barrow [397] and

Cook [398] have studied hydrogen bonding inter- actions of amino acids, with special interest in factors that promote zwitterion formation. As ex- pected, the results indicate the existence of hydrogen bonding interactions analogous to those occurring in the EtsN-AcOH system. However, the

comparative extent of intra- and intermolecular

bonding is determined in part by the molecular

geometry.

As a part of his investigation Cook examined the

EtaN-AcOH system (0.1 M) in five aprotic solvents: Benzene, carbon tetrachloride, chloroform, 1,2- dichloroethane, and dichloromethane. He also studied the reaction in a dipolar aprotic solvent

(acetonitrile) and in the absence of a diluent. In agreement with Barrow and Yerger [392] he deduced

that a stable 2:1 acid-base complex forms in aU

benzene solutions containing an excess of the acid

and that the continued addition of EtsN leads to the

replacement of the 2:1 complex by a 1:1 complex.

He reached the same conclusions using the other

five solvents, as well as in the absence of a solvent. However, Cook is in disagreement with Barrow

and Yerger as to the nature of the 1:1 complex. As we have noted. Barrow and Yerger postulated struc- ture (LI) for the 1:1 complex; Barrow [395] calls

this a "proton-transferred ion-pair species." In Cook's view the ion pair species (LI) is in tautomeric

equilibrium (eq. 42) with an isomeric hydrogen

bonded complex (LXV) in an inert solvent Uke CCLj:

Et2N...H0C(=0)CH^ Et^NHt.TOC(=0)CHj (42)

(LXV) (LI)

Like Barrow and Yerger, he considers that the

proton-transferred species becomes more stable

in a solvent that can hydrogen bond to carbonyl

[compare (LII)]. He has arranged the six solvents

used in his work in the order of increasing ability to promote proton transfer:

CeHe - CCL, < C2H4CI2 < CH2CI2

< CH3CN<CHCl3.

This order is based on the comparative ratios of the

intensities of the carboxylate and carbonyl peaks

in the different solvents. We referred earlier (sec. 4.3.3b) to Bruckenstein

and Saito's studies [352] of amine-carboxylic acid

interactions in benzene by the "DVP" method and

by infrared absorption spectroscopy. Examples of

acid-base systems for which they obtained IR ab- sorbance data are listed below.

Bases: Dodecylamine, dibenzylamine, piperidine,

A^,A^-dimethylbenzylamine.

Acids: Benzoic, trichloroacetic, trifluoroacetic. The stoichiometric molar concentrations, Cb and

Chxi were 0.01 n and 0.01 n', respectively, n and

n' usually having the value 1 but sometimes additional values in the range 0.5-5.0. The overaD results were interpreted as denoting

the formation of several species of salt, including

B+HX-, (BH+X-)2, BH+HX,7, and BH+X(HX)2, thus

being in partial agreement with the findings reported

above. The authors also postulate hydrogen bonding

in complex anions: for example, a structure similar

to the anionic portion of (L) is considered possible

for HXj. However, in pronounced disagreement

with the authors just referred to, Bruckenstein

and Saito concluded that their IR data provide no evidence of cation-to-anion hydrogen bonding,

Reactions between carboxylic acids and pyridine

in chloroform. As a sequel to Barrow and Yerger's

investigations of amine-acetic acid interactions in CCI4 and CHCI3, Barrow performed two types of

spectroscopic titrations to ascertain the associative

behavior of pyridine (Py) with seven aliphatic acids

(RCOOH) in CHCI3 [395]. The strengths of the acids

covered a span of more than four pK units. In the

order of increasing strength, as indicated by pK

values, the acids were: CH3COOH, CH2CICH2COOH, CH2ICOOH,

CH2CICOOH, CHCI2COOH,

CCI3COOH, CF3COOH.

In the first type of titration, Py was added gradually to RCOOH (maintained at 0.1 A/ concentration

throughout). Barrow deduced that where R=CF3

or CCI3 stoichiometric acid-base association oc- curred at Py concentrations up to 0.05 M, yielding

| the 1:2 salt Py(HA)2; to this product he assigned a

j hydrogen-bonded structure analogous to (L). With

weaker acids, formation of Py(HA)2 was less com- |! plete, 1:1 association being increasingly favored , until with R=CH3 or CH2CICH2 only a 1 : 1 adduct

j seemed to form. As evidence of the formation of

^ Py(HA)2 Barrow pointed out that the 6.97 p. absorp-

^ tion band of Py had been replaced by two bands r (a weak band near 6.12 p and a second band near

^ 6.72 p), both of which he assigned to PyH + ; fur- thermore, the "characteristic carboxylate band"

had appeared. According to Barrow's tabulated

results the CO^^ band varies in location, occurring

'

at 6.00 p in the formation of Py(HOCOCF3)2 and ^ gradually shifting to longer wavelengths with de- creasing acid strength until it appears at 6.41 p

for Py(HOCOCH2l)2.

Barrow was especially interested in the nature

of the 1:1 adducts. According to a commonly held

view, the association of a weak proton acceptor

such as Py {pKa ~ 5.2 at 20°) with a weak Br0nsted

acid like CH3COOH (pKa 4.76 at 25°) is properly

68
formulated as in eq 43. The product, symbolized

by B...H-A, is referred to as a "hydrogen-bonded

complex." The

B + H-A^B...H-A (43) H-A bond is considered to have retained its co- valent character, the hydrogen bond to B being the

result of electrostatic attraction. In contrast, a

similar interaction with a strong acid such as CCI3COOH ipK ~ 0.65) is considered to involve

incipient proton transfer and is therefore formulated

according to eq 44. In the product, which is dis- tinguished from that of

B + H-A^B-H+...A- (44)

equation 43 by the appellation "hydrogen-bonded

ion pair," the bond B-H+ is thought to have gained

a covalent character and the hydrogen bond to have become essentially electrostatic in nature. Barrow wished to ascertain whether the acid-base

associations represented by these two equations

are limiting examples of a single type of reaction

or, instead, represent two different reactions. The

first situation corresponds to a series of singleminimum potential energy curves for the proton,

with the minimum shifting gradually from a position near A to a position nearer B as HA increases

in strength. The second situation corresponds to a double-minimum potential energy curve. (See sec.

4.6; fig. 24.) Potential energy curves for different types of

proton-transfer reactions in aprotic solvents will be discussed more completely in sec. 4.6. In the

meantime it will suffice to mention that, if a doubleminimum curve apphes, one will expect there to be two potential troughs (wells) of unequal depth

in reactions involving either very weak (fig. 24(a)) or very strong (fig. 24(b)) acids. In the former case

the proton will remain near A, and in the latter case, near B. On the other hand, with acids of intermediate strengths one may expect to find reactions

where the two potential troughs are about equal in depth (fig. 24(c)); here, the proton may be able to

shift rapidly from one minimum (near A) to the

second minimum (near B) by leaping over or tunneling through the potential barrier. The problem of

whether the single-minimum or double-minimum

potential energy curve applies, apart from its theoretical interest, obviously has a practical importance in attempts to construct scales of relative acid and base strengths by determining equilibrium

constants or shifted spectral frequencies.

To ascertain the nature of the 1:1 adducts,

Barrow performed a second series of "titrations"

in which RCOOH was added gradually to Py in CHCI3 (maintained at 0.2 M concentration). He

concluded that his seven acids can be divided into three classes, as follows: Cla^s 1, consisting of the two weakest acids

(R=CH3 or CH2CICH2). These acids associate with

Py according to eq 43 — that is, they form the "nonproton-transferred species" Py...HOCOR. As

supporting evidence for this conclusion he noted

that an absorption band of Py found at 6.33 /jl in CCL4 shifts gradually to 6.24 /u, when the solvent is changed successively to CHCI3, n-BuOH, and

CH.iCOOH. In his view the gradual shift sig- nifies that nonproton-transferred species are formed in the three solvents (that is, Py...HCClt,

Py...HOBu, and Py...H0C0CH3), the differ- ing shifts in the 6.33-)U, band of Py indicating a difference in the degree of electrostatic attraction

between Py and the proton donor.

Class 2, consisting of the two strongest acids

(R = CF3 or CCI3). The product consists of "protontransferred" ion pairs (eq 44). The principal evi- dence is the appearance of the PyH+ bands near

6.12 (X and 6.72 /x, with approximately the same

intensity as in Py(HA)2. These two strong acids were found to combine stoichiometrically with

pyridine.

Class 3, composed of the three acids of inter- mediate strengths (that is, R = CH2l, CH2CI, or CHCI2). According to Barrow the double-minimum

concept was experimentally verified in these cases,

the products indicated by eqs 43 and 44 being

present simultaneously in tautomeric equilibrium:

R-& - R-/ , (45)

\ \ -

+

0-H...NC^H, O...H-NC^H^

5 5 5 5

(LXVI) (LXVII) The following observations were presented as evidence:

(1) The simultaneous appearance of the band near

6.24 ju, which he regards as characteristic of the

nonproton-transferred complex Py...HOCOR and

the two bands (near 6.12 and 6.72 p.) attrib- uted to PyH+ in proton-transferred ion pairs (for example, in CCl3COO"...H-NC5H5 and

+

CF3COO-...H-NC5H5).

(2) Intensification of the two PyH+ bands and

reduced intensity of the Py...HOCOR band with

increasing strength of the acid.

(3) Alterations in relative band intensities when

a mixture of Py and CHCI2COOH (both 0.4 M) in chloroform was heated from 25° to 55° by short

warming in an oven: the intensity of the 6.24-jU,

band increased by about 45 percent while that of

6.72-/X band decreased by about 16 percent. (However, the 6.12-/1 band appeared to increase in intensity.)

(4) Inability to detect any Py band that appeared

to assume continuously the character of PyH+ with

increasing strength of RCOOH.

In short. Barrow considers that his combined

results are most consistent with the hypothesis that a hydrogen-bonded complex and a hydrogenbonded ion pair are not limiting examples of a single

69
type of product. It should be noted, however, that he did not find completely satisfying explanations

for two observations: (1) The wavelength shift in the carbonyl-carboxylate band near 5.80 to 5.85 fx

is much smaller on ion-pair formation than would be

expected; (2) the band splitting expected in the case

of a tautomeric equilibrium is absent.

In the case of the pyridine-acetic acid system

equilibrium constants were determined for two

reactions (eqs 46 and 47) by performing "spectroscopic titrations" of 0.05 M acetic acid and combining the results with dimerization constants [265] for acetic acid in chloroform and carbon tetrachloride.

2Py + (CH3COOH)2^2CH3COOHPy (46) Py+CHgCOOH^CHgCOOHPy. (47)

The equilibrium constants (molar units) are as

follows: (1) For the reaction of acetic acid as the

dimer (eq 46), log K= 1.60; (2) for its reaction as the monomer (eq 47), log K values are 2.34 (CCI4 as

solvent) and 1.85 (CHCI3 as solvent). The last two

constants may be compared with the log K values

obtained for the EtsN — AcOH system: 2.90 (in CCI4) and 3.48 (in CHCI3) (see table 43). The re- versed solvent effects for the two systems provide,

in Barrow's opinion, an additional reason for believing that the Py-CHsCOOH and Et3N-CH3COOH

reactions yield different types of product. As readers

wiU recall, the greater value of log K for the

EtsN — CH3COOH system in CHCI3 was attributed

to stabilization of the carboxylate structure by hydrogen bonding of the solvent to carboxylate-carbonyl; see (LII). Evidently, when Py replaces Et3N as the

base the carboxylate-carbonyl does not become a

sufficiently good proton acceptor for hydrogen

bonding to CHCI3, and the role of this solvent

changes to that of a competitor of CH3COOH in the formation of hydrogen bonds to pyridine. These

examples illustrate well the inadequacy of the di- electric constant of a solvent as a single clue to its modifying effect on acid-base behavior.

Barrow's studies of 1:1 Py-RCOOH systems have

recently been extended by Johnson and Rumon

[399] to Py-CnHsCOOH and 17 additional adducts

in which the base was either pyridine, 3,5-lutidine,

2,6-lutidine, or 2,4,6-collidine, and the acid was

benzoic or a substituted benzoic acid; Py-CFsCOOH

and Py-CCl3C00H were also included. The adducts were examined as solids, and eight representative

adducts were also examined in acetonitrile solution. Johnson and Rumon were interested in the relation

between the type of hydrogen bonding (that is, B...II-A or B-H+...A-) and ^pKa(aq.) (that is, pK,i of base — pK,, of acid), and they were also con- cerned with the question of single-minimum

versus double-minimum potential energy curves.

In accord with common practice, the appearance

of two widely separated p-NH or v-OH bands was taken as evidence of a double-minimum

potential for the proton-stretching motion and of

a low barrier between the two minima, while a

single v-NH or i^-OH band below 1700 cm~' was

regarded as indicating a single-minimum potential function. From a detailed examination of

the absorption curves, with particular attention

to bands in the v-NH, f-OH, v-C^O, and v-CO

regions, they concluded that the adducts can be

classified in four groups: Classes A and B, which

are of type B...H-A, and classes C and D, which

are of type B-H+...A". Class A merges gradually

into Class B, and even Class D shows gradations in behavior. The following brief summary * indicates their line of reasoning: The spectra of Class A are closely similar to those of the parent

dimeric acids, for example, the CO2H band does 1 not indicate perturbation; however, two new, well- defined i^-OH bands indicate strong H bonding

between the base and the acid. In Class B, the bands

in the i^-C-0 region are stiU those characteristic

of RCO2H, but there is considerable broadening

of the v-C—O band and the two new v-OYi bands.

Among characteristics of Class C spectra are the

disappearance of certain bands attributable to i^-NH and v OW, and to v-C-0 as in the CO2H

group; a broadening and more complex character

of the v-C—0 band; and a broad and intense

absorption in the proton deformation region below

1200 cm"'. Class D adducts have distinct i^-NH

and V-CO2 bands, although they still show varying

amounts of spectral broadening.

The complete list of adducts is given in table 44,

together with values of Ap^a(aq.) and the type of hydrogen-bonded structure deduced (including ! inferences with respect to double- and single- i minimum potential energy functions). The following

conclusions should be emphasized:

(1) An abrupt change in the spectral characteristics of adducts occurs near ^pKa 3.75, signaling a

transition from B-II...A to B-H+...A-.

(2) In cases where a double-minimum function is indicated by two NH or OH bands it is believed that the band doubling is due to proton tunneling

through a low energy barrier. I, (3) No one of the single-minimum or double- I minimum systems showed a change in spectral 1 characteristics observed for the solid state after t being dissolved in MeCN. Evidently, a tautomeric / equilibrium between B...H-A and B-H+...A~, 1 which is not expected in the solid state, does not 1; occur in MeCN. j (4) Cooling samples to liquid nitrogen tempera- 1 ture caused little change in the spectra, thus ruling I out the possibility of attributing band doubling to a tautomeric equilibrium. 1

It is pertinent to mention at this point some ' dielectric polarization studies of Py • CH3COOH, ? Py • CH2CICOOH, and Py • CCI3COOH in benzene,

published in 1962 by Davies and Sobczyk [400]. it The polarization data yielded for the three com- r plexes the respective dipole moments (Debye units, s 20-21°) 2.97, 4,57, and 7.68. (An earlier study, which ( was cited in table 37, yielded slightly different

70
Table 44. The nature of hydrogen bonding in solid 1:1 pyridine-benzoic acid and

some related complexes, and its relationship to ApKa "

Base RCOOH Type of AddUct '' Class

Pyridine 2,4-Uimethylbenzoic 1.05 B...H-A (d) A

p-Toluic 0.86 B...H-A (d) A

Benzoic 1.03 *B...H-A (d) A

p-Chlorobenzoic 1.25 B...H-A (d) A

3,4-Dichlorobenzoic 1.62 *B...H-A (d) B

2,4-Dichlorobenzoic 2.55 B...H-A (d) B

4-Nitrobenzoic 1.81 B...H-A (d) B

3,5-Dinitrobenzoic 2.41 *B-H + ...A- (s) C

2,4-Dinitrobenzoic 3.81 *B-H+...A- (s) C

1 richloroacetic 4.58 *B-H+...A (d) Trifluoroacetic 5.00 *B-H+...A- (d)

3,5-Lutidine 3,5-Dinitrobenzoic 3.52 B-H+...A- (s) 2,4-Dinitrobenzoic 4.92 B-H+...A- (s)

2,6-Lutidine Benzoic 2.50 *B...H-A (d) 3,5-Dinitrobenzoic 3.88 B-H+...A- (d) 2,4-Dinitrobenzoic 5.28 B-H+...A- (d)

2,4,6-Comdine 3,4-Dichlorobenzoic 3.99 B-H+...A (d) D

2,4-Dichlorobenzoic 4.92 *B-H+...A- (s) C or D

3,5-Dinitrobenzoic 4.78 B-H+...A- (d) D

2,4-Dinitrobenzoic 6.18 B-H+...A- (d) D

" This table is a partial summary of conclusions of Johnson and Rumon from a

detailed study of IR absorption bands; samples were examined as muUs and also,

in some cases, in acetonitrile solution [399]. ApK„ = {pKa. base — pK„, acid)aq. The symbol (d) signifies that the hydrogen bond is believed to be a double- minimum type, and (s), that it is believed to be a single-minimum type. An asterisk

signifies that the adduct was examined in MeCN solution.

Reference:

{399] S. L. Johnson and K. A. Rumon, J. Phys. Chem. 69 , 74 (1965).

values.) These dipole moments are considered to indicate a definite hydrogen-bonding interaction in

all three cases, with the trichloroacetate having the

B-H+...A~ type of structure. The authors also concluded that in the frequency range up to 1700

MHz the solute particles behave like rigid-dipole

molecules, with no evidence of a "proton jumping

process." (However, they do not consider proton

jumping to be necessarily excluded in other acidbase complexes. Neither is it considered excluded

for the three above-mentioned complexes in a dif- ferent frequency range.)

Reaction of phenol-d, p-nitrophenol, and p- nitrophenol-d with amines in halohydrocarbons. To

obtain further evidence of possible tautoinerism in acid-base adducts, Barrow, in coauthorship with

[C. L. Bell [396], studied the association of 0.05 M

phenol-c? with A'^,A^-dimethylformamide, pyridine,

and triethylamine (apparently in large excess) in carbon tetrachloride and chloroform, by infrared

spectroscopy. As in Barrow's infrared studies of

Py-RCOOH systems [395], there seemed to be

band splitting. In addition, they determined the electronic ab- sorption curves of p-nitrophenol (6 X 10"'' M), mixed

with triethylamine (0.1 M), in the following four sol- vents: Cyclohexane, chloroform, ethylene chloride,

ethanol. The cyclohexane solution gave an absorption curve with a single strong band near 308 nm;

this was attributed to Et3N...HOC6H4-/>-N02.

An ethanol solution of p-nitrophenol made basic

with triethylamine (or KOH) also yielded an absorption curve with a single strong band, but the band was centered at 400 nm. However, when either

chloroform or ethylene chloride was the solvent, an

absorption band appeared in each of these regions.

This seeming "band splitting" in the ultraviolet was

interpreted as confirmation of Barrow's hypothesis

(based initially on infrared results) that a good many

1:1 salts, when dissolved in an aprotic solvent, are present as an equilibrium mixture of the nonproton-transferred species (B...H-A) and the

71
proton-transferred species {B-H+...A~). Tri- ethylammonium p-nitrophenolate and pyridinium

dichloroacetate are regarded as authentic examples

of such sahs.

A few analogous studies by others (for example,

measurements of the electronic absorption curves

of ethylene chloride solutions containing a fixed

amount of p-nitrophenol and varying amounts of

triethylamine [401]) have given comparable results and have received similar interpretations, and there seems to be a fairly wide acceptance of Barrow's

theory.

Comments on the double-minimum hypothesis as applied to 1:1 acid-base complexes in aprotic

organic solvents. It seems to this writer that there

is a serious obstacle to accepting Barrow's hypothesis: Band splitting in the ultraviolet seems to have

been observed only when the solvent was chloroform or ethylene chloride. No one seems to have

observed band splitting when the solvent was a hydrocarbon. For example. Bell and Barrow's ultra- violet absorption curve for triethylammonium p- nitrophenolate in cyclohexane contains only one

absorption band [396], and band splitting has also been found lacking in the ultraviolet curves for this

salt in toluene [404] and in benzene [17] (see fig. 25). How can this discrepancy be explained?

A clue can be found in table 32, which includes

examples of solvent effects on ion-pair dissociation constants of some tertiary ammonium picrates.

Table 32 indicates there to be an enormous differ- ence in the ion-pair dissociation of two closely

similar salts (t-AmsNHPi and BusNHPi) when one

of the two is dissolved in benzene, and the other is dissolved in ethylene chloride; that is, log K (where

K is the ion-pair dissociation constant) is —20.60 for t-AmsNHPi dissolved in benzene, and — 7.68 for BusNHPi dissolved in ethylene chloride. A few

other pertinent examples are to be found in table 33. These results of Kraus, Fuoss, and their associates,

like the still earlier conductance investigations of

Walden [322], have received inadequate recognition because attention has been focused for so long on water and water-Hke solvents. Actually, it is more than 50 years since Walden's experiments

provided clear-cut evidence that in the strictest sense practically none of the so-called "non-ionizing"

organic solvents is properly labelled as such. He

found that ionic dissociation of salts can occur in aprotic solvents like chloroform, dichloromethane,

1,2-dichloroethane (ethylene chloride), and even in solvents of dielectric constant as low as 2 to 3 (for instance, benzene, carbon bisulfide, carbon tetra- chloride, and toluene). The single absorption band

(near 400 nm) observed for triethylammonium p- nitrophenolate in ethanol is considered to indicate

that in this solvent the salt is present as solvated

ions, instead of as B...H-A. The overall reaction of triethylamine with p-nitrophenol can be

represented as:

EtsN + HOC6H4NO2 + (m + n)EtOH ^ (Et3NH+)mEtOH + (-OC6H4N02)nEtOH. (48)

In chloroform or ethylene chloride, which have only

limited ability to play a corresponding role, part of a salt can be expected to be present as the 1 : 1 acid- base adduct and part as ionic species. The proportion of ions should change with dilution.

If our supposition is correct, the two solvents can

be included in the large group of solvents that pro- mote partial, but not complete, ionic dissociation

of acid-base adducts. Among these solvents are

acetic acid [405, 406], acetonitrile [407], and nitro- benzene [24]. In acetic acid, for example, Kolthoff

and his associates have formulated 1:1 acid-base

interactions as: B + HA ^ BH+A- ^ BH+ + A-. (49) Ionization Dissociation

In their terminology [406], the shifting of the proton

from A toward B is referred to as "ionization," and' the separation of cations from anions through solva- tion is called "dissociation." To arrive at a satis- factory description of acid-base behavior in this solvent in numerical terms, equilibrium constants corresponding to both these processes have had to be determined. Analogous treatments have been

applied to acetonitrile and nitrobenzene solutions

(for example, see table 34). We venture to predict

that in future investigations ionic dissociation of

many acid-base adducts wiU be found to occur to a measurable, though perhaps only minute, extent in halogen-containing aliphatic hydrocarbons such as those mentioned above. Depending on the nature and concentration of the dissolved species, simple

ions, complex ions, or mixtures of these may be

formed (sec. 4.3.3a). This is a matter of considerable

importance in view of the extensive use of such

solvents in infrared spectroscopy.

Another conceivable explanation of the seeming

"band splitting" is the presence of unsuspected

impurities. The chemical instability of the aliphatic

halohydrocarbons was often overlooked in early

exploratory spectroscopic work, but a number of recent papers have included comments on their reactivity with organic bases; some examples were

given in section 3.4.6, and since that section was

written other examples have been encountered.

Furthermore, solvents like chloroform and ethylene

chloride react very readily with atmospheric water and oxygen when not protected from daylight

(apparently yielding hydrogen chloride as one of the products) [131], and it is not easy to maintain

them in a highly pure condition after the customary

stabilizer has been removed.

Another possibility for the introduction of an impurity into acid-base systems like those being

discussed is the spontaneous partial conversion of the tertiary amine to its oxide. It was discovered

during investigations of a related acid-base system

72
in benzene, in which the base was tribenzylamine,

that tribenzylamine oxide was formed on very brief

exposure of the solutions (at 25°) to air; even though

the yields were extremely small, the amine oxide —

which, like other tertiary amine oxides, is a very

good proton acceptor — affected spectrophotometric

results (in the visible region) [31]. Similarly, formation of tertiary amine oxides in scarcely more than

trace amounts is now believed to account for slow

drifts in the conductance of dilute solutions of tri- benzyl-, triethyl-, tri-n-propyl, and tri-n-butylammonium picrates in such solvents as toluene,

ethylene chloride, and o-dichlorobenzene [402, 403].

When such drifts were first observed, amine oxideformation was not known to provide a plausible

explanation, and they were tentatively attributed

to a slow tautomeric change of the kind indicated

by eq 45. Recent infrared investigations [390, 391]

have shown that various adducts of tertiary amine

oxides have characteristic absorption bands in this region. One would expect such bands to be located

rather close to corresponding bands for tertiary

amine-acid adducts, just as in the visible and ultra- violet regions (see figs. 14 and 15). On heating a pyridine-dichloroacetic acid mixture

in chloroform to about 55° Barrow observed certain

changes in relative band intensities which he

believed to be consistent with his hypothesis of a tautomeric equilibrium [395]. However, probable

temperature effects on equilibria of the types

indicated in eq 49 can account equally well for the

spectral shifts reported.

1 c. Results with wet aprotic solvents [408-414].

Throughout the experiments just described, careful

I efforts were made to exclude moisture. We shaU now

review very briefly a small group of experiments,

dating from about 1962, whose objective was to determine the nature of amine-acid adducts present

in aprotic (or dipolar aprotic) organic solvents after vigorous shaking of an amine-containing organic

solvent with an aqueous acid-containing solution. Such knowledge contributes to success in devising

: procedures for recovering, removing, and separat-

: ing acids. In some experiments attempts were made

i to remove any water carried into the organic layer

I by centrifuging or using desiccants, but rigorous

I drying conditions were not employed.

The reaction postulated for 1 : 1 association of an

, amine (R3N) with a monoprotic acid (e.g., HNO3,

. HCl) during the extraction process is indicated by

, eq 50. Equation 51 is a more general expression

f covering polyprotic acids (e.g., H2SO4, H3PO4) as

• well. The symbols "org" and "aq" indicate whether

Ijl a species is in the organic or the aqueous layer,

ii

' RsNorg + H+q + X-^ ^ RaNHX^rg (50)

j

i nR3Norg+ nH+q+ X-^ ^ [(R3NH+)„X-"]orff (51)

] Amines of relatively high molecular weight (e.g.,

I tri-ra-octyl- or tridodecylamine) are used to enhance

I solubility of the product in the organic layer.

In the experiments to be cited, the procedures

used for determining the composition and structure

of the extracted product(s) included not only conventional titrations for acid, base, and the water

content of the organic layer but also the recording

of IR spectra. The principal criteria relied on for detecting cation-to-anion hydrogen bonding were

shifts in t'(N-H) or i'(N-D) frequencies and evi- dence of deformation of the anion. Investigators were guided to some extent by IR band assignments

made previously in studies of solid salts. Table 45 gives examples of extraction experiments by several investigators [408-411] in which

the formation of hydrogen-bonded cation-anion

pairs in the organic layer was indicated by shifted

i'(N-H) frequencies. Benzene and chloroform are the principal solvents represented.

Table 46 lists examples of R:)N+-D...X~ bonding in which the solvent was carbon tetrachloride

and the base was n-OctsN [412]. The anions are

listed in the order of decreasing ability to function as deuterium bond acceptors. (The original

paper contains examples of R3N+-H...X~ bonding

as well; for example, see fig. 5.) The combined results given in tables 42, 45, and

46 permit the following generalizations:

(1) Conclusions about cation-anion hydrogen

bonding drawn from studies of solid specimens

generally agree closely with conclusions from studies

using aprotic solvents. (Here and elsewhere the

phrase "hydrogen bonding" is to be construed as including deuterium bond formation.)

(2) Anions of most common mineral acids (e.g., CI", NO3) are good hydrogen bond acceptors. This

statement applies also to acid anions such as HSOj and DSOj.

(3) As expected, the ability of anions to function

as hydrogen bond acceptors is greatest when the

negative charge is largely concentrated on a single

atom, as in simple ions like Cl~ or polarizable

anions like NO3 and Pi". When the negative charge

is diffuse, as in CIO4, anions are much poorer

hydrogen bond acceptors.

(4) However, even complex anions (for example,

Ph4B , ReOf, U02Clf, ZnClf) may exhibit some

ability to become hydrogen bonded to cations. It is not surprising that weak acceptors like CIO4 are

classified sometimes as hydrogen bond acceptors

and sometimes as nonacceptors.

There is no indication that small amounts of

water, likely to have been present in some of these

experiments, altered results significantly.

Table 47 presents examples of cation-to-anion

hydrogen bonding from Ryan's studies in the

visible and near-infrared regions [413-414]. In some experiments, the sohd salt was dissolved in the aprotic or dipolar aprotic solvent; in others, the

salt was formed in the organic layer by an extraction

or by metathesis. The anions listed are mainly

complexes of actinides: U(VI)02C1t, Np(VI)02Clj,

Pu(VI)02Cl4=, U(IV)Cl6=, and U(IV)Br6=. Apparently

C0CI4, also, can function as a hydrogen bond

73
Table 45. Some examples of R3N+-H...X~ formation in "wet" aprotic solvents, deduced from shifted v(N-li} frequencies ^'

R3N UX Solvents References

(n-C8H,7)3N HNO3 Benzene,'' chloroform, "modified" dodecane,*^ nitrobenzene

[408]

MeN(C«H,7)2N HNO3 " HCl " HBr « IJT d U/^l/^ e 111 rli^lU4

Cliloroform [409]

(«-C8H,7)3N HNO3, HCl, HCIO4,

H2S04, H3P04

Benzene [410]

(n-C,2H25)3N H2U02(N03)4f

HNO3

Benzene

Kerosene

[411]

The adjective "wet" signifies that the salt R3NHX was formed during extraction of HX from an aqueous solution into an RsN-containing organic solvent. Temperatures seem to have been in the range 20-25°. See articles cited for concentrations used.

^ Analogous results were obtained using DNO3.

"Modified" signifies that 6 vol.-percent of n-octanol was present. Strong hydrogen bonds.

' A weak hydrogen bond.

' Hypothetical formula. The acid-base reaction postulated is: 2R3NHNO3 (org) + U02(N03)2(aq) ^ (R3NH)2U02(N03)4(org).

References:

[408] J. M. P. Verstegen, Trans. Faraday Soc. 58, 1878 (1962). [409] J. I. Bullock, S. S. Choi, D. A. Goodrick, D. G. Tuck, and E. J. Woodhouse,

J. Phys. Chem. 68, 2687 (1964). [410] T. Sato, J. Appl. Chem. 15, 10 (1965). [411] T. Sato, J. Appl. Chem. 15, 92 (1965).

Table 46. Some examples of (n-Csliu)3^'^-^---^~ formation in "wet" CCI4, deduced from shifted i<N-D) frequencies

No. Anion No. Anion No. Anion

1 *ci- 6 Pi-, ThCl6=

' 11 *SCN2 *Br- 7 *CN-, Ph4B- 12 UOsCU"

3 *DCl2- 8 Re04- 13 ZnCU"

4 *N03- 9 DSO4- 14 CIO4-

5 CH3SO3- 10

^ The adjective "wet" signifies that the salts were formed

during extraction of DX from water into a CCI4 solution of

(/l-C8H,7).3N. The anions are listed in the order of decreasing ability to hydrogen bond to (C8Hi7)3ND+, as indicated by the IR shifts. This list is a combination of two sets of results presented in figures by W. E. Keder and L. L. Burger [412]; asterisks indicate anions included in their figure 4, the rest of the hst being taken from figure 6.

Reference:

[412] W. E. Keder and L. L. Burger, J. Phys. Chem. 69 , 3075

(1965).

acceptor. The cations listed are principally primary,

secondary, and tertiary ammonium ions; the hydrated hydronium ion and triphenylphosphonium

(a weak hydrogen-bond former) are also included.

In aU cases the cation is thought to be bonded to one of the anionic halogens. Cation-to-anion hydrogen bonding is strong enough in these examples to be detected when using dipolar aprotic solvents

(acetone, acetonitrile, nitromethane) as well as with an aprotic solvent (xylene).

It is noteworthy that bonding of anions to H,30+aq was detected in anion-exchange resins, as well as in solutions. In view of this finding Ryan believes that hydrogen bonding should be considered a factor in the adsorption of anionic metal complexes as well as

in their extraction into allcylamine-containing

organic solvents. As already mentioned, the results just discussed were obtained by spectrophotometry in the visible and near-infrared. The absorption curves of the

hydrogen-bonded species listed in table 47 showed

common similarities as well as distinctive differ- ences from the curves for corresponding quaternary

ammonium salts. The type of explanation offered

for such differences may be illustrated by the case of

74
Table 47. Some examples of hydrogen bonding of cations to MO2CI4 and U{IV)X|' in an aprotic or dipolar aprotic solvent, deduced by spectrophotometry in the visible and near-infrared

(25°) *

Cation Anion Solvents References

n-OctsNH* U02Clf Xylene [413] Et3NH+ U02Cl4= Acetone, acetonitrile. [413] nitromethane EtzNHj Nitromethane [413] n-OctNHj U02C1I Acetone [413] HaO+aq U02Cl4= Nitromethane [413] (C,2H25)3NH+ NpOzCU^ Xylene [413]

(Cl2H25)3NH+ PU02C1J Xylene [413] Et3NH+ PuO.Clf Nitromethane [413] R3NH+ UCl6= Acetone, acetonitrile. [414] nitromethane Et2NHj UClf Nitromethane [414] n-OctNHj UCli Acetone [414] HaO+aq UClf Acetonitrile ^ [414]

Ph3PH+ " UClf Acetone, acetonitrile. [414] nitromethane EtaNH? CoCl; ^ Acetonitrile [414] HsO+aq CoCl4=

Acetonitrile [414] Et3NH+ UBrf f Nitromethane [414] EtzNH:? UBre=

Nitromethane [414]

* M signifies U{VI), Np(VI), or Pu(VI); X" = CI' or Br".

^ The same, or a closely related, hydrogen-bonded species was

detected in anion-exchange resins (examined as mulls). No specific examples are given, but all three solvents seem

to have been used. A very weak hydrogen bond appears to form.

^ CoCl^ appears to be a fairly strong hydrogen bond acceptor.

' UBrf seems to be a weaker hydrogen bond acceptor than

UCl6=.

References:

[413] J. L. Ryan, Inorg. Chem. 2, 345 (1963).

[414] J. L. Ryan, Inorg. Chem. 3, 211 (1964). UCl^: The octahedral symmetry of the "simple"

anion, UC1|=, existing in quaternary ammonium salts

is destroyed on hydrogen bonding to R3NH+, etc.,' the distortion being sufficient to permit the occur- rence of electronic transitions that are forbidden

in the simple anion [414]. Ryan's studies of substituted ammonium salts containing UO2CI4 as the anion indicated that cation-anion pairs which seem to be hydrogen

bonded in aprotic solvents are not necessarily

hydrogen bonded in the solid phase [413]. For ex- ample, he observed that the spectrum of (Et3NH)2

U02CL4 as the solid is almost identical with the

spectrum of the quaternary ammonium salt (Pr4N)2

UO2CI4. Moreover, in studying the secondary am- monium salt (Et2NH2)2U02Cl4, which has two solid phases, he observed that the solid phase which is stable at 25° has an absorption spectrum very similar to that of "simple" UO2CI1', whereas the second

soHd phase, which is formed rapidly on heating to ~ 100°, has a spectrum somewhat similar to that of

the salt in a nitromethane solution (in which it is thought to be hydrogen-bonded—see table 47). Probably geometrical factors sometimes prevent

cation-to-anion hydrogen bonding in solid species.

(It also seems possible that in some instances multiple hydrogen bonding occurs, leading to a relatively

symmetrical anionic structure.)

From results cited in tables 42 and 46 it seems

possible that Ph4B~ is another anion that may

hydrogen bond to R3NH+ when in an aprotic solvent,

although not in crystalline salts. Ryan believes that hydrogen bonding of cations to chloro complex anions probably occurs widely. As

expected, UBrg appears to be a weaker hydrogen

bond acceptor than UCl^ [414]. Indications of

hydrogen bonding to 002(804)3" and U02(N03)i-2 were mentioned earlier in this section [389, 411]; in these cases oxygen is the proton-acceptor atom, of course.

d. Further discussion. There seems to be a general harmony in the infrared results with respect to cation-to-anion hydrogen bonding, at least for the

systems reported above, irrespective of whether the

materials were examined as soUds, as "dry" solu- tions in aprotic solvents, or as "wet" solutions. This

is worth emphasizing in view of the still frequent

assumption that acid-base behavior is always

strongly affected by water, even when the latter is present only in trace amounts. As a matter of fact, we concluded earlier (sec. 4.2) that traces of water

probably do have a considerable effect on the self- association of carboxylic acids. However, the

hydrogen bonds in carboxylic acid dimers and those

in hydrogen-bonded ion pairs are different in strength; the latter are often appreciably stronger

(see association constants and thermodynamic data

in sec. 5) and not easily ruptured through the

agency of a comparatively weak hydrogen-bond

former such as water.

It also seems desirable to take special note of the

striking parallelism in the behavior of the indicator

acid bromophthalein magenta (E or B) toward

amines, shown in visual tests and spectrophotometry

in the visible and near-ultraviolet (sec. 4.3.4a), and

reactions of acetic acid with amines, as studied by

infrared spectroscopy (sec. 4.3.56). Both sets of

investigations yielded strong evidence of cation-to- anion hydrogen bonding, as expressed, for example,

in structures (XLI) and (LI). The close parallelism in behavior can be understood on noting the similarity

of the structures written below for the indicator

(XXXVII) and the carboxylic acid (LXVIII). Some

years ago Madelung [415], and also Schwarzenbach

and associates [416], drew attention to analogies

between molecular resonance systems such as the

ions of various triphenylmethane dyes and simpler

ions like benzoate, carbonate, amidinium, and

guanidinium. For example, the anion of benzaurin,

C6H5-C(=C6H4=0)C6H40~, is very similar to the

benzoate ion except in having a phenyl group in280-613 OL-68—

75
serted between the central carbon atom and each

of the oxygens. Obviously, bromophthalein magenta

and related dyes, including the sulfonephthaleins,

may be regarded as derivatives of benzaurin. (Similar analogies are those between CO3 and the anion

of aurin, and between cations of well-known tri- phenylmethane dyes like crystal violet and malachite

green and the simpler cations, guanidinium and

amidinium.)

The two sets of investigations also showed the

readiness of a secondary ammonium ion to form an

additional hydrogen bond by using its "spare"

proton. Frequently this bonding leads to a dimeric

structure; for example, compare (XLIII) with (LV).

The cation-anion bond appears to be the stronger of

the two kinds of hydrogen bond; see (XLIV) and

(LVI) and related discussion.

It is evident that the dramatic color changes which

bromophthalein magenta undergoes in hydrogen

bonding reactions can be used to illustrate general

principles. It is a valuable teaching aid in either

lecture demonstrations or simple laboratory experiments, but for its fullest exploitation in such a manner, a higher degree of purity is desirable than

is generally required in using indicators for endpoint determinations.

4.3.6. Nuclear Magnetic Resonance Spectroscopy

[412, 417-422]

Applications of this newer technique to studies

of hydrogen bonding in ion pairs have given results

in substantial agreement with those obtained by

other methods.

a. Measurements of crystals. In an early application of NMR spectroscopy both the proton and the

fluorine magnetic resonance spectra of powdered

ammonium fluoride were measured (140-360 °K)

[417]. The data indicate the presence of very strong

N+—H...F~ bonds, in agreement with conclusions

from infrared [382] and neutron diffraction [383] data. In the case of NH4BF4, also, there is agreement in the resuhs of infrared [384-386, 388] and NMR

[418] investigations. It was inferred from the proton

and fluorine magnetic resonance spectra (20 and

90 °K) that only "exceptionally weak" hydrogen

bonding is possible. In the crystal the BF7 ion is packed with NH| in the same manner as with Rb+,

and it is impossible for any one NH^ ion to form four simultaneous linear bonds to fluorine.

b. Studies of solutions.

(1) In a study of the hydrochloride, hydrobromide, and hydroiodide of pyridine in CH2CI2,

CH3CN, and CH3NO2 it was concluded that the

three salts exist as hydrogen-bonded ion pairs in these solvents [419], in agreement with Cook's

conclusions from an infrared investigation of the same salts as sohds [423].

(2) Following earlier exploratory studies by

Keder and Wilson [420], Keder and Burger [412]

completed comparative infrared and proton magnetic resonance (PMR) investigations of hydrogen

bonding in "wet" CCI4 solutions of numerous tri- /i-octylammonium salts. Some of their IR results

(for deuterium bonding) have already been summarized (see table 46 and related discussion).

Table 48 is a summary of conclusions about hydrogen bonding, obtained by the PMR method.

Table 48. Some examples of (/i-CNHi7)3N+-H...X- formation in "wet" CCI4, deduced from chemical shifts o/N-H pro- tons

No. Anion No. Anion No. Anion

1 *ci- 6 Pi- 12 *CN2 CH3SO3- 7 *HCl2 13

3 ThCU- 8 *NOf 14 U02C1=

4 Ph^B 9 ReOj 15 *SCN5 *Br- 10 HSOj 16 CIO4

11 *I- 17 ZnCl;

The adjective "wet" signifies that the salts were formed

during extraction from water into a CCU solution of (n-C8Hi7)3N.

The probe temperature was ~ 33°. The anions are listed in the order of decreasing ability to hydrogen bond to (CsHi7)3NH+, as indicated by the chemical

shifts. This list is a combination of two sets of results presented

in figures by W. E. Keder and L. L. Burger [412]; asterisks indi- cate anions included in their figure 4, the rest of the list being taken from figure 6. To compare these NMR results with conclusions from

parallel experiments in which shifted KN-D) frequencies were

determined, see table 46.

Reference:

[412] W. E. Keder and L. L. Burger, J. Phys. Chem. 69, 3075

(1965).

Experimental procedures were analogous in the

two sets of investigations. It was found that two

fairly good hnear correlations of N-D stretching

frequencies and chemical shifts of N—H protons

exist: one for relatively simple anions and an- other for more complex anions (compare tables 46

and 48, in which the simpler anions are marked

with an asterisk). By both methods, CI" is the best

proton or deuteron acceptor among the anions

studied, while CIO ~ and ZnCl= are the two poorest

acceptors.

(3) Experimental results from a PMR investi- gation of the rates of proton transfer (35°) between

nine methyl-substituted ammonium salts and their conjugate amines in t-butyl alcohol (e=11.4 at 30°) have been interpreted as evidence of cationanion hydrogen bonding in this solvent [421, 422].

76
The cations represented were MeNH + , Me2NH + , and MesNH^; and the anions (X~) were p-CHsCslliSO2O-, Br-, CF3COO-, and CI". The solvent is considered to be a participant in the proton transfer

reaction, which may be represented by eg. 52

(where R signifies H or Me, and R' signifies t-Bu):

Me Me Me

R2N-H+...X- + O-H + NR2 > R2N

R' Me

+ H-0+-X...+H-NR.>. (52) R'

It is not known whether more than one molecule

of ?-BuOH is involved in the transfer of a proton

from salt to amine.

In analogous earlier experiments with water or methanol as the solvent, the proton transfer reaction

had been found to involve dissociated ions, proceeding either by a direct proton transfer, as indicated

by eq 53, or by an indirect transfer involving a molecule of solvent, as indicated by eq 54 (where

R" = H or Me):

Me Me ^ Me Me

R2N-H+ +NR2 —^ R2N + +H-NR2 (53)

Me Me , Me

R2N-H+ + 0-H + NR2 R2N

R"

Me

+ H-O + +H-NR2. (54) R"

The reactions proceed much more slowly in t-butyl alcohol than in the other two hydroxylic

solvents. Moreover, the proton-transfer rates in i-BuOH were found to follow second-order kinetics

(first order each in the substituted ammonium salt and the corresponding amine), whereas a mechanism involving dissociated cations would require

the rate to depend on the square root of the salt concentration. Still further, the rate constants in f-BuOH are substantially smaller for the trimethylammonium salts than for the mono- and dimethylammonium salts, and they also vary according to the anion. For such reasons, as well as because of

certain features of the NMR spectra, it was con- cluded that in ^-BuOH these substituted ammonium

salts all exist as hydrogen-bonded ion pairs, in which one NH proton is bonded to the anion and

any remaining NH protons are bonded to the solvent,

as suggested by (LXIX):

Me_N-Ht . .X

2|H

0-H

I t-Bu

(LXIX)

For the trimethylammonium salts, k2 increases in the following order:

CI- < CF3COO- < Br- < P-CH3C6H4SO2O-.

The NH chemical shift follows a slightly different

order for the di- and trimethylammonium salts,

increasing in the order:

CF3COO- < Cl< P-CH3C6H4SO2O-.

4.3.7. Concluding Discussion of Hydrogen-Bonded

Ion Pairs

It has been thought necessary to place considerable emphasis on cation-anion hydrogen bonding,

since the evidence — being scattered and, for the

most part, fragmentary — has been widely over- looked. Nevertheless, it is clear that abundant and

diverse evidence of cation-anion hydrogen bonding

has been accumulated in studies of incompletely

substituted ammonium salts, both as solutions in aprotic or dipolar aprotic solvents and as crystalline solids. In some cases the hydrogen bonds are strong enough to persist in moisture-containing

aprotic solvents or in a hydroxylic solvent like tert. butyl alcohol. Restating briefly some of the experimental re- sults, we note, for example, that incompletely substituted ammonium salts have been found to differ from their completely substituted analogs in the

following respects: (1) The ion-pair dissociation

(eq 32) is much less extensive in solvents like benzene, chlorobenzene, ethylene chloride, or nitrobenzene, although closely comparable in extent

in water, alcohols, or a basic solvent such as pyridine (for instance, see table 33). (2) The dipole

moments are several Debye units smaller (compare

table 36). (3) The tendency toward self-aggregation

with increasing concentration in an aprotic solvent,

as measured cryoscopically, is substantially less

(table 39, fig. 11, and related discussion). (4) Col- ored salts formed with indicators like BPM-E

and HPi are not so "deep" in color (that is, the

main absorption bands are shifted to higher fre- quencies)— see figures 14, 15, and 17. Although the

spectrochemical results furnish the most concrete evidence of cation-anion hydrogen bonding

in the incompletely substituted salts, the other

physicochemical data, especially conductance data,

provide strong "circumstantial evidence." Additional direct evidence is supplied by results ob- tained for incompletely substituted ammonium

salts using infrared and proton magnetic resonance

spectroscopy. In short, it is overwhelmingly evident

that many incompletely substituted ammonium salts

77
and their completely substituted analogs exhibit

strongly diflferentiating chemical and physical be- havior in aprotic media, which is satisfactorily

explained by the postulate that in the incompletely

substituted salts a strong hydrogen bond unites the

ammonium nitrogen and the anion. When the anion

is polyatomic (for example, the anion of a carboxyhc

acid or of bromophthalein magenta, nitric acid,

sulfuric acid, or a sulfonic acid), the nitrogen is hydrogen bonded to a specific atom, most commonly

oxygen.

That cation-anion hydrogen bonds are much

stronger than bonds of the better-known category

B...H-A is indicated by values of — AG and — A//

for reactions yielding B-H+...A". For example,

the overall change in enthalpy {minus A//) in the 1:1 association of diphenylguanidine with

benzoic acid or one of the toluic acids in benzene

appears to be about 16 kcal/mole (25°, molar units),

and for association of diphenylguanidine with three

stronger acids (o-, m-, and p-chlorobenzoic acids)

the enthalpy change appears to be in the vicinity

of 18 to 19 kcal/mole [33]. These results may be

contrasted with the values, ranging from about 4

to 7 or 8 kcal/mole, which formerly were regarded

as representative for hydrogen bonding reactions

[102]. (Thermodynamic constants obtained for other

acid-base reactions yielding hydrogen-bonded ion

pairs will be included in sec. 5.) The closely parallel behavior of primary, sec- ondary, and tertiary amines in leveling solvents

has led chemists to regard them as varying only in relative strengths, and until quite recently comparative investigations of their acid-base behavior

in aprotic media had not been performed. It is now apparent from experimental results cited above

that physicochemical behavior in aprotic media

sometimes sharply differentiates tertiary ammonium

salts from secondary and primary ammonium salts. To provide a brief recapitulation of the evidence,

the following examples may be mentioned: (1) Maryott found (see fig. 10 and related discussion)

that increasing the mole fraction (A'^2) of tributylammonium picrate in benzene produces essentially no change in the ratio Ae/A^^, signifying that BusNHPi does not form aggregates to a measurable

degree in the concentration range investigated.

Closely similar results were obtained for solutions

of n-OctNHgPi, n-Bu2NH2Pi, and /i-BugNHFi when

the solvent was dioxane. In contrast, Ae/A2 dropped

off sharply with increasing values of N2 in the case

of n-Bu2NH2Pi in benzene, the data being consistent with the postulate of a monomer-dimer

equilibrium; the dropping-off was still sharper for solutions of n-OctNHsPi in benzene. (2) From the

combined results of "dielectrometric titrations"

and "cryoscopic titrations" of benzoic acid with

amines in benzene, Gur'yanova and Beskina

reached analogous conclusions (see table 38 and

related discussion). (3) Applying the differential vapor pressure method in titrations of amines with

jcarboxylic acids in benzene, Bruckenstein and

Saito reached exactly comparable conclusions

(sec. 4.3.3b). (4) In visual and spectrophotometric ex- periments, using benzene and other aprotic solvents,

and with bromophthalein magenta and related

phenolsulfonephthalein derivatives as reference

acids, Davis and Schuhmann deduced that the

tertiary ammonium salts are monomeric, whereas

the secondary ammonium salts are dimerized

through dual hydrogen bonding of cations to anions;

the primary ammonium salts show clear but less well-defined indications of self-association through

hydrogen bonding (sec. 4.3.4). (5) Applying infra- red spectroscopy in studying the association of

acetic acid with triethylamine, diethylamine, and

n-butylamine in carbon tetrachloride. Barrow and

Yerger concluded that salts having two or three

protons attached to the ammonium nitrogen are extensively dimerized (sec. 4.3.5). Experiments with salts having a common cation

(R3NH+) — notably, Keder and Burger's studies,

by IR and NMR spectroscopy, of numerous n- Oct;iNH+ or n-0ct3ND+ salts in "wet" CCL4 [412]-

have shown that anions vary in their ability to become hydrogen bonded to cations (see tables 46

and 48). In 1939 Kilpatrick predicted [424] that this ability would be found to increase with increasing

basic strength of the anion, and his prediction has

proven to be correct. Hantzsch and co-workers

arranged various well-known acids in the order of decreasing strength, as established in several series

of semi-quantitative experiments employing static (indicator) or kinetic methods; in part of their work

the solvent was chloroform or acetic acid, while

other experiments were performed using concentrated aqueous solutions [19, 425, 426]. Three such

lists follow:

HCIO4 > HI > HBr > RSO3H

> HCl > HNO3 > CCI3COOH

HCIO4 > CCI3SO3H > CH3C6H4SO3H

> H2SO4 > HNO3 > CCI3COOH

CCI3COOH > CBr3C00H > CH2CICOOH

> CH2BrC00H > HCOOH > CH3COOH.

The first two lists indicate differences in the

strengths of acids that, at usual concentrations in water, are leveled to the same strength. On comparing the above lists with the extended lists in tables 46 and 48 and the fragmentary lists given elsewhere

in the discussion, it becomes evident that, with

few exceptions, the relative ability of anions to hydrogen bond to cations parallels ability to combine with protons to form neutral Br0nsted acids. For example, the anion of HCl, which is one of the

weakest in the group of strong acids, is widely

78
regarded as possessing the ability to become hydrogen bonded to cations, whereas C10~, the anion of

the strongest Icnown Br0nsted acid, exhibits Httle or no abihty to become hydrogen bonded to cations.

An understanding of the broad principles governing hydrogen bonding and its influence on acidbase behavior is valuable in many ways. For ex- ample, it might be thought that ability of anions to become hydrogen bonded to cations would promote

ease of extractability into amine-containing aprotic

solvents, but this is evidently not the case. From

the series of experiments cited above [412], Keder

and Burger concluded that "the most weakly hydrogen-bonded salts are those whose anions extract most strongly. This indicates that anion-cation

bonding in the organic phase is of little importance

in determination of the relative anion extractability. Rather, it appears that those anions which

are the best proton acceptors extract most poorly

because of their greater solvation in the aqueous

phase. . . ."

Factors determining the proton-acceptance ability

of anions are generally well-known — they include

charge, size, and polarizability, and possible steric factors must also be borne in mind [4, 427]. Obviously, anions will be better able to become hydrogen bonded to cations when their negative charge

is highly localized. The perchlorate ion is noted for

its small polarizability [427]. In contrast, the picrate

ion is regarded as highly polarizable, with its negative charge probably localized largely on the

phenolic oxygen atom; however, its high polarizability is partly offset by the bulkiness of the nitro groups present in the 2- and 6-positions, which

impede the close approach of any bulky cation

acid [4]. Our survey has shown that carboxylate

ions and the anions of bromophthalein magenta

and related indicators have a marked propensity

for hydrogen bonding to cations. This should be

expected, since anions of these chemical types

are known to be good proton acceptors.

In section 4.3.5a we cited the conclusion of

Sandorfy and associates that tertiary ammonium

ions form stronger hydrogen bonds to halide ions

than do primary and secondary ammonium ions;

this difference is understandable, since Et3NH+ can

form only one cation-anion hydrogen bond and the

cationic charge is highly localized. Davis and

Schuhmann [428] explained certain features of

absorption curves for mixtures of bromophthalein

magenta E with primary amines in benzene along

similar lines: that is, as signifying that multiple

hydrogen bonding by a single cation leads to in- dividually weaker cation-anion hydrogen bonds.

It should again be emphasized that it is likely to prove advantageous to use tertiary amines as ref- erence bases, in preference to primary and sec- ondary amines, when determining comparative

acid strengths in aprotic solvents or titrating acids

in such media.

4.4. Hydrogen Bonding of Cations to Neutral

Proton Acceptors

4.4.1. Introduction

Two types of hydrogen bonding remain to be

discussed: (a) Hydrogen bonding of cations (for example, H.,0+, Ph,GH+, PyH+, R,NH+, R.N^OH)

to neutral proton acceptors, which will be dis- cussed in this section; and (b) hydrogen bonding

of anions (ArO", CI", RCOO", etc.) to neutral proton donors, which will be discussed in section 4.5. Evidence will be presented that many cations of

the types mentioned (symbolized by BH+) can form

hydrogen bonds to their conjugate bases (B), giving B-H+...B, or to other neutral proton acceptors

(B') of widely varying acceptor abiUty, yielding B-H+...B'. Similarly, anions (symbohzed by A~)

can become hydrogen bonded either to their conjugate acids (HA), forming A~...H-A, or to other pro- ton donors (HA'), both strong and weak, giving

A"...H-A'. Such associations are not invariably

1:1, but this is the ratio most commonly encountered

It is often convenient to refer to such complex

cations or anions as ''conjugate ions," a term intro- duced by Kokhoff [429]. Kolthoff also coined the more specific designations "homoconjugate" and

"heteroconjugate"; these are convenient when

one wishes to indicate whether a conjugate ion is homogeneous in composition (that is, B-H+...B

or A~...H-A) or of heterogeneous composition

(B-H+...B' or A~...H-A'). Heteroconjugate ions

in which B' or HA' is a weak proton donor and/or

acceptor molecule and constitutes all or part of

the solvent are more famiUarly known as "solvated

ions." We shall adopt KolthofTs terminology for the hydrogen-bonded cations and anions about to be discussed, reserving Kraus and Fuoss's expression "triple ions" [326, 328] for non-hydrogenbonded ion aggregates, M+X"M+ and X~M+X-,

formed through electrostatic attractions (see sees. 4.3.1c and 4.3.3a).

4.4.2. Solvated Cations

a. Solvated hydronium ions. It is generally ac- cepted that the proton forms a very stable complex

with one molecule of H2O — namely, the hydronium

ion, H3O+. Furthermore, on the basis of experimental evidence such as that summarized by Clever

[34], the hydronium ion is now believed to exist

in most aqueous solutions (all but extremely con- centrated solutions) mainly as the trihydrated

species indicated below (LXX); this is jsrobably

in equihbrium with the mono- and dihydrated

species. Recent extraction experiments by Diamond,

Tuck, and others [430-435] have contributed additional evidence that H3O+ bonds to its conjugate

base, H2O, in the manner indicated by (LXX), and

also to other proton acceptors; this is referred to as "primary" ("first-shell") solvation. Further

("secondary") solvation of the primary unit is now

believed to occur [430]. An example of secondary

solvation will be given presently.

79
H H

I I H-°-.. ,-°-H

H, H

h/°-h

(lxx)

Experiments have demonstrated the feasibiUty

of extracting numerous mineral acids (for example,

HCIO4, H2SO4, HNO3, hydrogen hahdes) and

complex metal acids (HFeCl4, HAUCI4, HInBr4,

and others) into basic organic solvents. From this

list Diamond and coworkers selected HCIO4 (often

with tracer HRe04) for most experiments where

the main objective was to determine the composition of solvated hydronium ions present in the or- ganic layer after extraction procedures. They

recognized that the anions of these acids — which,

as we have just seen, do not bond readily to protons

or proton donors — were not likely to transport

enough water into the organic layer to confuse the

experimental results. Other acids used were HBr,

HAuCLj, and HAuBr4. At first, several basic organic

solvents were used as extractants — for example,

ethers, ketones, and esters (notably BU3PO4)

[430, 431] — but in later work the organic extractant consisted of an aprotic solvent (carbon tetrachloride, isooctane, or xylene) containing tributyl

phosphate [432-434] or trioctylphosphine oxide

[434, 435]. The compositions of extracted cations were determined by such means as measuring

volume changes, acid-base and Karl Fischer titra- tions, and ascertaining the water content of the

organic layer by IR spectroscopy.

As expected, the compositions deduced for ex- tracted species were found to depend on such vari- ables as the strengths and relative proportions of

competing proton donors and acceptors, and also on the steric availability of the acceptor atom. For

example, in the extraction of HCIO4 into dilute (<0.1 M) BU3PO4 in CCI4, the nature of extracted

species was found to depend on the stoichiometric

ratio BU3PO4/H+: When this exceeds 3, the only

two species thought to be extracted into the or- ganic layer are Bu3P04-H20 and ion pairs in which

the cation composition is 3Bu3P04 H30+ yH20,

y having a value between 0 and 1. However, when

BU3PO4/H+ exceeds 1 but is less than 3, several

solvated species may be present in the organic

layer. When the aqueous acid is very concentrated

(~ 10 M), no water is extracted into the organic

layer, and the extracted species is then believed

to be (LXXI) [432]:

(LXXI)

Structures considered reasonable for two types of

solvated hydronium ions are given below (LXXII,

LXXIII):

(BuO)jPO._ _.OP(OBu)^

*H H'

0

p(OBu)^

(LXXII)

(Primary solvation of hydronium ion)

(BuO),PO...H H H H...OP(OBu)

H H \ /

0

0—

I H

OP(OBu)^

(LXXIII)

(Note both primary and secondary solvation)

In (LXXII), each hydronium proton is bonded di- rectly to a phosphate oxygen, while in (LXXIII),

three H2O molecules act as intermediate bridges.

Structure (LXXI) is an ion pair with perhaps a weak cation-anion hydrogen bond.

Analogous structures were inferred for the products resulting from extractions of perrhenic acid

(containing radioactive ReOj) [432] and HBr [433] into a Bu3P04-containing aprotic solvent, as well

as from extractions of HCIO4 [435], HRe04 [435],

HAuCU [434], and HAuBr^ [434] (the latter three

acids being used in tracer amounts) into an aprotic

solvent containing OctsPO.

The following are some of the principal over-all conclusions of the authors cited, with respect to extractions of acids from aqueous solutions into or- ganic media of the types mentioned [433-435]: (1) When the acid is strong and only moderately con- centrated, the proton is extracted as H3O+ and be- comes trisolvated in the organic layer. (2) The ex- tent of hydration of H3O+ in the organic phase de- pends on the strength and concentration of the base

(B). (3) The cation composition H30+-3B-yH20

(0^ y ^ 3) should be found for any strong acid.

(4) OctsPO, being a stronger base than BU3PO4, can

be expected to compete better with H2O for the

hydronium ion, and is more likely to form the primary solvent layer, as in (LXXII). (5) If the anion is large and very weakly basic, it enters, and remains

in, the organic layer as an ion, being "essentially

ejected from the aqueous . . . phase because of

its disturbance of the hydrogen-bonded water

80
structure". The more hydrophobic, feebly basic,

and bulky the anion, the more easily is it ejected

into the organic layer [434]. (6) In aprotic organic

solvents the anions pair with solvated cations.

(Ion pairing does not occur, of course, in organic

extractants of high dielectric constant.) (7) Extraction of HCIO4 from concentrated aqueous so- lutions into a carbon tetrachloride solution of

OctsPO yields Oct3POH+...C104- [compare (LXXI)J

and perhaps also a hydrate such as (LXXIV):

0ct^P0H"t..0H2...C10^~

(LXXIV)

(8) If the extracted anion (e.g., NO3-, CCI3COO-,

CH3COO") is moderately basic, it will compete for the proton with water and the extractant, and if successful, will form the molecular acid (HA),

which can form only one hydrogen bond to the base

B. The anion will then appear in the organic layer as monosolvated HA — for example, BU3PO4 • HOCOCCI3, BU3 PO4 • HONO2, Oct3PO • HONO2.

Such extractions provide excellent examples of

practical situations in which the outcome depends

on which among various competing hydrogen bonding reactions are allowed to become dominant.

Brief mention should be made of Diamond's dis- cussion of the extractability of uni-univalent salts in which both the cation (for example, a quaternary

ammonium or phosphonium ion) and the anion

(such as a homolog of CH3C00~) are large and

hydrophobic (or at most poorly hydrated) [436]. In such instances ion pairing occurs in water, and

the ion pairs tend to be ejected completely from

water. Their extractability into aprotic solvents

is correspondingly facilitated, to an extent increasing with their bulkiness and hydrophobicity. Such

ion pairing is not considered to result from electrostatic attraction or to involve solvation, but is attributed principally to attempts of water to pre- serve its hydrogen-bonded structure, and it has been

christened water structure-enforced ion pairing."

Mention should also be made of Douglas and

McDaniel's brief review [437] of probable hydration

of H3O+ in solid acid hydrates, including HiSnCle

• 8H2O, the tri- and tetrahydrates of HAuCU, and

hydrates of HBr and HCl.

h. Miscellaneous solvated cations. Solvation of

incompletely substituted ammonium ions and re- lated cations, regarded for a long time as the probable cause of various chemical and physical phenomena, is very plausibly accounted for in terms

of hydrogen bonding. A number of examples of

cation solvation have already been mentioned.

For instance, we mentioned earlier (sec. 4.3.1a)

Walden and Birr's discovery that the mobilities of

Me4N+, Et2NH^, and r-BuNH^, although very similar in acetonitrile, are very different in water,

decreasing in the order named [324]. These authors

concluded that decreasing mobilities in water

probably correspond to increasing degrees of sol- vation, the closely similar mobilities in acetonitrile

indicating lack of solvation. Later studies by others

have added support to these conclusions as it has

become increasingly evident that MeCN is a much

poorer proton acceptor than water [438, 71 J. Using similar reasoning, Trotman-Dickenson

offered a plausible explanation of anomalies in the relative basicities of ammonia and primary,

secondary, and tertiary amines in certain solvents

[439, 440]. For example, the relative strengths of

NH3, IVIeNHa, MeaNH, and MesN in water, when

expressed as the pKa values of the corresponding

conjugate acids, are, respectively, 9.24, 10.62, 10.77,

and 9.80. If the relative strengths in water varied

strictly in accordance with the predicted inductive

effect of alkyl substitutions, Me3N would, of course, have the highest pKa value; instead, this base ap- pears to be weaker in water than the other two

methylamines. Trotman-Dickenson believes that such anomalies are caused mainly by differences

in the extent of stabilization of the various cations

by hydration, which obviously wiU depend on the

number of protons available for bonding to H2O.

This explanation, which has won general accept- ance, has been substantiated by later findings that mono-, di-, and tri-n-butylamine increase progressively in basicity — in accord with prediction —

when the basic strength is determined (by an in- dicator method) in aprotic solvents like chloro- benzene, bromobenzene, or fluorobenzene [354,

365, 441].

Turning to other types of physicochemical measurements, it is appropriate to recall four investi- gations, reviewed above in more detail, in which

patterns of acid-base behavior observed when

using an aprotic or dipolar aprotic solvent (or no

solvent) were found to be modified by adding or substituting a hydrogen bond-forming solvent:

(1) Experiments of Witschonke and Kraus (see table 35), in which the conductance of certain in- completely substituted ammonium salts in nitro- benzene increased on the addition of a basic solvent

(water, pyridine, piperidine, or triethylamine), the

increase being proportionate to the basicity and

concentration of the added solvent; (2) Maryott's

dielectric polarization experiments (see fig. 10), in which OctNHsPi, Bu2NH2Pi, and BugNHPi ex- hibited differentiating behavior in benzene (e=2.27)

but essentially uniform behavior in dioxane (e=2.21);

(3) visual and spectrophotometric investigations

of bromophthalein magenta and its salts in varied

organic solvents, by Davis and co-workers (sec.

4.3.4a); (4) Chenon and Sandorfy's infrared investi- gations [375], which led them to conclude that amine hydrohahdes form N+-H...X" bonds in the sohd state, but N+-H...0 bonds when in aqueous solutions (see sec. 4.3.5a). In each 01 these

investigations, the effects of addition of, or replacement by, a basic solvent were attributed to hydrogen bonding of cations to the basic solvent in question.

The examples just cited have all involved n-elec- tron donors. Evidence of cation solvation by tt81
electron donors, obtained in systematic extraction

experiments by Bucher and Diamond [442], should

also be reviewed briefly. In the experiments re- ferred to ~ 0.2 M HCIO4 (containing radioactive

HRe04 as tracer) was extracted from water into various organic diluents containing trilaurylamine

(~0.1 M), the principal objective being to deter- mine the role played by the diluent. Experimental

techniques included preparation of radioisotopic

Re and subsequent y-counting, forward and back

extractions, acid-base and Karl Fischer titrations,

freezing point measurements, and vapor pressure

measurements. Table 49 lists most of the diluents

Table 49. Comparison of dielectric constants and effectiveness of diluents used in extracting HC104 from water ^

Diluent LogX,(22°)

Triethylbenzene 2.26 4.38 Trimethylbenzene 2.28 4.52 Benzene 2.28 6.57 Tetrachloroethylene 2.30 4.52 Anisole 4.3.3 8.08 ChlorobenzenC' 5.62 7.62 o-Dichlorobenzene 9.93 8.18

^This is a slightly revised version of Table VI of J. J. Bucher

and R. M. Diamond, J. Phys. Chem. 69, 1565 (1965).

" Relative effectiveness is indicated by the magnitude of

Ki. the equilibrium constant for the following reaction (compare eq 50):

(C,2H25).3No,, + H:,-f ClOfac^ (C,2H2..),-,NH+-C104-„,g- "Aq" and "org" signify the aqueous and organic phases, respectively. As given in the reference cited. Reference:

[442] J. J. Bucher and R. M. Diamond, J. Phys. Chem. 69, 1565

(1965). used and their dielectric constants, and also indi- cates their comparative effectiveness in promoting

the extraction of HCIO4. The measure of effective- ness is the value of Ki (given here as its logarithm)

for the extraction corresponding to eq 50 (sec.

4.3.5c), where R = Ci2H25 and X- = C104-. The following are some of the conclusions

reached by Bucher and Diamond (by reasoning

that will not be restated here):

(1) Benzene solvates (Ci2H2.5)3NH+ by interaction

of its TT-electrons with the cationic proton, ren- dering it more stable in the organic phase.

(2) 5ym. -Trimethylbenzene (mesitylene) and

sym. -triethylbenzene interact similarly, but the extent of interaction, as expected, is reduced by

steric blocking by the substituent alkyl groups.

(3) Tetrachloroethylene, which has a single pair

of TT-electrons that are somewhat shielded by the

adjacent chlorine atoms, has an overall effective- ness very similar to that of triethylbenzene.

(4) Anisole is a more effective diluent on account

of being an n-electron donor as well ao a 7r-elec- tron donor.

(5) The relatively high effectiveness of chloro- benzene and o-dichlorobenzene stems from their

higher dielectric constants, which lead to "electrostatic solvation" of the ions — not to any specific chemical interaction with them.

A few other examples of cation "solvation"

wiU be encountered in the discussion of heteroconjugate catjons (sec. 4.4.5; see table 53). 4.4.3. Homoconjugate Cations

a. Studies using aprotic or dipolar aprotic sol- vents. Homoconjugate cation formation has been

studied in acetonitrile [444-446], benzene [5, 21],

o-dichlorobenzene [448, 449], and nitrobenzene

[24, 443, 447]. The evidence was obtained by three

types of measurement: (1) Absorption spectrophotometry in the visible region [5, 21]; (2) conductance

[24, 444, 448, 449]; and (3) potentiometry [445-447].

Miscellaneous examples are listed in table 50. Results of a comprehensive study of homoconjugate

cation formation are summarized in table 51. Homoconjugate cation formation was postulated

by Davis and Hetzer to account for color phenomena

observed in spectrophotometric studies of the

association of bromophthalein magenta E (usually

2.5X10-5M) with 1,3-diphenylguanidine [5] and with 1,3-di-o-tolylguanidine [21] in benzene.

On gradually increasing the concentration of the base they observed a rapid shift in color from

the yellow of the indicator to the magenta of

its 1:1 salt, followed by a slower shift toward

blue. The absorbance data were consistent with

the assumption of two slightly overlapping reactions:

(1) Association of tTie base (B) and acid (HA) in 1:1 ratio, yielding hydrogen-bonded ion pairs

(eq 44); and

(2) reaction of BHA with a second molecule of

base (eq 55): B + H-A^B-H+...A- (44) B-H+...A- + B^ (B-H+...B)A-. (55) 1,2,3-Triphenylguanidine showed similar behavior

[5] except that the ''primary'' (eq 44) and ""second- ary" (eq 55) reactions seemed to overlap more. It was thought that still other competing associations occur at high concentrations, but these were not

studied.

When 1,3-diphenylguanidine,

PhNH-C(=NH)NHPh,

was the reacting base the equilibrium constant (^bha) corresponding to eq 44 was found to be

2.2x10'^ (liter mole"*, 25°). Replacing one ortho

hydrogen in each phenyl group by methyl would be

expected to increase the basicity of the molecule,

whereas placing a phenyl group on the 2-nitrogen

should reduce the basicity. In agreement with

expectations, /^bha increased to 1.1 X 10^ when

diphenylguanidine was replaced by di-o-tolylguanidine, but dropped to 525 when it was replaced by

1,2,3-triphenylguanidine.

82
Table 50. Examples of homoconjugate cations (B-H+...B)

Base Solvent Anion Method of study"

Estimated

formation

constants

References

(PhNH)2C = NH (BPM-E)" VIS ^ 15.5 [5] CH3CN CH2CN- c "20 [444] CH:jCN C104- EMF None [446] (PhNH).C = NPh CeHfi (BPM-E) VIS ~ 22 [5] L'^J (o-MeC6H4NH)2C = NH CfiHg (BPM-E)-" VIS <^6.4 [21] Py PhNOi Pi- C "•^56 [24, 443] o-CgHjCL, Pi c < < 1 [448] o-CsHjCli Pi- c " 19,000 [448] CH3CN CIO7 EMF M [445, 446]

MeNHa; EtNHa PhNO^ CIO7 EMF'= d. h 10:<- 10^ [447] j-PrNH2 PhNO.f cio- EMF^^ d.h lO'l-lO" [447] MesNH; Et.NH PhNQf ClO; EMF« d.h lO-'-lO^ [447] Me,N, EtjN PhNQf C107 EMF'^ dt. 103-101 [447] Bu-)N Br~ Q [448] 0-C«H4Cl2 Pi C •^^25 [448] O-C6H4CI2 Pi C "560 [448] BU3NO o-CgH^CL. Pi- c " 1.9 X 107 [449]

^ Meaning of symbols: C, conductance; EMF, potentiometry; VIS, absorbance in the visible region. "(BPM-E) - signifies the anion of bromophthalein magenta E. This formation constant is for the reaction B-H + ... A- + B ^ (B-H+...B)A-; molar units, 25°. This formation constant corresponds to BH+ + B ^ BH + ...B (apparently molar units; 25° unless indicated otherwise). Calculated by KolthofT [443], using results of Witschonke and Kraus [24]. Traces of CH.jCOOH and (CH:iC0)20 vi^ere present. "The base, usually 0.001 M, was titrated with -0.034 M HCl.

" Experiments were done at ~ 20°. Aliphatic amines seemed to react slowly with PhN02.

References:

[5] M. M. Davis and H. B. Hetzer, J. Res. NBS 46, 496 (1951).

[21] M. M. Davis and H. B. Hetzer, J. Res. NBS 48, 381 (1952).

[24] C. R. Witschonke and C. A. Kraus, J. Am. Chem. Soc. 69, 2472 (1947). [443] I. M. KolthofT, D. St6cesoca, and T. S. Lee, J. Am. Chem. Soc. 75, 1834 (1953). [444] W. S. Muney and J. F. Coetzee, J. Phys. Chem. 66, 89 (1962). [445] J. F. Coetzee, G. R. Padmanabhan, and G. P. Cunningham, Talanta 11, 93 (1964). [446] J. F. Coetzee and G. R. Padmanabhan, J. Am. Chem. Soc. 87, 5005 (1964). [447] D. Feakins, W. A. Last, and R. A. Shaw, J. Chem. Soc, p. 2387 (1964). [448] E. K. Ralph, III, and W. R. Gilkerson, J. Am. Chem. Soc. 86, 4783 (1964). [449] W. R. Gilkerson and E. K. Ralph, III, J. Am. Chem. Soc. 87, 175 (1965). A large value for Kbha is indicative of a strong

interaction of B and HA. Hence, the "secondary"

equilibrium constants might be expected to vary

in the opposite sense (in the absence of strong steric

effects). This appears to be a correct prediction,

since the respective secondary equilibrium constants found for 1,3-di-o-tolylguanidine, 1,3-diphenylguanidine, and 1,2,3-triphenylguanidine are 6.4,

15.5, and ~ 22. Probably the first concrete evidence of homoconjugate cation formation was obtained by Witschonke and Kraus [24]. In an investigation already

referred to (for example, see table 35 and related

discussion) they found that the conductance of

pyridinium picrate in nitrobenzene can be augmented by adding surplus pyridine, although a

similar increase in conductance does not occur on

adding pyridine to a quaternary ammonium picrate

(for example, Me4NPi). Subsequently, Kolthofif and

coworkers [443] used data of Witschonke and Kraus

to calculate the equilibrium constant corresponding

to eq 56: BH+-f B^BH+...B. (56)

The value which they obtained is 56. In a stiU later calculation, Ralph and Gilkerson [448] obtained a

sHghtly different value (58). Gilkerson and Ralph have performed investigations parallehng those of Witschonke and Kraus,

83
but with o-dichlorobenzene as the solvent and Py,

n-BuijN, and /i-BusNO as the bases. In all cases, adding an excess of base to the picrate brought

about increased conductivity. They consider this

to be best explained by Witschonke and Kraus's

hypothesis of hydrogen bonding of BH^ to B

[448, 449].

Gilkerson and Ralph have suggested (LXXV) as a reasonable structure for (Bu3N)2HPi:

"3 H OC^H^CNO^)

(LXCT) The plane of the benzene ring in Pi~ is considered

to be perpendicular to the N-H...H plane, and

at least one butyl group on each molecule of amine

is believed to lack rotational freedom. Such a model, they think, permits almost as close approach

of cation and anion as in Bu3NH+..."OC6H2(N02)3.

The remaining examples listed in table 50 were deduced in potentiometric investigations, mainly

by Feakins and coworkers [447]. In their work the

solvent was nitrobenzene. When aliphatic amines were titrated with perchloric acid, inflections

appeared in the titration curves near the half-point.

The only reasonable explanation was BHB+ formation. Homoconjugate cation formation was inferred

for MeNH,,, EtNH2, i-PrNH., Me.NH, Et.NH,

MeijN and Et:jN, but was not detected for diphenylguanidine, piperidine, A^-methylmorpholine, di-A^- ethylaniline, or p-toluidine. Because of its mode

of preparation the perchloric acid solution con- tained traces of acetic acid and its anhydride. It should also be mentioned that the aliphatic amines

appeared to react slowly with the solvent. Feakins

and associates concluded that the amines listed form fairly stable homoconjugate cations in nitro- benzene, the estimated formation constant (eq 56)

in each case being about 10-^ to 10^. In their opinion,

a conjugate cation as unstable as PyH+Py would

not be detected in such experiments.

In comparing the formation constants listed in column 5 of table 50, it should be borne in mind

that some of these constants correspond to eq 55,

and others to eq 56. Moreover, all of the experiments were exploratory in nature, and in some cases

several competing equilibria had to be taken into account and simplifying assumptions had to be

made. We have already drawn some tentative conclusions with respect to the comparative behavior

of the three guanidine derivatives. Assuming the

other constants in table 50 to be approximately

correct, some additional tentative conclusions may

be drawn:

(1) PyH+...Py is much more stable in o-dichlorobenzene than in nitrobenzene.

(2) It is more stable than Bu3NH+...NBu;i in this solvent, but less stable than BU3NOH+...ONBU3.

(3) (Bu3NH+NBu3)Pi" appears to form more extensively in o-dichlorobenzene than does (PyH+Py)Pi~

(see eq 55). (Note: In reporting the equihbrium

constant for the reaction of PyH+...Pi- with Py

as being 1, we made an approximation; the actual

conclusion of Ralph and Gilkerson was that this equilibrium constant multiplied by [Py] is ^ 1, Py2H+ being the only complex formed to a significant extent in this system [448].)

By far the most comprehensive study of homoconjugate cation formation to date is a potentiometric investigation by Coetzee and Padmanabhan,

using acetonitrile as the solvent [445, 446], which

supplemented a conductimetric study by Muney

and Coetzee [444]. The solutions contained varying

amounts of free base together with a fixed concentration of its perchlorate or picrate. A wide assort- ment of amines was included, so as to determine

how their tendencies toward homoconjugation vary

with such factors as basic strength, steric hindrance,

and the number of cationic protons available for hydrogen bonding. 1,3-Diphenylguanidine and pyridine were also included, as well as some diamines.

Results obtained with monoamines are summarized

in table 51. The next to last column in table 51 contains values determined for Kbh+b (eq 56), and

the last column contains values for A^bh+B2 (eq 57).

BH+...B + B^BH+...B2. (57)

Formation of BH+ ...B2 is expected only when at least two cationic protons are available for hydrogen

bonding to B. The results given in table 51 clearly fulfill expectations. To begin with, the ^bh+b values for aliphatic

amines indicate a marked tendency for 1:1 bonding

to their conjugate acids. In contrast, pyridine, a much weaker base, exhibits less tendency toward

homoconjugation, and the stiU weaker bases,

aniline and p-toluidine, none at all. Secondly, it is apparent that homoconjugation

is impeded when bulky groups are attached to amine nitrogen. Of the tertiary aliphatic amines

studied, only MesN showed a measurable tendency

toward homoconjugation. Homoconjugation of sec- ondary aliphatic amines, also, is impeded noticeably

by substituents larger than Me, except for cyclic

amines (such as pyrrolidine, piperidine, and morpholine), where the substituent groups are "tied

back," leaving the amino nitrogen relatively

exposed. A cyclic tertiary diamine, triethylenediamine, was observed to form a 1:1 complex. (It would be interesting to study the behavior of the

cyclic tertiary monoamine, quinuclidine.) The

authors calculated [445] that a Kbh + b value near 30

(which seems typical for unhindered aliphatic

amines) corresponds to AC = — 2 kcal/mole.

As for diphenylguanidine, although it can bond

1:1 to its conjugate acid in benzene containing

bromophthalein magenta E (table 50), as well as

in acetonitrile without added acid (table 51), no

A^BH + B value was obtained in the potentiometric

studies.

84
Table 51. Approximate formation constants (25°) of some

homoconjugate cations in acetonitrile "

Base pKa in water

Anion

NH3 9.21 ClOj 11 2 MeNHs 10.62 ClOj 35 2 EtNH2 10.63 CIOt 25 2 n-PrNH2 10.53 ClOj 19 2 n-BuNHj 10.59 ClOj 26 2.5

t-BuNH. 10.45 Pi- 20 1 PhCH^NH. 9.34 Pi 15 HOCH.rH.NH, 9.50 Pi 24 1.5 PhNH2 4.58 CIOj; Pi- None

p-MeC6H4NH2 5.07 CIOj None

Me2NH 10.64 ClOj 31 0.5 Et.NH 10.98 ClOj 3 EtjNH 10.98 Pi- 1 n-Bu2NH 11.25 ClOj; Pi- None

i-Bu.NH 10.50 Pi- None

NH{CH2)3CH2

1

1

11.27 Pi 32 2

NH(CH2)jCH2

1

1

11.22 Pi 26

NH(CH2)20CH.CH2

.1 1

8.36 Pi 10

y.lO 0 EtaN 10.65 Pi- None

n-PisN 10.65 ClOj None

n-BujN 10.89 CIOj: Pi- None

i-AmsN Pi- None

(PhNH)2C = NH 10.00 ClOj None

(PhNH)2C = NH 10.00 CH2CN- ^20 Py 5.17 ClOj 4

^Formation constants (Kbh+b- ^bh + bJ were determined by

glass-electrode measurements in buffer solutions consisting of

a constant concentration (5x10"^ M) of the perchlorate or

picrate of the base and varying concentrations of free base

(usuaOy up to 0.1 M or higher). This table is an abridged and

rearranged version of Table I of J. F. Coetzee and G. R. Pad- manabhan, J. Am. Chem. Soc. 87, 5005 (1965). (Also see J. F. Coetzee, G. R. Padmanabhan, and G. P. Cunningham, Talanta 11, 93 (1964).)

" Kb„+b = [BHB+ ]/[B] [BH+] ; Kb„+b., = [BH-B2 )/[B] [BHB- J.

' Obtained in a conductometric study of diphenylguanidine in acetonitrile, by W. S. Muney and J. F. Coetzee, J. Phys. Chem.

66, 89 (1962), which indicated that at concentrations of the base > ~ 0.05 M the predominant ionization reaction is repre- sentable as: 2B + CH3CN ^ BH +B + CH2CN

References:

[444] W. S. Muney and J. F. Coetzee, J. Phys. Chem. 66, 89

(1962).

[445] J. F. Coetzee, G. R. Padmanabhan, and G. P. Cunningham,

Talanta 1 1 , 93 (1964).

[446] J. F. Coetzee and G. R. Padmanabhan, J. Am. Chem. Soc. 87, 5005 (1965).

Thirdly, values of /v^bh + B2 could be determined

for most of the primary aliphatic amines. As

expected, they are smaller than values for A^bh + b- Only two secondary aliphatic amines, dimethylamine and pyrrolidine, reacted measurably according

to eq 57. Piperidine formed only BH+B [445J. With

hydrazine, however, there was evidence of BH^B4

formation.

Coetzee and associates emphasized the signifi- cance of homoconjugation of bases in potentiometric

and conductimetric titrations, as well as in studies

of the relative strengths of bases [444-446].

They also emphasized that MeCN itself, although

a much weaker base than H2O, appears to solvate

ammonium ions to a significant extent. We shall see presently (sec. 4.4.4) that others have reached

the same conclusion.

b. Studies of solid salts of N-oxides. A number of recent studies have been concerned with homoconjugation of bases in their solid salts. Interest

in this topic began with work by Hadzi [390, 450],

who measured the IR spectrum of solid (PicO)2HBr

(where PicO signifies 2-picoline 1-oxide). On finding a resemblance to the spectra of potassium acid

maleate and certain other acid salts having homoconjugate anions (A...H...A)~, believed to contain a symmetrical hydrogen bond (see sec. 4.5.3a),

he suggested that the similarity can be explained

by postulating the presence of a symmetrical

homoconjugate cation in (PicO)2HBr, as indicated

by (LXXVI):

(PioO...H...OPic)"^ Br~

(LXXVI)

Inspired by Hadzi's discovery, several other investigators have measured the IR spectra of salts of

PicO with other monoprotic acids, as well as various

salts of related A^-oxides. Examples of such salts are listed in table 52. In the course of these studies, various A-oxides

(for instance, 2-, 3-, and 4-picoline 1-oxides, acridine

1-oxide) have been found to form two series of salts with monoprotic acids: (1) A ^'normaV series,

containing 1 equivalent of base to 1 of acid; and

(2) an ''abnormal'" series, containing 2 equivalents

of base to 1 of acid. In most normal salts of A-oxides

the cation (BO...H)+ is thought to be hydrogen

bonded to the anion, since the v{Oli) frequencies

have been found to vary with the proton affinity of

the anion (compare 4.3.5a). For example, Szafran

reported [454] that in normal salts of acridine

1-oxide with HCl, HBr, HI, and HCIO4 the ^OH)

band shifts increasingly toward lower frequencies

with increasing basicity of the anion — that is, the

t'(OH) shifts follow the order

Cl->Br->I->C104-.

In contrast, the t'(OH) band does not appear in the customary position (3700-1700 cm~') in the

spectrum of abnormal salts, and all IR spectra are

85
Table 52. Examples of normal and abnormal salts of pyridine 1 -oxide and related N-oxides

Base Normal salts Abnormal salts

Acid References Acid References

ryridine 1 -oxide [390]

T T \

¥-> [451] HBr [390] HSbFe [451] HNO3 [450] CCI3COOH [390]

2-Picoline 1-oxide [oVU, 40ZJ HAsFb " [451] HC1O4

[453] HPFfi [451]

2,6-Lutidine 1-oxide HCl [453] HI [453] HBr [453] HCIO4 [453] HSbCls [453]

Isoquinoline 1-oxide HCl [453] HBr"^ [453] HBr [453] HC104 [453] HSbCle [453]

2-Quinaldine 1-oxide [453] rlor [45oJ HBr [453] HI [453] HI [453] HCIO4 [453] CCI3COOH < [454] HSbCls [453]

Acridine 1-oxide HCl [454] HBr [454] HBr [454] HI [454] HI [454] rlL.lU4 (4o4J

Acridine 1-oxide HC104 [454] HBF4 [454] CC13COOH f [454] HSbCl« [454]

2,2'-Bipyridine 1,1 '-Dioxide ^ [455] HCl [455] HBr [455] HCIO4 [455] HSbCle [455]

^ Examined as solids, by IR spectroscopy except as indicated otherwise. In some cases deuterium analogs were also studied. Normal salts are defined [453] as those containing 1 equivalent of base to 1 equivalent

of a monoprotic acid, while abnormal salts are those containing 2 equivalents of base to

1 of acid. X-ray results [452] support the conclusions from IR data [390]. The corresponding salts of 3- and 4-picoline 1-oxide were also studied.

* Isolated as the monohydrate.

'This salt has an IR spectrum resembling that of abnormal salts, and is thought to have

the structure (BO...H...OB)+(A...H...A)-. "Even though 2,2'-bipyridine l,r-dioxide contains two basic groups, all of the salts isolated contained only one equivalent of acid. References:

[390] D. Hadzi, J. Chem. Soc, p. 5128 (1962). [450] D. Hadzi, Boll. Sci. Fac. Chim. Ind. Bologna 21, 23 (1963). [451] D. Cook, Chem. Ind. (London), p. 607 (1963). [452] H. H. MiUs and J. C. Speaicman. Proc. Chem. Soc, p. 216

(1963). [453] M. Szafran, Bull. Acad. Polon. Sci. Ser. Sci. Chim. 11,

111, 169, 497, 503 (1963). [454] M. Szafran, Bull. Acad. Polon. Sci. Ser. Sci. Chim. 12,

289,383(1964).

[455] M. Szafran, Spectrochim. Acta 19, 2132 (1963). 86
practically alike, whatever the anion present

(except, of course, for bands characteristic of anions)

[453]. Still further, replacing hydrogen by deuterium

(for example, in the abnormal iodide or perchlorate

of 2,6-lutidine 1 -oxide [453]) does not cause any

change in the IR spectrum. For such reasons, the

abnormal salts are believed to contain a homoconjugate cation analogous to (LXXVI) [390, 450-

455]. It is also believed that the homoconjugate

cation contains a symmetrical hydrogen bond.

Mills and Speakman, v^ho studied (PicO)2HBr by

x-ray diffraction at Hadzi's suggestion, agree with

this and other conclusions reached by him [452].

It should also be mentioned that the spectrum

of abnormal salts of 2,6-lutidine 1 -oxide was found

by Szafran to revert to the normal spectrum upon

solution of the salts in methyl sulfoxide [453]. From this he concluded that the homoconjugate

cation had been destroyed, by conversion to the

solvated cation, as indicated in eq 58:

(LutO...H...OLut)+ + Me2SO^

(LutOH...OSMe2)+ + LutO. (58)

He represents (LutOH...OSMe2)+ as being a resonant cation. Szafran concluded that the homoconjugate cation is stable toward chloroform.

The concept of homoconjugate cation formation can be seen to provide a reasonable explanation

for the existence of stable "abnormal" salts of

A^-oxides, and for their characteristic properties.

This explanation is further strengthened by Szafran's discovery [455] of the apparent inability of

2,2'-bipyridine l,r-dioxide (2,2'-dipyridyl l,r-dioxide) to form "normal" salts (see last item in table 52). Several structures are conceivable for the homoconjugate cation of this yV-oxide [455]; it should be

especially stable if it has a cyclic, intramolecularly

hydrogen-bonded structure. The postulate that the hydrogen bond in such

cations is symmetrical also seems reasonable, but

so far there does not appear to be any rigorous

proof of symmetry.

Quinaldine 1-oxide and acridine 1-oxide both

form stable 1 : 1 trichloroacetates, but the IR spectra

of these salts are abnormal [454]. In fact, they are

similar to the spectrum of a 1 : 1 mull of KH(CCl3-

C00)2 and (quinO)2HBr, as well as to spectra ob- tained from mulls in which the latter salt is replaced

by (quinO)2HC104 or a corresponding salt of

acridine 1-oxide. Szafran therefore thinks it likely that the structures of the 1 : 1 trichloroacetates are

as indicated by (LXXVII):

BO...H...OBj''" |^CCljC(=0)O...H...OC(=0)CCl2j

"

(LXXVII) These results with solid salts of A^-oxides are in harmony with the finding (see last item in table 50)

that w-BusNO forms an extremely stable complex

with its conjugate acid in o-dichlorobenzene.

4.4.4. MisceiianeouH Heteroconjugate Cations In section 4.4.1 we noted that /leteroconjugate

cations, B-H+...B', include the special cate- gory known as "solvated" cations. Examples of the

latter were discussed in section 4.2.2. This section

may be regarded as a continuation of that discussion.

The heteroconjugate cations now to be discussed

are listed in table 53, which emphasizes cases in which B' is a base with a measurable pKa value. Values of Kbh+b' (eq 59) are included where avail- able. Solvated cations in

BH+ + B'^BH+...B' (59)

which B' was EtOH, MeOH, MeCN, or H.O are also included to show analogies in behavior. The sol- vents principally used were benzene, two halo- genated aromatic hydrocarbons {chlorobenzene,

o-dichlorobenzene), and nitrobenzene. Only limited use could be made of the halogenated aliphatic

hydrocarbon, 1 ,2-dichloroethane, because of its reactivity with free aliphatic amines. It will be noted

that in most of the systems studied the anion

present was Pi~.

In all of the examples given the investigators were convinced that association of neutral molecules

with cations involved hydrogen bonding, from evi- dence similar to instances previously cited. It was

found, for example, in conductance investigations

that adding Py or MeCN to BusNHPi in o-dichlorobenzene greatly increased the conductivity, whereas

such additions to a solution of Bu4NPi had only a minor effect [448]. Furthermore, a comparison of the results cited

in table 50 with those in table 53 provides evidence

that cations BH+ which bond readily to their

conjugate bases (B) also bond readily to other

bases (B'). For example, the visual appearance and

absorption spectrum of the diphenylguanidinium

salt of bromophthalein magenta E in benzene solu- tions are altered similarly by ethanol, pyridine,

and an excess of diphenylguanidine [5]. In a differ- ent type of experiment, the conductance of tri-n- butylammonium picrate in o-dichlorobenzene was

found to be augmented by adding acetonitrile,

pyridine, 4-cyanopyridine, 4-picoline, or free tributylamine [448]. A third strong argument in support of the hydrogen bonding concept is provided by the comparative

values of Kbh"^ b (see table 53, 5th column)

determined in experiments in which three pyridinetype bases of known pKa values were added to n-BusNHPi in o-dichlorobenzene: a clear-cut parallelism between the strength of B' and the sta- bihty of BH+B' was revealed. Although the stability

of BusNH+NBus (table 50) is less than expected from

the basic strength of BuaN, this irregularity is at- tributable to steric hindrance. It is noteworthy that MeCN, though regarded as a far weaker base than

H2O, shows an appreciable tendency to associate

1 : 1 with BU3NH+.

87
Table 53. Examples of heteroconjugate cations (B—H+...B')''

B Solvent Anion B' " Estimated

'^BH + ...B'

References

(PhNH).C=NH CeHe (BPM-E)- " EtOH [5] (BPM-E)- " Py (5.17) [5] PhNOo Pi- Py (5.17) [24]

T> IVT rnIN(J2 n* — ri Py (5.17) * 6.2 [24, 448] n-Bu-,N PhCl Pi- Py (5.17) 820 [448] Pi- 4-MePy (6.02) 1830 [448]

n-BujN Br- Py (5.17) 1450 [448] 0-C(iH4Cl2 Pi- 4-CNPy (1.90) 40 [448]

r ) MeCJN 212 [448] Pi- Py (5.17) 1370 [448] Pi- 4-MePy (6.02) 3240 [448]

n-BujN C.2H4C12 ' Pi- MeOH 41 [448] MeaNO ^ PhN02 Pi- H.O [24] Pi- Py (5.17) [24] Pi Lt.iN (10.65) [24] Pi- CH.oNH [24] (11.22)

PhN(Me2)0 PhNOo Pi- Py (5.17) [24] n-BusNO 0-C6H4CI. Pi- Py (5.17) 1.2 X 10^ [449]

" In the first two examples tiie physical method used was absorbance in the visible region;

in all other cases it was conductance. pKn values from the literature are given in parentheses.

*^^BH+. .B' corresponds to BH+ + B' ^BH+...B' (25°); the values cited appear to be in molar

units.(BPM—E)- signifies the anion of bromophthalein magenta E.

* Calculated by Ralph and Gilkerson [448] from data of Witschonke and Kraus [24].

' Reaction of free amines with l,2-CiH4Cl> prevented their use as examples of B'.

^ Bases B' are Usted in the order of increasing effectiveness is promoting ion-pair dissociation (as indicated by increased conductivity).

References:

[5] M. M. Davis and H. B. Hetzer, J. Res. NBS 46, 496 (1951).

[24] C. R. Witschonke and C. A. Kraus, J. Am. Chem. Soc. 69, 2472 (1947).

[448] E. K. Ralph, III, and W. R. Gilkerson, J. Am. Chem. Soc. 86, 4783 (1964).

[449] W. R. Gilkerson and E. K. Ralph, III, J. Am. Chem. Soc. 87, 175 (1965).

Evidence of a similar kind came from Witschonke

and Kraus's conductance experiments, referred to earlier (see table 35), in which four different bases were added to MesNOHPi in nitrobenzene. The four

bases -HoO, Py, EtsN, and C^HioNH- were in- creasingly effective in the order listed in promoting

the ion pair dissociation of MesNOHPi, showing that effectiveness is proportional to proton affinity.

As yet, too few such studies have been performed

to permit many additional conclusions about homoand heteroconjugation of cations. However, the

following tentative conclusions seem justified by

the evidence given here: (1) /i-Bu.5N0H+ forms more

stable conjugate cations than n-BusN. (2) Both

Bu3NH+...Py and Bu,3NH+...4-]VIePy form more

extensively in o-dichlorobenzene than in chloro- benzene. (3) Bu:iNH + ...Py forms much less

extensively in nitrobenzene than in either of the

solvents just mentioned. It might also be thought

than Bu,3NH+ associates a little more extensively

with Py when the anion is Br~ than when it is Pi~

judging from a difference in values for Xbh+b'

(table 53); however, Ralph and Gilkerson do not regard the difference as exceeding the experimental

uncertainty [448].

4.5. Hydrogen Bonding of Anions to Neutral

Proton Donors

4.5.1. Introduction

We have just seen that homo- and heteroconjugate cations, B-H+...B and B-H+...B', exist very commonly in aqueous and nonaqueous systems.

Proton bonds between anions, A"...H-A or 88
A"...H-A', would be expected to form still more

readily. Experimental evidence of the correctness

of this view is much too extensive for complete

coverage here. However, numerous examples will be presented to illustrate the widespread occur- rence of homo- and heteroconjugate anions in aprotic

and dipolar aprotic solvents. Examples will also be

given of anion conjugation in solids. The principal subtopics, in the order of presentation, will be: (a) Homo- and heteroconjugation of 0H~ ions; (b) homoconjugation of aliphatic and

aromatic carboxylates, phenolates, halides, and

miscellaneous other types of anions; and

(c) heteroconjugation.

Illustrations will be given of the significance of

anion conjugation in some practical chemical

situations — for example, in performing conductimetric, photometric, and potentiometric titrations; interpreting IR spectra; and controlling the rates of synthetic reactions.

Conjugate anions, like conjugate cations, usually seem to consist of 1:1 aggregates. It should be

emphasized, however, that some anions form more

complex aggregates when environmental conditions

are favorable.

4.5.2. Homo- and Heteroconjugation of OH~ Ions In work briefly reviewed in section 4.4.2a [430—

435], Diamond and associates concluded that the

hydronium ion tends to be extracted into aprotic

solvents from aqueous solutions of acids either as the homoconjugate cation, H30+(H20)3, or as a heteroconjugate cation. Agarwal and Diamond

[456] believe that the OH" ion parallels H3O+ in its hydrogen-bonding ability. This conclusion is based

on experiments in which a tetraalkylammonium

hydroxide was extracted from water into benzene or nitrobenzene containing an alcohol, with the hydroxide at constant concentration and the alcohol at varying concentrations, or vice versa. Extraction

into benzene appeared to proceed according to eq 60,

with triply alcoholated ion pairs as the extracted

species, while with nitrobenzene the results were

similar except that in this solvent of much higher

dielectric constant the solvated hydroxide is present

as separate ions (eq 61).

R4N+<, + OH-q + 3R'OH„r, ^ R4N+ • "OH • SR'OHor, (60)

R4N:q + OH;o + 3R'OHn

-R4N+K + 0H- -SR'OHo,,. (61)

The extraction experiments are summarized in table 54. The suggested structure for triply alcoholated

OH- ions is indicated by (LXXVIII):

Table 54. Summary of extractions of qauternary ammonium

hydroxides (R,|NOH) from water into alcohol iR'OH} — contain- ing benzene or nitrobenzene

Solvent

Nitro- benzene

R'

Benzyl or decyl

Benzyl or 4- methyl-cyclohexyl.

Pentyl or hexyl

Propyl or

butyl

Composition of extracted species

R,N+ OH-(HOR'),

R.,N+ plus

0H-(H0R')3

" From B. R. Agarwal and R. M. Diamond, J. Phys. Chem.67,

2785 (1963) [456]. R'OH_

R'OH...b-H

r'oh'

(LXXVIII)

It is also believed that H2O molecules may form

bridges between OH" ions and R'OH molecules, as indicated in (LXXIX):

H r'oh...6h

H R'OH...iH...O-H

R'OH. .9HH

(LXXIX)

Structures (LXXVIII) and (LXXIX) are analogous to (LXXII) and (LXXIII), respectively.

From these results Agarwal and Diamond think

it probable that in aqueous solutions the OH" ion

holds a primary shell of three H2O molecules, by

very strong hydrogen bonds,- thus agreeing with

conclusions of other investigators.

4.5.3. Homoconjugate Aliphatic and Aromatic

Carboxylate Ions

a. Solid acid carboxylates. Ammonium, potassium, and sodium acid carboxylates have been

prepared and studied as solids for well over a century. For example, they are said [457] to have

been referred to in textbooks by Berzelius ("Lehrbuch der Chemie," 1847) and Gmelin ("Handbook

of Chemistry," 4th edition, 1853). By 1914, according to Pfeififer's estimate [458] some 60 to 70 acid

salts of aliphatic and aromatic carboxyHc acids

had been described.

Pfeififer himself reported the preparation of both

normal and acid salts of pyridine with 17 mono- or dicarboxyhc acids. The general method for preparing the normal salts consisted of dissolving the

acid in pyridine and then evaporating off all or part

of the excess of pyridine. Normal salts were con- verted into the acid salts on crystallization from

alcohol. Pfeiffer's work predated recognition of

89
the general importance of hydrogen bonding.

Nevertheless, reasoning from the well-known

tendency for RCOOH to form dimers, for which

he postulated the general structure (LXXX) (or, alternatively, the cyclic structure accepted nowadays), he believed the pyridinium acid salts to be

representable as (LXXXI):

R.C:O...HO.C.R „ OH 0 Py,HO-C:0. .,HO.C(=0)R

(LXXS) (LXXXI)

Smith and Speakman [457] investigated the acid

potassium salts of benzoic and phenylacetic acids

as possible pH standards. Since the late 1940's

solid acid carboxylates have been studied extensively by infrared and x-ray spectroscopy, and

occasionally by newer techniques, mainly to as- certain whether or not the RCOO-...HOCOR

hydrogen bond in a given crystalline salt is symmetrical. The principal acid carboxylates so studied

are listed in table 55, together with a few examples

from Pfeiffer's work. Although strong indications

of homoconjugation of anions in solid species do

Table 55. Examples of carboxylate homoconjugation, RC(=0)0~...HOC(=0)R, in solids"

Ciirboxylic acid Cation present Type of measurement ^ References

Acetic Na+ X-ray [462, 466] Benzoic K+ X-ray; IR; NMR [461, 463] 4-Chlorobenzoic NH+, K+, or Rb+ X-ray [468] Cinnamic NH; or K+ X-ray [467] Dibromohydrocinnamic PyH+ Chem [458]

Glycolic K+ or Rb+ X-ray [469] 2-Nitrobenzoic PyH+, K+, or Rb + Chem; x-ray [458, 465] 3-Nitrobenzoic PyH+ Chem [458] 4-Nitrobenzoic PyH+, K+. Na+, or Rb + Chem; x-ray; IR; NMR [458, 465, 463]

Phenylacetic or Rb+ X-ray; IR; ND; NMR [459, 460, 463, 464] Trichloroacetic (BO...H...OB)+ IR [454] Trifluoroacetic Cs+, K+. or Rb+ X-ray [470]

" In part of the examples cited a symmetrical hydrogen bond is believed by some authors to be

present. ''Meaning of symbols: Chem, chemical analysis; IR, infrared absorption spectroscopy; ND,

neutron diffraction; NMR, nuclear magnetic resonance spectroscopy; x-ray, x-ray diffraction., See table 52, footnote f.

References:

[454] M. Szafran, Bull. Acad. Polon. Sci. Ser. Sci. Chim. 12, 289, 383 (1964). [458] P. Pfeiffer, Chem. Ber. 47, 1580 (1914). [459] J. C. Speakman, J. Chem. Soc. p. 3357 (1949). [460] M. Davies and W. J. OrviUe Thomas, J. Chem. Soc, p. 2858 (1951). [461] J. M. Skinner, G. M. D. Stewart, and J. C. Speakman, J. Chem. Soc, p. 180 (1954). [462] J. C. Speakman, Proc. Chem. Soc, p. 316 (1959). [463] R. Bhnc and D. Hadzi, Spectrochim. Acta 16, 853 (1960). [464] G. E. Bacon and N. A. Curry, Acta Cryst. 13, 717 (1960). [465] H. N. Shrivastava and J. C. Speakman, J. Chem. Soc, p. 1151 (1961). [466] J. C. Speakman and H. H. Mills, J. Chem. Soc, p. 1164 (1961).. [467] R. F. Bryan, H. H. Mills, and J. C. Speakman, J. Chem. Soc, p. 4350 (1963). [468] H. H. Mills and J. C. Speakman, J. Chem. Soc, p. 4355 (1963). [469] L. Golic and J. C. Speakman, J. Chem. Soc, p. 2521 (1965). [470] L. Gohc and J. C. Speakman, J. Chem. Soc, p. 2530 (1965).

not necessarily signify that RCOO~...HOCOR

bonds will be stable enough to persist in aprotic

and dipolar aprotic organic solvents, it will be

seen presently that such bonds are generally

very stable. In fact, at least one such bond,

PhCOO~...HOCOPh, appears to persist even in aqueous solutions [471, 472, 472a].

Not all of the carboxylic acids tested by Pfeiffer yielded stable pyridinium salts at ordinary temperatures. For example, o-bromobenzoic acid,

cinnamic acid, and the isomeric toluic acids were

recovered unchanged from pyridine, and the pyridinium salt of benzoic acid was recovered from ethanol

as the free acid. The products clearly depend in

90
large measure on the extent of competing hydrogen

bonding reactions, such as chelation and bonding

to the solvent. The reader is referred to original papers for conclusions with respect to the symmetry of hydrogen bonds in salts listed in table 55.

b. Homoconjugation of carboxylates in solutions.

Experimental evidence of homoconjugation of a carboxylate ion in solution seems to have been

obtained first for the benzoate ion, in water. In the

early 1930's, the formation of anions corresponding

to double molecules of benzoic acid was postulated

to account for the "salting in" of benzoic acid by

sodium benzoate [471, 472]. The same conclusion was reached in a recent determination of the dis- tribution of benzoic acid between water and toluene

in the presence of sodium or potassium benzoate

[472a], in which the equilibrium constant for homoconjugation (eq 62) was found to be 1.0 (molar

units, 25°). CeHsCOO+CeHsCOOH

^ChH5COO-...HOCOCbH5. (62)

This value is fairly close to an earlier estimate

(0.45 to 0.77) of Kokhoff and Bosch [472]. Indirect evidence of acid carboxylate formation

in benzene was provided by cryoscopic data for mixtures of representative tertiary aliphatic amines

with acetic or myristic acid in benzene. Two species -R3N(R'COOH)2 and R3N(R'COOH)4-were

beheved to be formed [473]. We have already had occasion (sec. 4.3.5b) to discuss direct evidence, supplied by infrared

spectroscopy, that CH3C00"...H0C0CH.i is a product of the half-neutralization of acetic acid

by EtsN, Et2NH, or n-BuNH2 in carbon tetra- chloride or chloroform [392-394], as well as evi- dence that the acid carboxylates are intermediate

products in the neutrahzation of CCI3COOH and

CFsCOOH by pyridine in chloroform [395]. Still another infrared study was concerned with the

acid potassium salts of 4-chloro- and 4-nitro- benzoic acids in methyl sulfoxide [477]. Furthermore, in the titration of benzyl-A',A'-dimethylamine with CCI3COOH in benzene by the

differential vapor pressure (DVP) method [352] the shape of the titration curve disclosed that formation of the acid carboxylate had occurred

(see sec. 4.3.3b).

The above and additional examples of carboxylate

homoconjugation are hsted in table 56 or table 57. Table 56 is concerned with aliphatic and substi- tuted ahphatic acids, whereas table 57 deals with

benzoic and related acids. In 1947 Maryott pubhshed an account of experimental conductimetric titrations of acids (picric,

d/-camphorsulfonic, and trichloroacetic) and ali- phatic amines (triethyl-, di-/i-butyl, and n-heptyl-)

in three solvents, namely, benzene, dioxane, and

benzene containing 1 percent of methanol [341].

Both "forward" and "reverse" titrations were

performed, but we shaU concern ourselves only

with forward titrations, which by definition were

those in which the amine (~ 0.2 M) was added to the acid (~ 0.01 M). Representative titration curves

are presented in figures 19 to 21. Acids and amines of the types used would be

expected to have negligibly low conductances in dilute aprotic solvents like benzene. However,

as preceding discussion has shown (see tables 32

and 33 and context), partly substituted ammonium

salts (e.g., picrates, chlorides) are measurably

conducting in such media, even though only a tiny

fraction of the salt (a millionth of a percent or so)

is dissociated into free ions. Consequently, in the

1 : 1 reaction of an acid with an amine the conduct- ance could be expected to rise gradually until the

end-point and remain approximately constant

thereafter. Examples of this simple pattern of

behavior were obtained in titrating picric acid with

ra-Bu2NH or Et3N (see fig. 19). It will be noted that

-10 -

05 10 05 1.0

•BASE

Figure 19. Conductance titrations in benzene {solid lines), dioxane {dotted lines), and benzene containing 1 percent of methanol {dot-dash lines); ~ 0.2 M di-n-butylamine or tri- ethylamine was added to ~ O.OI M picric acid.

Adapted by permission from figures 2 and 3 of A. A. Maryott, J. Res. NBS 38, 527

(1947). " (Ref. No. 341.)

the titration curves in benzene and dioxane were

almost identical when dibutylamine was the titrant,

and that the curve for benzene-methanol had practically the same shape, although the conductance

values were higher throughout the titration. It is not certain what caused drifts in the conductance

during neutraUzation of picric acid with triethylamine in benzene or dioxane [341], but air was not excluded from the solutions and it therefore seems

possible that the anomalies resulted from the

presence of a httle EtsNO [31, 403].

The curves for conductance titrations of tri- chloroacetic acid in benzene are strikingly different

280-613 OL-68—

91
Table 56. Examples of 1 : homoconjugate anion formation from RCOOH, deduced in studies of

solutions ^

Solvent Identity of RBase or cation added Method *> References

Acetone CH3 R4NOH in Py EMF [475] CH2CI Et3N C [474] CHCI2 EtsN or ra-BusN C [474] Ft N fir n Rii N c

Acetonitrile CH2CI EtsN c [474] CHCI2 EtsN c [474] CCI3 EtsN c [474]

l^ri3 TRY v^ri 1 1*'j1 CCI3 R3N, BU2NH, or HpNH2 C, VPL [341, 352]

131127 R N PRY

C6H5OCH2 TDPNB VIS [482] C6H5C=C TDPNB ' VIS [482]

»^ri3 Ft N Ft NH nr RiiNH TR

CH3 Ft M Ft MtJ R..MtJ HitslN, il,t2l>rl, or t>ulNrl2 TRIK

CCi3 Py IR [395]

*^i"3 fy TRIn. [6\fi>\

Dioxane CCI3 EtjN or BU2NH C [341]

.C6H5CH2 K+ IR [477]

Toluene H BU4NOH in i-PrOH EMF [476] CH3 BU4NOH in t-PrOH EMF [476] CC13 BU4NOH in i-PrOH EMF [476]

^ RCOOH signifies an aliphatic or substituted aliphatic acid. See table 57 for examples of ArCOO ...HOCOAr. Meaning of symbols: C, conductance; CRY, cryoscopy; VPL, vapor pressure lowering; EMF,

IR, VIS, usual meanings.

TDPNB symbolizes the indicator base "p-Tolyl Di-n-propyl Nile Blue" (LXXXIV).

References:

[341] A. A. Maryott, J. Res. NBS 38, 527 (1947). [352] S. Bruckenstein and A. Saito, J. Am. Chem. Soc. 87, 698 (1965). [392] G. M. Barrow and E. A. Yerger, J. Am. Chem. Soc. 76, 5211 (1954).

[393] E. A. Yerger and G. M. Barrow, J. Am. Chem. Soc. 77, 4474 (1955).

[394] E. A. Yerger and G. M. Barrow, J. Am. Chem. Soc. 77, 6206 (1955). [395] G. M. Barrow, J. Am. Chem. Soc. 78, 5802 (1956). [473] S. Kaufman and C. R. Singleterry, J. Phys. Chem. 56, 604 (1952). [474] P. J. R. Bryant and A. W. H. Wardrop, J. Chem. Soc, p. 895 (1957). [475] H. B. van der Heijde, Anal. Chim. Acta 16, 392 (1957). [476] G. A. Harlow and D. B. Bruss, Anal. Chem. 30, 1833 (1958). [477] D. Hadzi and A. Novak, Spectrochim. Acta 18, 1059 (1962). [482j M. M. Davis and H. B. Hetzer, Anal. Chem. 38 , 451 (1966).

92
Table 57. Examples of 1 : 1 homoconjugate anion formation from ArCOOH,

deduced in studies of solutions'*

Solvent Substituent

in CeHs

Base or cation added Method References

Acetone 2-NO2 Et^N c [474] 3-NO2 Et,N c [474] 4-NO2 Et,N c [474] 3,5-(N02)2 Et,N c [474]

Acetonitrile None Anion-exchange resin DR [478] None Et4N+ EMF [481] 4-CH, Anion-exchange resin DR [478] 3-Br Et,N+ EMF [481] 3-Cl Anion-exchange resin DR [478] 2-OH BU2NH, BU3N, Na+, or C, S, EMF [407, 481] Et4N+ 4-OH Et,N+ EMF [481] 4-OCH3 Anion-exchange resin DR [478] 4-NO2 Anion-exchange resin DR [478] 4-NO2 Et4N+ EMF [481] 3,5-(N02)2 R:,N, R4N+, K+, Py-type C, S, EMF [479-481]

Benzene None R:,N, R2NH, or RNH2 CRY, P [347] 2-OH TDPNB " VIS [4821 3-NO2 TDPNB VIS [482]

Methyl sulfoxide 4-Cl K+ IR [477] 4-NO2 K+ IR [477]

Water None Na+ S [471, 472] None K+ or Na+ D [472a]

" ArCOOH signifies an aromatic or substituted aromatic acid. See table 56 for examples

of RCOO-...HOCOR.

''Meaning of symbols: DR, distribution of ArCOOH between its salt with an anion- exchange resin and the solvent; P, dielectric polarization; S, solubility of ArCOOH in the presence of its Na or K salt; other symbols as explained in footnote b of table 56. See table 56, footnote c.

References:

[347] E. N. Gur'yanova and I. G. Beskina, J. Gen. Chem. USSR Eng. Transl. 33, 914

(1963). [407] J. F. Coetzee and G. P. Cunningham, J. Am. Chem. Soc. 87, 2534 (1965). [471] E. Larsson, Z. Physik. Chem. Leipzig 153 , 466 (1931). [472] I. M. Kolthoff and W. Bosch, J. Phys. Chem. 36, 1685 (1932). [474] P. J. R. Bryant and A. W. H. Wardrop, J. Chem. Soc, p. 895 (1957). [477] D. Hadzi and A. Novak, Spectrochim. Acta 18, 1059 (1962). [478] J. E. Gordon, J. Phys. Chem. 67, 19 (1963). [479] I. M. Kohhoff and M. K. Chantooni, Jr., J. Am. Chem. Soc. 85 , 426 U963).

[480] I. M. Kohhoff and M. K. Chantooni, Jr., J. Am. Chem. Soc. 87, 1004 (1965). [481] I. M. Kohhoff and M. K. Chantooni, Jr., J. Phys. Chem. 70, 856 (1966). [482] M. M. Davis and H. B. Hetzer, Anal. Chem. 38, 451 (1966). [472a] A. Brandstrom, Acta Chem. Scand. 20, 1335 (1966).

93
from those for picric acid (see fig. 20, solid lines).

After an initial steep rise, which reaches a maximum

(-)

I 1 BASE=n-HpNH2

/"

/A

1

^BASE/ ^-CCIjCOOH

Figure 20. Conductance titrations in benzene (solid lines), dioxane (dotted lines), and benzene containing 1 percent of methanol (dot-dash lines). ~ 0.2 M n-Heptylamine, di-n-butyl- amine, or triethylamine was added to — 0.01 M trichloroacetic

acid.

Adapted by permission from figures 4 to 6 of A. A. Maryotl, J. Res. NBS 38, 527

(1947). (Ref. No. 341.)

a little beyond half neutralization of the acid, the

conductance drops sharply until the end-point,

after which there is a more gradual rise or a ten- dency to level off. The maximum is especially

pronounced in the curve for titration with triethylamine in benzene.

Analogous, though less pronounced, maxima

were also obtained with dioxane as the solvent

{dotted lines), and even in the benzene-methanol

mixture {dot-dash lines), some evidence of the midtitration maximum remained.

From the abnormally high conductance of solu- tions containing free acid as well as salt, Maryott

concluded that an acid salt forms as an intermediate

product in such titrations (eq 63):

BH+...A- + H-A^BH+...A-...H-A. (63)

Obviously, the cation-anion interaction will be

weaker in the acid salt than in BH+...A", thus

explaining the marked difference in conductance

of the two salt species. One way of representing

the acid anion is indicated by (LXXXII):

// C1,C-C

3 \

0)-CCl3

O-H...0

(LXXXII) However, Maryott pointed out that an acid carboxylate ion must have exceptional stability to exist in dioxane, a medium in which carboxylic acid dimers

do not seem to be stable, and he prefers the alter- native resonant structure indicated by (LXXXIII),

in which the proton is assumed to occupy a posi- tion such that it is shared by all four oxygen atoms,

which are located tetrahedrally around it.

Cl^C-C^

(LXXXIII)

It is apparent from figure 20 that substitution of

BU2NH for EtaN results in smaller increases in conductance during the titration of trichloroacetic

acid in benzene or dioxane. This finding is in har- mony with various results, cited in section 4.3,

which indicated that dimers of secondary am- monium carboxylates are quite stable in aprotic

media. Similar indications of salt dimerization

are apparent in the conductance curve obtained

by titrating trichloroacetic acid with n-HpNH^ in benzene (fig. 20) and comparative results obtained on titrating camphorsulfonic acid with EtsN and

BU2NH in dioxane (fig. 21).

-12

HA = dl-CAMPHORSULFONIC ACID

_j I i_ 0-5 1-0 0-5 1-0

"BASE

Figure 21. Conductance titrations in dioxane: —0.02 M Di- n-butylamine (see curves marked "2") or triethylamine (curves marked "3") ivas added to ~ 0.01 M trichloroacetic acid or dl-camphorsulfonic acid.

Adapted by permission from figures 4 to 6 of A. A. Maryott, J. Res. NBS 38, 527

(1947). (Ref. No. 341.)

These exploratory studies not only demonstrated

the feasibility of performing conductance titrations

in solvents of very low dielectric constant, but also showed that homoconjugation of a carboxylate or sulfonate ion is valuable in foreseeing an approaching end-point and locating it precisely. Furthermore,

they have inspired other systematic investigations

of conductance titrations of carboxylic acids in such solvents as acetone [474] and acetonitrile

[474, 407, 479-481].

In a particularly noteworthy investigation [474]

Bryant and Wardrop showed that similar sharply

defined mid-titration maxima are obtained in two

aprotic solvents of much higher dielectric constant than benzene, namely, acetone and acetonitrile. Using acetone as solvent for the most part, and a

tertiary amine (EtaN or n-BusN) as the titrant, they

94
obtained pronounced conductance maxima for the

following carboxylic acids:

CCI3COOH, CHCI2COOH, CH2CICOOH;

2-, 3-, and 4-N02C«H4COOH;

3,5-(N02)2C6H3COOH.

In agreement with Maryott, they attributed the

maxima to homoconjugation of anions. The maxima

disappear in ethanol, and the conductance titration curves then have the same shape as for picric acid. In further agreement with Maryott, the authors

concluded that homoconjugate carboxylate ions

must be exceptionally stable, since they are formed

in solvents in which the carboxylic acids are known

to be monomeric.

Although results obtained by titrations in acetone

and acetonitrile were similar, the conductivities were higher in the latter solvent. Acetic acid proved to be too weak for satisfactory

titration. Some results with nitrophenols will be

referred to in section 4.5.4. At about the time Bryant and Wardrop's work

was published it was also shown that reproducible

potential rises sometimes occur in the region of

hcdf neutrahzation in potentiometric titration curves

for carboxylic acids and other acidic compounds,

in solvents like acetone and toluene (see table 56). Such mid-titration inflections are believed to result from anion homoconjugation [475, 476].

Anion homoconjugation in acetonitrile has been

studied systematically by Kolthoff and Chantooni

[479-481] and Coetzee and Cunningham [407],

employing conductance titrimetry and in some cases

solubility measurements and potentiometry as

well. Using the combined data and making certain

reasonable assumptions, they were able to estimate

values of equilibrium constants for various competing reactions, such as: BHA;^BH+ + A- (17)

BHA^B + HA (18)

BH+ + B^BH+B (56)

BH+AHA- ^ BH+ + AHA- (64)

A- + HA^A-...HA. (65)

In this way they arrived at approximate values for the equilibrium constant /^a-...ha (eq 65). This

constant has been variously referred to as a "homoconjugation," "formation," or "stabihty" constant. They also applied the overall results in calculating

conductimetric titration curves, which agreed well

with actual titration curves. In a related type of study, Gordon determined the

comparative stabilities of the homoconjugate

anions of benzoic and four substituted benzoic

acids by measuring the distribution of each acid

between its salt with an anion-exchange resin and

the solvent acetonitrile (eq 66) [478]

ArCOOH (in MeCN)+ ArCOO-Res+

(ArCOO. . .HOCOAr), Res+. (66)

He found that his log K values are linearly related

to the pKa values of the acids within the experimental uncertainty:

log K = 3.48 - 0.200 j)Ka- (67) Gordon's five equilibrium constants are given'

in table 58 (column 6), along with the /^a-...h-a

values obtained by Kolthoff and Chantooni and

Coetzee and Cunningham; the pKa values of the

acids are included (column 2). It is appropriate not only to intercompare these values but also to compare them with corresponding dimerization constants (see table 24 and context). The following

tentative conclusions seern warranted:

(1) The stronger an acid of type ArCOOH, the more stable its homoconjugate anion.

(2) The statement just made has to be modified

for acids in which there is strong intramolecular

hydrogen bonding. For example, anion homoconjugation was found to be less extensive for 2- hydroxybenzoic (salicylic) acid than for three much

weaker acids, and was too small to measure in the case of 2,6-dihydroxybenzoic (y-resorcyhc) acid.

(3) Stabihties of homoconjugate anions of the

type ArCOO"... HOCOAr (in MeCN) vary inversely

with the stabilities of the corresponding acid dimers,

(ArC00H)2, in benzene.

(4) Conjugate anions of the type

ArCOO... HOCOAr

are considerably more stable than homoconjugate

cations of the ammonium type (compare table 51). Our final examples of apparent carboxylate

homoconjugation are from spectrophotometric

titrations of an indicator base known as ''p-Tolyl di- n-propyl Nile Blue" ("TDPNB") with several

types of acids in benzene [482]. The acids included

o-chlorobenzoic, m-nitrobenzoic, phenoxyacefic,

phenylpropiolic, salicyUc, and trichloroacetic

acids. The structure of the indicator base is repre- sented by (LXXXIV). This indicator is related to

(LXXXIV)

Nile Blue A anhydro-base, and its systematic name is "5-p-tolylimino-9-(di-n-propylamino)-benzo-

[a]phenoxazine" [482, 483]. The proton-acceptor

atom is considered to be the nitrogen atom at- tached to carbon No. 5. Examples of spectrophotometric titration curves

are given in figure 22, A and B. It will be seen that 95
Table 58. Comparative honwconjugation constants of some aromatic carhoxylate

ions in acetonitrile "

Carboxylic acid pKa in water EMF" References

4-Hydroxybenzoic 4.58 4-Methoxybenzoic 4.47

4-Toluic Benzoic 4.20 3-Chlorobenzoic 3.83 3-Bromobenzoic 3.81 4-Nitrobenzoic 3.43

Salicylic 3.00

3,5-Dinitrobenzoic 2.82 2,6-Dihydroxybenzoic 1.0

'800-1500 2,000

17,000 Very small

1.000

4,000

6,000

7,000 2,000

10,000

383

409

427

528

609

[481]

[478]

[478]

[481. 478]

[478]

[481]

[481, 478]

[407, 481]

[481, 479]

[407]

^ At 25°. Equation 65 applies except in method DR; that is. A" + HA ^ A". ..HA. "The acids were titrated conductimetrically with Bu.^N, BusNH, and re-BuNH2. Calculated from effects of added ArCOOH on the solubility of ArCOONa or ArCOOK.

" From EMF data for mixtures of ArCOOH and ArC00NEt4. This method is con- sidered more reliable than the solubihty method [481].

" In this method the equilibrium constant is for distribution of ArCOOH between

acetonitrile and its salt with an anion-exchange resin. That is: ArCOOH(in MeCN) + ArCOOrRes* ^ (ArCOO-...HOCOAr), Res+

'Values of /^'a ...HA varied with the titrant and the concentration of HA: Two

titrations with BujNH gave 800 and 900. and two with Bu.jN gave 1500 and 1000.

References:

[407] J. F. Coetzee and G. P. Cunningham, J. Am. Chem. Soc. 87, 2534 (1965). [478] J. E. Gordon, J. Phys. Chem 67, 19 (1963). [479] I. M. Kolthoflf and M. K. Chantooni. Jr., J. Am. Chem. Soc. 85, 426 (1963). [481] I. M. Kolthoff and M. K. Chantooni, Jr., J. Phys. Chem. 70, 856 (1966). some curves (such as that for trichloroacetic acid

in fig. 22A) show a continuous increase in absorb- ance at the beginning, followed by a gradual level- ing off. Such "simple" titration curves are obtained

when B and HA associate in 1:1 ratio (eq 44): B + H-A^B-H+...A-. (44)

The acids in this group are all strong. In contrast,

the titration curves obtained with moderately

strong carboxylic acids — for example, phenoxyacetic acid— iend to have an S-shape (see fig. 22B).

(The S-shape would be much more apparent, of course, if the horizontal scale in fig. 22B were not

so highly compressed.) Although the investigations were not continued beyond exploratory stages, it seems probable that such S-shaped curves result when there is overlapping production of BHA and

B(HA)2, with the former product predominating at very low acid concentrations, and the latter, at high acid concentrations [482]. For B-H+...A~ to be formed to an important

extent, a carboxylic acid must be reacting as the

monomer, and B must be a much better proton

acceptor than A". When A" and B are equally

good proton acceptors or A" is the better acceptor, and HA forms a stable dimer, formation of

B(HA)2 can be predicted. Probably B(HA)2 is

sometimes formed by direct interaction of B and

(HA)2 (eq 68), as postulated by Kaufman and

Singleterry when interpreting cryoscopic data for amine-carboxylic acid systems in benzene; a fur- ther association of B(HA)2 with (HA)2 was also thought to occur (eq 69) [473]:

B+ (HA)2^B(HA)2 (68) B(HA)2+ (HA)2^B(HA)4. (69) On the other hand, when the part of the carboxylic

acid is present as monomer, or as monomer loosely

bound to solvent molecules, eq 63 may be a preferable formulation:

BH^..A- + H-A^BH+...A-...H-A. (63)

While emphasizing homoconjugation in 1:1 ratio as the most frequently observed type, we

should not ignore occasional indications of conjugation in other ratios — for instance, there have

been reports of adsorption of carboxylic acids in excessive amounts on ion-exchange resins. No detailed structures can be offered to explain such

behavior. However, it should be possible for extra molecules of R'-COOH to hydrogen bond to a 96
T—I—I—I—I—I—r—1—I—I—I—I—

CONCN, OF ACID. EOUIVS.

Figure 22. Spectrophotometric titration curves obtained on

titrating tolyl dipropyl Nile Blue base with some strong {A)

and moderately strong {B) acids in benzene.

Reproduced from figures 9 and 10 of M. M. Davis and H. B. Hetzer, Anal. Chem.

38, 451TI966). (Ref. No. 452.)

carboxylate, for example, R'-COO~ • to yield aggregates like those suggested by (LXXXV)

and (LXXXVI):

R' - + I ,0 . NR. /O...H-0-C=0... / /•' R'-C R'-C'"."'"NR, \ \.

'

^0 . . .H-0-C=0 . . .H-0-C=0 ... ^0 . . .H-0-C=0 . .

R' R' R'

(LXKXV) (LXXXVI) From the examples listed in tables 55 to 58 it is evident that carboxylate homoconjugates are

stable enough to form under widely varying conditions. As an illustration, the homoconjugate of

CeHsCOO" forms in solvents as different as ben- zene and water and in the absence of any solvent. Furthermore, the homoconjugates form in the

presence of varied types of cations: R4N+, R3NH+,

R2NH^, PyH+, K+, Na+, and others. It was pointed

out in section 4.4.2a that Diamond and associates

used perchloric acid as the source of protons in studies of homo- and heteroconjugation of the

hydronium ion [430-435], their reason being that

this acid yields an anion with exceptionally low

proton affinity. Analogously, should one wish to maximize anion homoconjugation, the cation present ought to be one with minimal affinity for proton

acceptors; that is, it should belong to type R4N+

rather than to a type of cation (R3NH+, etc.) whose

interaction with anions is strengthened by hydrogen bonding, and K+ should be used if feasible,

in preference to Na+.

4.5.4. Homoconjugate Phenolate Ions We have seen that Maryott's conductance ti- tration curves for picric acid in benzene and di- oxane (fig. 19), unlike the curves for trichloroacetic

and camphorsulfonic acids (figs. 20 and 21), showed

no indications of anion homoconjugation. Extending Maryott's work to a larger group of carboxylic

acids and nitrophenols, with acetone or acetonitrile as the solvent, Bryant and Wardrop [474]

perceived that the exceptional behavior of picric

acid is shared by 2,6-dinitrophenol, and must be

the result of strong chelation. As readers will have

foreseen, a second impediment to homoconjugation is the presence of a bulky substituent in one or both ortho positions.

Tables 59 and 60 contain substantial documentation of these statements, and additional evidence

wiU be discussed briefly. Investigations yielding

positive evidence of phenolate homoconjugation are summarized in table 59. Table 60 summarizes ex- periments in which little or no tendency toward

homoconjugation was observed.

a. Examples of homoconjugation. The formation

of homoconjugate phenolate ions (usually 1:1 complexes) has been inferred in the following types of

measurement:

(1) Conductance titrations of phenols in aprotic

solvents (benzene, toluene, xylene, carbon tetra- chloride), in certain dipolar aprotic solvents

(acetone, methyl isobutyl ketone), and in the weakly

basic solvent, pyridine. The evidence of homoconjugation was a maximum, usually quite pronounced,

in conductance near the titration midpoint [474,

484].

(2) Potentiometric titrations of phenols (or, alternatively, determinations of EMF values for series of buffered phenol solutions) in a similar

range of solvents (benzene, toluene, gasoline,

pyridine, ethylenediamine, acetonitrile) [475, 476,

486, 487]. Here, the indication of homoconjugation was an additional potential rise near half-neutralization, which seems to have been reported first in work by van der Heijde [475]. He noted, for example,

a close parallelism in the shapes of titration curves

for acetic acid (in acetone) and phenol (in pyridine)

when titrating with Bu3N(Me)0H (in pyridine),

using a glass-calomel electrode combination. A

notable investigation of analytical aspects of such

titrations has been performed by Harlow and

Bruss [476]; several examples from their results are included in tables 59 and 60. Special mention

should also be made of the use of potentiometric

data in calculating . ha values for several

phenols in ethylenediamine [486] and acetonitrile

[487] (see table 59). 97
Table 59. Examples of homoconjugate phenolate ions (ArO ...HOAr)"

Solvents Type of measurement

Estimated

value of Refer

Acetone

Acetonitrile Benzene

Ethylenediamine

f-Propyl Alcohol

Pyridine Toluene

Toluene

Toluene ^ Pyridine Toluene ' Ethylenediamine

Pyridine Toluene

Ethylenediamine

Ethylenediamine Toluene '

i-Propyl Alcohol

Pyridine

Acetonitrile i-Propyl Alcohol

Pyridine Toluene

Acetonitrile Methyl sulfoxide

i-Propyl Alcohol

Acetonitrile Acetone

C

EMF

UV

EMF

UV

EMF, C

' EMF, C

C

C

C

C

EMF

EMF

EMF

EMF

EMF

EMF

UV

C

EMF

VIS C

EMF, C

EMF

IRVIS EMF

C

15

110

1400

100

[484]

[487]

[485]

[486] [485]

[475, 484]

[476, 484]

[484]

[484]

[484]

[484]

[486]

[475]

[476]

[486]

[486] [476]

[485]

[484]

[487]

[485]

[484]

[476, 484]

[487]

[391]

[485]

[487]

[474]

^ Usually the cation present was R4N+. but in a few cases it was K+. Na+, or Et3NH+. Meaning of symbols is the same as in the preceding tables. Some experiments utilized a series of buffered solutions; in other cases, titra- tions were performed.

"/^A- H-A applies to the reaction: A'-f H-A A-...H-A. ^The anion PhO"(HOPh)2 was also thought to form. 'A small amount of i-PrOH was present.

^ Thymol.

References:

[391] D. Hadzi, A. Novak, and J. E. Gordon, J. Phys. Chem. 67, 1118 (1963). [474] P. J. R. Bryant and A. W. H. Wardrop, J. Chem. Soc, p. 895 (1957). [475] H. B. van der Heijde, Anal. Chim. Acta 16, 392 (1957). [476] G. A. Harlow and D. B. Bruss, Anal. Chem. 30, 1833 (1958). [484] D. B. Bruss and G. A. Harlow, Anal. Chem. 30, 1836 (1958). [485] L. E. I. Hummelstedt and D. N. Hume, Anal. Chem. 32, 1792 (1960). [486] S. Bruckenstein and L. M. Mukherjee, J. Phys. Chem. 66, 2228 (1962). [487] J. F. Coetzee and G. R. Padmanabhan, J. Phys. Chem. 69, 3193 (1965). 98
Table 60. Examples of phenols having little or no tendency to form homoconjugate

anions^

Substituents Solvents Cation Type of References

in P H OH present measurement''

2-N02'^ Benzene'' n-Bu^N + VIS [485] Toluene'' n-Bu4N + EMF [476]

2,6-(N02)2 Acetone EtaNH* C [474]

2,4,6-(N02)3 Acetone Et3NH + C [474] Acetonitrile Ejlrir Acetonitrile Et4N + EMF' [488] Rii,NH„+ Q

3-Me-2,4,6-(N02)3 Benzene TDPNB" VIS [482]

2,6-Me2 Pyridine Na + or K + c [484]

2,4,6-Me3 Pyridine Na+ or K + c [484] Toluene"* ra-Bu4N + c [484]

2-Me-6-;-Bu Toluene'' n-Bu4N + EMF [476]

4-Me-2,6-;-Bu Pyridine Bu3NMe + EMF [475] Pyridine Na+ or K+ C [484]

2,4,6-t-Bu Toluene '' a-Bu4N + EMF [476]

Thiophenol Acetonitrile Bu3NMe + EMF [475] Toluene'' n-BujN + EMF [476]

"Except as stated otherwise, no indications of homoconjugation were detected. Meaning of symbols is the same as in the preceding tables.

'^ Compare with the result for an acetonitrile solution, cited in table 59.

'' A small amount of i-PrOH was present. This was an exploratory study. "Very slight homoconjugation" was detected, with /< a-... H-a estimated as being about 2.

" TDPNB symbolizes the indicator base "p-Tolyl di-n-propyl Nile Blue" (LXXXIV).

References:

[341] A. A. Maryott, J. Res. NBS 38, 527 (1947). [474] P. J. R. Bryant and A. W. H. Wardrop, J. Chem. Soc, p. 895 (1957). [475] H. B. van der Heijde, Anal. Chim. Acta 16, 392 (1957). [476] G. A. Harlow and D. B. Bruss, Anal. Chem. 30, 1833 (1958). [482] M. M. Davis and H. B. Hetzer, Anal. Chem. 38, 451 (1966). [484] D. B. Bruss and G. A. Harlow, Anal. Chem. 30, 1836 (1958). [487] J. F. Coetzee and G. R. Padmanabhan, J. Phys. Chem. 69, 3193 (1965). [488] I. M. Kolthoff and M. K. Chantooni, Jr., J. Am. Chem. Soc. 87, 4428 (1965). [485] L. E. I. Hummelstedt and D. N. Hume, Anal. Chem. 32, 1792 (1960).

(3) Photometric titrations of phenols. Phenolates

absorb at longer wavelengths than the corresponding

phenols, and a wavelength region can usually be

found in the near-ultraviolet or visible where the

absorbance of the phenolate increases sharply with

concentration, and where the phenol contributes

little or nothing to the absorbance. Taking advantage of this situation, a number of investigators

have successfully determined 1:1 acid-base association constants in aprotic solvents Uke benzene (for examples see refs. 14, 20, 32, 254a, 482) and chlorobenzene [354, 364, 441. 85]. The phenols selected

for such work (bromophthalein magenta E, bromophenol blue, 2,4- and 2,6-dinitrophenol) were very

unlikely to form homoconjugates, especially under

the conditions used.

Phenolate homoconjugation was very evident,

however, in a different type of investigation by

99
Hummelstedt and Hume, whose objective was to explore analytical possibilities of photometric titra- tions [485]. They titrated some 16 different phenols

(phenol itself, 1- and 2-naphthols, and halo- or

nitro-derivatives of phenol), usually 0.01 M in isopropyl alcohol, with n-Bu4N0H (~ 0.4 M) in the same solvent. In part of the experiments the iso- propyl alcohol was largely replaced by benzene.

With isopropyl alcohol as the sole solvent, the

titration curves for phenol, unchelated phenols,

and non-hindered phenols tended to have an S- shape resembling that illustrated earlier for certain

carboxylic acids (fig. 22B). However, as would be

expected, an anomalous curve shape was most

pronounced when most of the isopropyl alcohol was replaced by benzene. Agreeing with other

investigators, Hummelstedt and Hume concluded

that ArO"...HOAr formation occurs at the beginning of such titrations. An additional reasonable postulate was that ArO"...HOAr has spectral

absorption characteristics intermediate between

those of ArOH and ArO~. As they pointed out, undesirable spectral shifts occur during a photometric titration when using a mixture of solvents

(for example, benzene-isopropyl alcohol) whose

composition does not rernain constant. Hummelstedt and Hume also stressed the sensitivity of the photometric method as a detector

of homoconjugation: Whereas in toluene containing more than 1 percent of isopropyl alcohol neither

the conductimetric [484] nor the potentiometric

[476] method supplies indications of phenolate

homoconjugation, phenol homoconjugates are de- tectable photometrically even in pure isopropyl

alcohol. In their opinion the photometric method

offers unique advantages in analytical applications

(including differentiating titrations of up to four

weak acids), as well as for investigating complex

formation.

(4) Infrared spectroscopy. Hadzi, Novak, and

Gordon [391] isolated several acid phenolates,

including the sodium and the potassium acid salt of 4-nitrophenol, and measured their infrared ab- sorption spectra as mulls, as undercooled melts,

and as solutions in methyl sulfoxide. Indications

of ArO"...HOAr were the absence of a band in the region of the stretching vibration of free 0-H,

and the presence of an extremely broad absorption

band whose location and other characteristics were

considered consistent with its assignment to a hydrogen-bonded group, 0...H...0. These authors were concerned with the symmetry or asymmetry

of such hydrogen bonds. For our discussion, it is sufficient to note that homoconjugation was inferred

for specimens examined as solutions in methyl

sulfoxide, as well as for those examined as mulls

or melts. However, the hydrogen bonds thought to be symmetrical in certain intermolecular homoconjugates in the absence of a solvent became

asymmetrical on solution in methyl sulfoxide.

2,2'-Dihydroxybiphenyl and bis-(2-hydroxyphenyl)-

methane, as well as some of their phenyl-substituted derivatives, show marked readiness to form

conjugate anions [391, 476, 494]. For example, the

infrared spectra of the sodium acid salt of 2,2'- dihydroxybiphenyl (NaOC6H4C6H40H) — whether as

crystals, as an undercooled melt, or as a solution

in methyl sulfoxide — are all very similar, and the same holds true for the sodium acid salt of 2,2'- methylenebis-(4-chlorophenol), NaOC6H3(Cl)-CH2-

C6H3(C1)0H [391]. This similarity and other characteristics of the spectra are considered to indicate

that these salts contain a very stable A~...H-A

bond. The stability is understandable, since the

A"...H-A bonds, in all probability, form part

of a cyclic structure — a 7-membered ring in the

first-mentioned salt, and an 8-membered ring in the latter. Salts of these types display exceptional

titrimetric behavior. As an illustration, when

titrating various alkyl-substituted bis-(2-hydroxyphenyl)methanes in benzene-i-PrOH, Sprengling

observed [494] that one of the phenolic groups is much more acidic than a single phenolic group, whereas the second phenolic group is abnormally

weak. Harlow and Bruss [476] confirmed Sprengling's observations, finding one phenolic OH group

to be easily titratable, while the second OH could

not be titrated at all. Table 59 contains estimates of Ka . h a for phenol and three phenol derivatives in ethylenediamine [486], and for three nitro derivatives of phenol in acetonitrile [487]. While these results are obviously too few to support any broad generalizations, the values for the four phenol derivatives

in ethylenediamine suggest that, paralleling car- boxylate homoconjugation, the stronger the con- jugate acid ArOH, the more stable the homoconjugate ArO~...HOAr. The other three Xa ha

values will be discussed briefly in the next section.

b. Examples of non-homoconjugation. As mentioned a little earUer, homoconjugation is markedly

impeded when the phenol in question either (1) contains a stable chelate ring, or (2) is an appreciably

"hindered" phenol. At this point we shall elaborate

somewhat on these two topics.

Chelated Phenols. The role of chelation is well

illustrated by the differing conductance titration curves obtained for 3,5-dinitrophenol (listed in table

59) and its 2,6-isomer (see table 60) in acetone solutions: A mid-titration maximum appeared in the curve for 3,5-dinitrophenol, but was absent in that

for 2,6-dinitrophenol [474].

It is also instructive to compare the homoconjugation constants estimated from EMF data for 2- and

4-nitrophenols and 2,4-dinitrophenol in acetonitrile

(table 59) [487]: Ahhough 2-nitrophenol (pA:« = 7.23)

and 4-nitrophenol (pKa 7.15) are about equally

strong acids in water, it is plain that in the experiments cited homoconjugation was much more ex- tensive for the para- than for the ortho-isomer. The

orfAo-isomer is, undoubtedly, the example most

frequently cited — with extensive documentation —

in discussions of chelate ring formation in aprotic

solvents. Homoconjugation of 2,4-dinitrophenol ap100
pears, likewise, to be strongly inhibited by chelation: In spite of its much greater pKa value (4.11), its homoconjugation constant is much smaller than that of 4-nitrophenol, being, instead, about the same as

that of 2-nitrophenol.

Readers will probably have noticed that 2- nitrophenol is listed both in table 59 and 60. Its inclusion in the latter table arises from the fact that even a slight indication of homoconjugation seems to have been lacking in photometric titrations of this acid in' benzene containing a little isopropyl

alcohol [485], and also in EMF titrations in a similar

medium (namely, toluene containing a little iso- propyl alcohol) [476]. Picric acid, likewise, might

have been listed in both tables, since "very slight

homoconjugation", with Ka-. h-a ~2, was in- ferred in EMF measurements employing acetonitrile as the solvent [488].

Tables such as these, while not to be regarded

as rigorous classifications, have predictive value. We are not surprised, for example, to learn that in photometric titrations of tolyl dipropyl Nile Blue

base with trinitro-m-cresol (3-methylpicric acid)

in benzene, homoconjugation of this phenol derivative was not discernible. The absence of homoconjugation is indicated in figure 22A, as well as in figure 23. The latter figure shows representative

absorption curves for the indicator base before and

Wavenumber, cm-' x lO"' 25 20 15

60

1 1 1 1 1 I 1 1

: Tolyl Dipropyl Nile Blue

10-5 M in Benzene,

- plus Trinltro-m-cresol

CalcdN. 3ase, ji,.

1

' —1600 equlvs.J 100

40

0 / 10 3

// 5 \ / / 5 20

\v 1

0 1

1 0 400 500 600

Wavelength, nm

Figure 23. Spectral absorption curves of mixtures of tolyl dipropyl Nile Blue base {LXXXIV) with trinitro-m-cresol (methylpicric acid) in benzene; used in determining the equi- librium constant (Kbha) for the 1:1 acid-base association at 25°.

Reproduced from figure 7 of M. M. Davis and H. B. Helzer, Anal. Chem. 38, 451 (1966).

(Ref. No. 482.)

after adding trinitro-m-cresol in amounts ranging

from 1 to 1600 equivalents. Not only is there a

well-defined isosbestic point (near 556 nm), but

there is also excellent agreement between a cal- culated (dotted line) and experimental limiting curve. Data such as these yielded a precise value

(log .^BHA^4.13, molar units, 25°) of the 1:1 acidbase association constant [482]. Furthermore, with the benefit of our recently

acquired understanding of homoconjugation in carboxylates and phenolates, it is now possible to

offer reasonable explanations of some puzzling re- sults in early studies of acid-base behavior in aprotic

solvents. For example, Br0nsted and Bell's finding

[491] that the velocity of the addition of phenol to ethyl diazoacetate in benzene is strongly catalyzed

by carboxylic acids is perhaps attributable to en- hancement of acidity by Ae^eroconjugation (see

later discussion in section 4.5.7b). In another study,

concerned with the acid-catalyzed rearrangement of A^-bromoacetanilide in chlorobenzene. Bell [492]

derived a simple equation correlating the catalytic

effects in this solvent and the aqueous ionization constants of 2,4,6-trichlorophenol, pentachlorophenol, and the following eight carboxylic acids:

acetic; 2-chloropropionic; cinnamic; phenoxyacetic;

phenylpropiolic; and mono-, di-, and trichloroacetic

acids. Astonishingly, picric acid was far too poor

a catalyst to fit into this relation. Picric acid was

again a misfit in investigations by Bell and Caldin

[493] of the acid-catalyzed inversion of /-menthone

in chlorobenzene (99.4°); that is, its catalytic activity

in this reaction was anomalously low. Such results are in keeping with the failure of picric acid to undergo self-association or to readily form homoconjugate

anions. It is evident that the comparatively great

strength of picric acid in water (pKa = 0.33, according to recent spectrophotometric data [362]) is a misleading clue to its behavior in nonleveling sol- vents like benzene and chlorobenzene. Although,

as we saw earlier (fig. 19), picric acid is quantitatively titratable in such media with relatively strong

organic bases like di-n-butylamine and triethylamine, apparently the 0-H bond to an ortho nitro oxygen is too strong for there to have been extensive

intermolecular hydrogen bonding of picric acid to the weak organic bases involved in the catalytic

experiments just mentioned.

Hindered phenols [475, 476, 484]. A well-known

example of a hindered phenol is 6-methyl-2-feri.-

butylphenol (table 60). This phenol gave a "normal"

potentiometric titration curve in toluene, whereas

the curve for a less hindered isomer, 4-methyl-2-

ierf.butylphenol, showed an inflection at the midpoint of the titration [476]. Most of the phenols

listed in table 60 are substituted in both the 2- and

6-positions, but this does not mean that there is no inhibition of homoconjugation with only one

bulky ortho substituent. An illustration of this point can be found in Harlow and Bruss's paper

[476], where comparative potentiometric titration curves show decreasing evidence of homoconjugation with increasing shielding of the -OH group in the following series of phenols:

CeHsOH > 2-?-Bu-4-MeC6H30H

> 2-i-Bu-6-MeC6H30H > 2,4,6-(i-Bu)3C6H20H.

Further comparisons of potentiometric or conduct- ance titration curves for unhindered, hindered, and

partially hindered phenols are to be found in the

three papers cited.

101
Thiophenol [475, 476, 484]. In the same three

investigations thiophenol exhibited no tendency

toward homoconjugation in spite of its having a

distinctly more acid character than phenol. That

is, it was found to give "normal" curves in potentiometric or conductance titrations employing such

solvents as acetonitrile [475], pyridine [484], and

toluene containing a minute amount of isopropyl

alcohol [476]. This difference in the behavior of

phenol and thiophenol is attributed to the fact that sulfur is less able than oxygen to bond to hydrogen.

Possibly weak homoconjugation would be detected

in more favorable environments, that is, where

alternative hydrogen bonding reactions are more

completely excluded. This suggestion is supported

by a number of examples, summarij^ed by Pimentel

and McClellan [102], in which mercaptans and

other sulfur compounds have been found capable

of forming hydrogen-bonded complexes. Moreover,

thiophenol is now regarded as being somewhat

dimerized (or otherwise aggregated) in aprotic

media, although once considered to be incapable

of self-association. For example, Spurr and Byers,

employing infrared spectroscopy, deduced a dimerization constant (^1.2) of about 0.02 (CCI4, molar

units, 24°) [489], and Marcus and Miller, under

closely comparable conditions (PhCl, 28°), obtained

the value 0.011 by PMR spectroscopy [490]. These

are smaller than corresponding values reported

for phenol (see table 18).

4.5.5. Homoconjugation of Halide Ions At one period homoconjugation of halides was

thought to be limited to the acid fluoride ion, HF2,

but more recent work has shown that all four of

the common halide ions are capable of homoconjugation. Examples will be given of halide homoconjugation in aprotic and dipolar aprotic solvents,

and following that, some results with solid speci- mens wiU be summarized.

a. Homoconjugation in Solutions. It seems likely that Meerwein was the first to become aware of

chloride homoconjugation. In an article bearing

the title (in translation), "Enhancement of the

lonizability of Weak Electrolytes by Complex

Formation and Its Significance for the Catalytic

Process," he referred to "dimolecular acids, like dihydrochloric acid, [HCl2]H" as being, in commoVi

with "dimolecular organic acids," more strongly

ionized than the "monomolecular acids" [495].

(Ethyl ether seems to have been the solvent used.) In a major contribution appearing in 1954 [496],

Herbrandson and associates reported several reasons for postulating the existence of HCI2 in nitrobenzene containing both HCl and Et4NCl. For

example, in the mixtures the solubility of Et4NCl was substantially greater than in the absence of

HCl, the vapor pressure of HCl from the solution was less than in the absence of Et4NCl, and the

H-Cl stretching frequency at 2703 cm"' was

replaced by a broad absorption band at lower

frequencies. Results for acid-catalyzed reactions

also pointed to the same conclusion, yielding for the homoconjugation constant (see eq 70) the value

500 (PhNOa, molar units, 25°).

Cl- + H-Cl;:±HCl2". (70)

Furthermore, they isolated the acid salts Me4NHCl2

and Et4NHCl2 from nitrobenzene solutions.

During an investigation of triple ion formation

in benzene, Hughes, Ingold, and co-workers [497]

measured the conductance of solutions containing

HCl alone and in mixtures with BU4NCI, BU4NNO3,

and BU4NCIO4, obtaining the results shown in table 61. Numbers in the last column indicate the extent to which the conductance (30°) of any mixture

Table 61. Relative stabilities in benzene o/(ClHCl)-, (CIHNO3)-.

and (ClHClOj) , /rom conductance measurements (30°)^

Molar concentration Conductance''

HCl ra-BujNA HCl

only

n-Bu,NA

only Mixture

A- = C1-

0.40 X 10 "

3.0

3.3

3.4

3.6

3.8 4.0

1.92 X lO-'* 4.55 3.84

3.57

3.33

3.12 2.00

0.02

0.10

0.10

0.11

0.11

0.11 0.12

0.17

0.55

0.41 0.38

0.36

0.34

0.18

0.33

2.08

2.00

1.87 1.74

1.60 0.93

0.14

1.43

1.49 1.38

1.27

1.15 0.63

A- = N03-

10.0

9.7

9.5

2.55

2.5 2.4

2.55

2.63

2.83

0.25 0.50

0.01 0.02

0.12

0.41

A- = C104-

10.0 9.4

8.8

. 4.5

4.2

3.7

4.5

4.1 3.8

0.60

1.00

0.02 0.03 -0.12

0.07

See Table 4 of Hughes, Ingold, et al. [497]. The criterion of

stability is the magnitude of A; see accompanying discussion. ''The conductance (reciprocal resistance, 10" ohm~') was measured before and after adding various amounts of HCl to solutions containing n-Bu4NA, where A" was CI", NOj^, or ClOf

''A = The conductance of the mixture minus the sum of the conductances of the separate components.

Reference

[497] E. D. Hughes, C. K. Ingold, S. Patai, and Y. Pocker, J. Chem. Soc, p. 1206 (1957).

102
Table 62. Examples of homoconjugate anion (X~...H-X) formation, deduced in studies of solutions^

Solvent XClass of base or cation Method "

Acetonitrile Br" T d In ° VIS

Br- R4N+ IR, UV, PMR

ci- In" VIS

ci- R4N+ IR, UV, PMR

Cl" RjN* IR, Raman

Benzene Br" R^N^ IR

ci- R^N^ C, VPL, RATE

ci- R4N+' IR, Raman

Carbon tetrachloride F" R3N ' IR

Chloroiorm D — Br R:)lN

T \ u

ci- R3N D

r

r 1U

Dichloromethane ci- R4N+ IR, Raman

Ethyl ether ci- N.S.*^ RATE

Methyl sulfoxide Br- R4N+f IR, PMR

ci- R4N+ IR, PMR

Nitrobenzene Br- Py C

ci- R4N+ ' IR, RATE, S, VPL

ci- Py C

Nitromethane ci- R4N+ VIS

Estimated

value of

^x-...H-x'^

References

250

200

500

[500]

[502J

[500]

[502]

[504]

[503]

[497, 498]

[503, 504]

[381]

[409]

[409] [409]

[504]

[495]

[502]

[502]

[501]

[496]

[501]

[499]

^ In some cases X~...D-X was studied as v^ell. ''Significance of symbols: RATE, enhancement of acid strength by homoconjugation caused an increase in the rate of an acid-catalyzed reaction; other symbols are used as

in the preceding tables. At 25° in molar units. Kx-...h-x corresponds to X~ + HX ^ X"...H-X. Activity coeffi- cients are assumed to cancel or approximate 1. In this table "In" signifies a nitroaniline-type ("Hammett") indicator. Also see table 65.

^ Not specified. 'The salt used in experiments had been isolated and identified as B(HX)2. See table 63

for additional examples of soUd species of the composition B(HX)2.

References:

[381] A. A. Lipovskii and S. A. Nikitina, Russ. J. Inorg. Chem. Engl. Transl. 10, 93 (1965). [409] J. I. Bullock, S. S. Choi, D. A. Goodrick, D. G. Tuck, and E. J. Woodhouse, J. Phys.

Chem. 68, 2687 (1964). [495] H. Meerwein, Ann. Chem. 455, 227 (1927). [496] H. F. Herbrandson, R. T. Dickerson, Jr., and J. Weinstein, J. Am. Chem. Soc. 76,

4046 (1954). [497] E. D. Hughes, C. K. Ingold, S. Patai, and Y. Pocker, J. Chem. Soc, p. 1206 (1957). [498] S. F. Mok, cited in ref. [497], p. 1216.

[499] Y. Pocker, J. Chem. Soc, p. 240 (1958). [500] I. M. Kolthoff, S. Bruckenstein, and M. K. Chantooni, Jr., J. Am. Chem. Soc. 83,

3927 (1961). [501] I. R. Beattie and G. J. Leigh, J. Chem. Soc, p. 4726 (1962). [502] D. G. Tuck and E. J. Woodhouse, Proc Chem. Soc, p. 53 (1963). [503] J. C. Evans and G. Y-S. Lo, J. Phys. Chem. 69, 3223 (1965). [504] J. C. Evans and G. Y-S. Lo, J. Phys. Chem. 70, 11 (1966).

103
exceeded the sum of the conductances of the

separate components. From these results it was

concluded that (ClHCl)" is very stable in benzene,

(CIHNO3)- is moderately stable, and (CIHCIO4)-

does not exist. In harmony with these results, S. F. Mok [498]

found the vapor pressure of HCl (~ 0.001 M) from

benzene to be strongly depressed on the addition of

/i-BoiNCl in a comparable amount. Moreover, she

observed that the catalytic effect of HCl on the

methyl-alcoholysis of triphenylmethyl (trityl) chloride in benzene is largely destroyed by adding

an equivalent of BU4NCI, whereas HCl and n- Bu4NC104 act additively as catalysts in this reaction.

The examples just cited, along with numerous

others, are listed in table 62, where nine solvents

are represented. It will be seen that homoconjugation has been detected for F~, CI", and Br , employing numerous types of physical measurements (conductance; distribution; kinetics; IR, UV, visible,

PMR, and Raman spectroscopy; solubiUty measurements; and vapor pressure lowering). Indications

of homoconjugation need not all be discussed in

detail, since they are very much like those for homoconjugates already discussed. For instance,

in conductance titrations of HCl (0.01 M or 0.05 M)

in nitrobenzene the titration curves show a pro- nounced maximum near the mid-point, and similar

results are obtained in titrating HBr [501].

Naturally, there has been much speculation as

to the manner in which acids like HCl exert their

catalytic function in aprotic media. A pertinent

example is the interaction of phenol with trityl

chloride in o-dichlorobenzene to form p-tritylphenol.

Hart and Cassis, who studied this reaction and the

corresponding reaction of o-cresol with trityl

chloride, yielding p-trityl-o-cresol, concluded that the most likely role of HCl is in facilitating the

ionization of trityl chloride, as represented by the

following equation [505]:

PhaC-Cl + HCl ^ PhsC^Cr. . .H-Cl. (71)

b. Homoconjugation of X~ in solids [190b, 381,

502-504, 506-511]. During the past ten years a number of investigators have examined acid fluo- rides, chlorides, bromides, and iodides as the solids,

and they have unanimously inferred the presence of

X~...H-X (or, in some cases, X ...D-X). The nature

of such studies is briefly indicated in table 63. Mixed conjugates have also been detected in solids -for example, (CIHNO3)-, (CIDNO3)-,

(BrDCl)-, (IDCl)- [507], and (BrHCl)" [508]. There is wide agreement that FHF~ is the most

stable of all hydrogen-bonded species, and that, moreover, its hydrogen bond is symmetrical. In other words, it is thought that the proton in F~H^F~

is located in a single potential minimum, inter- mediate between the F~ ions [see fig. 24(d)].

Westrum and Pitzer [506] believe this conclusion

to have been established with certainty by their own careful investigations of the system KHF2 —

KF — HF, in which thermodynamic and dielectric

Table 63. Studies of homoconjugation of halide ions^ in solids

X Types of cations present Method* References

F- K+, R4N+, THERM, P, IR [506, 510, 190b, 511, R3NH + 381] ci- Cs +

, R4N + IR, Raman, [509, 503, 504, 507, THFRM OKJOj Br- THERM, IR [508, 502, 507] I- THERM, IR [508, 507]

aX-...H-X or X ...D-X.

^ "THERM" signifies the measurement of one or more thermo- dynamic properties such as heat capacity, entropy, and lattice energy.

References:

[190b] L. PauHng, The Nature of the Chemical Bond, 3d ed. (Cornell Univ. Press, Ithaca, N.Y., 1960). [381] A. A. Lipovskii and S. A. Nikitina, Russ. J. Inorg. Chem.

Engl. Transl. 10, 93 (1965). [502] D. G. Tuck and E. J. Woodhouse. Proc. Chem. Soc, p. 53

(1963). [503] J. C. Evans and G. Y-S. Lo, J. Phys. Chem. 69, 3223

(1965). [504] J. C. Evans and G. Y-S. Lo, J. Phys. Chem. 70, 11 (1966). [506] E. F. Westrum, Jr. and K. S. Pitzer, J. Am. Chem. Soc. 71,

1940 (1949). [507] J. A. Salthouse and T. C. Waddington, J. Chem. Soc. (A),

p. 28 (1966).

[.508] D. H. McDaniel and R. E. Vallee, Inorg. Chem. 2, 996

(1963). [509] T. C. Waddington, J. Chem. Soc, p. 1708 (1958). [510] T. C. Waddington, Trans. Faraday Soc. 54, 25 (19581. [511] S. A. Harrell and D. H. McDaniel, J. Am. Chem. Soc. 86,

4497 (1964).

data were obtained, as well as by infrared spectral

data from the literature. Another indication of the

exceptional stability of HF2" is found in Waddington's calculated value, 58 ±5 kcal/mole, for the

enthalpy change in the following reaction [510]:

HF2-(g)^HF(g) + F-(g). (72)

McDaniel and co-workers ascertained the comparative stabilities of HI^, HBrf, HCl^, and HF^

in "systems designed to minimize lattice energy

effects and specific cation-anion interaction"

[508, 511]. The reactions studied are representable

by the foUowing equation:

R4NX(s) + HX(g) ^ R4NXHX(s). (73)

For these four homoconjugates, in the order listed,

the respective values found for A// are: —12.4, — 12.8, — 14.2, and — 37 kcal/mole. The corresponding value obtained for the mixed species (BrHCl)"

is —9.1 kcal/mole; in the adoption of this value

the proton was presumed to be closer to Cl~ than

to Br-.

104
4.5.6. Homoconjugation of Other Organic

and Inorganic Acids

a. Perchloric acid. In work previously reviewed

CIO4 was found to have little if any tendency to become hydrogen bonded to cations R3NH+ (see

sec. 4.3.5c) or to HCl (sec. 4.5.5a) in aprotic media.

Consequently, it is not surprising that extraction

experiments have supplied no evidence of an ability

to form a homoconjugate anion. Such experiments

have shown that when concentrated aqueous mineral acids (HCl, HNO3, and certain others) are equilibrated with an RsN-containing aprotic solvent

the organic layer acquires more than one HA

molecule per molecule of R3N. However, after extractions of aqueous perchloric acid by a chloro- form solution of methyldioctylamine. Bullock and

associates [409] concluded that essentially no AHA" was present in the organic layer.

Similarly, Bucher and Diamond found that benzene, o-dichlorobenzene, and nitrobenzene solu- tions of trilaurylamine, following their use as extracting media, contained only the 1:1 perchlorate

[442]. Their experimental results are summarized

in table 64.

Table 64. Molar ratio YiQAOnltrilaurylamine m the organic

phase (22°) ^

Molarity of

aq. HCIO4

Molar ratio HC104/(C,2H25)3N in

Benzene o-Dichlorobenzene Nitrobenzene

0.495 1.00 1.00 1.00 0.991 1.00 1.00 3.92 1.00 1.00 6.37 1.00 1.00 1.03 10.00 1.10 1.07

"From Table I of J. J. Bucher and R. M. Diamond, J. Phys.

Chem. 69, 1565 (1965). The ratio was determined by appropriate

analytical procedures after benzene, o-dichlorobenzene, and

nitrobenzene solutions of trilaurylamine (—0.1 M) had been

equilibrated with aqueous perchloric acid of various concentrations by shaking for an hour. Corrected for HCIO4 extracted by the pure solvent.

Reference:

[442] J. J. Bucher and R. M. Diamond, J. Phys. Chem. 69,

1565 (1965).

b. Other acids. Table 65 lists conditions under

which the homoconjugate anions of several common

mineral acids (HBr, HCl, HNO3, H2SO4) and some

related organic acids (camphor- and methanesulfonic acids) have been observed to form. [Also see

table 62.] The di-(2-ethylhexyl) ester of orthophosphoric

acid is also beheved to form a homoconjugate

anion in toluene [513, 514]. After an x-ray examination of crystalline di-p-chlorophenyl hydrogen phosphate [517], Calleri and Speakman formed the

opinion that the hydrogen bond linking a molecule

of this acid with the anion of a second molecule

may be "genuinely symmetrical" (and probably,

therefore, comparatively strong). No corresponding

studies seem to have been made for a related acid

sah.

Estimates of Ka-...h-a for HBr, HCl, HNO3,

and H2SO4 in acetonitrile solutions [500, 516| are included in table 65. Apparently the first three of

these acids are moderately, and about equally,

homoconjugated in this solvent, and H2SO4 is considerably more so. The work reviewed has made clear that anion

homoconjugation is to be expected in solvents of

little or no proton-donor ability, the dielectric constant having, at most, a minor role, and it has

also extended our understanding of the role of water

and alcohols as ionizing media. These conclusions were aptly expressed by Hammett [515], following

a study of sulfuric acid-indicator reactions in nitro- methane: "The clear implication is that the reaction

of one molecule of acid with one molecule of base

as it is observed in waterlike solvents requires as an indispensable factor the solvation of the anion,

presumably by hydrogen bonding between the acidic

hydrogen of the solvent and the unshared electrons

of the anion. In the absence of solvent molecules

with the necessary properties the transfer of the

proton from the acid to the base may require the

drafting of an additional molecule or molecules of

the acid to perform the function fulfilled by the

solvent in waterlike systems."

4.5.7. Heteroconjugate Anions

The conclusions just quoted are documented still more strongly and directly by systematic investigations of heteroconjugation of anions. The majority

of such studies have been concerned with hydrogen

bonding to X" (halide) ions, especially Br" (see tables 66 to 69). A much smaller number (see table

70) have dealt with bonding to carboxylate, nitrate, or bisulfate.

a. Chemical behavior. The earliest observations were made without the aid of instruments. Using

a method devised by Hantzsch [18], in which the

comparative strengths of acids are determined by

ascertaining how much ether must be added to red

salts of the indicator base, methyl yellow, before the

color reverts to yellow, Meerwein observed [495] that for the hydrochloride this point was reached on

dilution to about 0.02 A'^. However, with benzoic acid

(or another organic acid), even in its saturated ether

solutions, the indicator remained yellow. Meerwein

then mixed the two yeUow solutions, whereupon the

color instantly became red, and considerably more

ether had to be added before the color was again

yellow. He concluded that the intensified acid

reaction was caused by the formation of a complex

of HCl and RCOOH, which he represented as [HCl • OCOR]H.

Dietzel and Paul [540] discovered that in titrations of alkaloids with p-toluenesulfonic acid in chloro- 105
Table 65. Homoconjugate anion formation from miscellaneous organic and inorganic acids, deduced in studies of solutions

Solvent Acid Base or cation present Method " A A~...HA References

Acetonitrile Hydrobromic In"^ VIS "280 [500, 516] Hydrochloric In"^ VIS " 170 [500, 516] Nitnc In*^ VIS 200 [500] Sulfuric In' VIS 1,150 [500] Sulfuric Na+ S 1,000 [516] Sulfuric n-BuNHj c [479]

Lhloroiorm Hydrocyanic K^iN D [409] Nitric R:,N D [409]

Dioxane c?/-Camphorsulfonic EtjN or Bu.NH C [341]

Nitrobenzene Methanesulfonic Py-type A// [512]

Nitro methane Sulfuric In'^'' VIS [5151

Toluene Di-(2-ethylhexyl) La+3 D, CRY [513, 514] phosphate

^ A// signifies heat of reaction. Other symbols have the same significance as in the preceding

tables.

*' Note that these values are slightly different from those cited in table 62. In this table 'Tn" signifies an indicator of the nitroaniline ("Hammett") type.

" In >0.1 M H2SO4 solutions an indicator saU of the composition InH+...OS(OH) (O...HOS03H)2 was thought to be formed.

References:

[341] A. A. Maryott, J. Res. NBS 38, 527 (1947). [479] I. M. Kohhoff and M. K. Chantooni, Jr., J. Am. Chem. Soc. 85, 426 (1963). [500] I. M. KohhofT, S. Bruckenstein, and M. K. Chantooni, Jr., J. Am. Chem. Soc. 83 , 3927 (1961). [512] H. C. Brown and R. R. Holmes, J. Am. Chem. Soc. 77, 1727 (1955). [513] D. F. Peppard, G. W. Mason, J. L. Maier, and W. J. Driscoll. J. Inorg. Nucl. Chem 4, 334 (1957). [514] D. F. Peppard, J. R. Ferraro, and G. W. Mason, J. Inorg. Nucl. Chem. 4, 371 (1957). [515] H. Van Looy and L. P. Hammett, J. Am. Chem. Soc. 81, 3872 (1959). [516] I. M. Kolthoff and M. K. Chantooni, Jr., J. Phys. Chem. 66, 1675 (1962). [409] J. I. Bullock, S. S. Choi, D. A. Goodrick, D. G. Tuck, and E. J. Woodhouse, J. Phys. Chem. 68,

2687 (1964).

form, with methyl yellow as the indicator, the endpoint could be made sharper by adding about

1 percent of phenol, but did not suggest an explanation. Here, it seems reasonable to postulate hydrogen

bonding of phenol to one or more sulfonate oxygens,

with consequent enhanced strength of the sulfonic

acid. Later, intensified acid reactions were attributed

by their discoverers to hydrogen bonding. For

instance, Bartlett and Dauben [518] found that the

strength of HCl in dioxane, measured by an indi- cator method, is increased on adding any of the fol- lowing proton-donor molecules: Methanol, ethanol,

isopropyl alcohol; phenol, o- and p-cresol, o- and

p-chlorophenol; 2,4,6-trichlorophenol; p-cyanophenol; acetic acid. They ascribed the enhanced

strength of HCl to bonding of CI to the proton donor,

and remarked that "hydrogen bond formation

appears to run parallel to the acid strength in the

series of phenols except for phenols with orthosubstituents capable of intramolecular hydrogen

bonding, which are far less efficient 'acceptors'

than would be expected from their acid strength."

Similar reasoning has been used in explaining

results of reaction rate studies. As an illustration. Swain [537] found that the conversion of trityl

chloride or bromide to trityl methyl ether by reaction with IVIeOH in benzene (a base such as Py or EtsN being present to combine with HX, the

by-product) is greatly accelerated by adding phenol,

and even more so by adding p-nitrophenol. After

eliminating other possible explanations. Swain

concluded that the increased rates are a conse- quence of hydrogen bonding to the halogen of

trityl chloride to form Ph3C-X...H0Ar, thereby

weakening the attachment of HX to the methane

carbon and facihtating its replacement by OlVIe (compare eq 71). Analogously, in the reaction of tert. -hutyl bromide

in nitromethane with various OH derivatives (water,

106
Table 66. Examples of heteroconjugate halide ion {X' ...H-A) formation, in several aprotic solvents "

Solvent HA" X References

Acetonitrile HCl, DCl Br- [529]

Benzene t>r HNOa" ci- [497] ROH" Br- [519] ROH Br-, ci- [521, 522] CH3CONHCH3 Br- [521, 522]

Carbon tetrachloride HOH« Br-, C1-, F-, I- [528] ROH Br-, C1-, F-, I- [520, 524, 525] ROH I-f [526] ArUo DT , r c on c oc

[5/U, ozoj cm-, Br-, CI- [525]

rlL- — * L.il2i5r Br , T-1 [5z5J R,NH, RNHAr, or Br- [520] Ar.NH

vjiilorotorni 1 1 V 7 1

RrDr , T1

Dichloromethane rH_l, UCl tJr [529] ROH Br-, C1-, F-, I- [525]

TDioxane ^ 151»]

Ethyl ether ^ RCOOH ci- [495]

Nitrobenzene ROH Br- [47]

Nitromethane ROH Br [519]

" Except when indicated otherwise, solutions contained a mixture of R4N+X and HA, and X-...H-A formation was deduced from shifted i'(H-A). R denotes an alkyl group and Ar, an aryl group. Both IR and Raman spectra were measured.

Formation of X-...H-Awas deduced from conductance data. 'RuaNH* and BU2NH2*, also, were used as cations.

'^Both IR and UV absorption were measured, the IR data showing the effect of I-...HOR formation on viO-H) and the UV data showing its effect on absorption by I-.

" From UV absorption data only. Enhancement of acid strength resulting from C1-...H-A formation caused the

indicator base methyl yellow to turn red.

References:

[47] H. Sadek and R. M. Fuoss, J. Am. Chem. Soc. 72, 301 (1950). [495] H. Meerwein, Ann. Chem. 455, 227 (1927). [497] E. D. Hughes, C. K. Ingold, S. Patai, and Y. Pocker, J. Chem. Soc, p. 1206

(1957). [518] P. D. Bartlett and H. J. Dauben, Jr.. J. Am. Chem. Soc. 62, 1339 (1940). [519] R. C. Miller and R. M. Fuoss, J. Am. Chem. Soc. 75, 3076 (1953). [520] H. Lund, Acta Chem. Scand. 12, 298 (1958). [521] J. Bufalini and K. H. Stern, Science 130, 1249 (1959). [522] J. Bufahni and K. H. Stern, J. Am. Chem. Soc. 83, 4362 (1961). [524] J. B. Hyne and R. M. Levy, Can. J. Chem. 40, 692 (1962). [525] A. Allerhand and P. von R. Schleyer, J. Am. Chem. Soc. 85, 1233 (1963). [526] M. J. Blandamer, T. E. Cough, and M. C. R. Symons, Trans. Faraday Soc. 60,

488 (1964). [527] M. J. Blandamer, T. E. Cough, and M. C. R. Symons, Trans. Faraday Soc. 62,

296 (1966). [528] S. C. Mohr, W. D. Wilk, and G. M. Barrow, J. Am. Chem. Soc. 87, 3048 (1965). [529] J. C. Evans and G. Y-S. Lo, J. Phys. Chem. 70, 20 (1966).

280-613 OL-68—

107
alcohols, phenols, carboxylic acids) the latter are considered to assist the ionization of the bromide.

Gelles, Hughes, and Ingold [538] found the effective- ness of phenol, water, and ethanol to decrease in the order named. Subsequently, Pocker [539]

extended these results, using a variety of alcohols

and phenols, and also acetic and benzoic acids. For alcohols, the catalytic order proved to be:

PhCHaOH > MeOH > EtOH > Ai-PrOH

> n-BuOH > t-PrOH > 5-BuOH > f-BuOH.

The catalytic order found for phenols was as indi- cated by the following list of substituents {pKa

values in parentheses):

P-NO2 (7.15) > p-CN (7.97) > p-I (9.30)

> p-Br (9.36) > p-Cl (9.42) > None (10.00)

> /n-Me (10.09) > p-Me (10.26)

> p-t-Bu > p-MeO (10.21).

To explain the comparatively low ability of the two

carboxylic acids to assist the ionization of f-BuBr

in nitromethane it was suggested that in this solvent

they are largely present as the dimers.

b. Physical measurements. The remaining examples of anion heteroconjugation were studied with

the aid of physical methods, principally conductance

and infrared or ultraviolet absorption spectroscopy.

Table 66 is a listing of examples in which F",

CI", Br", or I" is arbitrarily regarded as being the

proton acceptor, and the proton donor is one of the

following chemical types (all of which are known to be weaker acids than HX): Water, alcohols, phenols, carboxylic acids, an acetylene derivative

(propargyl bromide), chloroform-d, an amide, or a secondary amine (specifically, BU2NH, PhNHMe,

PhiNH, or indole). In addition there are examples of bonding of HCl or DCl to Br". Another inclusion —

the bonding of HNO3 to CI" — was mentioned pre- viously in the discussion of table 61. Continued study of such systems has been in- spired in part by conductance data indicating solute- solvent interaction. In two early examples, the

results for Bu4NBr and related salts in mixtures of

methanol with nitrobenzene [47], nitromethane

[519], or benzene [519] (supplemented by other perti- nent data, such as dielectric constants, viscosities,

and spectral absorption) were not consistent with a "continuum theory." Instead, the solute was

". . . seen to behave like two or three different

substances, depending on whether it is in solvent

A, solvent B, or a mixture of the two . . .," and the

authors concluded that ". . . this seeming absurdity can only be accounted for by postulating that the actual solute particles are different in the

three cases; this can only occur if there is specific

interaction between solute and solvent" [519]. In a preceding paper [47] the specific suggestion was

made that in MeOH — PhN02 mixtures the anion

associates with MeOH and the cation, with PhN02,

as represented by (LXXXVH):

Bu^N"^ . CO^Ii'^C^E^ ) .Br" ("''HO'CHj )

(LXmil)

Several papers dealing with infrared studies soon demonstrated that conductance behavior of the

sort just mentioned is indeed caused in part by

hydrogen bonding of proton-donor molecules to anions. In the first of these papers, Lund [520]

listed shifts in i^(O-H) values for a number of

proton donors caused by adding anions to their

solutions in carbon tetrachloride or chloroform.

For example, table 67 shows comparative v(0-li]

Table 67. Comparative shifts of i'(O-H) of p-cresol, 0.10 M

in CHCI3, produced by adding different anions as R4N+

salts ^- ^

Anion Apparent

(0.031 M) molar vol. Ai-, cm ' of anion

CIO4- 44.5 -180

I- 36.7 -340

NO3- 29.4 -400

Br" 25.1 -405

Pi" -405

ci- 18.0 -445

^Adapted from Table 2 of H. Lund, Acta Chem. Scand. 12,

298(1958). : In the chloride R4N+ was (Bu3NCH2Ph)+. but in all the other

salts it was Bu4N^. i Chloroform was used as solvent because of the low solubili- t ties of some of the salts in CCU: it was recognized as being a ! competing proton donor. The values cited are from ref. [530]. Units are cc per mole.

'

References: '

« [520J H. Lund, Acta Chem. Scand. 12, 298 (1958). S [530] K. Fajans and O. Johnson, J. Am. Chem. Soc. 64 , 668 s (1942). S

shifts for p-cresol in chloroform when the anion was

varied. Lund pointed out that the spectral shifts

| vary inversely with the apparent molar volumes of I the anions (in cc per mole), as computed by Fajans 1 and Johnson [530] for infinitely dilute solutions in

' water. Of the anions listed, ClOf, the anion of lowest electron density, is clearly the poorest pro- ton acceptor, the best proton acceptor being CI",

the anion of highest electron density.

Table 68, also based on results of Lund, compares

the effects of added Br" on viO-H) for a series of proton donors, the solvent in this experiment being

j

108
Table 68. Shifts in v(0-H) for some alcohols and phenols in ecu, produced by adding 0.031 M Bu4NBr^

Proton donor pK„(H^O) " Av, cm"'

(ert-Butyl alcohol -275

i^o-Propyl alcohol — 275

Ethanol "^16 -285

2,2,2-Trichloroethanol ' 12.24 -330

2,4,6-Trimethylphenol " 10.89 -365

o-Cresol 10.29 -400

p-Cresol 10.26 -402

Phenol 10.00 -405

2,4-Dibromophenol ''7.79 '-420

2,4-DichIorophenol 7.85 f-450

2,4,6-Tribromophenol «5.9 "6.24 -505

2,4,6-Trichlorophenol " 5.5 " 6.48 ' 6.46 -535

^ Adapted from Table 1 of H. Lund [520]. Except as noted otherwise, pKa values are from ref. [271], Appendix 12.1, Table 5. 'Ref. [531]. The pK,, value for ethanol was obtained by an

extrapolation.

" Ref. [532]. Communicated by R. A. Robinson [533].

' It will be recalled (see fig. 8 and related discussion) that the

t'(O-H) bands of various monomeric, mono-o-halophenols in CCLi have been found to have two peaks, one of which has been

assigned to the cis (intramolecularly hydrogen-bonded) isomer and the other, to the trans isomer. The v(0—H) shift cited here

is the one assigned [209] to the cis isomer.

" Ref. [534].

^ Calculated from pK,, values for the related monosubstituted

phenols [535] on the assumption of additivity of substituent

effects.

' Ref. [536].

References:

[209] N. A. Puttnam, J. Chem. Soc, p. 486, 5100 (1960).

[520] H. Lund, Acta Chem. Scand. 12, 298 (1958).

[531] F. A. Long and P. Balinger, in Electrolytes (Trieste Symposium, 1959, B. Pesce, ed.), p. 152 (Pergamon Press, New York, N.Y., 1962).

[532] H. C. Ko, W. F. O'Hara. T. Hu, and L. G. Hepler, J. Am.

Chem. Soc. 86, 1003 (1964).

[533] R. A. Robinson, unpublished work.

[534] A. G. Ogston, J. Chem. Soc, p. 1713 (1936).

[535] A. I. Biggs and R. A. Robinson, J. Chem. Soc, p. 388 (1961).

[536] G. T. Tiessens, Rec. Trav. Chim. 50, 112 (1931).

carbon tetrachloride. Unmistakably, increases

with increased strength of the proton donor. Recalling that some of the proton donors are chelated or hindered, we would not anticipate a perfect correlation. Lund's results for secondary amines as proton donors were not included in this table; they

ranged from 130 cm~^ (for dibutylamine) to about

285 cm~* (for indole).

Lund called attention to an analytical application

of these findings: A band near 3/a in the spectra of

quaternary ammonium halides in chloroform indi- cates the presence of ethanol, added as a stabilizer. No such band appears in the spectrum when the

solvent is pure chloroform or chloroform stabilized

with hexane.

Infrared results of Bufalini and Stern for benzene

solutions [521, .522] added strong support to the

overall conclusions of Lund. IMoreover, when combined with dielectric data of Richardson and Stern

[523| they indicated convincingly that in eq 74 the

value of n is 1.

Bu4N+Br- + nHOMe Bu4N+Br-. . .nHOMe.

(74)

The equilibrium constant obtained for this reaction

(expressed as log K, molar units, 25°) is 1.81. Values

approximating —6.6 and —14 were obtained for A^(kcal/mole) and AS (cal/mole deg), respectively.

Like Lund, they tested the effect of varying the

proton acceptor (using Pi", NO3, Br", or CI") and

the proton donor (IVIeOH, n-BuOH, f-BuOH,

CHsCONHMe), reaching the same overall conclusions almost in toto. Some conclusions of special

interest are:

(1) Anion "solvation demand" in benzene-alcohol

systems increases in the order:

Pi- <^ NOj < Br" < C1-.

In fact. Pi" did not seem to be solvated at all. This

conclusion may appear at variance with Lund's

result (table 67), but we must bear in mind that Lund used a much more active proton donor in his comparisons.

(2) A steric factor is very apparent in the comparative behavior of n- and f-butyl alcohols.

(3) The interaction of MeOH with PhN(Me2)HCl

is much less than with BU4NCI. The reduced hydrogen bonding of CI" to HOMe is beheved to be the

results of strong hydrogen bonding to PhN(Me2)H+.

(4) Although not shifted by added BU4NCIO4,

f(O-H) for methanol is shifted by AgC104.

(5) The shift caused by adding A^-methylacetamide indicates that anion solvation is not limited

to the 0-H group, and "the effect of even dilute electrolyte solutions on biologically important materials which exhibit hydrogen bonding may be

quite significant."

Observations of Hyne and Levy, who studied the

interaction of n-Bu4NBr with ^-BuOH in carbon

tetrachloride [524], are in general agreement with

those of the above-named authors. Minor respects

in which their results differ from those of Bufalini

and Stern are ascribed to the use of carbon tetra- chloride as solvent instead of benzene.

Another investigation of similar character but

greater scope was performed by AUerhand and

Schleyer [525]. Here, the inert diluent was carbon

tetrachloride or dichloromethane. Besides tetra- n-butylammonium fluoride, chloride, and bromide

and tetraheptylammonium iodide, the salts studied

included other kinds of "onium" compounds — for

109
example, A^-hexadecylpyridinium chloride, bro- mide, and iodide, A'^-methylquinolinium iodide, and

tetraphenylarsonium chloride. The types of proton

donors were also extended to include propargyl

bromide and three deuterohaloforms. Very briefly

summarized, the main conclusions were:

(1) The infrared spectral shifts observed for pro- ton donors are caused by hydrogen bonding to anions;

(2) the magnitudes of the shifts depend strongly on the nature of the anion X", being in the order

CI- > F- > Br- > I-;

(3) variations in the cation strongly affect the

spectral shifts when the anion is CI , but not when

it is Br- or l~;

(4) covalently bound halogen atoms are very

much weaker proton acceptors than X , and the

order of spectral shifts found is the reverse of that

in (2). A noteworthy aspect of investigations by Blandamer, Gough, and Symons [526, 527] is that bonding of l~ to HOR was studied with respect to the

effect on l~ (measured in the ultraviolet region) as well as the effect on v(O-li). In the first of these

papers, which dealt with the interaction of tetra- ra-hexylammonium iodide with methanol in carbon

tetrachloride, equilibrium constants were calcu- lated from both the ultraviolet and the infrared

data. Regarding the system as involving an equilibrium between "contact ion pairs" and "modified

contact ion pairs" corresponding to eq 74, they

derived values for n and K. From the ultraviolet

data, n = 0.93 and log K (molar units, 20° to 23°) = 1.199, and from the infrared data, «=1 and log

K=1.137. The ultraviolet absorption curves for different mixtures of methanol and carbon tetra- chloride with a fixed concentration of the salt showed sharply defined isosbestic points. In the second paper, ultraviolet spectra were

determined for solutions in dichloromethane,

with analogous results. Both methanol and tert.- pentyl alcohol were used as proton donors. Being

hindered, the latter alcohol solvated I- less completely than methanol.

Barrow and students [528] studied the association of water with F-, C1-, Br-, and I~ in carbon

tetrachloride. Like their predecessors, they con- cluded that the spectral shifts observed indicate

hydrogen bonding of the proton-donor molecule

to X-. Part of their experimental data are given in table 69. It will be seen that the order of anion effects indicated is

F- = CI- > Br- > I-,

which is a little different from the anion order found

by Allerhand and Schleyer. Table 69 also shows that values of Av become smaller when BU4N+ is re- placed by BU3NH+ or BU2NHJ, once again demonTable 69. Shifts in i'(0— H) of water on hydrogen bonding to X- of R4NX, R3NHX, or R2NH2X in carbon tetrachloride ^

Identity of salt V (free) V (bonded)

BU4NF 3679 3307 372 BU4NCI 3678 3308 370 BujNBr 3676 3376 300 HP4NI 3677 3422 255 BU3NHF 3682 3405 277 BU3NHCI 3679 3405 274 BU2NH2CI 3678 3447 231 BusNHBr 3678 3429 249

« From Table II of S. C. Mohr, W. D. Wilk, and G. M. Barrow,

J. Am. Chem. Soc. 87, 3048 (1965). The temperature was not

specified. "0.033 M.

' I' (free) minus F (bonded).

Reference:

[528] S. e. Mohr, W. D. Wilk, and G. M. Barrow, J. Am. Chem.

Soc. 87, 3048 (1965).

strating that X- is less easily bonded to a neutral

proton donor when already hydrogen bonded to a

cation.

Although the salt-water complex is believed to be re presentable by BU4NX...HOH at low salt concentrations, hydrogen bonding to a second molecule of BujNX is considered probable at high salt concentrations. Barrow also thinks two molecules

of HOH can hydrogen bond to F-. This investigation also included a study of hydrogen bonding of some neutral bases B to water. The

bases included acetone, acetonitrile, dimethylformamide, dioxane, ethyl acetate, ethyl ether,

4-picoline A^-oxide, pyridazine, and pyridine. Conclusions were that B...HOH is formed at low con- centrations of such bases, and B...HOH...B, at higher concentrations.

In a very recent pubUcation [529] Evans and Lo

reported a detailed study of quaternary ammonium

salts containing the anion (ClHBr)- or (ClDBr)~.

Acetonitrile, benzene, and dichloromethane were

used as solvents. Among other findings they con- cluded that in the mixed anion (ClHBr)- the H-Cl

bond is stronger, and the H-Br bond is weaker,

than in the respective homoconiugate anions.

There has been relatively little systematic in- vestigation of heteroconjugation involving anions

other than the halides. In the most complete such

investigation, Kolthoff and Chantooni [429] studied

the conjugation of 3,5-(N02)2C6H3COO- with p- bromophenol and with resorcinol in the solvent

acetonitrile. The conjugation of HSO4 with re- sorcinol was studied less completely. The conditions of the experiments are outlined in Table 70,

which also lists examples of conjugation of CH3COOwith chloroform in the same solvent (already dis- cussed at length in sec. 4.3.5b) and conjugation of HCOQ- and NO3" with methanol in benzene (in110
Table 70. Miscellaneous other examples of heteroconjugate anions (A ...H-A'), studied in solution

Anion, Proton Base or A- donor. Solvent cation Method References HA' present

Acetate CHCI3 CHCI3 EtjN IR [392] CHCI3 CHCI3 Et.NH IR [3931 CHCI3 CHCI3 n-BuNH2 IR [3941 Formate CH3OH BU4N+ IR [521. .522J 3,5-Dinitrobcnzoat€ v^in3V^i>( FtoN ll.l3iN r ^ArOH CH3CN S [429] ^ArOH CH3CN Et4N+ IR [429J Nitrate CH3OH CeHe Bu4N+ IR [521. .522] Bisulfate "ArOH CH3CN Py C [429]

^ ArOH = p-Bromophenol or resorcinol. See accompanying discussion of hetero- conjugate species formed under differing conditions. ArOH = resorcinol.

References:

[392] G. M. Barrow and E. A. Yerger, J. Am. Chem. Soc. 76, 5211 (1954). [393] E. A. Yerger and G. M. Barrow, J. Am. Chem. Soc. 77, 4474 (1955). [394] E. A. Yerger and G. M. Barrow, J. Am. Chem. Soc. 77, 6206 (1955). [429] I. M. Kolthoff and M. K. Chantooni, Jr., J. Am. Chem. Soc. 85, 2195 (1963). [521] J. BufaHni and K. H. Stern, Science 130, 1249 (1959). [522] J. Bufalini and K. H. Stern, J. Am. Chem. Soc. 83, 4362 (1961).

eluded in the studies of Bufalini and Stern). In

all of the examples cited the authors concluded,

from IR spectral shifts, that heteroconjugation in- volves hydrogen bonding of the proton donor to the anion. In the studies of Kolthoff and Chantooni a number of simultaneous equilibria had to be taken into account, including homo- and heteroconjugation

constants of anions and ion pair dissociation constants of homo- and heteroconjugate species of

salts. The experimental work involved conductance titrations, infrared absorption measurements,

and determinations of increases in the solubility

of potassium 3,5-dinitrobenzoate on additions of

p-bromophenol or resorcinol. Some of the principal

conclusions were:

(1) In acetonitrile heteroconjugation of the 3,5- dinitrobenzoate ion with p-bromophenol occurs

both in 1:1 ratio {Kp,- ha' = 360) and 1 : 2 ratio (^A-...2HA'=3900).

(2) Heteroconjugation with resorcinol likewise occurs in 1 : 1 ratio {Kp,- ha' = 340) and 1 : 2 ratio

(/^A-...2HA' = 5100).

(3) The presence of two OH groups in resorcinol

^makes possible the formation of a 2:1 complex,

also. The estimated value of /^a-...hoc6H40h...ais 500. The 2:1 complex is thought to form when

the molar concentration of dinitrobenzoate ex- ceeds that of resorcinol.

(4) By using the stability constants cited, in conjunction with other data and reasonable assumptions, conductance titration curves can be calculated that are in satisfactory agreement with ex- perimental titration curves.

(5) Heteroconjugation, like homoconjugation,

greatly increases ion pair dissociation of a salt BH+A", and consequently, the conductance.

Heteroconjugation is not necessarily beneficial,

however, from an analytical viewpoint. In contrast to homoconjugation, which causes a sharp

maximum to appear in conductance titration curves near half-neutralization of an acid, thus aiding in location of the end-point, heteroconjugation may

cause this maximum to be eliminated and the endpoint to become obscured.

However, heteroconjugation is probably sometimes beneficial in analytical titrations. For example,

Dietzel and Paul [540] and also Trautner and

Shaw [541], when titrating various amines and

alkaloids with 0.05 M toluenesulfonic acid in chloroform, with methyl yellow as the indicator

(an appHcation of Vorlander's "aminometric"

method [542-544]), found that the addition of 1 percent of phenol by volume led to a sharper color

change. As remarked earlier, hydrogen bonding of

phenol to the sulfonate ion seems a reasonable

explanation.

Kolthoff and associates have commented on the

comparative behavior of acetic acid and acetonitrile as solvents [500]. In spite of having a far lower dielectric constant than acetonitrile. acetic

acid is a much better medium for ion pair dissocia111
tion, undoubtedly because it hydrogen bonds to anions such as HSOj and CI", while acetonitrile

cannot do so. As a final example of probable heteroconjugation through hydrogen bonding, mention may be

made of spectrophotometric studies by Steigmar)

and Lorenz [545], using mixtures of a phenolic

indicator (2,4-dinitrophenol or bromophthalein

magenta E) with dialkylammonium or quaternary

ammonium salts in benzene. The authors ascribed

spectral evidence of increased acid-base inter- actions to the formation of "unsymmetrical anionic

triple ions." Their interpretation of results was made somewhat difficult by a competing hydrogen

bonding reaction (dimerization of secondary am- monium salts) and interaction of quaternary am- monium salts with the glass container. Such a reaction with glass has been attributed to ion- exchange [546].

4.6. Concluding Discussion of Hydrogen

Bonding

4.6.1. Definition of a Hydrogen Bond.

Systems Which Form Hydrogen Bonds

It may seem illogical to have postponed defining a hydrogen bond until the end of this lengthy

discussion, but it has been assumed that readers

already have some knowledge of hydrogen bonding.

Moreover, as recently as 1957, at the Ljubljana

Symposium on Hydrogen Bonding, the assembled

experts considered it worthwhile to devote their

final discussion to a possible clarification or re- framing of the definition of a hydrogen bond.

For example, it was suggested that it might be well

to distinguish between hydrogen bonding to an

n -electron donor and bonding to a Tr-electron donor. However, there seemed to be a consensus in favor of Pimentel's "operational definition,"

according to which a hydrogen bond is said to exist when the following two criteria have been met:

(1) there is evidence of molecular association;

and

(2) there is evidence that this association in- volves bonding to a hydrogen atom already bonded

to another atom [102, 547].

An alternative definition based on such arbi- trary criteria as a minimum energy of interaction

or the magnitudes of spectral shifts would ob- viously not be satisfactory.

If Pimentel's definition is adopted, it is clear

that his dual criteria are met in the systems given

special emphasis here — symbohzed by B-H+...A~,

B-H+...B, B-H+...B', A-...H-A, and A-...H-A'—

and from one standpoint this review can be regarded

as a supplement to Pimentel and McClellan's com- pendium [102] on hydrogen bonding, which deals

mainly with bonding of type B...H-A. Cor- respondingly, their book has been regarded as a supplement to our discussion, and it has not been

thought necessary or desirable to discuss formally

examples of hydrogen bonding given thorough

treatment by them — for example, self-aggregation of alcohols and amines, and the bonding of

alcoholic OH to the oxygen of esters, ethers, and

ketones.

The adoption of Pimentel's operational definition of the hydrogen bond implies a simultaneous ac- quiescence in the view that hydrogen bonding is of far wider occurrence than was formerly supposed.

For example, one accepts the view that H-C ini haloforms and mono-substituted acetylenes is a proton-donor group, and that 7r-electron donors

should be grouped with n-electron donors.

The evidence regarding hydrogen bonding that has been brought together in this review, in combination with the evidence compiled by Pimentel

and McClellan, may be likened to an assembled

jigsaw puzzle. There can no longer be doubt that hydrogen bonding plays an extremely vital role in acid-base behavior, leading to the formation of many types of aggregates, that vary widely in stability. Moreover, it seems quite probable that some

of these aggregates exist in water, as well as in aprotic solvents. However, water is not a suitable

solvent for determining their characteristics be- cause it is involved so actively in hydrogen bonding reactions. Thus, as Br0nsted foresaw [13], systematic investigations of individual acid-base sys- tems in aprotic solvents have yielded a more comprehensive and deeper insight into acid-base be- havior than that resulting from studies employing

water and waterlike solvents. Clearly, the overall composition of systems consisting of aprotic diluents mixed with a variety

of proton-acceptor and proton-donor species will depend on such factors as the number of inter- acting species, their relative concentrations, the

relative proton donor-acceptor affinities, and

particle shapes and sizes. To attain desired re- sults in practical situations it is advisable not to add

solvents and reactants when they are not known

to be essential or beneficial. Numerous values of hydrogen bonding equilibrium constants have already been given — mostly in the tables, but also, in some cases, in the text. Section 5 will contain tables of values for log

Xbha in which a variety of bases, acids, and aprotic

solvents are represented. When available, values

of A// and AS will also be given. It will be seen

that these results represent merely a good start toward the accumulation of hydrogen bonding [ equilibrium constants. More, as well as better, data

are much needed. Interesting possibilities for additional studies of hydrogen bonding in aprotic

solvents are suggested by Leffler and Grunwald's

brief discussion of "lifetimes of solvation complexes" [548]. Examples referred to are a benzenepicric acid complex with a lifetime greater than

5 X 10"" sec (20°) and determinations of rate con- stants for the dissociation of dimeric benzoic

acid in carbon tetrachloride and toluene.

112
4.6.2. Some Recent Concepts of Hydrogen Bonding

The theory of hydrogen bonding is certainly rele- vant to our discussion — in fact, it was necessary

in section 4.3.5b to anticipate some of the discussion

of this section. However, there is agreement among

theoreticians that a theory accounting satisfactorily

for all hydrogen bonding phenomena has not yet

been developed. Writing on this subject around

1960, Pimentel said [102]: "Today the theory re- mains a subject of considerable controversy, its qualitative predictive power is limited, and it has

almost no quantitative predictive power. It seems

certain that the full impact of the H bond upon our

general theory of the chemical bond has not yet

been realized."

There is a need for improved communications

between theoreticians and those concerned with

chemical and physical manifestations of hydrogen

bonding. Not infrequently, theoreticians seem to have based their attempts to develop a theory on

a too restrictive concept of the scope of hydrogen

bonding interactions and the range of energy

changes involved, whereas experimentalists have

sometimes been over-ready to adopt a theoretical

explanation without eliminating alternative explanations. It is hoped that theoreticians will be aided

by the foregoing summary of experimental results. Our remaining discussion will represent an attempt

to aid experimentalists wishing to have a brief outline of some modifications which the theory has

undergone and to be directed to more extensive

and authoritative reviews.

Since the formal presentation of Latimer and

Rodebush's theory in 1920 the theory of the hydrogen bond has undergone considerable alteration. Their conception of a weak hydrogen bond was a "hydrogen nucleus held between 2 octets," and

they believed there could be continuous gradations

in hydrogen bonding, ranging from the case of

associated water molecules, where the bonding

hydrogen was thought to be "held quite firmly to the original water molecule," to ammonium chloride, where complete transfer of hydrogen from

chlorine to ammonia was considered to occur [549]. Before long, however, Pauling pointed out that

the hydrogen atom, having only one stable orbital

(the Is orbital) can form only one covalent bond;

hence, in an aggregate of the type B...H-A the

hydrogen bond to B must be largely electrostatic

in character [190b].

As time passed, more and more theoreticians ex- pressed the view that, while electrostatic attraction

must be a factor in hydrogen bonding, it does not account for all hydrogen bonding phenomena. For

example, if it were the sole factor, there should be

a correlation between the dipole moment of the

electron donor and the strength of a hydrogen bond,

whereas Gordy and Stanford found there to be a complete absence of such a correlation in their

infrared studies of the interaction of CH3OD with

electron-donor liquids of varied types [550, 551,

102]. Furthermore, increases in the intensity of the

t'(O-H) absorption band on hydrogen bond for- mation are sometimes far in excess of the increase

expected for a purely electrostatic model [551,

102]. In Coulson's opinion [551 1, the excessively

large increases "imply that during the motion of

the proton there are larger fluctuations of charge

when the hydrogen bond is formed than when it is not formed. This seems to be possible only if charge can move on to and away from the farther atom

(e.g., O2 in O1-H...O2) during the vibrations

of H. Thus effects due to the enhanced intensity of

vibrational transitions in the infrared and the

change of vibration frequency on compression seem

to require what is now often referred to as a delocalization of electrons away from and on to the

Oi-H region from the O2 region . . ." Attempts

have been made to estimate the relative importance

of electrostatic and covalent contributions in hydrogen bonding interactions (for examples, see

refs. [27, 551, 552, 102, 375, 376]), on tlie assumption

of contributing valence bond structures analogous

to those shown below [102]:

A H B (Covalent A-H bond) A" H+ B (Ionic; no charge transfer) A~ H B+ (Covalent H-B bond with charge transfer) A+ B (Ionic; no charge transfer) A H" B+ (Covalent A-B bond with charge transfer).

Pimentel has commented on some defects of this approach, and foresees a probable "turn toward the

molecular orbital approach [102]."

MuUiken's ideas about hydrogen bonding, mentioned in a recent very concise review [553], will undoubtedly be enlarged on in a book said to be in preparation.

A continuing divergence of opinion is indicated

by comments by Longuet-Higgins [554] and Ubbelohde [555] at a "Symposium on Molecular

Interactions in the Liquid Phase" (during which

MuUiken's paper [553] was presented). LonguetHiggins expressed the view that "broadly speaking one may distinguish weak hydrogen bonds, which

occur very widely, from strong hydrogen bonds,

which are much less common . . . ." On the other

hand, Ubbelohde, who believes that resonance

effects become progressively more prominent as hydrogen bonds become shorter and stronger, re- gards it as "unlikely that a reaUy sharp distinction can be made between weak and strong hydrogen

bonds in the sense suggested by Longuet-Higgins . . . ." Instead, he thinks "a whole spectrum

of hydrogen bonds appears to exist." In the last decade or two, increasing emphasis

has been placed on the mobility of the proton within

hydrogen bonds (for example, see refs. [102, 25,

395, 396, 450, 551-552, 556-567] and additional

references cited therein). This is well illustrated by Zimmermann's discussion of "Protonic States

in Chemistry" [562]. He begins with a reminder

that any elementary particle may be described

113
as a wave or as a corpuscle, adding that the wave

picture has been used to advantage by chemists

in describing the behavior of electrons, whereas

the proton, with more than 1800 times the mass of

the electron, is generally represented as a cor- puscle located in a more or less fixed position

in a molecule. He then points out that the proton

is smaller than other atoms by about five orders of

magnitude, concluding that "it therefore seems

reasonable to base the description of protons in chemistry on the wave picture as well. According

to this view, states are supposed to exist for pro- tons, much as in the case of electrons, in which

protons are delocalized" [562]. Applying this idea to intermolecular hydrogen

bonding, Zimmermann states that a hydrogen bond

represents the joining of a Br0nsted acid-base

pair to give B...H-A, and it is reasonable to pre- sume that there can be proton transfer within

the bond to yield a polar form, B+-H...A~. Some

limiting conditions which can be visualized for the potential energy and for protonic states in hydrogen bonds are indicated by the four hypothetical potential energy curves in figure 24. The

different situations represented are as foUows:

-A B —H"" 'A B A B

INTERNUCLEAR DISTANCE

Figure 24. Hypothetical potential energy curves for four classes of hydrogen bonding, indicating limiting conditions for proton

states; suggested by Fig. 5 of H. Zimmermann, Angew. Chem.

Intern. Ed. Engl. 3, 157 (1964). (Ref. No. 562)

(a) A weak proton donor HA is bonded to a weak

proton acceptor B. The proton remains near A

in its ground state. The complex, B...H-A, is referred to as a "hydrogen-bonded complex."

(b) The proton donor and acceptor are both

strong. The proton shifts from A toward B. The

product, symbohzed by B-H+...A~, is now com- monly called "a hydrogen-bonded ion pair."

(c) The proton donor and acceptor are both in- termediate in strength. The proton is about equally

hkely to be located at a covalent bond length from

B and a covalent bond length from A. The potential

energy troughs are about the same in depth, and

the barrier separating them is low, hence the proton can shift toward B or toward A by "leaping over" the barrier or "tunnehng through." This means that B...H-A will be in equilibrium with

B-H+...A-.

(d) B and A are both extremely good electron

donors, and the hydrogen bond is very short. In

its ground state the proton stays about mid-way be- tween B and A. The homoconjugate anion HF^^ is considered the prime example (recall sec. 4.5.5). The "tunnel effect," not possible according to classical theory but quite plausible according to the quantum-mechanical picture, has become an

important concept of modern chemistry. It is thought to have been demonstrated experimentally

in studies of the infrared and Raman spectra of

imidazole [187, 562]. It has provided an explanation for the "inversion" of the ammonia molecule,

is frequently invoked in discussions of the kinetics

of proton transfer processes iior illustrations, see

refs. [559, 563, 567]), and it may have major bio- logical implications [561]. The concept represented by figure 24(c), which

Zimmermann refers to as "protomerism" [562], seems to have been suitably applied in his studies

of the imidazole spectrum. As we saw in section

4.3.5b, Barrow and Bell [395, 396] interpreted in- frared and electronic spectral absorbance data as demonstrating that triethylammonium p-nitrophenolate and pyridinium dichloroacetate are au- thentic examples of salts which exist in aprotic

solvents as equilibrium mixtures of the two pro- tomers, B...H-A and B-H+...A-. Furthermore,

they believe that many other salts probably be- long in the same category. The writer has given

a few specific reasons for unwilhngness to accept

these particular conclusions, and beheves that the point at issue is too important for final ac- ceptance of Barrow's interpretation without more

extensive experimental documentation and systematic elimination of alternative explanations.

Appropriate additional information would include

acid-base equihbrium constants and conductance

data.

It is appropriate to note here that since 1960

proton transfer reactions of excited molecules have

received increasing attention. Weller and his associates [568, 571] and Mataga, Kawasaki, and

Torihashi [569, 570] are prominent among those

pioneering in the field. Solvents used in the in- vestigations include acetonitrile [569, 571], ben- zene [569], n-butyl chloride [571], chlorobenzene [569], o-chlorotoluene [571], ethylene

chloride [569], methylcyclohexane [568, 571],

toluene [571] and water [568].

5. Acidity and Basicity Scales in Aprotic Organic Solvents

5.1. Early Attempts to Determine Relative

Acidities

New ideas about acid-base behavior which were propounded by Hantzsch and by Br0nsted, starting

in the early 1920's, made a strong appeal to many

young chemists of that period. With the widening

recognition that the presence of water is not a prerequisite for acid-base behavior there began

varied attempts to determine relative strengths of

114
acids dissolved in solvents like benzene, chloro- benzene, chloroform, and ether. Hantzsch and his associates, using two kinetic methods and a static method, demonstrated that acids which appear

equally strong in aqueous solutions (including

perchloric, hydrochloric, and nitric acids) display

markedly different strengths in organic solvents

such as those named (for examples and discussion, see refs. [18, 426, 572, 19, 22, 425]). One

criterion of relative strength was the effectiveness of an acid in accelerating the decomposition of ethyl

diazoacetate. The static method usually employed

methyl yellow, which forms fairly stable red salts with the stronger acids. The relative strength of an acid was denoted by the quantity needed for conversion of the original yellow color to red, or, alternatively, by the volume of an "indifferent"

solvent that had to be added to effect dissociation

of the red salt into the constituent base and acid. Weissberger and Fasold [22] tried to determine the

equilibrium constant for 1:1 association of methyl

yeUow with trichloroacetic acid in benzene and in chloroform (and also studied the influence of ad- denda on this equilibrium in benzene). The results did not correspond to a simple 1:1 equilibrium.

With the benefit of present knowledge, summarized

in earlier sections, it is clear that this was not a

suitable acid-base system for demonstrating the

validity of Hantzsch's basic concepts. Moreover,

at the time of the experiments spectrometric equipment enjoyed by present-day chemists was not

available. Investigators of the Br0nsted school also per- formed static experiments using indicators and

studied acid-catalyzed reactions in such solvents

as benzene and chloroform (for illustrations, see

refs. [10, 491, 357, 573, 492, 493, 355]; conductance

and potentiometry were also employed in attempts

to derive acidity scales [340]. Reference was made

in section 4.5.4b to experiments in which the catalytic effects of a series of carboxylic acids and

halophenols, dissolved in chlorobenzene, were

related to their aqueous strengths by a simple

equation (into which, unaccountably, picric acid

did not fit). In other experiments, employing

2,6-dinitrophenol and bromophenol blue (the

change from colorless to yellow [29]), Griffiths also found [355] that the strengths of a series of car- boxylic acids in chlorobenzene closely paralleled

their strengths in water. Somewhat later, Bodforss [574] had the novel

idea of using an optically active base (brucine)

and measuring the changes in optical rotation on

its reaction with a large group of aliphatic and aro- matic carboxylic acids in benzene and in chloro- form. "Empirical acid constants" indicated a general, though imperfect, parallehsm of strengths

in the aprotic solvent and in water. Bodforss con- cluded that the equihbria did not appear to con- form to the mass action law.

The first successful demonstration of the cor- rectness of Hantzsch's claim that the fundamental

acid-base reaction in an aprotic solvent is an

association seems to have been obtained in a spec- trophotometric investigation of reactions of the

indicator acid bromophthalein magenta E in benzene with bases such as triethylamine or 1,3- diphenylguanidine [14|. Equilibrium constants were calculated for the 1 : 1 association of B with

HA to give BHA. Shortly thereafter, similar equilibrium constants were obtained for the reaction of tribenzylamine with picric acid [23, 20| and with

trinitro-m-cresol [20] in benzene. The equilibrium involving tribenzylamine, picric acid, and

tribenzylammonium picrate was studied both

dielectrometrically [23] and spectrophotometrically

[20], with results in excellent concordance. One

reason why these systems were suitable for testing

Hantzsch's postulate is that the reference acids are monomeric in an aprotic solvent. There are no competing hydrogen bonding reactions of significance

at the concentrations employed; moreover, any

dissociation of the salt into ions is negligible in extent. Other examples soon followed. Among the first was Bell and Bayles' survey [354] of the reactivities of various aliphatic and aromatic amines with

2,6-dinitrophenol or bromophenol blue (first step, colorless to yellow) in chlorobenzene.

5.2. Log /<LBHA Scales of Acidity and Basicity

Let us remind ourselves that the reason why pH

values can be determined for aqueous solutions is that water can function both as a proton donor

and as a proton acceptor and interacts in the ap- propriate capacity with dissolved acids and bases.

Although chemists have been known to inquire

about the "pH of an oil" it becomes clear, on re- flection, that there is no such thing as the pH

(—log a//+) of, say, a benzene solution, and neither can there be pKa and pKb values in the same sense

as in water. In fact, in a solvent of this type self- association of acids and bases is more likely to assume significance than their ionic dissociation.

Since the only way to determine the strength

of an acid is by ascertaining its readiness to react with a base, a suitable reference base must, obviously, be added when the solvent is of an aprotic

type. Similarly, a suitable reference acid is required

for determining the strength of a base.

The simplest kind of acidity or basicity scale

consists of a series of equilibrium constants cor- responding to the reaction

B + HA ^± BHA, (30) base acid salt

with B or HA constituting the reference compound

as the case may be. Then

/i:BMA=[BHA]/([B][HA]). (75)

In the tables of 1:1 acid-base association constants

to be introduced presently, the values of ^bha will

115
be seen to fall mostly in the range 10- to 10*^; hence,

log /^BHA is generally in the range 2 to 7. The relation of such an acidity scale to the

aqueous pK„ scale is easily comprehended by

supposing the interaction of an acid HA with water

to be representable as a two-stage reaction (eqs

20 and 76): H20 + HA^H20H+...A- (20)

base acid

H2OH+ . . . A- + 2H2O ^ H2OH+ . . .OH2

+ A-...HOH. (76)

(It will be understood, of course, that eqs 20 and 76

are intended to represent the hydrogen bonding

role of water in the simplest possible terms.) A

crude representation of the overall equilibrium is

HAa^ + aq^H.OH+q + Aaq. (77)

As we know, aqueous ionization constants for organic acids are commonly in the range 10"^ to 10" and the symbol pKa (—log Ka) has been

adopted to eliminate the inconvenience of having

to use exponents and negative signs. A low aqueous

pKa value will therefore correspond to a high log

iCgHA value. It will be obvious that various log Kbha

scales for acids and bases in any aprotic solvent

cannot be merged into a single scale corresponding

to the aqueous pKa scale. The symbols "'pK" (as in "pKrei") "p" (as in PHCIO4, PRNH2, etc.) are sometimes encountered

in discussions of acidity and basicity scales in aprotic solvents [575-579]. For example, Tremillon

and associates [575-578] refer to numbers here

called "log A^bha" values as "pK"" values, arriving

at their terminology in the following manner:

(1) The acid-base reaction (in chloroform, let us

say) is formulated as an association:

B (base) + HA (acid) ^ BHA (sah). (30)

(2) Next, the equihbrium constant for the acidbase reaction is formulated as a dissociation:

/^ = (B)(HA)/(BHA) (78)

(3) Finally, Kaissocn., calculated according to eq

78, is converted into its negative logarithm, or pK value. We do not see any advantage in extending the use of the symbol "p" except when by so doing a negative coefficient can be avoided. In fact, its use

for both aprotic and amphiprotic solvents is likely

to lead to confusion, because the ''pK"' (that is, log A^bha) values in an aprotic solvent become

larger with increasing acidic strength instead of

smaller, as in water. A more serious objection is the likelihood that the fundamental relationship

between 1:1 acid-base association constants in aprotic solvents and ionic dissociation constants

in water and like solvents will become obscured.

A further argument in favor of using log A^bha for aprotic solutions is that proton donor-acceptor

interactions involving strong acids and bases are thereby placed in a proper relationship with weak

hydrogen bonding interactions, and also with

electron donor-Lewis acid interactions.

A complication encountered in solution chemistry

is that in the computation of equilibrium constants

the units of concentration vary with the preference

of the investigator. One worker may employ molar

units, a second may use molal units, while a third

may adopt the mole-fractional system. To differentiate the three sets of equilibrium constants, the

symbols Kc, Km and (or Kn) are ordinarily used.

(See sec. 4.1 for useful approximate relations

between the three equihbrium constants.) For

measurements involving aprotic solvents the molar

scale is by far the most popular.

Substitution of the expression "equilibrium

quotient"" for "equilibrium constant"" has occasionally been advocated for systems in aprotic

solvents, because activity coefficients are not incorporated in the equations. However, few workers have adopted this recommendation. For one

thing, the solutions studied are very often much

more dilute than aqueous solutions characterized

as "dilute," and it is considered justifiable to regard the activity coefficients as not differing measurably from 1. It should also be emphasized that under proper experimental conditions acid-base

equilibrium "quotients" have been found to be

satisfactorily "constant" over a substantial range

of concentrations (for examples, see refs. 32 and

254a). An extrapolation procedure is not used.

(Among conditions essential in "properly per- formed" experiments are care to exclude impurities

and maintain a constant temperature, selection of reactants having neither very great nor very low

affinity for each other, and avoidance of a large

imbalance in molar ratios of reactants.) See further

discussion of this topic in ref. 254a.

Izmailov has determined the relative strengths

of acids in such solvents as benzene, chlorobenzene,

carbon tetrachloride, and ethylene chloride potentiometrically [579], according to a procedure de- vised by LaMer and Downes [340]. The relative strengths are indicated by numerical values of

"p/Crei", where p/Crei = -log (/^HA/HA°)- The identity

J of HA° is not always the same — for example, in some experiments it was picric acid, and in others,

benzoic acid. It should be noted that such experiments sometimes employ solvents which, though

referred to as "aprotic," actually contain significant

quantities of other types of solvents.

5.3. Methods Used in Determining Values of

Log Kbha

Whereas infrared absorption spectroscopy has , provided the principal means for studying weak r hydrogen bonding interactions, absorption spectros- ^ copy in the visible and ultraviolet has been the <

116
principal method for studying interactions of strong

or moderately strong acids and bases. Figure 23

(introduced earlier) shows some of the absorption curves obtained during determination of log /^bha

for the association involving trinitro-m-cresol and

tolyl dipropyl Nile Blue base in benzene. In these

experiments successively larger amounts of trinitro- m-cresol were added to tolyl dipropyl Nile Blue

(which was 10 M in each mixture), until there was no further change in the absorption curve. The

experimental curve for a mixture containing 1600

equivalents of the acid was very similar to the

calculated "limiting curve." Moreover, there was a

well-defined isosbestic point, supporting the

assumption that only two deeply colored species were present [482J. A related example is shown in figure 25, in which

the reference compound was the acid (p-nitroWavenumber, cm-' x 10-3 35 30 25 1

1 ' 1 1 1 1 1 1 1 1 1 1330

300 350 400

X, nm Figure 25. Spectral absorption curves of mixtures of p-nitro- phenol (10"'* M) with triethylamine in benzene; used in deter- mining the equilibrium constant (Kbha) for the 1:1 amine-phenol

association (25°) [17].

phenol). Here again, the experimental and theoretical limiting curves were the same, and there was a sharply defined isosbestic point. Some other indi- cators also used as reference acids in benzene are represented in figures introduced earlier: For example, bromocresol green (fig. 13); bromophthalein

magenta E (fig. 14); picric acid (fig. 15); and 2,4-

dinitrophenol (fig. 17). A well-defined isosbestic point cannot be regarded

as proof that a solution contains only two colored

species. In some instances, a more extensive study

of the system will indicate that secondary changes

are occurring which have little effect on the spec- trum except when one of the reactants is present

i at a high concentration. The case of methyl yellow

\ deserves special mention, since this indicator

(LXXXVIII) has probably been used more extensively than any other indicator in studying acid-base

equihbria. The yellow color of the base is attributed

to the azo structure, and the red color seen on the

addition of acids can be explained by the resonance

hypothesis, as indicated by (LXXXIX, a and b). If a proton should add to the amino nitrogen (XC) the

yellow color would not be affected.

(LXXXVIII)

(CHj)2N-C^lI^-N=l'l-C^Hj ^ (CH^)2N=C^H^=N-A-C^K5

(LXXXIXa) (LXXXIXb)

(CH3)2A-C^H^-N=N-C^H5

(XC)

Through systematic investigations of Brode and

co-workers [582, 583] it has become well estabhshed

that in solvents like benzene the stable trans form

of methyl yellow is partly converted to the cis form,

in light. Moreover, there is extensive experimental

evidence that Hantzsch and Burawoy were correct when, in 1930, they asserted that the red and yellow

species of salts both form when an acid is added to methyl yellow [584-586]:

Davis and Hetzer [632] studied the interaction of

methyl yellow with diphenyl phosphate in benzene,

obtaining the system of curves shown in figure 26.

Wovenumber, cm"' « lO"^ 30 25 20

•

'—I • ' '

1

' ' q

300 400 500 600

Wavelength, nm

Figure 26. Spectral absorption curves of mixtures of methyl

yellow (2.5X10 * M) with anhydrous diphenyl phosphate in benzene [632].

Great pains were taken to exclude daylight, and it seemed evident that the intensity of the spectrophotometer source was too low to cause measurable

isomerism to the cis form. Although two seemingly

very sharp isosbestic points were obtained, and a value for log /Crha (1-57) was calculated by assuming the double-peaked band near 540 nm to be due

solely to a red salt, irregular changes in the absorb- ance near 400 nm indicated that part of the indicator was probably converted to the yellow salt. In

117
short, this indicator shows complicated behavior

and does not seem well suited for exact work when

used in an aprotic solvent. Moreover, methyl yellow

is not a very strong base, and with most organic

acids an extremely large excess is required for a measurable change towards red. Dimerization and

homoconjugation of acids may then occur. Sometimes it is assumed that "electronic spectra

are incapable of showing any appreciable difference

between free ions, ion-pairs, or any other aggregates

containing the same absorbing species" (see ref. [85] and literature cited). Section 4.3.4 presented

substantial evidence that this generalization does

not hold for aprotic solvents. These were by no means isolated examples, and it should constantly

be borne in mind that aprotic solvents are "differentiating" media. Whenever uniformity of salt behavior is assumed, the results will be, at best, only

approximately correct, and an opportunity for extending knowledge will be forgone.

Not infrequently the limiting curve for an acidbase association cannot be determined experimentally, because one of the reactants is only

sparingly soluble or the association constant is very

small. In some of these cases it is useful to apply

a method proposed by Rose and Drago [580] for calculating the limiting absorbance and the association constant. Another method, originally devised

by Benesi and Hildebrand [72] for determining the

formation constant of the 1:1 complex of benzene

with iodine, has been widely used for weak complexes, but some applications of this method have

been criticized; Trotter and Hanna [581] have contributed one of the most recent discussions of this topic.

Successful use of the dielectrometric method for studying the acid-base dissociation of tribenzylammonium picrate and determining 1/A^bha [23] was

mentioned a little earlier. A few other methods

occasionally used will be mentioned in discussing

the tables of log Kbha values.

5.4. Tables of Log /^bha. A//, and AS Obtained

Using Aprotic Solvents

About twenty tables of values of log A^bh^a and —

when available — values of A// and AS, constructed

after an extensive (but not exhaustive) survey of

the literature, will be presented in this section and

discussed briefly. Results for only one or a few com- pounds were not included unless they helped to validate general conclusions. Most of the available

values are expressed in molar units, but some values

in mole-fractional units are also included. Except

when noted otherwise, AH and AS were calculated

from data for log K^ha at two or more temperatures.

In all cases AH is in kcal/mole and A5 is in entropy

units. In the present state of knowledge the tabulated values should be regarded as approximate, and

in some cases, as provisional. However, it seems

probable that many of the log /^bha values are as accurate as the available pKa values for the acid

and base concerned.

The numerical magnitude of values of log ^bha^

AH, and A5 would be expected to increase with

the ability of the reactants to donate or accept

protons. Most of the tables include pKa values for aqueous solutions to facilitate comparisons of acid

and base strengths in water and the aprotic solvent

concerned.

Examples are included for solutions in benzene,

anisole, chlorobenzene, miscellaneous aromatic

solvents, heptane, carbon tetrachloride, chloroform,

and a few other organic solvents. The tables will be presented according to the sequence of solvents

just given.

5.4. 1. Aromatic and Substituted Aromatic Hydrocarbons

a. Benzene. Table 71. The values of log ^bha were obtained

using three reference bases: pyridine (Py), 1,3- diphenylguanidine (DPG), and triethylamine (EtsN).

The strengths of the bases in water increase in the

order named. The acids represented are phenol,

p-nitrophenol, the six dinitrophenols, and picric acid. In part, the stabilities of complexes parallel the

aqueous strengths of the constituent bases and

acids, but there are noteworthy exceptions. For

example, according to Swain and Brown's results (using a dynamic vapor pressure method [588]), 2,4- dinitrophenol has far less ability than p-nitrophenol

to hydrogen bond to pyridinp in benzene, in spite of its much smaller pKa value. Furthermore, on

turning to the last two columns it will be seen that lower constants were obtained with triethylamine as the reference base than with diphenylguanidine,

although according to the pKa values of these bases

triethylamine is the stronger base in water. An investigation of the relative strengths of the

isomeric dinitrophenols in benzene in terms of association with triethylamine [16] showed that the

log Kbha values are not linearly related to the pKa

values of these acids, probably because internal

hydrogen bonding (chelation) in orf/io-nitrophenols

interferes with intermolecular hydrogen bonding.

A similar conclusion can be drawn from the parallel

investigation in which diphenylguanidine was the ? reference base [17]. When the log ^bha values for < B — diphenylguanidine are plotted (as y) against

the corresponding values for B = triethylamine all

' of the points except that for HA = 2,6-dinitrophenol

^ fall on or very close to a straight line. Excluding

the data for 2,6-dinitrophenol, the least squares

equation is j

7=0.596 + 0.936 jc. (79) j

The point for 2,6-dinitrophenol would fall on the line

if log A^bha for B = Et.iN were greater. This acid is "hindered," and its association with triethylamine seems likely to be obstructed by "F-strain," as in

j various reactions of triethylamine studied by H. C. , Brown and associates [635]. ^ 118
Table 71. Log ^lbha for phenol and some related acids in benzene {molar units, 25°) ^

Acid ipKa")

Phenol (10.00"^) p-Nitrophenol (T.IS"") 3,5-Dinitrophenol (6.69)

3.4- Dinitrophenol (5.42)

2.5- Dinitrophenol (5.21)

2.3-Dinitrophenol (4.96)

2.4- Dinitrophenol (4.11)

2.6-Dinitrophenol (3.71) Picric Acid (0.33 ^)

Reference Base:

Py (5.17)

" 1.26 " 1.56 "2.04

0.1

'3.62

DPG(IO.O)

"3.49

'3.92 "3.06

'4.28 "3.83

'5.77

Et.,N (10.87)

' 2.46 " 2.42

'3.07

'3.61 '2.62 '3.87

- 3.49 ' 3.50 ^3.47

'4.94

^ Three reference bases were used, namely, pyridine (Py), 1,3-diphenylguanidine (DPG), and triethylamine. Numbers in parentheses are pK,, values from the literature.

*" With the exceptions noted, the pKa values are as given in

ref. [587].

' From ref. [271]. From ref. [588]; a dynamic vapor pressure method was used. From ref. [589]; obtained from dielectric data at 25° and 40° which also yielded AH — — 9 and AS = — 24. 'See ref. [17].

^ See table 75.

" See ref. [590]. 'See ref. [16]. ^See ref. [591].

X See ref. [362]. 'See ref. [592].

References:

[16] M. M. Davis, J. Am. Chem. Soc. 84, 3623 (1962).

[17] M. M. Davis, unpublished results.

[271] R. A. Robinson and R. H. Stokes, Electrolyte Solutions, 2d ed.. Academic Press, Inc., New York, N.Y., 1959.

[362] M. M. Davis and M. Paabo, J. Res. NBS 67A, 241 (1963).

[587] R. A. Robinson, M. M. Davis, M. Paabo, and V. E. Bower,

J. Res. NBS 64A, 347 (1960).

[j [588] C. G. Swain and J. F. Brown, Jr., .|. Am. Chem. Soc. 74,

2691 (1952).

[589] R. J. Bishop and L. E. Sutton, J. Chem. Soc, p. 6100 (1964).

[590] M. M. Davis, M. Paabo, and N. Matheny, unpublished

results.

[591] R. G. Pearson and D. C. Vogelsong, J. Am. Chem. Soc. 80, 1038 (1958).

[592] T. Jasinski, T. Misiak, and T. Skarzynska, Roczniki Chem.

39, 1549 (1965).

In summary, it seems reasonable to assert that strengths of acids and bases in terms of values of

log Xbha tend to parallel the strengths in water

indicated by pKa values, but that deviations occur when there is a strongly competing hydrogen bonding reaction such as chelation and when one or both

of the reactants is hindered.

Table 72. Having failed to find a suitable indicator

base for determining relative strengths of aromatic

and aliphatic carboxylic acids in benzene Davis

and Schuhmann [360] devised an indirect procedure

in which diphenyl- or di-o-tolylguanidine is the

reference base and the carboxylic acid and bro- mophthalein magenta E are competing proton

donors. Davis and Hetzer developed the method

further and applied it in determining values of

log ^BHA for some 40 aromatic carboxylic acids [32],

and later the measurements were extended by

Davis and Paabo to additional aromatic acids and some aliphatic acids [254a|. The 1:1 complex of

bromophthalein magenta E (yellow) and diphenylguanidine (colorless in the visible region) is magenta

in benzene solution, with \max near 540 nm; for

this association log Kbha = 5.41 (see table 73). The

color shifts toward yellow, of course, on the addition

of a carboxylic acid (colorless in the visible region).

From absorbance values for solutions containing the

three components in varying, but not greatly dif- ferent ratios, the log A^bha values given in table 72 were calculated. In some cases a correction was

applied for the content of carboxylic acid as the

dimer, but for the most part competing hydrogen

bonding reactions were thought to be negligible

in extent. Four linear relations emerged when values of log /^assocn. (log /^bha) in benzene were plotted

against pKa in water. These were for meta-, para-, mono-ortho-, and di-ort/io-substituted benzoic acids. Equations for the first three lines are given below.

The equation for di-ori/io-substituted acids was

regarded as only a rough approximation [32].

roeJa-Substituted acids [254a]: Log Kbha= 14.37-2.17pA:„ (80)

orfAo-Substituted acids [254a]: Log A:bha= 10.06- 1.30pA:„ (81)

para-Substituted acids [32]: Log /^bha = 12.96- 1.80pA:„ (82) These three linear relations are shown in figure 27.

7

-

Log Kassocn. (DPG, Benzene)

o-Benzoylbi

2-Naphthoi

trans-Cinnomic

Aliphatic Acids

enzoic

zoic

6 5 4 3 2 pKa (Wafer)

Figure 27. Relative strengths of some carboxylic acids in benzene and water at 25°. The strength in benzene is expressed

as the value of log Kbha (log Knssorn.) where B is 1 ,3-diphenylguanidine. Equations for the three lines are given in the text. Open circles, results from ref. 32. Half-filled and filled circles, results given in ref. 254a.

The dotted, unlabeled line is the one obtained for para-substituted acids. The relation for metasubstituted acids is considered "normal," and the

existence of separate relations for the isomeric

119
Table 72. Comparative values of pKg (water, 25°) and log Kbha {benzene, B = DPG, molar units, 25°) for some carboxylic acids ''

Name of acid Log Kt, Name of acid Log Kb

Benzoic acid series'*

Benzoic

o-Acetoxybenzoic (aspirin).... o-Aminobenzoic (anthranilic), m-Aminobenzoic

p-Aminobenzoic

o-Benzoylbenzoic'' o-Bromobenzoic m-Bromobenzoic

p-Bromobenzoic

o-Chlorobenzoic

m-Chlorobenzoic

p-Chlorobenzoic

m-Cyanobenzoic

p-Cyanobenzoic

m-Dimethylaminobenzoic

o-Fluorobenzoic m-Fluorobenzoic

p-Fluorobenzoic

o-Hydroxybenzoic (salicylic). m-Hydroxybenzoic

p-Hydroxybenzoic

o-Iodobenzoic m-Iodobenzoic

p-Iodobenzoic

o-Methoxybenzoic

(o-anisic)

m-Methoxybenzoic

(m-anisic) p-Methoxybenzoic

(p-anisic)

4.Hp..iO 90

3.48 5 72

4.95 4.94

4.75 (4.35) 4.93 4 45

3.54 6.44

2.85 6.17

3.81 6.06

3.97 5.86

2.94 6.08

3.83 6.06

3.98 5.82

3.60 6.56

3.55 6.53

5 1 r4 'M\ 4.89

^ 97 77

3.86 5.94

4.14 5.61 3.00 7.45

d. OR O.oD

4.58 4.94

2.86 6.22

3.85 6.05

3.98 5.85

4.09 3.7?

4.09 5.38

4.47 4.92

o-Methylbenzoic

(o-toluic) m-Methylbenzoic

(m-toluic) p-Methylbenzoic

(p-toluic)

o-Nitrobenzoic m-Nitrobenzoic

p-Nitrobenzoic

2,4-Dichlorobenzoic

2,6-Dichlorobenzoic

3,4-Dichlorobenzoic

2,6-Dimethoxybenzoic.

2.4- Dimethylbenzoic^ .

2.5- Dimethylbenzoic'' .

2.6- Dimethylbenzoic

(2,6-xylic)

3,5-Dimethylbenzoic

(3,5-xylic, mesitylenic)... 3,5-Dinitrobenzoic 2,4,6-Trimethylbenzoic (yS- isodurylic, mesitoic)

Miscellaneous other acids'"

Acetic

Isobutyric

Trimethylacetic (pivalic).. Laurie

Stearic

frawi-Cinnamic

2-Furoic

1- Naphthoic

2- Naphthoic

3.91

4.27

4.37

2.17 3.49

3.42

2.76

1.82

3.64

3.44

4.18

3.98

3.25

4.30

2.82

3.44

4.76

4.86

5.05

4.44

3.17

3.69 4.16

4.94

5.13

5.03

7.44

6.82 6.80

6.55

7.34

6.48

4.92

4.72 (4.75*) 4.86

5.24

5.03

8.23

5.00

4.45*

4.44*

4.46* 4.35*

4.37*

4.86 (4.91*) 6.07

5.31 (5.34*) 5.39 (5.41*)

"Log ^LBHA corresponds to B + HA^BH+...A , where the reference base, B, was diphenylguanidine (DPG); bromophthalein magenta E was present as a reference acid (HA')- See accompanying discussion. pK„ values in parentheses are "corrected" values, derived by substituting values of log /^bha in eq 9 of ref. [32]. Log K^BHA values marked with an asterisk were derived by

applying a correction for the monomer-dimer equilibrium of HA, as explained in the paper cited.

** Except as indicated otherwise, all log ^bha and pK,, values

are taken from Table 1 of M. M. Davis and H. B. Hetzer, J. Res. NBS 60, 569 (19,58). (Ref. No. 32.)

^' Log A^BH^ and pK,, values are from Tables I to III of M. M.

Davis and M. Paabo, J. Org. Chem. 31, 1804 (1966). (Ref. [254a].)

References:

[32] M. M. Davis and H. B. Hetzer, J. Res. NBS 60,569(1958).

[254a] M. M. Davis and M. Paabo, J. Org. Chem. 31, 1804(1966).

ortho- and para-substituted acids is attributed partly

to such factors as chelation and steric hindrance,

and partly to solvation of substituents and anomalously enhanced strengths of orfAo-substituted

acids in water. Readers are referred to the original

papers for details and further discussion.

Table 73. At the top of this table are values of

log Kbha, ^H, and AS for some of the reactions

represented in the preceding table, namely, those

involving bromophthalein magenta E, benzoic acid, ii the three toluic acids, and the three monochloro- i: benzoic acids. At the bottom of the table are corre- i sponding values for association of triethylamine \ with bromophthalein magenta E and with benzoic f acid. Points worthy of special comment are: ( (1) In benzene, most values of A// are in the range — 14 to — 19 kcal/mole, whereas the corresponding t enthalpy changes for ionization of these aromatic \

120
Table 73. Values of log Kbhai ^H, and AS {molar units, 2.5°) for some 1:1 asso- ciations in benzene"

Acid (pKa) Log /lbha A///0.298 Refer- AS ences

Reference Base == 1,3-Diphenylguanidine

BPM-E''(4.15'^) 5.34 [14, 5) "5.34 "-15.9 "-28.8 1.85 [594J

4.1 — 110.0ft 9ft 9 [32, 33] 5.41 -14.1 -22.6 2.1 [15] CeHsCOOH (4.20) 5.30 -16.2 -30.1 1.8 [32, 33] 5.32 -14.1 -22.9 2.1 [15] 0-CH3C6H4COOH (3.91) 4.97 -16.5 -32.5 1.7 [32, 33] m-CHaCeH^COOH (4.27) 5.18 -16 5 -31.5 1.8 [32, 33] P-CH3C6H4COOH (4.37) 5.13 -16.0 -30.1 1.8 [32, 33] 0-C1C6H4COOH (2.94) 6.10 -18.8 -35.1 1.8 [32, 33] m-ClC6H4COOH (3.83) 6.08 -19.0 -36.0 1.8 [32, 33] P-CIC6H4COOH (3.98) 5.84 -18.0 -33.7 1.8 [32, 33]

Reference Base = Triethylamine

BPM-E'^ (4.15'^) 4.36 [14] 4.44 -15.3 -30.9 1.7 [15] Benzoic (4.20) ^3.58 -11.0 -20.7 1.8 [15]

^ AHc and ASc signify that these constants were derived from data expressed in molar units of concentration. As is customary, A//<. is in the units kcal/mole and

ASc is in entropy units. "BrM — E signifies bromophthalein magenta E (3',5',3",5"-tetrabromophenol- phthalein ethyl ester).

' An estimated value; see footnote 29 of ref. [32].

" Originally expressed in mole fraction units; conversion to molar units was made

by this author.

* An almost identical value was obtained using 2,5-dinitrophenol (in place of bromophthalein magenta E) as the reference indicator acid.

References:

[5] M. M. Davis and H. B. Hetzer, J. Res. NBS 46, 496 (1951).

[14] M. M. Davis and P. J. Schuhmann, J. Res. NBS 39, 221 (1947).

[15] M. M. Davis and M. Paabo, J. Am. Chem. Soc. 82 , 5081 (1960).

[32] M. M. Davis and H. B. Hetzer, J. Res. NBS 60, 569 (1958).

[33] M. M. Davis and H. B. Hetzer, J. Res. NBS 65A, 209 (1961). [594] Results of A. R. Anderson and W. F. K. Wynne-Jones, as given in the Ph. D.

thesis of A. R. Anderson, The University of Newcastle upon Tyne, 1954, and communicated by Lord Wynne-Jones to M. M. Davis in June, 1964.

acids in water are positive in sign (with one exception) and their magnitude is only about 0.02 to 0.3 kcal/mole (see ref. [33], table 3).

(2) In benzene, AG is negative in sign, whereas

AG for ionization in water is positive.

(3) In benzene, the ratio ^H|T^S (where T=298

°K) is substantially constant for the reactions

represented. This is not true for the ionic dissociation of the acids in water. Interestingly, the ratio A///0.298AS for dimerization of the same or closely

related aromatic acids in benzene, just as for their association with diphenylguanidine, is close to 1.8. (See further discussion in ref. [33].) A linear relationship of A// and "A5 in a series of

association reactions is customarily held to signify that with increasing strength of the bond between a proton donor and acceptor (or between a Lewis

acid and an electron donor) there is increased re- straint on motions of the component parts (for example, see ref. [102]). These results support our

121
assertion that log K^ha values for an aprotic solvent

like benzene are likely to be a better index of

intrinsic strengths of acids and bases than aqueous

pKa values. Furthermore, it is evident that the

maximum enthalpy change in hydrogen bonding in- teractions has for a long time been underestimated.

Attention should also be directed to the striking

finding [15] that in benzene log ^bha values for the

reaction of bromophthalein magenta E with di- phenylguanidine (5.41) and the reaction of benzoic

acid with diphenylguanidine (5.31) are about the

same, paralleling the similarity in the pKa values

(4.15 and 4.20, respectively) of the acids. In contrast, benzoic acid seems much weaker than bromophthalein magenta E when the reference base is

[596] estimated the isothermal heats of neutralization by hydrogen bromide in acetonitrile (24°) for a large group of aliphatic and aromatic amines from

the "initial slopes" of thermometric titration curves, and subsequently. Mead [597] performed analogous

triethylamine (log K^ha 3.58 versus log K^ha 4.44). To explain this finding, it has been suggested that the diphenylguanidinium ion, (PhNH)2C=NH^,

can form two liydrogen bonds to a carboxylate ion, one to each oxygen (LX), whereas Et3NH+ can form only one bond (LIX) [15]. The plausibility of

this hypothesis is enhanced by the observation of

Streuli and Miron [593] that there is an "enhancement of acidity . . . relative to carboxylic acids"

when phenols are transferred from water to the

solvent pyridine. Its correctness could be tested

by comparing the behavior of 1,3-diphenylguanidine

and its 2-methyl-derivative.

Table 74. This table contains values of A// ob- tained by a different procedure. Forman and Hume

s

8

experiments (28°) with benzene as the solvent and

trichloroacetic acid as the titrant. A representative

titration curve obtained in the latter series of ex- periments is shown in figure 28. It is apparent that

in each series the magnitudes of A// parallel roughly

Table 74. Comparative heats of neutralization of different types of organic bases by

hydrogen bromide in acetonitrile {24°) and by trichloroacetic acid in benzene (28°) ^

Base AH, MeCN^ A//,, CeHe" A//2, CsHe"

n-Butylamine 10.60 -25.8 -29.3 -0.7

sec-Butylamine 10.56 -25.8 -29.6 -0.3

j-Butylamine 10.42 -25.5 -29.1 -0.7

Di-i-propylamine 11.05 -24.4 -29.4 -1.2

Di-n-butylamine 11.3 -26.2 -33.3 -1.9

Di-sec-butylamine 11.0 -24.1 -28.5 -1.3

Morpholine ^8.49 -25.5 -25.0 -2.3

Tri-?i-butylamine 9.93 -23.9 -23.0 -4.9

1 ,3-Diphenylguanidine 10.0 -20.7 -23.8 -1.9

Aniline 4.60 -14.9 -16.4

p-Anisidine 5.29 -17.7 -18.8 0.0 p-Chloroaniline 3.99 -12.3 -9.8

m-Nitroaniline 2.46 -8.2 -0.3

A'-Methylaniline 4.85 -13.2 -13.8

A',A'-Dimethylaniline 5.06 -14.9 -11.3

A',A'-Diethylaniline 6.56 - 15.6 -15.3

Pyridine 5.19 -14.4 -13.6

"Estimated from initial slopes of thermometric titration curves; AH is in kcal/mole and

is based on molar concentrations. See accompanying text.

" Most pKsH+ values are from ref. [271] or ref. [595]. From results of E. J. Forman and D. N. Hume, J. Phys. Chem. 63, 1949 (1959).

" From results of T. E. Mead, J. Phys. Chem. 66, 2149 (1962).

^ See H. B. Hetzer, R. G. Bates, and R. A. Robinson, J. Phys. Chem. 70, 2869 (1966).

References:

[271] R. A. Robinson and R. H. Stokes, Electrolyte Solutions, 2d ed. (Academic Press,

Inc., New York, N.Y. 1959).

[595] N. F. Hall, J. Am. Chem. Soc. 52 , 5117 (1930). [596] E. J. Forman and D. N. Hume, J. Phys. Chem. 63, 1949 (1959). [597] T. E. Mead, J. Phys. Chem. 66, 2149 (1962). [598] H. B. Hetzer, R. G. Bates, and R. A. Robinson, J. Phys. Chem. 70, 2869 (1966).

122
TCC)

Figure 28. Representative thermometric titration curve for

reaction of tertiary aliphatic amines with trichloroacetic acid

in benzene.

Reproduced from T. E. Mead. J. Phys. Chem. 66, 2149 (1962), with the permission

of the author and the publisher [597].

the strengths of the bases. The two acids repre- sented are a good deal stronger than the acids

represented in table 73, and this explains the

greater magnitudes of A//.

Because of the large number of bases represented

in both series of experiments. A// for the series in acetonitrile was plotted against A// values for the

reactions in benzene, yielding figure 29. It is evident

that there is a roughly linear correlation of the two

sets of values.

Table 75. This table contains values of logio/^w,

^H%, and AS^ for a large group of acid-base associations in benzene, taken from the Ph. D. thesis (1954)

of A. R. Anderson, and made available through the

kindness of Lord Wynne-Jones [594]. The last column contains log /Cbha values on the molar basis

(calculated by this writer). If expressed in molar

units of concentration, — A// would have been

[ smaller by 0.5 kcal or less. To illustrate, for triethylI amine with 2,4-dinitrophenol (25°) ^Hc is —13.95

I kcal/mole, and for diphenylguanidine with broil mophthalein magenta E, A//c is — 15.87. The value

1 1 1 I >i 9

- / dpg

/*—Slope = 1

V. —1 1 1 I 1 11 1 1 1 I i_ 0 10 20 30 -AH. CsHe

Figure 29. Heats of neutralization (— AH) of some bases with hydrogen bromide in acetonitrile (24°) plotted against the heats of neutralization of the same bases with trichloroacetic acid

in benzene (28°). Based on data reported by E. J. Forman and D. N. Hume, J. Phys. Chem. 63, 1949

(1959) and by T. E. Mead, J. Phys. Chem. 66, 2149 (1962). (Refs. 596 and 597, respectively.)

of — AS on the molar basis is greater by about 3 to 5 entropy units. A few values from other investigations are in- cluded in table 75. It is not feasible to discuss this interesting table in detail. It will suffice to note that the results support the previous general conclusion

that log Kqha, ^fi-: and AS tend to increase in magnitude with the strength of the proton donor

or acceptor. Some of the irregularities can be

attributed to steric hindrance.

b. Anisole and chlorobenzene.

Table 76. Bell and Bayles [354] compared the

strengths of a large group of aromatic amines as well as a few pyridine-type bases in terms of log K

for their association with bromophenol blue (first stage, colorless to yellow). Their results are shown

in table 76, together with accepted values of pK„

for the bases concerned. They concluded that the

indicator constants are about the same in the two

solvents. Furthermore, the indicator constants for amines in anisole were found to correlate well with

the catalytic constants in the decomposition of

nitramide in the same solvent. In previous work of

Bell and Trotman-Dickenson [623], where the dis- sociation constants in water had been the criteria of basic strengths, separate linear relationships

with the catalytic coefficients in anisole had been

found for primary, secondary, tertiary, and heterocycUc amines.

280-613 OL-68—

123
Table 75. Values of log Kbhai ^H, and AS {mole fraction units, 25°) for some 1:1 associations in benzene ^

Base Log XloKiv" ^m Log Kc'^

Reference acid"= p-Nitrophenol

ra-Propylamine 3.860 -8.214 -9.9 2.81 Diethylamine 3.812 -8.896 -12.4 2.76 Triethylamine 3.464 -8.799 - 13.7 2.42 "2.46

Reference acid = 2,4-Dinitrophenol

re-Propylamine 3.017 -11.220 -23.8 1.97 n-Butylamine 3.068 - 12.830 -29.0 2.02 re-Butylamine " 2.962 -11.86 -26.2 1.91 Diethylamine 4.279 - 14.760 -29.9 3.23 '3.09 Dipropylamine 4.064 - 14.920 -31.5 3.02 Dibutylamine 4.064 - 12.830 -24.4 3.02 Dibutylamine * 4.098 -13.13 -25.3 3.05 Triethylamine 4.535 - 14.450 -27.7 3.49 '3.47 ^3.50 Tripropylamine 3.765 -12.570 -24.9 2.72 Tributylamine 3.831 -11.440 -20.9 2.78 '2.79 Tributylamine 3.871 -11.97 -22.4 2.82

Reference acid = 2,4-Dinitro-o-cresol

Diethylamine 3.794 - 12.240 -23.7 2.74

Reference acid = Picric acid

Tribenzylamine 4.248 - 10.638 -16.3 3.20 "•'3.20

Reference acid = 2,4,6-Trinitro-m-cresol

Tribenzylamine 3.701 - 10.058 -16.8 2.65 '2.65

" Except as indicated otherwise, the values cited are results of A. R. Anderson and W. F. K. Wynne-Jones, as given in the Ph. D. thesis of A. R. Anderson, the University of Newcastle upon Tyne, 1954; they were communicated to M. M. Davis by Lord Wynne-Jones in June, 1964. As usual, A//° is in kcal/mole and AS° is in entropy units. Log Ku)K\ signifies log Khha in mole fraction units. Log Kc signifies log Khha in molar units. Unless otherwise

indicated the values in this column were computed by this author from the results given in the second column.

« See ref. [17].

« See ref. [441]. 'See ref. [591].

" See ref. [16].

" See ref. [23].

' See ref. [20]. References:

[16] M. M. Davis, J. Am. Chem. Soc. 84, 3623 (1962).

[17] M. M. Davis, unpublished results.

[20] M. M. Davis and E. A. McDonald, J. Res. NBS 42 , 595

(1949).

[23] A. A. Maryott, J. Res. NBS 41, 7 (1948). [441] J. W. Bayles and A. F. Taylor, J.-Chem. Soc, p. 417 (1961). [594] Results of A. R. Anderson and W. F. K. Wynne-Jones, as given in the Ph. D. thesis of A. R. Anderson, the University of Newcastle upon Tyne, 1954, and communicated by

Lord Wynne-Jones to M. M. Davis in June 1964.

[591] R. G. Pearson and D. C. Vogelsong, J. Am. Chem. Soc. 80,

1038 (1958). Table 76. Comparative values of log Kbha (molar units, 25°) for some aniline- and pyridine-type bases in anisole and chloro- benzene

Log Kbha

Base pKa" Anisole Chloro- benzene

(4.33) (5.62)

Aniline 4.58 1.0 1.0 o-ChloroanUine 2.62 0.1 m-Chloroaniline 3.52 0.0 p-Chloroaniline 4.0 0.0 o-Toluidine 4.39 1.0 1.26

p-Toluidine 5.07 1.75 1.8 A^-Methylaniline 4.85 1.74 1.8 A'-Methyl-o-toluidine 4.59 1.68 1.80 yV-Methyl-p-toluidine 5.33 2.51 2.66 A'-EthylaniUne 5.11 2.43 2.48

A^-Ethyl-p-toluidine 4.93 3.04 3.26

A'.A'-Dimethylaniline 5.06 2.87 2.80

A',A'-Dimethyl-o-toluidine 5.86 3.61 3.51 A'.A^-Dimethyl-p-toluidine 5.5 3.48 3.53

yV.A'-Diethylaniline 6.56 4.4 4.5

Pyridine 5.37 3.75 <^3.52 Quinoline 5.0 3.99 4.0 Isoquinoline 5.24 4.15 4.2

" Taken with minor changes from Table 6 of R. P. Bell and

J. W. Bayles, J. Chem. Soc, p. 1518 (1952). Values shown in parentheses are dielectric constants. Throughout, the reference

acid was bromophenol blue. To obtain these values the pKi, values as given in the paper

cited were subtracted from 14. In a similar measurement Popovych obtained the value 3.40

[85]. Some of his other results were cited earlier (see table 6).

References:

[85] O. Popovych, J. Phys. Chem. 66, 915 (1962). [354] R. P. BeU and J. W. Bayles, J. Chem. Soc, p. 1518 (1952).

Table 77. Additional values of log Kbha in chloro- benzene are given in this table. It is obvious that in chlorobenzene, as in benzene, the stabihty of the

complex with a given reference acid increases with

the strength of the base. We saw earlier (see table 6) that pyridine-type bases listed here associate more

extensively with bromophenol blue in chlorobenzene

124
than in benzene, but less extensively in toluene

[85]. A reasonable explanation is that the solvents,

as TT-electron donors, associate somewhat with the

indicator acid. Toluene is the best electron-donor of

the three solvents, and chlorobenzene is the most

inert. Reference should be made to an investigation, by

Bayles and Evans [634], of the "effect of chain

length and freedom of ring rotation" on the tendency

of various nitrogen bases to associate with 2,4-

dinitrophenol in chlorobenzene.

Table 77. Log Kbha (molar units, 25°) for some 1:1 associations

in chlorobenzene

Reference acid 1 t^fT K . '^BHA Refer- ences

2,4-Dinitrophenol n-Hexylamine 1.90 [5911 Dimethylamine 3.03 [591) Diethylamine 3.40 [5911 Di-n-butylamine 3.40 [5911 Trimethylamine 3.14 [591] Triethylamine 3.99 [5911 Tri-n-butylamine 3.58 [5911

Picric acid Pyridine "4.30 [5921

Bromophenol blue Pyridine 3.40 [851 3-Picoline 4.16 [851 4-Picoline 4.46 [851 2,4-Lutidine 5.64 [851

" Ezell and Gilkerson [599] obtained a closely similar value

(4.24) using o-dichlorobenzene as the solvent instead of chloro- benzene.

First color change (colorless to yellowr). References:

[85] O. Popovych, J. Phys. Chem. 66, 915 (1962).

[591] R. G. Pearson and D. C. Vogelsong, J. Am. Chem. Soc. 80,

1038 (1958).

[592] T. Jasiiiski, T. Misiak, and T. Skarzynska, Roczniki Chem.

39, 1549 (1965). I

[599] J. B. Ezell and W. R. Gilkerson, unpublished results cited

in ref. [449].

c. Miscellaneous aromatic solvents. Table 78. This table presents log Kn, A//jv, and

ASn (and also the corresponding log Kc, calculated

by this writer) for 2,4-dinitrophenol with mono-, di-, and tri-n-butylamine in six aromatic solvents and

also in n-butyl ether [441]. As in the work just

discussed, log K is greater in chlorobenzene than in benzene, but smaller in toluene. Interested readers

will wish to consult the original paper for more

details and discussion.

Added in proof: Recently Caldin and Crooks [660],

using a microwave temperature-jump technique,

determined at room temperature the

and "reverse" rate constants, kj- and

sponding to eq 44,

B + H-A«^B-H+...A-,

"forward"

kr, corre-

(44)

where H-A was 2,4-dinitrophenol, B was mono-,

di-, or tri-rt-butyiamine, and the solvent was cliloro- benzene. For each of the three systems the (experimental value of A;/ proved to be much low(?r than the value calculated on the assumption of a diffusion- controlled reaction. To explain these low values the

authors suggest that the forward reactions require an activation energy of the order of 1 kcal/mole

(molar units), the activation energy being connected

with the necessity of breaking the intramolecular

hydrogen bond in 2,4-dinitrophenol. [Compare

sees. 4.5.4b and 5.4.1a.]

Table 78. Effect of solvent on log; Kbha. '^H, and A.S for mono-,

di-, and tri-n-butylamine {HA= 2,4-dinitrophenof, mole frac- tion units, 25°) "

Solvent (e) Loj A//n Log Kr

Base = re-Butylamine

Benzene (2.27) 2.96 -11.86 — ZO.Z Toluene (2.38) 2.79 -11.81 -26.8 1.83 n-Butyl ether (3.06) 4.00 -15.39 -33.3 3.23 Anisole (4.33) 3.64 -12.11 -24.0 2.68 Bromobenzene (5.40) 2.98 - 13.23 -30.8 2.00 Fiuorobenzene (5.42) 3.07 -13.23 -30.3 2.05 Chlorobenzene (5.62) 3.21 -11.80 -24.9 2.22

Base = Di-n-butylamine

Benzene (2.27) 4.10 -13.13 -25.3 3.05 Toluene (2.38) 3.92 -13.04 -25.8 2.96 n-Butyl ether (3.06) 4.21 -15.22 -31.8 3.44 Anisole (4.33) 4.62 -13.75 -24.9 3.66 Bromobenzene (5.40) 4.46 -12.91 -22.9 3.48 Fiuorobenzene (5.42) 4.37 -12.72 -22.6 3.35 Chlorobenzene (5.62) 4.30 -11.89 -20.2 3.31

Base = Tri-fi-butylamine

Benzene (2.27) 3.87 - 11.97 -22.4 2.82 Toluene (2.38) 3.66 -11.71 -22.5 2.70

/!-Butyl ether (3.06) 3.35 -11.24 -22.4 2.58 Anisole (4.33) 4.42 -12.57 -21.9 3.46 Bromobenzene (5.40) 4.64 - 13.49 -24.0 3.66 Fiuorobenzene (5.42) 4.64 -13.81 -25.1 3.62 Chlorobenzene (5.62) 4.68 -14.16 -26.1 3.69

^ Results of J. W. Bayles and A. F. Taylor, J. Chem. Soc, p. 417

(1961). Log signifies log Krha in mole fraction units whereas

log Kc (derived by this author) is log ^bha molar units. Dielectric constants of solvents are from ref. [35]. References:

[441] J. W. Bayles and A. F. Taylor, J. Chem. Soc.. p. 417 (1961).

[35] A. A. Maryott and E. R. Smith, Table of dielectric constants of pure liquids, NBS Circ. 514, Aug. 10, 1951.

125
5.4.2. Aliphatic Hydrocarbons

Table 79. Heptane was used as solvent for determining the log A^BHA values in this table, both for its inertness and its transparency in the ultraviolet region. In one case (pyridine with phenol) a dielectrometric procedure was used [589]. Again there is a parallelism between strength of acid and magnitude

Oflog/^BHADenyer, Gilchrist, and co-workers have deter- mined log /^BHA for trimethylamine with four of the same phenols (phenol, p-cresol, and 1- and 2-naphthol) in cyclohexane [603]. AU values are of the same

order as those given in table 79 for their association

with triethylamine in heptane, and in some instances

the agreement is very close. Here, the method was

partition of one reactant between the vapor phase

Table 79. Log Kbha {molar units, 25°) for some base-phenol

associations in heptane ^

Reference

acid (pKa) "

Base

Py BuNHs BuzNH Et3N BujN

p-Cresol

(10.26) m-Cresol

(10.09) Phenol (10.00) 2-Naphthol

(9.57, 19°)"= 1-Naphthol

(9.39) '

2,4-Dinitro- phenol (4.11)

1.83

1.93 2.06

2.11

2.16

1.92

1.99 2.11

2.14

2.21

"2.56

1.74

1.82

1.93 "1.92

2.01 "2.01

2.07 "2.08

M.90

"2.00

^ With the exceptions noted log Kbha values are from M.

Bonnet and A. Julg, j. Chim. Phys. 59, 723 (1962).

" Most pKa values are for 25° and from ref. [271]. ^From R. J. Bishop and L. E. Sutton, J. Chem. Soc, p. 6100

(1964); A// = -6, AS = -11.

" See S. Nagakura and M. Gouterman, J. Chem. Phys. 26, 881

(1957). •'See F. Kieffer and P. Rumpf, Compt. Rend. 2 38, 360, 700

(1954). 'L. K. Creamer, A. Fischer, B. R. Mann, J. Packer, R. B.' Richards, and J. Vaughan, J. Org. Chem. 26, 3148 (1961).

^ See R. G. Pearson and D. C. Vogelsong, J. Am. Chem. Soc. 80, 1038 (1958). References:

[600] M. Bonnet and A. Julg, J. Chim. Phys. 59, 723 (1962). [271] R. A. Robinson and R. H. Stokes, Electrolyte Solutions, 2d ed. (Academic Press, Inc., New York, N.Y., 1959). [589] R. J. Bishop and L. E. Sutton, J. Chem. Soc, p. 6100 (1964). [353] S. Nagakura and M. Gouterman, J. Chem. Phys. 26, 881

(1957). 1601] F. Kieffer and P. Rumpf, Compt. Rend. 2 38, 360, 700

(1954). [602] L. K. Creamer, A. Fischer, B. R. Mann, J. Packer, R. B. Richards, and J. Vaughan, J. Org. Chem. 26, 3148 (1961). [591] R. G. Pearson and D. C. Vogelsong, J. Am. Chem. Soc. 80,

1038 (1958).

and the solution [603]. Results of Bellon [604] for triethylamine with these same phenols in cyclohexane, obtained by ultraviolet absorption spectroscopy, are also in good accord; Bellon's values are expressed in mole-fractional units. He pointed

out that the log ^bha values in cyclohexane are

related linearly to pKa values of the phenols in water.

5.4.3. Carbon Tetrachloride

As would be anticipated, carbon tetrachloride has

very often been the solvent for determinations of

log A^BHA by infrared absorption spectroscopy.

Table 80. This table illustrates how much smaller

values of log K^ha are when HA is an alcohol. The

enthalpy and entropy changes (see footnotes) are correspondingly small. It will be noted that replacement of hydrogen by phenyl on the primary carbon

atom, which increases the acidic character of an

alcohol, leads to slightly enhanced association of the

alcohol with pyridine. As would be expected, phenol

associates more extensively with pyridine than any

of the alcohols listed.

Table 80. Log Kbha (molar units, 20—27°) for association of

pyridine with some alcohols in carbon tetrachloride

Alcohol 20° 22° 25° 27° Refer- ences

Methanol

Ethanol

1-Propanol

Propanol

1- Butanol

2- Butanol

/-Butyl alcohol Pentanol Hexanol

Octanol

Benzyl

alcohol

4= !i= + Phenol <^

0.78

0.43

^0.47

"0.38

^^0.30

[135, 606]

[607, 606]

[609]

[608] [605, 608]

[605, 608]

[605, 608]

[608]

[608]

[608]

[608]

0.40

0.35 0.04

0.32 -0.03

-0.24 "0.16

0.32

0.42

0.45

0.72

1.72-1.78 1.62, 1.65 1.67

^A// = -3.88 (kcal/mole), AS = -10.8 (entropy units). ''A// = -3.66, A5 = -10.5.

<^A// = -4.3. A5 = -11.9. "A// = -3.98, AS = -12.6.

" See values of log K^ha listed in tables 82 and 84. See tables 83 and 84 for A// and A5. References:

[605] T. Zeegers-Huyskens. L. Lamberts, and P. Huyskens,

J. Chem. Phys. 59, 521 (1962). [135] T. Gramstad, Acta Chem. Scand. 16, 807 (1962). [606] E. D. Becker, Spectrochim. Acta 17, 436 (1961). [607] I. Prigogine, J. Chim. Phys. 45, 17 (1948). [608] A. M. Dierckx, P. Huyskens, and T. Zeegers-Huyskens, h. J. Chim. Phys. 62 , 336 (1965). [609] T. J. V. Findlay and A. D. Kidman, Australian J. Chem. 18, { 521 (1965).

126
Table 81. Once again, it is apparent that

log KbhA'. a//, and AS are greater for a phenol than

for an alcohol, and, moreover, these constants are

greatest for pentachlorophenol. Collidine (trimethylpyridine), being a stronger base than pyridine, as- sociates more extensively with all of the proton

donors listed.

Table 81. Log Kbha> AH, and AS (molar units, 20°) for association of pyridine and collidine with ROH or ArOH in carbon

tetrachloride ^

Base Log KsHA AH AS

ROH = Methanol

Pyridine CoUidine0.78

0.88 -3.2

-3.6

-7.4

-8.2

ROH = Ethanol

Pyridine 0.43

ArOH = 1-Naphthol (pKa = 9.39)

Pyridine

Collidine

1.92 2.30 -5.7

-6.7

-10.5

-12.2

ArOH = Pentachlorophenol {pKa = 4.82)

Pyridine CoUidine

2.05

2.35' -5.8

-7.6

-10.4

-15.1

^ With the exception noted values are from T. Gramstad,

Acta Chem. Scand. 16, 807 (1962); AH in kcal/mole, AS in

' entropy units.

I "See I. Prigogine, J. Chim. Phys. 45, 17 (1948).

tj References:

j [135] T. Gramstad, Acta Chem. Scand. 16, 807 (1962).

i [607] I. Prigogine, J. Chim. Phys. 45, 17 (1948). Table 82. Here, pyridine is the reference base

throughout. The increase in stability of the complex

with increase in proton-donor ability of the acid,

noted in previous tables, is again evident.

Table 83. In this table, phenol is the reference acid

throughout and the behavior of pyridine-type bases

of varying strength is investigated. The same general trends as before are observable.

Table 84. Comparisons of pyridine-type bases,

with phenol as the reference acid, are continued in

j this table.

li Table 85. This table is appropriately compared

with table 82. Aniline, being a much weaker base

than pyridine, associates much less extensively with

phenol and substituted phenols.

Table 82. Log Kbha (rnolar units, 20-27°) for association of

pyridine with phenol and related acids in carbon tetrachloride

Acid "

. 20"^

2|o,. 25" 27" "

p-Cresol (10.26) p-Methoxyphenol (10.21) m-Cresol (10.09) Phenol (10.00) p-Fluorophenol (9.91)

p-Chlorophenol (9.42) 1-Naphthol (9.39)

p-Bromophenol (9.36) p-Iodophenol (9.30) OT-Chlorophenol (9.13)

m-Bromophenol (9.03) m-Nitrophenol (8.38) Pentachlorophenol (4.82)

Picric acid (0.33)

1.62

1.62

1.63 1.77

1.53

1.51

1..55 1.67

1.84

2.05

" 1.78 ' 1.62

2.10

1.92 2.04

1.90 2.00

2.09

2.34

2.15 2.24

2.05

' 3.73

Numbers in parentheses are pK„ values (25°). With exceptions noted they are from A. I. Biggs and R. A. Robinson, J. Chem. Soc,

p. 388 (1961). "Results of T. Gramstad, Acta Chem. Scand. 16, 807 (1962). Results of J. Rubin, B. Z. Senkowski, and G. S. Panson,

J. Phys. Chem. 68, 1601 (1964). "Obtained by A.M. Dierckx, P. Huyskens, and T. Zeegers- Huyskens, J. Chim. Phys. 62, 336 (1965).

''H. Dunken and H. Fritzsche, Z. Chem. 1, 249 (1961), ob- tained 1.72. 'R. J. Bishop and L. E. Sutton, J. Chem. Soc, p. 6100 (1964), obtained 1.65 by dielectric polarization measurements. See L. K. Creamer, A. Fischer, B. R. Mann, J. Packer,

R. B. Richards, and J. Vaughan, J. Org. Chem. 26, 3148 (1961).

" See R. A. Robinson and R. G. Bates, J. Res. NBS 70A, 553

(1966). 'M. M. Davis and M. Paabo, J. Res. NBS 67A, 231 (1963).

' T. Jasinski, T. Misiak, and T. Skarzyriska, Roczniki Chem.

39, 1549 (1965).

References:

[535] A. I. Biggs and R. A. Robinson, J. Chem. Soc, p. 388 (1961). [135] T. Gramstad, Acta Chem. Scand. 16, 807 (1962). [610] J. Rubin, B. Z. Senkowski, and G. S. Panson, J. Phys.

Chem. 68, 1601 (1964). [608] A.-M. Dierckx, P. Huyskens, and T. Zeegers-Huyskens,

J. Chim. Phys. 62, 336 (1965). [611] H. Dunken and H. Fritzsche, Z. Chem. 1, 249 (1961). [589] R. J. Bishop and L. E. Sutton, J. Chem. Soc, p. 6100 (1964). [602] L. K. Creamer, A. Fischer, B. R. Mann, J. Packer, R. B. Richards, and J. Vaughan, J. Org. Chem. 26, 3148 (1961). [612] R. A. Robinson and R. G. Bates. J. Res. NBS 70A, 553

(1966). [362] M. M. Davis and M. Paabo, J. Res. NBS 67A, 231 (1963). [592] T. Jasinski, T. Misiak, and T. Skarzyriska, Roczniki Chem.

39, 1549 (1965).

Table 86. This table should be compared with

tables 83 and 84. The association of aniline-type

bases with phenol is far less than that of pyridinetype bases, in accord with expectation.

Table 87. This table is appropriately compared

with table 81; there does not appear to be much

127
Table 83. Log Kbha- ^H, and AS for association of some pyridine-type bases with phenol in carbon tetrachloride {molar units, 20°) "

Base" Log A:bha AW A5 Log A:bha

(25°)

Pyridine (5.17) 1.78 — 7.0 — 15.6 1.62

Derivatives of pyridine

3-Methyl (5.68) 1.86 -6.9 -15.0 1.78

2-f-Butyl (5.76) 2,6-Di-t-butyI 2-Methyl (5.97)

1.28 0.48

1.87 -6.9 -15.0 1.80 4-Methyl (6.02) 1.90 -6.6 -13.9 1.92 3,5-Dimethyl (6.19)* 2.03 2,6-Dimethyl (6.74)' 1.98 -6.9 -14.4 1.89 2,4-Dimethyl (6.82)'' 2.02 -6.8 -14.1

2,4,6-Trimethyl (7.59) ' 2.14 -7.5 -15.7

Related bases ^ Quinoline (4.85) 1.76 -7.2 -16.4

Isoquinoline (5.14) 1.79 -6.9 -15.4

Quinaldine (5.42)

1.90 -7.1 -15.6

Acridine (5.60) 1.83 -6.6 -14.1

"Values at 20° are from T. Gramstad, Acta Cham. Scand. 16, 807 (1962); AW is in kcal/mole and A5 is in entropy units. Log Kbha'^I 25° (see last column) is from

A. Halleux, Bull. Soc. Chim. Belg. 68, 381 (1959). Additional values at 20° for some

of the bases, and also values of log Kbh\^^ 27°, are given in table 84.

" pKa values in water (mainly thermodynamic values at 25°) are given in parentheses. See H. C. Brown and D. H. McDaniel, J. Am. Chem. Soc. 77, 3752 (1955). Unless stated otherwise pKa values for this group are results of H. C. Brown and

X. R. Mihm, J. Am. Chem. Soc. 77, 1723 (1955). The pKa value was calculated assuming that effects of alkyl substituents on the pKa of pyridine are additive. 'See H. C. Brown, S. Johnson, and H. Podall, J. Am. Chem. Soc. 76 , 5556 (1954). ^pKa values (at 20° with the exception noted) are from H. C. Brown, D. H. McDaniel, and O. Hafliger, Chapter 14 ("Dissociation Constants") in Determination of Organic Structures by Physical Methods (Academic Press, Inc., New York, N.Y.,

1955). (E. A. Braude and F. C. Nachod, eds.)

^ At 25°. References:

[135] T. Gramstad, Acta Chem. Scand. 16, 807 (1962). [613] A. Halleux, BuU. Soc. Chim. Belg. 68, 381 (1959). [614] H. C. Brown and D. H. McDaniel, J. Am. Chem. Soc. 77, 3752 (1955). [615] H. C. Brown and X. R. Mihm, J. Am. Chem. Soc. 77, 1723 (1955). [616] H. C. Brown, S. Johnson, and H. Podall, J. Am. Chem. Soc. 76, 5556 (1954). [617] H. C. Brown, D. H. McDaniel, and O. Hafliger, Chapter 14 ("Dissociation Constants") in Determination of Organic Structures by Physical Methods (Aca- demic Press, Inc., New York, N.Y., 1955). (E.A. Braude and F. C. Nachod, eds.) difference in the behavior of pyridine and triethyl- Readers should perhaps be reminded that table 43

amine as proton acceptors, when the proton donor contains log Kbha values for some amines with

is an alcohol or 1-naphthol. acetic acid in carbon tetrachloride and chloroform.

Table 88. Four tertiary aliphatic amines, an

alcohol, and three phenohc acids are represented 5.4.4. Chloroform

in this table. Tribenzylamine, a much wealcer base

than triethylamine, associates less readily with Table 89. As would be predicted, log Kq^a for proton donors of the types represented. The excep- pyridine with phenol in chloroform is less than in tionally low value oi log A;bha for tribenzylamine with carbon tetrachloride (compare table 82). Table 89

pentachlorophenol is noteworthy — presumably this may be compared with table 75, which contains

is to be attributed to steric hindrance. Values of log A^bha values for some of the same amine-phenol

A// and AS for associations involving tribenzylamine combinations in benzene. No clear-cut conclusions

are all exceptionally low. are possible.

128
Table 84. Log Kbha. ^H, and AS for association of some

pyridine-type bases with phenol in carbon tetrachloride (molar

units, 20°) "

Substituent in Pyridine * Log Kbha A// AS Log Kbha (27°)

2-Cl (0.72) 2-Br (0.90)'' 3-CN (1.36) <" 4- CN (1.90) " 3-Br (2.84)

3-Cl (2.84)

^• H (5.17) <^ 3-CH., (5.68)

< 2-CH3 (5.97)

<• 4-t-Bu (5.99)

4-CH3 (6.02)

4-C2h.5 (6.02)

0.95

0.90

1.15 i.OB

1.18

-4.9

-3.2

-4.5

-11.5

-6.0

-10.0 1.17

1.17

1.67 1.79

1.80

1.77

1.81 -6.5

-6.7

-14.0

-14.5

1.92

1.91 1.89

-7.1'

-6.4

-5.9

-12.0

-13.0

-11.5

1.92

* Some values of log ^bha at 27° are also given (see last column);

these are from A.-M. Dierckx, P. Huyskens, and T. Zeegers- Huyskens, J. Chim. Phys. 62 , 336 (1965). The other results are from J. Rubin and G. S. Panson, J. Phys. Chem. 69 , 3089

(1965). Numbers in parentheses are pKa values (at 25°. where

specified). Determined by H. C. Brown and D. H. McDaniel, J. Am.

Chem. Sec. 77, 3752 (1955). "See S. F. Mason, J. Chem. Soc, p. 1247 (1959); temperature not specified. ''See H. C. Brown and X. R. Mihm, J. Am. Chem. Soc. 77,

1723 (1955).

References:

[624] J. Rubin and G. S. Panson, J. Phys. Chem. 69 , 3089

(1965).

[608] A.-M. Dierckx, P. Huyskens, and T. Zeegers-Huyskens,

J. Chim. Phys. 62, 336 (1965).

[614] H. C. Brown and D. H. McDaniel, J. Am. Chem. Soc. 77,

3752 (1955).

[618] S. F. Mason, J. Chem. Soc, p. 1247 (1959).

[615] H. C. Brown and X. R. Mihm, J. Chem. Soc. 77, 1723

(1955).

5.4.5. Miscellaneous Aprotic Solvents Table 90. This is a collection of values, mainly

for pyridine, tertiary amines, and their A'-oxides,

with a phenol or iodine as the reference acid.

It will be noted that the amine oxides seem to form more stable adducts than the corresponding ter- tiary amines. Figures 14 and 15 indicated that in amine oxide adducts the proton of the donor is less displaced towards the proton acceptor.

Before leaving the topic of log A!^bha values it seems desirable to call attention to some derivatives

of phenol that might prove more useful for aprotic

media than phenols that form stable chelate rings.

A list of phenols covering the pKa range from about

0.3 to 4.8 is given in table 91. It would be interest- ing to ascertain log /^bha values for these phenols

Table 85. Log Kdha (molar units, 27°) for association of aniline with phenol and some of its monosubsliluted derivatives

in carbon tetrachloride "

AClu \pi\a )

1 f\if iC . '-'"B "BHA

p-Cresol (10.26) 0.44

/^-ItI li l\/If»tVi/ivvnh#»nnl cllKl MO 91^ Ay [Jlltrliui \ i\J.^ I ) fl d9

ffi-KjrKzsoi \ ixj.yjy) PK^nril I Ilt~IMH MO(H^.V/l/^ 00^ 0 ^0

luorophcnol (9.91)

p-Chlorophenol (9.42) 0.79 p-Bromophenol (9.36) 0.75 p-Iodophenol (9.30) 0.68 m-Chlorophenol (9.13) 0.75 m-Bromophenol (9.03) 0.80 m-Nitrophenol (5.38) 1.13

" Results of A.-M. Dierckx, P. Huyskens, and T. Zeegers- Huyskens, J. Chim. Phys. 62, 336 (1965). Obtained by A. O. Biggs and R. A. Robinson, J. Chem. Soc,

p. 388 (1%1). References:

[608] A.-M. Dierckx, P. Huyskens, and T. Zeegers-Huyskens, J. Chim. Phys. 62, 336 (1965). [535] A. I. Biggs and R. A. Robinson, J. Chem. Soc, p. 388 (1961).

Table 86. Log Kbha (molar units, 27°) for association of some

aromatic bases with phenol in carbon tetrachloride

Aromatic base Log Kbha

o-Bromoaniline (2.53) 0.32 o-Chloroaniline (2.65) 0.33

o-Fluoroaniline (3.20) 0.29 m-Fluoroaniline (3.59) 0.37 p-Bromoaniline (3.86) 0.51 p-Chloroaniline (3.98) 0.51

o-Toluidine (4.45) 0.56 Aniline (4.60) "^0.60 p-Fluoroaniline (4.65) 0.58 m-Toluidine (4.73) 0.66 p-Toluidine (5.08) 0.75 p-Anisidine (5.34) 0.87

A',A'-DiethylaniHne (6.56") •^0.42 (20°)

* Results of A.-M. Dierckx, P. Huyskens, and T. Zeegers- Huyskens, J. Chim. Phys. 62, 336 (1965). Values in parentheses are thermodynamic pK„ values (25°). All but the last were determined by A. I. Biggs and R. A. Robinson, J. Chem. Soc, p. 388 (1961). The same value was obtained at 25° by A. Halleux, Bull. Soc. Chim. Belg. 68, 381 (1959).

" See N. F. Hall and M. R. Sprinkle, J. Am. Chem. Soc. 54,

3469 (1932). ''See T. Gramstad, Acta Chem. Scand. 16 , 807 (1962); combined results for 20° and 50° yielded A// = - 1.3 and AS = -2.7.

References:

[608] A.-M. Dierckx, P. Huyskens, and T. Zeegers-Huyskens, J. Chim. Phys. 62, 336 (1965). [535] A. I. Biggs and R. A. Robinson, J. Chem. Soc, p. 388 (1961). [613] A. Halleux, Bull. Soc. Chim. Belg. 68 , 381 (1959). [619] N. F. Hall and M. R. Sprinkle, J. Am. Chem. Soc. 54,

3469 (1932). [135] T. Gramstad, Acta Chem. Scand. 16, 807 (1962).

129
Table 87. Log Kbha. '^H, and AS (molar units, 20-25°) for as- sociation of triethylamine with some alcohols and phenols

in carbon tetrachloride

Proton donor

{pKa)

Log Xbha" AH A5 Refer- ences

Methanol 0.81 (20) -3.8 -9.2 [135] Ethanol 0.65 (22) "-4.83 [605, 621] 1-Butanol 0.50 (22) "-4.86 [620, 621] 2-Butanol 0.23 (22) "-4.98 [605, 621] Phenol (10.00) 1.87 (20) -8.35 -20.0 [622] Phenol 1.96 (20) -7.8 -17.7 [135] Phenol 1.95 (25) -9.2 -21.7 [137] 1-Naphthol (9.39) 2.04 (20) -6.6 -13.1 [135]

^ Temperature of measurement (°C) is given in parentheses.

" Determined calorimetrically at 20°.

References:

[620] T. Zeegers-Huyskens, Bull. Soc. Chim. Beige 69, 282

(1960). [621] L. Lamberts and T. Zeegers-Huyskens, J. Chim. Phys. 60,

435 (1963). [622] H. Dunken and H. Fritzsche, Z. Chem. 1, 127 (1961). [135] T. Cramstad, Acta Chem. Scand. 16, 807 (1962). [137] M. D. Joesten and R. S. Drago, J. Am. Chem. Soc. 84,

3817 (1962). [605] T. Zeegers-Huyskens, L. Lamberts, and P. Huyskens, J. Chim. Phys. 59, 521 (1962).

in solvents like benzene and chlorobenzene, using

the same reference base or bases, and see how

they relate to the aqueous pK„ values.

Several investigators, including Weller [568],

have determined aqueous pKa values for phenolic

acids such as 1- and 2-naphthol in the excited state, finding that on excitation the acid strength is much

enhanced. Weller has also determined the equiHbrium constants (20°) for hydrogen bonding of

3-hydroxypyrene with pyridine and with 2-chloro- pyridine in some aprotic solvents, making use of

both absorption and fluorescence data, and thus

deriving both K° and K*, where K° denotes the

equilibrium constant in the ground state, and K*

is the corresponding constant for hydrogen bonding

in the excited state. Some of his results, expressed

as the logarithms of K° and K*, are given in table 92. These results support the following conclusions:

(1) In aprotic solvents, as in water, a phenol is a

better proton donor in the excited state than when

in the ground state.

(2) Pyridine, which is a much stronger base than

2-chloropyridine, associates more extensively with

3-hydroxypyrene in the excited state as well as in the ground state.

(3) The extent of hydrogen bonding is not related

to the dielectric constant of the solvent, but rather

to its electron-donor properties, since hydrogen

Table 88. Log Kbha, ^H, and AS {molar units, 20°) for asso- ciation ofsome tertiary aliphatic amines with ROH or ArOH

in carbon tetrachloride ^

Tertiary amine Log KbhA AH A5

ROH = Methanol

Triethyl (10.87) 0.81 -3.8 -9.2

Tripropyl 0.53 -2.7 -6.9

Tributyl 0.61 -3.5 -8.8

Tribenzyl (5.60) 0.53 -1.5 -2.5

ArOH = Phenol (10.00)

Triethyl (10.87) 1.96 -7.8 -17.7

Tripropyl 1.35 -5.9 -14.1

Tributyl 1.47 -6.9 -16.9

Tribenzyl (5.60) 0.43 -1.6 -3.4

ArOH = 1-Naphthol (9.39)

Triethyl (10.87) 2.04 -6.6 -13.1

Tripropyl 1.49 -6.1 -14.2

Tributyl 1.65 -5.9 -12.6

Tribenzyl (5.60) 0.28 -1.3 -3.4

ArOH == Pentachlorophenol (4.82)

Tribenzyl (5.60) 0.11 -3. -10.0

"Results of T. Gramstad, Acta Chem. Scand. 16, 807 (1962). Numbers in parentheses after the name of a phenol or amine

are aqueous pK,, values. Reference:

[135] T. Gramstad, Acta Chem. Scand. 16, 807 (1962).

Table 89. Values of log Kbha (molar units, 25°) for some

acid-base associations in chloroform ^

Reference acid Base Log/^BHA

2,4-Dinitrophenol MeNH2 L54

n-BuNH2 L65

n-HxNH2 1.72 MejNH 3.09 Et.NH 3.24 n-Bu2NH 3.41 MeaN 3.43 EtjN 4.20 n-BusN 3.96 Phenol Py " 1.23

"All values of log Kbha except the last one are from R. G. | Pearson and D. C. Vogelsong, J. Am. Chem. Soc. 80, 1038 I (1958). , " A result of J. Rubin and G. S. Panson, J. Phys. Chem. 69,

3089 (1965); t = 20°.

References:

[591] R. G. Pearson and D. C. Vogelsong, J. Am. Chem. Soc. 80,

1038 (1958). [624] J. Rubin and G. S. Panson, J. Phys. Chem. 69, 3089

(1965).

130
Table 90. Log Kbha. AH, and AS (molar units) for several tertiary amines and

tertiary amine N-oxides in some aprotic solvents "

Base Ref. Acid Solvent t "C Log ^BHA

Pvry HPi 0-v-(jn4V^l2 „ p 14 PI A OA PyO h CCI4 23 1.89 Me:,NO I. CH2CI2 22 3.75 MesNO C^HsOH CH2CI2 20.5 3.57 MeaNO I-CoHtOH CH2CI2 20.7 3.80 rH.n. 9n 7 o.oz Et3N I-CoHtOH CH2C12 20.7 1.74 EtjN 2-C,oH70H CH2C12 20.7 1.78 BuaN HPi 0-C6H4C12 25 7.04 BujN HPi CH2CICH2CI 25 7.15

BuaNO HPi 0-C«H4Cl2 25 7.89 (PhCH2)3N BPM-E " CeHe 25 -0.89 (PhCH2)3NO BPM-E

CsHc 25 5.68 (PhCH2)3NO I2 CH2CI2 20 3.54

A// A.S Refer- ences

-5.85

-10.0 -7.9

-11.0

-16.9 -10.6

-10.5 -19.6

[599]

[625] [6251

[6261

[626]

[626]

[626]

[626]

[449] [448]

[449]

[31]

[31] [625]

* A// in kcal/mole, AS in entropy units. ''BPM-E symbolizes bromophthalein magenta E (3',5',3",5"-tetrabromophenol- phthalein ethyl ester).

References:

[31] M. M. Davis and H. B. Hetzer, J. Am. Chem. Soc. 76, 4247 (1954). [448] E. K. Ralph, III, and W. R. Gilkerson, J. Am. Chem. Soc. 86, 4783 (1964). [449] W. R. Gilkerson and E. K. Ralph, III, J. Am. Chem. Soc. 87, 175 (1965). [599] J. B. Ezell and W. R. Gilkerson, unpublished results cited in ref. [449]. [625] T. Kubota, J. Am. Chem. Soc. 87 , 458 (1965). [626] T. Kubota, J. Am. Chem. Soc. 88, 211 (1966).

bonding is most extensive in the most inert solvent

(methylcyclohexane) and is least in the best elec- tron-donor solvent (benzene).

In section 4.6.2 reference was made to a recent symposium in which Ubbelohde expressed dis- agreement with the opinion that weak hydrogen

bonds may be distinguished from strong hydrogen

bonds [554], asserting, instead [555], that "a whole

spectrum of hydrogen bonds appears to exist." The

tabulated values of log Kbha and A// in aprotic

solvents, just presented, provide much support for Ubbelohde's view.

131
Table 91. Values o/ pKg (water, 25°) for some halogen- and

nitro-derivatives of phenol

Acid pKa

Picric acid

Xrinitro-^-cresol "0.81

Trinitro-3.5-xylenol 1.38 2-Chloro-4,6-dinitrophenol "2.10

4-Chloro-2,6-dinitrophenol ^2.97

4-Chloro-2,6-dinitro-m-cresol "3.24

2,6-Diiodo-4-nitrophenol ''3.32 2,6-Dibromo-4-nitrophenol "3.39

2,6-Dichloro-4-nitrophenol "3.55

2,6-Dinitrophenol '3.71

2,4-Dinitro-6-phenylphenol "3.85

2,4-Dinitro-3-methyl-6-isopropylphenol "3.96

2.4-Dinitrophenol 1^4.11 Bromophthalein magenta E ^ ~ 4.15

2,6-Dinitro-p-cresol " 4.23

4,6-Dinitro-o-cresol "4.47

4.6-Dinitro-2-i5opropylphenol "4.54

2,4-Dibromo-6-nitrophenol "4.71 Pentachlorophenol '4.82

^ M. M. Davis and M. Paabo, ref. [362]. "M. M. Davis and M. Paabo, ref. [627]. M. M. Davis. M. Paabo, and R. A. Robinson, ref. [628].

" R. A. Robinson, unpublished result. E. E. Sager, R. A. Robinson, and R. G. Bates, ref. [636]. 'G. Kortiim and H. Wilski, Z. Physik. Chem. 2, 256 (1954);

J. F. J. Dippy, S. R. C. Hughes, and J. W. Laxton, J. Chem.

Soc, p. 2995 (1956). R. G. Bates and G. Schwarzenbach, ref. [629]. M. M. Davis and H. B. Hetzer, ref. [32], footnote 29.

' R. A. Robinson and R. G. Bates, ref. [612].

References:

[32] M. M. Davis and H. B. Hetzer. J. Res. NBS 60, 569 (1958). [362] M. M. Davis and M. Paabo, J. Res. NBS 67A, 241 (1963). [533] R. A. Robinson, unpublished work.

[612] R. A. Robinson and R. G. Bates, J. Res. NBS 70A, 553

(1966). [627] M. M. Davis and M. Paabo, J. Res. NBS 64A, 533 (1960). [628] M. M. Davis. M. Paabo, and R. A. Robinson. J. Res. NBS

64A, .531 (1960). [629] R. G. Bates and G. Schwarzenbach, Helv. Chim. Acta 37,

1069 (1954). [630] G. KortQm and H. Wilski. Z. Physik. Chem. 2, 256 (1954). [631] J. F. J. Dippy, S. R. C. Hughes, and J. W. Laxton. J. Chem.

Soc.. p. 2995 (1956). [636] E. E. Sager. R. A. Robinson, and R. G. Bates, J. Res. NBS 69A, 305 (1964).

Table 92. Log Kbha (molar units, 20°) for hydrogen bonding of 3-hydroxypyrene to pyridine and to 2-chloropyridine in the ground and excited states "

Solvent (e, 20° '^)

Proton acceptor ''

Pyridine (5.17) 2-Chloropyridine

(0.72)

Log K° Log K* Log K° Log K*

Methylcyclohexane (2.02) 2.34 2.99 1.27 1.78 Benzene (2.28) 1.76 2.43 O.8., 1.26 o-Chlorotoluene (4.45) 1.89 2.63 0.9 1.40

" See A. Weller. in Progress in Reaction JCinetics, Vol. 1 (G. Porter, Ed.), p. 189 IT. (Pergamon Press, New York, 1961). pKi, at 25° (see table 84) is given in parentheses.

' See ref. [35]

References: •

[35] A. A. Maryott and E. R. Smith, Table of dielectric con- - stants of pure liquids, NBS Circ. 514. Aug. 10. 1951.

'

[568] A. Weller, in Progress in Reaction Kinetics, Vol. 1 (G.

| Porter, Ed.), p. 189 ff. (Pergamon Press, New York, 1961). K

!

i

t

E

I

;

!

i

I'

132
6. Acid-Base Titrations in Completely Aprotic Solvents

6.1. Introduction

The assigning of an exaggerated importance to water is nowhere better illustrated than in the

history of acid-base titrations. The beUef still persists that water must be present, at least in trace amounts, if acid-base re'actions are to occur. However, organic chemists have recognized that water is not needed, and as mentioned earlier,

Vorlander discovered more than 60 years ago that anihne is titratable with hydrogen chloride in benzene, using methyl yellow as the indicator [542]. Moreover, about a decade later, Folin and Flanders [637] demonstrated that carboxylic acids can

be titrated more satisfactorily in mixtures of an

alcohol with benzene, carbon tetrachloride, chloro- form, or toluene than in alcohol-water mixtures; in their experiments the titrant was sodium ethylate

or amylate in the corresponding alcohol, and phenoiphthalein was considered the best of the available

indicator dyes.

Fohn and Flanders' paper was generally over- looked for about 35 years, but eventually alcohol- hydrocarbon mixtures became popular for acidbase titrations — especially mixtures of benzene or toluene with isopropyl alcohol (a medium adopted

for ASTM standard methods) [638], ethylene glycol

[639], or methanol [50]. A small amount of water is considered necessary when the glass-calomel

electrode system is used for locating an end-point,

but not in conductance or indicator titrations. As

time has passed, analysts have tended to reduce

the alcohol content to the point where it is a very

minor component of the medium, introduced only

as part of the titrant. (For examples, see refs. [476, 484, 485]; several examples are also given

in ref. [640]). In titrations requiring a strong base,

the presence of a Uttle alcohol and/or water has

been unavoidable, since no satisfactory substitute

for quaternary ammonium hydroxides (or alkoxides)

has yet been found. However, procedures have

been devised for reducing their content to a very

small percentage of the total solvent. On the whole, there has been a strong inclination to use mixed solvents in nonaqueous titrimetry.

Not infrequently, several solvents are present, and

it has been comparatively rare for a single non- aqueous solvent to be used for titrations — especially an aprotic solvent. However, titrations in completely

aprotic solvents are entirely feasible when a proper

combination of reactants is selected. Our remaining discussion will consist of a brief survey of such

titrations, in which various indicator dyes or physical measurements were used for end-point location.

I We shall not dwell on modern improvements in inI strumentation, since these are widely recognized

I and taken advantage of. Emphasis will be placed,

instead, on the need for further modernization of

chemical aids, such as reference acids and bases.

Comparatively little work has been done in this

'I

I;

area, and the aid of organic chemists could well be enlisted — in particular, in developing improved

reference bases. Further developmental work with

indicator dyes is also needed, since titrations with

indicators are still preferred by many because of

their speed and simplicity. However, for the most

fruitful use of indicator dyes, as much effort should

be expended in gaining insight into the structural changes responsible for color changes as is ex- pended in learning about the construction and

proper use of instrumental aids.

6.2. Titrations With Indicator Dyes in

Completely Aprotic Solvents

Possibilities in this category are indicated in table 93. Bromophthalein magenta E, tamarack

green base, and Victoria blue B anhydro-base

are examples of indicators intended especially

for aprotic solvents. The latter two are bases.

p-Naphtholbenzein, an acid weaker than bromophthalein magenta E, is sometimes useful in benzene. Diphenyl phosphate was found to be a very

satisfactory reference acid in the examples given

in the table. It can readily be converted from its dihydrate to the anhydrous state, does not absorb

moisture readily, is easily soluble in solvents like benzene, and is colorless in the visible region. In water it appears to be as strong as hydrochloric

acid [313]. Some suggestions about reference com- pounds will be given in section 6.4.

6.3. Instrumental Procedures for Detecting

End-points in Completely Aprotic Solvents

6.3.1. Conductance Titrations Examples from Maryott's work [341] were shown

in figures 19 to 21. It will be recalled that he ti- trated aliphatic amines in benzene with such acids

as picric acid, trichloroacetic acid, and (fZ-camphorsulfonic acid. The work was suggested by earlier exploratory titrations of LaMer and Downes [340],

who did not continue their investigations after ob- serving anomalies in behavior.

6.3.2. Dielectrometric Titrations Gur'yanova and Beskina [347] measured changes

in the dielectric permeability of a solution of benzoic

acid in benzene on the addition of n-butylamine,

various secondary amines (i-PrjNH, BU2NH,

Et2NH, piperidine), or a tertiary amine (EtsN,

AmaN). Their objectives were to determine stoi- chiometry and relative strengths, rather than to develop titrimetric procedures.

Bryant and Wardrop's conductimetric studies of

acid-base interactions in acetone and acetonitrile

[474] included dielectrometric titrations of tri- chloroacetic acid (about 0.001 M) with triethylamine

(about 0.2 M) in benzene and also in dioxane. The
Table 93. Some examples of acid-base titrations in completely aprotic solvents, using indicator dyes

C 1 oolvcnts Acids T A- A Indicfltor dyes R f i\eie Fences

Benzene; HCl; ArSOsH " Amines; drugs; " [540-543, chloroform ^ alkaloids 641-643] Chloroform CH3COOH; HPi 'C2G 'p-m [360] Benzene CCl3C00H;HPi ^ 0-T2G " BPM-E [14] Benzene ' RSO3H ^C2G; ^Ph2G BPM-E [14] Benzene " Ph2HP04 ^o-T2G; iPhsG "BPM-E; ' "TGB [313]

^ Also carbon tetrachloride, dichloromethane, tetrachloroethane, ligroin. ''Ar = C6H5 or P-CH3C6H4.

A',A^-Dimethyl-p-azoaniline (p-dimethylaminoazobenzene); usually called methyl

yellow (but occasionally, "dimethyl yellow" or "butter yellow"). This group of papers was reviewed briefly in ref. [14].

* 1,3-Dicyclohexylguanidine; see ref. [14]. 'p-Naphtholbenzein.

^ 1,3-Di-o-tolylguanidine. Bromophthalein magenta E (3',5',3",5"-tetrabromophenolpthalein ethyl ester).

' In this table RSO3H signifies d-lO-camphorsulfonic acid.

' 1,3-Diphenylguanidine.

Diphenyl phosphate (anhydrous).

' Tamarack green base (anhydro-A',A''-diphenyl-p,p'-diamino-o"-chlorotriphenylcarbinol). Victoria blue B anhydro-base, prepared from the commercial salt, which has

Color Index No. (1st ed.) 729, was another satisfactory indicator.

References;

[14] M. M. Davis and P. J. Schuhmann, J. Res. NBS 39, 221 (1947). [313] M. M. Davis and H. B. Hetzer, J. Res. NBS 54, 309 (1955). [360] M. M. Davis and P. J. Schuhmann, unpubhshed work.

[540] R. Dietzel and W. Paul, Arch. Pharm. 273, 507 (1935); 2 76, 408 (1938). [541] E. M. Trautner and F. H. Shaw, Austrahan Chem. Inst. J. & Proc. 12, 405 (1945). [542] D. Vorliinder, Chem. Ber. 36, 1485 (1903). [543] D. Vorliinder, J. Fischer, and F. Wildner, Chem. Ber. 66, 1789 (1933). [641] D. Vorlander, Chem. Ber. 67, 145 (1934). [642] E. M. Trautner and C. E. Neufeld, Austrahan Chem. Inst. J. & Proc. 13, 70

(1946). [643] G. Ya. Khalt, Farmatsiya 8, 26 (1945); see Chem. Abstracts 41, 2206 (1947).

dielectric constant increased gradually until the

equivalence point, then became essentially constant

(see their figure 5). A "simulated" dielectrometric titration curve for picric acid (~ 0.002 M) with triethylamine in ben- zene [644] is shown in figure 30. The points were

0 0 5 1,0 :.5 2.0 Moles of Triethylamine per Mole of Ptcnc Acid

Figure 30. Curve for the simulated dielectrometric titration of

picric acid (~ 0.002 M) with triethylamine in benzene.

Based on unpublished results of R. Megargle, D. Rosenthal, and G. L. Jones, Jr., Clarkson College of Technology, 1966.

obtained by measuring the dielectric constant of a

series of solutions containing a fixed amount of the

acid but varying amounts of the base. In the ex- ample shown, the calculated and experimental

end-points agreed within 1.7 percent. In this pro- cedure the volume of solvent did not change during

the titration. The authors have titrated other acid- base systems in benzene with comparable results. Ishidate and co-workers (for example, see ref. [648])

have recently explored analytical dielectrometric

titrations in dioxane.

6.3.3. Photometric Titrations What are in effect photometric titrations have

been performed frequently in spectrophotometric

investigations intended mainly to determine acid- base stoichiometry and relative strengths in solvents

like benzene (for examples, see refs. [5, 20, 21, 14] and [482]).

Davis and IMcDonald [361] titrated picric acid

with such bases as diphenyl- and di-o-tolylguanidine

spectrophotometrically, obtaining curves very

134

I
similar to figure 30. Here again, the main objectives were to determine stoichiometry and relative strengths. Such work is a necessary stage, of course,

in the development of titration methods.

6.3.4. Thermometric Titrations Here we may refer again to the titrations of bases

with hydrogen bromide in acetonitrile, performed

by Forman and Hume [596], and to the titrations of

similar bases with trichloroacetic acid in benzene

by Mead [597]. Mead's principal objective was to determine values of A// and correlate them, if possible, with relative strength of the bases. (See

fig. 28.)

6.3.5. Titrations by the DVP Method [352]

The reader is referred to figure 12 and related

discussion. Here, again, is a method used hitherto

for determining stoichiometry and relative strengths, not as an analytical procedure.

6.3.6. Cryoscopic Titrations Gur'yanova and Beskina [347] performed cryo- scopic titrations of benzoic acid in benzene with

amines like AmsN, EtsN, and piperidine, as an

adjunct to dielectrometric titrations, in studying the

association of benzoic acid with amines (see sec.

4.3.2). More recently, Bruckenstein and Vanderborgh

[645] have titrated bases in benzene with trifluoro- acetic and trichloroacetic acids, using an experimental apparatus for recording continuously the

change in the freezing point depression during a

titration. They obtained curves which can be in- terpreted analogously to differential vapor pressure

data. The cryoscopic method is useful when the

differential vapor pressure method is not appUcable (for example, when volatile solutes are

involved).

6.3.7. Other Possible Instrumental Procedures

Other physical methods that have been found

to indicate changes in composition resulting from

acid-base interactions in aprotic solvents are:

a. Refractive index. Giles, McKay, and Good [647]

determined changes in the refractive index of carbon

tetrachloride solutions containing such mixtures

as benzene-phenol, chlorobenzene-phenol, and

hexachlorobenzene-phenol. V-shaped curves were

obtained for chlorobenzene-phenol on plotting the

change in the refractometer reading against the

composition (expressed as mole fractions).

b. Density. Changes in density on the addition

of tripentylamine to benzoic acid in benzene appear

to be very abrupt near the end-point [347]. Changes

in density upon varying the proportion of piperidine

to propionic acid in benzene have also been measured [646].

c. Optical rotation. Bodforss' studies of the

change in optical rotation on adding an acid to

brucine in benzene or chloroform were mentioned

a Httle earher [574]. Although these physical methods may not all be suitable for analytical titrations, this list serves

to remind us that there are doubtless many untried

types of measurements that could be applied to advantage in determining the composition of acid- base complexes, relative strengths, and possibly

also in analytical titrimetry in aprotic media.

6.4. Reference Acids and Bases for

Aprotic Organic Solvents

6.4.1. Acids

We shall comment briefly on five chemical

types that have been used hitherto.

a. Hydrogen chloride. This was used in the

pioneering aminometric titrations of Vorlander

and associates [542, 543, 641]. Its proton-donor

ability makes it suitable for many titrations, but

standard solutions are troublesome to prepare

and have to be restandardized frequently. Also,

the gas escaping from solutions is likely to con- taminate other materials and cause corrosion of

metal objects.

h. p-Toluenesulfonic acid. This is a useful solid reference acid, so far supplied commercially

only as the monohydrate. Trautner and Shaw [541]

described a method for dehydrating it. Maryott

used c?MO-camphorsulfonic acid for conductance

titrations of amines in dioxane [341]. The optically

active isomer, c?-10-camphorsulfonic acid, is some- what more easily soluble in benzene, and was used

for titrations of bases in this solvent in the presence

of indicator dyes [14].

c. Ficric acid. Although this acid is easily obtained and purified, is conveniently handled, and

is a good proton donor, the strong yellow color of

its salts is sometimes objectionable. In section 5

a list of possible replacements (for some pur- poses) was suggested (see table 91).

d. Trichloroacetic acid is a good proton donor

and readily soluble. It is somewhat volatile and

therefore difficult to dry completely.

e. Diphenyl phosphate. This acid has been in- vestigated [313] as a reference acid in benzene

(as well as in benzene-methanol mixtures). See

comments on table 93.

6.4.2. Bases The discussion will be confined to bases of at least moderate strength.

a. Aliphatic amines. We shall only discuss ter- tiary amines, which are preferable to primary and

secondary amines for reasons indicated earlier. Triethylamine is still used extensively, although it seems to be giving place to amines of higher

molecular weight — for example, A'.A'-dimethylbenzylamine. Tribenzylamine has admirable

physical properties, but is not a very strong base.

135
It is rather surprising that the cyclic tertiary

amine, quinuclidine (XCI) has not been made

H

I

N

(XCI)

available commercially. Brown and Eldred [649]

described its preparation "in good yield starting

with the commercially available 4-pyridineethanol,"

and later, S. Leonard and Elkin [650] described

another procedure for obtaining it from the same

starting material. This interesting base is a very

close relative of triethylamine, but the ethyl groups

are "tied back," thus leaving the nitrogen atom more exposed, and thereby more reactive. As an

illustration, Brown and Sujishi [635] found quinuclidine to form the most stable known complex

with the Lewis acid, trimethylboron, whereas

Brown and Taylor [651] found the association of

triethylamine with trimethylboron to be too small

to measure. The model of quinucHdine shows it to have a globular structure. Although easily subhmed, it is a sohd at ordinary temperatures (m.p.

156° in a sealed tube [650]).

There may also be potentially useful reference

bases among derivatives of pyrrolizidine (XCII).

H C I CH \ /

HgC-^ ^CHj

(XCII)

Various bases of this class have been found by

Adams and co-workers [652] and by N. J. Leonard

and Beck [653] to be more basic than analogous

acyclic amines.

b. Amidines. Substituted amidines,

R-C(-NR')NR^',

where the R's are alkyl or aryl, should probably

be mentioned as possible reference bases, but

they seem less promising for titrimetry than derivatives of guanidine, which will be discussed next.

c. Derivatives of guanidine. We have already re- ferred to 1,3-diphenylguanidine, an easily obtained

solid base (much used in the rubber industry), and

stated that in benzene it is a stronger base than

triethylamine. Its use as a primary standard in acid-base titrations was proposed by Carlton in 1922 [654], and it has been used in such varied

nonaqueous solvents as acetic acid, acetonitrile.

alcohols, and benzene. Davis and co-workers used , both diphenyl- and di-o-tolylguanidine successfully

for titrations in benzene [14, 313]. The somewhat

stronger base, 1,3-dicyclohexylguanidine, whose

hydrochloride at one time was available commer- ' cially as a research material, was not considered

soluble enough in benzene for practical use; it has

adequate solubility in chloroform, but reacts with

this solvent to some extent [14]. Di- and tri-alkylguanidines have not been obtain- ; able commercially. In fact, 1,3-dialkylguanidines

have not been very extensively investigated, perhaps i because for almost twenty years they were generally

believed to be rather weak bases [655, 190b], and

this error has not been universally corrected. \ However, there is now substantial evidence [656,

657] that partial or complete alkylation of guanidine

does not diminish the strength of this very strong ' base, but, instead, probably increases the strength | somewhat. | In the past, derivatives of guanidine have usually I been investigated as salts, and comparatively few

| seem to have been converted into the free bases, 'i However, Joshua [658] has described a preparation r of l,3-diphenyl-2-methylguanidine, and reported the fj melting point as 107°. On the other hand, Braun ? and Randall [659] evidently isolated 1,3-dibenzyl- i guanidine and 1,2,3-tribenzylguanidine only as •; hydrochlorides. ! It should be very interesting to compare the

behavior of 1,3-diphenyl- and l,3-diphenyl-2- ' methylguanidine, as this would test a suggested ex- planation [15] for the observation that diphenyl- I guanidine combines much more extensively than

triethylamine with carboxylic acids in benzene, ' although it is the weaker base in water. The tend- ency toward self-cissociation observed for 1,3- diphenylguanidine, and attributed to N-H...N

bonding, should be absent in its 2-methyl derivative. I Obviously, whatever the method used for locating

the titration end-point in an aprotic solvent, the

degree of success will depend on having satisfactory j primary standards. The search for improved standards, with aid from organic chemists, should be , an interesting and rewarding pursuit for analytical

and physical chemists wishing to improve titration i procedures.

The author thanks the investigators and pub- li lishers who graciously approved the reuse of

| published figures (see captions to figures 6, 8-12, [ 19-21, 28), and gives special thanks to the following ( investigators for contributing previously unpub- S lished results: A. R. Anderson and W. F. K. WynneJones (see table 75), R. A. Robinson (tables 68 and '

91), and R. Megargle, D. Rosenthal, and G. L. | Jones, Jr. (see figure 30).

136
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(1955). [616] H. C. Brown, S. Johnson, and H. Podall, J. Am. Chem.

Soc. 76, 5556 (1954). [617] H. C. Brown, D. H. McDaniel, and O. Hafliger, Ch. 14

("Dissociation Constants") in Determination of Organic

Structures by Physical Methods (E. A. Braude and

F. C. Nachod, Eds.) (Academic Press, New York, 1955). [618] S. F. Mason, J. Chem. Soc, p. 1247 (1959). [619] N. F. Hall and M. R. Sprinkle, J. Am. Chem. Soc. 54,

3469 (1932). [620] T. Zeegers-Huyskens, Bull. Soc. Chim. Belg. 69, 282

(1960). [621] L. Lamberts and T. Zeegers-Huyskens, J. Chim. Phys. 60,

435 (1963). [622] H. Dunken and H. Fritzsche, Z. Chem. 1, 127 (1961).

[623] R. P. Bell and A. F. Trotman-Dickenson, J. Chem. Soc,

p. 1288 (1949). [624] J. Rubin and G. S. Panson, J. Phys. Chem. 69, 3089 (1965). [625] T. Kubota, J. Am. Chem. Soc. 87, 458 (1965). [626] T. Kubota, J. Am. Chem. Soc. 88, 211 (1966). [627] M. M. Davis and M. Paabo, J. Res. NBS 64A, 533 (1960). [628] M. M. Davis, M. Paabo, and R. A. Robinson, J. Res. NBS

64A, 531 (1960). [629] R. G. Bates and G. Schwarzenbach, Helv. Chim. Acta 37,

1069 (1954). [630] G. Kortum and H. Wilski, Z. Physik. Chem. 2, 256 (1954). [631] J. F. J. Dippy, S. R. C. Hughes, and J. W. Laxton, J. Chem.

Soc, p. 2995 (1956). [632] M. M. Davis and H. B. Hetzer, unpublished results. [634] J. W. Bayles and B. Evans, J. Chem. Soc, p. 6984 (1966). [635] H. C. Brown and" S. Sujishi, J. Am. Chem. Soc. 70, 2878

(1948). [636] E. E. Sager, R. A. Robinson, and R. G. Bates, J. Res. NBS

69A, 305 (1964). [637] O. Folin and F. F. Flanders, J. Am. Chem. Soc. 34, 774

(1912). [638] L. Lykken, P. Porter, H. D. Ruliffson, and F. D. Tuemler,

Ind. Eng. Chem., Anal. Ed., 16, 219 (1944). [639] S. R. Palit, Anal. Chem. 18, 246 (1946). [640] J. Kucharsky and L. Safah'k, Titrations in Non-Aqueous

Solvents, 2d ed. (Elsevier, New York, 1965).

[641] D. Vorlander, Chem. Ber. 67, 145 (1934). [642] E. M. Trautner and C. E. Neufeld, AustraUan Chem. Inst.

J. & Proc. 13, 70 (1946). [643] G. Ya. Khalt, Farmatsiya 8, 26 (1945); see Chem. Abstracts 41, 2206 (1947). [644] R. Megargle, D. Rosenthal, and G. L. Jones, Jr., Clarkson

College of Technology, unpublished results (1966). [645] S. Bruckenstein and N. E. Vanderborgh, Anal. Chem. 38,

687 (1966). [646] E. B. R. Prideaux and R. N. Coleman, J. Chem. Soc, p. 462

(1937). [647] C. H. Giles, R. B. McKay, and W. Good, J. Chem. Soc,

p. 5434 (1961). [648] M. Ishidate, H. Nishizawa, H. Sano, and I. Horikoshi,

Yakugaku Zasshi 81, 1303 (1961), translated by I. Oiwa.

[649] H. C. Brown and N. R. Eldred, J. Am. Chem. Soc. 71,

445 (1949)

[650] S. Leonard and S. Elkin, J. Org. Chem. 27, 4635 (1962). [651] H. C. Brown and M. D. Taylor, J. Am. Chem. Soc. 69,

1332 (1947). [652] R. Adams, M. Carmack, and J. E. Mahan, J. Am. Chem.

Soc. 64, 2593 (1942). [653] N. J. Leonard and K. M. Beck, J. Am. Chem. Soc. 70,

2504 (1948). [654] C. A. Carlton, J. Am. Chem. Soc. 44, 1469 (1922). [655] T. L. Davis and R. C. Elderfield, J. Am. Chem. Soc. 54,

1499 (1932). [656] B. Neivelt, E. C. Mayo, J. H. Tiers, D. H. Smith, and

G. W. Wheland, J. Am. Chem. Soc. 73, 3475 (1951). [657] S. J. Angyal and W. K. Warburton, J. Chem. Soc, p. 2492

(1951). [658] C. P. Joshua, J. Sci. Indust. Res. (Delhi) B21, 588 (1962). [659] C. E. Braun and W. M. Randall, J. Am. Chem. Soc. 56,

2134 (1934). [660] E. F. Caldin and J. E. Crooks, J. Chem. Soc. [B], p. 959

(1967).

144
8. Author Index

Reference numbers (see preceding bibliography) are given in brackets after the name of each author. Other numbers designate

figures, tables, and pages in which references are cited or discussed. A

Accascina, F., [56]: figure 3; p. 8 Adams, R., [652]: p. 136 Adolph, H. G., [373]: p. 61 Affsprung, H. E., [267, 273]: table 23: pp. 31-32, 34

Agarwal, B. R., [iS6]: table 54: p. 89

Akerlof, G., [51]: figure 2: p. 8

Aldo, M., [1.56]: p. 20

Allan, E. A., [229]: p. 29

Allen, G., [250]: tables 24, 25: pp. 31-34, 36-37

Allerhand, A., [104. .525]: table 66: pp. 13-14, 16, 107, 109-110

AUred, A. L., [88, 119, 120]: tables 11, 12: pp. 11, 13, 15, 41 Ames, S. R., [45]: p. 7 Anderson, A. R., [594]: tables 73, 75: pp. 121, 123-124, 136 Anderson, D. M. W., [171]: table 17; pp. 20-24

Anderson, M. L., [39b]: table 2: p. 7 Anderson, R., [99]: p. 13 Andrews, L. J., [77. 78]: table 4: pp. 10, 20

Angyal, S. J., [657]: p. 136 Antony, A. A., [69]: figure 5: p. 8

Arnett, E. M., [71]: pp. 9, 81 Arnold, M. H. M., [247]: pp. 31-32

Ashby, E. C, [310]: pp. 38, 40 ASTM (Amer. Soc. Testing Materials), [44]: pp. 7, 133 Avery, R. A., [60]: figure 3: p. 8B

Baba, H., [401]: p. 72

li Bacon, G. E., [464]: table 55: p. 90

I Badger, G. M., [220]: p. 28

I Badger, R. M., [238]: p. 31

! Baes, C. P., Jr., [305. 306]: p. 38

j Babe, L. W.. [36a]: table 2: p. 7

I Bailey, G. F., [226]: table 21: p. 28

j Baker, A. W., [221-223]; table 22; pp. 28-30

i| Baker, H. T., [306]: p. 38

,1 Baker, W., [246]: p. 31

J Balinger, P., [.531]: table 68; p. 109

Bambagiotti, M., [380]; p. 63

j Banes, P. W.. [251a]; tables 24, 25: pp. 31. 33-34, 36-37

Barrow, G. M., [118, 265. 392-397, 528, .5.56-558]; tables 11, 43,

i 56, 66, 69, 70; pp. 15. 31, 64-73, 78, 91-92, 107, 110-111,

j 113-114

I Bartlett, P. D., [518]: table 66; pp. 106-107

Barton, B. C., [248]; pp. 31, 34

Bates, R. G., [598, 612, 629. 636]: tables 74, 82, 91; pp. 122,

] 127, 132

Batson, F. M., [348]; p. 50

I Bauer, S. H., [238]: p. 31

i; Bayles, J. W.. [354. 364. 441. 634]; tables, 75, 76, 78; pp. 54, 60,

81, 99, 115. 122, 124-125

i Beattie, I. R.. [501]; table 62: pp. 103-104

Beck, K. M., [6.53]; p. 136

Becker, E. D.. [122, 199, 606]: table 80; pp. 14, 25, 126 Beecher, L., [295]: p. 38

Beens, H., [571]; p. 114

Beichl, G. J.. [139]: p. 17

Bell, C. L.. [396, .5.57]: pp. 64, 71-72, 113-114

Bell, R. P., [8, 9, 247, 354, 404, 440, 491-493, 559, 567. 623];

table 76; pp. 2, 31-32. 54, 72. 81, 99, 101, 113-115. 123

Bellamy, L. J., [95, 210, 286, 295]; figure 7; tables 14, 20; pp. 13, 18-19, 26-27, 37-38

Bellon, L., [604]; p. 126

Benesi, H. A., [72]: pp. 9, 118 Berns, D. S., [61]; figure 4; p. 8 Bernstein, H. J.. (quoted in [90]); p. 12 Bertocci, U., [291]; table 27; pp. 37-38

Beskina, I. G., [347]; tables .38, 57; pp. 48-49, .52, 78, 93, 1.33, 1.35 Biczo, G., [565]; p. 113

Biggs, A. I., [215, .535]: tables 68, 82, 85, 86; pp. 27, 109, 127, 129

Birr, E. J., [323, 324]: pp. 42, 81 Bishop, R. J., [589]; tables 71, 79, 82; pp. 119. 126-127 Blandamer, M. J., [526, .527]: table 66: pp. 107, 110

Blinc, R.. [463]; table 55; p. 90

Blumer, D. R., [279]; p. 35 Bock, E., [419]: p. 76

Bodforss, S., [574]; pp. 115, 135 Bolander, B., [293]: p. 38

Bonnet, M., [600]; table 79; p. 126 Booge, J. E. [28]; p. 4

Borucki, L., [262]; p. 31 Bosch, W., [472]: table 57; pp. 90-91, 93 Bouck, J., [Ill]: p. 14 Bourassa-Bataille, H., [232]; pp. 29-30

Bower, V. E., [587]: table 71; p. 119

Bradley, K., [223a]: table 22; pp. 28-29

Brady, J. D., [80]; pp. 10-11

Briindstrom, A., [472a]: table 57; pp. 90-91, 93 Branch, G. E. K., [245]; p. 31 Brand, J. C. D., [127]; p. 16 Brandenberger, M., [416]; pp. 75-76

Braun, C. E., [659]; p. 136

Briegleb, G., [70]: p. 9

Briggs, A. A., [150]; p. 18

Brissette, C.. [376]: pp. 62-63, 113

Brode, W. R., [582, 583]: p. 117 Br0nsted, J. N., [10-13, 491]; table 3; pp. 2-4, 8-9, 101. 112, 114-115

Brooks, C. J. W., [244]: pp. 31, 35-36

Brown, A. C, [52]: figure 2: p. 8 Brown, H. C., [80. 512, 614-617, 635. 649, 651] ; tables 65, 83, 84;

pp. 10-11, 106, 118, 128-129, 136 Brown, I., [231]: pp. 29-30

Brown, J. P., Jr., [588]: table 71; pp. 118-119

Bruckenstein, S., [352, 406, 486, .500. 645]: figure 12: tables 56,

59, 62, 65: pp. 50-51, 68, 72, 78, 91-92, 97-98, 100. 103, 105-

106, 111, 13.5-136 Bruss, D. B.,[476, 484] ; tables 56, 59, 60: pp. 92, 95, 97-102, 133 Bryan, R. P., [467]; table 55; p. 90

Bryant, P. J. R., [474]; tables 56, 57, 59, 60; pp. 92-95, 97-100, 133-134

Bucher, J. J.. [442]; tables 49, 64; pp. 82, 105 Buckley, P., [147]: p. 18

Bufalini, J.. [521, 522]: tables 66, 70: pp. 107, 109, 111 Bullock, J. I., [409]; tables 45, 62, 65: pp. 73-74, 103, 105-106 Burawoy, A., [504]: p. 117

Burger, L. L., [412]: tables 46, 48; pp. 73-74. 76. 78-79

Burgess, D. S., [332]: table 32: pp. 43-44

Buswell, A. M.. [110]; p. 14 Byers, H. P., [489]; p. 102

Bystrov, V. P., [158, 183, 213]: table 17: pp. 20, 22-23. 26 c

Caldin, E. P.. [250, 493, .567, 660]; tables 24. 25; pp. 31-34,

36-37. 101, 114-115, 125

Calleri, M., [517]: p. 105

145
Calvet, E., [261]; p. 31

Calvin, M., [245b, 366]: pp. 31, 60-61

Campbell, R. J., [55, 58, 67]: figures 3, 5: p. 8 Cannon, C. G., [552]: p. 113 Cardinaud, R., [236]: p. 30

Carlton, C. A., [654]: p. 136 Carmack, M., [652]: p. 136 Carpenter, D. R., [123]; p. 14

Cassis, F. A., [505]; p. 104 Chantooni, M. K., Jr., [429, 479-481, 488, .500, 516]: tables 57, 58, 60, 62, 65, 70; pp. 79. 93-96, 99, 101, 103, 105-106, 110-111

Chariot, G., [578]; p. 116 Charman, H. B., [128]: p. 16 Chenon, B., [375]; pp. 62-63, 81, 113 Chetwyn, A., [364]; pp. 60, 99

Chittenden, R. A., [311]; p. 38

Choi, S. S., [409] : tables 45, 62, 65; pp. 73-74, 103, 105-106

Christian, S. D., [267, 273]; table 23; pp. 31-32, 34

Clever, H. L., [34]; pp. 5, 79

Cocivera, M., [421, 422]; pp. 76-77

Coetzee, J. F., [407, 438, 444-446, 487]; tables 50, 51, 57-60;

pp. 72, 81-85, 93-101

Coffen, D. L., [269]; pp. 31, 35

Coggeshall, N. D., [194, 202, 203]; tables 18-20; pp. 25-27 Coleman, R. N., [646]: p. 135

Collings, A. J., [276]; p. 35

Collins, R. F., [133]: p. 17 Colwell, J. E., [139]; p. 17 Conocchioli, T. J., [435]: pp. 79-81, 89, 97

Cooic, D., [423, 451]: table 52; pp. 76, 86-87

Cook, D. B., [398]; p. 68 Copenhafer, D. T., [350]; table 39; pp. 50-51

Coplan, M. A., [49]; p. 7 Copley, M. J., [105-107, 124]; pp. 13-15, 20-21

Coppinger, G. M., [136]; p. 17 Corse, J., [226]; table 21: p. 28

Cote, G. L., [384]; table 42; pp. 63, 76 Coulson, C. A., [551]: p. 113 Cox, B., [385]; table 42; pp. 63, 76 Creamer, L. K., [602]; tables 79, 82: pp. 126-127

Creswell, C. J., [88, 119, 120]; tables 11, 12; pp. 11, 13, 15, 41

Critchfield, F. E., [37, 53]; figure 2; table 2; pp. 7-8

Cromwell, N. H., [138]; p. 17 Crooks, J. E., [404, 660]; pp. 72, 125

Crossley, M. L., [483]: p. 95 Cunningham, G. P., [407, 445]; tables 50, 51, 57, 58; pp. 72, 82-85, 93-96

Curran, C, [3] ; p. 2 Curry, N. A., [464]; table 55: p. 90

D

Daasch, L. W.. [294]; p. 38 Dannhauser, W., [36a]; table 2: p. 7 Danyluk, S. S., [339]; p. 46 D'Aprano, A., [66]; figure 5; p. 8 Dauben, H. J., Jr., [518]: table 66; pp. 106-107

Davies, M. (or M. M.), [153a, 240, 242, 259, 281, 283, 344, 400, 460]; tables 23, 36, 55; pp. 20, 31-32, .35, 47, 70-71, 90

Davis, M. M., [5, 14-17, 20, 21, 29-33, 2.54, 2.54a, 313, 356, 359-362, 428, 482, .544, 546, 587, 590, 627, 628, 632]; figures 1, 13-18, 22, 23, 25-27; tables 24, 40, 50, 53, 56, 57, 60, 71-73.

75, 82, 90, 91. 93; pp. 2-5, 8-9, 20, 25, .30-31, 33-34, 36-37,

39, 47, 52-61, 67, 72-73, 78-79, 81-83, 87-88, 91-93, 95-97,

99, 101, 111-112, 115-122, 124, 127, 131-136

Davis, T. L., [655]; p. 136 Davison, J. A., [200]; p. 25 Dearden, J. C, [172, 235]; table 18; pp. 22, 25-26, 30 Decroocq, D., [267b]; pp. 31, 34

Deitz, v., [336]; table 33; pp. 43-44

de Maine, M. M., [150]; p. 18 de Maine, P. A. D., [150]; p. 18 Denison, J. T., [38]; table 2; p. 7 Denyer, R. L., [603]: p. 126

Detoni, S.. [315]; p. 40

Diamond, R. M., [430-436, 442, 456]; tables 49, 54, 64; pp. 79-82, 89, 97, 105 Dickerson, R. T., Jr., [496]: table 62: pp. 102-103

Dickinson, R. C, [193]: table 18; p. 26

Dierckx, A.-M., [608]: tables 80, 82, 84-86: pp. 126-127, 129

Dietzel, R., [540]; table 93; pp. 105-106, 111, 134 Dippy, J. F. J., [631]: table 91; p. 132

Dirksen, H. W., [41]; table 2; p. 7 Dischler, B., [262]: p. 31 Dodd, R. E., [174]; p. 22 Donohue, J., [377]; p. 62 Doub, L., [369]; pp. 60-61

Douglas, B. E., [437]; p. 81 Downes, H. C, [340, 357, 573]; pp. 46, 58, 115-116, 133 Drago, R. S., [137, 580]: table 87; pp. 17, 118, 130

Drain, L. E., [417]: p. 76 Dreisbach, P. F., [483]: p. 95

DriscoU, W. J., [513]: table 65; pp. 105-106 Duncan, J. L., [171]; table 17; pp. 20-24 Dunken, H., [249, 2.52, 2.58, 611, 622]: tables 24, 25, 82, 87; pp. 31, 33. 35-37, 127, 130 Dyrssen, D., [296-298]; tables 28, 29; pp. 38-40 Dyumaev, K. M., [213]; p. 26 E

Earp, D. P., [97]; p. 13 Edelson, D., [402]; p. 73

Eglinton, G., [127, 231, 244]; pp. 16, 29-31, 35-36 Ehrmann, K.. [132]; table 13: p. 17 Eigen, M., [567]; p. 114

Eisenberg, H., [280]: p. 35 Ekberg, S., [298]: table 29; pp. 38, 40

Elderfield, R. C, [655]; p. 136

Eldred, N. R., [649]; p. 136

Elkin, S.. [6.50]: p. 136 ' ElHott, J. H., [405]: p. 72 Emerson, M. T., [116]; tables 11, 12; pp. 14-15

Evans, B., [634]; p. 125 Evans. J. C, [155, 503, 504. 529]; tables 62, 63, 66; pp. 17, 20, 103-104, 110 EzeU, J. B., [.599]; tables 77, 90; pp. 125, 131 F

Fajans, K., [427, 530]; table 67; pp. 79, 108

Fasold, K., [22]; pp. 4, 115 Feakins, D., [447]; table 50; pp. 82-84

Feeney, J., [157]; pp. 17, 20 Feneant, S., [287]: p. 37 Ferguson, L. N.. [367]: p. 60

Ferraro, J. R., [299-302, 514]; table 65; pp. 38-39, 105-106 Few, A. v., [173]: p. 22

Finch, J. N., [230]; figure 8; pp. 29. 136

Findlay. T. J. V., [609]; table 80: p. 126

Finegold, H., [Ill]; p. 14

Fischer, A., [602]; tables 79, 82; pp. 126-127

Fischer, J., [.543]; table 93; pp. Ill, 134-135

Flanders, F. F.. [637]; p. 133

Flood, S. H., [82]: table 5: p. 10

Folin, O., [637]: p. 133 Forbes, W. F., [235, 269, 278]; pp. 30-31, 35 Forman, E. J., [596]; figure 29: table 74; pp. 122-123, 135

Foster, P. W., [138]; p. 17 Foster, R. W., [134]; p.l7 Fox. C. J.. [267c]; p. 34

Fox, I. R., [145]; p. 17 Fox, J. J., [282]: p. 35 Fredenhagen, K., [272]; p. 32 French, C. M., [337]; table 34; p. 45

Friedel, R. A., [205]; p. 26

Fritz, J. S., [50]; pp. 7, 133

Fritzsche, H., [611, 622]; tables 82, 87; pp. 127, 130

Fryer, C. W., [380]; p. 63

Fuller, W.. [.378]; p. 62

146
Fuoss, R. M., [47-49, 61-66, 326, 330, 331, 336, 402, 403, S19];

figures 4, 5; tables 32, 33, 66; pp. 7-8, 42-45, 56, 72-73, 79, 91, 107-108

Fuson, N., [84, 103, 166, 175, 177]; table 17; pp. 10, 13-14,21-23

G

Geddes, J. A., [4]; pp. 2, 79

Gelles, E., [538]; pp. 106, 108

Gergely, J., [565]; p. 113 Gibson, J. A., Jr., [37, .53]; figure 2; table 2; pp. 7-8

Gilchrist, A., [603]; p. 126

Giles, C. H., [647]; p. 135

Gilkerson, W. R., [448, 449, .599] ; tables 50, 53, 77, 90; pp. 82-84,

87-88, 125, 131 Ginsberg, E., [107]; p. 14, 20

Glasstone, S.. [96, 97]; p. 13 Goates, J. R., [98]; p. 13 Goddu, R. P., [207, 208]; p. 26 Goldenson, J., [Ill]; p. 14

Golic, L., [469, 470]; table 55; p. 90 Good, W., [647]; p. 135 Goodrick, D. A., [409]; tables 45, 62, 65; pp. 74, 103, 106 Gordon, J. E., [391. 478]; tables 57-59; pp. 64, 73, 93, 95-96, 98, 100 Gordy, W., [108, 109, 126. 1.53, 164, 165, 5.50]; table 17; pp. 14-15, 18, 21, 23, 113 Gormley, J. F., [39]; table 2; p. 7 Gough, T. E., [526, 527]; table 66; pp. 107, 110 Gould, J. H., [582, 583]; p. 117 Gouterman, M., [353]; tables 41, 79; pp. 53, 61-62, 126 Gramstad, T., [135]; tables 80-83, 86-88; pp. 17, 126-130

Grandberg, I. I., [183]; table 17; pp. 22-23

Greenfield, B. F., [289, 304]; tables 27, 28; pp. 37-39

Grellmann, K. H., [571]; p. 114

Griffiths, D. C, [355]; pp. 54, 115

.Griffiths, D. M. L., [281, 283]; p. 35

Gross, P. M., [256, 257]; pp. 31, 34-35, 41 Grunwald, E., [421, 548]; pp. 76-77, 112 Gurr, M., [571]; p. 114 Gur'yanova, E. M., [347]; tables 38, 57; pp. 48-49, 52, 78, 93,

133, 135 Gutowsky, H. S., [212]; table 20; pp. 26-27 H

Hadzi, D., [315, 390. 391. 450. 463. 477]; tables 52. .55-57. 59; pp. 40, 64, 73, 85-87. 90-93, 98. 100. 113

Hafliger. O., [617]; table 83; p. 128 Hagger, O., [416]; pp. 75-76

Hall, J. L., [37, 53, 55, 58, 67]; figures 2. 3. 5; table 2; pp. 7-8

Hall. N. F.. [595. 619]; tables 74, 86; pp. 122, 129 Hallam, H. E., [95, 152]; figure 7; table 14; pp. 13, 18-19

Halleux, A., [613]; tables 83, 86; pp. 128-129

Halpern, E., [Ill]; p. 14 Hammer, R. N., [39b]; table 2; p. 7 Hammett, L. P., [227-a, b, 515, 572]; table 65; pp. 28, 105-106,

115 Hanna, M. W., [581]; p. 118 Hantzsch, A., [18. 19. 425. 426, 584] ; pp. 4, 78, 105, 114-115, 117 Happe, J. A.. [178]; table 17; pp. 22-23

, Hardy, C. J., [289. 303. 304]; tables 27, 28; pp. 37-39

j Harlow, G. A., [476, 484]; tables 56, 59, 60; pp. 92, 95, 97-102.

133

^ Harrell, S. A., [511]; table 63; p. 104

Harris, J. T. Jr., [264]; table 23; pp. 31-32

i Hart, H., [505]; p. 104

Hartley, H. E. R., [311]; p. 38

Hatton, J. v., [87]; p. 11 Haurie, M., [275]; p. 35 Hay, L. D. (also see Liem, D. H.), [297]; table 28; pp. 38-39

Heaton, L. [318]; table 30; pp. 40-41

Hendricks, S. B., [216. 228]; pp. 27-28 Henne, A. L.. [267c]; p.34 Henrich. C., [372]; p. 61 Hepler, L. G., [532]; table 68; p. 109

Herbrandson, H. F., [496]; table 62; pp. 102-103

Hetzer, H. B., |5, 21, 31-33, 313, 482, .598. 6.32); figures 22, 23. 26, 27; tables 24, .50, .53, .56, .57, 60, 72-74, 90, 91, 93; pp. 2. 4.

5, 8-9. 20. 31. .33-.34, .36-37, .39. 52, .56-.59, 72-73, 78, 82-83.

87-88, 91-93, 9.5-97. 99, 101, 116-117, 119-122. 131-136

Hiibert, G. E., [228]; p. 28

Hildebrand, J. H., [72]; pp. 9, 118 Hirsch, E., [63]; figure 4; p. 8 Hobbs, M. E., [2.56, 257, 264]; table 23; pp. 31-.32, 34-.35. 41

Hfigfeldt, E., [292. 293] ; p. 38 H0jendahl.K., [146]; p. 17

Hiipfner, A., [160. 161]; tables 15, 16; pp. 20-21

HSpfner, H. M., [161]; tables 15, 16; pp. 20-21 Hoffsommer, J. C, [373]; p.61 Hofmann,C. M., [483]; p. 95

Holley, C. E., Jr., [124]; p. 15 Holmes, R. R., [512]; table 65; p. 106 Hoover, T. B., [39a]; table 2; p. 5 Horikoshi, I., [648]; p. 134 Hornig, D. F., [382]; pp. 63, 76 Howard, B. B., [116]; tables 11, 12; pp. 14-15 Hu, T., [.532]; table 68; p. 109 Huggins. C. M., [123, 195]; table 18; pp. 14, 2,5-26 Hughes, E. D., [497, .538] ; tables 61, 62, 66; pp. 102-104, 106-108 Hughes, S. R. C, [631]; table 91; p. 132 Hume, D. N., [485, 596]; figure 29; tables 59, 60, 74; pp. 98-101,

122-123, 133. 135 Hummelstedt. L. E. I.. [485]; tables 59, 60; pp. 98-101, 133 Hunt, C. K., [201]; p. 25 Hunter, L., [184. 206]; table 17; pp. 22-24. 26-28 Huong. P. V.. [129]; p. 16 Huyskens, P., [605, 608]; tables 80, 82, 84-87; pp. 126-127, 129-130 Hyne, J. B., [196, .524]; tables 18, 66; pp. 25-26, 107, 109

I

I'Haya, Y., [267a]; table 23; pp. 32. 34

Ingold, C. K., [81, 497, 538]; tables 61, 62, 66; pp. 10, 102-104, 106-108

Ingold, K. U., [211]; pp. 26, 28 Ingraham, L. L., [226]; table 21; p. 28

Ishidate, M.. [648]; p. 134

Ito, M., [197, 268]; table 18; pp. 26, 31, 35

Ives, D. J. G., [.52]; figure 2; p. 8

Izmailov, N. A., [26, 579]; pp. 2, 4, 116

J

Jiiger, G., [252]; tables 24, 25; pp. 31. 33. 35-37

Jaffe, H. H.. [225. 371]; pp. 28. 60-61

Jan, J., [307]; table 29; pp. 38-40

Janz, G. J., [339]; p. 46

Jasiriski, T., [.592]; tables 71, 77, 82; pp. 119, 125. 127

Joesten, M. D., [137]; table 87; pp. 17, 130 Johnson, B., [373]; p. 61 Johnson, J. R.. [218]; p. 28

Johnson, O.. [.530]; table 67; p. 108 Johnson. S., [616]; table 83; p. 128 Johnson, S. L., [399]; table 44; pp. 70-71

Jones, D. A. K., [92, 233]; tables 9, 10; pp. 12, 29-30

Jones, G. L., Jr., [644]; figure 30; pp. 134, 136

Jones, P., [259]; table 23; pp. 31-32

Jones, R. N., [365]; p. 134

Joop, N., [188]; table 17; pp. 23-24

Joshua, C. P., [6.58]; p. 136

Josien, M.-L.. [83. 84, 103, 129. 166, 175, 177. 275]; table 17; pp. 10, 13-14. 16, 21-23, 35

Julg, A., [600]; table 79; p. 126 Jumper, C. F., [116]; tables 11, 12; pp. 14-15 K

Kablukoff, I., [338]; p. 46

Kaeding, W. W., [222]; pp. 28-30

147
Kagarise, R. E., [101, 117]; table 11: pp. 13-15

Kamlet, M. J., [373]: p. 61 Katchalsky, A., [280]: p. 35 Kaufman, S., [473]: table 56: pp. 91-92, 96 Kawasaki, Y., [569, 570]: p. 114 Kearns. E. R., [114]; p. 14 Keder, W. E., [412, 420]; tables 46, 48; pp. 73-74, 76, 78-79

Keefer. R. M., [77, 78]; table 4; pp. 10, 20

Kehiaian, H. V., [316. 317]; p. 40 Kempter, H.. [193a]; table 18; p. 26

Kendall, J., [28]; p. 4 Kenttamaa, J., [40, 54]; figures 2, 3; table 2; pp. 7-8

Khart, G. Ya.. [643]: table 93; p. 134 Kidman, A. D., [609]; table 80; p. 126

Kieffer, P., [601]; table 79; p. 126

Kilpatrick. M., [181, 182, 285, 424]; table 17; pp. 22-24, 36, 78 Kirkwood, J. G., [149]; p. 18 Kitchen, L. J., [204]: pp. 26-27

Klingensmith, G. B., [145]; p. 17 Knight, A. R., [269, 278]; pp. 31, 35

Ko, H. C., [532]; table 68; p. 109 Kochergina. L. A., [284]: pp. 35-36 Kokubun. H.. [401]; p. 72

Kolthoff, I. M., [6, 7, 358, 406, 429. 438. 443, 472, 479-481, 488. 500, 516]; tables 50, 57, 58, 60, 62, 65, 70; pp. 2, 58, 72, 79, 81-83, 90-91, 93-96, 99, 101, 103, 105-106, 110-111

Kortiim, G.. [630]; table 91; p. 132

Kosolapoff, G. M., [113, 309, 310]; pp. 14. 38, 40 Kotowycz, G., [419]; p. 76

Kraihanzel. C. S., [130]; p. 16 Krasovec, F., [307]; table 29; pp. 38-40

Kraus, C. A., [4, 24, 248, 326-329, 331-335, 342. 348-351];

figures 9, 11; tables 32-36, 39, 50, 53; pp. 2, 4, 31, 34, 42-47,

49-51, 54, 56, 72, 79, 81-83, 88, 136 Kreevoy, M. M., [128]; p. 16

Krell, M. W., [167]; p. 21 Krueger, P. J., [93]; p. 13 Kubota, T., [625, 626]; table 90; p. 131 Kucharsky, J., [640]; p. 133 Kuhn, S. J., [82]: table 5: p. 10 Kuzina, M. G., [389]; table 42: pp. 63-64, 75 L

Ladik, J., [565]; p. 113 Lady, J. H., [112, 168]; table 11: pp. 14-15, 21-22

Lake, R. P., [286]; p. 37 Lamb, R. W., [277]: p. 35 Lambert, L. D., [154]; p. 20 Lamberts, L., [605, 621]: tables 80, 87; pp. 126, 130 La Mer, V. K., [340, 357, 573]: pp. 46, 58, 115-116, 133 Lang, E. M., [203]; p. 26 Langbein, W., [425]; pp. 78, 115 LaRochelle, J. H., [57]; figure 3: p. 8 Larsson. E., [471]: table 57; pp. 90-91, 93 Lascombe, J., [115, 129, 275]; pp. 14. 16. 35

Lassettre, E. N., [189, 193]; table 18; pp. 24, 26

Last, W. A., [447]; table 50; pp. 82-84

Latimer, W. M., [549]; p. 113 Laxton, J. W., [631]; table 91; p. 132 Leader, G. R., [39]; table 2; p. 7 Lee, M. K. T., [90]: table 8: pp. 12-13

Lee, T. S., [443]; table 50: pp. 82-83

Leermakers, J. A., [217]; p. 28 LePevre, R. J. W., [2.55]; table 24; pp. 31, 33-34

Leffler, J. P., [.548]; p. 112 Leicknam, J.-P., [103]: pp. 13-14

Leigh, G. J., [501]: table 62; pp. 103-104

Leonard, N. J., [6.53]; p. 136 Leonard, S., [650]; p. 136 ' Levy, H. A., [383]; pp. 63. 76 Levy, R. M., [524]; table 66; pp. 107, 109 Lewis, G. E., [585]; p. 117 Lewis, G. N., [2, 366]; pp. 2, 60-61

Lezina, V. P., [158, 213]; pp. 20, 26

Licata, S. B., [45]; p. 7 Liddel, U., [94, 216, 228]; pp. 13, 27-28

Liem, D. H. (also see Hay, L. D.), [298]; table 29; pp. 38, 40

Lien, A. P., [79]; p. 10

Lifson, S., [280]; p. 35 Lindberg, J. J., [40, 54, 68, 266]; figures 2, 3, 5; tables 2, 23; pp. 7-8, 31-32, 35

Linnell, R. H., [156, 176, 179]; table 17; pp. 20, 22-23

Lipovskii, A. A., [381, 389]; tables 42, 62, 63; pp. 63-64, 75, 103-104

Lippincott, E. R., [230, 564]; figure 8; pp. 29, 113, 136

Lo, G. Y-S., [503, 504, 529]; tables 62, 63, 66; pp. 103-104, 107, 110 L5wdin, P.-O., [561]; pp. 113-114

Long, P. A., [531]; table 68; p. 109 Longuet-Higgins, H. C, [554]; pp. 113, 131 Lord, R. C., [374]; p. 62 Lorenz, P. M., [545]; p. 112 Lovelace, M. E., [29]; figure 13; pp. 4, 25, 53-54, 58, 115, 117

Liittke, W., [219]; p. 28

Liittringhaus, A., [41]; table 2: p. 7 Lumbroso, H., [180]; table 17; pp. 22-23

Lund, H., [520]; tables 66-68; pp. 107-109

Lutskii, A. E., [284]: pp. 35-36

Lykken, L., [638]: p. 133 M

McAlonie, G. E., [150]; p. 18 McCaulay, D. A., [79]; p. 10 McClellan, A. L., [102]; pp. 13-14, 37, 78, 102, 112-113, 121 McCullough, J. P., [113]; p. 14 McDaniel. D. H., [437, 508, 511, 614, 617]; tables 63, 83, 84;

pp. 81, 104, 128-129 McDonald, E. A., [20, 361]; figure 16; table 75; pp. 4, 25, 47, .53, 56, 59, 99, 115, 117., 124, 134-135 MacDougall, P. H., [279]; p. 35 McDowell, M. J., [333]: table 32; pp. 43-44

Mcintosh, R. L., [330]; table 32; pp. 43-44 McKay, R. B., [647]; p. 135 McKinney, P. C, [556]: p. 113 McKinnon, L R., [100a]: p. 13 Madelung, W., [451]; pp. 75-76

Maguire, M. M., [198]; table 18: pp. 2.5-26 Mahan, J. E., [652]; p. 136 Maier, J. L., [513]; table 65; pp. 105-106

Maier, W., [262, 263]; tables 24, 26; pp. 31-33, 37, 39

Majani, C, [266]; table 23; pp. 31-32, 35 Malmberg, C. G., [36]; tables 1, 2; pp. 6-7 Mann, B. R., [602]; tables 79, 82; pp. 126-127 Manogg, P., [262]; p. 31 Marcus, S. H., [490]; p. 102

Marriott, J. A., [184]: table 17; pp. 22-24

Martin, A. E., [282]; p. 35 Martin-Smith, M., [231]: pp. 29-30

Marvel, C. S., [105-107]; pp. 13-14, 20-21

Maryott, A. A., [23, 35, 147, 257, 341, 343]; figures 10, 19-21;

tables 1, 2, 14, 36, 37, 56, 60, 65, 75, 78, 92; pp. 4, 6-7, 18, 31, 34, 41. 46-48, 60, 78, 81, 91-92, 94-95, 97, 99, 106, 115, 118, 124-125, 132-133, 135-136

Marzocchi, M. P., [380]; p. 63 Mason, G. W., [299, 300, 513, 514]: table 65; pp. 38-39, 105-106 Mason, S. P., [365. 618]; table 84; pp. 60, 81, 129 Mataga, N., [569, 570]; p. 114 Matheny, N., [.590]; table 71; p. 119 Mathieu, J. P., [387]; p. 63 Matsuyama, A., [401]: p. 72 Mayo, E. C.. [656]; p. 136 Mead, D. J., [330, 331]: table 32; pp. 43-44, 72 Mead, T. E., [597] : figures 28, 29; table 74; pp. 122-123, 13.5-136 Mears, W. H., [285]; p. 36 Mecke, R., [193a, 219]; table 18: pp. 26-28

Meerwein, H., [495]: tables 62, 66; pp. 102-103, 105, 107 Megargle, R., [644]; figure 30; pp. 134, 136 Melnick, A. M., [267]: table 23: pp. 31-32, 34

Merrifield, R. E., [374]; p. 62

Mettee, H. C., [93]: p. 13 Mihm, X. R., [615]; tables 83, 84; pp. 128-129

148
Miller, J. G., [139]; p. 17

Miller, R. C, [519]; table 66; pp. 107-108

Miller, S. I., [490]; p. 102 MiUs, H. H., [452, 466-468]; tables 52, 55; pp. 86-87. 90

Miron, R. R., [.593]; p. 122

Misiak, T., [.592]; tables 71, 77, 82; pp. 119, 125, 127 Mizushima, S., [239, 241]; p. 31 Moede, J. A., [3]; p. 2 Moelwyn-Hughes, E. A., [2.59]; table 23; pp. 31-32

Mohr, S. C, [528]: tables 66. 69; pp. 107, 110 Mok, S. F., [498]; table 62: pp. 103-104

Morgan, K. J., [276]; p. 35

Morino, Y., [239]: p. 31 Morman, J. F., [127, 244]; pp. 16, 31, 3.5-36 MiiUer, E., [132]; table 13; p. 17 Muggleton, D. F., [337]; table 34; p. 45 Mukherjee, L. M., [486]; table 59; pp. 97-98, 100

Mulliken, R. S., [73-76, .553]; pp. 9. 113 Muney, W. S., [444]; tables 50, 51; pp. 82-85

Murray, F. E., [274]; p. 35 Murthy, A. S. N., [288]; p. 37 N

Nagakura, S., [353]; tables 41, 79; pp. 53, 61-62, 126 Nageswara Rao, B. D.. [288]; p. 37

Neivelt, B., [656]; p. 136 Nenitzescu, C. D., [316]: p. 40

Neufeld, C. E., [642]: table 93; p. 134

Nikiforov, G. A., [213]; p. 26

Nikitina, S. A.. [381]; tables 62, 63; pp. 63, 103-104

Nikuradse, A., [121]; p. 14 Nishizawa, H., [648]; p. 134 Nissema, A., [54]; figure 3; p. 8 Novak, A., [391, 477]; tables 56, 57. 59; pp. 64, 73, 91-93, 98. 100

Nuttall, R. H., [388]; table 42; pp. 63-64, 76 Nyquist, R. A., [237]; p. 30 o

Ogimachi, N., [78]; pp. 10, 20

Ogston, A. G., [534]; table 68; p. 109 O'Hara, W. F., [532]; table 68; p. 109 Olah, G. A., [82]; table 5; p. 10 Onsager, L., [148]; p. 18 Orchin, M., [371]; pp. 60-61

O'Sullivan, D. G., [186]; table 17; pp. 23-24

Ott, G. H., [416]; pp. 75-76 On, J. B., [98]; p. 13 P

Paabo, M., [15. 254, 254a, 362, 587, 590, 627, 628]; figure 27; tables 24. 71-73, 82, 91; pp. 3-4, 31, 33-34, 52, 54, 59, 67, 99, 101, 116, 119-122, 127, 132, 136 Pace, R. J., [286]; p. 37

Packer, J., [602]; tables 79, 82; pp. 126-127 Padmanabhan, G. R., [445, 446, 487]; tables 50, 51. 59, 60; pp. 82-85, 97-101

Palit, S. R., [639]; p. 133

j Palm, K., [258]; table 24; pp. 31, 33

I Panson, G. S., [610, 624]; tables 82, 84. 89; pp. 127, 129-130

i.j Paoli, C., [261]; p. 31 Parker, A. J., [141]; p. 17

I Parker, R. P.. [483] ; p. 95

Patai, S., [497]; tables 61, 62, 66; pp. 102-104, 107

Patnaik, D., [259]; table 23; pp. 31-32

Paul, W., [540]; table 93; pp. 105-106, 111, 134

Pauling, L.. [190-a, b]; table 63; pp. 25, 27-29, 31, 104, 113, 136 Pearson, R. G., [591] ; tables 71, 75, 77, 79, 89; pp. 119, 124-126,

130 Pegg, J. A., [603]; p. 126 Pendred, D., [418]; p. 76 Peppard, D. F., [299-302, 513, 514] ; table 65; pp. 38-39, 105-106

Person, W. B., [76, 99a]; pp. 9, 13

Peterson, .S. W., [.383]; |)p. 63, 76

Petrucfi, S., [56]; figure 3; p. 8

Pfeiffer, P., [4.58]; tal)le .55; pp. 89-90

Fhilbrick, F. A., [192]; table 18; pp. 2.5-26 Fimentei, (;. C, [102, 195, .5471; table 18; pp. 1.3-14, 25-26, 37,

78, 102, 1 12-113, 121 Pineau, P., [84, 177|; table 17; pp. 10, 22-23

Fitzer, K. S., [506J; table 63; p. 104 Flourde, G. R., [319]; table 31; p. 41 Flumb, R. €., [382]; pp. 63, 76 Focker, Y., [497,499,539]: tables 61 , 62, 66; pp. 102-104, 106-108

Podall, H., [616]; table 83; p. 128

Pohl, H. A., [256]; pp. 31, 35 Popov, A., [39b]; table 2; p. 7 Popovych, O.. [46, 85]; tables 6, 76, 77; pp. 7, 10-1 1 , .54. 99. 118. 124-125

Porter, P., [638]; p. 1.33 Poulet, H., [387]; p. 63

Powell, D. L., [90]; table 8; pp. 12-13

Powell, J. S., [309]; p. 38

Powell, R. L.. [166]; table 17; pp. 21, 23

Pratt, A. L.. [143]; p. 17 Prausnitz, J. M., [99]; p. 13

Price, E., [145]; p. 17 Prideaux, E. B. R., [646]; p. 135

Prigogine, I., [607]; tables 80, 81: pp. 126-127

Proskauer, E. S., [131]; pp. 16, 72 Pschenichnov, E. A., [560]; p. 113 Puttnam, N. A., [209]; table 68; pp. 26-28, 109 R

Raab, F., [156]; p. 20 Ralph, E. K., Ill, [448, 449]; tables 50, 53, 90; pp. 82-84, 87-88,

131 Ramsey, J., [38]; table 2; p. 7 Randall, W. M., [659]; p. 136 Rao, C. N. R., [224, 288. 370]; pp. 28, 37, 60-61

Reeves, L. W., [229, 270]; pp. 29, 31 Reynolds. M. B., [334]; table 32: pp. 43-44

Richards, R. B., [602]; tables 79, 82; pp. 126-127

Richards, R. E., [87, 418]; pp. 11, 76 Richardson, E. A., [523]; p. 109

Riddick, J. A., [131]; pp. 16, 72 Rieger, M., [214]; p. 27

Rieseberg, H., [262]; p. 31 Rieux, A., [576]: p. 116

Ritchie, C. D., [143. 144]; p. 17 Roberson, W. E., [167]; p. 21 Roberts, R. M., [185]; table 17; pp. 22-23

Robinson, R. A., [271, 533, 535, 587, 612, 628, 636]; tables 23. 24, 68, 71, 72, 74, 79, 82, 85, 86, 91; pp. 32-33. 109, 119. 122.

126-127. 129. 132. 136 Rodebush, W. H., [110, 549]; pp. 14, 113

Rolandi, G., [291]; table 27; pp. 37-38

Rose, N. J., [580]; p. 118

Rosenthal, D.. [644]; figure 30: pp. 134, 136 Rossmy, G., [234]; p. 30

Rossotti, F. J. C., [169-171]; table 17; pp. 20-24

Rossotti, H., [169, 170]; table 17, pp. 20-21. 23-24

Rothrock, D. A., Jr., [349]; p. 50 Rouse, P. E., [251b]: tables 24. 25; pp. 31. 33. 36

Roy, M. F., [110]: p. 14 Rubin, J., [610, 624]; tables 82, 84, 89: pp. 127. 129-130

RudakofF. G., [249]: p. 31 Rudolph, J., [563]; pp. 248-249

Ruhffson, H. D., [638]: p. 133 Rumeau. M., [575. 576]; p. 116 Rumon, K. A.. [399]; table 44; pp. 70-71 Rumpf. P.. [601]; table 79; p. 126 Ryan, J. L., [413. 414]; table 47: pp. 53, 73-75

Ryzhikova. L. A.. [59]; figure 3; p. 8

S

Sacks, F. M., [48]; p. 7 Sadek,H., [47, 62, 64, 65]; figure 4; table 66; pp. 7-8. 17. 107-108

149
Safah'k, L., [640]; p. 133

Sager. E. E.. [636]; table 91; p. 132

Saier, E. L., [194]; tables 18, 20; pp. 25-27

Saito, A., [352]; figure 12; table 56; pp. 50-52, 68. 78, 91-92, 135-136

Salthouse, J. A., [507]; table 63; p. 104

Sandorfy. C, [232, 375, 376]; pp. 29, 62-63, 79, 81, 113 Sano, H., [648]; p. 134

Sato, T., [410, 411]; table 45; pp. 73-75

Saunders, M.. [196]; table 18; pp. 25-26

Sauvageau, P.. [232]; p. 29 Sawyer, D. W., [210]; p. 25

ScargiU, D., [289, 303]; tables 27, 28; pp. 37-39

Schaefer, T., [419]; p. 76 Schleyer, P. von R.. [90, 91, 104, 525]; tables 8, 66; pp. 12-14,

16, 107, 109-110

Schlundt, H., [42]; p. 6 Schmidt, U.. [163]; p. 20

Schneider, W. G., [86, 243]; figure 6; pp. 11, 31, 136 Schroeder, R., [564]; p. 113 Schuhmann, P. J., [14, 29, 360, 428, 544, 546]; figures 13, 14; tables 40, 73, 93, pp. 3-4, 25, 30, 39, 52-59, 78-79, 99, 111-112,

115, 117, 119, 121, 134-136 Schwarzenbach, G., [416. 629]; table 91; pp. 75-76, 132

Sears, W. C., [204]; pp. 26-27

Senkowski, B. A.. [610]; table 82; p. 127 Sharova, G. I., [183]; table 17; pp. 22-23

Sharp, D. W. A., [388]; table 42; pp. 63-64, 76 Sharpe. A. G.. [385]; table 42; pp. 63, 76 Sharpe, A. N., [100]; p. 13 Shatenshtein, A. I., [25]; pp. 2, 4, 9, 16, 56, 113 Shaw, F. H., [541]; table 93; pp. Ill, 134-135 Shaw, R. A., [447]; table 50; pp. 82-83

Shibuya, T.. [267a]; table 23; pp. 32, 34

Shoolery, J. N., [195]; table 18; pp. 25-26

Shrivastava, H. N., [465]; table 55; p. 90

Shulgin, A. T., [223]; table 22; pp. 28-30

Singleterry, C. R., [473J; table 56; pp. 91-92, 96 Skarzynska, T., [592]; tables 71, 77, 82; pp. 119, 125, 127

Skavinski, E. R., [308]; p. 38

Skinner, J. M., [461]; table 55; p. 90

Smith, D. C., [294]; p. 38

Smith, D. H., [656]; p. 136 Smith, E. R., [35]; tables 1, 2, 14, 78, 92; pp. 6-7, 18, 125, 132 Smith, J., [603]; p. 126 Smith, J. W., [173, 566]: pp. 22, 113 Smith, N.. [457]; pp. 89-90

Smith, T. L., [405]; p. 72 Smyth, C. P., [43, 69, 146]; figure 5; pp. 6, 8, 17 Sobczyk, L., [345, 346, 400]; table 37; pp. 48-49, 70-71

Sokolov, N. D., [27, 560]; pp. 4. 113 Somers, B. G., [212]; table 20; pp. 26-27

Sourisseau, G.. [83]; p. 10 Speakman, J. €., [452, 457, 459, 461, 462, 465-470, 517]; tables

52, 55; pp. 86-87, 89-90, 105 Spinrad, B. L., [402]; p. 73 Spivak, L. L., [579]; p. 116 SprengHng, G. R., [494]; p. 100

Sprinkle, M. R.. [619]; table 86; p. 129

Spurr, R. A., [260a, 489]; table 23; pp. 31-32, 102

Stanford, S. C., [126, 165. 550]; table 17; pp. 15, 21, 23, 113

Starobinets, G. L.. [59]; figure 3; p. 8

Starobinets, K. S., [59]; figure 3; p. 8 Staschewski, D., [159]; p. 20 Steigman, J., [545]; p. 112 Stephenson, G. W., [174]; p. 22

Stern. K. H., [521-523]: tables 66, 70; pp. 107, 109, 111 Stevenson, D. P., [136]; p. 17 Stewart, G. M. D., [461]; table 55; p. 90

Stieglitz, J., [1]; pp. 1-2

StiUson, G. H., [201]: p. 25'

Stitt, F., [226]; table 21; p. 28

Stocesoca, D., [443]; table 50; pp. 82-83

Stokes, R. H., [271]; tables 23, 24, 68, 71, 74, 79; pp. 32-33, 109,

119, 122, 126

Stone, H. W., [308]; p. 38

Streuli, C. A., [593]; p. 122

Strong, E. D. T., [154]; p. 20

Sujishi, S., [635]; pp. 118, 136 SuUivan, R. J., [98]: p. 13 Sundaram, S.. [274]; p. 35

Sutchffe, L. H.. [157]; pp. 17, 20

Sutherland, G. B. B. M., [240]; p. 31 Sutton, L. E., [589, 603]; tables 71, 79, 82; pp. 119, 126-127

Swain, C. G., [537, 588]; table 71; pp. 106, 118-119 Symons, M. C. R., [526, 527]; table 66; pp. 107, 110

Syrkin, J. K., [345]; table 37; pp. 48-49

Szafran, M., [453-455]: tables 52, 55; pp. 85-87, 90

Szigeti, B., [363]; p. 60 T

Taft. R. W.. [145]: p. 17

Taylor, A. F., [441]; tables 75, 78; pp. 81, 99, 124-125

Taylor, E. G., [335]; table 32"; pp. 43-44

Taylor, M. D.. [320, 651]; pp. 41, 136

Taylor, S. A., [273]: p. 32 Templeman, M. B., [320]; p. 41 Thomas, L. €., [311]: p. 38 Thomas, W. J. 0., [460]; table 55; p. 90 Thompson, H. W., [384]: table 42; pp. 63, 76

Thyrion, F., [267b]; pp. 31, 34

Tiers, J. H., [656]; p. 136

Tiessens, G. T., [536]: table 68: p. 109 Tocher, M. I., [434, 435]; pp. 79-81, 89, 97 Tomlinson, T. E., [603]; p. 126 Toops, E. E., Jr., [131]; pp. 16, 72

Torihashi, Y.. [569. 570]; p. 114

Trautner, E. M., [541, 642]: table 93; pp. Ill, 134-135

Tremillion, B., [575-578]; p. 116 Trotman-Dickenson, A. F., [439, 623]; pp. 81, 123

Trotter, P. J., [581]: p. 118 Tuck, D. G., [409, 431, 502]; tables 45, 62, 63. 65; pp. 73-74,

79-80. 89, 97. 103-106 Tuemler, F. D.. [638]: p. 133 Turner, R. J.. [483]; p. 95 u

Ubbelohde. A. R.. [555]; pp. 113. 131

Ulbrich. R.. [121]; p. 14 Ungnade, H. E.. [277]; p. 35

Utterback, E., [166]; table 17; pp. 21, 23 V

Vallee. R. E.. [508]: table 63; p. 104 Vandenbelt. J. M.. [369. 372]; pp. 60-61 Vanderborgh. N. E., [645]; p. 135 van der Heijde, H. B.. [475]: tables 56, 59, 60; pp. 92, 95, 97-99, 101-102 Van Looy. H.. [515]: table 65; pp. 105-106 Vaughan. J., [602]; tables 79. 82; pp. 126-127

Venkateswarlu, P., [288]; p. 37 Vernon, A. A., [57, 60]; figure 3; p. 8 Verstegen, J. M. P. J., [290, 408]; tables 27, 45; pp. 37-38, 73-74

Vinard, D. R., [128]: p. 16 Vine, H.. [255]: table 24; pp. 31, 33-34

Vinogradov, S. N., [176, 182]; table 17; pp. 22-23

Vogelsong, D. C., [591]: tables 71, 75, 77, 79, 89; pp. 119, 124-126,

130

Voigt, W., [19]: pp. 4, 78, 115 von Halban, H., [363]; p. 60 von Keussler, V., [234]; p. 30

Vorliinder, D., [542, 543, 641]; table 93; pp. Ill, 133-135 w

Waddington, T. €., [386, 388, 507, 509, 510]; tables 42, 63; pp. 63-64, 76, 104 Wahr, J. C., [155]; pp. 17, 20 Walden, P., [322-324]; pp. 41-42, 45, 72, 81 Waldron, R. D., [379]; pp. 62-63 150
Walker, S., [100]; p. 13 WaO, F. T., [251-a, b]: tables 24, 25; pp. 31, 33-34, 36-37 WaUing, C, [318]; table 30; pp. 40-41 Warburton, W. K., [657]; p. 136 Wardrop, A. W. H., [474]; tables 56, 57, 59, 60; pp. 92-95, 97-100, 133-134

Watkinson, J. G., [92, 233]; tables 9, 10; pp. 12, 29-30

Weinstein, J., [496]; table 62; pp. 102-103

Weissberger, A., [22, 131, 217, 253]; table 24; pp. 4, 16, 28, 31, 33-34, 72, 115

Weller, A. H., [568, 571]; table 92; pp. 114, 130, 132 Wenograd, J., [260-a, b]; table 23; pp. 31-32 Wesson, L. G., [142]; p. 38 West, R., [89-91, 130, 198]; tables 7, 8, 18; pp. 11-13, 16, 25-26,

41 Westheiiner, F. H., [214]; p. 27 Westrum, E. F., Jr., [506]; table 63; p. 104 Whatley, L., [90]; table 8; pp. 12-13 Wheeler, M. W., [138]; p. 17 Wheland, G. W., [191, 368, 656]; pp. 25, 60-61, 136 Whetsel, K. B., [101, 112, 167, 168]; table 11; pp. 13-15, 21-22

White, N. E., [181]; table 17; pp. 22-24 Whitmore, F. C, [125]; p. 15 Whitney, D. C, [432-434]; pp. 79-81, 89, 97

Wildner, F., [543]; table 93; pp. Ill, 134-135

Wilk, W. D., [528]; tables 66, 69; pp. 107, 110

Williams, G., [344]; table 36; p. 47

Williams, H., [140]; p. 17 WiUiams, R. L., [95, 151, 210]; figure 7; tables 14, 20; pp. 13, 18-19, 26-27

Williamson, A. G., [100a]; p. 13 Wilski, H., [630]; table 91; p. 132

Wilson, A. S., [420]; p. 76 Wimette, H. J., [179J; table 17; pp. 22-23

Witschonke, C. R., [24]; tables 32-35, 50, 53; pp. 4, 43-45, 56,

72, 81-83, 88

Wolff, H., [159-163]; tables 15, 16; pp. 20-21 Woodhouse, E. J., [409, 502]; tables 45, 62, 63, 65; pp. 73-74, 103-106

Wright, W. G., [314]; p. 40

Wull, O. R., [216, 228]; pp. 27-28 Wyman, G. M., [.582, 583]; p. 117 Wyman, J., [60]; figure 3; p. 8 Wynne-Jones, Lord (W. F. K.), [325, 594]; tables 73, 75; pp. 42-43,

45, .54, 121, 12.3-124, 1.36 Y

Yabroff, D. L.. [245a]; p. 31 Yerger, E. A., [118, 265, 392-394]; tables 11, 43, 56, 70; pp. 15, 31, 64-68, 70, 78, 91-92, 111 Young, H. S., [351]; table 39; pp. 50-51 z

Zadorozhnyi, B. A., [284]; pp. 35-36

Zeegers-Huyskens, T., [605, 608, 620, 621]; tables 80, 82, 84-87;

pp. 126-127, 129-130

Zellhoefer, G. F., [105, 106]; pp. 13-14, 21 Zimmermann, H., [187, 188, 562, 563]; figure 24; table 17; pp. 23-24, 113-114

Zollinger, H., [586]; p. 117 Zwolenik, J. J., [403]; pp. 73, 91

151

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