Chapter 1 Nuclear magnetic resonance of organic charge-transfer complexes

CHAPTER 1

NUCLEAR MAGNETIC RESONANCE OF ORGANIC CHARGE-TRANSFER

COMPLEXES

ROY FosTmt and COLIN A. FYFES

Chemistry Department, Queen’s College, University of St. Andrews, Dundee, Scotland

CONTENTS

Glossary of Symbols

1. Introduction 1.1. General, Electron Donors and Electron Acceptors 1.2. Mulliken’s Valence Bond Treatment

2. NuclearMagneticResonanceMeasurements ofFast Equilibria in theLiquid Phase

2.1. Introduction 2.2. Intermolecular Interactions Involving Hydrogen Bonding

2.2.1. Hydrogen bonds formed froma-H 2.2.2. Hvdroeen bonds formed from N -H 2.2.3. Hydrogen bonds formed from S-H 2.2.4. Hydrogen bonding through acetylenic protons 2.2.5. Hydrogen bonds formed between =C-H protons and aromatic

rr-systems 2.2.6. Hydrogen bonds formed between EC--H protons and n-electron

donors 2.3. Carboxylic Acid Associations 2.4. Intermolecular Interactions Involving Boron Trihalides 2.5. Interaction of Amides and of Other Carbonyl Compounds with Aromatic

Compounds 2.5.1. Amide-aromatic compound interactions 2.5.2. Carbonyl compound-aromatic compound interactions

2.6. Charge-transfer Associations 2.6.1. Aromatic-aromatic interactions 2.6.2. Recognized charge-transfer complexes

3. The Determination ofAssociation Constants ofOraanicCharae-transfer Complexes in DiluteSolution

” _

3.1. General: the Effects of Concentration Scales. Solvent Comnetition, Specific Solvation, Isomeric Complexes

3.2. The Determination of Association Constants for Charge-transfer Complexes in Dilute Solution from Chemical Shift Data and Evidence for Their Correctness

2

4 4 6

9 9

10 10 14 16 18

19

20 22 24

27 27 29 34 34 36

38

38

41

tPresent address: Chemistry Department, University of Dundee, Dundee, Scotland. *Present address: Chemistry Department, University of British Columbia, Vancouver,

B.C.. Canada.

ROY FOSTER AND COLIN A. FYFE

3.3. Results Obtained from Chemical Shift Measurements of Charge-transfer Equilibria in Dilute Solution 50 3.3. I. Summary of results SO 3.3.2. A general discussion of association constant values 50 3.3.3. A discussion of h,, values 62 3.3.4. Systems ofbiological interest 67

3.4. Other Physical Methods of Measuring Association Constants of Charge-transfer Complexes in Dilute Solution 70

3.5. A Comparison between Association Constants Evaluated from Chemical Shifts and from Electronic Spectral Data 73

3.6. Evaluation of Optical Absorptivity of Charge-transfer Complexes from a Combination of Chemical Shift and Optical Data 77

4. NuclearMagnetic Resonance andNuclear Quadrupole Resonance Studies of Solid Charge-transferComplexes 78

Acknowledgements 82

6

SC,

6,

A

A0

A,

A:

V

VI1

px

A D

chemical shift for an equilibrium mixture with respect to a reference standard separate from the reactants. chemical shift for a pure complex with respect to a reference standard separate from the reactants. chemical shift for an equilibrium mixture at infinite dilution with respect to a reference standard separate from the reactants. chemical shift for an equilibrium mixture with respect to one of the reactants. chemical shift for a pure complex with respect to one of the reactants. chemical shift for an equilibrium mixture at infinite dilution with respect to one of the reactants. chemical shift difference for a solute in a dilute solution in solvent Y compared with a solution in solvent S. frequency of an NMR absorption. frequency of NMR absorption of an n-mer. relative population at a site X. electron acceptor, or hydrogen donor. electron donor (hydrogen acceptor).

[Cl] equilibrium concentration of a species. [O],, total (free and complexed) concentration of a reactant species. Nx total (free and complexed) mole fraction of a species X.

Mx molecular weight of species X.

& density of species X.

GLOSSARY OF SYMBOLS

AH AG AS

f AD

~‘max

A

A* AD

E.4

K KAD

KC K.V K,

ORGANIC CHARGE-TRANSFER COMPLEXES

enthalpy of formation subscript c indicates a molar- free energy of formation concentration scale; subscript N entropy of formation I indicates a mole-fraction scale. oscillator strength. frequency in cm-l of the maximum intensity of an optical charge- transfer transition for a complex AD. transition dipole for an optical intermolecular charge-transfer transition. wavelength. absorptivity (optical density) at wavelength X. molar absorptivity (extinction coefficient) of a species AD at a wavelength A. association constant. association constant for the formation of the species AD. association constant on a molar concentration scale. association constant on a mole-fraction scale. association constant on a molal concentration scale.

4 ROY FOSTER AND COLIN A. FYFE

1. INTRODUCTION

1.1. General, Electron Donors and Electron Acceptors

The term “charge-transfer complex” is here used to describe the product of the relatively weak interaction between electron-donor and electron-acceptor species. Although the term “charge-transfer complex” is widely used it can be a misleading description. There is often very little contribution from charge-transfer forces in the ground state of the complex. Even in complexes where stabilization through charge transfer is significant, there may be larger contributions to the ground state interaction through normal van der Waals forces. The complexes dissociate, at least partially, in the vapour phase and in solution. It must be recognized that the limits of this type of interaction are arbitrary. There is a continuous gradation of the degree of interaction which, as it increases, eventually assumes the character of a full chemical bond.

This chapter is generally limited to a discussion of the complexes formed betweep organic electron acceptors and electron donors.? Typical of the stronger electron acceptors which have been used in such studies are: polynitroaromatics, aromatic nitriles, quinones (particularly those containing electron-attracting groups), acid anhydrides, tetracyanoethylene and tetracyanoquinodimethane. Many inorganic species may act as electron acceptors to form similar weak complexes. Although these are mainly outside the scope of this chapter, the properties of such complexes, particularly those of iodine,“) have been central in the development of the theory of charge-transfer complexes. Most of the stronger organic acceptors involve T*-orbitals and may be classified as n-acceptors. Some organic species such as carbon tetrachloride act as electron acceptors through the involvement of cr*-orbitals. These organic cr-acceptors are usually weak and contrast with some inorganic a-acceptors such as iodine and iodine monochloride which are very effective acceptors. The various organic electron donors which have been used include: aromatic hydrocarbons, aromatic and aliphatic amines, phenols, alcohols, ethers, sulphides, and various heterocyclic systems. In some complexes, for example those involving aromatic hydrocarbons, donation is from a n-orbital (n-donor), whilst with other donors, the interaction is through an electron from a non-bonding orbital (n-donor), as in the case of sulphides, alcohols, ethers and aliphatic amines. In some

PHowever, some discussion of hydrogen-bonded complexes is given for reasons stated below.

ORGANIC CHARGE-TRANSFER COMPLEXES 5

molecules the donor function may be dependent on the particular acceptor species. Thus aromatic amines behave as r-donors towards m-acceptors such as 1,3,5trinitrobenzene, whereas they appear to act as n-donors in complex formation with iodine. c2) In any case the terms “electron donor” and “electron acceptor” are relative. Some molecules such as the nitro- terphenyls can assume the role of both donor and acceptor.c3) Other species can form self-complexes; thus the species I, has been considered to be a charge-transfer complex of two iodine molecules.(4*5)

The interaction of an electron donor and an electron acceptor to form a charge-transfer complex will often involve an enthalpy change of the order O-4 kcal mole-‘, although examples of interactions with enthalpy changes of about 12 kcal mole-l are known. The complex is usually characterized by the appearance of one or more electronic absorption bands which are not observed in either component alone. The presence of such an absorption has been taken by some authors as diagnostic of a charge-transfer complex. To make this a definition of a charge-transfer complex could lead to difficulties since the transition may be of such energy as to be experimentally undetectable through the presence of other overlapping bands. For example, the silver ion-benzene interaction has been thought of as a charge-transfer complex for 15 years, although only recently has the charge-transfer band been observed.@) In many cases solid complexes with a simple stoichiometry may be isolated from solution or from melts. The existence of such products both in the solid phase and in solution has been known for many years.c7) Mulliken@) has given a widely used theoretical description of such interactions. This is described briefly in the next section.

Charge-transfer forces do not act to the exclusion of other forces in these weak interactions. In particular, for appropriate systems there may be significant contributions from hydrogen bonding. Because of the lack of a sharp division between types of complexes, and because of the numerous excellent analyses of shift data developed for hydrogen bonded and other systems, the methods described are applicable to all fast-exchange reactions; these are included in Section 2.

A number of general reviews on charge-transfer complexes have been published$“-17) as well as others more specifically concerned with the energetics,(18) spectra,(1g-23) reactions of charge-transfer complexes,(24) and of their relationship with biochemistry.c25) The reader is referred to these for details of other properties of organic charge-transfer complexes.


1.2. Mulliken’s Vulencr Bond Trerrtment

MullikeiP) has described the relatively weak interaction between an electron donor (D) and an electron acceptor (A) in terms of a wave function of the form

$J~(DA) = a$,,(D.A)+b$,(D+-A-). (1)

The acceptor and the donor may in general be molecules, molecule-ions, atoms or atom-ions with the restriction that they are both in their totally symmetric ground states. The wave function & has been termed by Mulliken the “no-bond” function. It represents the structure in which the binding between the two parts of the complex results from dipole-dipole, dipole-induced-dipole interactions, London dispersion forces. hydrogen- bonding and other similar forces. The wave function $r has been termed by Mulliken the “dative” function. It corresponds to a structure in which one electron has been transferred from the donor to the acceptor whilst remaining paired, so that the singlet-spin multiplicity is maintained. This amounts to a covalently bonded structure, but the binding energy is relatively low because of the large separation of A and D. Where the interaction is weak the no-bond structure is the main contributor to the ground state: that is, n + h in equation (1).

There will be an excited state $E corresponding to I,!J~ where:

$E = N*$,(D+-A-) --b*&(D,A) (2)

where (I = CI” and b = b*. The optical absorption characteristic of the complex as a whole. corresponds to the transition tj~:~ -+ +!J~. It is an intermolecular charge-transfer transitiont

In order to obtain a finite contribution from the dative structure there must be an overlap of the filled molecular orbital in D which contains the electron to be donated, with the vacant orbital in A which is to receive the electron. It has been assumed that there will be a tendency for the two components of the complex to orientate themselves in such a way that this overlap is a maximum (overlap and orientation principle).

The integrated intensity of the charge-transfer band, or the oscillator strength,f, is given by the expression:

where ZJ,~):~ is the frequency in cm-’ of the maximum intensity of the band and pfi;V is the transition dipole which may be shown to be related approx-

+Muitiple transitions are observed in some complexes.


e+A+ D+ --------

/ / A-+D+

/ /

/ hv

I I

A*+ D

____------ G+D

2-E------- CD

Intermolecular distance

FIG. 1. Energetics for a weak charge-transfer complex between a donor(D) and an acceptor (A). W,V represents the energy corresponding to the ground-state wave function p,v. and W, represents the energy corresponding to the excited-state wave function VE. W, and W, represent these energies if there were no resonance stabilization. The energy of the optical intermolecular charge-transfer transition is marked

hv. (After R S Mulliken.‘x’) . .

imately to the dipole moments p,, and p, of the no-bond and dative structures respectively. Thus:

PE.V= u*bh-p,,). (4)

It has been argued that as the resonance interaction increases (i.e. an increase in b) so pE9 and consequently f should increase. Some con- sequences of this argument are discussed in Section 3.5.

In this valence bond description it is clear that charge-transfer forces are not exclusive. The overall binding in the ground-state of the complex includes in the no-bond structure other non-covalent forms of binding. Although the principle of the charge-transfer model accounts for the electronic absorption spectrum of the complex, it does not prove that there is a large contribution from charge-transfer forces to the binding in the ground-state of the complex. lndeed it is not necessary for a donor- acceptor pair to have any binding energy in the ground state for there to be a charge-transfer absorption band, if one accepts the principle of contact charge-transfer.@“)

;y


In systems where the dative structure makes a significant contribution to the ground-state, the more effective the two molecules are in their roles as electron donor and electron acceptor, the more stable the complex should be. This effectiveness will be dependent on several factors including an absence of those steric effects which would prevent the application of the overlap and orientation principle. General steric effects and varied contributions by other forces to the stability of the ground-state of the complex can lead to orders for the association constants and enthalpies of formation, which do not follow the orders of donor and of acceptor strength of the two components.

--------- A+e+D+

WE*

I I I I IO I 5 20 25 3.0 co

Intermolecular distance

FIG. 2. Energetics for a strong interaction between an electron donor and an electron acceptor, for example BF:, + NMe,. The symbols have the same significance as in

Fig. 1. (After R. S. Mulliken.@))

Recently, molecular orbital treatments of the problem have yielded similar results.C21,27-31’ M OS used is a Hiickel molecular orbital treatment t to obtain ionization potentials and/or values of hv and to correlate these with experimental values. Very recently values for AH for complex formation have been obtained by more exact treatments.(32-34)


2. NUCLEAR MAGNETIC RESONANCE MEASUREMENTS OF FAST EQUILIBRIA IN THE LIQUID PHASE

2.1. Introduction

The appearance of the resonance of a magnetic nucleus which can exist in more than one chemical environment depends on the life-times in the different environments.(35-42) If the life-time is long compared with the transition time between the magnetic states then separate sharp lines with intensities proportional to the populations in the different environments are observed. Wlen the life-times are short, however, a single sharp line at a position corresponding to the population-weighted average of the nucleus in the various environments is obtained. In the general case where the magnetic nucleus may exist in several chemical environments I, II, III . . . , the absorption will appear at a position SH where:

where P,, Pii, PIII . . . are the fractional populations of the nucleus in the environments I, II, III, . . ., and 6i, 6ii, SIri. . . are the chemical shifts of the measured nucleus in these respective environments.

When the transition time between the magnetic states and the mean life-times in the chemical states are comparable, then a single broad line is observed.

For the case of a nucleus which can exist in two chemical environments I and II, and for which the life-times in the two environments are equal, the coalescence of the two lines occurs when the life-time in either environment is equal to V2(7~8~,i~H)-~ where 81,11H is the frequency difference in c/s between the two absorptions if each appeared singly(37*42) (see Fig. 3).

Since the rate of exchange of the magnetic nucleus between the two environments is temperature dependent, the signal shape of the coalesced line will also be a function of temperature. Very successful investigations of fast equilibria, e.g. proton-transfer, electron-transfer, restricted rotation, ring inversion and cation hydration, have been made by nuclear magnetic resonance measurements of the changes in signal shape with temperature. The subject has been well reviewed.(43-47)

In nearly all the weak interactions considered presently, the rates of exhange between different environments, at the experimental temperatures, fall within the category of very fast reactions. Consequently the absorption will, for singlet absorptions, appear as a single sharp line. Thus factors which affect the position of the chemical equilibrium,


FIG. 3. Changes in the shape of the resonance for an increasing exchange rate between two positions, A and B, with equal populations. (After H. S. Gutowsky

and C. H. Holm.‘37))

such as changes in temperature and concentration, will alter the observed line position of the nucleus concerned. Nevertheless, at sufficiently low temperatures or in systems where the binding approaches that of a normal covalent bond (for example in the interaction of n-type electron donors with vacant-orbital electron acceptors) chemical exchange rates comparable with the rates of exchange between magnetic states of the nucleus may obtain (see Section 2.4). In such cases line broadening or separate lines will be observed.

2.2. Intermolecular Interactions Involving Hydrogen Bonding

2.2.1. Hydrogen bonds formedfrom O-H

The temperature and concentration dependence of the NMR line position of hydroxyl protons was noted by Arnold and Packard.(4s*49) Liddel and Ramseyc5”) suggested that the temperature dependence(4x) could be explained by self-association. Cohen and Reidc51) investigated the dilution curves (NMR line position against concentration of solute) for various alcohols, phenols and water in various solvents and suggested


that they were consistent with the formation of trimers and tetramers, although a quantitative analysis was not given. Two different quantitative interpretations of such curves have been made. Shoolery and his co- workers(5”,53) have measured the dilution curves of phenols and ethanol in carbon tetrachloride solution. At low concentrations, the limiting slope (&/dx),=, is equal to 2K,A, =(&imer- &,,onomer) and KN = Ndrmer/ (N ltlO”OmW )’ is the formation constant for dimerization in mole-fraction units. In general a non-zero slope for the dilution curve at infinite dilution is indicative of the presence of dimeric species. Rao and his co-workers(54) subsequently used this method to study the self-association of various alcohols in carbon tetrachloride. Similarly Somers and Gutowsky(55j have investigated the self-associations of phenols. Qualitative studies of the latter systems have been made by Bystrov et LZ~.(“~-~~) There is, however, a serious difficulty in this type of analysis: only the product KG,(or KA,) can be obtained directly. Any evaluation of K depends on an estimated value of a0 (or A,) which cannot be determined directly.

An alternative treatment of the dilution curve has been described by Saunders and Hyne. (WACO) They simplified the problem of the quantitative estimation of n-mer formation by assuming that only one n-mer species is significant. The observed frequency (v) will be the weighted average of the monomer and n-mer populations:

and v = (v,M + nv,K,Mn) / (M + nK,M”) (6)

C = M + nK,M”, (7)

where v1 and V, are the frequencies of the monomer and n-met-, K, is the association constant for the formation of the n-mer from the monomer, M is the monomer concentration at equilibrium, and C is the total, free and complexed, concentration of the solute as monomer. For arbitrary values of M,, both v and C may be obtained by introducing the appropriate constants, which may then, as theoretical curves, be used to match the experimental data. This includes the determination of IZ. The method works for any value of n. Through an application of Buckingham’s II theorem ~1 it can be shown that any equation of this type which contains the variables V, v,. v,, C and K must be of such a form that C and K may be replaced by the dimensionless term CK(lpn)ln. Thus, if a theoretical curve of frequency vs. concentration is constructed for a specific value of K, the curve for any other value of K may be obtained by multiplying the C-scale by an appropriate factor. This variation is further facilitated by using a logarithmic scale for C, since the multiplication may then be achieved merely by sliding the log C scale.


Extrapolation of the curve to low values of log C tends to a constant value of frequency which may be equated to the monomer frequency (v,). The central near-linear portion of the curve (see Fig. 4) is proportional to (I+ -v,). Thus Y, may be determined and the best fit for any value of II is taken to be the curve which with the optimum value of K for this n-mer provides the most acceptable set of values of K, v, and II. The application of the method for t-butanol in carbon tetrachloride is shown in Fig. 4.

TETRAMER - - - - -

TRIMER -

DIMER -.- - -

OBSERVED l

- ‘\

\ \

75

FIG. 4. Theoretical curves for monomer-dimer, monomer-trimer and monomer- tetramer equilibria which best fit the observed data for the self-association of t-

butanol at 40 MC/S in carbon tetrachloride. (M. Saunders and J. B. Hyne.‘9

The application of this method in the case of t-butanol association has been criticized by BeckeP) who points out that the data of Saunders and Hyne gives a non-zero limiting slope at infinite dilution which indi-


cates dimer formation. In the particular case cited this would correspond to 10-l 5% dimerization. In their turn, Saunders and Hyne@O) have treated the data of Becker, Liddel and Shoolery(53) for the self-association of ethanol in carbon tetrachloride and have shown that there is no pre- ponderance of a dimer over the rest of the concentration range, although there is no clear distinction between trimeric and tetrameric species. More recently, Feeney and Walked 63) have applied the method of Saunders and Hyne to the system methyl cellosolve in carbon tetrachloride. They found that a monomer-dimer equilibrium fitted the dilution’ curve at low and intermediate concentrations, but departures from the theoretical curve at higher concentrations indicated the formation of higher polymeric species. From measurements of the chemical shift of the proton involved in the hydrogen bond at various temperatures and their extrapolation to low temperatures Feeney and Walker@3) were able to obtain an independent estimate of A0 (the chemical shift difference between the monomer and n-mer). This was based on the supposition that at low temperature the solute is entirely in the form of the n-mer. In this way they overcame the difficulty that the near-linear central portion of the dilution curve is not very sensitive to the value of ho. Reinvestigation of the original systems of Saunders and Hyne(5s,60) by this method confirmed the original analyses and in the case of ethanol resolved the nature of the polymeric species as the tetramer.

Martin(64-66) has investigated the effect of solvent on the dimerization equilibrium and has interpreted (65) the temperature and concentration dependence of the hydroxyl proton chemical shift of 3-ethylpentan-3-01 in carbon tetrachloride in terms of the formation of both dimer and tetramer species.

Because of the complication of the presence of more than a single polymeric species, it is unreasonable to compare directly results obtained from very dilute solutions with results obtained in higher concentration ranges where a different n-mer may predominate. However, the remarkable fitting of experimental data obtained by Saunders and Hyne(5s,60) and Feeney and Walker, (63) based on the principle of a single monomer- n-mer equilibrium, is indicative of the strong predominance of a particular n-mer species. The hydrogen bonding of hydroxyl compounds to other species as opposed to self-association has been investigated by several workers.(67-76) The dilution curves of the alcohol in a complexing solvent have normally been fitted by assuming that a 1 : 1 association occurs. Corrections for self-association have not been made in all cases. Lussan et aE.(6g) studied the interactions of t-butanol with acetone in carbon tetrachloride and of t-butanol with dioxan in carbon tetrachloride. They did allow for self-association of the butanol and found that the association


with the second species was somewhat concentration dependent if a 1 : 1 association is assumed. Takahashi and Li’72,73’ examined the hydrogen bonding of alcohols and phenols to acetamides in carbon tetrachloride solution, and obtained values for the formation constant of an assumed 1 : 1 association together with the enthalpy of formation. kin et u/.(~“)

suggested that a trimeric self-association of water occurred in 1,2- dichloroethane compared with a 1 : 1 association between water and acetone in dichloroethane solution. However, Takahashi and Li(74) interpreted the concentration dependence of the chemical shift of the hydrogen bonding proton in the systems water + acetone and tetrahydrofuran + dimethylacetamide in terms of a termolecular complex involving one molecule of water and two molecules of the electron donor.

A qualitative study of the association of methanol with pyridine has been made using both ‘H and 14N nuclear magnetic resonance measurements.(76)

2.2.2. Hydrogen bonds formedfrom N-H Dilution of aliphatic amines with an inert solvent causes a shift of

the N-H proton resonance to higher fields. This has been interpreted (77-83) in terms of the breaking of N-H...N hydrogen bonds. The treatment generally used for the quantitative interpretation is that already described for O-H hydrogen bonded systems by Saunders and Hyne (WOO) where a dilution curve is fitted on the assumption that there is one predominant n-mer present so that the equilibrium may be described by a single formation constant K, = [A,]/[A]“, or K.v = NA,/(NA)ll. By this method successful curve fitting can give valuable information on the size of favoured aggregates. Feeney and Sutcliffeo7) measured the dilution curves of a variety of amines in carbon tetrachloride (Fig. 5). The curves were extrapolated to infinite dilution to obtain the chemical shift for the monomeric species. These valueso@ were used to fit the dilution curves of diethylamine and of ethylamine in carbon tetrachloride, using the technique of Saunders and Hyne. (5y,60) An excellent fit was obtained on the assumption of a monomer-tetramer equilibrium. The values obtained at 25” were K,. = 2.5 X lo-“Me3 for diethylamine; and K,. = 2.5 X IO-“M-” for ethylamine. From measurements at -37”, a value of AH” = -6.8 kcal moleP’ was obtained for the diethylamine association. This corresponds to an enthalpy change of -1 a7 kcal per hydrogen bond (based on two temperatures).

The same method was used by Crook and Schug(7”) in an investigation of the self-association of hydrazine and of substituted hydrazines in cyclohexane solution. They found that the data were best fitted on the assumption of monomer-dimer equilibria, with the constants Ks = l-8


2C

15

2 ‘5-

IC

r:

,-

,-

,-

FIG. 5. The variation of the observed chemical shift differencefwith the logarithm of the concentration C of diethylamine in carbon tetrachloride, on the basis of a monomer-tetramer equilibrium at (A) 25” and (B) -37°C. The full lines have been calculated by taking the values of K to be 2.5 X 1O-4 M-~ and 2.5 X IO-” M-~. respec-

tively. (J. Feeney and L. H. Sutcliffe.‘78’).

(N,N-dimethylhydrazine), K,%, = 1 .O (N,N’-dimethylhydrazine), K.h' = 0.8 (N,N,N’-trimethylhydrazine). The dimerization of these bi-functional compounds is in harmony with the formation of tetramers by the mono- amines. Bystrov and Lezina cso) found that in carbon tetrachloride piper- idine, hexamethylenimine, methylamine and dimethylamine each formed tetramers in agreement with the observations of Feeney and Sutcliffe$77) but that ethyleneimine, trimethylamine and pyrrolidine formed trimeric aggregates. Bystrov and Lezina further noted that for these solutions the gradient of the dilution curve tended to zero as the concentration of the solute approached zero. This suggests that dimeric species are absent. However, in chloroform solution finite limiting slopes were observed, which suggests that C l,CH...Nc hydrogen bonds are formed. From


these slopes the degree of association of each of the systems was calcu- latedi5*) on the assumption that a 1 : 1 complex was formed.

This method of limiting slopes’52) has been used by Happe’*l) to interpret the dilution curve of pyrrole in cyclohexane in terms of a monomer-dimer equilibrium. However, deviations from the calculated curve occur at concentrations greater than 0.08 mole-fraction which suggests that higher polymeric species are present. The dilution curve for pyrrole in pyridine was treated in terms of a pyrrole-pyridine 1: 1 adduct. Springer and Meek@*) have used the method of Saunders and Hyne(5gT60) to obtain the constant for tetramerization of diethylamine in cyclohexane, namely K, = 1.75 x lop2 MW3, and claim good agreement with Feeney and Sutcliffe.‘78’

LaPlanche et LID. have given a general treatment for the cases of (a) an associating solute in a non-associating solvent, and (b) an associating solute in an associating solvent. They have applied the method to a study of the self-association of N-mono-substituted amides in carbon tetrachloride and in the associating solvents: dioxan, diethyl ketone and dimethylsulphoxide.

2.2.3. Hydrogen bonds formedfrom S-H

From the small shifts observed on dilution of benzene solutions of thiophenol (benzenethiol) and thiobenzoic acid, Rao and his co- workers(84*85) concluded that there was negligible self-association of the solute. More recently Marcus and Miller,‘86’ using essentially the method of Saunders and Hyne,(5s,60) have determined that the association of this phenol in carbon tetrachloride is primarily due to formation of the dimer. Because the degree of association was small in all cases, the point of inflection of the dilution curve occurred at relatively high concentrations, consequently an iterative procedure was used to evaluate K. The hydrogen bonding of thiols to other molecules has been investigated.‘87-s1) The hydrogen bonding of thiophenol to various donor molecules has been studied(87,88) by measuring the S-H proton chemical shift as a function of the excess donor concentration at a fixed concentration of thiophenol. Under these conditions the following relationship obtains:

1 .l 1 1 1 -=-=- - V-V~ A KcA;[Dl,+S, (8)

where v is the observed frequency of the thiol proton, i.e. the weighted average of the characteristic frequencies for this proton in the free thiophenol (v& and in the complexed thiophenol (vc), K, is the association constant for the 1: 1 complex, A0 = ( vc - vf) and [D10 is the total, i.e. free and complexed, concentration of the electron donor. From equation (8)


a plot of l/A vs. l/[D], yields a straight line from which A,, and K, may be obtained (Fig. 6). See also Section 3.2 for a more detailed discussion.

25 I I -1

N

0 X

I s- - , >

0 ’ I J

0 2 4

FIG. 6. Plots of l/(u-I+) (= l/A, see equation (8)) against l/[D], (where [D], is the total concentration of donor) for a series of benzenethiol-dimethylformamide solu-

tions in carbon tetrachloride at: 6, -12”; l ,8”; 0,26”; @, 48”; 0,68”. R. Mathur, E.D. Becker, R. B. Bradley and N. C. Li.@‘))

From plots of log K, vs. T-l, values for the enthalpy of formation of the complexes were obtained. A satisfactory feature of this analysis is that the values of A,, are temperature independent. The results are summarized in Table 1.

TABLE 1. THERMODYNAMIC CONSTANTS FOR HYDROGEN BOND FORMATION

BETWEEN BENZENETHIOL AND VARIOUS HYDROGEN ACCEPTORS IN CARBON TETRACHLORIDE AND RELATED CHEMICAL SHIFT DATA (R. Mathur, E. D. Becker, R. B. Bradley and N. C. Li ;‘*” R. Mathur. S. M. Wang and N. C. LicT8))

Hydrogen acceptor

Pyridine

~~

a At 60 MC/S. “At 22”.


Rousselot(sY) has studied the dilution curves for thiols in chloroform. In agreement with previous work cyO) he concludes that a 1 : 1 association R-S . . . . H-CC& occurs. The temperature variation of the equilibrium

has been studied.(“l)

2.2.4. Hydrogen bonding through acetylenic protons A number of workers(Y2-103) have demonstrated the sensitivity of the

chemical shift of acetylenic protons to the nature of the solvent and to concentration. Whipple et al. cy2) found a correlation between the chemical shift extrapolated to infinite dilution and the infrared C=C stretching frequency for propargyl bromide in a series of solvents. Richards and Hatton(y3) measured the chemical shift at infinite dilution for phenylacetylene and propargyl chloride in a series of solvents. The values obtained were explained in terms of solute self-association and solute- solvent association. Petrov and his co-workers(y7-100) investigated qualitatively the solvent effect for various acetylenes including vinyl- acetylene and diacetylene. Kreevoy et al. (‘(‘l) attempted a quantitative interpretation of the system phenylacetylene + pyridine in carbon tetrachloride by using a value of K from infrared data in order to calculate the chemical shift for the pure complex in solution. The value so obtained was constant in solutions containing up to 20% solute. However, they admit that any value of K, up to unity would give such a correlation. A direct treatment of nuclear magnetic resonance data has been given by Brand rt a/.(‘02) They investigated the interaction of benzoylacetylene with each of benzene, ioluene, xylene, mesitylene and di-n-butyl ether, in cyc!ohexane as a diluting solvent. The following relationship was used:

N,A-’ = N,jA,,-‘+ K,,,A,,-’ (9)

where N, is the total (free and complexed) mole-fraction of the donor solvent (D), K, is the mole-fraction association constant for a 1 : 1 complex between D and the benzoylacetylene, A is the observed shift of the ethynyl proton and A, is the shift due to the complex, both these shifts being measured relative to the shift of “free” ethynyl resonance as measured in pure cyclohexane (see Section 3.2.). Corrections for the anisotropy of aromatic solvents were made. From plots of NJA against ND, straight lines were obtained from which K, and A, were evaluated. The constants so obtained are given in Table 2.

Tyrrell(“‘“) has investigated the association of benzene with propargyl chloride in cyclohexane using the same technique. He obtained very similar values for K., from measurements of the shift of both the ethynyl


TABLE 2. THERMODYNAMIC CONSTANTS FOR HYDROGEN BOND

FORMKTION BETWEEN BENZOYLACETYLENE AND VARIOUS HYDROGEN ACCEPTORS IN CYCLOHEXANE AND RELATED

CHEMICAL SHIFT DATA (J. C. D. Brand, G. Eglinton and J. Tyrrell”02’)

Hydrogen acceptor 1 “v(‘~“) 1 (p$m) j (kcalt2k’)

Benzene

Toluene n-Butyl ether

and the methylene protons: (a) from the methylene protons: KN = I *O * O-02, A,, = 63$& 1.3 c/s (60 MC/S); (b) from the ethynyl proton: Ks = 0.96+-O-06, A,, = 37.0~2.4 c/s (60 MC/S). The fact that the methylene protons have a higher A0 value than the ethynyl proton led to the conclusion that the methylene protons are closer to the benzene ring.

2.2.5. Hydrogen bonds formed between $C-H protons and aromatic 7i--systems

Several workers(104-8) have studied the chemical shifts of protons attached to a saturated carbon atom involved in hydrogen bonding to aromatic r-systems. Reeves and Schneidetilo4) measured the proton chemical shift of chloroform as a function of concentration in solutions in benzene and in various substituted benzenes and obtained values for the shifts at infinite dilution. They concluded that these shifts were “indicative of specific complex formation”. Support for this conclusion was obtained from phase diagrams which showed definite eutectics. In a subsequent paper SchneidePos’ examined the proton shifts of chloroform and of acetonitrile in benzene solution. The behaviour of these systems was attributed to the formation of complexes which have the structures I and II respectively.

Abraham(“‘+‘) studied the effects of temperature on solutions of cyclohexane, methyl iodide and iodoform in toluene solution over the range -60” to +lOO”. He found that whereas the chemical shift for cyclohexane


was temperature independent, the shifts for methyl iodide solutions and iodoform solutions showed a significant variation.

The chloroform-benzene interaction in cyclohexane has been studied by Creswell and Allred.~los~ They took a fixed, low concentration of chloroform and various excess quantities of benzene. The measured shift (6) is related (see Section 3.2.) to the shift of the complex aAD and the shift of the chloroform aA by equation (IO):

+h!$Q (&4D-K4) --*, h

(10)

where NAD is the mole-fraction of the 1 : 1 complex and NA is the mole- fraction of the total (free and complexed) chloroform. Arbitrary values of KN enable hypothetical ratios N,,/N, to be calculated; only the correct value of K, will give ratios N,,/N, which vary linearly with 6. By this technique and by studying the temperature variation of the shift, the following values were obtained: K,V(2.5”) = I.06 & 0.3, AHN = I.97 2 0.35 kcal mole-‘, AS,V(250) = -6.5 5 O-5 cal mole-’ deg-l and A0 = 1 a91 & 0.4 ppm. The effects of corrections for anisotropy were considered.

2.2.6. Hydrogen bonds formed berrveen>C-H and n-electron donors

The interactions of >C-H protons with a variety of n-electron donor

molecules have been mvestigated.(110-27) Huggins, et al.(‘lO) obtained dilution curves for chloroform in acetone and in triethylamine. From the values at infinite dilution, A, was obtained by assuming a value for K. These were then used to try to fit the whole of the dilution curve. Only the correct value of K will give a complete fit. Creswell and Allred,~los~ using the same method of analysis as for the system chloroform-benzene,uos) have studied the association of chloroform with triethylamine, and of various trihalomethanes with tetrahydrofuratP8’ (Table 3). McClellan et al.(120r applied the method of Huggins et al.“lO’ to the system chloroform-dimethylsulphoxide, but found that the results could not be fitted by a simple 1 : 1 equilibrium. Self-association of the chloroform and termolecular complexes involving two molecules of chloroform to one of dimethylsulphoxide had to be invoked to provide an adequate explanation of the observations. Howard et al. (121) have made similar studies of chloroform complexes with diethyl ether, acetonitrile, acetone, triethylamine and pyridine. Takahashi and LI ‘(w have used the method of Mathur et al.(87,88) (see equation (8)) to investigate the association between chloroform and N-substituted amides. This was in order to generalize their treatment of


TABLE 3. THERMODYNAMIC CONSTANTS FOR HYDROGEN BOND FORMATION BETWEEN THE HALOFORMS AND TETRAHYDROFURAN IN CYCLOHEXANE AND

RELATED CHEMICAL SHIFT DATA (C. J. Creswell and A. L. Allred”‘8))

CHFB CHCl, CHBr, CHI,

AH, kcal mole-’ -24kO.2 -3.6kO.4 -2.6-tO.2 - 1.6kO.4

AG,v kcal mole-’ (25”) -0.90 -1.04 -0.93 -0.91

ASs cal mole-’ deg-’ (25 ) av.A forXCH. .OCH

-5.2’-cO.6 -8.6r’rO.4 -5.6-cO.4 -2.3kO.6

ppm’ 3 . 4x, -0.69kO.03 -0.85 eO.03 -0.83 50.02 - 0.43 k0.03

*CHX3 dilute,” ppm -4.800 -5.653 -5.257 -3.383

“Chemical shift with C,H,, as internal reference.

amine-amide interactions when an “active” diluting solvent (namely chloroform) was used. Equation (8) was utilized by Berkeley and Hanna(122-3) to investigate the associations between chloroform and the donors: pyridine, acetonitrile, ethylidene-isopropylimine and N-methyl- pyrrolidine.

An interesting treatment of chemical shift data was devised by Alley and ScotP7) in their investigation of the association of 1 -hydropentadeca- fluoroheptane with acetone and with amines. This involves re-expressing the shift values as fractions of the difference (a,,,, so,ute - a,,,, solvent) so that when this function is plotted against mole-fraction of solute, both axes range from zero to unity. The resulting curves can be fitted by the correct choice of KN alone as shown in Fig. 7. The value of A,, can be obtained from the shift at infinite dilution (A,) in conjunction with K, viz.

A, = A,$,( 1+ KN) -’ (11)

or if the shifts are measured with respect to the signal of a non-partici- pating species:

(1 la)

Strong support for this analysis is provided by the observation that the A0 values so obtained are temperature independent.

Contact shifts of chloroform proton resonance when the chloroform is acting as a solvent in the presence of Co(fI) or Ni(II) have been inter- pretedoz7’ in terms of hydrogen bonding to the complex ions. The shifts show a linear temperature dependence and this has been taken as evidence for an integral coordination number.


.6-

0 .I .2 .3 4 .5 .6 .7 .8 .9 I.

MOLE-FRACTION ACETONE

FIG. 7. Experimental dilution curve for C,F,,H in acetone----. together with calculated curves for KV values of 2.6 and 1.4-----. at 30”. This indicates a best fit for

K, = 2 * 0.6. (S. K. Alley, Jr. and R. L. Scott.““‘)

2.3. Carboxylic Acid Associations

Hydrogen bonding in carboxylic acids has been investigated by a number of workers’52~128-44) using nuclear magnetic resonance spectroscopy. Huggins rt al. (52) found no shift in the hydroxyl proton signal of acetic acid when it was diluted in carbon tetrachloride. They suggested that there was no appreciable dissociation of the dimeric acid units down to the limit of the chemical shift measurements (0.2 M). They observed a large shift, however, when acetone was used as solvent, which they interpreted in terms of competition of the acetone as a hydrogen-bonding base, causing the dissociation of the acid dimers at higher concentrations than is the case when carbon tetrachloride is the solvent. Reeves and Schneider(l’s) investigated the dilution curves of acetic acid in a series of “non-interacting” solvents and found that on dilution there was first a shift to low field, but on further dilution the shift was to higher fields compared with pure acetic acid. They explained these results in terms of a


polymer-dimer equilibrium at high acid concentrations and a monomer- dimer equilibrium in dilute solutions. They also measured the variation of the hydroxyl proton signal in pure acetic acid with temperature. The results were consistent with the formation of a very stable dimer. When the work was extended to other carboxylic acids!12”)it was suggested that, whilst acetic acid formed polymeric hydrogen bonded associations in concentrated solutions, the longer chain acids showed only dimer associations in the same concentration regions. Later work(130) gave evidence for the formation of weak acid-solvent complexes in dichloromethane and 1,2-dichloroethane. Similar studies have been made by other workers.(131-6) More recently, however, Muller and Rose(‘37’ have pointed out that experimental dilution curves measured by other workers(52*128.136) do not agree with each other, and that interpretations involving a large degree of dimerization at low or moderate concentrations are in disagreement with data from other sources. (145) They suggest that serious errors may be introduced by the presence of a small aqueous impurity in the solvent, and they have shown that this would lead to large spurious shifts (Fig. 8). They have investigated the dilution of acetic acid in acetic

-5oo- I , , I I

0 0.05 t I1 I I I1 I I I I I

0.10 0.15 Mole-Fraction AcOH

FIG. 8. Chemical shift (in c/s from acetic anhydride at 56.4 MC/S) of the hydroxyl proton as a function of acetic acid concentration: lowest curve shows the results of Muller and Rose”37’ for acetic acid in acetic anhydride; middle curve, calculated shifts for acetic acid in “hypothetical wet solvent”; upper curve shows the results of

Reeves”*“‘for acetic acid in acetone. (N. Muller and P. I. Rose.(13’))


chemical shift for boron trichloride + diethyl ether in dichloromethane as a diluting solvent. For other related systems curves were obtained. For series of solutions containing a constant concentration of the ether (- 3 mole%) and varying concentrations of the boron trihalide (up to 60 mole%) the plots of chemical shift vs. concentration of boron trihalide were used to obtain the association constants (K,) by a curve-fitting procedure (Fig. 9). Their results are given in Table 4. It was found possible to explain the results only in terms of a 1 : 1 complex, which suggests that in the systems studied, under the experimental conditions used, no significant amounts of the complex A,D are formed. The large value of K for the one adduct (BCI:;Et,O) which did not show significant deviations from linearity on a Craig-Richards type plot”““’ provides some vindication for Gore and Danyluk’s conclusions. It also sets a limit to equilibria which may be treated in this way. The differences in the chemical shifts and association constants for BF:,$CH,),O and BF,,.(C,H,),O between those reported by Craig and Richards(14”) and those reported by Gore and Danyluk(‘“‘) may be attributed (lJ”) to the effect of the solvent in the

30 t

20 I I 1 _1___1__ 0 10 20 30 40 50 60

CONCEflTRATlON (mole % 6X,)

DIETHYL ETHER COMPLEXEG

FIG. 9. Variation of the chemical shift of the methylene protons of diethyl ether against BX,, cont. for diethylether-boron halide in dichloromethane so!ution at

23”: 0. e. measured shifts: -. calculwted shifts. (E. Gore and S. S. Danyluk.“““)

ORGANiC CHARGE-TRANSFER COMPLEXES 27

TABLF. 4. EQUILIBRIUM CONSTANTS AND CHEMICAL SHIFTS FOR ETHER-B• K~N

TRIHALIDE COMPLEXES IN DICHLORO-

ME-~HANE AT 93” (E. Gore and S. S. ,&ny,u~“““‘)

--- T_---_

Complex

“Relative to internal dichloromethane. at 60 MC/S.

latter case. This may include hydrogen bonding between the solvent and the ether.

Finch and his co-workers(‘““) have found qualitative evidence from the magnetic resonance absorption of “B in solutions of boron tribromide in benzene for a solvent-solute association in this system.

2.5. lntrraction ofAmides and of” Other Carbony Compounds with

Aromatic Compounds

By contrast with the very strong interaction between boron trihalides and n-electron donors, the interaction of amides and ketones with aromatic compounds appears to be very weak. Nevertheless, both may be circumscribed by Mulliken’s (x) charge-transfer valence bond description.? For systems which are the subject of this section, it has been clearly demonstrated by several groups of workers that specific interactions occur and that for non-conjugated compounds the structure of the complex probably involves dipolar forms of the amide, or carbonyl function which interacts with the r-electron system of the aromatic donor component.

2.5.1. Amide-uromatic compound intrrrrctions

Hatton and Richards”‘“+“) observed that the dilution of dimethylformamide with benzene caused a large upfield shift of the methyl resonances. They also noticed that the signal due to the methyl group trans to the carbonyl function shifted much more than the signal due to the cis-methyl group. These observations have been substantiated by later workers.(“)5.‘56) The shifts have been interpreted in terms of

Wnce this chapter was written, arguments against a charge-transfer description for such interactions have been given, e.g.: J. RONAYNE and D. H. WIILIAMS, J. C/ICJO~. SM. (B). 540 (I 967).


specific solute-solvent interaction involving a structure III in which the association is between the positive charge on the nitrogen, in the dipolar structure of the amide, and the r-electrons in benzene. The oxygen atom containing the negative end of the amide dipole is orientated as far away from the benzene ring as possible. Further evidence for a spec$c

_ cH3\;=C/o CH,(; ‘H

I ,

solvent-solute interaction includes (a) the large temperature coefficients of the observed chemical shifts,(156-8) and (b) the fact that the temperature at which intramolecular free rotation of the amide occurs, together with its associated activation energy for rotation, is solvenP7) and concentration(158) dependent. Support for the specific structure III proposed by Hatton and Richards(154) comes from the observed shifts for solutions of N-methyl-N-cyclohexylacetamide~*5g~ and of methylformamide(160) in benzene solution. In the latter case, a much larger shift is shown by the methyl group when in the trans position than when in the cis position. This is explicable if the complexes have the structures IV and V respectively, where in both cases the molecules are arranged with the oxygen atom as far away from the ring as possible, as was suggested for the case of dimethylformamide by Hatton and Richards.u5”)

In their original paper, Hatton and Richards(154) extrapolated the measured chemical shifts to a value for an infinitely dilute solution (6,). They found that 6, for dimethylformamide in various methylbenzenes as solvents decreased as the number of methyl groups in the solvent species was increased. This observation was explained in terms of steric interference by the aromatic methyl groups, thus reducing the intermolecular interaction. However, as has been pointed out (equation (1 I)), 6 (or A) is a function of both the association constant (K) and the shift for the pure complex (a,,) (or AJ. For this reason the association of


dimethylformamide with methylbenzenes in a diluting solvent was studied by Sandoval and Hanna.(lG1) Hanna and Berkeleyuz2) and Takahashi and Li(72,73) had previously applied this method to hydrogen-bonded systems. Sandoval and Hanna found that K increased as the degree of methylation of the aromatic electron-donor was increased. It is the larger rate of decrease in 6, which gives rise to the decrease in 6, with increased methylation of the aromatic electron-donor observed by Hatton and Richards.(154) The question as to why a,, should decrease with increased methylation of the donor still remains. It could be the result of steric effects as suggested by Hatton and Richards.(154’

From dilution studies of N-methyllactams in benzene Moriarty and Kliegman(162) have proposed that specific solvent-solute interaction occurs in these systems. The structures which they propose, namely VI and VII, are in accordance with the observed shifts and are completely analogous to those proposed by Hatton and Richards(154) for the dimethylformamide-benzene complex.

The solvent effect of pyridine on these amides is not the same as that of benzene and its homo10gues,(155,159*162) and it has been suggested(15s) that the structure of the N-methyl-N-cyclohexylformamide-pyridine complex has the structure VI I I.

--CH/\O- -/\C,, _ 2

CH,, N-C”

o-

2.5.2. Carbonyl compound-aromatic compound interactions The interactions of the carbonyl function with aromatic compounds

appear to be similar in principle to those involving amides which have been described in the preceding paragraphs. However, whereas in the amides the bonding is through the nitrogen atom and the carbonyl group is removed from the centre of the interaction, in the case of many esters and ketones, interaction appears to be through the carbon atom of the

30 ROY FOSTER 4ND COLIN A. FYFE

carhonyl group. These latter systems have been the subject of qualitative rather than quantitative investigation. This stems,from the fact that the

shifts involved were to a large extent originally used in order to simplify nuclear magnetic resonance spectra of carbonyl-containing solutes by comparison of their spectra in aromatic and non-aromatic solvents. Differences in these spectra have enabled different t.~pvprs of protons in carbonyl-containing compounds to be identified. Nevertheless, work. particularly by Bhacca. Williams and their co-workers”‘;:‘+7i’ and by Connolly and McCrindle. “i2-7n) has demonstrated clearly that a specific interaction with the aromatic donor molecule occurs. Some details of the geometry of these complexes have been deduced.

In an investigation of the nuclear magnetic resonance spectra of some steroidal ketones and acetates, Bhacca and Williams(lf’:S) observed diifer-

ences in the chemical shifts of some of the protons of the carbonyl compounds when dissolved in benzene, compared with the shifts when the solvent was deuterated chloroform. They concluded that a solute- solvent complex is formed, in which the n-electrons of the benzene ring interact with the partial positive charge on the carbonyl carbon atom so that they are as far away as possible from the partial negative charge on the oxygen. The structure was formulated schematical!y as IX. Their conclusions were confirmed when the observations were extended to a large number of cyclic ketones with rigid structures. It was found”““)

that the relative solvent shifts (&.,,(.,:, - SC.l;E,e) were approximately constunt

for protons in N purticulur position relative to the carbonyl group. A reference plane for predicting the solvent shifts was proposed by Williams and Bhacca,“““-“I and independently by Connolly and McCrindle”7’P:S’ (see structure X). These solvent shifts have been used to determine the conformation of 2,2,6,6-tetramethylcyclohexanone””7’ and to reassign the methyl group signals in camphor. (liZ) Considerations of a number of

proton resonances in 1 I-oxo-5cu-steroids when dissolved in various solvents including benzene have enabled the geometry of some of the complexes with benzene to be determined. (IfiS) Thus Sn-androstan- 1 1 -one

does not appear to have an approximately planar association with

ORGANIC CHARGE-TRANSFER COMPLEXES Jl

benzene, corresponding to the structures of simpler ketone-benzene complexes suggested by other workers (156*168) but rather a structure XI in which the plane of the benzene ring is almost at right angles to the overall plane of the steroid molecule. (I’jx) A further discussion of this problem has also been given by Ronayne and Williams.(‘““)

All the above treatmenis are of a qualitative nature since the shift, even in very dilute solutions in the associating solvent, is a function of both K and A,,(or a,,), and for an infinitely dilute solution A, # A,, except when K,v = 00 (it has been pointed out already (equation ( I I ) that for finite values of K,,. A, = A,K,\( 1 + K v)--‘). Recause associations of this type are re!atively weak, A will not approach the value of A,,. Consequently unless K and A,, vary in a parallel manner, comparisons of a single function of both K and A,, may not be meaningful, although in series of closely related compounds some correlation appears to obtain. It must be emphasized that such criticisms do not apply to the relative shifts of atoms in the same molecule; these will all be functions of the same K, and the measured shift values for dilute solutions will be proportional to the corresponding A,, values.

Laszlo and Williams(‘iO) have measured the relative solvent shifts ( 6czzrc,,x - i&,J for methyl and methoxy protons in substituted p-benzo- quinones and found them to be much larger than for saturated ketones. Ronayne and Williams”““) concluded that 2: 1 complexes with structures of the type XII are formed. Other authors, however, have considered such interactions to involve a 1 : 1 stoichiometry (structure XIII) in which the aromatic ring interacts with the whole of the delocalized n-system of the quinone (Ii41 (see Section 2.6). Satisfactory quantitative analyses have been obtained on the basis of such an hypothesis (see Section 3.3). If this model of the complex is correct, then much larger values of K would be expected than on the basis of structure XIII. This would explain the relatively large observed shifts.

In an attempt to relate the shift differences on a more quantitative basis, Laszlo and Williams(““J measured the shifts of steroidal ketones in toluene solution as a function of temperature. The value of the chemical


shift obtained by extrapolation to absolute zero was considered to be equal to A,,. The enthalpy change on complex formation was estimated on the assumption that a 1 : 1 complex is formed. Their value for Sa-androstan- 1 1 -one, namely AH.%, = -0.65 -+ 0.15 kcal mole-‘, although claimed only to represent the order of magnitude of AH, nevertheless does indicate the weakness of this type of interaction. However, no exact quantitative treatment of associations of carbonyl compounds with n-aromatic donor molecules has yet been made. Until such determinations have been made there can be no completely satisfactory comparisons with other systems.

Timmons(17j’ has measured the shift difference (SCHCls- CjCsHJ for a&unsaturated k e t ones, for which he has proposed reference planes as in structures XIV and XV. Such interaction is discussed in more detail

-- Deshlelding

2 ~ R-

--R

in Section 2.6 where the interaction will be considered as being derived from the delocalized r-electron system rather than localized on the carbonyl group.

The nuclear magnetic resonance spectra of axial and equatorial steroidal acids in the solvents chloroform, benzene and pyridine have been measured.(176) Cyclopropanones substituted with various electronegative groups have also been examined and correlations with (i&C13-&,Hs) have been found.“77-g’ Similar solvent shifts have been observedosO’ for y-lactones although in these cases it was concluded that it was not possible to predict the exact geometry of the complex. Grigg et a/.(181) have investigated solvent effects in coumarins, proposing the reference plane indicated in structure XV I. This is another example where the interaction may be much more delocalized. Anderson oR2) has investigated steric effects in the (benzene) solvent-induced shifts of 1,3-dioxans. He concludes that very weak 1: 1 complexes are formed with the probable geometry XVII or XVIII.

In some of the papers quoted in this section,(i6R,17”-7) comparisons of


Shielding 8

ODeshielding

solvent-induced chemical shifts in solute molecules dissolved in pyridine rather than in benzene have been made. Other examples may also be cited (183--E) In 1960 Slomp and MacKellaP*) examined the solvent shifts

G&l3 - 6 cJHaN) for thirty carbonyl-containing steroids and concluded that the pyridine preferred certain sites for coordination and that the large shifts were due to the anisotropy effect of its ring current. Johnson ef aLoB5) investigated the pyridine-induced solvent shifts in a number of model cyclohexanones, and used the correlation to deduce the stereo- chemistry of the cyclohexanone ring in cycloheximide and its stereo- chemical isomers naramycin-B, isocycloheximide and neocycloheximide. Connolly and McCrindleo86’ have investigated the effects of both pyridine and pyrrole as solvents. WilliamP7) suggested, from an investigation of pyridine shifts, the reference plane XIX for the pyridine complex, compared with X for the benzene complex. Hampel and Kraemer(lE8) have examined a large number of steroids in pyridine-d, solutions and have interpreted the shifts in terms of solvent-solute complex formation.


2.6. Charge-transferAssociutionst

This section is in two parts. The first part concerns work in which the interactions have not been specifically identified by their authors as charge-transfer complexes but which would appear to come well within this category. The second section discusses work where the interactions have been recognized as involving charge-transfer complexes.

2.6.1. Aromatic-aromatic interactions

Very large dilution shifts have been observed for some aromatic compounds when dissolved in benzene. Corio and Dailey(ls”) measured the shift at infinite dilution for various substituted benzenes in several solvents including benzene. In a series of papers(‘““-2) in which an investigation of the effects of solvent on the spectra of a large number of substances was carried out, Schneider and his co-workers found large shifts for molecules such as nitrobenzene, p-benzoquinone and I-X- substituted ethylenes in benzene, compared with solutions in neopentane and concluded that specific solute-benzene complexes existed. Structures such as XX were proposed for the complexes.(156~1”‘-3) Hatton and SchneideP”) also found that the shifts showed very large temperature coefficients, which strongly supports the idea of complex formation. Diehl,(i9”-s) however, from his examination of solvent effects of substituted benzenes, concluded that the benzene-induced shifts were due only to anisotropy and that 1 : 1 complex formation was unlikely since, amongst other reasons, he found it impossible to fit the dilution curves from the parameters obtained from the temperature behaviour; also the energies of formation would be the same for all the solutes he used.

Benzene-induced solvent shifts have also been found for maleic anhvdride (ly6) the methoxy protons of methoxy-substituted benzenes(lg7) and’substiiuted anilines.‘l”xJ From the variation of the aldehydic and aromatic protons of substituted benzaldehydes, (lug~zoo) Klinck and StotherPoo) deduced that there was a specific solute-solvent interaction in benzene solutions, and that differences in orientation occurred depending on the electron distribution, for example as in XXI and XXII. Similarly Wein- berger and his co-workers!““‘) found large shifts in the spectra of substituted benzalmalononitriles in benzene solution, compared with the absorption bands in chloroform solution. Structures of the type XXIII were suggested.

Taft et ~1.‘“~~ found that the lgF resonances of 4-fluoronitrobenzene and 4-fluoronitrosobenzene were very concentration dependent in

:Although many of the interactions described in Section 2.5 may be described in the broadest terms as charge-transfer complexes. the term is commonly restricted to interactions described in this section.


carbon tetrachloride solution. Using a computer technique they fitted the experimental dilution curves with the best values of K, and A,, (Fig. 10). Their results are summarized in Table 5.

Several workers have investigated associations which involve hetero-

10.50

10.40

10.30

10.20

10.10

E, 10.00

” 9.90 =: e 9.80

9.70

9.60

9.50

9.40

9.30

0 1.0 2.0 3.0 4.0 5.0 6.0 7.0 8.0 9.0 Molarity of p-fluoronitrobenzene

FIG. 10. Shielding parameter (/~~““‘) = ( A) for p-fluoronitrobenzene in binary

mixtures with carbon tetrachloride; 0. experimental values: 0. calculated values. The fit of concentration dependence by the assumption of dimer formation from K, = [dimer]/ [ monomer] 2 = 0.266 M-’ and lIm =-11.14ppm (=A,)). (R. W.Taft.

G. B. Klingensmith and S. Ehrenson.lLOY’)


TABLE 5. ASSOCIATION CONSTANTS (K,) AND RELATIVE CHEMICAL SHIFTS FOR MONOMER-DIMER ASSOCIATIONS IN CARBON TETRA- CHLORIDE AT 25” (R. W. Taft, G. B. Khngensmith and S. Ehrenson(*02))

~

a Shielding parameter/unit fluorobenzene derivative. b Shielding of 4-F relative to 3-F.

cyclic aromatic compounds. From differences in the shifts of the pure liquid and of solutions in benzene. Murrell and GiPo3) suggested that methyl substituted pyridines undergo self-association in the pure liquid and specific association with the solvent in benzene. They proposed such structures as XXIV and XXV. Dhingra et al.@04) made a quantitative investigation of the self-association of 2-amino-3-methylpyridine and 2-amino-6-methylpyridine. The proposed m-n-interaction in the dimerization of pyrrole(81) has already been discussed (Section 2.2).

2.6.2. Recognized charge-transfer complexes Charge-transfer complexes involving iodine have had an important

place in the historical development of charge-transfer complexes both in the experimental optical studies and in theoretical studies (see Section I). Several workers@05-‘0) have investigated the nuclear magnetic resonance spectra of these systems. The results have been somewhat disappointing, mainly due to the different characteristics of the resonances, compared with optical absorption bands. The former depend on the ratio [AD]/[D] whilst the latter depend only on [AD].

Matsuoka and his co-workers@05-6) examined the spectra of various solvents containing dissolved iodine, but found no shifts which were dependent on [I,] except in the case of mesitylene. Fratiello(207) measured the chemical shifts of solvents containing high concentrations (1 M) of iodine. No change in the shifts was observed in the case of benzene, xylenes or nitro compounds. Changes were observed, however, for


amines including various aliphatic amines, aniline and pyridine. The amino proton absorption in the case of aniline moved by 1.27 ppm whilst the ring proton absorptions showed only a slight movement. He concluded that the bonding was localized on the nitrogen atom- a conclusion drawn earlier by Tsubomura c2) for N,N-dimethylaniline on the basis of optical data. Koch and Zollinger( =O) also observed only small changes in the chemical shift when iodine was added to aromatic solvents. A somewhat larger change was observed in the case of pyridine and they have concluded that this involved an interaction located at the nitrogen atom. From the discussion given in Section 3.2, it will be seen that only very small changes in the chemical shift should be expected since for these systems the association constant will not be large and the iodine is not in excess, so that the fraction of the donor which is complexed can only be small.

Wyant and his co-workers(211) observed no difference in the spectrum of benzene when it was saturated with tetracyanoethylene. Singer and Cramc212) obtained a similar result from a study of the nuclear magnetic resonance absorption of paracyclophane in dichloromethane solution and in carbon disulphide solution in the presence and in the absence of tetracyanoethylene. Here again, the amount of tetracyanoethylene which will dissolve may not be sufficient to produce a significantly large concentration of complexed donor molecules relative to the concentration of free donor molecules; alternatively A, may be small.

Changes in the chemical shift were, however, observed when conditions obtained such that a reasonable fraction of the observed species was complexed; for example, from solutions in chloroform-d of the 1: 1 adduct: 1,3,5-trinitrobenzene-hexamethylbenzene and similarsystems.‘210) Similarly Perkampus and Kriiger’213’ observed changes in the chemical shift of p-benzoquinone in a 0.25 M solution in carbon tetrachloride when high concentrations (- 0.5 M) of various methylbenzenes or naphthalene were introduced. Neither of these two groups of workers separated out the parameters K and A0 for the complex from the observed changes in chemical shift.

Determinations of K and A0 on a recognized charge-transfer system were first made by Hanna and Ashbaugh,c214) in a study of the complexes between 7,7,8,%tetracyanoquinodimethane and methylbenzenes in dioxan solution. They applied a method they had previously used to study hydrogen-bonded systems.022’ This method is discussed in detail in Section 3.2. They noted that their results were dependent on the particular concentration scale used. More recently, they have ascribed at least a part of the observed shift to interaction with the solvent.(215) Nehman and PoPov(~~~) have applied the same method to the complex between penta-


methylenetetrazole and 1,35trinitrobenzene in carbon tetrachloride. In this case the solvent-solute interaction would be expected to be very much smaller, and the results they obtain show reasonable agreement with the association constant obtained from optical data.

Larsen and Allred measured the equilibrium between iodine and p-methoxythioanisole in carbon tetrachloride at various temperatures. They took a fixed concentration of the thioanisole and variable amounts of iodine and measured the shift of the methyl protons in the donor. K was evaluated by the method of Creswell and Allred(““‘) using equation (10) (see also Section 3.2). A vindication of their analysis is that their value of the shift in the complex relative to the shift in the free donor was found to be temperature independent. Larsen and Allred(Z18’ also investigated the interaction of iodine with 2,4,6_trimethylpyridine (2,4,6- collidine) in mixtures of nitrobenzene and carbon tetrachloride. They rationalized their results in terms of an ionic and a (non-ionic) charge- transfer interaction. Two separate sets of methyl signals were observed resulting from (a) the ionic pyridinium structure and (b) the “average” value between free pyridine and its charge-transfer complex. The association constant for the ionic product was estimated from the band areas. This constant having been evaluated, they were able to compute the association constant for the charge-transfer complex.

The association constants and chemical shifts of the pure complex have been determined for a range of n-donor r-acceptor systems by Foster, Fyfe and their co-workers. (21s--83) The method and the results which have been obtained are discussed in detail in Sections 3.2 and 3.3.

3. THE DETERMINATION OF ASSOCIATION CONSTANTS OF ORGANIC CHARGE-TRANSFER COMPLEXES IN DILUTE

SOLUTION

3.1. General: the Effects of Concentrution Sales, Solvent

Competition, Specific Solvation, Isomeric Complexes

In Section 2 the experimental determinations of some association constants were outlined. These included binary systems involving either self-association of one species or complex formation between the two species A and D (52 -3 .3 C . ,a. s5’. S) In some of these experiments the concentrations ranged from pure A to pure D. The evaluation of the association constants was then generally achieved by curve-fitting techniques. By contrast, interactions of two species, A and D, have also been studied using dilute solutions with a third species acting as the diluting solvent. Some examples have been mentioned. (72,73) A discussion of nuclear magnetic resonance determinations for such equilibria, together with a


comparison with the results obtained using other physical methods, is the subject of this section.

In the first place it will be assumed that the two species A and D interact reversibly by very fast reactions to form only a single 1 : 1 complex AD, and that the activity coeffici&ts for the species A, D and AD, namely yA, yo, yAD, are such that the quotient yan/yAyo is unity (where YA, YD and YAD are appropriate to the concentration units used). Some implications of non-ideality will be discussed in Section 3.2. For solutions ideal on a molar concentration scale, an equilibrium constant KPD may be written where:

K$D = [AD]/[A] [D]. (17)

Some workers have preferred to express the equilibrium constant as a dimensionless quantity K$=r’, either by expressing all the solute species in mole-fraction units or, more usually, by expressing the concentration of one component species, normally D, if it is present in large excess over the second component species, as a mole-fraction:

K;?” = [ADl/[Al ND = NAD/N.SNIP (18)

where NA, ND and NAD are the mole-fractions of these species respectively in the total solution mixture. In dilute solutions for systems in which the association constants are not too small (KtD > - 1 M-l), KC” and KtD are simply related by the molar volume, in litres, of the solvent

(us) : K?D = u.K”” .z h cs . (19)

However, when the association constants are small in solutions where one solute is in large excess, say [Dlo ZQ [A],,, the two constants will be related by equation (20):(215,234)

(20)

where MD, M, are the molecular weights of donor and solvent respectively, dn, d, their densities (i.e. u,~ = 1000 L&/M,). and [A],, and [D],, are the total, free and complexed, concentrations of A and D respectively. Alternatively, the association constant may be expressed on a molal scale, viz. KA”. For relatively large values of association constant: 111

KAD = d,KAD m A (‘ . (21)


When the association constants are small (K,AD < - 1 M-‘) for solutions where CD],, Z+ [A],, the two constants are related by the expression:

KAD = K,AD d, - [DldMDds 1000 d.q > ’

(22)

Comparisons between experimental values of KfiDP, K,AD, K,AD are made in Section 3.2.

In cases where a diluting solvent is used there will normally be some competition by the solvent for one or other component in the complex by acting itself as either an electron donor or an electron acceptor.(235-7) Although the appropriate formation constant may be small, the large concentration of the solvent species may have a profound effect on the apparent association constant for the complex AD. For example, if the solvent (S) shows a donor behaviour then it will reduce the number of free acceptor molecules which are available to complex with the solute donor. The measured formation constant, K:, is related to the constant K,AD, which would obtain if there were no complexing, by the relationship:

K;= Kp (l+K;[S])-‘, (23)

where Kz is equal to the quotient [SA]/[A] [S] for the equilibrium A+S ti AS. Very few association constants published are corrected for such competition. The data listed in this chapter will similarly not take account of this effect.

lsomeric Complexes. A further assumption which is tacit in determinations of association constants of charge-transfer complexes is that a single 1 : 1 complex is formed. If in fact several isomeric 1: 1 complexes are present then the experimentally determined value of the association constant K,,, is such that@)

K en, = 7 KiAD9 (24)

where KAD is the association constant for the ith 1 : 1 complex. Orgel and MulliketP) have proposed that optical absorption charac-

terized by the charge-transfer band(s) may arise not only through the TN + qE transition(s) in the charge-transfer complex but also by a similar intermolecular charge-transfer process when molecules of A and of D are in collision. They have termed this absorption “contact charge- transfer”. It was argued that such an effect would mean that the measured molar absorptivity as calculated by, say, the Benesi-Hildebrand method(‘) (see below) would not be equal to the true molar absorptivity of the complex AD. Mulliken and Orgel (=) also stated that (a) the presence of


contact charge-transfer would not affect the calculated value of the association constant for the complex; (b) the formation constant for collisional “contact charge-transfer” interactions is zero. These conclusions have been questioned. Prue (238) has suggested that such collisional interactions should give rise to a small but finite association constant, but this conclusion may be incorrect.

3.2. The Determination ofAssociation Constants for Charge-transfer Complexes in Dilute Solution from Chemical Shift Data and Evidence

for Their Correctness

In principle, any of the methods already discussed for the determination of association constants of equilibria involving fast reactions could be applied to charge-transfer complex formation. However, mainly because of the low solubilities of most organic electron-acceptor species, the experimental determinations of association constants in these cases are usually made on solutions in which the donor component is in large excess.

The methods to be described can be, and have been, applied to suitable non-charge-transfer equilibria.

For the simple case of an equilibrium in dilute solution

A+D=AD (25)

let one component, say A, have for simplicity a singlet absorption at a position aA relative to some reference signal when the concentration of D is zero; let 8*,, be the chemical shift of the same nucleus in the A moiety of the pure charge-transfer complex relative to this same reference. From equation (5) it is seen that the measured chemical shift (8) relative to the reference signal for a solution containing A, D and AD in fast equilibrium is

6 = PASA + PADGAD, (26)

where P, and PDA are the fractions of A which are uncomplexed and complexed respectively. Since (PA + PAD) = 1, equation (26) may be written:

6 = P.&‘&D - 6,) + 8*. (27)

Association constants have been evaluated from this relationship. A plot of PAD against 6 should be linear. The method requires the insertion of an equilibrium constant,? e.g. KAD as a variable parameter which can be

tFor conciseness the following discussion is given in terms of an association constant KAD without specifying the particular concentration unit. For convenience many of the K values have been calculated in the units kg solution mole-l. This is effectively equivalent to KkD (see below).


adjusted until values of P,, obtained from it vary linearly with the corresponding value of 6. The method is obviously suitable for solution by using a computer programme.

However, under suitable conditions KAD may be evaluated by a simple linear relationship. Let (6A -SAD), namely the shift for the pure compIex in solution relative to the shift for the pure acceptor in solution, be represented by A,. Let (6, -S), the observed shift for the system in equilibrium relative to the shift for the pure acceptor in solution, be represented by A. The relationship of these terms is shown in Fig. 11.

I

Acceptor alone

+ A0 ---A I I

I I I

Pure complex (not meosureable)

I

-1

I

FIG. 11. The relationship between the terms A. A,. 6 and 6,); see equation (28).

Equation (27) may now be rearranged:

8*-6 [ADI A pAD=-=_=_ 6A - 8AD [Alo A,’

(28)


For an ideal solution in which ID],, >> [Al,:

[ADI KAD = [Dlo ([Alo - [ADI) ’

From equations (28) and (29) we obtain:‘73~74~87~122~ 161,214)

1 1 1 1 _=- - A K”“*A;[DI+$. 0 0

43

(29)

(30)

A plot of l/A against l/ [Dlo should be linear. The gradient of the line yields the term KADAo, whilst the intercept with the ordinate yields a value for l/Ao. Thus KAD may be evaluated.?+ The form of equation (30) is analogous to the Benesi-Hildebrand equation’” (discussed in Section 3.3) which has frequently been used in the determination of association constants from the analysis of optical data. The application of equation 30) suffers a similar disadvantage to the Benesi-Hildebrand plot: namely that an extrapolation to solutions of high concentration has to be made. Equation (30) was applied to hydrogen-bonded complexes(73*74,87) (see Fig. 6) and later to charge-transfer complexes(21s-6) (see Fig. 12). Parti- cularly in systems where the association constants are small, the extrapolation may lead to incorrect values of A,, and consequently incorrect values of KAD. However, values consistent with optical data have been obtained.‘*16) Equation (30) may be rearranged in the following form:(21g)

A/ [Dlo = - AK*” + AoKAD . (31)

Thus, for a series of solutions in which [Dlo s [A],, a plot of A/[D], against A should be linear, the gradient of the line being equal to -KAD. The value of A,, may be obtained from the intercept with the ordinate. By contrast with equation (30) this requires an extrapolation to an infinitely dilute solution, and the evaluation KAD is itself obtained without recourse to an extrapolation. Equation (31) is the nuclear magnetic resonance analogue of the relationship between absorbance and [D], derived from optical data (see Section 3.4).

The application of equation (3 1) to a series of methylbenzene-fluoranil complexes by measurements of the ‘“F shift in fluoranil (A) for a series of solutions where the concentration of fluoranil is low compared with the concentration of the particular methylbenzene (D) is shown in Fig. 1 3.(224)

lRelations similar to equation (30) and to equation (3 1) (see below) may be written for the shift of donor nuclei for solutions in which [A], B [D],.

*Equation (30) is identical with equation (8).


32

20

24

20

16

12

8

4

L I I I I I I I 0 2 4 6 6 IO 12 14

_!_ (kg mole-‘)

mi

FIG. 12. Plots of l/A,,,, vs. llmDo (where M Do is the total molal concentration of the donor) for ‘H chemical shifts in 7,7,8,84etracyanoquinodimethane using as donors: mesitylene, 0; durene, 0; pentamethylbenzene, 0; and hexamethylbenzene, A; in

dioxan as solvent at 37°C. (M. W. Hanna and A. L. Ashbaugh.(*14))

The method yields, within experimental error, the same value of KAD whether a series of solutions in which [D],, B [A],, and measurements of a nucleus in the species A are used, or the condition that [D], 4 [A],, and measurements of a nucleus in the species D are used (see Table 6).(22g)

For a component in low concentration which contains two or more non-equivalent protons, the measurements of chemical shift for these nuclei in the presence of large excesses of the second component also yield consistent values for KAD (see Table 7).(225*227)

Further support for the nuclear magnetic resonance method is obtained from measurements of different nuclear species in a given complex. The association constant for hexamethylbenzene (D) 1,3-difluoro-2,4,6- trinitrobenzene (A) in chloroform has been determined(220) under the condition [D], % [Alo from measurements of the chemical shift of the ‘H nucleus in the acceptor and also of the two equivalent lsF nuclei in the same molecule. The results are summarized in Table 8.

3-

O-

f

c

3


F

\

\

D

C

1 \)i, A \ ‘1

50 100 150

n

FIG. 13. Plots of A/[D], (where [D], is the total concentration of the donor in moles per kg of solution) against A for the complexes of fluoranil with: (A) benzene; (B) toluene; (C) p-xylene; (D) mesitylene; (E) durene; (F) pentamethylbenzene; and (G) hexamethylbenzene in carbon tetrachloride as solvent at 334°C. (From the data

of N. M. D. Brown, R. Foster and C. A. Fyfe.‘Z24’)

TABLE 6. ASSOCIATION CONSTANTS ( KAD), IN KG MOLE-', AND THE CHEMICAL SHIFTS RELATIVE TO THE PURE COMPONENT (A,)FoR THE 1,3,5-TR~NITROBENZENE COMPLEXES OF MESITYLENE AND OF DIJRENE

IN 1,2-DICHLOROETHANE SOLUTION AT ~~.~"UNDER THE CONDITIONS [A], * [D] AND [A],, Q [D],, (C. A. Fyfe(229))

Relative cont.

IAl, + KG

Plo * [Al,

Mesitylene Durene

KAD &la KAD &la (kg mole-‘) (c/s) (kg mole-‘) (c/s)

0.26 2 0.07 73~25 0.4OkO.l 712 18

0.28 -t 0.05 59-t 8 0.33 f 0.06 71 -c 10

a At 60.004 MC/S.


TABLE 7. VALUES OF ASSOCIATION CON- STANT (IP) AND RELATIVE CHEMICAL SHIFT (A,)FROM THE SHIFTS OF H, AND

~~~~~~~~~~~~~~~

hexamethylbenzene in Ccl, (33.5")

(M. I. Foreman.IzL7))

Nucleus KAD AlI”

(kgmole-I) (c/s)

HLI 1.4, 54

HP 1'5" 98

a At60.004 MC/S.

TABLE 8. VALUES OF ASSOCIATION CONSTANT (KAD) AND RELATIVE CHEMICAL SHIFT (A,) FROM

THE'H AND~~FCHEMICALSHIFTSIN I.3-DIFLUORO-

2,4,6_TRINITROBENZENE IN THE COMPLEX WITH

HEXAMETHYLBENZENE IN CHLOROFORM AT 33.5" (R. Foster and C. A. Fyfe'*20J)

'H 40.004 0.4710.02 93.2

1°F 37.664 0.4510.02 140.7

Where a choice of acceptor is possible, it is obviously desirable to make measurements on nuclei which yield singlet absorptions. Com- pounds which have this feature and also are good acceptors include the proton-containing molecules: 1,4_dinitrobenzene; 1,3,5_trinitrobenzene; 2,6-dichloro-p-benzoquinone; 2,5-dichloro-p-benzoquinone and 2,3- dicyano-p-benzoquinone. The use of lsF resonances, as opposed to ‘H, has certain advantages.(220~224) There is no interference from other proton absorptions by either component or solvent. The l”F shift differences also


tend to be considerably larger than the ‘H shift differences so that the accuracy of the determination of K,AD is increased. Easily accessible fluorine-containing electron acceptors include: 2,3,5,6-tetrafluoro-p- benzoquinone (fluoranil) and 1,4-dicyano-2,3,5,6_tetrafluorobenzene.

Equation (31) may be extended to apply to systems containing several acceptors Al, A’, A”, . . . simultaneously in the presence of a single donor.‘223’ The equilibrium constants corresponding to the various complexes A’D, A2D, A3D may be evaluated from the data for a single set of solutions. Apart from the assumption, already made, that the reactions involved are very fast, the only condition which has to be

200 -

150-

0

E

2

100 -

50- A4m , /// '

{/,',,'

/,/ ' //'/'

15:' /9'

/F I I I

0 20 40 60

n

FIG. 14. Plots of Al[D], (where [D], is the total concentration of the donor in moles per kg of solution) against A for the complexes of: (A’) 1,3,5+initrobenzene; (AZ) 2,5-dichloro-p-benzoquinone; (A3) 1,4_dinitrobenzene and (A4) p-benzoquinone with hexamethylbenzene in carbon tetrachloride as solvent at 33.5”. These measurements are on solutions containing all these acceptors together. (From the data of

R. Foster and C. A. Fyfe.‘223’)

P

00

TABLE

9. VALUES OF ASSOCIATION CONSTANT

(KA

D) IN KG MOLE-'ANDTHE

CHEMICAL

SHIFTFORTHE

PURE COMPLEXRELATIVE

TO THE ACCEPTOR

(A,) FOR COMPLEXES

OF HEXAMETHYLBENZENE

WITH:(A')I,~,~-TRINITROBENZENE:(A')

2,5-DICHLORO-p-

BENZOQLJINONE: (A3) 1,4-DINITROBENZENE: (A4) p-BENZOQUINONE

IN CARBON

TETRACHLORIDE

AT 33.5", MEASURED

SIMUL-

TANEOUSLY,IN

PAIRS AND SEPARATELY(R.

Foster and C. A. Fyfe(z*3')

Acceptor

A'

A2

A3

A4

Combination

K A

D

I I

&la

K

AD

I

I A

oa

KA

D

I I

&I”

K

AD

&

I”

(kgmole-')

(c/s)

(kgmole-')

(c/s)

(kgmole-I)

(c/s)

(kgmole-')

I (c/s)

Pairs A', A3; AZ, A"

5.2

65

2.1

87

0.9

96

0.7

81

A1,A2,A3,A4fogether

4.8

66

2.1

86

0.9

97

0.6

90

Each s

ingl

y 5.1

65

1.9

91

1.0

91

0.7

89

aAt60.004 MC/S.


fulfilled is that the total concentration of the various species of acceptor is very small compared with the concentrations of donor used. The problem of overlapping bands does not often arise because the lines are narrow and they usually move in the same direction when the donor concentration is increased. The technique provides a useful method for comparing the complexing abilities of different acceptors in a common solvent-solute environment. Plots of A/ [III0 against A for a series of hexamethylbenzene solutions containing several different acceptors are shown in Fig. 14. The values of KAD obtained from these plots show good agreement with the values obtained from determinations involving each acceptor singly or pair-wise (see Table 9).

Trotter and Hanna(*15’ have emphasized recently the effect of different solvent scales on the optical evaluation of association constants and extinction coefficients for the measured complexes. This effect has been referred to already in Section 3.1. The problem had been discussed previously by Scott.(234’ For an assumed ideality on any one scale, deviations from the simple relationships expressed by equations (19)

TABLE 10. ASSOCIATION CONSTANTS AND A, VALUES FOR 2,3,5,6-TETRAFLUORO- p-BENZOQUINONE METHYLBENZENE COMPLEXES IN CARBON TETRACHLOR~DE

CALCULATED FROM THE ORIGINAL EXPERIMENTAL DATA ON THE ‘<KG SOLUTION PER MOLE SCALE” (KAD), MOLE-FRACTION SCALE (Kit’), AND MOLAL (KkD)

SCALE (R. Foster, C. A. Fyfe and J. W. Morris’Z32’)

Donor KAD KV A KAD m KcD/KAD =

Toluene p-Xylene Mesttylene

Durene Pentamethyl-

benzene Hexamethyl-

benzene

0.96 5.9 I.1 6.2

1.5 9.3 1.7 6.2

2.2 14.4 2.4 6.6

4.9 30.1 4.7 6.1

7.9 49.6 8.0 6.3

15.4 100.6 16.1 6.5

I

KA”,KAD b m

1.1

1.1

1.0, 0.9,

1.0,

1.0,

I-

Donor

Toluene p-Xylene Mesitylene Durene

Pentamethyl- benzene

Hexamethyl- benzene

A, A0 (kg solution per mole)’ (mole-fraction)’

(e/s) (e/s)

85 90

108 112

115 117 139 141

152 154

174 169

(mo4ual)

(e/s)

77 100 108 137 150

167

aFor ideal dilute solutions ratio is 6.5. bFor ideal dilute solutions ratio is 1.0. ‘Measured at 37.664 MC/S.


and (2 I ) become serious at low values of association constant. However. for values of K”j.’ above - I K’, this effect does not appear to be important in relatively dilute solutions. CalculationP2 ah inirio of KAr),

K‘t” and K:,” for some systems have been made from the necessary

experimental data. Examples are given in Table 10. The values obtained are consistent with equations ( 19) and (2 1).

Abraham”“” has used the variation of the measured shift, A, with

temperature to obtain the enthalpy of formation (AH) for solvent-solute interactions of dilute solutions. If P is the fraction of the solute species which is complexed then:

&=expgiexp(-g). (32)

At a given temperature P = A/A,, hence a plot of In A vs. l/7‘ will give AH directly providing that A,, is temperature independent (recent work”:‘:”

has shown that for many charge-transfer complexes the ‘H A,, values are in fact temperature independent). This method for obtaining AH is obviously applicable to charge-transfer complexes in dilute solution.

3.3. Results Ohtrrinrdfrom Chemical Shift Merrsurements oj

Chcrrge-trrrnsfer Equilihriu in Dilute Solution

3.3. I. Summary oj’results

Over the past two years a number of chemical-shift determinations of association constants for typical organic charge-transfer complexes have been made.(ll-c.rl(i,sl!‘3”’ A selection of the results is given in Tables 1 l-20. The experimental plots obtained for the determinations using equation (31) are similar to those given in Fig. 13. The good iinear correlation gives strong support for the basic assumptions made (see Section 3.2). and argues against the presence of termolecular species.

3.3.2. A general discussion ofassociation constant values

There have been many discussions of the relationship of optically determined values of K, for groups of complexes, with various properties of the donor, acceptor and solvent species. A particularly detailed study has been made by Brieg1eb.c”) A few observations on such relationships involving values of K*” obtained from chemical shift measurements will now be made.

(a) Dependence of ussociution constant on the donor component. The association constants of complexes of a given acceptor with a group of structurally-related donors, for example the methylbenzenes, increase as the Lewis basicity of the electron donor increases,(214s”“) as would be


TABLE Il. ASSOCIATION CONSTANTS ( KAD), IN KG SOLUTION PER MOLE, AND ‘H CHEMICAL SHIFTS FOR THE PURE COMPLEX REI.ATIVE TO THE

ACCEPTOR (A,,) FOR COMPLEXES OF VARIOUS DONORS WITH 1,3,5-TRI-

NITROBENZENE IN CARBON TETRACHLORIDE Sor LITION AT 33.5”

Donor

Benzene Benzene-d, Toluene p-Xylene Mesitylene Durene Pentamethylbenzene Hexamethylbenzene 1,3,5-Triethylbenzene Hexaethylbenzene Ethylbenzene n-Propylbenzene Isopropylbenzene n-Butylbenzene tert-Butylbenzene n-Propylbenzene wHexylbenzene n-Heptylbenzene n-Octylbenzene rt-Nonylbenzene Aniline N-methylaniline N-ethylaniline N-n-propylaniline N,N-dimethylaniline N,N-dimethylWbromoaniline N,N-diethylaniline o-Toluidine N,N-dimethyl-o-toluidine N.N-diethyl-o-toluidine m-Toluidine p-Toluidine 2-Ethylaniline 4-Ethylaniline 2,4-Xylidine N,N-dimethyl-2,4-xylidine 2,6-Xylidine

“At 60.004 MC/S.

K””

(kg mole--’ )

0.5,) 0,5,, 06, 0.9, 1.2, 2.1, 3.0,, 5.1, 0.8,

< 0.05 0.7,, O,S,, 0.6, 0.6,

0.4, 0.6, 0.5, 0.5,, 0.5,;

0.60 2.2, 2.7,, 3.1; 2,8;, 3.2,; 2.9, 3.6,

3.0,) I .h, I .4, ;2,, 7,2,,

?,9,, 3.2,

4.1,) 2.1,) 3.5,

A,,>’ (C/S)

___

75 75 71 68 61 62 66 65 59 -

62 63 60 74 12 72 70 73 76 76 54 59 56 58 61 49 60 48 53 56 53 54 56 53 57 S6 61

Refs.

219 239 219 219 219 219 219 219 226 226 226 230 230 230 230 230 230 230 230 230

219.226 226 226 226

219.226 226

!19,?26

226 226 226 226 226 226 226 226 226 226


TABLE 12. ASSOCIATION CONSTANTS (KAD), IN KG SOLUTION PER MOLE, AND IH CHEMICAL SHIFTS FOR THE PURE COMPLEX RELATIVE TO THE ACCEPTOR (A,) FOR COMPLEXES OF VARIOUS DONORS WITH 1.35TRI-

NITROBENZENEINCHLOROFORMAT~~.~

Donor

Durene Pentamethylbenzene Hexamethylbenzene Indole 2-Methylindole 3-Methylindole 7-Methylindole 3-Ethylindole 3-Dimethylaminomethylindole Benzothiophen Thiophen Pyrrole

aAt 60.004 MC/S.

K AD (kg mole-‘)

0.4, 0.6, 0.8, 1.6, 2.1, 2.4, 1.9, 1’7, 1.3, 0.8, 0.1, 0.2,

-_ Aoa

(c/s) Refs.

80 224 79 224 80 224 99 225 99 225 98 225 99 225

102 225 84 225 99 225

128 225 69 225

expected. A linear relationship of log KAD with the ionization potential of the donor species is not observed. Such a relationship has been reported on occasion for KAD values obtained from optical data. However, in terms of the valence bond descriptions of the ground-state of a weak charge- transfer complex (see equation (2) in Section 1.2), the binding energy in the ground-state will not be solely influenced by the dative structure. Con- sequently a linear variation of log KAD (as a measure of the free energy of

TABLE 13. ASSOCIATIONCONSTANTS(K*~).INKGSOLUTIONPER MOLE, AND 19F CHEMICAL SHIFTS FOR THE PURE COMPLEX RELATIVE TO THE ACCEPTOR(A&FORASERIESOF COMPLEXES OF VAUIOUS DONORS WITH 2,3,5,6-TETRAFLUORO-P-BENZO-

QUINONE(FLUORANIL)INCARBONTETRACHLORIDEAT 33.5”

Donor KAD

(kg mole-‘) Refs

Benzene Toluene p-Xylene Mesitylene Durene Pentamethylbenzene Hexamethylbenzene Ethylbenzene n-Propylbenzene n-Butylbenzene

“At 37.664 MC/S.

0.7, 66.5” 0.9, 85.5” 1.5 107.5a 2.2 115” 4.9 139a 7.9 152.5=

15.4 174= 1.1, 124” 0.7, 129b 0.9, 105b

bAt 56.458 MC/S.


TABLE 14. ASSOCIATION CONSTANTS (KAD), IN KG SOLUTION PER MOLE, AND 19F CHEMICAL SHIFTS FOR

THE PURE COMPLEX RELATIVE TO THE ACCEPTOR (AJ FOR COMPLEXES OF VARIOUS DONORS WITH 2,3,5, 6-TETRAFLUORO-p-BENZOQUINONE (FLUORANIL) IN CHLOROFORM AT 33.5” (N. M. D. Brown, R. Foster

and C. A. Fyfe’224))

Donor KAD

(kg mole-‘) A,a

(c/s)

Benzene 0.16

Toluene 0.2, p-Xylene 0.4, Mesitylene 0.68 Durene 1.3 Pentamethylbenzene 2.0 Hexamethylbenzene 3.9

=At 37.664 MC/S.

formation of the complex) with ionization potential of the donor molecule should not in general be expected. However, plots of log KAD for such a series of complexes of a given acceptor are proportional to the values of log KAD for the corresponding complexes of a second acceptor in many caseP4) as shown in Fig. 15 (see also Fig. 16). Some very weak donors do not fit this type of correlation. In such cases charge-transfer is not a dominant factor in the ground state.

The primary steric effect between donor and acceptor has been observed previously in, for example, the large difference between the

TABLE 15. ASSOCIATION CONSTANTS (KAD), IN KG SOLUTION PER MOLE, AND 19F CHEMICAL SHIFTS FOR THE PURE COMPLEX RELATIVE TO THE ACCEPTOR (AO) FOR COMPLEXES OF VARIOUS DONORS WITH 1,4-

DICYANO-2,3,5,6-TETRAFLUOROBENZENZENE IN CARBON TETRACHLORIDE AT 33.5” (N. M. D. Brown, R. Foster

and C. A. Fyfe’224J)

Donor KAD

(kg A,” (c/s)

“At 37.664 MC/S.


TABLE 16. ASSOCIATION CONSTANTS (Kr) and ‘H CHEMICAL SHIFTS FOR THE PURE COMPLEX RELATIVE

TO THE ACCEPTOR (A,,) FOR COMPLEXES OF VARIOUS

DONORS WITH 7,7,8,8-TETR.~CYANOQUINODlMETHANE

IN DIOVAN AT 374t0.5” (M. W. Hanna and A. 1.. Ashbaugh’2’1’)

p-v

Donor A,, a

@pm)

Benzene Toluene o-Xylene ‘I Mesitylene Durene Pentamethylbenzene Hexamethylbenzene

L( Dependent on solvent concentration scale used, the values quoted are based on the molal scale. I’ Identical values for o-. w- and p-xylene.

K AD for the complex 1,3,5-trinitrobenzene- hexamethylbenzene in carbon tetrachloride is 5.1 kg moleP1, compared

TABLE 17. ASSOCIATION CONSTANTS (K""). IN KG SOLUTION PER MOLE, AND IH OR 19F

CHEMICAL SHIFTS FOR THE PURE COMPLEX RELATIVE TO THE ACCEPTOR (A,,) FOR

COMPLEXES OF VARIOUS ACCEPTORS WITH HEXAMETHYLBENZENE IN CARBONTETRA-

CHLORIDE AT 33.5”

Acceptor Nucleus

1,4_Dinitrobenzene 1,3,5-Trinitrobenzene 3,5-Dinitrobenzonitrile 3.5-Dinitrobenzonitrile 3,5-Dinitrobenzotrifluoride 3,5-Dinitrobenzotrifluoride 3,5_Dinitroanisole 3,5-Dinitroanisole N,N-dimethyl-3,5-dinitroaniline N,N-dimethyl-3,5-dinitroaniline 1,4-Dicyano-2,3,56tetrafluorobenzene

p-Benzoquinone 2.5-Dichlot-o-p-benzoquinone Fluordnil

Kh”

(kg mole-‘)

A,,;[

(c/s) Refs.

1.0, 91 219 5.12 65 219 3.5, 69 227 3.3, 75 227

1.4, 54 227 I .5,, 98 227 l.li 48 227 I .o, 62 227

1.4, 38 227 I .O!, 40 227 5.2 I I2 224 0.6,; 89 222 1.9, 91 222

15.4 174 224

,*Proton shifts at 60-004 MC/S, fluorine shifts at 37,664 MC/S.


TABLE 18. ASSOCIATION CONSTANTS ( ICkD). IN KG SOLUTION PER MOLE, AND ‘H OR lSF CHEMICAL SHIFTS FOR THE PURE COMPLEX RELATIVE TO THE ACCEPTOR (A,,) FOR COM- PLEXES OF VARIOLIS ACCEPTORS WII‘H HEXAMETHYLBENZENE IN 1,2-DKHLOROETHANE

OR CHLOROFORM AT 33.5”

Acceptor

1,4-Dinitrobenzene I ,3,5_Trinitrobenzene I .2,3,5_Tetranitrobenzene I ,3,%Tricyanobenzene 1,2,4,5-Tetracyanobenzene 2,3,5.6-Tetrafluoro-1,4-dicyannbenzene p-Benzoquinone 2.5.Dichloro-p-benzoquinone 2.3-Dicyano-p-benz.oquinone 2.3,5-6-Tetrafluoro-p-benzoquinone 7,7,8&Tetracyanoquinodimethane I -Chloro-2,4,6_trinitrobenzene I-Bromo-2,4.6-trinitrobenzene 1 -Iodo-2,4,6-trinitrobenzene 2.4.6-Trinitroanisole

;olvent”

D D D D D D D D D D D C C C C

Nucleus

‘H ‘H ‘H !H ‘H

‘!‘F

‘H ‘H ‘H

1°F

‘H ‘H ‘II ‘H ‘H

k’“D

kg mole-‘1

0.15 0,5$, 0.6,; 0.1, 0.4, 0.7, 0.1, 0.6, 2.8” 3.6,, 0.7,: O.ZL 0.2, 0.2, 0.2,

--

&,h C/S)

177 85 99

292 115 128 125 100 72

125 64

156 143 165 125

~

Refs.

219 219 222 222 222 224 222 222 222 224 222 226 226 226 226

“Solvents: D = 1,2-dichloroethane, C = chloroform. “Proton shifts at 60.004 MC/S, fluorine shifts at 37.664 MC/S

with a value of less than 0.05 kg mole-’ for the corresponding hexaethylbenzene complex. In the complexes formed between I ,3,5trinitrobenzene and straight chain alkylbenzenes C,H,R where R ranges from CH,- to n-C,H,,-, the equilibrium constant rises to the toluene complex, then drops to successively lower values for the ethyl and n-propylbenzene complexes (see Table 11). The value for n-propylbenzene appears to be a limiting value for the higher homologues, for which all the K”D values are the same within experimental error. w”) Similar observations have been made on fluoranil complexes of the alkylbenzenes up to n-propylbenzene (see Table 11). The effect of side-chain branching has been examined(230) for complexes of 1,3,_5-trinitrobenzene (see Table I 1). The association constants increase on replacement of H by methyl, then follows a continual decrease as the substituent is altered to ethyl, i-propyl, t-butyl. This follows the expected order.

Steric effects are observed in complexes with substituted anilines(226) (see Table 11). Thus N,N-diethylaniline complex with 1,3,5-trinitrobenzene has a smaller association constant than has the N,N-dimethylaniline complex. This is probably a primary steric effect. The 1,3,5_trinitrobenzene complexes of N ,N-dimethyl-o-toluidine and of N,N-diethyl-o-toluidine are considerably weaker (see Table I I). This is


TABLE 20. ASSOCIATION CONSTANTS (KAD), IN KG SOLUTION PER MOLE, AND CHEMICAL SHIFTS FOR THE PURE COMPLEX RELATIVE TO THE ACCEPTOR (A,J FOR COMPLEXES OF VARIOUS DRUG DONOR MOLECULES WITH p-DINITROBENZENE AT 33.5” (R. Foster and

C.A. FyfeoZ1))

Donor Solvent

Chlorpromazine CHCI, Promethazine CHCI, Diethazine CHCI, Ethopropazine CHCI, Promazine CHCI, Trimeprazine CHCI, IO-Methylphenothiazine CHCl, Chlorpromazine ccl, Promethazine CCI, Diethazine CCI, Ethopropazine CCI, Promazine ccl, Trimeprazine CCI,

aAt 60.004 MC/S.

KAD (kg mole-‘)

AO” (c/s)

0.2” 143 0.2, 153 0.2, 132 0.2, 167 0.2, 135 0.2, 146 0.36 124 1.19 59 1.32 72 1.2, 63 1.2, 76 1.4” 71 1.2, 77

probably the result of the ortho steric effect which prevents the molecule from assuming a planar structure, thus reducing the delocalization.

The association constant of 1,3,5_trinitrobenzene with N,N’-dimethyl- N,N’-di-p-tolylhexamethylene diamine (XXVI) has been comparedzz8’ with the association constant of N-methyl-N-propyl-p-toluidine; the latter being effectively “half” the molecule XXVI. The association constant of XXVI with 1,3,5_trinitrobenzene is twice that of N-methyl- N-propyl-p-toluidine (see Table 11). This may be accounted for on a statistical basis if molecules of XXVI present two alternative sites for simple 1: 1 complexing with the acceptor. It would not be expected if both 7r-systems in XXVI were involved in complexing with a single acceptor molecule.


(b) Dependence of association constant on the acceptor component. Some regularities in values of KAD are observed when the acceptor component is altered,‘222~224~ as for example in Fig. 15, where there is a linear relationship between the log K-‘D values of various sets of complexes of different acceptors. Some data show a linearity between the various

l-2

0.6

-0-6

FIG. 15. Plots of log,, K4” for complexes of a series of methylbenzene donors with the acceptor fluoranil against the values of log,, KAD for the complexes of these donors with 1.3.Strinitrobenzene, (A) in carbon tetrachloride solution,(B) in chloroform solution, all at 33.5”. (N. M. D. Brown, R. Foster and C. A. Fyfe.(?

KAD values themselves (see Fig. 16). These correlations lend support to the method of evaluating association constants from chemical shift data. The linearity of the correlations may not always occur. Indeed, in the cases illustrated in Fig. 16 there must be some deviations from linearity at low values of K since these plots do not pass through the origin.


I I I I I 0 I 2 3 4 5

K~~(TNS)

FIG. 16. Plots of KAD (in kg solution mole-‘) for a series of methylbenzene donors with the acceptors: (A) p-benzoquinone; (B) 1,4_dinitrobenzene; (C) 2,5-dichloro- p-benzoquinone, plotted against the KAD values for the corresponding complexes of I .3Strinitrobenzene in carbon tetrachloride at 33.5”. (R. Foster and C. A. Fyfe.‘***))

The K values for complexes of 2,5-dichloro-p-benzoquinone are less than those of the corresponding complexes of 1,3,5_trinitrobenzene. Like- wise the K values for complexes of p-benzoquinone are less than those for the complexes of 1,4dinitrobenzene. By contrast, the values of the frequencies (z$,D,,) of the optical charge-transfer bands show the opposite trend. Thus even in non-sterically hindered systems it is not possible to arrange electron acceptors in an order of increasing “strength” which will be the same whether vCT or K (or -AC) is used as a measure of strength. Briegleb@) had shown previously a similar lack of correlation based on purely optical data although a few positive correlations for small groups of complexes have been reported. The effect of the nature of the solvent, which may alter the order in the values of K for a series of complexes of a given donor with various acceptors (see below) without alter- ing the order of the respective vgX values, is a further reason why it is not possible to set up an unambiguous order of acceptor strengths based on K values. Nevertheless, some gross effects may be noted. For related compounds the introduction of electron-withdrawing groups generally results


in higher K values. In this respect it is interesting to note that no changes in chemical shift are observed in hexafluorobenzene solution in Ccl, when large excesses of hexamethylbenzene are introduced. This implies that either K or A0 or both are very small for this system. Evidence for some type of complex in solution comes from measurements of the excess heats and volumes of mixing of hexafluorobenzene in the lower methylbenzenes which are liquids at or near room temperature,‘241-2’ and from infrared measurements.(243)

It is of interest that, whereas p-benzoquinone is a much weaker base than fluoranil, and 1,4_dicyanobenzene is very much weaker than its pet-fluoro analogue (see Table 17), 1,3,5trinitrobenzene is a considerably stronger acceptor than 2,4-difluoro- 1,3 ,5-trinitrobenzene.(224) This difference may be explained by the steric inhibition of resonance of the nitro groups by the fluorine atoms in the latter compound. Further examples of intramolecular steric effects in the acceptor component are illustrated by the values of the association constants of hexamethylbenzene with picryl fluoride, picryl chloride and picryl iodide compared with 1,3 ,5-trinitrobenzeneczz6) (see Table 17). A similar steric effect is probably operating in 1,2,3,Stetranitrobenzene.

(c) Effect of solvent. In a series of complexes with methylbenzenes the effect of solvent on the relative values of the association constants of the complexes with a given acceptor appears to be nearly independent of the particular donor used (see Fig. 17). Although for complexes with a given acceptor, KAD decreases as the solvent changes in the order Ccl, 9 CHCl, > CH,C lCH,C 1 = CH,C12, the amount of this change varies from acceptor to acceptor. It is therefore not possible to arrange un- ambiguously the acceptors in an order of “strength” based on K values obtained from complexes with a common donor species unless the solvent is specified.

The order of solvents in terms of their effect on K for a given complex as cited above (see Table 19) reflects the expected solvent competition for the acceptor species by these various solvents (see equation (23). In systems where dioxan has been used as solvent (see Table 16) it has been suggested that specijc solvation of both free and complexed molecules of the acceptor occurs.(215)

(d) Effect of temperature. Preliminary investigations have been carried out on the variation of KAD with temperature for the series of complexes of 1,3,5_trinitrobenzene and of fluoranil with various methylbenzenes.(233) Values of the enthalpy (AH) and entropy (AS) of formation of the complexes are given in Table 2 1. Good van’t Hoff plots were obtained. In the case of the 1,3,5_trinitrobenzene complexes, A,, was found to be independent of temperature within the experimental error (see Section 3.3.3 for a


FIG. 17. Plots of KAD (in kg solution mole-‘) for complexes of a series of methylbenzene donors with: (A) fluoranil; (B) 1,4-dicyano-2,3,5,6-tetrafluorobenzene in carbon tetrachloride solution against K AD for the same complexes in chloroform solutions,

all at 33.5”. (N. M. D. Brown, R. Foster and C. A. Fyfe.czz4))

discussion of the A0 values). For the fluoranil complexes, the shifts were observed to be somewhat temperature dependent. However, the relative shift of the free fluoranil is itself temperature dependent to about the same extent when measured against the internal reference, fluorotrichloro- methane. These variations may be the result of changes in shift of the reference standard.

The values of AH and AS for the 1,3,54rinitrobenzene complexes agree in sign and in order of magnitude with the values obtained from optical data,(244) but there is disagreement on the relative magnitudes of AH within a given series. The nuclear magnetic resonance data would suggest that values of AH increase regularly for the various methylbenzene complexes of 1,3,5_trinitrobenzene up to and including hexamethylbenzene within the present admittedly large error (see Table 21). The same trend is observed in the fluoranil complexes (see Table 21). Results from optical data for 1,3,5_trinitrobenzene complexes show a less


TABLE 21. ENTHALPIES AND ENTROPIES OF FORMATION OF 1 ,3,5-TRINITROBENZENE AND

OF FLUORANIL CHARGE-TRANSFER COMPLEXES FROM NMR DATA ( IN Ccl,) (R. Foster, C. A. Fyfe and M. 1. Foreman’233’)

Acceptor I ,3,5_Trinitrobenzene Fluoranil

Donor -AH -AS

I

-AH -AS (kcal mole-‘) (cat mole-’ deg-I) (kcal mole-‘) (cal mole-’ deg-‘)

\ 1 1

Benzene 1.9, 7.8 2.0, 74 Toluene 2.1, 7.9 2.3,) 7.5 p-Xylene 2.5, 8.2 2.7,) 8.0 Mesitylene 2.8, 8.8 3.0, 8.1 Durene 3.0, 8.5 3.9, 9.6 Pentamethylbenzene 3.3, 8.6 4.4, 10.1 Hexamethylbenzene 3.6, 8.8 5.4, 12.1

regular trend in AH values. (244) Further experimental work is required to provide more accurate values of AH. In this respect, the method of Abraham(‘O’) (see Section 3.2) would be of use in systems where A, has been shown to be temperature independent as it would enable a rapid collection of experimental data to be made, and it would minimize effects resulting from experimental inaccuracies in making up solutions. It is likely that direct calorimetry would provide a more accurate estimate of AH.

3.3.3. A discussion ofA, values A0 is the chemical shift difference of a nucleus between the free and

complexed forms of one of the components (see Fig. 11). Attempts to interpret the precise chemical shifts of simple molecules have not been wholly successfu1.(245-fi) It should therefore not be surprising that the magnitudes of A,, for the various complexes are not understood completely. However, several characteristics of charge-transfer interactions are likely to contribute to A,,: the protons in an acceptor species might be affected by the ring current of the donor molecule, by transfer of charge in the complex and by differences in anisotropy of polar groups in the acceptor molecule caused by the transfer of charge. Although no quantitative description can be given, several broad trends in the values of A0 emerge from the data quoted in Tables 1 l-20.

(i) Hydrogen A,, values. For single-ring to single-ring interactions, the A0 values are, in general, within the range 60-80 c/s at 60 MC/S. For a set of complexes of a given acceptor with a series of donors, for example the methylbenzene- I .35trinitrobenzene complexes (see Table I I).


the values of A, are constant. Because A,, is not dependent on the nature of the donor, it would seem that transfer of charge is not an important factor, since A, should increase as the ionization potential of the donor decreases. Further evidence against the importance of transfer of charge is that the actual degree of electron transfer appears to be small, and the w:ork of Leto et ~11.(*~~) and of Fraenkel et u1.(248) suggests that the ‘H shifts change by on/y 10 ppm for a one-electron charge on an adjacent carbon atom. The smallness of the range of A, values and the apparent independence of the particular donor species suggest that ring current effects may be important. Several observations support this general hypothesis.

(a) For systems which have been studied under the conditions [A],, 4 [D],, and [A], ZG [D],, the A, values for the donor and the acceptor ring protons are the same within experimental error, and are both in the same direction, namely upfield from the absorption bands of the parent molecules. This latter observation is further evidence against any theory which would give importance to transfer of charge.

(b) The A,, value of 1,3,5_trinitrobenzene when complexed with excess naphthalene in dichloromethane is much larger (138 c/s at 60 MC/S) than when complexed with alkylbenzenes in dichloroethane (- 80 c/s at 60 MC/S). This would be expected on a ring-current argument since the acceptor protons are closer to the centre of the donor ring-current in the case of the naphthalene complex.

(c) When 1,3S_trinitrobenzene is added to naphthalene in 1,2-dichloroethane, the structure of the multiplet changes (see Fig. 18). Analysis

- - L---

IOC/E, lOC/S IOC/S

FIG. 18. Proton magnetic resonance spectra at 60 MC/S of naphthalene in 1,2- dichloroethane and the theoretical spectra; (A) 0.05 M naphthalene (JA = 0 c/s; J, = 6 c/s; J’ = 1.4 c/s; J = 8.6 c/s, u,6 = 21 c/s); (B) 0.05 M naphthalene plus 0.89 M

1,3,S+initrobenzene (~8 = 15.9 c/s); (C) 0.05 M naphthalene plus 2.37 M 1.3,5-

trinitrobenzene (v,S = 12.1 c/s). (C. A. Fyfe.‘22y’)


shows that although the coupling constants remain approximately unaltered, the chemical shift difference between the H, and H, protons of the naphthalene decreases,(225) Koch and ZollingePO) have made the same experimental observation.

This implies that the H, proton resonances are shifted more than the H, proton resonances, in agreement with the ring current hypothesis involving the symmetrical structure XXVII in which the H, protons are closer to the acceptor ring current than the H, protons. Both H, and HP A, values (A0 multiplet centre - 30 c/s at 60 MC/S) are less than the A0 value for 1,3,5trinitrobenzene in the same complex, as would be expected. They are also smaller than the corresponding value for a single ring donor, again in agreement with the ring current hypothesis. The nuclear magnetic resonance spectra of anthracene in the presence of excesses of 1,3,5trinitrobenzene (see Fig. 19) show a similar behaviour and may be explained in terms of structure XXVIII.

(d) The fact that the hydrogen A0 values are temperature independent (see Section 3.3.2) is in agreement with the ring current hypothesis.

Although the observed results can be rationalized qualitatively in this way, there are problems. The normal intermolecular distances in weak charge-transfer complexes, as determined by X-ray diffraction, are 3 to 3.5 A, and quantitative ring current calculations, for example those of Johnson and Bovey,(24g) would predict maximum shifts of about 20 c/s at 60 MC/S at these distances. Further, although no correlation of the lH A, values with donor strength has been noted in the work quoted in Tables 11 and 12, such a correlation has been recorded by Hanna and his co-workers’214-*5’ (see Table 16). In this case A0 decreased as the degree of complexing of 7,7,8,8-tetracyanoquinodimethane with a series of alkylbenzenes increased. This was attributed to an alteration in the paramagnetic contribution to the acceptor proton shifts.

(ii) Fluorine A, values. The A0 values for lgF resonance are often at least twice as large as those for ‘H. This fact coupled with the relatively large association constants of many complexes of fluorine-containing acceptors makes lgF measurements very attractive. The lsF values are


- 12 c/s

(d)

- I2c/s

FIG. 19. ‘H NMR spectra at 60 MC/S of anthracene ,in 1,2-dichloroethane and the theoretical spectra; (A) 0.37 M anthracene (JA = 0.2 c/s, J, = 6.5 c/s, J’ = 1.4 c/s, J = 8.6 c/s, v,ci = 33.1 c/s); (B) 0.37 M anthracene plus 1.02 M 1.3,Strinitrobenzene

(~“6 = 27.2 c/s); (C) 0.37 M anthracene plus 1.47 M 1,3,5-trinitrobenzene (q,S = 24.4 c/s); (D) 0.37 M anthracene plus 2.42 M 1,3,5_trinitrobenzene (~6 = 23.5 c/s). (C. A.

Fyfe.‘229’)

discussed separately from ‘H values because they show some quite different characteristics. For complexes of a given acceptor with a series of donors, A0 increases markedly as the ionization potential of the donor decreases (see Tables 13- 15). The trend is maintained at various temperatures.(233) The lgF chemical shifts are known to be much more dependent on electron density than ‘H shifts,(250) and it would seem‘ that at least part of the upfield shift comes from the transfer of charge in the, interaction. If this is the case, measurements of lgF in donor molecules should~ show downfield, or at least greatly reduced upfield, shifts. This has been confirmed by qualitative measurements of the lgF resonance in 4-fluoraniline in the presence of excess 1,3,5trinitrobenzene in 1 ,2-dichloroethane.(22g) The point is well illustrated in complexes of 4-fluorobenzonitrile acting as a donor with boron halide acceptors. (251) For the very stable complexes (K, 2 5000) with the acceptors B2C14, BC13, BBr,, the chemical exchange was sufficiently low (i.e. there was a significant activation energy) to cause the appearance of separate lines for the free and complexed species. For the weaker complexes with the acceptors B(CH3)3, BF3 and p- FCGH4BC12, lines corresponding to the population average of free and complexed species were observed. The lgF resonance in the donor always


moved downfield on complex formation. In the case of the interaction XXIX the ‘“F resonance of the donor moved downfield on complex formation (A0 = -10.9 ppm) whilst the lgF resonance in the acceptor moved upfield by about the same amount (A, = +I2 -t 1 ppm) (see Fig. 20).

Acceptor (6-ve) Donor (6 + ve 1

-220- / , I / I I

10 20 30 40 50 60 70 80 %

mole ocld/mole base

FIG. 20. Double-label investigation of donor-acceptor complexing between p- fluorophenylboron dichloride and p-fluorobenzonitrile in dichloromethane at 25”; theoretical curves based on K$” = 6.0510.6 M-l and indicated limiting shifts.

(R. W. Taft and J. W. Carten.““l’)


3.3.4. Systems of biological interest

67

There has been considerable discussion concerning the possible role of charge-transfer complexes in biological systems.(8~25’252--62) and many suggestions have been made in specific instances where this might occur. It seems reasonable to suppose that if, for example, a powerful donor is involved, and the biological receptor contains potentially electron- attracting sites, then some form of charge-transfer complexing might be involved, although the contribution to the overall binding of the drug- substrate complex might be relatively small compared with other forces such as hydrogen bonding. Several systems have been investigated from the point of view of determining the characteristics of the charge-transfer interaction of physiologically active donor species by studying their behaviour with non-physiologically active acceptors.

(a) Phenothiazines. Phenothiazines (XXX) are known(257’263) to have

low ionization potentials and when mixed with electron acceptors such as 1,3,Strinitrobenzene or chloranil show optical absorptions characteristic of charge-transfer complexes. (25,258,264) Allison and Nash(265)

have suggested that a combination of charge-transfer complexing from the phenothiazine moiety and hydrogen bonding by the amino group is a possible means of attachment of the phenothiazine-based drug chlorpromazine at a biological receptor site.

Chemical shift measurements of the relative complexing abilities of lo-methylphenothiazine and a series of phenothiazine drugs (as the free bases) with 1,4_dinitrobenzene in chloroform and in carbon tetrachloride have been made.‘221’ Association constants were calculated by an application of equation (3 1). Appropriate plots are given in Figs. 21 and 22 and the results are listed in Table 20. The general correspondence between Fig. 21 and Fig. 22, and the resulting values of KAD and A,, in Table 20, suggest that any specific solvent interaction must either be insignificant or be common to all complexes.

Comparison of the values of K AD for various drug donors with that for N,N-dimethylaniline shows the former to be considerably stronger complexing agents in this type of interaction. despite the fact that N.N-dimethylaniline is itself recognized to be a strong electron


2o10 I I I

20 30 40

A c/s

FIG. 21. Plots of A/[D], against A for the interactions of: (A) chlorpromazine; (B) promethazine; (C) diethazine; (D) ethopropazine; (E) promazine; (F) trimeprazine; (G) 1 0-methylphenothiazine, with 1,4-dinitrobenzene in chloroform at 33.5”.

(R. Foster and C. A. Fyfe.(221))

donor.@) Amongst the values obtained for the complexes in chloroform solution, it is seen that lo-methylphenothiazine has a larger KAD value than any of the drug molecules (see Table 20). Since all these drugs have long side chains attached at position 10, this observation suggests that the reduction in KAD may be the result of a primary steric effect such as has been discussed in Section 3.3.2. The lowering of KAD from promazine to chlorpromazine is probably the result of the high electronegativity of the chlorine atom.

(b) Zndoles. A similar study has been made of indole and substituted indoles,(225) models for physiologically-active compounds such as tryp- tophan, serotonin, bufotenine and psilocine. The results supported previous optical measurements which had suggested that indoles could form charge-transfer complexes with electron acceptors.(254-6.266-7) Results for complexes with 1,3,5_trinitrobenzene in chloroform solution


60

FIG. 22. Plots of A/[D],, against the interactions of: (A) chlorpromazine; (B) promethazine; (C) diethazine; (D) ethopropazine; (E) promazine; (F) trimeprazine, with 1,4-dinitrobenzene in carbon tetrachloride at 33.5”. (R. Foster and C. A. Fyfe.(221))

are given in Table 22. It is seen that indole forms a more stable complex than N,N-dimethylaniline (see Table 19). The introduction of methyl groups into the 2 and 3 positions in indole greatly increases KAD, whereas methylation in the 7 position does not have such a marked increase. The introduction of groups larger than methyl at the 3 position causes a lowering of KAD even though the substituent may significantly increase the electron density of the m-system. There can obviously be no centro- symmetrical arrangement of the molecules in a complex with indole, and the results may indicate some tendency for the association to be orientated towards the hetero-ring. Such a suggestion had previously been made for an indole-iodine complex.@54-6)

In solutions containing indole and 1,3 ,Qinitrobenzene, with the latter in excess, the H3 proton shift moves much more than the aromatic ring proton shifts (see Table 22). If the main contributing factor to the shift


TABLE 22. VALUES OF ASSOCIATION CONSTANTS

WAD) AND RELATIVE CHEMICAL SHIFTS (A,)

FROM MEASUREMENTS OF VARIOUS ‘H SHIFTS IN THE COMPLEX INDOLE-I ,~.~-TR~NITR~~JEN~E~~E IN I.2-DICHLOROETHANE AT 33.5” (R. Foster and

C. A. Fyfe.‘225’)

Proton measured KAD A,?

H, in indole

H,,,,, in indole” H in 1,3,Strinitrobenzene

“At 60.004 MC/S. maximum height.

b Measured for the band of

positions is the ring current effect (see Section 3.3.3) these shifts could be explained by the preferential orientation of the acceptor towards the heterocyclic ring of indole. Recent molecular orbital calculations(31~zfi8) using the frontier electron density principle(26”-70) in conjunction with experimental data(267) for vgX are in harmony with such a structure.

3.4. Other Physical Methods of Measuring Association Constants of Charge-transfer Complexes in Dilute Solution

A variety of other physical methods has been used for the experimental evaluation of KAD. They include: partition of one component of the complex between two immiscible solvents,‘271--2) solubility’273-5’ and vapour pressure measurements(276-~) including gas-liquid chromatography,‘278-*~) polarography,‘2R5-6’ kinetic measurements of the further irreversible chemical reactions of the components and/or the charge-transfer complex, (287-O) calorimetry,‘2OO-l) ultraviolet and visible spectrophotometry, (234,292-300) intensity measurements of infraldd absorption bands,‘301-s’ and of Raman bands. (306) It is not possible to review here the various techniques involved in these approaches or to detail the results which have been obtained. Nevertheless, it is necessary to give some


indication of the methods and results obtained by the most widely used of these techniques, namely ultraviolet and visible absorption spectrophotometry of the intermolecular charge-transfer band of the complex, because of certain anomalous results which have been obtained by such procedures.

In these optical determinations of association constant it is assumed that the intensity of absorption measured by the absorbance A (optical density) is proportional to the concentration of the complex, the value of A having been corrected for absorption due to free A, free D and solvent. Then absorbance measured as log,, (III,), the logarithm of the rates of the intensity of the incident and transmitted radiation at the wavelength of measurement. will be for a path-length of 1 cm:

Ah = l tD[AD] (33)

where efD is the molar absorptivity (molar extinction coefficient) of the species AD at the wavelength of measurement (h). For solutions where one component, say D, is in large excess, in which only a single 1 : 1 complex AD is formed, and in which the various species behave ideally:

KAD = [ADI c

[Dlo([Al, - [ADI)

on a molarity scale from equations (29) and (33):

(29)

(34)

where [A], and [D10 represent the “weighed out” concentrations of A and D respectively. Equation (34) is generally known as the Benesi-Hilde- brand equation.“’ A plot of [A],/A, against l/[D], should be linear. The gradient of the line is equal to 1 I( K,AD E fD ) and the intercept with the ordinate equal to 11~;~. Unfortunately this intercept requires an extrapolation to high concentrations of D, and its numerical value is small. A similar criticism of the equivalent nuclear magnetic resonance expression (equation (30)) has already been made (see Section 3.2). Alternative represen- tations of equation (34) which avoid this type of extrapolation have been used. For example:‘234’

[AloP _ 1 ) UXJ Ah KAD AD

c Eh AD’ Eh

(35)


From a plot of [A],,[D],/A, vs. [D&,, K,AD and efD may be obtained. Alternatively:(2g4)

A* - = -K,AD Ah + K,AD [Alo EfD. @lo

A plot of AI[D], against A,, should be linear, the gradient of the line being equal to -KtD. This equation is the optical analogue of equation (3 1).

Alternative expressions have been derived for systems where the components themselves show significant absorption at the wavelength of measurement.(2g3) Other expressions have been derived(2gs,307-8) for determinations on solutions where the condition that one component species of the complex is not in large excess obtains (including the condition where [A], = [Dlo).

Comparison of the values of K AD obtained for a given system show that in some cases, for example iodine-n-electron donor complexes there is good agreement between values of K obtained from measurements on the charge-transfer band of the complex(30g) and the values obtained from intensity measurements of the iodine band. (310) However, for at least some complexes involving n-donors and m-acceptors there are some remarkable inconsistencies amongst values of association obtained from optical data.‘311’ For example, the apparent value of the association constant obtained from measurements of the intensity of the charge-transfer band for the complex between 1,3,5trinitrobenzene and N,N-dimethylaniline in chloroform is dependent on the concentration condition [Dlo s [A],,, [Dlo = [Alo or [Dlo 4 [Alo (see Table 23).‘311’ This behaviour was first noted by Ross and Labes. (307) They suggested that the differences could be accounted for by the presence of significant concentrations of the termolecular species AD, and A,D in solutions where [Alo + [Dlo and [A], + [Dlo respectively. However, more recent datac311) suggest that if this were the explanation of the difference and if, under the condition [A], = [D], only the species AD is significant, then the association constant for termolecular complex formation would be unrealistically large in some cases. For yet other systems the experimental data lead to the impossible conclusion that the formation constant of the termolecular species is negative. Another anomaly within the optical determinations is observed in some complexes which have an optical charge-transfer absorption which has more than a single maximum. Determinations of the association constant for such systems from absorbance measurements made over a range of wavelengths can show a strong wavelength dependence if the condition [D],, % [A], is used, as is normally the case.(312-13) By contrast, determinations made under the experimental condition [Dlo = [A], yield values of K,AD which are wavelength indepen-


TABLE 23. COMPARISON OF DATA FOR CHARGE-TRANSFER EQUI- LIBRIA FROM OFWCAL MEASUREMENTS AND FROM CHEMICAL SHIFT MEAWREMENTS UNDER VARIOUS CONCENTRATION CONDITIONS AT 33.5” (P. H. Emslie, R. Foster, C. A. Fyfe and I. Horman;‘311’

C. A. Fyfe’229’)

Method Rel. cont. / ;t; / EX10-3 / K&X10-3

N,N-dimethylaniline-1,3,5-trinitrobenzene in CHCI,

N,N,N’,N’-tetramethyl-p-phenylenediamine-l,3,5-trinitrobenzene in CH,CICH,CI

Optical Optical Optical NMR

[Al, * [Dlo I .9, 1.1, [Al, Q [Dlo 1.4, 1.6, [Al, = @lo 1% 2.2, [Al, +z tD1, 1.0, -

Hexamethylbenzene-fuoranil in Ccl4

2.31 2.3* 2.3,

Optical [Al, * WI, 16.5 2.9, 49.4 Optical IAl, = [Dlo 11.5 4.2, 49.4 NMR [Al, + [Dl, 9.7 -

aK:D values calculated on the assumption of ideal behaviour.

dent.‘312) From independent observations on similar systems Johnson and BowenC313) suggested that, under the condition [DIO % [A],, significant concentrations of the termolecular species AD, could account for the apparent anomaly. Synthetic data have been used to demonstrate the effect of such complexes.‘313’ However, relatively large values for the constants of formation of termolecular complexes and/or large values of extinction coefficient for the termolecular species have to be used in some cases if the experimental results are to be explained in these terms.‘22g)

3.5. A Comparison between Association Constants Evaluatedfrom

Chemical Shifts andfrom Electronic Spectral Data

Because of the conflicting values of association constant obtained purely from optical data (referred to in Section 3.4), it is of interest to compare such results with values of association constants obtainable from


chemical shift data where this is possible. Some of these comparisons are shown in Table 23. The remarkable feature is that, in all the cases where measurements have been made, where there are significant differences in the optical determinations under the various concentration conditions, agreement with the value obtained from chemical shift data is only obtained when the optical determinations are made under the condition [D], = [A],,. The fact that the chemical shift data are obtained under other concentration conditions suggests that these differences are not the result of the presence of termolecular complexes. This conclusion is further supported by the results given in Table 6, which indicates that, in these cases at least, the evaluation of the 1 : 1 association constants is independent of the condition [D10 % [A],, or [D],, G [A],,.

Various possible explanations of these differences are now considered: (a) Presence of termolecular complexes. The effect of the presence of

(hypothetical) termolecular complexes on the evaluation of association constants from chemical shift data may be estimated. Let us consider a series of solutions in which [D],, + [A],,. Let P,, PAD, PAD2 be the fractions of the total population of A in the form of the free species, in the form of the complex AD and in the form of the complex AD, respectively. Let

8A21 SAD and aAD. be the chemical shifts of each of these species if each alone were the sole solute. The measured chemical shift (6) of a given nucleus in the acceptor relative to some standard will be:

6 = P.4SA + PAD SAD + PAD2 SAD2. (37)

If K:D = [AD]/[A][D] and K$D2 = [AD,]/[AD][D], then for solutions where [DJO + [Al,,:

- KPD (It [DloKtD’)A + I+ (38)

where A = (6, - 6) ; A$” = (6, - 6,D) ; AiD* = (6, - 6.&.

Equation (38) corresponds to equation (3 1) and becomes identical with it when KtDz = 0. Synthetic data’22”’ suggest that in fact K,AD would have to be very small to account for the high linear correlation such as is observed in Fig. 13. In none of the systems measured could KtDZ exceed K$D/10 and in the majority of the systems it would have to be less than - K,AD/20. The postulate that significant concentrations of termolecular species are present in such solutions would therefore appear to be untenable.

(b) Non-idea& of the so1utions.t Let us suppose that on some con-

tThe Authors are grateful to Dr. P. G. Wright (University of Dundee) for his comments and suggestions on matters in this section.


centration scale, the activity coefficients yA, yn, yAD appropriate to this concentration scale of the species A, D and AD respectively are such that the quotient yAD/yA yD is not unity. The equilibrium constant KAD may be written:

K*D = [ADI YAD ~ -

[Al PI * YAYD ’ (39)

From equations (28) and (39) it is seen that

or

(40)

. (41)

The high correlation of experimental points to a linear relationship between A/[D], and A implies that: either the quotient YAD/YAYD is unity over the concentration used; or YAD/YAYD is a linear function of [Dlo. If we write:

zD= 1 +BIDIO, (42)

where B is independent of [D],,, but could vary with [A],,, then equation (3 1) becomes

A - = - ( KAD + B)A + KAD&. @lo

(43)

This would lead to an “incorrect” value (K4D+ B) for the association constant and for Ao, although the product K”DA, would be “correct”. The “incorrect” value of (K”” + B) corresponds to Guggenheim’s’“14’ “sociation constant”.

By the same arguments the optical relationships corresponding to equation (34) and equation (36) for such non-ideal solutions are:

and

[AI,= 1 1 flIKAD+B ___- A* E;DKAD. [D], ,fD\ KAD (44)

[Al,, -=- Ah

(KAD+B)Ah+KAD~fDIA]O. (45)

This would therefore also lead to an “incorrect” value for the association constant of (KAD + B) but a “correct” value for the product (KAD . l ).


Therefore although both methods may lead to a “sociation constant” rather than a true thermodynamic association constant, the differences between the measured values of the formation constant depending on whether the chemical shift method or the charge-transfer band optical method is used cannot be explained in terms of deviations from ideality.

(c) Specific s&don. Carter et al. c315) have suggested that the solute species have well-defined solvent shells so that the equilibrium previously expressed as:

A +D=AD (25) should be written as:

AS, + DS, = mSp+qS, (46)

where n, m and p are the numbers of solvent molecules (S) in association with the species A, D and AD respectively, and (n + m) = (p + q). For such an equilibrium it may be shown that the experimental values of the association constant K,,, are related to values of KAD from optical data:

K exp = KAD+q(m+ l)/C% (47)

where E,,~ is the experimentally obtained molar absorptivity and [S], is the total solvent concentration when [D],, = 0. Such specific solvation may well occur in systems in which the solvent is a strongly interacting compound, for example, an alcohol or an ether. However, specific solvation would affect the value of association constant determined from chemical shift data in a similar manner. It therefore cannot provide an explanation of the differences between the optical and nuclear magnetic resonance determinations such as those given in Table 23.

(d) Apparent deviations from Beer’s ZLZW.‘~~~’ From the discussion above it is difficult to avoid the conclusion that the differences exempli- fied in Table 23 are the result of some false assumption in the optical determinations, where one or other component is in excess, which assumption does not have to be made in the corresponding chemical shift determinations. The one assumption which involves only the optical determinations based on the charge-transfer absorption is that the measured absorbance is proportional to the concentration of the complex AD (when allowance is made for the absorption by other species present), i.e. it is assumed that Beer’s law for AD is obeyed. Some years ago Barbc317’ considered the possibility of a deviation from Beer’s law. More recently the implications of a concentration dependence of .$” have been discussed.‘311’ The method of Liptay’2gs) for evaluating KAD was claimed to demonstrate the concentration independence of l fD. However, it has


been pointed ouU318) that a satisfactory evaluation by this method only implies that the shape of the absorption curve is the same for all solutions. Furthermore, SOme sets of solutions cannot be evaluated by Liptay’s method because they do show a significant variation of band shape with concentration (see Section 3.4).

It may be that, owing to the weak binding between the donor and acceptor moieties in the complex, the transition dipole (see Section 1.2) is affected by a non-specific clustering of molecules of the species in excess around complexed DA pairs in solutions containing a large excess of one component, say D. This effect would be dependent on [Dlo. This is consistent with the observation that for solutions where [D10 = [A],, where neither concentration is large, nor does either vary so much as in systems where one component is in large excess, no anomaly is observed. An explanation in these terms is still conjectural. Obviously further work on this problem is required.

3.6. Evaluation of Optical Absorptivity of Charge-transfer Complexes from a Combination of Chemical Shift and Optical Data

The majority of molar absorptivity values for charge-transfer complexes corresponding to the maximum of the charge-transfer band (or bands), namely z&tX, have been obtained purely from optical data using the Benesi-Hildebrand or similar equations.

For a set of complexes of a given acceptor with a series of structurally related donors in which there are no large steric effects, the intensity of the band is expected to increase as the donor-acceptor interaction increases.(*) The theoretical argument for this has already been given in Section 1.2. In fact for many such series the intensity, as measured by e&r&, shows the opposite trend. (~3 315-16) This observation, among others, has featured in the theory of contact charge-transfer by Orgel and Mulliken,‘26’ of energy borrowing from the excited states of the donor or the acceptor by the charge-transfer transition,‘316’ and of the specific solvation hypothesis of Murrell and his co-workers.(315)

Orgel and Mulliken’26’ also showed that if isomeric 1: 1 complexes were formed (see Section 3.1) then in a Benesi-Hildebrand type of determination, the experimental value of the molar absorbance (eexp) would be a weighted average for the various complexes:

where KPD is the association constant for the ith 1: 1 complex and l fD is the molar absorptivity of the ith 1 : 1 complex at the wavelength of

(49)


measurement. It is difficult to prove or to disprove the presence of more than a single 1: 1 complex in solution. The experimental evidence, such as it is, does not provide positive evidence for such isomers.‘“Os’

One interesting observation noted by several groups of workers’“OT, ‘111,319) is that although different conditions may yield different values of K”” and l tD for a given system at a given wavelength, the product (KADeiD) is constant (see, for example, Table 23). Indeed, a large part of the problem of the analysis of optical data centres around the problem of the separation of these two terms. If it is assumed that the optical data yield a correct value for this product (KADefD), and the chemical shift data provide reliable estimates of KAD, extinction coefficients may be calculated’219’ from the quotient (K”“~~“)~~~~~~~/(K”“)~~~. A set of EE, values calculated in this fashion is given in Table 24. Comparison with &,E values obtained from Benesi-Hildebrand type experiments from optical data show that in this series the values for E$& are altered significantly. The values derived using ( KAD)NMR increase as the Lewis basicity increases in the manner originally anticipated by Mulliken.‘“’ It should be emphasized that although this may be the true order of eki, in the series cited, there may be other reasons why such an order will not always be observed even if the correct values of &,& are used. Thus && is very sensitive to steric interference by the close approach of suitably orientated AD pairs, for example in acceptor-hexaethylbenzene complexes. (226,236.320-l)

TABLE 24. VALUES OF K$%(~PTIcAL); KAgMR,; l (OPTICAL); AND E’ = {K$ E (OPTICAL)/K~!&,~+} FOR COMPLEXES OF METHYLATED BENZENES WITH 1,3,5- TRINITROBENZENE IN CAREZONTETRACHLORIDE AT 33.5”(P. H. Emslie.R.Foster.

C. A. Fyfe and I. Herman’:““)

KfD. l X 1O-3 (optical)

1 K%m 1 e (optical) 1 l ’

Mesitylene I .74 0.79, 3000 2180 Durene 3.33 1.32, 2390 2520 Pentamethylbenzene 5.69 1.93, 2440 2950 Hexamethylbenzene 9.92 3.30, 2500 3100

“Optical maxima for the benzene. toluene and xylene complexes could not be resolved.

4. NUCLEAR MAGNETIC RESONANCE AND NUCLEAR QUADRUPOLE RESONANCE STUDIES OF SOLID

CHARGE-TRANSFER COMPLEXES

Nuclear magnetic resonance measurements on charge-transfer complexes in the solid phase have been used to investigate the problem of whether the component molecules are rigidly held in the crystal lattice,


or whether they can undergo rotation. Gilson and McDowe11(322’ have measured the proton magnetic resonance of the solid benzene-silver perchlorate complex at 298°K and 77°K at 40 MC/S. From the spectra they obtained values of 0.9 G* (298°K) and 6.10 G* (77°K) for the second moments of these systems. From the structure of the complex derived from X-ray diffraction by Smith and Rundle,‘323’ Gilson and McDowell calculated a second moment of 5.9 G2 for the complex. The agreement between this value and the measured value at 77°K led them

Temperature, “K

FIG. 23. Temperature dependence of the line width and second moment of the proton resonance in the complex C,H,.C,F,. Calculated line width dependences obtain

from the following parameters. (D. F. R. Gilson and C. A. McDowell.‘~27’)

( B (gauss) 1 E (kcal mole-‘) ( q,(c/s)

Curve A 8.0 1.9 2.3 x 10” Curve B 9.0 1.7 1.2 x 10” Curve C 10.0 1.6 1.2 x 109


100 200

Temperoture, “K

FIG. 24. Temperature dependence of the line width and second moment of the fluorine resonance in the complex C,H,. C,F,. Calculated line width dependences obtained from the folIowing parameters. (D. F. R. Gilson and C. A. McDowell.‘327))

E (kcal mole-‘) avo (c/s)

Broad line 1.5 1.0 x 107 Narrow line 1.3 1.1 x 106

to conclude that at this temperature the molecule is more or less rigidly held in the crystal lattice. For “free” rotation about the sixfold axis, the intermolecular contribution to the calculated second moment is greatly reduced. Under. these conditions an estimated total second moment of l-01 GZ was obtained, in good agreement with the experimental value at 298°K. From this, Gilson and McDowe11~322~ concluded that at room temperature the benzene molecule is rotating more or less freely about the sixfold axis; however, they pointed out that the nuclear magnetic resonance will give the appearance of rotation if reorientation occurs at


frequencies of the order 1 04- lo5 c/s or greater, whereas the X-ray diffraction data will show a stationary molecule if the molecule spends most of its time at the bottom of a potential energy well for the rotational motion. Gilson and McDowe11’325’ also suggest that if the forces were great enough to distort the benzene ring (a proposition made by Smith and Rundle), then it is likely that these forces are sufficient to hold the molecule in a rigid lattice. However, Smith (324) still maintained that the crys- tallographic evidence does not support a freely rotating model for the structure at room temperature. Gilson and McDowell have quoted the infrared evidence of Daasch’326’ in support of their structure.

Gilson and McDowe11’327’ have extended their work to many other complexes, including the hexafluorobenzene complexes of benzene, mesitylene and durene, for which they observed both the lH and lgF resonances (Figs. 23 and 24). The tetrabromomethane complexes of benzene and durene were also studied. The results again suggest that at room temperature there is molecular rotation. Estimates of the barriers to rotation were obtained from the temperature dependence of the line- widths. For benzene in the benzene-hexafluorobenzene complex the barrier to rotation ranged from 1.6 to 1.9 kcal mole-l. For the hexafluorobenzene molecule the barriers were calculated from both the broad and narrow components of the resonance line: values of 1.3 and 1.5 kcal mole-’ respectively were obtained. For the benzene-carbon tetrabromide complex the barrier was estimated to be within the range 2.1-2.3 kcal mole-‘, the actual value being dependent on the value chosen for the rigid- lattice line width. A similar investigation of molecular motion in the solid hexamethyfbenzene-chloranil complex led Anderson’328’ to conclude that there is rapid molecular reorientation of hexamethylbenzene about the sixfold axis. The fact that the nuclear quadrupole resonance of chlorine in the same complex shows two resonances(32g) suggests that the chloranil molecules reorientate very slowly, if at all. Douglass’“2g’ has also estimated that, us an upper limit, there is only a 5-10% transference of charge in the ground state of this complex. This suggestion is based on the assumption that the nuclear quadrupole moment of the chlorine depends on the ionic character of the C-Cl bond, and thus on the electron density on the carbon atoms. Although this result depends mainly on estimates of extraneous crystal field effects, nevertheless, the results are in harmony with the earlier estimates of ionic character in the complex based on dipole measurements.(244)

Douglass’32Z” observed no 35Cl nuclear quadrupole resonance absorption for the complexes picryl chloride-hexamethylbenzene andp-chloroaniline -1,35trinitrobenzene. He suggested that this might be the result of disordered structures in the solids.


Attempts to rationalize the shifts in the XIBr nuclear quadrupole resonance between pure component and sohd complex have been made for various bromide-aromatic solid complexes’“:“‘-l) Other measurements have been made on solid complexes involving inorganic electron acceptors which are outside the scope of this present chapter.‘““2p”‘.

ACKNOWLEDGEMENTS

The following figures have been reprinted by permission of their authors and the copyright owners. Figure 4, reprinted from J. Chem. Phys. 29,

253 (1953) by the American Institute of Physics. Figure 5, reprinted from J. Chem. Sac. 1 I23 (1962); Figs. 15 and 17, reprinted fromJ. Chem.

Sac., (B), 406 (1967). all by the Chemical Society (London). Figure 6, reprinted from J. Phys. Chem. 67, 2 190 (1963); Fig. 7, reprinted from J. Phys. Chem. 67, I182 (1963); Fig. 8, reprinted from J. Amer. Chem. Sac. 85, 2173 (I 963); Fig. 9, reprinted from J. Phys. Chem.

69, 89 (1965); Fig. 10, reprinted from J. Amer. Chem. Sot. 87, 3620

(1965): Fig. 12, reprinted from J. Phys. Chem. 68, 8 1 1 (1964); Fig. 20, reprinted from J. Amer. Chem. Sac. 86, 4199 (1964). all by The American Chemical Society. Figure 16. reprinted from Truns. Faraday

Sac. 62, 1400 (1966) by the Faraday Society. Figures 2 I and 22, reprinted from Biochem. et Biophys. Acta, 112, 490 ( 1966). by Elsevier Publishing Company (Amsterdam). Figures 23 and 24, reprinted from Can. J. Chem.

44,945 ( 1966) by the Chemical Institute of Canada.

REFERENCES

1. 2. 3. 4. 5.

6. 7. 8. 9.

10.

I I.

12. 13. 14.

H. A. BENESI and J. H. HILDEBRAND. J. Amer. Chem. Sec. 71,2703 (1949). H. TSJBOMURA, J. Amer. Chrm. Sot. 82,40 (I 960). R. L. HANSEN. .I. Phys. Chem. 70, 1646 (1966). R. M. KEEFER and T. L. ALLEN,J. Chem. Phys. 25,1059 (1956). M. M. DE MAINE. P. A. D. DE MAINE and G. E. MCALONIE. J. Mol. Spectroscopy, 4.27 1 (1960). J. N. MURRELL and S. CARTER.J. Chem. SW. 6 185 (1964).

P. PFEIFFER. Orgcrnisc~he Mo/ekii/r;erbindunRen. 2nd ed.. Enke. Stuttgart, 1927. R. S. MIJLLIKEN,J. Pkys. Chem. 56,801 (1952). G. BRIEGLEB. Elektrorrrrr-Doncrtor-A~c.eptor-Ku777ple.xe. Springer Verlag. Berlin. 1961. L. J. ANDREWS and R. M. KEEFER, Molecultrr Compiexrs in Or,qunic Chrmistrv, Holden-Day, San Francisco, 1964. W. B. PERSON and R. S. MUI.LIKEN. Molecrrlor Complexes: A Lecture trnd Reorinr

Volume, Wiley. New York (to be published). L. J. ANDREW% Chem. Rev. 54,J I3 (1954). J. A. A. KETELAAR. Chem. Weekblad. 52,218 (1956). L. J. ANDREWS and R. M. KEEFtR,AdvufI. Inorg. Chem. Rudiochem. 3.9 I ( I96 1).


15. R. S. MULLIKEN and W. B. PERSON.A~~. Rev. Phys. Chem. 13, 107 (1962).

16. G. CAUQUIS and J. J. BASSELIER, Ann. Chim. (France) 7,745 (1962). 17. V. P. PARINI, Uspekhi. Chim. 31,822 (1962); Russ. Chem. Rev. 31,408 (1962). 18. S. P. MCCLYNN, Chem. Rev. 58, 1113 (1958).

19. J. R. PLATT,A~~. Reu. Phys. Chem. 10,367 (1959). 20. W. C. PRIcE,A~I. Rev. Phys. Chem. 11, 140 (1960). 21. J. N. MURRELL, Ouart. Rev. 15, 19 1 (196 1).

22. G. BRIEGLEB andJ. CZEKALLA,AngeH’. Chem. 72,401 (1960). 23. G. BRIEOLEB, PureAppl. Chem. 4, 105 (1962). 24. E. M. KOSOWER, Progress in Physical Organic Chemistry, Ed. S. G. COHEN. A.

STREITWIESER~~~ R. W. TAFT, Vol. 3, Interscience, New York, 1965, p. 8 1. 25. A. SZENT-GYORGYI. Introduction to a Submolecular Biology, Academic Press, 1960. 26. L. E. ORGEL and R. S. MULLIKEN.J.Amer. Chem. Sot. 79.4839 (1957). 27. M. J. S. DEWAR and A. R. LEPLEY. J. Amer. Chem. Sot. 83.4560 ( I96 I).

28. A. R. LEPLEY, J. Amer. Chem. Soc.84,3577 (1962). 4 29. M. NEPRAS and R. ZAHRADN~K. Tetrahedron Letters. 57 (1963): Co/l. Czech. Chem.

Comm. 29, 1545, 1555 (1964). 30. D. C. MuIcHEIuEE and A. K. CHANDRA, J. Phys. Chem. 68,477 (1964).

3 1. J. P. GREEN and J. P. MALRIEU, Proc. Nat. Acad. Sci. 54,659 (1965). 32. R. L. FLURRY. JR., J.Phys. Chem. 69, 1927 (1965). 33. T. A. ROURKE, unpublished work. 34. S. IWATA, J. TANAKA and S. NAGAKURA, J. Amer. Chem. Sot. 88,894 (1966). 35. H. S. GUTOWSKY, D. W. MCCALL and C. P. SLICHTER, J. Chem. Phys. 21,279 (1953).

36. H. S. GUTOWSKY and A. SAIKA, J. Chem. Phys. 21,1688 (1953). 37. H. S. Gu-rowsKY and C. H. HoLM,J. Chem. Phys. 25.1228 (1956). 38. E. GRUNWALD, A. LOEWENSTEIN and S. MEIBOOM. .I. Chem. Phys. 27. 630, 641

(1957). 39. A. LOEWENSTEIN and S. MEIBOOM. 1. Chem. Phys. 27, 1067 (1957).

40. M. TAKEDA and E. 0. STEJSKAL, J. Amer. Chem. Sot. 82,25 (1960). 41. H. M. MCCONNELL, J. Chem. Phys. 28,430 (1958). 42. J. A. POPLE, W. G. SCHNEIDER and H. J. BERNSTEIN. High Resolution Nuclear

Magnetic Resonance, McGraw-Hill, New York, 1959, p. 222. 43. Ref. 42, p. 38 1. 44. R. F. M. WHITE. Phy.~ical Methods in Heterocyclic Chemistry. Ed. A. R. KATRITZKY.

Vol. 2. Academic Press. Londong, 1963. chap. 9.

45. E. F. CALDIN, Fast Reactions in-Solution, Biackwell, Oxford. 1964, chap. 1 1. 46. L. W. REEVES, Advances in Physical Organic Chemistry, Ed. V. GOLD. Vol. 3.

47

48 49. 50.

51. 52.

53. 54.

55. 56.

57.

58.

59. 60.

Academic Press, London, 1965, pi. 187. ”

J. W. EMSLEY, J. FEENEY and L. H. SUTCLIFFE, High Resolution Nuclear Magnetic Resonance Spectroscopy, Vol. 1, Pergamon Press, Oxford, 1965.

J. T. ARNOLD and M. E. PACKARD, Phys. Rev. 83.2 IO (195 I). J. T. ARNOLD and M. E. PACKARD, J. Chem. Phys. 19, 1608 (195 1). U. LIDDEL and N. F. RAMSEY, J. Chem. Phys. 19,1608 (195 I).

A. D. COHEN and C. REID, J. Chem. Phys. 25.790 (I 956). C. N. HUGGINS, G. C. PIMENTEL and J. N. SHOOLERY. J. Phys. Chem. 60,13 11 (1956) E. D. BECKER, U. LIDDEL and J. N. SHOOLERY, J. Mol. Spectroscoov 2.1 (1958). J. C. DAVIS. JR., K. S. PITZER and C. N. R. RAO, J. Phys. Chem. 64; i744 (i 960): B. G. SOMERS and H. S. GUTOWSKY. J. Amer. Chem. Sot. 85,3065 (1963). V. F. BYSTROV, K. M. DYUMAEV, V. P. LEZINA and G. A. NIKIFOROV. Doklady Akad. NaukSSSR 148, 1077 (1963). V. F. BYSTROV and V. P. LEZINA, Opitika i Spectroskopiya 16, lCO4 (1964): Optics and Spectroscopy 16,542 (1964). V. F. BYSTROV, V. V. ERSHOV and V. P. LEZINA. Opitika i Spectroskopiya 17, 538 (1964); Optics and Spectroscopy 17.290 (1964). M. SAUNDERS and J. B. HYNE. J. Chem. Phys. 29,253 (1958). M. SAUNDERS and J. B. HYNE, J. Chem. Phys. 29. 13 19 (1958).


61. P. W. BRIDGMAN, Dimensional Analysis, Yale University Press, New Haven, 1922, p. 40.

62. E. D. BECKER, .I. Chem. Phys. 31,269 (1959). 63. J. FEENEY and S. M. WALKER,J. Chem. Sot. (A), 1148 (1966).

64. J. MARTIN, Compt. Rend. 256d, 365 1 (1963). 65. M. MARTIN and M. QUILBEUF, Compf. Rend. 252d, 4151 (1961).

66. M. MARTIN, J. Chim. Phys. 59,736 (1962). 67. G. MAVEL. Compt. Rend. 248d, 1505 (1959).

68. G. MAVEL. J. Phys. Radium 21,37 (1960). 69. C. LUSSAN, B. LEMANCEAU and N. SOUTY, Compt. Rend. 254d, 1980 (1962). 70. J. R. HOLMES. D. KIVELSON and W. C. DRINKARD, J. Amer. Chem. Sot. 84, 4677

(1962). 71. A. FRATIELLO and D. C. DOUGLASS,J. Mol. Spectroscopy 11,465 (1963). 72. F. TAKAHASHI and N. C. LI, .I. Phys. Chem. 68,2136 (1964). 73. F. TAKAHASHI and N. C. LI. J. Phys. Chem. 69,1622 (1965). 74. F. TAKAHASHI and N. C. L1.J. Amer. Chem. Sot. 88,1117 (1966).

75. T. F. LIN, S. D. CHRISTIAN and H. E. AFFSPRUNG,J. Phys. Chem. 69.2980( 1965). 76. H. SAITB. K. NUKADA, H. KATO, T. YONEZAWA and K. FUKUI, Tetrahedron Letters

111 (1965). 77. J. FEENEY and L. H. SUTCLIFFE, Proc. Chem. Sot. 118 (1961).

78. .I. FEENEY and L. H. SUTCLIFFE, J. Chem. Sot. 1123 (1962). 79. J. R. CROOK and K. SCHUG, .I. Amer. Chem. Sot. 86,427l (1964). 80. V. F. BYSTROV and V. P. LEZINA, Opitiku i Spectroskopiyu, 16, 790 (1964); Optics

and Spectroscopy, 16,430 (I 964). 81. J. A. HAPPE, J. Phys. Chem. 6572 (1961). 82. C. S. SPRINGER, JR. and D. W. MEEK, J. Phys. Chem. 70,48 I (1966).

83. L. A. LAPLANCHE, H. B. THOMPSON and M. T. ROGERS, J. Phys. Chem. 69, 1482 (1965).

84. A. S. N. MURTHY. C. N. R. RAO, B. D. NAGESWARA RAO and P. VENKATESWARLU, Trans. Faraday Sot. 58,855 (1962).

85. B. D. NAGESWARA RAO, P. VENKATESWARLU, A. S. N. MURTHY and C. N. R. RAO,

C’U~. J. Chem. 40,963 (1962). 86. S. H. MARCUS and S. I. M1LLER.J. Amer. Chem. Sot. 88,37 19 (I 966). 87. R. MATHUR, E. D. BECKER, R. B. BRADLEY and N. C. LI, J. Phys. Chem. 67, 2190

(1963).

88. R. MATHUR, S. M. WANG and N. C. LI, J. Phys. Chem. 68,2 140 (1964). 89. M. M. ROUSSELOT and M. MARTIN, Compt. Rend. 262c, 1445 (1966). 90. M. M. ROUSSELOT, Compt. Rend. 262c, 26 (1966). 9 1. M. M. ROUSSELOT, Compt. Rend. 262c, 49 (1966). 92. E. B. WHIPPLE, J. H. GOLDSTEIN, L. MANDELL, G. S. REDDY and G. R. MCCLURE,

J.Amer. Chem. Soc.81, 1321 (1959). 93. J. V. HATTON and R. E. RICHARDS, Trans. Famduy Sot. 57,28 (1961). 94. A. A. PETROV, V. B. LEBEDEV and 1. G. SAVICH, Zhur Obschchei Khim. 32, 658

(1962); J. Gen. Chem. U.S.S.R. 32,655 (1962). 95. A. A. PETROV and N. V. ELSAKOV, Zhur. Obschchei Khim. 33, 319 (1963); J. Gen.

Chem. lJ.S.S.R.33,314(1963). 96. N. V. ELSAKOV and A. A. PETROV, Opitika i Spectroskopiya 16, 148 (1964); Optics

und Spectroscopy 16,77 (1964). 97. N. V. ELSA~OV and A. A. PETROV. Opitiku i Spectroskopiyu 16, 797 (1964); Optics

and Spectroscopy, 16,434 (1964). 98. A. A. PETROV, N. V. ELSAKOV and V. B. LEBEDEV, Opitiku i Spectroskopiyu 16,lO 13

(1964); Optics und Spectroscopy 16,547 (1964). 99. A. A. PETROV, N. V. ELSAKOV and V. B. LEBEDEV. Opitikn i Spectroskopiya 17. 679

(1964); Optics and Spectroscopy 17,367 (1964). 100. A. A. PETROV, N. V. ELSAKOV, V. S. ZAVGORODNII and V. B. LEBEDEV, Teor. i

Ehsperim. Khim., Ahtrd. Ntruk Ukr. S.S.R. 1, 697 (1965).


101. M. M. KREEVOY, H. B. CHARMAN and D. R. VINARD, J. Amer. Chem. Sot. 83,

1978 (1961).

102. J. C. D. BRAND, G. EGLINTON and J. TYRRELL, J. Chem. Sot. 5914 (1965).

103. J. TYRRELL, Can. J. Chem. 43,785 (1965).

104. L. W. REEVES and W. G. SCHNEIDER, Can. J. Chem. 35,251 (1957). 105. W. G. SCHNEIDER. J. Phys. Chem. 66,2653 (1962). 106. R. J. ABRAHAM, J. Chem. Phys. 34, 1062 ( 1961). 107. R. J. ABRAHAM, Mol. Phys. 4,369 (1961). 108. T. MATSUO and Y. KODERA, J. Phys. Chem. 70,4087 ( 1966).

109. C. J. CRESWELL and A. L. ALLRED, J. Phys. Chem. 66,1469 (1962). 110. C. M. HUG~;INS, G. C. PIMENTEL and J. N. SHOOLERY, J. Chem. Phvs. 23, 1244

(1955). 1 I 1. G. J. KORINEK and W. G. SCHNEIDER, Can. J. Chem. 35. I 157 (1957).

112. C. M. HUGGINS and D. R. CARPENTER. J. Phys. Chem. 63.238 (1959).

113. Z. PAJAK and F. PELLAN, Compr. Rend. 251d, 79 (1960). 114. L. W. REEVES, E. A. ALLAN and K. 0. STROMME, Can. J. Chem. 38,1249 (1960). 115. C. M. HUGGINS,J. Phys. Chem. 65,188l (1961). 116. D. L. ANDERSEN, R. A. SMITH, D. B. MYERS, S. K. ALLEY, JR., A. G. WILLIAMSON

and R. L. SCOTT, J. Phys. Chem. 66,621 (1962).

117. S. K. ALLEY, JR. and R. L. SCOTT, J. Phys. Chem. 67,1182 (1963). 118. C. J. CRESWELL and A. L. ALLRED, J. Amer. Chem. Sac. 85, 1723 (I 963). 119. R. KAISER, Can. J. Chem. 41,430 (1963).

120. A. L. MCCLELLAN, S. W. NICKSIC and J. C. GUFFY, J. Mol. Spectroscopy 11, 340 (1963).

121. B. B. HOWARD, C. F. JUMPER and M. T. EMERSON, J. Mol. Spectroscopy 10, 117 (1963).

122. P. J. BERKELEY, JR., and M. W. HANNA, J. Phys. Chem. 67,846 (1963). 123. P. J. BERKELEY, JR., and M. W. HANNA, J. Chem. Phys. 41,253O (1964).

124. P. BISCARINI, L. LUNAZZI and F. TADDEI, Boll. Sci. Fat. Chim. Ind. Bologna 22. 67 (1964).

125. A. L. MCCLELLAN and S. W. NICKSIC, J. Phys. Chem. 69,446 (1965). 126. F. TAKAHASHI and N. C. LI, J. Phys. Chem. 69,295O (1965).

127. M. F. RETTICI and R. S. DRAGO, J. Amer. Chem. Sot. 88.2966 (I 966). 128. L. W. REEVEsandW.G. SCHNEIDER, Trans. Faraday Soc.54,314 (1958). 129. L. W. REEVES, Trans. Faraday Sot. 55, 1684 (1959). 130. L. W. REEVES, Can.J.Chem. 39,171 l(l961).

131. S. S.DHARMATTI,G.GOVIL.C. R. KANEKARandY. P.VIRMANI,A~~~. Sci. (Geneva), 13,479 (1960).

132. A. CAPALBI, C. FRANCONI, B. PESCE and M. L. TOSATO, Sot. Ital. Prog. Sci. Sci. Tec.5, 125 (1961).

133. G. MAVEL,~~P~. Poudres43, 130(1961).

134. A. PEROTTI and M. COLA, Guzz. Chim. Ital. 91, 1153 (196 1). 135. A. PARMIGIANI, A. PEROTTI and V. RIGANTI, Gazz. Chim. Ital. 91, 1148 (1961).

136. B.N. BHAR,Arkiv. Fysik. 12, 171 (1957). 137. N. MuLLERand P. I. ROSE, J.Amer. Chem. Soc.85,2173 (1963). 138. N. MuLLERand P. 1. RosE, J. Phys. Chem. 69,2564 (1965). 139. N. MULLER~~~ 0. R. HUGHES, J. Phys. Chem. 70,3975 (1966).

140. J. C. DAVIS, JR. and K. S. PITZER, J. Phys. Chem. 64,886 (1960). 141. T. C. CHIANG and R. M. HAMMAKER, J. Mol. Spectroscopy, 18, 110 (1965). 142. T. C. CHIAN~; and R. M. HAMMAKER, J. Phys. Chem. 69,27 15 (1965).

143. J. R. FERRARO and D. F. PEPPARD, J. Phys. Chem. 67,2639 (1963). 144. E. ODEBLAD. Arkiv. Kemi 2567 (1964). 145. G. ALLEN and E. F. CALDIN, Quart. Rev. (Lond.) 7,255 (1953). 146. R. A. CRAIG and R. E. RICHARDS, Trans. Faraday Sot. 59,1962 (1963). 147. M. OKADA, K. SUYAMA and Y. YAMASHITA. Tetrahedron Letters 2329 (1965). 148. M. OKADA, K. SUYAMA and Y. YAMASHITA. Kogyo Kagaka Zasshi (J. Chem. Sot.

Japan)@& 1431 (1965).


149. A. C. RUTENBERG, A. A. PALKO and J. S. DRURY, .I. Amer. Chum. Sot. 85, 2702 (1963).

150. A. C. RUTENBERG, A. A. PALKO~~~ J. S. DRURY,~. Phys. Chem. 68,976 (1964). 151. A:C. RUTENBERG and A. A. PALKO, J. Phys. Chem.69,527 (1965). 152. E. GORE and S. S. DANYLUK, J. Phys. Chem. 69,89 (1965). 153. A. FINCH. P. G. GATES and D. S. STEELE, Trans. Faraday Sot. 61,2623 (1965). 154. J. V. HATTON and R. E. RICHARDS, Mol. Phys. 3.253 (1960). 155. J. V. HATTON and R. E. RICHARDS. Mol. Phys. 5,153 (1962). 156. J. V. HATTON and W. G. SCHNEIDER, Can. J. Chem. 40,1285 (1962).

157. A. G. WHITTAKER and S. SIEGEL, J. Chem. Phys. 42.3320 (i965). 158. A. G. WHITTAKER and S. SIEGEL, J. Chem. Phys. 43,1575 (1965).

159. R. M. MORIARTY, J. Org. Chem. 28, 1296 (1963). 160. L. A. LAPLANCHE and M. T. ROGERS, J. Amer. Chem. Sot. 86,337 (1964).

161. A. A. SANDOVAL and M. W. HANNA, J. Phys. Chem. 70, 1203 (1966). 162. R. M. MORIARTY and J. M. KLIEGMAN, J. Org. Chem. 31,3007 (1966). 163. N. S. BHACCA and D. H. WILLIAMS, Tetruhedron Letters 3 127 (1964). 164. D. H. WILLIAMS and N. S. BHACCA, Tetrahedron21, 1641 (1965). 165. D. H. WILLIAMS and N. S. BHACCA, Terruhedron 21,202 I (1965).

166. N. S. BHACCA and D. H. WILLIAMS, Applications of NMR Spectroscopy in Organic Chemistry, Holden Day, San Francisco, 1964, chap. 7.

167. S. BORY, M. F~TIZON. P. LASZLO and D. H. WILLIAMS, Bull. Sot. Chim. France 2541 (1965).

168. D. H. WILLIAMS and D. A. WILSON, J. Chem. Sot. (B) 144 (1966). 169. J. RONAYNE and D. H. WILLIAMS, Chem. Comm. 7 12 (1966). 170. P. LASZLO and D. H. WILL1AMs.J. Amer. Chem. Sot. 88,2799 (1966). 171. J. H. BOWIE. D. W. CAMERON, P. E. SCHULTZ and D. H. WILLIAMS, Tetrahedron 22,

1771 (1966). 172. J. D. CONNOLLY and R. MCCRINDLE, Chem. and Ind. 379 (1965). 173. J. D. CONNOLLY and R. MCCRINDLE. Chem. andlnd. 2066 (1965). 174. See, for example, K. N AKAMOTO, J. Amer. Chem. Sot. 74, I739 (1952). 175. C. J. TIMMONS, Chem. Comm. 576 (1965). 176. C. R. NARAYANAN and N. K. VENKATASUBRAMANIAN, Terruhedron Letters 3639

(1965). 177. J. SEYDEN-PENNE, T. STRZALKO and M. PLAT. Tetrahedron Letters 4597 (1965). 178. P.S. WHARTON~~~T. I.BAIR,J. Org.Chem.30, 1681 (1965). 179. J. SEYDEN-PENNE, T. STRZALKO and M. PLAT, Tetrahedron Letters 36 11 (1966). 180. G. DI MAIO, P. A. TARDELLA and C. IAVARON, Tetrnhedron Letters 2825 (1966). 181. R. GRIGG, J. A. KNIGHT and P. ROFFEY. Tetruhedron 22,330l (1966). 182. J. E. ANDERSON, Tetrahedron Letters 47 13 (1965). 183. G. F. H. GREEN, J. E. PAGE and S. E. STANIFORTH, J. Chem. Sot. 7328 (1965). . 184. G. SLOMP and F. MACKELLAR, J. Amer. Chem. Sot. 82,999 (1960). 185. F. JOHNSON, N. A. STARKOVSKY and W. D. GUROWITZ, J. Amer. Chem. Sot. 87,

3492 (1965). 186. J. D. CONNOLLY and R. MCCRINDLE, J. Chem. SOL.. (C), 1613 (1966). 187. D. H. WILLIAMS, Tetrahedron Letters 2305 (1965).

188. B. HAMPEL and J. M. KRAEMER. Tetruhedron 22, 1601 (1966). 189. P. L. CORIO and B. P. DAILEY. J. Chem. Phys. 25, 1291 (1956). 190. T. SCHAEFER and W. G. SCHNEIDER, J. Chem. Phvs. 32.1224 (1960). 19 1. T. SCHAEFER and W. G. SCHNEIDER, J. Chem. Ph;s. 32, 12 I8 (1960). 192. W. G. SCHNEIDER, J. Phys. Chem. 66,2653 (1962). 193. A. D. BUCKINGHAM. T. SCHAEFER and W. G. SCHNEIDER. J. Chem. Phys. 32, 1227

( 1960). 194. P. DIEHL. He/u. Chim. Actrc 45.568 (1962).

19.5. P. DIEHL, J. Chim. P/IFS. 61. 199 (1964). 196. B. A. ARBLJZOV. Y. Y. SAMITOV and A. P. KONOVALOV, IZC. Aktrd. Nouk S.S.S.R.

Ser. Fi;. 27, 82 ( 1963).


197. J. H. BOWIE, J. RONAVNE and D. H. WILLIAMS,J. C’hem. Sot. (B) 785 (1966). 198. H. SUHR, Mol. Phys. 6, 153 (1963). 199. R. E. KLINCK and J. B. STOTIIERS. Can. J. Chem. 40.1071 (1962).

200. R. E. KLINCK and J. B. STOTHERS, Can. J. Chem. 40,2329 (1962). 201. M. A. WEINBERGER, R. M. HEGCIE and H. L. HOLMES, Can. J. Chem. 43, 2585

(1965). 202. R. W. TAFT, G. B. KLINGENSMITH and S. EHRENSON, .I. Amer. Chcm. Sot. 87,

3620 (1965). 203. J. N. MURRELL and V. M. S. GIL. Trans. Faraday Sot. 61.402 (1965).

204. M. M. DHIIZGRA, G. GOVIL, C. L. KHETRAPAL and K. VENKATA RAMIAH, Proc. Indian Acad. Sci. 62A, 90 (I 965).

205. S. MATSUOKA, A. MORI and S. HATTORI, SC;. Repts. Kannzawa Univ. 8,45 ( 1962). 206. S. MATSUOKA and S. HATTORI,J. Phys. Sot. Japan 17, 1073 (I 962). 207. A. FRATIELLO, J. Chem. Phys. 41,2204 (1963).

208. R. G. SATTERFIELD, Dissertation Abstr. 25,3293 (1964). 209. J. W. BRAUNER. Dissertation Abstr. 24, 1846 (1963). 210. W. KOCH and HCH. ZOLLINGER, Helv. Chim. Acta 48,554 (1965). 2 I 1. R. E. WYANT. E. J. POZIOMEK and R. H. POIRIER. Ann/. C’him. Acfa 28,496 ( 1963).

2 12. L. A. SINCER’and D. J. CRAM..~. Amer. Chem. Sot. 85. 1080 ( 1963). 2 13. H. H. PERKAMPUS and U. KRUGER. Z. Physik. Chem. (Frankfurt), 48,379 (1966). 2 14. M. W. HANNA and A. L. ASHBAUGH. J. Phys. Cheat. 68,8 11 (1964). 215. P. J. TROTTER and M. W. HANN.4.J. Amer. Chem. Sot. 88. 3724 (I 966). 216. T. C. NEHMAN and A. I. POPOV, J. Phys. Chem. 70,3688 (1966). 2 17. D. W. LARSEN and A. L. ALLRED, ./. Amer. Chem. Sot. 87,12 16 (1965). 218. D.W.L~~~~~andA.L.A~~~~~,J.Amer.Chem.Soc.87,1219(1965). 219. R. FOSTER and C. A. FYFE, Trans. Faraday Sot. 61, 1626 (I 965). 220. R. FosTERand C. A. FYFE, Chem. Comm. 642 (1965). 22 1. R. FOSTER and C. A. FYFE, Biochem. et Biophys. Acfa 112,490 (I 966). 222. R. FOSTER and C. A. FYFE, Trans. Faraday Sot. 62, 1400 (1966). 223. R. FOSTER and C. A. FYFE, Nature 213,59 1 (1967). 224. N. M. D. BROWN, R. FOSTER and C. A. FYFE, J. Chem. Sot. (B), 406 (1967). 225. R. FOSTERING C. A. FYFE, J. Chem. Sot. (B), 926 (1966). 226. J. W. MORRIS. unpublished work. 227. M. I. FOREMAN, unpublished work. 228. H. A. S. PAYNE, unpublished work.

229. C. A. FYFE, unpublished work. 230. P. H. EMSLIE, unpublished work. 23 1. 1. HORMAN, unpublished work.

232. R. FOSTER, C. A. FYFE and J. W. MORRIS, unpublished work.

233. R. FOSTER, C. A. FYFE and M. 1. FOREMAN. Chem. Comm. 913 (1967).

234. R. L. SCOTT. Proc. of the International Conf. on Coordination Compounds. Amster- dam, 1955,p.265; Rec. Trav. Chim. 75,787 (1956).

235. R. FOSTER and D. LL. HAMMICK.J. Chem. Sot. 2685 ( 1954). 236. R. E. MERRIFIELD and W. D. PHILLIPS, J. Amer. Chem. Sot. 80,2778 (1958). 237. M. TAMRES. J. Phys. Chem. 65,654 ( I96 I ). 238. J. E. PRUE. J. Chem. Sot. 7534 (1965). 239. L. J. ANDREWS and R. M. KEEFER, J. Amer. Chem. Sot. 74,450O (1952). 240. R. FOSTER. D. LL. HAMMICK and B. N. PARSONS, J. C/tern. Sot. 555 (1956). 241. F. L. SWINT~N. Chem. Sot. Anniversary Meeting, Exeter (1967). 242. W. A. DUNCAN, J. P. SHERIDAN and F. L. SWINTON, Trans. Faraday Sot. 62, 1090

( 1966). 243. D. S. STEELE, P. N. GATES and W. WHEATLEY, Chem. Sot. Anniversary Meeting,

Exeter (I 967). 244. G. BRIEGLEB and J. CZEKALLA. Z. Electrochem. 59, 184 (1955).

245. N. JONATHAN. S. GORDON and B. P. DAILEY, J. Chem. Phys. 36,2443 (1962). 246. G. G. HALL and A. HARDISSON, Proc. Roy. Sot. A 268,328 (1962).


247. J. R. LETO, F. A. COTTON andJ. S. WAUGH, Nature 180,978 (1957). 248. G. FRAENKEL, R. E. CARTER, A. MCLACHLAN and J. H. RICHARDS, J. Amer. Chem.

Sot. 82,5846 (I 960). 249. C. E. JOHNSON and F. A. BOVEY.J. Chem. Phys. 29. 1012 (1958). 250. A. SAIKA and C. P. SLICHTER,J. Chem. Phys. 22,26 (1954). 25 1. R. W. TAFT and J. W. CARTEN, J. Amer. Chem. Sot. 86,4199 (I 964). 252. R. S. MULLIKEN.J.Amer. Chem. Sot. 74,811 (1952). 253. L. D. WRIGHT and D. B. MCCORMICK, Experientin 20,501 (1964). 254. A. SZENT-GYBRGYI and I. ISENBERG. Proc. Nutl. Acad. Sci. 46, 1334 (1960). 255. A. SZENT-GY~RGYI, 1. ISENBERG and J. MCLAUGHLIN, Proc. Natl. Acad. Sci. 47,

1089 (1961). 256. B. SMALLER, I. ISENBERC~ and S. L. BAIRD, JR., NutUre 191, 168 (1961). 257. G. KARREMAN, I. ISENBERG and A. SZENT-GY~~RGYI,S~~~~~~ 130,1191 (1959). 258. R. BEUKERS and A. SZENT-GY~~RGYI, Rec. Tram. Chim. 81,255 (1962).

259. D. R. H. GOURLEY, Progress in Drug Resenrch, Ed. E. JUCKER, Birkhauser. Basle and Stuttgart, 1964, Vol. 7, pp. 36,4 I.

260. M. E. WOLF, W. Ho and R. KWOK, .I. Med. Chem. 7,577 (1964). 26 1. D. R. KEARNS and M. CALVIN../. Amer. Chem. Sot. 83.2 110 ( 196 1). 262. A. PULLMAN and B. PULLMAN, Quantum Theory of Atoms, Molecules and the

SolidStute, Ed. P. 0. LOWDIN. Academic Press, New York. 1966. pp. 345. 263. A. PULLMAN and B. PULLMAN, Quantum Biochemistry, Wiley, New York, 1963,

p.804: D. R. KEARNS and M. CALVIN, J. Chem. Phys. 34.2026 (1961); L. E. LYONS and J. C. MACKIE, Nature 197,589 (1963).

264. R. FOSTER and P. HANSON, Biochem. et Biophys. Acta 112,482 (1966). 265. A. C. AL.LISON and T. NASH, Biochem. Ph&m&ol. 12,601 (1963). 266. G. CILENTO and P. GIusTI.J.Amer. Chem. Sot. 81.3801 (1959). 267. R. FOSTER and P. HANSON, Trans. Faruduy Sot. 60,2189 <1964). 268. J. P. GREEN, Mechanisms of Release of Biogenic Amines, Ed. B. UVNAS, Pergamon

Press, London, 1965. 269. K. FUKUI, T. YONEZAWA, C. NAGATA and H. SHINGU, J. Chem. Phys. 22, 1433

(1954). 270. K. FUKUI, T. YONEZAWA and C. NAGATA, J. Chem. Phys. 26,83 1 (1957). 27 1. T. S. MOORE, F. SHEPHERD and F. GOODALL, J. Chem. Sot. 1447 (193 I). 272. H. D. ANDERSON and D. LL. HAMMICK, J. Chem. Sot. 1089 (1950). 273. L. J. ANDREWS and R. M. KEEFER, J. Amer. Chem. Sot. 71,3644 (1949). 274. L. J. ANDREWS and R. M. KEEFER, J. Amer. Chem. Sot. 72,3 113 (1950). 275. L. J. ANDREWS and R. M. KEEFER, J. Amer. Chem. Sot. 72,5034 (1950).

276. H. C. BROWN and J. D. BRADY, J. Amer. Chem. Sot. 74,357O (1952). 277. H. C. BROWN and J. J. MELCHIORE, J. Amer. Chem. Sot. 87,5269 (1965). 278. M. A. MUHS and F. T. WEISS. J. Amer. Chem. Sot. 84,4697 (1962). 279. A. GIL-Av and J. HERLINC~, J. Phys. Chem. 66, 1208 (1962). 280. R. J. CVETANOVI~, F. J. DUNCAN and W. E. FALCONER, Can. J. Chem. 42, 2410

(1964). 281. R. J. CVETANOVI&, F. J. DUNCAN, W. E. FALCONER and R. S. IRWIN, J. Amer. Chem.

Sot. 87, 1827 (1965). 282. R. J. CVETANOVI~, F. J. DUNCAN, W. E. FALCONER and W. A. SUNDER, J. Amer.

Chem. Sot. 88, 1602 (I 966). 283. A. R. COOPER. C. W. P. CROWNE and P. G. FARRELL, Trans. Faraday Sot. 62,

2725 (1966). 284. A. R. COOPER, C. W. P. CROWNE and P. G. FARRELL, Trans. Faraday Sot. 63,447 (1967). 285. M. E. PEOVER, Proc. Chem. Sot. 167 (1963). 286. M. E. PEOVER. Trans. Faraday Sot. 60,417 (1964). 287. S. D. Ross and I. KUNTZ, J. Amer. Chem. Sot. 76,300O (1954). 288. F. M. MENGER and M. L. BENDER, J. Amer. Chem. Sot. 88,13 1 (1966). 289. A. K. COL.TER, S. S. WANG. G. H. MERGERLE and P. S. OSSIP, J. Amer. Chem. Sot.

86,3 106 ( 1964).


290. T. F. BOLLES and R. S. DRAGO,J.A~~~. Chem. Sot. 87,5015 (1965).

291. T. F. BOLLES and R. S. DRAGO,J. Amer. Chem. Sot. 88,3921 (1966). 292. L. J. ANDREWS and R. M. KEEFER, J. Amer. Chem. Sot. 73,4169 (195 1).

293. .J. A. A. KETELAAR, C. VAN DE STOLPE, A. GOUDSMIT and W. DZCUBAS, Rec. Trap. Chim. 71,1104 (1952).

294. R. FOSTER, D. LL. HAMMICK~~~ A. A. WARDLEY,J. Chem. Sot. 3817 (1953). 295. R. FOSTER, Nature, 173,222 (1954).

296. P. A. D. DE MAINE,J. Chem. Phys. 26, 1042 (1957). 297. C. P. NASH, J. Phys. Chem. Chem. 64,950 (1960).

298. W. LIPTAY,~. Electrochem. 65,375 (1961).

299. N. J. ROSE and R. S. DRAGO,J. Amer. Chem. Sot. 81,6 138 (I 959). 300. R. S. DRAGO and N. J. RosE,J. Amer. Chem. Sot. 81,614l (1959).

301. H. YAMADA and K. KOZIMA, J. Amer. Chem. Sot. 82, 1543 (1960). 302. J. MORCILLO and E. GALLEGO, Anales de la Real Sot. Espafi. de Fis. y Quim. 56B,

263 (1960). 303. R. E. KAGARISE, Spectrochim. Acta 19,629 (1963). 304. J. LAURANSAN, P. PINEAU and J. LASCOMBE,J. Chim. Phys. 63.635 (1966). 305. A. I. POPOV, R. E. HUMPHREY and W. B. PERSON, .I. Amer. Chem. Sot. 82, 1850

(1960). 306. G. MICHEL and G. DUYCKAERTS, Spectrochim. Acta 21,279 (1965). 307. S. D. Ross and M. M. LABES, J. Amer. Chem. Sot. 79,76 (1957). 308. R. FOSTER and I. B. C. MATHESON, Spectrochim Acfa, 23A. 2037 (1967). 309. M. TAMRES and M. BRANDON,J. Amer. Chem. Sot. 82,2134 (1960). 310. M. BRANDON, M. TAMRES and S. SEARLES, JR., J. Amer. Chem. Sot. 82,2129 (1960). 311. P. H. EMSLIE, R. FOSTER, C. A. FYFE and I. HORMAN, Tetrahedron 21.2843 (1965). 312. R.FOSTERANDI. HoRMAN,J.C~~~.SOC.(B), 171 (1966). 3 13. G. D. JOHNSON and R. E. BOWEN, J. Amer. Chem. Sot. 87,1655 (1965).

3 14. E. A. GUGGENHEIM, Trans. Faraday Sot. 56,1159, (1960). 3 15. S. CARTER, J. N. MURRELL and E. J. ROSCH, J. Chem. Sot. 2048 (1965). 316. J.N.M~~~~~~,J.Amer.Chem.Soc.81,5037(1959). 317. W. G. BARB, Trans. FaradavSoc.49,143 (1953). 318. R. G. SATTERFIELD, Ber. Bunsenges. Physik. Chem. 69, 88 (1965); cf. W. LIPTAY,

Ber. Bunsenges. Physik. Chem. 69,89 (1965). 3 19. W. B. PERSON. J. Amer. Chem. Sot. 87,167 (1965). 320. N. OGIMACHI, L. J. ANDREWS and R. M. KEEFER, J. Amer. Chem. Sot. 77, 4202

(1955).

321. R. FOSTER, D. LL. HAMMICK and B. N. PARSONS, 1. Chem. Sot. 555 (1956). 322. D. F. R. GILSON and C. A. MCDOWELL, J. Chem. Phys. 39,1825 (1963). 323. H. G. SMITI~ and R. E. RuNDLE,J.Amer. Chem. Sot. SO,5075 (1958). 324. H. G. SMITR, J. Chem. Phys. 40,2412 (1964). 325. D. F. R. GILSON andC. A. MCDOWELL, J. Chem. Phys. 40,2413 (1964).

326. L. W. D.%ASCH, Spectrochim. Acta 9,726 (1959). 327. D. F. R. GILSON and C. A. MCDOWELL, Can. J. Chem. 44,945 (1966). 328. J. E. ANDERSON. J. Phys. Chem. 70,927 (1966). 329. D. C. DOUGLASS, J. Chem. Phys. 32,1882 (1960). 330. M. READ, R. CAHAY, P. CORNIL andJ. DUCHESNE, Compt. Rend. 257d 1778 (1963). 331. P. CORNIL, J. DUCHESNE, M. READ and R. CAHAY, Bull. Classe Sci. Acad, Roy.

Belg. 50,235 (1964).

332. D. BIEDENKAPP and A. WEISS, Z. Naturforsch. I9a, 15 18 (1964). 333. V. S. GRECAISHKIN and I. A. KYUNTSEL, Soviet Phys. Solid State 5,694 (1963). 334. V. S. GRECHISHKIN and I. A. KYUNTSEL, Opitika i Spectroskopiya 15, 832 (1963);

Optics and Spectroscopy 15,453 (1963). 335. V. S. GRECHISHKIN and I. A. KYUNTSEL, Zhur. Struct. Khim. 7, 119, (1966). 336. V. S. GRECHISHKIN and I. A. KYUNTSEL, Opitika i Spectroskopiya 16, 161 (1964);

Optics and Spectroscopy 16,87 (1964).

Documents

Chapter 1 Nuclear magnetic resonance of organic charge-transfer complexes