Advances in Physical Organic Chemistry, Volume 11

Advances in Physical Organic Chemistry

This Page Intentionally Left Blank

Advances in Physical Organic Chemistry Volume 11 Edited by

V. Gold Department of Chemistry King’s College University of London Associate Editor

D. Bethel1 The Robert Robinson Laboratories University of Liverpool England

1975

Academic Press London New York A Subsidiary o f Harcourt Bruce Jouunouich, Publishcrs

San Francisco

ACADEMIC PRESS INC. (LONDON) LTD 24-28 Oval Road, London NWI

U S . Edition published by ACADEMIC PRESS INC. 111 Fifth Avenue, New York, New York 10003, San Francisco

Copyright

0 1975 By Academic Press Inc. (London) Ltd

All Rights Reserved

No part of this Book may be reproduced in any form by photostat, microfilm, or any other means, without written permission from the publishers

Library of Congress Catalog Card Number: 62-22 125 ISBN 0-12-033511-5

PRINTED IN GREAT BRITAIN BY WILLIAM CLOWES & SONS LIMITED LONDON, COLCHESTER AND BECCLES

Contributors to Volume 11

J. Cornelisse, Department of Organic Chemistry, Gorlaeus Laboratories, Leiden University, The Netherlands

G. P. de Gunst, Department of Organic Chemistry, Gorlaeus Laboratories, Leiden University, The Netherlands

D. G . Farnum, Department of Chemistry, Michigan State University, East Lansing, Michigan, U.S.A. T. H. Fife, Department of Biochemistry, University of Southern California, Los Angeles, California 90033, U.S.A. E. Havinga, Department of Organic Chemistry, Gorlaeus Laboratories, Leiden University, The Netherlands G. M. Kramer, Corporate Research Laboratories, Exxon Research and Engineering Co., Linden, New Jersey 07036, U.S.A.

M. Liler, School 0.f Chemtitry, The University, Newcastle upon Tyne, England

This Page Intentionally Left Blank

Preface

With the appearance of Volume 11, Advances in Physical Organic Chemistry enters the second decade of its life. Many exciting developments of the subject have taken place since the publication of the first volume. It is a matter of some pride to be able to state that the contents of the first ten volumes have not only reflected these advances but, in many cases, have pointed the way and had a seminal influence on later work. Over the same period, Physical Organic Chemistry has become “respectable”, as is shown by the titles of several books and by the institution of professorships and other academic appointments in the sub.ject. The value of physical organic methods has become widely appreciated in industry. Biennial International Conferences on Physical Organic Chemistry have been started under the auspices of the International Union of Pure and Applied Chemistry, who have also appointed a Commission on Physical Organic Chemistry. No doubt, these tokens of recognition give some satisfaction to many scientists who have found the search for glimpses of a quantitative understanding of organic chemistry the most fascinating field of study. Yet this very respectability has its dangers. What made Physical Organic Chemistry a dynamic movement in science was that some outstanding men refused t o be type-cast as physical or organic chemists and instead pursued the investigation of absorbing problems irrespective of the nature of the techniques required. Physical Organic Chemistry as an institutionalized “discipline” runs the risk of losing this spirit. It will be a continuing objective of our series not t o allow this to happen, by encouraging the publication of contributions which do not as yet conform to established notions of the scope of the field. It is natural that a living subject should at times produce quite extreme divergences of opinion. This series will continue t o include

... Vlll

PREFACE

contributions which may present a topic from an unfashionable point of view. At other times, complementary accounts of controversial issues may be published, with the dual aim of pinpointing problems and resolving confusion. Thus, the present volume contains a reappraisal of the structure of the norbornyl cation by George Kramer, and in a subsequent one the problem of the norbornyl cation and of its significance in organic chemistry will be discussed in an article by George Olah, Eric Nordlander and Paul Schleyer. Beginning with the present volume, the series will have the benefit of the specialist knowledge of Dr. Donald Bethell who is joining me as Associate Editor. At the same time, some changes have been made in the production of volumes, so as to improve lay-out and speed up publication. It is a pleasure to acknowledge the service to the subject provided by the publishers, Academic Press, and especially the indefatigable assistance given to the Editor over the years by Messrs R. S. Lawrence and R. Adams. My main thanks go to the contributors who have so readily come forward with ideas or responded to suggestions for timely reviews. Dr. Bethel1 and I hope that we shall continue to receive such cooperation, and will always welcome expressions of views on the series and its contents. V. Gold London, November 1974.

Con tents

Contributors to Volume 11 .

.

v

Physical Organic Model Systems and the Problem of Enzymatic Catalysis - Thomas H. Fife

1. Introduction . 2. Enzymatic Catalysis: General Principles . 3. General Base, Nucleophilic Catalysis: a-Chymotrypsin 4. Metal-Ion Catalysis: Carboxypeptidase . . 5. General Acid Catalysis: Lysozyme 6. Conclusion .

. .

1 5 . 29 . 64 . 81 .115

Charge Density-NMR Chemical Shift Correlations in Organic Ions - D. G. Farnum

1. Introduction

.

2. Theory. 3. Empirical Observations. 4. Conclusions .

.

. 123 . 126 .135

. 172

The Norbornyl Cation: A Reappraisal o f its Structure Under Stable Ion Conditions - G. M . Kramer

1. Introduction

.

.177

CONTENTS

X

2. 3. 4. 5. 6.

7. 8. 9. 10. 11.

Possible Structures of the Norbornyl Cation Solvolytic Background. . Theoretical Status . The Search for a Protonated Cyclopropyl Ring ESCA . 13C-nmr . 'H-nmr. . Raman Spectra . Related Ions. Summary .

. 179

. 180

. 192 .194 .199 .202 f211 .215 . 218 .221

Nucleophilic Aromatic Photosubstitution - J. Cornelisse, G . P. de Gunst, E. Havinga

1. Introduction . .225 2. The Excited State from which the Nucleophilic Aromatic Photosubstitution Starts. Kinetics . .236 3. Orientation Rules in Nucleophilic Aromatic Photosubsti.245 tution . .253 4. Investigations on Intermediates . 5. Epilogue . .261

Alternative Protonation Sites in Ambident Conjugated Systems - M. Liler

1. 2. 3. 4.

Introduction . . Methods of Investigation Cation Stability and Solvation . Protonation Sites in Conjugated Molecules

.

.267 .270 .287 .300

Physical Organic Model Systems and the Problem of Enzymatic Catalysis

THOMAS H. FIFE

Department of Biochemistry, University of Southern California, Los Angeles, California 90033, U.S.A. 1. 2.

3.

.

Introduction Enzymatic Catalysis: General Principles Intramolecular Catalysis Multifunctional Catalysis . Conformation and Strain Effects . General Base, Nucleophilic Catalysis: a-Chymotrypsin Imidazole Catalysis Hydroxyl Group Catalysis . Mechanistic Conclusions Metal-Ion Catalysis: Carboxypeptidase. . Metal-ion Catalysis of Ester and Amide Hydrolysis Carboxyl Group Catalysis . Mechanistic Conclusions General Acid Catalysis: Lysozyme General Acid Catalysis . Intramolecular Catalysis Nucleophilic Catalysis Acylal Hydrolysis. Mechanistic Conclusions Conclusion References

.

.

.

.

4.

5.

.

. .

6.

. .

.

.

.

.

1

.

5

.

.

. . . . . . . . . . . . .

. .

. .

5 19 28 29 30 39 60 64 66 73 80 81 81 90 105 108 112 115 117

1. INTRODUCTION Enzymes are the best catalysts known. They catalyse chemical reactions in aqueous solution near neutral pH and at mild temperatures with attainment of great velocities. Although suitable standards 1

2

T. H. FIFE

of comparison are often lacking, still it can be calculated (Bruice and Benkovic, 1966; Jencks, 1969; Bender, 1971) that certain enzymes will give rate enhancements by factors of the order of 1010-1014 compared with similar non-enzymatic reactions. Therefore, knowledge of the mechanism of action of enzymes is of great importance, not only because enzymes are of biological interest, but also for the greater chemical understanding of catalysis in aqueous solution that would result. A complete description of the mechanism of action of an enzyme requires knowledge of a number of factors, among which are: (a) The structure of the active site and of enzyme-substrate complexes. (b) The specificity of substrates and their ability to bind to the enzyme. (c) The rate constants for the various steps in the reaction and the chemical mechanism of each step, ie., a description of all of the intermediates and transition states along the reaction coordinate. (d) An explanation for the magnitude of the rate constants. Hence, in approaching the problem of enzymatic catalysis, infonnation must be obtained from a number of experimental techniques. Being proteins, enzymes are extremely complicated molecules, and the mechanistic interpretation of kinetic data is generally difficult. It is this complexity that in large part explains why, although kinetic studies have been conducted with many enzymes since the end of the nineteenth century, there is still n o enzyme for which a mechanism has been definitely established. Extensive effort has been expended in recent years on mechanistic investigations of model reactions, chemical reactions similar t o a post-binding enzymatic reaction, but where the number of variables is limited and strong evidence can be obtained pertaining to mechanism. Model work can begin only after enzymologists have provided information about the structure of the enzyme and the identity of groups in the active site. Once this basic information is available, the next step is to attempt t o gain understanding of the chemistry of the process. In some cases it has not been possible to begin with sophisticated models, but rather, chemical work has been necessary to determine whether mechanisms suggested for an enzyme were even chemically feasible. Work on acetal hydrolysis (Fife, 1972) related t o the mechanism of action of lysozyme is an illustration of this point.

PHYSICAL ORGANIC MODEL SYSTEMS: ENZYMATIC CATALYSIS

3

In addition to determining the mechanism of action of an enzyme, it is important t o be able t o explain, in quantitative terms, the rapid rate of the enzyme reaction. That is, if factors A, B, and C are important, we would like to analyse the observed rate of the reaction in terms of these factors according t o equation 1 and give numerical values to the coefficients. To date this has not been possible for most log kobsd = aA + bB

+ cC + . . . nN

enzymes and possible only to an approximation with a-chymotrypsin (Bender et al., 1964). It is apparent that this type of information can best come from detailed chemical studies where the various factors are examined in isolation. Chemical models then fulfil a twofold purpose: first, in providing reasonable mechanisms for the enzyme and casting doubt on others, thereby suggesting experimental approaches having a heightened chance of being fruitful; and second, in leading t o explanations of the observed rates in terms of structure and mechanism [items (c) and (d)]. The number of different functional groups that an enzyme can use in the catalytic process is quite limited. Among them are the imidazole ring, aliphatic and phenolic hydroxyl, carboxyl, sulphhydryl, and amino groups. A reasonable speculation is that enzymes which catalyse similar reactions will have mechanistic features in common. This supposition is supported by the fact that several different esteratic enzymes having serine at the active site have an identical sequence of amino-acids around the active serine (Bruice and Benkovic, 1966; Bender, 1971). The amino-acid sequences for chymotrypsin and subtilisin (Wright et ~ l . ,1969) are dissimilar, but the catalytically important groups are the same and are arranged in nearly the same geometrical relationship, implying identical mechanisms of action. Also, it seems reasonable that common mechanistic features may extend to enzymes which catalyse different reactions if the same catalytic group can function. For example, if an aliphatic hydroxyl group acts as a nucleophile in two different reactions, at least some common factors will probably be important in promoting catalysis. Consequently, from studies of catalysis by relatively few functional groups in carefully selected reactions it may be possible t o formulate general principles for catalysis in aqueous solution including enzyme catalysis. The purpose of the present review is to summarize the pertinent data concerning several well-studied reactions, to draw reasonable conclusions, and to point out some of the questions remaining to be

4

T. H. FIFE

answered. T o restrict the discussion to a manageable level, concern will be mainly with hydrolysis reactions relevant to the research interests of the author. Many important mechanism studies must therefore be omitted. For the most part only those reactions will be considered in detail which bear direct analogy to corresponding enzyme reactions and can therefore be considered true models. The enzymes themselves will not be discussed except to give necessary background information pertaining t o postulated mechanisms. Chemical model investigations will be described which relate to three hydrolytic enzymes, a-chymotrypsin, carboxypeptidase, and lysozyme. Numerous excellent review articles have appeared concerning these enzymes, and their reactions (see for example, Boyer, 1970) and the mechanisms of the enzyme-catalysed reactions (Westheimer, 1962; Bruice and Dunn, 1973; Kaiser and Kaiser, 1972). These enzymes serve as perhaps the best examples of the model approach because detailed structural information is available. In each case the complete amino-acid sequence of the enzyme is known, and the three-dimensional structure has been determined by x-ray crystalloresolution (for three dimensional structural graphic analysis at 2 drawings, see Dickerson and Geis, 1969). Also, the basic chemistry of the types of reactions catalysed has been extensively studied over a long period of time, and there is a wealth of background chemical information (Bender, 1960; Bruice and Benkovic, 1966; Johnson, 1967; Jencks, 1969; Cordes, 1967; Fife, 1972). The enzymes have the aliphatic hydroxyl group of serine, the imidazole ring of histidine, and carboxyl groups as catalytically important functional groups at the active sites. The problem resolves itself into determining how such functional groups can participate in the hydrolytic reactions and how the rates of the enzymatic reactions can be accounted for in mechanistic terms. Model studies have made a major contribution towards an understanding of the mechanistic possibilities which lysozyme might employ. Indeed, the model approach is perhaps the method with the greatest chance of success in attempts t o comprehend the complex reactions of enzymes such as lysozyme where the natural substrates are not suitable for detailed kinetic work. a-Chymotrypsin can be studied with a variety of ester and amide substrates of low molecular weight, and much mechanistic information is available from direct work on the enzyme. Chemical models have been chiefly useful in promoting understanding of the chemistry of the reactions and in providing a basis for comprehension of the magnitude of the rate

a

PHYSICAL ORGANIC MODEL SYSTEMS: ENZYMATIC CATALYSIS

5

constants in terms of individual mechanistic factors. At the other extreme, crucial model systems have not yet been investigated in relation to carboxypeptidase, but that enzyme illustrates types of questions that well-designed chemical models might answer. Information is currently available which allows more definite conclusions than have previously been possible. Rate enhancements have been obtained in several simple chemical reactions that are of similar magnitude to those observed in analogous enzyme-catalysed reactions, and the goal of analysing the individual factors that can give such large rate accelerations is now perhaps within reach. The recent work will be stressed in this review. Deeper insight into complex enzyme reactions has resulted from the study of simpler chemical models; but also, fundamental and novel observations concerning catalysis in aqueous solution have been made. The chemical work is therefore o f interest in its own right, apart from any applicability t o enzymatic catalysis. It.is the author’s. opinion that this has not been stressed t o a sufficient extent in previous discussions of enzymatic catalysis. Apart from the practical advantages that might result from understanding how enzymes function, the primary scientific importance would appear to be greatly increased knowledge of chemical catalysis. From this inverted viewpoint it is the chemistry that is elucidated that will prove to have ultimate scientific value, whether or not it directly applies to enzymes.

2. ENZYMATIC CATALYSIS: GENERAL PRINCIPLES Intramolecular Catalysis Several fundamental aspects of enzymatic catalysis must be considered in any discussion of the chemistry of enzymatic reactions. First, an enzyme-catalysed reaction proceeds with formation of an E+S

kl

k- 1

ES

k cat

E+P

(2)

enzyme-substrate complex which then breaks down t o product and free enzyme [equation (2)]. The evidence for this is in most cases only suggestive (saturation kinetics), but an ES complex has been observed in some reactions (Chance, 1943, 1947), and it is generally assumed that such a complex is formed in most, if not all,

6

T. H. FIFE

enzyme-catalysed reactions. Second, it is held that reaction takes place in a limited, specific area of the protein surface (the active site) with functional groups of amino-acid residues participating in the reaction. These ideas have been highly advantageous in regard to the development of chemical catalysis in aqueous solution. If the above concepts are correct, then an enzymatic reaction proceeding through an enzyme-substrate complex with the substrate bound close to appropriate functional groups is quite analogous to a chemical intramolecular reaction. Substantial effort has therefore been expended on the study of such reactions in attempts better to comprehend enzyme catalysis (Bruice, 1970; Kirby and Fersht, 1971). Intramolecular reactions usually occur rapidly in comparison with similar bimolecular reactions, and a number of explanations have been put forth concerning the efficiency of intramolecular catalysis. Westheimer (1962) pointed out that intramolecular reactions should have a more favourable entropy of activation than their bimolecular counterparts because translational entropy of the catalyst is not lost as it is in the bimolecular reaction. To determine the efficiency of an intramolecular reaction, the first-order rate constant is often compared with the second-order rate constant of the corresponding bimolecular reaction proceeding by the same mechanism. This ratio has units of molarity ( s - ~ /-M‘ s - ’ ) and is taken to be the “effective molarity” of the neighbouring group, i.e. the concentration of bimolecular catalyst required to give a pseudo-first order rate constant of the magnitude observed in the intramolecular reaction. An intramolecular dimethylamino-group has an effective molarity of 1000-5000 M for attack at the ester carbonyl (Bruice and Benkovk, 1963) and there is a more favourable AS* value than in the intermolecular reaction between trimethylamine and aryl acetates (see the data in Table 1). Page and Jencks (1971) have calculated that favourable changes in AS* in intramolecular reactions could be large enough to explain the rapid rates of some enzymatic reactions, and effective concentrations of the order of l o 8 M may be accounted for without introducing new chemical concepts. Intramolecular nucleophilic reactions could also be facilitated over their intermolecular counterparts if the reaction centre and the nucleophile are compressed in the ground state. Part of the van der Waals repulsion energy could thereby be overcome in the ground state, resulting in a more favourable AH* value. Solvation

TABLE 1 Activation Parameters (kcal mole-')ofor Nucleophilic Displacement by the Dimethylamino-Group at 25 (Bruice and Benkovic, 1963)

m 0 2

m-NOz p-CI H P-CH3

12.3 12.1 12.5 12.9

-6.3

-8-0 -9-1 -9.4

11.9 11.5 15.9 12.5 13.7

-1.9 -4.3 -2-2 -5-7 -5.1

11.5 11-8 13.8 12.3 14.4

-2.6 -4.4 -4.1 -6.4 -5-5

m z N

4

5

=!

n

T. H. FIFE

8

factors might also be of great importance. Before a nucleophile can attack it must be desolvated, requiring expenditure of substantial energy. However, an intramolecular nucleophile would be less heavily solvated in the ground state than an intermolecular nucleophile in dilute solution if the neighbouring group and the reaction centre are immediately adjacent in the same molecule so that water molecules cannot fit between. As an intramolecular nucleophile is more rigidly held with respect to the reaction centre, the rate of the reaction increases as illustrated in Table 2. Bruice and Pandit (1960b) concluded that the rate increases were due t o restriction of unfavourable rotamer distribution. The most energetically favourable ground state conformation would have the carboxyl group extended into the solvent, viz. 0

II -0-c-c-o

0

"

-0

where it could not attack the carbonyl. Removal of rotational degrees of freedom would therefore greatly enhance the rate of reaction. As seen in Table 3, similar rate enhancements have been observed in lactonization reactions as the reacting groups are held in proximity t o each other (Storm and Koshland, 1970, 1972a). A rate enhancement of 10'' M is observed in lactonization of [ l ] with respect to the bimolecular esterification in equation (4) (Milstien and Cohen, 1972). The methyl group in the 3-position of [ l ] fits between the geminal methyl groups in the side chain and locks the system. One function of an enzyme is t o bring the substrate by binding at the active site into proximity with functional groups of the enzyme. 0

5--. \

+ HzO

CH3

(3)

9

PHYSICAL ORGANIC MODEL SYSTEMS: ENZYMATIC CATALYSIS

TABLE 2 Effect of Structure on Carboxyl Group Participation in Hydrolysis of Phenyl Esters (Bruice and Pandit, 1960b) Compound

Relative rate

1

19.3

230

10,300

0

53,000

Koshland (1962) has calculated, however, that such a “propinquity” effect will not explain the large rate enhancements observed with enzymes unless there are more than two functional groups involved; with utilization of five functional groups ( 2 substrates and 3 catalytic groups) a rate increase of 1 0 l 8 would be possible. Such multifunctional catalysis would, of course, be impossible to demonstrate

T. H. FIFE

10

TABLE 3 Acid Catalysed Lactonization at 25' in 20% EtOH-H20, pa = 0.4 (Storm and Koshland, 1972a) Correctedb rel. rate

klac

'

(M- min- ')

Substrate

c::

EtOH + CH3C02H

(Co2H

VH

0.00 109

krel

1.0

1

0.086

80

413

0.127

117

200

0.344

316

17

6620

1660

1,030,000

18,700

I

CI13

&COzH CH2 OH

7.23

1120

'Ionic strength. Corrected for torsional strain and proximity effects and normalized t o a value for the bimolecular reaction of EtOH and HOAc.

in chemical tions. Koshland reason why an increase

models and appears very improbable in enzymatic reac-

(1962) argued that if proximity effects were the only intramolecular reactions are favoured, that is, because of in local concentration of the catalytic function, then

PHYSICAL ORGANIC MODEL SYSTEMS: ENZYMATIC CATALYSIS

11

effective molarities of 5-50 M would be observed. In an intramolecular reaction [equation (5)] each A must have B as a nearest neighbour. In a bimolecular reaction [equation ( 6 ) ] A and B will A + B

kl

AB

(6)

encounter each other randomly. Only a fraction of A molecules will have a B nearest neighbour depending upon the concentration of B. The ratio of rate constants (k /k, ) should be the concentration of B necessary to give A one nearest neighbour B. The upper limit in concentration would be liquid B. For a reaction in water (where B is H,O) this ratio should be no greater than 55.5 M . Assuming that reactant species are about the size of a water molecule and that in a dilute solution of A there would be about 12 nearest neighbour B molecules, a theoretical ratio of rate constants of 55*5/12= 4.6 M is obtained. Effective molarities of 5-50 M have been observed but much larger ratios have also been measured ( l o 8 M ). For a detailed discussion see: Bruice, 1970. From studies of intramolecular lactonization reactions, where rate enhancements of l o 3 t o lo6 over the bimolecular counterpart were found (Table 3 ) , Koshland proposed the concept of “orbital steering” (Storm and Koshland, 1970). Since collision of a nucleophile and receptor atom will be effective on only a small portion of the surface of each, the reaction velocity should be sensitive to proper orientation; steric factors which caused favourable alignment of orbitals might give rise to large rate increases. Bruice et al. (1971) have argued forcibly that some of the rate enhancements seen in intramolecular reactions are much larger than expected from orientation effects. After correcting for proximity effects, orientation factors of 106-107 would be necessary. Bruice argues that this would require covalent bonds to be approximately 100 times more resistant toward bending deformations than spectroscopic data and molecular orbital theory indicate. Consequently, “orbital steering” will not account for rate enhancements of up t o 10’. Koshland has recently replied to the criticisms that have been levelled at the concept of orbital orientation as an important factor in intramolecular reactions (Storm and Koshland, 1972a, b). The data in Table 4 were presented showing effects on lactonization of minor changes in structure, such as changing the carbon superstructure, changing orientation by the presence of methyl groups, and changing the nucleophile from oxygen to sulfur. The lactonization rate is sometimes promoted while in other cases rate decreases are brought

T. H. FIFE

12

TABLE 4 Effect of Ring Structure, Methyl Groups, and Nucleophile on Rates of Lactonization (See Table 3 for Reference Compounds; Storm and Koshland, 1972b) at 25O in 20% EtOH-H20 (p = 0.4)

Compd.

4 f4

&lac (M-'min-')

Rate relative to unsubstituted compound

10.0

C02 H

OH

0.950

CO2H

152

&c02H CH2OH

2.88 x 10-4

4

0

1.50 x

4700'

3.0 x 10-3

936'

2 H CH2 SH

71.0 OH

90'

CO2H

0.63 x lo-'

PHYSICAL ORGANIC MODEL SYSTEMS: ENZYMATIC CATALYSIS

13

TABLE 4-continued Compd.

&CH3 OH

OH

k h c (M-' min-')

0.30

Rate relative to unsubstituted compound

2.7 x 10-4

CO2H

107.0

113

100

1

CO2H

d? COzH

522

5 2.2

a CH3. CH2. SH + CH3C02H: k~ = 3.20 x 10-6M-' min-': re1 rate = 1.0

about by structural modifications. Koshland argues that such changes would be expected if orbital orientation was of great importance. This controversy forms part of the large general question of why intramolecular reactions are so favourable. Increases in the rate of anhydride formation due t o bulky substituents in the 3-position of mono-/3-bromophenyl glutarates (the 3,3-diisopropyl derivative cyclises 1030 times faster than the unsubstituted compound) are accompanied by more positive AS*-values (Bruice and Bradbury, 1968). Tetramethylsuccinanilic acid [ 21 cyclises rapidly at 25" t o the anhydride with carboxyl group participation which is 1200 times more favourable than with succinanilic acid (Higuchi et al., 1966). Furthermore, rate enhancement is associated with a more favourable AH*,the value of AS* actually being more negative. The activation parameters were rationalized by Kirby and Fersht (1971) by assuming that the geminal methyl groups force

14

T. H. FIFE

the carboxyl and carbonyl groups into such close alignment in the ground state that part of the van der Waals repulsion energy is 0

17)

CH3 0

overcome. The unfavourable effect on AS* was ascribed to greater solvation of the transition state than of the ground state. Thus, structural effects on activation parameters in intramolecular carboxyl group-catalysed ester and amide hydrolysis can be manifested in either AS* or AH*, depending upon the compound or type of reaction. The ester hydrolysis reactions studied by Bruice and Bradbury (1968) are more easily interpretable examples since the amide cyclization probably involves a pre-equilibrium protonation step. Effects of structure on AH* and AS* could arise from effects on either the pre-equilibrium step or the rate-determining nucleophilic attack. Bunnett and Hauser (1965) found that lactonization of [3] proceeded 300 times faster than lactonization of 2-hydroxy-. methylbenzoic acid. Steric compressional effects are probably responsible, and again the rate increase is characterized by a more favourable AH*.

Bruice and Benkovic (1964) and Bruice (1970) have noted that conversion of a bimolecular reaction into an intramolecular reaction corresponds t o a reduction in kinetic order. Comparison of reactions of varying kinetic order (Table 5 ) reveals that a change of 45 kcal mole-' in TAS* accompanies reduction in order by one. These reactions include 17 displacement reactions on phenyl esters and 4 on thiol esters with an average value of -TAS*/(kinetic order)

PHYSICAL ORGANIC MODEL SYSTEMS: ENZYMATIC CATALYSIS

15

TABLE 5 Comparison of the Value of TAS*/Kinetic Order in Displacement Reactions of Phenyl and Thioiesters (Bruice, 1970) Reactants

-TAs* /kinetic order (kcal mole-')

Ref.

4-0-6

U

1

4

b

1

3

b

1

5

C

5

d

4

d

Kinetic Order

II

0

0

1

T. H. FIFE

16

TABLE 5-continued ~~

Reactants

+ HzO

X

X

G

O

A

c + (CH3)3N

e OAc

+ -OH

~

~~~~~~

Kinetic Order

-TAS*/kinetic order (kcal mole-')

Ref.

2

4

e

4

e

2

3

f

2

5

f

4

d

4

g

3

d

PHYSICAL ORGANIC MODEL SYSTEMS: ENZYMATIC CATALYSIS

17

TABLE 5-continued Kinetic Order

Reactants

o OAc + 2(NH2)2

X

3

N ~ N H X

o

O

A

c

+ 2

L J

3

;;1...(1

3

-TA.S*/kinetic order (kcal mole-')

Ref.

4-5

g

5

g

5

h

5

h

5

h

5

h

a Bruice and Bradbury, 1968.

Gaetjens and Morawetz, 1960. Thanassi and Bruice, 1966. Bruice and Benkovic, 1963. Felton and Bruice, 1969. f Fersht and Kirby, 1967. g Bruice and Benkovic, 1964. Fedor and Bruice, 1964.

being 4-4 f 0-8 kcal mole-'. At 25" a rate decrease of about l o 5 will result for each additional species incorporated in the transition state. Bruice (1970) points out that only to an approximation can kinetic and potential energies of activation be associated with observed

18

T. H. FIFE

values of AS* and AH*. Also, since AH* and AS* are not completely independent, compensation may occur, favourable changes in one being accompanied by unfavourable changes in the other. Therefore, the treatment’s success with a fairly large body of data is striking. Undoubtedly the effects observed by Koshland and earlier by Bruice are produced by the same cause. Whether orientation effects are due to “orbital steering” or elimination of conformations unfavourable for the reaction makes a conceptual difference only in that “orbital steering” would include a transition-state effect whereas the latter would arise from restriction of possible ground states. What is unmistakably clear from available data is that constraining an intramolecular nucleophile close to the reaction centre can produce large rate increases. Thus, from the data of Bruice and Pandit (1960) in Table 2, p-bromophenyl exo-3,6-endoxo-A-4-tetrahydrophthalate cyclizes 53,000 times faster than p-bromophenyl glutarate. In the bicyclic compound the carboxyl group is constrained adjacent to the carbonyl and degrees of freedom have been removed, whereas in the glutarate ester free rotation is possible about two single bonds. Bruice and Pandit note that removal of each degree of freedom will increase the rate about 230 times. It is doubtful whether unfavourable conformations could be eliminated more effectively than in the bicyclic ester without steric compression effects becoming important. Bruice’s viewpoint is that the advantage of intramolecular reactions is entropic because of the freezing out of degrees of freedom in the ground state. This results in favourable changes in translational and rotational entropy of activation compared with corresponding bimolecular reactions. This is considered to be entirely a proximity effect. In the intramolecular reactions studied by Bruice and Koshland and their co-workers, proximity effects (reduction in kinetic order and elimination of unfavourable ground state conformations) and orientation effects might give rate accelerations of l o 7- l o 8 . Hence, these effects can by themselves account for the enhancements seen in most intramolecular reactions. However, a factor of 107-108 is less than the rate acceleration calculated for many enzyme reactions and certain intramolecular reactions, for example, hydrolysis of benzaldehyde disalicyl acetal (3 x 10’) (Anderson and Fife, 1973) and the lactonization reaction of [ 11 (10’ ) where a trimethyl lock has been built into the system. If hydrolysis of tetramethylsuccinanilic acid (Higuchi et al., 1966)represents a steric compression effect ( l o 3 rate acceleration), then proximity, orientation, and steric compression

PHYSICAL ORGANIC MODEL SYSTEMS: ENZYMATIC CATALYSIS

19

might give rate increases of 10' '-10' '. Thus there can be a factor of 102-104 in the rate constants n o t accounted for by th e three effects above; transition state stabilization and release of ground state strain are likely candidates. Diamines of varying structure show rate enhancements of 20- 200 fold, compared to monofunctional aliphatic amines, in nucleophilic reactions with N-acetylimidazole (Page and Jencks, 1972). These were attributed to intramolecular general base catalysis of proton removal from the attacking nitrogen, viz., -0

The low effective molarity of -1.0 M for the catalysing base and lack of sensitivity t o diamine structure suggested a loose transition state with minimal geometric requirements. General base catalysis of aminolysis of N-acetylimidazole approaches the lower limit 01 the range of rate accelerations in intramolecular reactions. If a bimolecular reaction takes place with a tight transition state, loss o f substantial translational an d rotational cntropy of the reactants (40-50 C.U. for typical molecules in solution) will be required. In a n intramolecular reaction o r intracomplex enzymatic reaction, a significant portion of this entropy has already been removed so that little further entropy loss nccd occur, resulting in a large rate enhancement (1O8 M). If, however, the transition state has only limited geometrical requirements, little cntropy nccd be lost in the bimolecular case and the rate acceleration in the intramolecular reaction will be small.

hilultifunctionul Catalysis Another fundamental idea that has been invoked to explain enzymatic catalysis is that such reactions utilize bifunctional or multifunctional catalysis; that is, several functional groups in the active site are properly aligned with the substrate so that concerted catalysis may occur. Mutarotation of tetramethyl glucose is frequently cited as an example o f bifunctional catalysis. Lowry and

20

T. H. FIFE

Faulkner (1923; 1927) reported that mutarotation occurred much more readily in a mixture of pyridine and cresol than in either by itself. Swain and Brown (1952) extended this work and found that a-pyridone, where an acidic and a basic group are incorporated in the same molecule, is an excellent catalyst, 0.05 M concentration giving rise to a reaction 5 0 times faster in benzene solution than a solution of 0.05 M phenol and 0-05 M pyridinc. A 0.001 M solution of a-pyridone was 7000 times more effective than equivalent concentrations of phenol and pyridine. A concerted mechanism was proposed [4].

The advantage of bifunctional over monofunctional catalysis could arise from reduction in the number of species in the transition state, which would presumably affect AS*, and elimination of high-energy intermediates which might affect AH*. Rony (1968) measured activation parameters for the mutarotation of tetramethylglucose in benzene solution and found that 2-pyridone has a low activation enthalpy relative to pyridinium-phenoxide. The enhanced catalytic ability of 2-pyridone was attributed to the fact that it can exchange two protons without forming a high-energy dipolar ion. Zeffren and Hall (1973) have commented that, since reactions with polar transition states in nonpolar solvents can be accelerated by several orders of magnitude by the presence of low concentrations of salts (Winstein e t al., 1959), the rate enhancement of tetramethylglucose mu taro tation provided by the presence of acid-base pairs such as phenol and pyridine may be due t o formation of ion pairs in benzene solution. Salts which do not act as acids and bases catalyse mutarotation of tetramethylglucose in aprotic solvents (Eastham et a/., 1955; Blackall and Eastham, 1955; Pocker, 1960). The efficiency of enzymatic catalysis could arise largely from electrostatic catalysis

PHYSICAL ORGANIC MODEL SYSTEMS: ENZYMATIC CATALYSIS

21

in a non-polar region of the enzyme. The presence of ion pairs could be extremely important in such a region. Examples of catalysis involving concerted cyclic proton transfer in water have recently been reported in the hydrolysis of N-phenyliminotetrahydrofuran (Cunningham and Schmir, 1966, 196 7). There is no effect of H,PCT4, HCO;, and CH,COOH on the rate of hydrolysis, but the type of product obtained, butyrolactone or y-hydroxybutyranilide is dependent on their concentration. The following scheme was proposed [equation (8)].

?

+

?

I

I

The buffer species were suggested to increase the rate of formation of butyrolactone by the cyclic concerted mechanism [ 5 ] . Hydrolysis of yhydroxybutyranilide is also catalysed by H2 PO, and HCO; but not by imidazole. A similar mechanism has been postulated in the hydrolysis of trifluoroacetanilide [ 61 (Eriksson and Holst, 1966).

&OJ F3

There are few, if any, examples of intramolecular bifunctional catalysis in aqueous solution. Maugh and Bruice (1971) examined ester hydrolysis reactions where two functional groups are present in

22

T. H. FIFE

the ester. In all cases the bell-shaped pH-rate constant profiles were shown to be the result of participation by one functional group and a substituent effect by the second group. Higuchi et al. (1971) reported an example of possible bifunctional catalysis in the hydrolysis of the monosuccinate ester of hexachlorophene [ 71 . The monoacetate ester hydrolyses 500 times faster than the diacetate below pH 8, and the monosuccinate hydrolyses

61

C1

[71

lO’-fold faster than the diacetate below pH 5. These relative rates plus a bell-shaped pH-rate profile for the succinate ester were taken to indicate intramolecular nucleophilic-general acid catalysis. However, Oka and Bruice (personal communication) have recently found that methylation of the phenolic group decreases the rat& of carboxylate anion attack by only 3-2-fold. They ascribe the bellshaped profile to inhibition of attack of the succinyl carboxylate anion by ionization of the phenolic hydroxyl group and conclude that no genuine examples of bifunctional catalysis in ester hydrolysis are at present known. Hydrolysis of benzaldehyde disalicyl acetal is characterized by a bell-shaped pH-rate constant profile (Anderson and Fife, 1971a,

yCD 0

II

e H : : .

‘’

-0 : ‘ O D C ‘

PHYSICAL ORGANIC MODEL SYSTEMS: ENZYMATIC CATALYSIS

23

1973). The large enhancement in kObd of 2.7 x l o 9 in comparison with the dimethyl ester establishes that carboxyl-group participation is occurring. The bifunctional mechanism shown in [ 8J involving intramolecular general acid catalysis and electrostatic stabilization of a developing carbonium ion by the carboxylate anion is an attractive possibility for hydrolysis of the monoanionic species, but the monoanion rate constant is only 65 times greater than that for the completely unionized species. Part of this difference is due t o changes in the inductive effect produced by ionization of the carboxyl group. This reaction will be considered in more detail in Section 5 in connection with acetal hydrolysis and the mechanism of action of lysozyme. Jencks (1972) has concluded that concerted bifunctional acid-base catalysis is rare or nonexistent because of the improbability of meeting simultaneously at two sites on reactant and catalyst the conditions of the rule which he has proposed for concerted reactions. The rule states that concerted general acid-base catalysis of complex reactions in aqueous solution can occur only (a) at sites that undergo a large change in pK in the course of the reaction, and (b) when this change in pK converts an unfavourable to a favourable proton transfer with respect to the catalyst, i.e., the pK-value of the catalyst is intermediate between the initial and final pK-values of the substrate site. It is clear that bifunctional catalysis does not necessarily represent a favourable process in aqueous solution even when a second functional group is held sterically in proper position to participate in the reaction. Caution should then be used in assuming that most enzymes are utilizing bifunctional or multifunctional catalysis. Nevertheless, this idea has played a leading role in concepts of enzymatic catalysis. Several examples of intramolecular catalysis of bimolecular attack by a nucleophile have been reported. Bromilow and Kirby (1972) reported an interesting example in hydrolysis of salicyl phosphate dianion. In opposition to the previous mechanistic explanations forthe relatively rapid reaction of the dianionic species (40-200 fold greater than the para-carboxyl derivative) of nucleophilic attack (Chanley and Feageson, 1955) and carboxylate ion stabilization of a proton on the leaving group (Bender and Lawlor, 1963), Bromilow and Kirby proposed intramolecular general acid catalysis by the carboxyl group with substantial P-0 bond breaking but little proton transfer in the transition state. Mechanism [9] was arrived at from

24

T.

€3. FIFE

consideration of the following data: the reaction shows no significant solvent deuterium isotope effect, AS* is close to zero, and substituent effects in the 4-and 5-positions are consistent. The latter data 0

were analysed by means of equation (9) from a plot of o;'log k/ko o z / o l ,giving p1 (carboxyl) = 0.99 f 0.18; and p z (phenol) = 1.74,

us.

values thought impossible to reconcile with the mechanism of Bender and Lawlor (1963). The sensitivity to the leaving group is greater than found for hydrolysis of monoaryl phosphate monoanions where proton transfer is complete and is closer to that for the dianions, showing that transfer of a proton is not complete in the transition state for salicyl phosphate hydrolysis. The p-value for the carboxyl group, -0.99, was taken as equal to the Bronsted coefficient for general base catalysis by the carboxylate anion corresponding to a Bronsted 01 of 0 for the carboxyl group acting as a general acid. Graphical analysis employing a transformed equation can lead to erroneous results with the extended Hammett equation. The proper procedure is by multiple regression analysis of the four-parameter equation (E. Anderson, personal communication) giving p (carboxyl) -0.76. The multiple regression coefficient so obtained is not high (0-87), signifying poor fit of the data to the extended Hammett equation. This is not surprising because the treatment ignores possible substituent effects on both the phenolic oxygen and the carboxyl group. Of great interest, however, was the finding that the reaction of substituted pyridines with the salicyl phosphate dianion was subject t o general acid catalysis by the carboxyl group [ 101 with the rate enhancements as high as 108-fold. Reaction of n-butylamine with methyl salicylate in dioxane is second-order in amine (Snell e t al., 1967), but reaction with methyl p-hydroxybenzoate could not be detected. The aminolysis of phenyl

PHYSICAL ORGANIC MODEL SYSTEMS: ENZYMATIC CATALYSIS

25

salicylate in acetonitrile occurs 132 times faster than reaction with the corresponding methoxy compound; general base catalysed attack of n-butylamine assisted by the neighbouring hydroxyl group [ 111 was proposed (Menger and Smith, 1969). H.

Examples of possible intramolecular general acid-base catalysis were reported by Kupchan et al. (1962). The methanolysis of coprostanol acetate and coprostane 3/3,5/3-diol 3-monoacetate [ 121 in aqueous methanol was conducted in triethylamine-triethylammonium acetate buffer. The rates of methanolysis at constant

buffer ratio were proportional to the concentration of the amine base. The presence of a neighbouring hydroxyl group in [ 121 gave a 300-fold rate enhancement. The following mechanism [ 131 was

26

T. H. FIFE

1131

suggested. Concerted intramolecular general base-general acid catalysed solvolysis of complex alkaloids having a tertiary nitrogen was claimed to proceed by mechanism [ 141 (Kupchan et al., 1966%b).

1141

Phenoxide ion is released relatively rapidly from tetrahydrofurancis-2,3-diol phenyl phosphate [15] at 50' (Usher et al., 1970; Oakenfull et aL, 1967). A cyclic phosphate diester intermediate is formed with pronounced general species catalysis. The pH-rate

0

\

OH

o=eo/

b constant profile for spontaneous hydrolysis (reaction at zero buffer concentration) shows apparent hydronium ion and hydroxide ion catalysis and, of most interest, a pH-independent region from

PHYSICAL ORGANIC MODEL SYSTEMS: ENZYMATIC CATALYSIS

27

approximately pH 4- 6. The apparent hydroxide ion catalysed reaction was thought to be nucleophilic attack of the oxide ion on the phosphate monoanion as shown in (10).

+

eo-

The pH-independent reaction could involve either of two kinetically equivalent possibilities [ 16a and b] . From the observed rate constants it was calculated that, if the mechanism is possibility (a),

then protonation of the phosphate oxygen gives a rate enhancement of l o s [compared to the apparent hydroxide ion catalysed reaction of equation ( l o ) ] . If the mechanism is possibility (b), then nucleophilic attack by an oxygen anion on the phosphate monoanion is lo7 times more favourable than attack by an unionized hydroxyl group. Mechanism (a) was thought most likely. If this mechanism is operative then it is clear that general acid catalysis by a group in the active site of ribonuclease, partially protonating a phosphate oxygen, could be a factor of great significance. Although a detailed report of the buffer-catalysed reaction has not appeared, it is likely that the mechanism involves partial proton abstraction from a neighbouring alcohol hydroxyl group as the

28

T. H. FIFE

oxygen attacks phosphorus [ 171, analogous to a proposed mechanism for ribonuclease. Conformation and Strain Effects

A corollary to the idea that binding to enzyme may bring the substrate into proximity with functional groups in the active site is the concept of “induced fit” which assumes that binding of a specific substrate may induce a conformational change in the flexible enzyme that brings functional groups into proper relationship with the substrate, thereby enhancing the catalytic process (Koshland, 1958). A portion of the binding energy of the substrate would be utilized in bringing about the conformational change. There is evidence suggesting that this idea is correct for certain enzymes. X-ray crystallographic studies of lysozyme (Blake et al., 1965; Johnson and Phillips, 1965) and carboxypeptidase (Hartsuck and Lipscomb, 197 1) have demonstrated movement of groups when a substrate or inhibitor binds to the active site. Conformational changes in the substrate or the enzyme have also been invoked in attempts to explain how an enzyme might increase the ease of bond-breaking or bond-formation by introduction of steric strain. For an excellent discussion of strain and conformational effects in enzyme reactions see Jencks (1969). The rack hypothesis of Eyring was an early example (Eyring e t al., 1954). This proposed that, after binding of the substrate to the enzyme, the bond to be broken is stretched by a conformational change of the enzyme. More recently, it has been suggested (Phillips, 1969; Lowe e t al., 1967) but not proved, that when substrates for lysozyme bind to the enzyme, the hexose ring in subsite D is distorted from the normal chair conformation to a half-chair resembling the conformation of the transition state, thus promoting bond breaking.

PHYSICAL ORGANIC MODEL SYSTEMS: ENZYMATIC CATALYSIS

29

To devise analogues of such effects of conformational changes in chemical models is difficult. Almost invariably studies have been directed at elucidating the mechanism of bond-breaking and -making whilst ignoring possible conformational or strain effects. This is necessary in our present state of knowledge of enzymes and of solution chemistry, but for a complete description such effects must be taken into account. A recent example of a chemical study showing how strain effects could be important in an enzymatic reaction, dealt with the hydrolysis of benzaldehyde di-t-butyl acetal [ 181 (Anderson and Fife, 1971b). As shown by a Stuart-Briegleb model, substantial groundstate strain is present which would be partially relieved in the

transition state as the C-0 bond breaks. The rate of the hydronium ion-catalysed reaction is facilitated in comparison with the rate of hydrolysis of the unstrained benzaldehyde diethyl acetal, but more important, general acid catalysis by buffer acids occurs in contrast with the lack of such catalysis in hydrolysis of benzaldehyde diethyl acetals. Ease of bond breaking is the predominant effect in giving rise to general acid catalysis of acetal hydrolysis (Fife, 1972), and in [18] the bond breaking process has been sufficiently enhanced that general acid catalysis is observable. Distortion of the substrate could, therefore, have important mechanistic consequences as well as an effect on the rate in enzyme-catalysed reactions.

3. GENERAL BASE, NUCLEOPHILIC CATALYSIS: a-CHYMOTRYPSIN a-Chymotrypsin is a proteolytic and esteratic enzyme upon which a monumental amount of experimental effort has been expended (Bender et al., 1964; Bruice and Benkovic, 1966; Jencks, 1969).

30

T. H. FIFE

Chymotrypsin catalysis takes place through a three-step process, equation ( l l ) ,where ES is an enzyme substrate complex which breaks down to give an acylated enzyme intermediate, ES' and P1,

the alcohol or amine portion of the ester or amide substrate. Deacylation then occurs to regenerate the free enzyme and to liberate the carboxylic acid. There is conclusive evidence for involvement of serine-195 and more ambiguous but still convincing evidence for participation by histidine-57 in the catalytic process. This evidence has been extensively reviewed (Bender and Kkzdy, 1964; Bruice and Benkovic, 1966; Bender, 1971). Imidar ole Catalysis

The early problem in regard to the mechanism of action of a-chymotrypsin was to determine how serine and histidine could participate with maximum effectiveness in the hydrolysis of esters and amides. In 1957 it was found that imidazole would act as a nucleophile towards p-nitrophenyl acetate (Bender and Turnquest, 1957a; Bruice and Schmir, 1957). When the pK,-value of the conjugate acid of the phenolic leaving group exceeds that of imidazolium ion by 3 pK,-units, general base assistance of imidazole attack becomes important (Bruice and Benkovic, 1964; Kirsch and Jencks, 1964). With aliphatic esters, imidazole general base catalysed attack of water is the only observable mechanism. The influence of the pK,-value of the leaving group on mechanism reflects the necessity of an attacking species with a basicity comparable to that of the leaving group. Bruice and Sturtevant, (1959) and Bruice, (1959) found extremely facile intramolecular nucleophilic attack by neighbounng imidazole in the hydrolysis of p-nitrophenyl y-(4-imidazoyl)butyrate [ 191. The rate constant for imidazole participation (release of p-nitrophenolate) in this reaction is nearly identical with the rate constant for a-chymotrypsin catalysed release of p-nitrophenolate ion [ 190 min-' ; K 2 in equation (11) at pH 7 and 25'1 from p-nitrophenyl acetate. Comparison of the rate constant for intramolecular imidazole participation t o that for the analogous bimolecular reaction (imidazole attack on p-nitrophenyl acetate) (s-' /M s - l )

-'

PHYSICAL ORGANIC MODEL SYSTEMS: ENZYMATIC CATALYSIS

31

gives a value of 9-4 M for the effective molarity of the neighbouring group (Bruice and Benkovic, 1963). Neighbouring imidazolium ion participation was also observed in the hydrolysis of the corresponding amide (Bruice and Sturtevant, 1959), proceeding with

nucleophilic attack by the basic imidazole group on the protonated amide. However, there was no intramolecular participation in hydrolysis of the methyl ester. This cast serious doubt upon the ability of imidazole to act as a nucleophile in acylation reactions of chymotrypsin, which is a good catalyst for hydrolysis of methyl and ethyl esters of the aromatic amino-acids, tyrosine, phenylalanine, or tryptophan. The y(4-imidazoy1)butyrate system does not afford maximum opportunity for participation since the imidazole ring is not held adjacent to the carbonyl group; nevertheless, nucleophilic participation by imidazole has not to date been observed in cases where the pK,-value of the leaving group is much greater than that of the attacking base. Imidazole will function as a general base in the hydrolysis of acyl-activated esters such as ethyl dichloroacetate (Jencks and Carriuolo, 1961) and esters where the pK,-value of the leaving group is 2-3 units lower than that of ethanol and methanol such as

NHC * CH3

/ \

H

H

It

0

32

T. H. FIFE

2,2-dichloroethyl acetate (Bruice et al., 1962a) or N,O-diacetylserinamide (Anderson et al., 1961; Milstien and Fife, 1968). The mechanism in these examples involves proton transfer in the critical transition state as shown by ratios of rZY2 '/k:Ao of 2-3. Thus, the most likely mechanism is [ 201 or a kinetic equivalent. Intramolecular general base catalysis has been detected in the case of 2-(4-imidazoyl phenyl acetate) (Felton and Bruice, 1968, 1969), mechanism [21a]

being proposed on the basis of the magnitude of TAS' (-8.9 kcal mole-' ) and the D, 0 solvent isotope effect (kH o/kD = 3.23). Intramolecular general base catalysis of hydrolysis (21a) was unexpected since the ester has a phenolic leaving group. Felton and Bruice (1968, 1969) reasoned that, if nucleophilic attack occurred, the leaving phenolate ion group would be properly positioned to attack the intermediate acylimidazole and thereby reverse the reaction. The normally less efficient general base reaction then becomes the favoured pathway, as in hydrolysis of acetyl salicylate (see Section 4). Likewise, Fife and McMahon (1970) explained bimolecular general base catalysis by imidazole (21b) in hydrolysis of 0-(4-nitrophenylene) carbonate (k:; /I$!,' = 3.49) by reversibility of the nucleophilic reaction by neighbouring phenoxide ion. There is now convincing evidence that an acyl chymotrypsin intermediate is formed from both specific and non-specific substrates (Bender and Ktzdy, 1964; Bender e t a/., 1964). This intermediate is undoubtedly an acylserine. Acyl- and phosphorylserine derivatives have been isolated and identified. In view of evidence such as a D20 solvent isotope effect (kH o / k D ) of 2- 3 for both acylation and deacylation (Bender and Hamilton, 1962), alcohol and amine nucleophiles showing little dependence on the pK,-value of the nucleophile in reaction with furoyl enzyme (Inward and Jencks, 1965), and the influence of increasing steric bulk in the acyl group (Fife and Milstien, 1967; Milstien and Fife, 1968, .1969), consistent

'

PHYSICAL ORGANIC MODEL SYSTEMS: ENZYMATIC CATALYSIS

33

with a general base mechanism, it is now generally believed that histidine is functioning as a general base in both acylation and deacylation [equation (13)] or by a kinetically equivalent process. The reaction in equation (13) has been depicted as proceeding with formation of tetrahedral intermediates. 0

II

RC-OR'

0

+ R'OH

The pH-dependence of chymotrypsin-catalysed hydrolysis of formylphenylalanine formylhydrazide and semicarbazide substrates implies the existence of an intermediate between the Michaelis complex and the acyl enzyme which is probably a tetrahedral intermediate (Fersht and Raquena, 1971; Fersht, 1972). A difference was found in the pK-value of the Michaelis complex determined from V , a x and K , measurements. The Iower apparent pK,-value of kcat (6.08) was attributed to a pH-dependent change in rate-determining step for acyl enzyme formation. At low pH the rate-determining step in acylation was considered to be formation of the intermediate from the Michaelis complex, whereas at high pH breakdown of this intermediate to acyl enzyme is rate-limiting. Such a change in rate-determining step is consistent with that found by Blackburn and Jencks (1968) in aminolysis of methyl formate, the microscopic reverse of the dcoholysis of an amide, and Cunningham and Schmir (1967) for intramolecular alcoholysis of y-hydroxybutyranilide. In contradiction, Lucas et al. (1973) have recently reported that, in the case

34

T. H. FIFE

of N-formylphenylalanine formylhydrazide, pK-values influencing V m a X and K , are almost identical. In an investigation of achymotrypsincatalysed hydrolysis of acetyl tyrosine and acetyl tryptophan anilide substrates, Caplow (1969) favoured a scheme in which build-up of a tetrahedral intermediate occurs when an electron-withdrawing substituent is present. A nitrogen isotope effect (1*006,1-010,and 1.006 at pH 6.73, 8.0, and 9-43)has been observed in the chymotrypsin-catalysed hydrolysis of N-acetyl-L-tryptophanamide which requires the C-N bond of the amide to be broken in the rate-determining step (O’Leary and Kluetz, 1972). The isotope effect is similar to that observed for the reaction of amides with hydroxide ion which is known to proceed through a tetrahedral intermediate. The Hammett p-value for a-chymotrypsin-catalysed acylation (k ) of phenyltrimethyl acetates is +1.4 (Bender and Nakamura, 1962), whereas that for hydrolysis of substituted anilides of N-acetyl-Ltyrosine is - 2 0 (Inagami et al., 1965). These values are most easily interpreted as indicating that phenyl esters containing a good leaving group form tetrahedral intermediates in the rate-determining step, but that amide substrates, where the leaving group is poor, hydrolyse with rate-determining breakdown of the tetrahedral intermediate and require general acid catalysis by the imidazolium ion. Protonation in the rate-limiting step would give rise to the negative value of p observed with anilides. Philipp et al. (1973), however, found little dependence of k c a t on pK, for the leaving group for N-acetyl-Ltyrosinamides. All of this evidence supports the existence of tetrahedral intermediates in a-chymotrypsin-catalysed reactions, but it should be noted that 0-exchange with water is not observed in deacylation of cinnamoyl-’ 0-chymotrypsin, in contrast with the hydrolysis of 0-cinnamoyl-N-acetylserinamidewhere such exchange is detected (Bender and Heck, 1967). Lack of exchange in the enzyme reaction could reflect interactions of the tetrahedral intermediate with the protein. Acylation and deacylation in equation (13) proceed through similar transition states. If deacylation occurs through attack of an alcohol molecule R’OH rather than water on the carbonyl carbon atom, then deacylation is the microscopic reverse of acylation. Bender and coworkers (Bender and Kizdy, 1965) have demonstrated the symmetry of the reaction about the acyl enzyme in reactions in which reversibility can be observed.

PHYSICAL. ORGANIC MODEL SYSTEMS: ENZYMATIC CATALYSIS

35

Hubbard and Kirsch (1972) have recently proposed that histidine may act as a nucleophile in a-chymotrypsin acylation reactions of esters having a good leaving group (p-nitrophenol). This suggestion was based on a similarity in p-value for acylation by p-substituted nitrophenyl and dinitrophenyl benzoates and nucleophilic attack on these compounds by imidazole, in contrast with less positive p-values for hydroxide ion catalysis. Hammett p-values for hydrolysis of substituted phenyl esters are given in Table 6 and show little apparent trend. The values for hydroxide ion and alcoholate ions are TABLE 6 Rho Values for Hydrolysis of Phenyl Esters Reaction

+

Imidazole OHtris(Hydroxymethyl) aminomethane anion Pentaerythritol NH3 Hydrazine Trimethylamine

p

Ref.

1.9

n

1.0

0.64 0.98 2.1 2.9 2.6

1-3

b

1.7

c

22

d

1.46

f

0

HN

0

\ / II

0

+

Enz ( K z ) ~

36

T. H. FIFE

TABLE 6-continued

p

Ref.

2.1

g

1.21 0.97 2.01

h, i

1.73 Enz ( k ~ l K $ ~ 1.60 OH2.20

h, i

Reaction

+

Imidazole Enz ( k z / K d e OH-

+ Imidazole

\ /

~ - @ + b N 0 2

\ /

Bruice and Benkovic, 1964. Bruice and Sturtevant, 1959. Bruice and Pandit, 1960a. Caetjens and Morawetz, 1960. orChymotrypsin. f Bender and Nakamura, 1962. g Caplow and Jencks, 1962. Hubbard and Kirsch, 1972. f 33%CH3CN-HzO;r = 0.3. Bruice and Benkovic, 1963. a

*

less than for any other catalyst and may reflect a transition state with little bond breaking. The p-values for the less basic catalysis are more positive, and it will be noted that p for intramolecular carboxylate is comparable or greater than the values for the more basic nitrogen nucleophiles. Substituent effects in the leaving group are much larger in cyclization of monophenyl glutarates and succinates than in acetate ion catalysed hydrolysis of phenyl esters p = 1-1 (Gaetjens and Morawetz, 1960). The p-value will, of course, depend on transition state structure which can be altered upon going from a bimolecular reaction

PHYSICAL ORGANIC MODEL SYSTEMS: ENZYMATIC CATALYSIS

37

to an intramolecular reaction. Thus, there is little basis for drawing mechanistic conclusions from comparisons of p in bimolecular reactions and the intracomplex reactions of an enzyme. Nucleophilic attack on the acyl enzyme p-nitrophenoxycarbonylchymotrypsin, where the leaving group has a pK,-value comparable to that of histidine, takes place by a group in the active site which is most probably histidine (Fife et a1.,1972; Hutchins and Fife, 1972; Bender and Wedler, 1972), showing that nucleophilic attack is sterically possible [equation (14)]. p-Nitrophenolate ion is released

b’

+ -0

HN-N

and a partially inactive enzyme is produced. Nucleophilic attack is, however, markedly hindered; the effective molarity of histidine in the nucleophilic reaction is only 0.2 M in comparison with the rate constant calculated for imidazole attack on the acyl enzyme in a bimolecular reaction (Hutchins and Fife, 1972). In contrast, the effective molarity of neighbouring imidazole in substituted phenyl y-(4-imidazoyl)butyrates ranges from 9-4 M to 33 M (Bruice and Benkovic, 1963). Deacylation also takes place to a nearly equal extent by a mechanism not giving an inhibited enzyme which must be the normal general base mechanism. When the leaving group and the catalysing base are of comparable pK,, nucleophilic attack is the only reaction usually observed in non-enzymatic reactions. If histidine is the nucleophile in deacylation of p-nitrophenoxycarbonylchymotrypsin then it must be located in an environment which seriously restricts its ability to function as a nucleophile. X-ray crystallographic analysis of a-chymotrypsin reveals that the carboxyl group of aspartic acid-102 is hydrogen bonded to histidine-57 in the interior of the protein (Blow et aL, 1969). Blow has proposed involvement of aspartate-1 02 in the catalytic process by means of what was termed a “charge relay” system [22] which would produce a negative charge on the serine oxygen. The pK,value of a carboxyl group in a hydrophobic environment in the

38

T. H. FIFE

P21

interior of a protein might be quite high so that the anionic species could be a powerful general base catalyst, but there is no evidence that this is actually a feature of chymotrypsin catalysis. Polgar and Bender (1969) have instead suggested that the function of aspartic acid-102 is to form a hydrogen bond to histidine-57 and thereby to stabilize on it the proton which is received from serine in the initial step of the reaction. This proton could then be donated to the leaving group when the tetrahedral intermediate partitions to products. Hunkapiller et al. (1973) have recently inferred from NMR data that the pL-value of ca. 7 in serine proteases is actually that of aspartic acid while histidine has a lower pK, ( < 4 ) so that at pH-values from 3-6-7 these residues will be neutral. A mechanism was proposed, utilizing the asp-102-hist-57-ser-195 charge relay system which would avoid charge separation. Rogers and Bruice (1974) have measured rates of hydrolysis of aryl acetates containing a 2-imidazole group with a carboxyl group substituent within hydrogen bonding distance of the imidazole nitrogen [23]. Hydrolysis is by pathways previously determined for esters of this type without carboxyl group substitution [ Z l ] . N o 0 + N

1231

PHYSICAL ORGANIC MODEL SYSTEMS: ENZYMATIC CATALYSIS

39

transfer of the acetyl group was detected. Only a %fold rate facilitation can be attributed to the carboxyl group in the imidazole general base catalysed attack of H2 0. In 94% CH, CN-6% H2 0, no enhancement of catalysis is observed. Unequivocal assignment of pK, -values in both water and organic mixed solvents shows no inversion of carboxyl and imidazolyl base strengths. Numerous attempts have been made to explain why the serine oxygen is such a potent nucleophile in a-chymotrypsin. The “charge relay” mechanism of Blow e t al. (1969) is only one of these attempts. What all such explanations have not taken into account is the enormous effective molarity of an oxide ion nucleophile in an intramolecular reaction ( l o 6 - l o 8 M ; Hutchins and Fife, 1973a, b). Oxide ions are therefore normally powerful intramolecular nucleophiles, and even at the low concentrations determined by their pK, (13.6 in N-acetyl serineamide; Bruice e t al., 1962b) and the pH of the solution, the reaction will be rapid. It is probable that the rates of a-chymotrypsin reactions will eventually be accounted for in terms of the normal chemical properties of the species involved without invoking extraordinary features due to the protein. In view of the reduced ability of histidine to function as a nucleophile even when the leaving group in the reaction is p-nitrophenol and the great efficiency of a hydroxymethyl group as an intramolecular nucleophile (see below), it is not surprising that the role of histidine-5 7 in a-chymotrypsin catalysed reactions should be that of a general base catalyst or its kinetic equivalent.

Hydroxyl Group Catalysis Bender and Glasson (1959), in studies of alcoholysis and hydrolysis of alkyl esters in aqueous alcohol, found that the rate of disappearance of ester is decreased by increasing alcohol concentration. However, product analysis led to the conclusion that both methanolysis and ethanolysis are faster than hydrolysis in alcoholwater mixtures. It was calculated that in pure water attack by hydroxide, methoxide and ethoxidr ions would occur at about the same rates. Bruice and Lapinski (1958) reported that logarithms of secondorder rate constants for reaction of p-substituted phenoxide ions with p-nitrophenyl acetate were a linear function of the pK,-value of the phenol with a slope of 0.8. Phenolate ions cannot displace

T. H. FIFE

40

alkoxide ion from an alkyl ester. Tni-(hydroxymethy1)aminomethane (Jencks and Carriuolo, 1960; Bruice and York, 1961) will displace p-nitrophenolate from p-nitrophenyl acetate. Jencks and Carriuolo favoured a mechanism in which the amino-group acts as an intramolecular general base for transesterification. Bruice and York, however, determined that the reaction involving the amino group was simple aminolysis (It; * '/kFzo = 1.0). Transesterification, k F 2 ° } k t 2 0 = 0.55, occurs by a specific base catalysed pathway CH2 0-

CHz OH

I

H,N--C--CHzOH 1

CH20H

+-OH

I

S NHz-C-CH20H + H2O I CH2 OH

(15)

0

+

[equation (15)]. Pentaerythritol (pK, = 14) also transesterifies p-nitrophenyl acetate in an apparent hydroxide-catalysed reaction. The reaction of oxyanions with p-nitrophenyl acetate was reported to be characterized by a linear plot of log I t , us. pK, of the nucleophile with a slope of 0.78 (Bruice et al., 1962b). N-acetylserinamide, a model for the active site serine of a-chymotrypsin, displays normal reactivity for an alcohol of pK, 13.6. Likewise N,O-diacetyl serinamide undergoes base-catalysed hydrolysis at a rate in accord with the pK,-value of the leaving group and a slope of -0.3 in a plot of log koH vs. leaving group pK, for a series of acetate esters. Jencks and Gilchrist (1962) reported a non-linear plot of log kT us. pK, for reaction of oxyanions with p-nitrophenyl acetate. Jencks and Gilchrist (1968) also reported curved plots of the logarithms of the rate constants us. pK, of oxyanion nucleophiles for esters with varying leaving groups. Curvature is most pronounced with the esters having the poorer leaving groups e.g., phenyl acetate. The slope changes from -1.0 to -0.3 as the basicity of the nucleophile increases. Thus, the sensitivity of the reaction to basicity is greater with weakly basic anions and esters with poor leaving groups.

PHYSICAL ORGANIC MODEL SYSTEMS: ENZYMATIC CATALYSIS

41

These results arise either from a change from rate-determining formation to breakdown of a tetrahedral intermediate or from a continuous change in a concerted reaction in which there is no intermediate. Attack of strongly basic oxygen anions on esters with good leaving groups is characterized by rate constants proportional t o those obtained for hydroxide ion catalysed hydrolysis. The rate constants increase with increasing pK, of the nucleophile and decreasing pK, of the leaving group. Very weakly basic nucleophiles (acetate ion) act predominantly as general base catalysts. The types of transition states shown in [ 241 are in accord, where [a] represents the transition 6-

0

1:

0

6-

1:

6-

6-

RO-------C---OR'-

RO---C-------OR' I

I

R

R

state when pK, (nucleophile, OR-) < pK, (leaving group, OR'-) and [b] the case pK, (nucleophile) > pK, (leaving group). If tetrahedral intermediates are formed (a) would then represent rate-limiting breakdown ( / z - ~> k 2 ) and (b) rate limiting formation (k2 > k - I ) of the intermediate in equation (16). 0 RO-

+

II

RC-OR'

0kl k-1

0

I

k2

I

k -2

RC-OR'

II

RC-OR

+ -OR' (16)

OR

Oakenfull and Jencks (1971) have investigated reaction of oxygen anions with the model amide N-acetylimidazole and N-acetyl imidazolium ion. The pKa-value of the leaving group of the neutral compound is 14.2, and that of the protonated species is 7.0. The reactions are uncomplicated by general acid-base catalysis, unlike corresponding aminolysis reactions. The plot of log k, for reaction of the protonated species us. pK, of the nucleophile shows pronounced curvature with a slope varying from 0 to 1.7. There is little sensitivity of the reaction to base strength for nucleophiles of high pK,. Phenolate and trifluoroethoxide have almost identical reaction rates, in spite of their 100-fold difference in basicity. This indicates that there can be little bond formation in the transition state (25a) (rate-determining attack, if there is a tetrahedral intermediate). A

T. H. FIFE

42

slope of 1.7, for weak bases is identical to the slope of the graph of the logarithms of the equilibrium constants for transfer of the acetyl group to acyl acceptors us. the pK, of the acceptor. Hence, there must be considerable C-N bond breaking in the transition state [25b] for the sensitivity of the rate constants and equilibrium con-

stants to pK, of the nucleophile to be the same. The plot of log k, is reasonably linear with neutral N-acetylimidazole for all nucleophiles and has a slope of 1.3, signifying a transition state resembling products with expulsion of imidazole anion. Toward neutral N-acetylimidazole, trifluoroethoxide is 1O5 times more reactive than phenolate ion. Oakenfull and Jencks conclude that reactions of strong bases with neutral and protonated species proceed by a concerted pathway in one or both cases, or, alternatively, that there is a tetrahedral addition intermediate whose lifetime is too short for it t o reach equilibrium with respect t o proton transfer. Since serine-195 is the site of acylation, a good deal of recent work has been directed towards determining the efficiency of neighboring alcohol and phenol groups as intramolecular nucleophiles in ester and amide hydrolysis. In comparison with acetamide and butyramide, the hydrolysis of y-hydroxybutyramide [equation (17)] in the alkaline and neutral pH-range is accelerated 800-fold (Bruice and Marquardt, 1962). These reactions are attack of alkoxide ion on the neutral and protonated amide function,

us. pK,

Jt

-k

H+

43

PHYSICAL ORGANIC MODEL SYSTEMS: ENZYMATIC CATALYSIS

respectively. Both general base and general acid catalysis was reported (Belke et al., 1971) in the lactonization of 2-hydroxymethylbenzamide [equation (18)] . The Bronsted 0-value for general base catalysis is 1.0, indicating that proton transfer is diffusioncontrolled (Eigen, 1964). A Bronsted plot of log k, us. pKa of the 0

I1

-

-

"H2

~

CH2 OH

0

II

W O + NH3 \ (18)

catalysing base shows apparent curvature, the slope changing t o 0.2 at high pK,. Belke et al. (1971) interpreted this as signifying a change in rate-determining step at about pH 8 from decomposition of a tetrahedral intermediate at low pH to formation of a tetrahedral intermediate at more alkaline values. Okuyama and Schmir (1972) have come t o the opposite conclusion from studies of the decomposition of l-benzylimino-l,3-dihydroisobenzofuran as a function of pH [equation (19)]. A tetrahedral intermediate identical to that of equation (18) is formed. 2-Hydroxymethylbenzamide is the major product at high pH whereas phthalide is the product at low + HNR H+

+

1IHz0

0

I1 C-NHR

HO

e 0

(19)

0

NHR

+ RNH2

CH2 OH

0

II

-0 NHR

-

4

\

d \

o+

RNH2

+ -OH

44

T. H. FIFE

pH, showing preferential amine expulsion from the tetrahedral intermediate at low pH. The conclusion that the Bronsted plot for cyclization of 2-hydroxymethylbenzamide is curved rests heavily on the position of the point for hydroxide ion. A reasonably straight line can be drawn through the other points. It is well known that hydroxide ion generally shows negative deviations in Bronsted plots. Also, Okuyama and Schmir (1972) observed curved plots of k o b s d us. bicarbonate buffer concentration for both reactions (18) and (19), indicating a change in rate-limiting step with increasing buffer concentrations and thereby casting doubt on the Bronsted plot reported by Belke et al. and its mechanistic interpretation.

[26] R = H [27] R = C H 3

I

R

Hutchins and Fife (1973a) studied cyclization of the carbamate esters (26) and (27) having a neighbouring phenolic hydroxyl group. Phenoxide ion is released in the reaction and a benzoxazolinone is produced in quantitative yield. These reactions are very rapid, requiring stopped-flow rate measurements, in contrast with the great stability normally exhibited by carbamate esters. Compound [ 271 cyclized approximately 10 times more rapidly than [ 261 . Aromatic carbamate esters with hydrogen on nitrogen can undergo elimination to give an isocyanate at rates lo6 times greater than for hydrolysis of N-substituted carbamates where such elimination is precluded (Bender and Homer, 1966). In view of the similarity in rate constants for cyclization of [ 261 and [27], it is likely that the mechanism is the same for both, i.e., nucleophilic attack by the oxygen anion on the ester carbonyl. The pH-rate constant profile for cyclization of the N-methylated ester [27] is sigmoid (Figure 1) with pK,,, = 8.9. Hydroxide ion catalysis is not observed, even at pH-values as high as 13. In the case of [26], however, apparent hydroxide ion catalysis is

PHYSICAL ORGANIC MODEL SYSTEMS: ENZYMATIC CATALYSIS

45

detectable at high pH. Buffer catalysis is not observed in any of the cyclization reactions. Therefore the reactions must involve preequilibrium ionization of the phenolic hydroxyl to phenoxide ion as in equation (20).

PH

Figure 1. Plot of kobsd for ring closure of phenyl N-(2-hydroxypheny1)-N-methyl carbamate to N-methyl-2 benzoxazolinone and phenoxide ion at 25" in HzO with /J = 0.5 M (with KCl).

Intermolecular alcoholysis of carbamate esters will also take place, but the reactions are very slow. In comparison with intermolecular nucleophilic attack by a phenoxide ion of the same pK, on the unsubstituted ester [ 281, the effective molarity of the neighbouring phenoxide ion of [27] is 3 x lo8 M. Thus, a phenoxide ion is an

extremely powerful intramolecular nucleophile for attack at the ester carbonyl. A comparable effective molarity ( l o 8 M ) has been determined for neighbouring carboxylate in carboxyl group assisted phenolic ester hydrolysis (Bruice and Turner, 1970). These are the largest effective molarities that have been determined to date for neighbouring groups in ester hydrolysis reactions. T h e largest effective molarity of a neutral nitrogen base is 5 x lo3 M found for the dimethylamino-group of p-nitrophenyl ydimethylaminobutyrate (Bruice and Benkovic, 1963).

46

T. H. FIFE

Part of the great efficiency of the intramolecular reactions of [ 261 and [27] is undoubtedly due to the correct alignment of the rigidly held nucleophile and carbonyl group. Molecular models show that in one of the conformations of [27] in which steric interactions are minimized, the phenoxide ion is immediately adjacent to the carbonyl group and in an excellent position for perpendicular attack (Bender, 1960); but other factors must also be important. Correct orientation would not explain why anionic nucleophiles are superior to neutral nucleophiles. Extensive studies have not been carried out with nitrogen nucleophiles in carbamate ester hydrolysis, but Hegarty and Frost (1972) found that carbamate [29] underwent elimination to an isocyanate. This can be contrasted with the

nucleophilic attack occurring with [26] and in reaction of the analogous carbamate with an ortho-carboxyl group which proceeds with nucleophilic attack by the carboxylate anion (Frost and Hegarty, 1973). Release of phenoxide is approximately 1O6 times faster than in the case where the carboxyl group is in the paraposition. The difference in basicity of the phenoxide nucleophile of [26] and the amine group of [29] makes firm conclusions difficult, but an aromatic amine and a carboxyl should have similar pK,values. A possible explanation for the great efficiency of anionic nucleophiles in intramolecular reactions is that desolvation of the nucleophile is not as energetically unfavourable as in corresponding bimolecular cases where the anion would be highly solvated in the ground state. Desolvation of anions in solution is energetically difficult (Gibson and Scheraga, 1967). Bruice and Turner (1970) found that the effective molarity of neighbouring carboxylate in ester hydrolysis is not altered appreciably on changing the solvent from H20 to lhrl H,O-DMSO, but it is not certain that a carboxylate anion would be completely desolvated in the latter case: Hydroxide ion is highly solvated by HzO in that solvent (Goitein and Bruice, 1972). Any explanation of the efficiency of anionic nucleophiles in intramolecular reactions must at the present time be

PHYSICAL ORGANIC MODEL SYSTEMS: ENZYMATIC CATALYSIS

47

speculative, but the observation of effective molarities of 10' M is striking and undoubtedly of importance in regard to the reactivity of the serine hydroxyl that is acylated in reactions of serine proteases. Participation by neighbouring alkoxide ion is observed in the cyclization of p-nitrophenyl and ethyl esters of 2-hydroxymethyl-Nmethylcarbanilic acid [30] -[33] (Hutchins and Fife, 1973b). As with the carbamates having a neighbouring phenolic hydroxyl, ring R O

I 1I

0

N-C-OR'

[30] R = H;

R' = p-nitrophenyl R' = p-nitrophenyl

CH2 OH

[ 311 R = CH3 ; [32] R = H; [33] R = C H 3 ;

R' = -Et R'=-Et

closure occurs t o give a quantitative yield of the corresponding cyclic compound, and buffer catalysis is not observed in any case. The unmethylated nitrophenyl ester [ 301 releases p-nitrophenoxide ion 1O4 times more rapidly than does the corresponding N-methylated ester [31] , so that it is probable that an isocyanate intermediate is being produced [equation (21)]. Cyclization still occurs quite rapidly with [31], koH , the second-order rate constant for hydroxide H O

a N = ' = O

~301

I

+

- O e N 0 2

CH2 OH

ion catalysis, being 3 x lo5 times greater than k O H for the hydrolysis of p-nitrophenyl N-methyl carbanilate. The value of koH for [31] is lo5 times greater than the second-order rate constant for transesterification of p-nitrophenyl N-methylcarbanilate by pentaerythritol (pK, of the hydroxyl groups of pentaerythritol is 14). Ring closure also occurs rapidly with the ethyl esters [32] and [33]. The k O H values for these reactions are nearly the same and

48

T. H. FIFE

1.3 x lo6 times greater than koH for the hydrolysis of ethyl Nmethylcarbanilate. Both reactions must involve intramolecular nucleophilic attack by the neighbouring alkoxide ion on the ester carbonyl. Plots of log kobsd us. pH for cyclization of these compounds are linear with slopes of 1.0, indicating apparent hydroxide

ion catalysis. This fact, plus the lack of buffer catalysis, shows that pre-equilibrium ionization of the hydroxymethyl group is occurring as in equation (22). According to the scheme of equation (22), K O

is given by equation (23), where K, is the dissociation constant of the hydroxymethyl group and Kw is the ionic product of water. The effective molarity of the neighbouring hydroxymethyl groups in compounds [30]-[33] cannot be calculated, since this would require evaluation of k I from equation (23). This cannot be accomplished because K , cannot be directly measured. However, from the linear pH-rate constant profiles, it is apparent that pK, is no lower than 14. Assuming a pK,-value of 14 for the N-methylated ethyl ester, a lower limit of the effective molarity of the neighbouring group of l o 5 M can be calculated from the comparison with the hydroxide ion-catalysed hydrolysis of ethyl N-methylcarbanilate. Considering that a kinetically less favoured 6-membered ring transition state is being formed, compared to a 5-membered one with the compounds having neighbouring phenoxide, it appears that neighbouring alkoxide and phenoxide ions are comparable in

PHYSICAL ORGANIC MODEL SYSTEMS: ENZYMATIC CATALYSIS

49

catalytic efficiency. Bruice and Pandit (1960b) found that intramolecular carboxylate attack on phenyl esters was 230 times faster with succinate than glutarate monoesters (Table 2). Thus, as with phenoxide ions, alkoxide ions are extremely powerful nucleophiles towards esters but only in intramolecular reactions. Buffer catalysis is observed in the intramolecular tranesterification of ethyl 2-hydroxymethyl benzoate to phthalide [equation (24)] (Fife and Benjamin, 1973). There is general base catalysis by 0

cikEt <\

CHzOH

N+NH

I1

+

+ EtOH CH2

(24)

+ Im

[341

imidazole and a variety of other bases. The D,O solvent isotope effect for imidazole catalysis (kzo/kfm20) is 3-46, showing that proton transfer is occurring in the critical transition state. The mechanism could therefore be closely analogous to the mechanism for acylation of a-chymotrypsin by ethyl esters [equation (13) or a kinetic equivalent]. A Bronsted plot of log k, for general base catalysis us. pK, of the catalysing base is linear (Figure 2) with a slope (p) of 0.87, showing that proton transfer is nearly complete in the transition state. The cyclization reaction to phthalide is quite rapid. The second-order rate constant for apparent hydroxide ion catalysis, k O H , is l o 5 times greater than koH for the hydrolysis of ethyl benzoate (Bender, 1951). Buffer catalysis is not seen in cyclization of the carbamate esters [26], [27], and [30] -[33]. This is possibly because the acyl group is strongly deactivated by the adjoining nitrogen, necessitating nucleophilic attack by a fully developed negative charge. The acyl group of the benzoate ester [34] is not as deactivated and nucleophilic attack will proceed much more readily, allowing proton transfer t o become part of the rate-limiting step. General base catalysis of the cyclization of ethyl 2hydroxymethyl Pnitrobenzoate [ 3 5 ] , where the acyl group is further activated by electron withdrawal by the nitro-substituent, is characterized by a Bronsted coefficient of 0-97, i.e. unity within the limits of error, suggesting that proton transfer is a diffusion-controlled process (Fife

50

T. H. FIFE

4 -

2-

c

I

u

%

c

-5

O-

m 2

OI

-0 -2 -4

CH3COO-

5

7

9

11

13

15

PK, Figure 2. Bronsted plot of log kg us. pK. for the catalysing base in the cyclization of ethyl 2-hydroxymethylbenzoate to phthalide at 30" in H 2 0 (u= 0-5).

and Benjamin, unpublished data). Considering the observed rate constants for the reaction, it is apparent that, if proton transfer is diffusion-controlled, then it must be preceded by an equilibrium step. Thus, the mechanism is general base catalysed decomposition of a tetrahedral intermediate [equation ( 2 5 ) ] or the kinetically equiva-

0

II

+ EtO-

+ BH*

PHYSICAL ORGANIC MODEL SYSTEMS: ENZYMATIC CATALYSIS

51

lent process of equation (26). The rate-determining step for intramolecular carboxylate ion attack on phenyl esters may be breakdown of a tetrahedral intermediate (Goitein and Bruice, 1972), whereas in bimolecular hydroxide ion catalysis formation of a tetrahedral intermediate is rate-determining (Gaetjens and Morawetz, 1960; Bruice and Pandit, 1960b). The attack step in an intramolecular reaction is facilitated in comparison with bimolecular counterparts, and departure of the leaving group may be rate-limiting in cases where the nucleophile is less basic than the leaving group. 0

0

+ H+

0

II

+ EtOH + B From the second-order rate constant for imidazole-catalysed cyclization of the ethyl ester (34) (8 x l o m 3M -ls-l) and the rate constant for acylation of a-chymotrypsin by N-acetyltyrosine ethyl ester (1600 s - l ) , it can be calculated that, in order to attain a rate constant of the magnitude seen in the a-chymotrypsin reaction, a neighbouring imidazole would have to possess an effective molarity of 200,000 M . An effective concentration of this magnitude is not unreasonable, but it is probable that other factors are also important in the enzymatic reaction. Aryl esters of 4-hydroxybutyric acid, 5-hydroxyvaleric acid, 2hydroxyphenylacetic acid, and 3-(2-hydroxyphenyl)propionic acid lactonize with rate constants proportional to lopH - p K w (Capon et al., 1973). The second-order rate constant at 30" for lactonization of phenyl 4-hydroxybutyrate is cu. 3000 times greater than k, for hydrolysis of phenyl acetate at 25". Lactonization of phenyl 4hydroxybutyrate is catalysed by acetate and phosphate buffers in

52

T. H. FIFE

what was considered to be general base-catalysed reactions although no further evidence was presented. Imidazole catalysis of phenoxide ion release proceeds with nucleophilic attack by imidazole at the carbonyl of the ester. In this system therefore, where the leaving group is a phenol, general base-catalysed intramolecular alcoholysis is

not sufficiently advantageous to compete with nucleophilic attack by imidazole. It would be quite important to determine whether this is the case when the alcohol group is rigidly held close to the carbonyl in contrast with the 4-hydroxybutyrate system where degrees of rotational freedom exist. Belke et al. (1971) reported general base and general acid catalysis in cyclization of 2-hydroxymethylbenzamide [equation (18)] . However, with 2-hydroxymethyl-6-aminobenzamidestrict general base catalysis by buffer bases is observed with a Bronsted coefficient of 0 3 9 (Fife and Benjamin, unpublished data). In contrast with the unsubstituted amide, the Bronsted plot is nicely linear. An aminogroup in the 6-position might assist decomposition of a tetrahedral intermediate as in [37a, b] or a kinetic equivalent. The pH-rate constant profile for spontaneous cyclization at zero buffer concentra-

tion is very similar in shape to that of 2-hydroxymethylbenzamide. There is hydronium ion catalysis, hydroxide ion catalysis, and a pH-

PHYSICAL ORGANIC MODEL SYSTEMS: ENZYMATIC CATALYSIS

53

independent reaction, but the values of kobsd for the pH-indepdndent reaction at 30" are 10 to 100 times less than those for 2-hydroxy__ rnethylbenzamide at 40'. The value of k o H is lo3 timesless than that for the 4-amino-derivative (0 = 0.43) where the effect of the substituent group is entirely electronic. Thus, a 6-amino-group greatly reduces the rate instead of promoting the reaction.

2

4

6

8

1 0 1 2

PH

Figure 3. Plot of log hex for 3-amino-24tydroxymethybenzamide (rate constant obtained by extrapolation to zero buffer concentration) us. pH at 30" and p = 0.5 in 50% dioxane-H2O (v/v) 0 and 50%dioxane-D2O 0 . The rate constants are for appearance of product (4aminophthalide).

The neighbouring amino-group of 3-amino-2-hydroxymethylbenzamide facilitates the pH-independent cyclization between pH 3.5 and 10 (Fife and Benjamin, 1974), as seen in Figure 3. That reaction proceeds lo3 times faster than in the case of the 6-amino-, 5-amino-, or unsubstituted derivatives. There is no clearly defined pH-independent reaction with the 4-amino-substituted amide because of relatively rapid hydroxide and hydronium ion catalysis. The point of intersection of these limbs on the pH-rate constant profile, which gives a maximum rate constant for a possible pHindependent reaction, is s-l at 30°, l o 3 times less than for the 3-amino-derivative. An amino-group in the 3-position could interact only with the hydroxymethyl group. Intramolecular general base catalysis is indicated by the D20solvent isotope effect ( k ~ 2 0 / k ~ z o = 2.82) and the lack of bimolecular general base catalysis observed with the other 2-hydroxymethylbenzamides in the series. The effective molarity of the neighbouring group, in comparison with catalysis by an amine of pK, 3.5 in cyclization of the 4-amino-amide, is 15 M . A Stuart-Briegleb model shows that the stenc situation is

54

T. H. FIFE

not favourable for the concerted process shown in [38]. Substantial restriction of rotational freedom would be required, as reflected in the AS* of -23 e.u. for pH-independent cyclization, thus explaining

the small effective molarity of the intramolecular general base. Such restriction is perhaps an important reason why bifunctional mechanisms are uncommon. General base catalysis through one or more solvent molecules [39] would be in accord with the highly negative AS* and would avoid the steric complications of mechanism [38]. The steric situation would

be more favourable if the 3-amino group of [38] is replaced by a 3-methylamino-group. A seven-membered ring for proton transfer would be required, but since the hydroxyl group must approach the carbonyl carbon in a concerted reaction the resulting steric fit would be excellent. The 3-methylamino-compound is at present being studied in our laboratory. Neighbouring phenoxide ion will act as an intramolecular general base in situations where nucleophilic attack is precluded. Bender et al. (1963) found such catalysis in the hydrolysis of p-nitrophenyl 5-nitrosalicylate. Mechanism [401 was favoured over the kinetically equivalent hydroxide ion attack on the neutral species, the reason

PHYSICAL ORGANIC MODEL SYSTEMS: ENZYMATIC CATALYSIS

55

being that charged nucleophiles, such as fluoride or azide, the action of which might be aided by general acid catalysis by the unionized hydroxyl group but not by general base catalysis by the ionized

species, do not have an enhanced rate constant for reaction with the neutral species. Similar evidence and reasoning was employed in assigning an intramolecular general base mechanism t o hydrolysis of ethyl 2-hydroxy-5-nitrophenyl carbonate [4 11 (Fife and Hutchins, 1972). Intramolecular general base catalysis in these systems

produces only moderate rate increases in comparison with spontaneous hydrolysis of the neutral compound, 30-fold with p-nitrophenyl 5-nitrosalicylate and 50-fold in the case of the carbonate ester. Bruice and Tanner (1965) have also reported intramolecular general base catalysis in hydrolysis of o-hydroxybenzamide [ 421.

The number of well-established intramolecular general basecatalysed reactions is quite small, but from the available examples a tentative conclusion is possible that such a mechanism is not greatly

56

T. H. FIFE

favoured over bimolecular counterparts and will not afford large rate increases. The largest is l o 3 in hydrolysis of [Zla] . Involvement of solvent in the general base-catalysed reaction may minimize the entropic advantage of an intramolecular reaction. It would appear that rate enhancements in ester hydrolysis reactions comparable to those seen with enzymes will require intramolecular nucleophilic attack with the leaving group breaking away from the remainder of the molecule so that AS* will be favourable. Numerous suggestions have been made that enzymes might owe part of their catalytic efficiency to the opportunity they afford for stabilization of intermediates or transition states by hydrogen bonding to functional groups near the active site. For example, in the case of &-chyrnotrypsin this might be represented as in [43] where O R

0

II I

-N-C-CH-pC-

II

H

[431

the transition state for acylation of chymotrypsin by a tyrosyl ester is being stabilized by hydrogen bonding from amide groups of the protein chain (Henderson, 1970; Robertus et al., 1972). Several examples of facilitation of alkaline hydrolysis of esters by a neighbouring aliphatic alcoholic hydroxyl group have been reported (Henbest and Lovell, 1957; Kupchan et al., 1959; Kupchan and Johnson, 1956; Kupchan and Narayanan, 1959; Kupchan et al., 1960; Zachau and Karau, 1960; Bruice and Fife, 1962). Bruice and Fife (1962) concluded that the neighbouring-group effect on alkaline hydrolysis of acetate esters of cyclopentanediols is due to stabilization of the transition state [44] by internal hydrogen bonding. There was no correlation between the observed rate and ground state hydrogen bonding either with regard to the

PHYSICAL ORGANIC MODEL SYSTEMS: ENZYMATIC CATALYSIS

57

position (carbonyl or alkoxyl oxygen) or to the strength of the hydrogen bond determined by high dilution infra-red studies. The effect must therefore be on the transition state of the reaction. The

influence of the neighbouring hydroxyl group is to make AS* much more favourable, whereas AH* is actually unfavourable in comparison with reaction of the reference compounds cyclopentyl acetate and truns-2-methoxycyclopentylacetate. A similar change of the activation parameters is noted when an ester hydrolysis is transferred from the solvent water to alcohol-water mixtures. Thus the effect of the neighbouring hydroxyl group in the cyclopentanediol acetates can be looked upon as an internal solvation. The alcoholic hydroxyl group would not solvate the transition state as well as water, resulting in a less favourable AH*,but fewer water molecules would have to be restricted in the transition state giving a much more positive AS*. Rate enhancements in the cyclopentanediol series are fairly small. In comparison with hydrolysis of cyclopentyl acetate, cis-l,2-cyclopentanediolmonoacetate hydrolyses 33 times and the truns-isomer 1 7 times more rapidly. trans-2-Methoxycyclopentyl acetate hydrolyses 3 times faster than cyclopentyl acetate. Inductive effects of neighbouring hydroxyl and methoxyl groups are nearly identical (Taft, 1956). While small, the rate increases due to a neighbouring hydroxyl group are significant and show that transition state stabilization is chemically reasonable and quite possibly a factor in enzymatic catalysis. The magnitude of the acceleration that might be obtained from such an effect will, of course, depend upon the nature of the transition state. Transition state stabilization would be of the greatest importance in reactions having highly solvated polar transition states. This is the case in ester and amide hydrolysis reactions when bond-making with a nucleophile or bond breaking in decomposition of a tetrahedral intermediate has progressed to a substantial extent.

T. H. FIFE

58

A neighbouring hydroxyl group provides little rate enhancement in the general base-catalysed hydrolysis of acyl-activated esters (Bruice et al., 1962a). The imidazole-catalysed hydrolysis of trans-2hydroxycyclopentyl dichloroacetate is only 2 times faster than that of the 2-methoxy-substituted ester. A consequence of the strong acyl activation is that in the transition state for formation of the tetrahedral intermediate there would be little bond-making with the attacking nucleophile and therefore little charge development on the carbonyl oxygen. Acyl group activation would greatly facilitate attack of a nucleophile. Cyclodextrins are cyclic polymers of glucose having 1 , k t - D glucopyranose units in a staggered chair conformation (Hybl et al., 1965). The molecules are cylindrical in shape. The 2- and 3-hydroxyl groups are located around the top of the cylinder and the 5-hydroxy methyl groups at the bottom. The cyclodextrins are of interest because of their ability to form inclusion compounds within the cavity with a variety of organic molecules. With suitable substrates such binding may be followed by reaction of the substrate with a hydroxyl group of the cyclodextrin. Thus, these compounds can be considered unique model enzymes. Spectral shifts due to binding of dyes indicate that the interior of the cavity is much less polar than H, 0. The hydrolysis of substituted phenyl acetates has been studied in the presence of cyclodextrins (Van Etten et al., 1967a, b). No correlation was found between the rate constants for hydrolysis and 0 for the substituent group. Specificity was directed towards metasubstituents. m-t-Butylphenyl acetate hydrolyses 240 times faster in the presence of 0.01 M cyclohepta-amylose. Comparison of spectral shifts upon inclusion of p-t-butyl and m-t-butylphenol indicated that benzene rings of p-substituted phenols are included within the cavity of cyclodextrins [45] but that the benzene ring of the meta-isomer )

PHYSICAL ORGANIC MODEL SYSTEMS: ENZYMATIC CATALYSIS

59

[46] extends into the aqueous solvent. 4-Nitrophenol binds to cyclos-l (Cramer et at., dextrin with a rate constant of about 10' M 1967), whereas the rate constant for binding of 1,3-substituted compounds is only about 10 M s-'. A double displacement reaction (27) was shown to occur with aryl benzoates (Van Etten, 1967b).

-'

-'

J

9 NO*

0

-

+

fJ

In accord with the above scheme, variation of the meta-substituent altered the rate of acylation but not that of deacylation. The reaction of equation (27) is similar to reaction of a-chymotrypsin with phenolic esters.

60

T. H. FIFE

Brass and Bender (1973) have recently studied reaction of bis (ib-nitrophenyl) carbonate with cyclohepta-amylose. MichaelisMenten kinetics were observed, and two equivalents of p-nitrophenol released. Release of the first equivalent of p-nitrophenol is ratedetermining, the second one being released in a fast step. The latter

+ -

O

G

t

X

-

O

O

x

process is an intramolecular reaction of a second adjacent hydroxyl group on the cycloamylose forming a cyclic carbonate diester, demonstrating once more the effectiveness of a hydroxyl group as an intramolecular nucleophile.

Mechanistic Conclusions The mechanism given in equation (13) for a-chymotrypsin is almost certainly correct in its essential details. The evidence is strong and for the most part convincingly in favour of general base catalysis by histidine in both acylation and deacylation or a kinetic equivalent. Wang and Parker (1967) have argued for a kinetic equivalent possibility in which the serine anion is the nucleophile, attacking a partially protonated substrate, but Williams (1970), Fersht (1971), and Phillip et al. (1973) have discounted this mechanism. The only serious chdlenge to the picture of equation

PHYSICAL ORGANIC MODEL SYSTEMS: ENZYMATIC CATALYSIS

61

(13) in recent years is the suggestion by Hubbard and Kirsch (1972) that acylation of the enzyme by phenolic substrates entails nucleophilic attack by histidine. Comparative p-values for the enzyme reaction and for imidazole attack on substituted nitrophenyl and dinitrophenyl benzoates do not however provide compelling evidence. Nucleophilic attack does take place in deacylation of p-nitrophenoxycarbonyl-a-chymo trypsin, but is much less favourable than in corresponding intramolecular reactions (Hutchins and Fife, 1972). In view of the extremely large effective molarity of a neighbouring alcohol group in an intramolecular reaction, it is likely that the serine hydroxyl is the attacking nucleophile, regardIess of the nature of the leaving group. A problem in the chymotrypsin mechanism concerns the nature of general base catalysis. Does histidine partially abstract a proton from the serine hydroxyl? Does it catalyse decomposition of a tetrahedral intermediate formed by nucleophilic attack of the serine oxygen anion? Or does histidine catalyse both nucleophilic attack by serine and decomposition of the tetrahedral intermediate? When the leaving group is poor as with amides and esters of aliphatic alcohols, catalysis of the decomposition of a tetrahedral intermediate to product may be required (as in the general base-catalysed intramolecular alcoholysis of the ester [35] ). When the leaving group is a phenol, it is probable that formation of the tetrahedral intermediate will be rate-determining, with acylation assisted by catalysis of nucleophilic attack. The positive p-values for acylation by aryl acetates and trimethylacetates do not require electrophilic catalysis, but Williams (1970) found little sensitivity of k c a t and k c a t / K mto substituents in the leaving group of aryl hippurates. Bender et al, (1964) made an attempt to explain the rate constants attained in a-chymo trypsin reactions as follows for N-Acetyl-Ltryptophan amide: Rate constant for hydroxide ion-catalysed hydrolysis

3x

1 0 - 4 ~ u r - ls-l

(1) Conversion to an intermolecular general base-catalysed reaction involving imidazole (1.6 x 4.8 x 1 0 - ' o ~ - ' s - l

(2) Conversion to an intramolecular general base-catalysed reaction involving imidazole 4.8 x 10-9 (10 M )

s-l

62

T. H. FIFE

(3) Change in rate determining step ( l o 2 )to an alcoholysis (4) Freezing of substrate specificity ( l o 3 ) (5) General acid catalysis by imidazole ( l o 2 ) Experimental rate constant

4.8 x 10-7 s-l 4.8 x 10-4 s - l 4.8 x s-’ 4-4x s-1

This type of analysis is desirable in enzymatic reactions. However, at the time of its presentation, only very rough estimates could be given of the effect of some of the presumed factors on the rate. From work on ester and amide derivatives of 2-hydroxymethyl benzoic acid, more accurate numbers can now be ascribed to these factors. Thus, conversion to an intramolecular alcoholysis will give a rate enhancement of lo5 times in comparison with hydroxide ion-catalysed ester and amide hydrolysis, rather than the factor lo2 estimated by Bender. The effective molarity for an amino-group acting as an intramolecular generd base in intramolecular dcoholysis of [38] is only 15 M. Bender’s estimate of the effective molarity of histidine-57 is therefore possibly close, but the unfavourable alignment of groups in [38] for a concerted reaction leads to the conclusion that much more efficient reactions would occur if orientation were favourable as in the active site of an enzyme. Nevertheless, on the basis of model work, there is a partly unexplained factor of 10-lo2 in the rates of acylation of a-chymotrypsin by amide substrates. Bender and coworkers (1964) found the difference in rate constants for deacylation of acetylchymotrypsin and Nacetyltyrosylchymotrypsin to be entirely in a more favourable AS* for deacylation of the acyl enzyme derived from the specific amino-acid substrate. It was postulated that the aromatic acyl group binds in a specific site so that the resulting steric situation is more favourable for the ensuing reaction. Bender termed this “freezing of substrate specificity”. The effect appears to be similar to that of geminal substitution on intramolecular nucleophilic attack by the carboxyl group in glutarate monoesters which is also manifested in a more favourable ASq (Bruice and Bradbury, 1968). Thus achymotrypsin exerts an effect on specific substrates that is akin to proximity effects in simple intramolecular reactions and which enhances the rate by a factor of approximately lo3 in comparison with nonspecific substrates. In like manner, increased steric bulk in the acyl group of acyla-chymotrypsin progressively retards the rate of deacylation [ 6 in the

PHYSICAL ORGANIC MODEL SYSTEMS: ENZYMATIC CATALYSIS

63

Taft equation, log ( k 3 / k 0 ) = p*o* + 6 E s , is -1.0; Fife and Milstien, 19671 , but hexanoylchymotrypsin displays an enhanced rate of deacylation. The point for the hexanoyl group lies well above the line in a plot of log ( k 3 / k 0 ) us. E s , probably because of binding of the hexanoyl group in a hydrophobic site. Furthermore, the more rapid reaction of hexanoyl- than acetyl-chymotrypsin is again entirely due to a more favourable AS*. Compensation may take place in deacylation of acylchymoirypsins with straight-chain acyl groups, AH* becoming more unfavourable as AS* becomes less negative. A plot of AH* us. AS* is linear (Fife and Milstien, 1967). This was confirmed by Martinek et al. (1972) who also found a parallel linear plot in deacylation of N-acyl amino-acid acyl enzymes. The difference in AS* for corresponding acyl groups in the .two series is 1 2 e.u. Lumry and Rajender (1972) have observed compensation between AH* and AS * in deacylation of N-acetyltyrosylchymotrypsin as the pH is varied. Among the questions of importance remaining to be answered concerning the chemistry of the a-chymotrypsin reactions are the following: (1) What is the effective molarity of imidazole as an intramolecular general base in intramolecular alcoholysis and hydrolysis? This is perhaps the most important single piece of information that is currently missing for analysis of the rates of a-chymotrypsin reactions in terms of specific catalytic effects. (2) What is the explanation for the large effective molarity of a neighbouring alcohol function in ester and amide hydrolysis? (3) Of what magnitude are orientation effects due to binding of substrate to enzyme? Certainly no more than a factor of lo3 need be ascribed to such effects. (4)What contribution does transition state stabilization make? From the chemical work cited, it is probable that such an effect is small. ( 5 ) Does the charge relay mechanism play an important role and, if so, what rate enhancement would such a mechanism provide? It appears that it will not be necessary to invoke a charge-relay mechanism t o account for the rates of achymotrypsin reactions in terms of known chemistry. The presence of aspartic acid in the interior of serine proteases could, of course, have structural rather than mechanistic significance.

T. H. FIFE

64

4. METALION CATALYSIS: CARBOXYPEPTIDASE A Carboxypeptidase A is a hydrolytic enzyme containing zinc bound at the active site. Carboxypeptidase A catalyses hydrolysis of ester and amide linkages of N-acyl-a-amino-acids and 0-acyl-a-hydroxy acids adjacent to free terminal carboxyl groups (Hofmann and Bergmann, 1940; Snoke, et al., 1948). The complete amino-acid sequence is known (Bradshaw et al., 1969) as well as the threedimensional structure from x-ray crystallographic studies at 2 A resolution (Lipscomb et al., 1968; 1970; Lipscomb, 1970). Such studies were also carried out with glycyl-L-tyrosine bound t o the active site. The following conclusions were drawn: (I) the zinc ion is coordinated to the carbonyl group of the substrate; (2) the carboxyl group of glutamic acid-245 is located in close proximity and is presumably involved in the catalytic action of the enzyme; (3) the phenolic hydroxyl group of tyrosine-248 is also located close t o the substrate. Structure [47] can then represent the active site of carboxypep tidase.

glu-245-c

/

0-

O \

[471

Lipscomb has commented that glutamic acid-245 might act either as a general base or a nucleophile. The available mechanistic information has been reviewed by Kaiser and Kaiser (1972), who postulated that the carboxylate anion of glutamic acid-245 acts as a nucleophile forming an anhydride intermediate [equation (29)]. The divergent DzO solvent isotope effects, k,"$o/k:a$o = 1.33 for the peptide substrate N-(N-benzoylglycy1)-L-phenyl alaninate

PHYSICAL ORGANIC MODEL SYSTEMS: ENZYlL4TIC CAT.4LYSE

63

and 2.0 for the specific ester substrate O-(trans-cinnamoyl)-L-& phenyl-lactate, were explained by rate-limiting formation of the anhydride in the case of the peptide substrate and rate-limiting breakdown of the anhydride with the ester substrate. It is known that spontaneous hydrolysis of anhydrides proceeds with a sizable D20 solvent isotope effect (Butler and Gold, 1961), whereas rate-limiting nucleophilic attack should give a rate constant ratio in the two solvents near unity. It must, of course, be borne in mind that mechanistic interpretation of D20 solvent isotope effects in enzymatic reactions is difficult (Jencks, 1963). Co(II1) carboxypeptidase., which /R2

I,:="[ exchanges ligands slowly, is fully active in ester hydrolysis but has lost all peptidase activity (Kang and Storm, 1972). Thus, mechanisms for ester hydrolysis through a substrate-metal ion complex are quest ionable. There are three fundamental questions concerning the mechanism of action of carboxypeptidase: ( 1 ) How do metal ions affect ester and amide hydrolysis? (2)Under what conditions will a neighbouring carboxyl group participate in ester and amide hydrolysis and what is the mechanism of such participation? (3) How will a metal ion affect

66

T. H. FIFE

intramolecular carboxyl group participation in ester and amide hydrolysis? These are questions pertaining to the chemistry of ester and amide hydrolysis which should be answerable from model studies. There is considerable information concerning questions ( 1) and (Z), but the experimental difficulties in attempts to answer question (3) are formidable, for example, finding a system where carboxyl-group participation is possible but where the metal ion does not bind the carboxyl group. Metal-ion Catalysis of Ester and Amide Hydrolysis

Metal-ion catalysis has been extensively reviewed (Martell, 1968; Bender, 1971). It appears that metal ions will not affect ester hydrolysis reactions unless there is a second co-ordination site in the molecule in addition to the carbonyl group. Hence, hydrolysis of the usual types of esters is not catalysed by metal ions, but hydrolysis of amino-acid esters is subject to catalysis, presumably by polarization of the carbonyl group (Kroll, 1952). Cobalt (11), copper (11), and manganese (11) ions promote hydrolysis of glycine ethyl ester at pH 7-3-7-9 and 25", conditions under which it is otherwise quite stable (Kroll, 1952). The rate constants have maximum values when the ratio of metal ion to ester concentration is unity. Consequently, the most active species is a 1 :1 complex. The rate constant increases with the ability of the metal ion to complex with amines. The scheme of equation (30) was postulated. The rate of hydrolysis of glycine ethyl R

H0

R'CH-C

EtOH+

I

1

'.+ p

H2N,

M

#'

-

R -0

I

I I

CH-C-OEt

/ H2N.

0

',+ /'

M

ester and phenylalanine ethyl ester in glycine buffer depends on the concentration of glycine co-ordinated to Cu++. Bender and Turnquest (1957b) found ratios of khYdrolysis/kexchange = 3.9 for

PHYSICAL ORGANIC MODEL SYSTEMS: ENZYMATIC CATALYSIS

67

Cu++-catalysed hydrolysis and 5.4 for alkaline hydrolysis of 180-labelled ester. Formation of the tetrahedral intermediate in the scheme of equation (31) is indicated by the observed 80-exchange. However, interactions of the carbonyl oxygen and the metal ion

hydrolysis

exchange

might not permit the tetrahedral intermediate to become completely symmetrical so that the ratio of 12hydro,ysis/12exchange need not have been expected. An alternative mechanism could involve intramolecular attack of Cu++-bound hydroxide ion (Bruice and Benkovic, 1966). There is a difference of approximately lo6 in rate between the Cu+ -catalysed and alkaline hydrolyses of dl-phenylalanine ethyl ester. Comparison of rate constants for hydrolysis of phenylalanine ethyl ester at 25" and pH 7-3 for catalysis by hydronium ion, hydroxide ion, and Cu++ (0-0775 M ) is of interest in showing the facility of the reaction promoted by the metal ion. Under these conditions the respective rate constants are 1-46 x lo-' s - l , 5.8 x s - l , and 2.67 x s-l (Bender and Turnquest, 195713). The following relative second-order rate constants have been obtained for hydroxide ion-catalysed hydrolysis: glycine ethyl ester, 1; protonated glycine ethyl ester, 41; and the cupric ion complex of glycine ethyl ester, 1.3 x l o 5 (Conley and Martin, 1965). The large effect of the cupric ion cannot be due entirely to electrostatic effects, but rather to catalysis by direct co-ordination with the ester function. Metal-ion catalysis in these systems could take place by coordination of the metal ion with the amino-group and either the +

68

T. H. FIFE

carbonyl oxygen [48a] or the alkoxyl oxygen [48b]. In either case the metal ion would polarize the carbonyl group, thereby promoting attack of hydroxide ion. Coordination to the carbonyl oxygen would

also stabilize the transition state for attack of hydroxide ion. Catalysis might also be explained if the metal ion binds both substrate and hydroxide ion, as suggested for the hydrolysis of ethyl oxalate and ethyl malonate catalysed by calcium, barium and thallium(1) (Hoppe and Prue, 1957). Metal ions will catalyse hydrolysis of a number of amides where the metal ion can complex to another functional group besides the amide group (Meriwether and Westheimer, 1956). Metal ion effects are apparently not as large in the hydrolysis of amides as in that of esters; glycinamide hydrolyses only 30 times faster in the presence of 0 0 2 M copper(I1) ion at pH 7-9-9-35. This must be due to the different leaving groups. If the effect of metal ion is to facilitate nucleophilic attack at the carbonyl group, then the observed rate enhancement could be less in amide hydrolysis where the poor leaving group makes breakdown of a tetrahedral intermediate ratedetermining, the tetrahedral intermediate partitioning preferentially to starting material. Ni++-ion produces a 400-fold enhancement in the hydroxide ion-catalysed hydrolysis of 1,I O-phenanthroline-2carboxamide (Breslow, et al., 1967). Much larger rate accelerations are seen in the hydration of phenanthroline carbonitrile, Ni++giving lo', C u + + lo9, and Z n + + 104-fold enhancement. The large rate increases are due entirely to more favourable AS* values. In all of these studies the free metal ion is in equilibrium with the complex species and ligand exchange can occur. Stable cobalt(II1) ion complexes with which ligand exchange does not occur were prepared by the scheme of equation (32) (Alexander and Busch, 1966; Buckingham et al., 1967a, b, c), and were isolated and characterized. Large rate enhancements in hydrolysis are observed in comparison with the uncomplexed esters for attack of H z O at pH 1-4, but it is of greatest interest that the reactions are markedly catalysed by general bases. It has been observed that

PHYSICAL ORGANIC MODEL SYSTEMS: ENZYMATIC CATALYSIS

c1

C1

69

0

I

catalysis by both oxygen and nitrogen bases takes place in the hydrolysis of cobalt-chelated glycine isopropyl ester (Buckingham et al., 1970a, 1972; Alexander and Busch, 1966). In view of the different Bronsted 0-values (0.4 for oxygen bases and 0-8 for nitrogen bases) and the large D20 solvent isotope effect for catalysis by chloroacetate, ( K D o / k , 0 = 5 , it was originally concluded that catalysis by nitrogen bases followed a nucleophilic mechanism but that oxygen bases functioned as classical general bases. However, it was later shown by competition experiments that both types of bases were nucleophiles (Buckingham, et al., 1972). It should be noted that the metal ion in these systems is not exerting a catalytic effect since the metal ion is incorporated during synthesis of the compound to give a highly reactive molecule, tantamount to synthesis of a protonated ester. Glycine amide chelated through its carbonyl oxygen and terminal nitrogen atoms and peptide complexes of the type C CON^ (glyNR1R 2 ) ] show lo4 to 106-fold rate increases for hydroxide ioncatalysed hydrolysis compared to the unco-ordinated amides (Buckingham e t al., 1970b). Amides in this system hydrolyse l o 5 - lo6 times more slowly than esters. Monodentate amides and esters show only small rate enhancements. Buckingham calculates that the carboxypeptidase-catalysed hydrolysis of benzoylglycyl-r.phenylalanine is -lo4 times faster than would be afforded by C o + + + activation. It is unlikely that divalent Z n + +would be more effective at polarizing the carbonyl group of a peptide than Co+++.Although direct activation of a carbonyl group by metal ions can result in large + + +

70

T. H. FIFE

rate enhancements for hydrolysis (-lo6), the enzyme must also be exerting an extra effect. Hydrolysis of the cis [Co(en);?OHglyNH2)]++ion [49] requires a rate at least l o 7 times greater than hydrolysis of the unco-ordinated glycine amide (Buckingham et al., 1 9 7 0 ~ )Thus . amide hydrolysis via coordinated hydroxide is rapid. Buckingham points out that, if attack of bound hydroxide is rate-determining, (1) the rate of hydrolysis will be maintained at biological pH-values since the complex will exist mainly in the hydroxo form (pK, 6); and (2) despite the greatly reduced basicity of coordinated hydroxide (-lo8), it is a much more efficient nucleophile than hydroxide ion from the solvent. If loss of ammonia from a tetrahedral intermediate were rate-determining, a second-order rate constant 1O7 times greater than that for the hydrolysis of glycine amide would be required. The Co +++-induced intramolecular hydrolysis reaction is much more efficient than direct metal-ion polarization of a carbonyl group. A increases the effective zinc-hydroxide mechanism in which Zn

-

+

0

-

II

+%NHzCH2 C-NH2 (en)2Co,0H

+

+ 3 / NH2 \

(en)2Co

b-c

CH2

I

+

[491 0

(33)

\o

NH3

concentration of bound hydroxide was considered as a possibility for carb oxypep tidase . The Ni++- and Zn++-complexes of phenyl and salicyl esters of pyridine-2,6-dicarboxylic acid [ 50a and b] hydrolyse much faster than the uncomplexed esters (Breslow and McAllister, 1971); the rate of hydroxide ion catalysis of the Ni++-complexesincreases 9300

and 3100 times, respectively. The Ni++-complex of the salicyl ester hydrolysed 1.7 times more rapidly when the salicyl carboxyl group was ionized than when it was unionized. On this basis it was

PHYSICAL ORGANIC MODEL SYSTEMS: ENZYMATIC CATALYSIS

71

suggested that the carboxylate anion was exerting a weak catalytic effect. A metal ion complex of pyridine carboxaldoxime [51] is an excellent catalyst for deacetylation of 8-acetoxyquinoline-8-sulfonate

(Breslow and Chipman, 1965). The metal-nucleophile combination is quite reactive toward a coordinating substrate compared with p nitrophenyl acetate. Reaction of the m-nitrophenyl ester of pyridine-2,5-dicarboxylic acid with cyclodextrin (see Section 3) gives a picolinate ester [52] of a cyclodextrin secondary hydroxyl group (Breslow, 197 1; Breslow and Overman, 1970) which will bind metal ions or a metal ionpyridine carboxaldoxime complex. Such a complex will catalyse hydrolysis of p-nitrophenyl acetate bound within the cyclodextrin cavity leading to a rate constant approximately 2000-fold greater at

pH 5.11 than in the absence of catalyst. Binding to cyclodextrin gives only a modest rate enhancement factor of 4 in comparison with the Ni++-pyridine carboxaldoxime-complex acting alone. A Cu++cyclodextrin picolinate ester will cause p-nitrophenyl glycinate to hydrolyse about four times faster than in the presence of C u + + alone. Although showing modest effects, these studies are of interest

72

T. H. FIFE

as attempts to demonstrate the influence of a metal ion in systems where an ester is hydrophobically bound in proper steric relationship with a nucleophilic group. Sigman et al. (1972) have reported an example of metal-ion facilitated phosphorylation of an alcohol by adenosinetriphosphate (ATP). Complexing of the metal ion, ATP, and 2-hydroxymethyl-

phenanthroline (1:1:1) brings the terminal phosphoryl group of ATP and the hydroxymethyl group into the proper relationship for the ensuing reaction. Reaction takes place through the alcoholate anion and is analogous to an intramolecular alcoholysis. Sigman reasons that complexing takes place as shown in [53a] or [53b]. The metal

A-O-

ion serves several apparent functions: (1) it binds Z-hydroxymethylphenanthroline and ATP together; (2) it promotes attack of the alcoholate ion by polarizing the substrate or by screening the negative charge of the phosphoryl group; and (3) it lowers the pK,-value of the alcohol to 7.5 so that an appreciable fraction is ionized at experimental pH-values. Acetylation of the alcohol by p-nitrophenyl acetate in the presence of a metal ion was also observed. This system is not, of course, a model for the carboxypeptidase reaction, but it does illustrate nicely the possible effects of a chelated metal ion.

PHYSICAL ORGANIC MODEL SYSTEMS: ENZYMATIC CATALYSIS

73

Carboxyl Group Catalysis More information is available on carboxyl as an intramolecular catalyst in ester and amide hydrolysis than on any other functional group. It was earIy found that aspirin [54] hydrolysed readily and gave a pH-rate constant profile signifying participation by the carboxyl group in the ionized form (Edwards, 1952; Garrett, 1957). A nucleophilic mechanism was originally preferred [equation (35)] but several pieces of evidence have cast doubt on this interpretation. The value of AS* is -24-7 e.u. (Garrett, 1957), considerably more negative than expected in a unimolecular reaction, and the reaction is 0

0

I I

(35)

slower in D 2 0 than HzO ( J Z H 2 0 / J Z D 2 0 = 2-2). A D 2 0 solvent isotope effect near unity is expected for a nucleophilic reaction (Butler and Gold, 1960). Substituent effects in the 4- and 5-positions of the benzene ring were interpreted to be in accord with a general base process, [55], as are the AS* value and Dz O solvent isotope effect (Fersht and Kirby, 1967). Furthermore, it was found, contrary

to the report of Bender et al. (1957), that 8 O is not incorporated into the salicylic acid product when the reaction is run in water

T. H. FIFE

74

enriched in " 0 (Fersht and Kirby, 1967). Lack of 180-incorporation is not consistent with a nucleophilic mechanism. General acidcatdysed attack of hydroxide ion is unlikely because of the magnitude of the calculated rate constants for such a reaction and because general acid catalysis of the attack of other oxygen nucleophiles is not observed. The 3- and 5-nitrosalicylates also hydrolyse with general base participation by the carboxyl group, even though the leaving group is a nitrophenol (Fersht and Kirby, 1968a). The carboxyl group does act as a nucleophile in hydrolysis of acetyl 3,5-dinitrosalicylate [ 561, as shown by incorporation of l 8 0 into the 3,5-dinitrosalicyIate product. However, the D 2 0 solvent isotope effect is 2-0 and AS* is -20-6 e.u. Fersht and Kirby reason that an intermediate anhydride is formed, the decomposition of which is rate-determining and catalysed by the adjacent phenoxide ion in a general base mechanism.

[561

products

A bell-shaped pH-rate constant profile is obtained in hydrolysis of 3-acetoxyphthalic acid [ 571 although bifunctional catalysis is not possible (Fersht and Kirby, 1968b). Kirby concludes that both carboxyl groups are involved in the reaction in a stepwise manner 0

PHYSICAL ORGANIC MODEL SYSTEMS: ENZYMATIC CATALYSIS

75

which he terms “series nucleophilic catalysis”. The second carboxyl group causes the hydrolytic rate constant of [57] to be 6.3 x l o 3 times larger at pH 3-8 than that for ionized acetylsalicylate. 3Hydroxyphthalic anhydride was ascertained to be an intermediate in the reaction [equation (37 ) ]. Carboxyl group participation has also been investigated in the hydrolysis of phthalate monoesters. Bender et al. (1958a) reported a pH-rate constant profile for hydrolysis of monomethyl phthalate showing maximum participation by the dissociated carboxyl group. An anhydride intermediate and nucleophilic attack by the carboxylate ion on the neutral ester were postulated. It was later reported that the undissociated carboxyl group is the catalytic species in the hydrolysis of monomethyl 2,3-di-t-butyl succinate, and monomethyl 3,6-dimethylphthalate, which give anhydride intermediates (Eberson, 1964). In contradiction of the work of Bender, Thanassi and Bruice (1966) found a pH-rate profile for hydrolysis of monomethyl phthalate showing maximum participation by the carboxyl in the undissociated form. The carboxylate ion is probably attacking the

eYH+

0

0

I

+

- a) OH

0

p c H 3

CH30H

+

o y y II

II

0

0

protonated species [equation (38)]. The corresponding phenyl ester, where the leaving group is of low basicity, did hydrolyse with participation by the ionized carboxyl group [equation (39)]. The leaving

-

a\: ; o II

0

+

(\- t o/ -

(39)

76

T. H. FIFE

group of N-acetyl serinamide or propargyl esters is of intermediate basicity (pK,-values of the alcohols are 13.6 and 13.5 respectively), and the hydrolysis is pH-independent over the entire range studied. Hence, the two mechanisms are energetically comparable with. esters having such leaving groups. The pH-rate constant profile for hydrolysis of phthalamic acid showed participation by the undissociated carboxyl group, and an anhydride intermediate was detected (Bender, 1957; Bender et al., 1958b). Hydrolysis was l o 5 times faster than in the case of p-carboxybenzamide. A four-center mechanism [ 581 was postulated.

II

0 ~581

The D2O solvent isotope effect for such a reaction should be large since proton transfer is occurring in the transition state, but the = 1-5).As suggested reaction is actually faster in D2 0 ( k , /KH by Bender, nucleophilic attack by the carboxylate ion on the protonated amide [equation (40)] may be the actual mechanism. Thus, the

II

0

II

0

I

0

mechanisms for the amide and the methyl ester are probably similar, requiring protonation of the poor leaving group before nucleophilic attack by the carboxylate anion can occur.

PHYSICAL ORGANIC MODEL SYSTEMS: ENZYMATIC CATALYSIS

77

While a carboxylate anion is a potent nucleophile in an intramolecular reaction it is not powerful enough to displace methoxide or an unprotonated amine, unless the nucleophile and the carbonyl group are held even more rigidly than in the phthalate system. Kirby and Lancaster (referred to by Kirby and Fersht, 1971) have found that such displacement can occur in disubstituted maleate monoesters and amides. The estimated rate constant for cyclization of N-methyl dimethylmaleamic acid [equation (41)] is 16,000 times greater than that for the unsubstituted compound. Below pH 5.6 hydrolysis of the cyclic anhydride becomes rate-determining. 0

0

II

I1

-

0

0

-

products

(41)

0 + CH3 NH2

An ionized carboxyl group can displace methoxide ion from methyl hydrogen diisopropyl maleate (Aldersley et al., 19 72). There is a plateau in the pH-rate constant profile from pH 8-14 which was ascribed to reaction catalysed by the carboxylate anion [equation (42)i. 0

i-prIco2 i-Pr

/

i-Pr

co2-CH3

i-Pr

I1 0

- 120 i-Pr

i-Pr

II

4

(42)

y o

+ CH30products

Nucleophilic carboxyl group participation is much more favourable with the phthalate, maleate, glutarate, and succinate monoesters than the acetylsalicylic acid derivatives. Monophenyl phthalate hydrolyses more than lo6 times as fast as aspirin. The carboxylate ion acts as a nucleophile with monophenyl phthalate but as a general base in the case of acetylsalicylic acid, even though the leaving group

78

T. H. FIFE

from both compounds is a phenol. This is probably because with monophenyl phthalate the leaving group breaks away from the rest of the molecule (one particle breaks down to two). A favourable entropy change results which would not occur if the carboxyl group of acetylsalicylic acid acted as a nucleophile, since the leaving group would then still be attached to the remainder of the molecule; much of the advantage of intramolecular nucleophilic attack would thus be 0

I/

0

lost. Furthermore, the neighbouring phenoxide ion of the intermediate compound would itself be a potent nucleophile and should attack the anhydride carbonyl group, leading to marked reversibility in the reaction [equation (43)]. As a consequence, general base catalysis becomes the most favourable mechanism. It can be concluded that, for maximum effectiveness of intramolecular nucleophilic attack, the leaving group must actually break away. Morawetz and Oreskes (1958) postulated that intramolecular general acid catalysis occurs in concert with nucleophilic attack by carboxylate anion in the hydrolysis of the monoanion of succinyl salicylate. At pH 4 the observed hydrolytic rate constant was found to be lo4 times greater than that for hydrolysis of acetylsalicylate. A 0

II

0

PHYSICAL ORGANIC MODEL SYSTEMS: ENZYMATIC CATALYSIS

79

bell-shaped pH-rate profile was in accord with bifunctional catalysis, but Maugh and Bruice (1971) ruled out concerted bifunctional catalysis after examining the para-carboxyl derivative. The profile is similar to that for 3-acetoxyphthalate and a similar mechanism involving series nucleophilic catalysis [equation (44)] is possible. A carboxylate anion is an efficient nucleophile in succinate and glutarate monoesters (Section 2). Gaetjens and Morawetz (1960) varied substituents in the phenolic leaving group. A linear plot of the logarithms of the rate constants us. (T is obtained, with a p-value of +2-2, close to that for ionization of phenols. Furthermore, differences in rate within each series are due to differences in AS'. These results were interpreted as indicating a concerted mechanism [59] with considerable bond breaking in the transition state. They are equally consistent with rapid formation of a tetrahedral intermediate whose breakdown is rate-limiting.

A carboxylate anion will function as an intramolecular nucleophile even when there is a large difference in pK, between the attacking species and the leaving group. Thus, with a p-methoxy-substituent in the phenyl succinate and glutarate series ApK, will be 5-6, and catalysis is still nucleophilic, in marked contrast to corresponding intermolecular reactions where a ApK, of 2- 3 is sufficient to change the mechanism to general base catalysis. The acetate-ion catalysed hydrolysis of p-nitrophenyl acetate proceeds in large part by a general base mechanism (Oakenfull et al., 1966; Gold et al., 1968). The ability of an intramolecular carboxylate anion to act as a nucleophile in spite of a poor leaving group is a reflection of the extremely large effective molarity of such a group. The effective molarity of the carboxyl group of p-methoxyphenyl 3,6-endoxo-A-4tetrahydrophthalate is 10' M in comparison with acetate-ion catalysed hydrolysis of p-methoxyphenyl acetate (Bmice and Turner, 1970).

80

T. H. FIFE

Mechanistic Conclusions It is apparent from the above discussion that much less can be said in a conclusive manner about the carboxypeptidase mechanism than that of a-chymotrypsin. While the general features of carboxypeptidase catalysis are understood (Kaiser and Kaiser, 1972), the critical experiments that would allow a quantitative explanation of the rates of carboxypeptidase-catalysed reactions have not been performed. What is needed specifically is knowledge of the effect of a metal ion on intramolecular carboxyl group participation in ester and amide hydrolysis. The important work of Buckingham and coworkers on hydrolysis of Co+++-complexed amino acid esters comes close, in that oxygen bases were found to catalyse hydrolysis. However, in those reactions the metal ion is built into the ground state of the molecule, and catalysis is not intramolecular. Consequently, the effects are striking but do not provide the necessary information. The metal-ion complexes of the salicyl ester of pyridine-2,6-dicarboxylic acid might be taken as models for carboxypeptidase action (Breslow and McAllister, 1971), but the rate increase associated with the salicyl carboxylate ion is too small to permit conclusions. Kaiser and Kaiser (1972) concluded that the most likely role for the carboxyl group of glutamic acid-245 is that of a nucleophile. In view of the presence of a metal ion (super acid) and possible general acid catalysis by a tyrosine hydroxyl group, it is probable that, in the active species, glutamic acid-245 is ionized. A nucleophilic mechanism is more probable than general base catalysis because, when they are sterically possible, intramolecular nucleophilic processes are much more effective. The few examples of intramolecular general base catalysed-reactions occur in systems where nucleophilic attack is hindered because of an unfavourable steric situation or because the leaving group does not break away from the rest of the molecule (“endocyclic displacement”). A factor of great importance in promoting a nucleophilic mechanism could be that a carboxylate ion is involved since the effective molarity of such a group in an intramolecular or intracomplex reaction ranges up to 10’ M. The original postulation of general acid catalysis by tyrosine-248 rested on closeness of that group to the amide function of the substrate in the crystal. However, the conformation in the crystal and solution states may be different in regard to the position of tyrosine248 (Riordan and Muszynska, 1974). Thus, mechanistic inferences concerning that group may not be warranted.

PHYSICAL ORGANIC MODEL SYSTEMS: ENZYMATIC CATALYSIS

81

5. GENERAL ACID CATALYSIS: LYSOZYME

General Acid Catalysis Lysozyme is a glycosidic enzyme of particular interest since the complete amino-acid sequence is known (Jollks, 1964; Canfield, 1963), as is the three-dimensional structure from x-ray crystalloresolution (Blake e t al., 1965; Johnson and graphic studies at 2 Phillips, 1965). Substrates for lysozyme include polysaccharides composed of alternate repeating units of N-acetylglucosamine (NAG) and muramic acid (NAM). Polysaccharides made up entirely of N-acetylglucosamine repeating units will also serve as substrates. It is likely that carboxyl groups from glutamic acid-35 and aspartic acid-52 are part of the active site. Several mechanisms have been suggested for lysozyme (Phillips, 1966; Lowe e t al., 1967; Raftery and Rand-Meir, 1968). All of these mechanisms involve glutamic acid-35 acting as a general acid, partially donating a proton to the leaving group in the transition state. A proposal [60] recently receiving support entaiIs general acid catalysis by glutamic acid-35 and electrostatic stabilization of a developing carbonium ion by the carboxylate anion of aspartic acid-52 (Raftery and Rand-Meir, 1968). A from the glycocarboxyl oxygen of glutamic acid-35 is located 3 sidic oxygen of the leaving group in the reaction, allowing proton

a

a

I I

CHzOH

H

NHCCH?.' .,

II

0

-c

I

transfer readily to occur. The pK,-value of this carboxyl group may be abnormally high so that a significant percentage would be in the un-ionized form at the pH-values at which the enzyme displays maximum velocity. Macroscopic pK,-values of 4-5 and 5.9 have recently

82

T. H. FIFE

been assigned to aspartic acid-52 and glutamic acid-35 respectively, ir 0 1 5 M KC1 and at 25" (Parsons and Raftery, 1972), although this work has been challenged (Banerjee et al., 1973). Kinetically determined values of pK,,, are 3.8 and 6-7 for k c a t and 4-2 and 6-1 for k c , , / K , (Banerjee et al., 1973) in hydrolysis of the P (1-+ 4) linked hexasaccharide of N-acetylglucosamine at 40" and p = 0.1. A role for aspartic acid-52 has been difficult to ascertain. It could participate in the anionic form as a nucleophile or by electrostatic stabilization of an incipient carbonium ion. A nucleophilic mechanism has been looked upon with disfavour because of the 3 i% distance between the carboxyl oxygen of aspartic acid-52 and C-1 of the hexose unit undergoing cleavage (Vernon, 1967), and no evidence for formation of a glucosyl-substituted enzyme has been found. The evidence that ratios of transfer products (with various alcohols) to hydrolysis products are constant suggests that a carbonium ion intermediate is produced, but is not conclusive (Rupley et al., 1968). Also, upon substitution of deuterium for hydrogen at C-1, adjacent to the bond undergoing cleavage, a k ~ / ratio k ~ of 1.11 was obtained, consistent with formation of a carbonium ion intermediate in lysozyme-catalysed hydrolysis of phenyl-P-NAG-glucoside at pH 3.1, 5.5, and 8.3. Acid-catalysed hydrolysis of phenyl 0-D-glucoside, which goes through a carbonium ion intermediate, has k H / k D = 1.13 while base catalysis gives k H j k , = 1-03 (Dahlquist et al., 1969). Nucleophilic participation by the neighbouring acetamido-group in the substrate has also been postulated, but substrates lacking this group have been developed which give enzymatic rates (Raftery and Rand-Meir, 1968). It is therefore not necessary to invoke such participation to explain the rates attained with the enzyme. The binding site of lysozyme will accommodate six hexose residues (A-F). The carboxyl groups of glutamic acid-35 and aspartic-52 are located between sites D and E. The lactyl groups at C-3 of the NAM residues cannot fit in sites A, C, and E. Therefore NAM residues must be located at sites B and D. Cleavage occurs at the reducing end of the NAM residue in site D (the bond being broken is between residues D and E). In addition to x-ray elucidation of the structure of the crystalline enzyme, the structure of a crystalline complex of lysozyme and tri(N-acetylglucosamine) was determined (Phillips, 1966). The trisaccharide occupied sites A, B, and C. Assuming that binding of a hexamer (adding hexose residues D, E, and F) would not change the conformation of the enzyme, the conformations of the substrate at

PHYSICAL ORGANIC MODEL SYSTEMS: ENZYMATIC CATALYSIS

83

sites D, E, and F were deduced. It was concluded that the ring occupying site D must be distorted to a half-chair because of interartions between the 6-hydroxymethyl group and residues 58 and 108 of the enzyme and the N-acetyl group of the hexose residue bound in site C. It must be emphasized that this conclusion was based on model building and pot direct experimental evidence. Distortion of the substrate in site D from the favoured chair conformation to a half-chair should greatly favour bond breaking at C-1 because a half-chair should resemble the transition state, assuming that a cyclic glycosyl carbonium ion intermediate is being formed. Chipman and Sharon (1969) have estimated that AG for binding in sites A, B, C, E, and F is negative (-5.5 kcalmole-’ in site C) but positive (‘3 to +6 kcal mole-’) in site D, supporting the conclusions based on model building. At the time the mechanism shown in [ 601 was proposed it was not known whether it was even chemically reasonable. Bimolecular general acid catalysis involving proton transfer in the transition state [equation (45)] had never been observed in the hydrolysis of glycosides or simple acetals. Bronsted and Wynne-Jones (1929) had

RCH /OR‘ ‘OR‘

+HA

-

H OR’ RCH::6+ ‘*OR’

(45)

I

found hydrolysis of the orthoesters, ethyl orthoacetate, ethyl orthopropionate, and ethyl orthocarbonate to be catalysed by general acids but had not detected such catalysis with acetals. In subsequent years a number of investigators made a diligent search for general acid catalysis of acetal hydrolysis without success. Capon (1963) and Capon and Smith (1965) had reported that a neighbouring carboxyl group in o-carboxyphenyl 0-D-glucoside [61] and o-methoxymethoxy benzoic acid [62] facilitated the rate of hydrolysis by factors of lo4 and 600 in comparison with the para-substituted

a4

T. H. FIFE

compounds, but the mechanism of the carboxyl group participation remained obscure because of kinetically indistinguishable possibilities.

Prior to 1967 acetal hydrolysis had been found to be a specificacid catalysed reaction with the accepted mechanism [ equation (46)] involving fast pre-equilibrium protonation of the acetal by hydronium ion, followed by unimolecular rate-determining decomposition of the protonated intermediate to an alcohol and a resonance stabilized carbonium ion (Cordes, 1967). An A-1 mechanism was supported by an extremely large body of evidence, but it appeared unlikely that such a mechanism could explain the

RCH,,OR' OR'

+ H30'

a

H RCH
+ H20

rapid rates obtained in the hydrolysis of glycosides catalysed by glycosidic enzymes functioning near neutrality. A free-energy reaction-coordinate diagram for the specific acidcatalysed A-1 hydrolysis of simple acetals might appear as shown in Figure 4(a) (Anderson and Fife, 1969). Decomposition of the protonated acetal is normally the rate-determining step. To obtain partially rate-determining protonation by hydronium ion [Figure 4(b)] and general acid catalysis, either basicity must be reduced, increasing the height of the peak for the protonation step, and/or the ease of C-0 bond breaking must be increased, decreasing the height of the peak for the bond breaking step. If the two peaks are of comparable height, the reaction might become concerted.

PHYSICAL ORGANIC MODEL SYSTEMS: ENZYMATIC CATALYSIS

85

Employing this as a guide, the author and his coworkers were successful in conclusively demonstrating general acid catalysis in the hydroIysis .of three different types of acetals and ketals, 2-@substituted phenoxy) tetrahydropyrans having electron withdrawing

~

Reaction coordinate Figure 4. Free energy-reaction co-ordinate diagrams for acid-catalysed hydrolysis of simple acetals

substituents in the leaving group [63] (Fife and Jao, 1968; Fife and Brod, 1970), tropone diethyl and ethylene ketals [64] (Anderson and Fife, 1969; Fife and Anderson, 1971b), and benzaldehyde di-tertbutyl acetals [18] (Anderson and Fife, 1971b). With all of these

1631

~ 4 1

compounds the rate constants for buffer acid catalysis are less in D 2 0 than H, 0, indicating that proton transfer occurs in the critical transition state, the reactions proceeding by classical general acid catalysis. This work has recently been reviewed in detail (Fife, 1972), so only a summary of the pertinent conclusions will be given here. Electron withdrawal by the para-substituent in the leaving group of the tetrahydropyran derivatives both lowers basicity and promotes ease of bond breaking. The positive value of the Hammett p-factor for formic acid catalysis (+0*9)in comparison with the negative value for hydronium ion catalysis (-0.9) shows that ease of bond breaking becomes more important as the catalysing acid becomes weaker, i.e., bond breaking has progressed further in the transition state when the catalyst is a weak acid. This conclusion has been supported by secondary deuterium isotope effects (Bull et al., 1971). The

86

T. H. FIFE

reaction is therefore best represented as a concerted process [ 6 5 ] , a representation further supported by the Bronsted coefficient of 0.5 for general acid catalysis in water. The proton should be closest to

the weakest base in the transition state. Consequently, to account for an a-value of 0-5 (implying that the proton is halfway between catalyst and substrate) it is necessary to postulate that basicity is increased in the critical transition state, i.e., the C-0 bond is breaking. That bond breaking is quite facile can be seen from the fact that spontaneous hydrolysis of the p-nitro-derivative is pHindependent at pH-values greater than 4. The D 2 0 solvent isotope effect ( k H /kD 0 = 0.9) and the AS* (-2.2 e.u.) indicate that the pH-independent reaction is a unimolecular decomposition [ 661 made possible by the good leaving group and the moderately stable

carbonium ion intermediate produced in the reaction (Fife and Brod, 1970). Anderson and Capon (1969) similarly showed that the hydrolysis of benzaldehyde aryl methyl acetals is subject to general acid catalysis. The same structural features are undoubtedly important as in the case of phenoxytetrahydropyrans, and again, p is positive in the acetic acid-catalysed reaction. A phenolic leaving group is not by itself sufficient for general acid catalysis to occur if the intermediate carbonium ion is not of sufficiently great stability. Thus, general acid catalysis is not observed in the hydrolysis of phenolic glycosides (Piszkiewicz and Bruice, 1967). A glycosyl carbonium ion is quite unstable owing to inductive electron withdrawal by the ring hydroxyl groups. The tetrahydropyranyl carbonium ion, however, is of moderate stability.

PHYSICAL ORGANIC MODEL SYSTEMS: ENZYMATIC CATALYSIS

87

It is primarily this difference in carbonium ion stability that makes 2-methoxytetrahydropyran 3 x 1O7 times more reactive than methyl a-D-ghcopyranoside (Dyer et al., 1962). A phenolic leaving group and moderate carbonium ion stability must be present together for general acid catalysis to occur; the hydrolysis of 2-ethoxytetrahydropyran is specific acid-catalysed (Fife and Jao, 1968). Low basicity, even in combination with a reasonably stable carbonium ion intermediate, will not permit general acid catalysis if the leaving group is not sufficiently good. Thus, hydrolysis of benzaldehyde methyl S-phenyl thioacetaIs [67] is not subject to general acid catalysis (Fife and Anderson, 1970) even though

hydrolysis of the exactly analogous oxygen acetals is general acidcatalysed. It was shown conclusively that thiophenol is the leaving group in the hydrolysis of the thioacetals. It was important to find general acid catalysis in the hydrolysis of acetals having poor (aliphatic) leaving groups because the natural substrates for lysozyme are of that type. Tropone diethyl ketal 164-1 has a poor leaving group and its basicity is presumably normal, but ease of bond breaking has been greatly enhanced by the extraordinary stability of the intermediate carbonium ion. That such high carbonium ion stability is necessary for general acid catalysis to occur with poor leaving groups was shown by the study of a series of acetals having similar (aliphatic alcohol) leaving groups and carbonium ion intermediates of progressively increasing stability (Fife and Anderson, 1971b). In the series benzophenone diethyl ketal, p , p’-dimethoxybenzophenone ethylene ketal, 2,3-diphenylcyclopropenone diethyl ketal, ferrocene diethyl acetal, and tropone diethyl and ethylene ketal, general acid catalysis was observed only with the tropone ketals. I t is concluded that bimolecular general acid catalysis will be observed in cases where the leaving group is an aliphatic alcohol only when carbonium ion stability approaches that of an alkoxytropylium ion. Benzaldehyde di-t-butyl acetal [ 181 has an aliphatic alcohol leaving group, but the bond breaking process is aided by relief of ground

88

T. H. FIFE

state strain [ 681. The general acid catalysis observed with t-butyl acetals (a = 0.6) (and not detectable with the analogous diethyl acetals) is attributable primarily to greater ease of bond breaking.

General acid catalysis by H,POi and imidazolium cation has recently been detected in hydrolysis of 2-methoxy-3,3-dimethyloxetane [69] (Atkinson and Bruice 1974). Such catalysis must be due t o relief of ground state strain as the ring breaks to give an open-

chain carbonium-ion intermediate [ 701. Initial bond cleavage was accelerated by a factor of at least 5 x l o 5 compared to hydrolysis of simple acyclic acetals by an A-1 mechanism. Relief of ground state

strain can be sufficient, therefore, to allow general acid catalysis when the leaving group is poor and the carbonium ion is not highly stabilized. It should be noted that the open-chain carbonium ion

PHYSICAL ORGANIC MODEL SYSTEMS: ENZYMATIC CATALYSIS

89

would be less stable than the alkoxybenzyl ion from [ 181 but should be more stable than a glucosyl carbonium ion. Giudici and Bruice (1971) concluded that planarity of an acetal would not by itself permit general acid catalysis by buffer acids since none was observed in the hydrolysis of [71]. Consequently, relief of

strain must occur for a change in mechanism when the leaving group is an aliphatic alcohol and the carbonium ion intermediate only moderately stable. From this work the structural features in an acetal or ketal that will give rise to general acid catalysis are clear. General acid catalysis will be seen in hydrolysis of an acetal when the leaving group is good (a phenol) in cases where the intermediate carbonium ion is of moderate stability. General acid catalysis will also occur with acetals having a poor leaving group (an aliphatic alcohol) when the intermediate carbonium ion is of great stability, approaching that of an alkoxytropylium ion, or when relief of steric strain facilitates bond breaking. In all cases ease of bond breaking is the key feature with basicity considerations being of secondary importance. Jencks (1973) proposed that concerted general acid catalysis will be detected only when pK,-values for the ground state and the product are widely separated, with the basicity of the transition state being intermediate. Proton transfer then becomes concerted with bond breaking to avoid a high-energy protonated intermediate. Acetal hydrolysis appears to fit this criterion nicely. A reasonable estimate of the pK,-value of 2-(p-nitrophenoxy)tetrahydropyran would be -10 while that of the product, p-nitrophenol, is 7. The Jencks postulate appears to be generally useful, but nevertheless it is conceptually misleading when applied to acetal hydrolysis because it stresses low basicity as a principal cause of general acid catalysis, whereas ease of bond breaking is the key factor. Spontaneous

90

T. H. FIFE

hydrolysis of acetals subject to general acid catalysis, such as 2-@-nitrophenoxy)tetrahydropyran, is a pH-independent unimolecular decomposition over much of the pH-range. General acid catalysis can then be looked upon as a genuine facilitation; it does not arise simply to avoid a protonated intermediate. While this viewpoint and the Jencks postulate are formally equivalent, more insight is provided into the mechanism of acetal hydrolysis by laying proper stress on bond breaking rather than basicity considerations. After reflecting upon the structural features in an acetal necessary for general acid catalysed hydrolysis, it might be concluded that the lysozyme reaction would not proceed by that mechanism since the natural substrates have poor leaving groups and an intermediate carbonium ion in the reaction would be quite unstable. However, relief of ground state strain in the transition state will promote bond breaking to an extent sufficient for general acid catalysis to occur; and an integral feature of postulated mechanisms for lysozyme is distortion during binding of ring D to a half-chair. Estimates obtained from unstrained lactone substrates show that relief of ground state strain could increase the rate of lysozyme reactions by 103-104 (Secemski and Lienhard, 1971; see also the calculation on p. 104). Also, a simple consideration of structure does not take into account the large effective molarity of an intramolecular catalyst which might be sufficient to bring that type of mechanism into operation.

In tram olecular Catalysis Intramolecular carboxyl group catalysis has been found in a number of salicyl acetals. Capon et al. (1969) preferred intramolecular general acid catalysis as the mechanism of carboxyl group participation in the hydrolysis of o-carboxyphenyl P-D-glucoside and o-methoxymethoxybenzoic acid although the kinetically equivalent possibilities were not convincingly eliminated in the case of the glucoside. Capon e t al. considered three possible mechanisms [equation (47)] for hydrolysis of o-methoxymethoxybenzoic acid. Mechanisms (a) and (b) involve nucleophilic attack by the carboxylate anion on the protonated acetal, and mechanism (c) proceeds with intramolecular general acid catalysis. Capon was able to rule out mechanisms (a) and (b) since the synthetically prepared intermediate in (a) is stable under the reaction conditions. That in (b) leads to products at

PHYSICAL ORGANIC MODEL SYSTEMS: ENZYMATIC CATALYSIS

91

a faster rate than the overall reaction, but there is no UV absorbance due to the intermediate at the isosbestic point between reactant and H

+ CH30H

II

0

+ CH2 0

I1

0

product, whereas such absorbance would have been expected if the intermediate in (b) were indeed being produced. Capon prefers mechanism (c)-intramolecular general acid catalysis. The usual A-1 mechanism was not considered because of the observed facilitation of the rate of hydrolysis by the ortho-carboxyl group. Dunn and Bruice (1970, 1971) presented evidence that the mechanism of hydrolysis of methoxymethoxybenzoic acid is actually A-1. Evidence included a Bronsted coefficient for intramolecular catalysis of 1.0, based upon points for the unsubstituted compound and the 5-N02 derivative, and p*-values that were the same (-3.0) for both the methyl ester and the carboxyl derivatives of [ 721 when

92

T. H. FIFE

R was varied. Since the mechanism for the methyl ester is certainly A-1 and since intramolecular general acid catalysis should give a different transition state structure and therefore a different p* value, it was concluded that the mechanism was A-1 in both cases. The rate enhancement provided by the carboxyl group substituent was ascribed to electrostatic catalysis whereby a proton is stabilized on the acetal oxygen, thus lowering the dissociation constant of the conjugate acid. Complete protonation of methoxymethoxybenzoic acids might be required because of the unstable carbonium ion intermediate. An A-1 mechanism for hydrolysis of methoxymethoxybenzoic acids and electrostatic stabilization are reasonable postulates to explain the fairly small rate enhancements (-lo3) seen in comparison with reference compounds. Such a mechanism is similar to the electrostatic mechanism invoked to explain the modest increase in rate (-100 times) of hydrolysis of salicyl phosphate in comparison with p-carboxyphenyl phosphate (Bender and Lawlor, 1963), but it is debatable whether electrostatic stabilization of a proton on oxygen could account for the large rate enhancement factors of 10' O - 10' seen in enzymatic reactions. There does not appear to be convincing evidence in the literature that a neighbounng carboxyl group has an extraordinary effect on the pK,-values of neutral bases. For example, the pK,-values of 0-and p-aminobenzoic acids are about the same. A hydrophobic environment for the active site might make electrostatic catalysis quite effective, but pK, for aspartic acid-52 in lysozyme is normal and the pK,-value of glutamic acid-35 is close to 6.0, corresponding roughly to that of acetic acid in 50% dioxane-water, a solvent that is still quite polar. The active site of lysozyme must then be in a reasonably aqueous environment despite previous suggestions to the contrary. The effect of carboxyl group substitution is not greatly different for acetal hydrolysis in water and 50%-dioxanewater (Fife and Anderson, 1971a; Anderson and Fife, 1973). Craze and Kirby (19 74) have recently proposed that hydrolysis of methoxymethoxybenzoic acids proceeds with intramolecular general acid catalysis with little proton transfer from the carboxyl group and alarge amount of bond breaking in the transition state. This suggestion was, however, based solely on substituent effects and a graphical analysis using the extended 4-parameter Hammett Equation (see Section 2). In view of the problem of kinetic equivalence it was important to study carboxyl-substituted acetals where bimolecular general acid

PHYSICAL ORGANIC MODEL SYSTEMS: ENZYMATIC CATALYSIS

93

catalysis had been observed in hydrolysis of the unsubstituted compounds. It is a reasonable assumption that the mechanism of the intramolecular reaction will be the same as that of the corresponding

II

[ 731

0

[ 741

intermolecular reaction, i.e., general acid catalysis. Fife and Anderson (1971a) therefore examined hydrolysis of 2-(o-carboxyphenoxy) tetrahydropyran [ 731 and benzaldehyde methyl o-carboxyphenyl acetal [74]. 40

0

PH Figure 5. Plot of kobsd for hydrolysis of 2-(o-carboxyphenoxy)-te~ahy&opyranin 50% dioxane-HzO at 15" us. pH.

The pH-rate constant profile shown in Figure 5 was obtained for hydrolysis of (73) in 50% dioxane-water at 15'. A large plateau in the profile is to be noted. As in the other cases of intramolecular catalysis, kinetically equivalent possibilities exist, and the curve in Figure 5 can be calculated from either equation (48) or equation (49) with appropriate values of the rate constants, where kl is the

94

T. H. FIFE

second-order rate constant for hydronium ion-catalysed hydrolysis of the unionized acetal, k2 is the second-order rate constant for hydronium ion-catalysed hydrolysis of the ionized species, k, is the rate constant for intramolecular general acid catalysis, and K, is the dissociation constant of the carboxyl group. The neighbouring carboxyl group greatly accelerates hydrolysis. The value of k2 is 6-1x 10' times greater than kH+ for the corresponding ethyl ester and 3-8 x lo' greater than k H + for the unsubstituted compound. These are maximum differences in k,bsd for hydrolysis of the compounds at any pH (k2/kH+). The k2-value for the p-carboxylsubstituted compound could not be determined since its rates of hydrolysis are much too slow to measure accurately at the required pH at 15". However, from the measured value of K , and a reasonable assumption as to the difference between k, and k,, based on the known p-value for hydronium ion-catalysed hydrolysis of 2-(4substituted phen0xy)tetrahydropyrans and the a-values for C 0 2 H and CO;, it was calculated that K 2 @ m u ) must be only 1.6 x times the magnitude of k 2 (ortho). These differences in the rate constants are many times larger than expected on the basis of inductive effects. The usual method for establishing partidly rate-limiting proton transfer, determination of the rate constants in D 2 0 , would give ambiguous results (Bruice and Piszkiewicz, 1967). However, intramolecular general acid catalysis [equation (48)] is the preferred mechanism in view of the intermolecular buffer acid catalysis observed with the unsubstituted compounds. General acid catalysis [ 751 should therefore be favoured in the intramolecular reaction.

The ratio of rate constants for intramolecular general acid catalysis of the hydrolysis of [ 731 and intermolecular formic acid-catalysed hydrolysis of 2-phenoxytetrahydropyran is 580 M , a minimum value since the intramolecular reaction was studied at 15" while the bimolecular rates were measured at 50". The ratio would be much larger if comparisons could be made at the same temperature

PHYSICAL ORGANIC MODEL SYSTEMS: ENZYMATIC CATALYSIS

95

(-10,000 M ). The calculated ratio represents the concentration of formic acid in the bimolecular reaction required to give a pseudofirst-order rate constant comparable to that obtained in the intramolecular reaction. It is evident that intramolecular catalysis is greatly favoured, and that an increase in the local concentration of the carboxylic acid catalyst cannot explain the great efficiency of the reaction relative to intermolecular catalysis. The ratio of 580 M may reflect the nearly optimal geometry of the system in [73] for proton transfer between the two oxygen atoms. The magnitude of the rate constant k 2 for hydrolysis of benzs - l ) rules aldehyde methyl o-carboxyphenyl acetal (1.2 x l o 7 M out nucleophilic attack of a carboxylate anion on the protonated substrate or any mechanism involving a completely transferred proton, if the basicity of the acetal is normal. It can be calculated that L l , the rate constant for transfer of a proton from the conjugate acid of the substrate to water, would then necessarily have to be greater than that for a diffusion-controlled process. A decrease of l o 5 in the dissociation constant of the conjugate acid would be required to account for the rapid rate of hydrolysis since k, for (74) is 1.9 x lo5 times greater than kH+ for the corresponding methyl ester. As in the case of [73], the most likely mechanism is intramolecular general acid catalysis. These acetals then provide relatively unambiguous examples of intramolecular general acid catalysis. In hydrolysis of the analogous thioacetal, benzaldehyde methy1 S-(0-carboxyphenyl thioacetal [ 7 61 , carboxyl group participation cannot be observed. There is a difference of only 3 0 between k and

-'

II

0

k , , supporting the contention that the same structural features will permit both intramolecular catalysis and buffer acid catalysis. In view of the low basicity of sulphur, intramolecular general acid catalysis would be a favourable process if low basicity were of predominant importance in permitting general acid catalysis. A most important

96

T. H. FIFE

feature of these reactions must be that the salicyl anion is a better leaving group than the thiosalicyl anion. Release of salicylic acid from benzaldehyde disalicyl acetal gives a bell-shaped pH-rate constant profile (Figure 6 ) analogous to lysozyme-catalysed reactions (Anderson and Fife, 1973). The maximum

Figure 6. Plot of kobsd us. pH for hydrolysis of benzaldehyde disalicyl acetal in 50% dioxane-H20 at 25".

enhancement in kobsd in comparison with the dimethyl ester is a factor of 3 x l o 9 , is a rate enhancement of the magnitude seen in enzymatic reactions. It is probable that intramolecular general acid catalysis [ 771 is occurring. General acid catalysis by buffer acids is observed in hydrolysis of the dimethyl ester, and, as in the case of

II

0

PHYSICAL ORGANIC MODEL SYSTEMS: ENZYMATIC CATALYSIS

97

[ 731 and [ 741 , it is a reasonable supposition that the intramolecular reaction proceeds by the same mechanism. Likewise, if the reaction were A-1,transfer of the proton from the conjugate acid to water would necessarily have a rate constant greater than that of a diffusion-controlled reaction. The monoanionic species is most reactive, but its associated rate constant for intramolecular general acid catalysis is only 65 times greater than that for the unionized species. Most of the large rate enhancement in comparison with the dimethyl ester is due to participation by one carboxyl group, as is the case with the unionized acetal [ 771. If the carboxylate anion of the monoanionic species is electrostatically stabilizing the incipient carbonium ion in the reaction [ 81 , its effect on the rate must be small. Part of the rate difference between the monoanion and the unionized species is due to a changing inductive effect caused by ionization of the carboxyl group. The magnitude of the inductive effect was determined by studying the p-nitrobenzaldehyde acetal with one of the carboxyl groups in the para-position. The pH-rate constant profile for p-nitrobenzaldehyde disalicyl acetal is bellshaped with a rate factor of 260 between the monoanion and unionized species. When one carboxyl group is in the para-position the profile is still bell-shaped, but the magnitude of the bell is greatly reduced; there is a factor of only 5 between the rate constants and this must result from the changing inductive effect. The rate constants for the two unionized compounds are about the same. Thus, a factor of -50 in the rate constant for the monoanion of p-nitrobenzaldehyde disalicyl acetal cannot be accounted for by an inductive effect. This could represent the effect of carbonium ion stabilization by the carboxylate anion, but an effect of this size could also be steric. It is known that ortho-substituted phenyl glycosides and phenyl methyl formals hydrolyse slightly faster than the corresponding compounds substituted in the para-position (Nath and Rydon, 1954;Dunn and Bruice, 1970). The leaving group of the salicyl acetals, where intramolecular carboxyl group participation has been observed, is quite good. A question arises, therefore, concerning the ability of an intramolecular carboxyl group to catalyse hydrolysis of an acetal with a poor (aliphatic alcohol) leaving group. Intermolecular buffer acid catalysis is not normally observed with such compounds (with the exceptions noted above). In view of the large effective molarity for an intramolecular general acid, however, it is theoretically possible that intra-

98

T. H. FIFE

molecular general acid catalysis could be taking place. Benzaldehyde diglycolic acid acetal [78] gives a pH-rate constant profile for hydrolysis which is quite similar in shape to that found for the monocarboxyl substituted compound o-carboxyphenoxy tetrahydropyran (Fife and Anderson, unpublished data). The data do not give a good ,OCH2COOH m c H \ 0 C H 2

COOH

fit to the rate equation for the hydrolysis of an acetal with two carboxyl groups unless it is assumed that the difference in rate constants for the monoanion and the unionized acetal is small. Thus, a carboxylate anion is not facilitating the reaction. The maximum factor between k o b s d for [78] and the dimethyl ester is 390, which is larger than expected on the basis of inductive or steric effects. It is probable that a carboxyl group is participating, perhaps as a general acid [ 791. While the observed rate enhancement is small, it is still significant in showing that intramolecular catalysis can take place when the leaving group is an aliphatic alcohol. I t must be noted that pKa for the leaving group of [ 781 is probably less than that for ali-

phatic alcohols such as ethanol because of the inductive electronwithdrawing ability of a carboxyl group. This could explain why only weak carboxyl-group catalysis is observed. For large rate enhancements to occur when the leaving group is poor, the carboxyl-group catalyst and the reaction centre will presumably have to be held much more rigidly, in the proper steric relationship, than is possible with [ 781. A model in which the leaving group is an aliphatic alcohol and where a carboxyl group is held rigidly in position for participation is

PHYSICAL ORGANIC MODEL SYSTEMS: ENZYMATIC CATALYSIS

99

benzaldehyde methyl cis-2-carboxycyclohexyl acetal [ 801 , but carboxyl group participation is not observed (Fife and Anderson,

[go1

unpublished data). Compound [ 801 hydrolyses only 10 times faster than the methyl ester. The data in Table 7 reveal that leaving-group ability and not carbonium ion stability is the key factor permitting intramolecular TABLE 7 Hydrolysis of Carboxyl Substituted Acetals and Ketals Compd.

1.

Ref.

Rate C02H enhancement participation factof

C+

Leaving group

/OEt (CH\oEt

b

no

Stable

Poor

no

Stable

Poor

Stable

Poor

Stable

Poor

jC02H

2.

no

3.

HoocKx: C

no

4. Yes

OH

lo4

Unstable Good

100

T. H. FIFE

TABLE 7-continued Ref.

Rate COzH enhancement participation factora

d

yes

600

Unstable

e

yes

105-106

Moderately Good stable

e

yes

lo5

Stable

f

Yes

3~ 109

Moderately Good stable

~cH\OCH2C02H g

390 Yes

Moderately Interstable mediate

Compd.

9.

C+

Leaving POUP

Good

Good

/OCHzCO2H

10. Stable

a Rate enhancements are in comparison with the para carboxyl substituted or

unsubstituted compounds or corresponding methyl esters. Fife, unpublished work. Bruice and Piszkiewicz, 1967. Capon et a l , 1969. Fife and Anderson, 1971a f Anderson and Fife, 1971a, 1973. Fife and Anderson, unpublished work.

Poor

PHYSICAL ORGANIC MODEL SYSTEMS: ENZYMATIC CATALYSIS

10 1

carboxyl group participation. Thus, when the leaving group is phenolic (salicyl), participation occurs when the carbonium ion intermediate is unstable (compounds 4 and 5), moderately stable (compound 6), or stable (compound 7). On the other hand, when the leaving group is poor (an aliphatic alcohol), intramolecular participation is not observed, even when the intermediate carbonium ion is reasonably well stabilized (compounds 1, 2, 3, and 10). As discussed above, the weak carboxyl group participation in hydrolysis of benzaldehyde diglycolic acid (compound 9) can be explained by a reduced pK,-value of the leaving group in comparison with ethanol and cyclohexanol. With acetals having a normal aliphatic alcohol of high pK, as leaving group, bond breaking must be so difficult that intramolecular general acid catalysis or a kinetic equivalent does not take place even in cases where the carbonium ion is an alkoxybenzyl ion; specific hydronium ion catalysis is then the only observable mechanism. If the intermediate carbonium ion could be made exceedingly stable, as for example with the tropone ketal [ 8 1 ] , bond breaking might be sufficiently easy for carboxyl group participation even with an aliphatic alcohol leaving group, but this possibility is at present untested.

~-O-@-D -Glucopyranosyl-L -gulonic acid [ 821 hydrolyses more than 100 times faster than 4 - 0 - p --glucopyranosyl-D ~ -glucitol [ 831 at pH 4 and exhibits a plateau in its pH-rate profile (Smidsrod et al., 1966), although a complete profile was not obtained. Intramolecular catalysis [84] was proposed and likewise for hydrolysis of alginate

102

T. H. FIFE

which contains uronic acid residues, but more evidence is necessary before these can be considered legitimate examples of general acid catalysis. In particular, it is necessary to assess the changing inductive effect.

--....op “o Note should be made of the fact that all of the mecha&rxls for lysozyme in which a carbonium ion intermediate is formed are depicted as proceeding through a cyclic glycosyl 0x0-carbonium ion. In the acid-catalysed hydrolysis of glucosides the carbonium ion formed is probably cyclic rather than an open-chain ion (Capon, 1969), even though the cyclic ion would necessarily have to adopt a

+

R‘OH

(50)

HO

V30+

G

HO

\

+

H=OR OH

I

products

/

strained half-chair conformation [ 851 (strain energy amounting to 8-10 kcal mole-’ ; Eliel, 1962). Formation of a cyclic carbonium ion would be preferred because: (1) AS* should be favourable for a

PHYSICAL ORGANIC MODEL SYSTEMS: ENZYMATIC CATALYSIS

103

process in which one particle decomposes into two; and (2) reversal of the formation of the ion would be more favourable for an

open-chain carbonium ion because of intramolecular reaction with a neighbouring hydroxyl group [equation (SO)] . Hydrolysis of benzaldehyde- 1,3-dioxolanes, where re-closure of the ring can similarly take place after carbonium ion formation, proceeds about 35 times more slowly than that of analogous benzaldehyde diethyl acetals (Fife and Jao, 1965; Fife and Hagopian, 1966). The difference in rate is due entirely to an 8-10 e.u. AS * . p-Methoxyphenyl-4,4,5,5-tetramethyl1,3difference in dioxolane [86] hydrolyses 40,000 times more slowly than the corresponding diethyl acetal (Fife, 1967; Fife and Brod, 1968). There is evidence for a change in mechanism to A-2 with the tetramethyl- 1,Sdioxolanes. It is likely that the great facility of ring closure due to the geminal methyl substitution makes an A-1 mechanism unfavourable and thereby leads to a normally less favoured mechanism not involving a discrete carbonium ion intermediate. Geminal substitution would restrict rotation of the

hydroxyl group away from the carbonium centre and might lead to faster reaction with the neighbouring hydroxyl than with a water molecule. Thus, the factors of AS’ and reversibility have a profound influence on rate and mechanism in hydrolysis of cyclic acetals. It is reasonable that a cyclic glycosyl carbonium ion would be formed in lysozyme-catalysed reactions, but this expectation must be viewed in the light of the postulated straining of the substrate during

104

T. H. FIFE

binding. Pronounced strain in the hexose ring would prevent reclosure of the ring if an open-chain carbonium ion were formed. Thus, a major factor in cyclic carbonium-ion formation would be removed. An open-chain carbonium ion would not only be more stable than a cyclic ion, but cleavage of the ring would release all strain energy in the ring. The bond breaking process might then be quite facile in spite of the poor leaving group so that general acid catalysis by glutamic acid-35 could occur. The question that must be asked is whether reversibility or AS * considerations are more important in regard to the position of bond cleavage. If the AS* difference of 8-10 e.u. for benzaldehyde diethyl and ethylene glycol acetals (Fife and Jao, 1965) accurately represents the change expected for cleavage of an acetal ring as compared to an acyclic bond, then reversibility must be the predominant influence. D Glyceraldehyde dimethyl acetal hydrolyses at 25" with a secondM-' s-' (Capon order rate constant for acid catalysis of 1 - 7 7 x and Thacker, 1965), whereas methyl p-D -glucopyranoside has a M s-' at 60" (data corresponding rate constant of 3.5 x from Capon, 1969, Table 9). If one makes a reasonable allowance for the temperature difference, the glycoside hydrolyses approximately lo4 times more slowly. The rate difference must be primarily caused by strain in the cyclic carbonium-ion intermediate of the glucoside. Transglycosidation with retention of configuration (Chipman and Sharon, 1969) would be more difficult to explain if an open-chain carbonium ion were formed in lysozyme reactions, necessitating an equilibrium reaction with free aldehyde. It seems unlikely, however, that a cyclic carbonium ion intermediate could have a sufficiently long lifetime to react with a saccharide molecule that can bind to the enzyme only after displacement of the leaving group in a fairly aqueous environment (above discussion). Therefore, the concept of a cyclic carbonium ion also presents difficulties for interpretation and should not be accepted uncritically. Reaction of the carbonium ion with water could be reduced if overlap occurred with the carboxylate anion of aspartic acid-52 either during or after the glycoside-cleavage step. Since the carboxylate anion would be held adjacent to the carbonium ion in the active site, equilibrium should be far to the side of the acylal. Reaction of acylal with H,O would then very probably be ratedetermining in the forward direction. Evidence has been obtained that the solvent is directly involved in the hydrolysis of the cyclic acylal 2-(p-nitrophenoxy)phthalide where steric factors are similar

-'

PHYSICAL ORGANIC MODEL SYSTEMS: ENZYMATIC CATALYSIS

\O-

\O-

NHCOCH3

105

NHCOCH3 (52)

(Fife and De, 1974). Hence, the above scheme [equation (52)] would lead to build-up of a detectable concentration of acylal intermediate which is contrary to fact.

Nucleophilic Catalysis Piszkiewicz and Bruice (1967) found both hydronium ion and hydroxide ion catalysis in the hydrolysis of p-nitrophenyl 2-acetamido-6-D -glucoside and also, pH-independent hydrolysis from pH 2-1 1. Such a reaction is not seen in the hydrolysis of the corresponding a-derivative, and at neutrality the p-isomer hydrolyses 1O3 times more rapidly. There is also a pH-independent hydrolysis (pH 3-7) with p-nitrophenyl p-D -glucoside, but the reaction is 300 times more favourable with the P-acetamido-derivative. The large rate difference indicated that the acetamido-group was participating in the reaction. A nucleophilic mechanism [87] involving attack of the neutral amide at C-1 of the glycoside was favoured since a kinetically equivalent mechanism in which the negatively charged amide oxygen attacked the protonated glycoside would require unreasonably large rate constants. The p-value for nucleophilic catalysis is 2.6. Thus, catalysis becomes less effective as the leaving group becomes worse.

106

T. H. FIFE

Acid-catalysed hydrolysis of corresponding a- and P-methyl 2acetamido-D -glucosides was studied with similar results (Piszkiewicz and Bruice, 1968a). The a-derivative has a second-order rate constant ( k H + ) one-fourth of that of the fi-anomer. When log k,+ values for O-D -glycosides of N-acetylghcosamine are plotted us. log k H + for the corresponding p-D-glucopyranosides, a linear relationship exists. The point corresponding to the methyl glycosides deviates from this plot and can be accounted for as a 50-fold enhancement in the rate of hydrolysis of methyl 2-acetamido-fi-D-glucoside. Mechanism [ 881 was proposed in which the amide oxygen attacks the protonated gly coside.

CH3 [881

A neighbouring carboxyl group in o-carboxyphenyl-2-acetamidogives further rate enhancement, hydrolysing 8.1 x lo4 times faster than expected from a Hammett up plot

p-D-glucopyranoside [ 891

(89)

of log k H + for a series of substituted phenyl 2-deoxy-2-acetamidoglycopyranosides us. u (Piszkiewicz and Bruice, 1968b). Concerted intramolecular nucleophilic attack by the acetamido-function and general acid catalysis by the ortho-carboxyl group was suggested. oCarboxyphenyl 2-deoxy-2-acetamidoglucopyranoside hydrolyses only 7 times faster than o-carboxyphenyl 0-D-glucoside so that catalytic effects of the two groups are not additive. Incorporation of a second functional group within the molecule should lower M* but would entail additional restrictions in the transition state. Orientation of the second group would make TAS" more unfavourable, assuming that M* is a measure of potential

PHYSICAL ORGANIC MODEL SYSTEMS: ENZYMATIC CATALYSIS

107

energy changes and AS* measures kinetic energy changes. The observed values of the activation parameters (kcal mole-' ) for o-carboxyphenyl 0-D-glucopyranoside (AH*,28.4; TAS*, leg), ocarboxyphenyl 2-acetamido-2-deoxy-P-~ -glucopyranoside (M, 24.6; TAS*, 0-0), and o-nitrophenyl-2-acetamido-2-deoxy-~-~-glucopyranoside (AH*, 27.0; TAS*, 0-2)are in accord. The decrease in TAS' necessary to orient the second catalytic group is about one-half of that predicted if the second group were acting as a bimolecular catalyst. These studies with acetamidoglycosides provide support for the plausibility of mechanisms for lysozyme in which nucleophilic attack by the neighbouring group takes place. The synthetic substrates developed for lysozyme which lack this group (Raftery and RandMeir, 1968) render unnecessary any postulate of involvement of the 2-acetamido-group, but, of course, do not rule out the possibility. Perhaps the greatest importance of the work of Bruice and Piszkiewicz is in showing that intramolecular nucleophilic attack HOCH2

I

-

HOCH2

HOCH2 I H

I

H

H+

OH

+

OH

CH30H

CH3S 'CH2' CH

/OEt 'OEt

H30'

, -

EtOH + CH3S.CH2-C<

CHjS.CH2.CH

I

H ,OEt +

'OEt

0

H20

+ EtOH

+ -

H

+

CH2--CH--OEt

(54)

108

T. H. FIFE

by a suitable nucleophile will take place in the hydrolysis of glycosides. Capon and Anderson have also found an example of intramoIecular nucleophilic attack with o-carboxybenzaldehyde diethyl acetal in 82% dioxane-water (referred to in Capon, 1969). The carboxyl group does not participate in the reaction at dioxane concentrations less than 82%. In 82% dioxane, pK, for the neighbouring carboxyl group would be approximately 11-12 so that it would be a potent nucleophile under such conditions. Nucleophilic participation has also been suggested in the reactions (53) and (54) on the basis of small rate enhancements in comparison with reference compounds (Capon and Thacker, 1965; Speck, et al., 1965).

Acylal Hydrolysis

If aspartic acid-52 acts as a nucleophile in lysozyme reactions a glycosyl enzyme intermediate will be formed [60]. There is no evidence, kinetic or otherwise, for substituted enzyme intermediates, but rapid breakdown might preclude attainment of detectable concentrations. Formation of a substituted enzyme could explain the observed retention of configuration at the anomeric carbon in transglycosidation reactions, provided backside attack in a subsequent reaction is chemically reasonable. It has therefore been important to attempt t o understand the chemistry of acylal hydrolysis so as to assess the properties that would be expected of an acylal intermediate in reactions catalysed by the enzyme. y-Ethoxy-y-butyrolactone undergoes both hydroxide- and hydronium ion-catalysed hydrolysis but, in addition, there is a pHindependent rate (Figure 7) from pH 4-9 (Fife, 1965). The hydronium ion-catalysed reaction was considered to proceed by an A-1 mechanism. The AS* of -6.9 e.u. is slightly more negative than is usual for A-1 reactions but is of the magnitude obtained in other A-1 ring cleavage reactions (Fife and Jao, 1965; Fife and Hagopian, 1966). Acylals have generally been observed to undergo acidcatalysed hydrolysis by an A-1 mechanism (Salomaa, 1957; Salomaa and Laiho, 1963; Salomaa, 1964); although Salomaa (1957) suggested that methoxymethyl formate hydrolyses by both A-1 and A-2 mechanisms and Weeks et al. (1968) concluded that methyl pseudo2-benzoylbenzoate [ 9 11 hydrolyses by an A-2 mechanism (ratedetermining attack of H, 0 on the protonated substrate). The pH-independent reaction of yethoxy-y-butyrolactone is a

PHYSICAL ORGANIC MODEL SYSTEMS: ENZYMATIC CATALYSIS

109

unimolecular S N 1-type decomposition [92] similar to that in hydrolysis of the acetals 2-(pnitrophenoxy)tetrahydropyran [ 631 and

tropone diethyl ketal [64] where bond breaking is quite facile. y-Ethoxyr-butyrolactone possesses structural features similar to 2-(p-nitrophenoxy)tetrahydropyranin that the leaving group is quite good and the intermediate carbonium ion of moderate stability. A unimolecular mechanism was supported by the D, 0 solvent isotope effect ( k D 2 0 / k H 0 = 0.9) and lack of detectable general base catalysis even at very high concentrations of acetate or imidazole. If water acts as a nucleophile in the pH-independent reaction, much stronger bases should also function. Brown and Bruice (1973) found a similar pH-independent reaction from pH 3- 5 in hydrolysis of glucosylbenzoates. a-Acetoxy p-nitrophenyl benzyl ether [ 931 and 2-(p-nitrophenoxy)phthalide [ 941 also have large pH-independent regions in their pH-rate constant profiles (Fife and De, 1974). The AS*-value for reaction of the cyclic acylal is -47-1 e.u. and k f z 0 / k ; 2 O = 1-84

/o-No2 DNO'

o""-. 0

\

I

CH3 [931

aFO \

tl [941

so that the reaction undoubtedly involves solvent, in contrast with the unimolecular decomposition of y-ethoxy-y-butyrolactone. The rate constants for hydroxide ion catalysis and for the pH-independent reactions are similar for [93] and [94]. It is therefore likely that the mechanisms are the same, namely nucleophilic attack at the carbonyl group. Catalysis by imidazole and formate ion takes place

T. H. FIFE

110

in the hydrolysis of [93] and [94] in the pH-region where the spontaneous reaction is pH-independent. Unimolecular decomposition may not represent a favourable mechanism because the stabilization of a carbonium ion intermediate by an adjoining p-nitrophenoxy group would be weak. The pH-independent unimolecular hydrolysis of glucosyl benzoates (pH 3-5) is without buffer catalysis (Brown and Bruice, 1973). While the rates of alkaline and pH-independent hydrolysis of [93] and [94] are similar, hydronium ion catalysis is 2500 times less favourable in the case of the cyclic acylal. The large rate difference in comparison with the open chain analogue may reflect the fact that, if a carbonium ion intermediate with a carboxyl group held rigidly adjacent to the carbonium carbon atom were formed in an A-1 reaction, rapid reclosure of the ring would result from recapture of the carbonium carbon by the carboxyl group [equation (55)]. A greatly diminished rate of hydrolysis would result. Perhaps there is a

II 0

1

0

change in mechanism t o A-2 (which is supported by a Bunnett w-value of +2-3)so that a carbonium ion intermediate is not involved. This has consequence for the lysozyme reaction because, if an acylal intermediate were formed by nucleophilic attack or carbonium ion capture by aspartic acid-52, the steric situation would be similar to that of a cyclic acylal. Unimolecular decomposition would result in a carbonium ion and a carboxyl group held in juxtaposition [equation

(5211. Weeks and Crane (1973) have recently studied acid-catalysed hydrolysis of 3-methoxyphthalide in aqueous sulfuric acid. Values of AS* (-3.1 e.u.), k H /k, (0-51) and a linear plot of log kobsd us. -H, with a slope of 0.96 were consistant with an A-1 mechanism. A cyclic carbonium ion intermediate was proposed on the basis of substituent effects at the 3-position, the following reactivities being observed: H > CH3 > Et Ph. It was thought that A-1 ring opening

-

PHYSICAL ORGANIC MODEL SYSTEMS: ENZYMATIC CATALYSIS

11 1

would not explain this order. However, the acid catalysed ring opening of 2-substituted-2-phenyl-1,3-dioxolanes, proceeding by an A-1 mechanism, displays a similar order of reactivity (Fife and Hagopian, 1966). A mechanism utilizing a cyclic carbonium ion can

OH

I

be criticized on the basis that, ( 1 ) an open-chain carbonium ion with an adjacent methoxy-group would be stabilised to a greater extent by resonance interaction with oxygen; (2) the carboxyl group would be the best leaving group in the reaction; and (3) if formation of a cyclic carbonium ion is energetically favourable, it would be formed in the pH-independent reaction of compound [941 , but that reaction unquestionably involves solvent participation. Slopes of log kobsd us. -H0 deviate greatly from unity and AS* values become quite negative with increasing size of substituents at the 3-position; the 3-phenyl derivative gives a slope of 0.67 and AS* is -20.6 e.u., a value more in accord with an A-2 mechanism. Acid-catalysed hydrolysis of 3-methoxyphthalide occurs approximately 1000 times faster than reaction of 39-nitrophenoxyphthalide. A rate factor of lo3 is not consistent with a mechanism in which a cyclic carbonium ion is formed as an intermediate. Electron withdrawal in the leaving group has only a small effect on the rate in the hydrolysis of glycosides and tetrahydropyran acetals (Hammett p-values are -0.66 and -0.9 respectively), and in both cases the p-nitrophenoxy- and methoxy-derivatives hydrolyse at comparable rates. The reactivities of the phthalide acylals probably reflect differing stabilities of a ring-opened ion resulting in a difference in mechanism [A-2 in the

T. I€ FIFE

112

I

I

I

I

I

I

I

I

I

L

I

I

J

1 2 3 4 5 6 7 8 9 1 0 1 1

PH Figure 7. Plot of log k o h d us. pH for the hydrolysis of rethoxy-y-butyrolactone in HzO 0 and D2O Q (pH = pD) at 30" and p = 0.25 M.

case of [ 941 1. The influence of reversibility would be greater with the less stable carbonium ion intermediate since equilibrium would lie further on the side of protonated acylal.

Mechanistic Conclusions The problem of the mechanism of action of lysozyme can be considered in regard to possible effects by glutamic acid-35 and aspartic acid-52, the two amino-acids implicated as participants in the reaction. Glutamic acid-35 might function as an intracomplex general acid, partially transferring a proton to oxygen of the leaving group in the transition state or by carboxylate anion stabilization on oxygen of a completely transferred proton. General acid catalysis of acetal hydrolysis must now be regarded as a perfectly reasonable chemical mechanism since it has been unambiguously demonstrated as bimolecular buffer acid catalysis in a number of systems and relatively unambiguously as intramolecular carboxyl group catalysis of hydrolysis of types of acetals subject t o buffer acid cataIysis. A point of concern is that structural features in an acetal shown t o be necessary for general acid catalysis (a good leaving group or an extraordinarily stable carbonium ion intermediate) are not present in normal substrates for lysozyme. Intramolecular catalysis does not occur in model compounds where the leaving group is an aliphatic alcohol of high pK, even with an acetal, such as benzaldehyde methyl cis-2-carboxycyclohexyl acetal [ 801 , where the carbonium ion is a well-stabilized methoxybenzyl ion. A plausible viewpoint is

PHYSICAL ORGANIC MODEL SYSTEMS: ENZYMATIC CATALYSIS

113

that relief of ground state strain in the transition state of lysozyme reactions enhances the ease of bond breaking to the point where general acid catalysis can take place, i.e., proton transfer only need be partial. This possibility has been demonstrated in benzaldehyde di-t-butyl acetals and 2-methoxy-3,3-dimethyloxetane, but it has not yet been investigated in intramolecular systems. Stabilization of a proton on oxygen by a carboxylate anion (the mechanism suggested by Dunn and Bruice for methoxymethoxybenzoic acids) is a doubtful possibility for lysozyme. It is unlikely that such a mechanism would give large rate enhancements, and indeed those observed with methoxymethoxybenzoic acids are relatively small. The electrostatic stabilization mechanism cannot be generally favourable, for it is not seen when the leaving group is an aliphatic alcohol or with thioacetals. While it can be explained that intramolecular general acid catalysis does not take place with such compounds because of the poor leaving group, it is difficult to see why carboxyl group stabilization of a protonated intermediate might not be important, particularly with thioacetals where the low basicity of sulfur would allow only a small concentration of conjugate acid. General acid catalysis by glutamic acid-35 represents at present the mechanism for lysozyme best able to explain the kinetic and structural data. For it to occur, however, distortion of the hexose ring in subsite D to a half-chair must take place so that relief of strain in the transition state will make bond breaking sufficiently easy. A question that must be answered for this picture to be tenable is whether relief of strain can so greatly facilitate bond breaking when the carbonium ion is a glycosyl ion [see also the discussion in Fife (1972) and Atkinson and Bruice (1974)l. In regard to possible mechanistic effects by the carboxylate anion of aspartic acid-52, the proposal that it acts as a nucleophile forming an acylal intermediate in concert with general acid catalysis by glutamic acid-35 is unsupported. From the chemical work on hydrolysis of the acylal 3-@-nitrophenoxy)phthalide it can be concluded that, when the carbonium ion is not highly stabilized and a carboxyl group is held adjacent to it, as in the active site of lysozyme, the rate of acylal hydrolysis will be slow and may involve solvent or nucleophile participation. Reversal of the enzyme reaction through the same transition state as in the forward direction would require an SN 2-type attack of alcohol at C-1 of the hexose ring if an acylal intermediate were formed. There is no indication from studies of acylal hydrolysis that this represents a feasible mechanism. On the

114

T. H. FIFE

other hand, reactions catalysed by lysozyme are readily reversible, as shown by the occurrence of transglycosidation. The simplest explanation is that an acylal intermediate is not produced in the rate-determining step. Formation of an acylal after the ratedetermining step by carbonium ion capture by the aspartate anion is an attractive possibility in explaining transglycosidation but should lead to detectable concentrations of acylal. The alternative role for direct involvement of aspartic acid-52 is electrostatic stabilization of an oxocarbonium ion, a reasonable expectation in view of the steric situation in the active site, the poor internal stabilization of a glycosyl carbonium ion, and the apparent need for some kind of stabilization if general acid catalysis by glutamic acid-35 is t o occur. Nevertheless, this assignment is questionable since electrostatic stabilization has not been demonstrable with o-methoxymethoxyisophthalic acid (Dunn and Bruice, 1970) and benzaldehyde disalicyl acetals (Anderson and Fife, 1973). In the former example, such stabilization should be favourable because the carbonium ion intermediate is unstable. For general acid catalysis to occur, bond breaking must be easy. The transition state will then be reached earlier along the reaction coordinate than with an acetal where bond breaking is difficult. Electrostatic stabilization, however, will be of maximal importance when the carbonium ion is unstable and the transition state closely resembles the carbonium ion. Thus, structural features in an acetal that facilitate one type of catalysis will tend to inhibit the other. It is not surprising that bifunctional catalysis of the type postulated for lysozyme has not been demonstrated in acetal hydrolysis. The best opportunity for observation of concerted bifunctional catalysis might be provided by an acetal where bond breaking in the transition state is midway between protonated acetal and carbonium ion, but in such a case electrostatic catalysis would not be highly efficient. Bell-shaped pH-rate profiles are obtained in lysozyme reactions (Rupley et al., 1967) which are consistent with direct involvement of two groups in the reaction. However, bell-shaped pH-rate constant profiles are also observed in the hydrolysis of benzaldehyde disalicyl acetals, and in the case of p-nitrobenzaldehyde o-carboxyphenyl p-carboxyphenyl acetal only one carboxyl group can participate. One should then take care in postulating bifunctional catalysis in the lysozyme reaction, since the observed kinetics and the rate enhancements are explicable in terms of a chemically simpler mechanism (general acid catalysis by glutamic acid-35 along with release of ground state

PHYSICAL ORGANIC MODEL SYSTEMS: ENZYMATIC CATALYSIS

115

strain). It is possible that aspartic acid-52 influences the reaction by stabilizing a particular conformation of the enzyme or by some other indirect effect without direct involvement in the critical transition state.

6. CONCLUSION The three hydrolytic enzymes that have been discussed, a-chymotrypsin, carboxypeptidase A, and lysozyme, cover a wide range of substrate types and mechanistic possibilities. Formulation of principles which might apply to enzymatic catalysis in general is difficult from such a small sampling, but certain features of the enzymatic and model reactions warrant some comment. First and most important, it is clear that enzymatic rates can be approached in simple chemical models where the mechanism involves monofunctional catalysis or a slightly more complicated possibility in which two intramolecular groups act in a stepwise manner. In no case has concerted intramolecular general acid-general base bifunctional catalysis been observed in aqueous solution even though systems have been studied which were specifically designed to afford maximum opportunity for demonstration of such catalysis. As a consequence, it is not invariably necessary to postulate chemically complex mechanisms for enzymes catalysing analogous reactions. The chemical difficulties of bifunctional mechanisms are great. Perhaps by ridding ourselves of the idea that they are essential in order to attain enzymatic rates of reaction, greater progress can be made. A more specific generalization that might be made is that nucleophilic attack in enzyme reactions will usually be by nucleophiles bearing a negative charge or at least a partial negative charge, although the great efficiency of such nucleophiles in intramolecular and presumably intracomplex reactions cannot at present be completely explained. The role of neutral nitrogen bases would appear to be preferentially that of general bases either assisting nucleophilic attack or facilitating the reaction by catalysing breakdown of stable intermediates. Intracomplex general acid catalysis or its kinetic equivalent may be necessitated when the bond breaking process is difficult, as seen in the mechanisms suggested for all three enzymes. The presence of other factors may be required, however, before general acid catalysis will be effective.

116

T. H. FIFE

The importance of factors such as strain and orientation of reacting groups in enzymatic reactions is difficult to assess in quantitative terms even when the precise geometry of the active site is known from x-ray crystallographic studies. It is probable that these factors are important not only in promoting bond making and bond breaking but in allowing certain mechanisms t o take place. Likewise, the influence of orientation of reacting groups on mechanism goes beyond mere accessibility to the reaction centre, as seen, for example, in the general base catalysis by imidazole observed in cyclization of 2-hydroxymethylbenzamide. The hydroxyl group is rigidly held adjacent to the carbonyl group, whereas in the cyclization of y-hydroxybutyramide and anilide the groups are not rigidly held and imidazole catalysis does not take place. The greater ease of nucleophilic attack when the groups are properly aligned may permit proton transfer steps to become partially rate-limiting. Thus, mechanism, structure of the active site, and the steric situation resulting when the substrate is bound in the active site are intimately interrelated. It may be that a single “perfect” model for any enzyme is not possible. Nevertheless, by studying the chemical features of enzyme reactions and by designing models which mimic the action of enzymes as closely as possible, greater understanding may be gained of enzymatic catalysis and, in general, of chemical catalysis in aqueous solution. Certainly, models of ever-increasing sophistication will be forthcoming in the future. REFERENCES Aldersley, M. F., Kirby, A. J., and Lancaster, P. W. (1972). Chem. Commun. 834. Alexander, M. D., and Busch, D. H. (1966).J. Amer. Chem. SOC. 88, 1130. Anderson, B. M., Cordes, E. H., and Jencks, W. P. (1961). J. BioL Chem. 236, 455. Anderson, E., and Capon, B. (1969). J. Chem. SOC. ( B ) , 1033. Anderson, E., and Fife, T. H. (1969).J. Amer. Chem. SOC. 91, 7163. Anderson, E., and Fife, T. H. (1971a). Chem. Comm. 1470. Anderson, E., and Fife, T. H. (1971b).J. Amer. Chem. SOC. 93, 1701. Anderson, E., and Fife, T. H. (1973). J . Amer. Chem. SOC.95, 6437. Atkinson, R. F., and Bruice, T. C. (1974).J. Amer. Chem. SOC. 96, 819. Banerjee, S. K-, Kregar, I., Turk,V., and Rupley, J. A. (1973).J. Biol. Chem. 248, 4786. Belke, C. J., Su, S. C. K., and Shafer, J. A. (1971). J. Amer. Chem. SOC. 93, 4552. Bender, M. L. (1951).J. Amer. Chem. SOC. 73, 1626. Bender, M. L. (1957).J. Amer. Chem. SOC. 79, 1258.

PHYSICAL ORGANIC MODEL SYSTEMS: ENZYMATIC CATALYSIS

117

Bender, M. L. (1960).Chem. Rev. 60,55. Bender, M. L. (1971).“Mechanisms of Homogeneous Catalysis from Protons t o Proteins”, Wdey-Interscience, New York, N.Y. Bender, M. L., Chloupek, F., and Neveu, M. C. (1958a).J. Amer. Chem. SOC.89,

5384.

Bender, M. L., Chow, Y., and Chloupek, F. (195813).J. Amer. Chem. SOC. 80,

5380.

Bender, M. L., and Hamilton, G. A. (1962).J.Amer. Chem. SOC. 84, 2570. Bender, M. L., and Heck, H. d’A. (1967).]. Amer. Chem. SOC.89, 1211. Bender, M. L., andHomer, R. B. (1965).J.Org. Chem. 30,3975. Bender, M. L., and KCzdy, F. J. (1964).J . Amer. Chem. SOC. 86,3704. Bender, M. L., KCzdy, F. J., and Gunter, C. R. (1964).J.Amer. Chem. SOC. 86,

3 7 14.

Bender, M. L., Kizdy, F. J., and Zerner, B. (1963).J. Amer. Chem. SOC.85,

301 7.

Bender, M. L., and Lawlor, J. M. (1963).J.Amer. Chem. SOC. 85,3010. Bender, M. L., and Nakamura, K. (1962)..I. Amer. Chem. SOC. 84, 2577. Bender, M. L., and Turnquest, B. W. (1957a).J. Amer. Chem. SOC. 79, 1652,

1656.

Bender, M. L., andTurnquest, B. W. (1957b).J. Amer. Chem. SOC. 79, 1889. Bender, M. L., and Glasson, W. A. (1959).J.Amer. Chem. SOC. 81, 1590. Bender, M. L., and Wedler, F. C. (1972).J.Amer. Chem. SOC. 94, 2101. Blackall, E. L., and Eastham, A. M. (1955).J. Amer. Chem. SOC.77, 2184. Blackburn, G. M.,and Jencks, W. P. (1968).J . Amer. Chem. SOC.90, 2638. Blake, C. C. F., Koenig, D. F., Mair, G. A., North, A. C. T., Phillips, D. C., and Sanna, V. (1965).Nature 206, 757. Blow, D. M., Birktoft, J. J., and Hartley, B. S. (1969).Nature 221, 337. Boyer, P. D. (1970).Ed., “The Enzymes”, 3rd ed., Academic Press, New York. Bradshaw, R. A, Ericsson, L. H., Walsh, K. A., and Neurath, H. (1969).Proc. Natl. Acad. S c i U.S. 63, 1389. Brass, H. J., and Bender, M. L. (1973).J.Amer. Chem. SOC.95, 5391. Breslow, R. (1971).I n “Bioinorganic Chemistry”, Advances in Chemistry Series, R. P. Gould, Ed., American Chemical Society, Washington, D.C. Breslow, R., and Chipman, D. J. (1965).J. Amer. Chem. SOC.87,4195. Breslow, R., Fairweather, R., and Keana, J. (1967).J. Amer. Chem. SOC. 89,

2135.

Breslow, R., and McAllister, C. (1971).]. Amer. Chem. SOC. 93, 7096. Breslow, R., and Overman, L. E. (1970).J.Amer. Chem. SOC. 92, 1075. Bromilow, R. H., and Kirby, A. J. (1972).J.C.S. Perkin I1 149. Bronsted, J. N., and Wynne-Jones, W. F. K. (1929).Trans. Faraday SOC.25, 59. Brown, A., and Bruice, T. C. (1973).J.Amer. Chem. SOC.95, 1593. Bruice, T. C. (1959).]. Amer. Chem. SOC. 81,5444. Bruice, T. C. (1970). In “The Enzymes,” Third Edition, P. D. Boyer, Ed., Academic Press, New York, Vol. 11, p. 217. Bruice, T. C.,and Benkovic, S. J. (1963).J.Amer. Chem. SOC.85, 1. Bruice, T. C., and Benkovic, S. J. (1964).J. Amer. Chem. SOC. 86,418. Bruice, T. C., and Benkovic, S. J., (1966). “Bio-organic Mechanisms”, W. A. Benjamin, New York, N.Y. Bruice, T. C., and Bradbury, W. A. (1965).J. Amer. Chem. SOC. 87,4838. Bruice, T. C., and Bradbury, W. A. (1968).J. Amer. Chem. SOC. 90,3808. Bruice, T. C., Brown, A., and Harris, D. 0. (1971).Proc. Nut. Acad. S c i , US.

68, 658.

Bruice, T. C., and Dunn, B. (1973).Adv. Enzymol. 39, 1.

118

T. H. FIFE

Bruice, T. C., and Fife, T. H. (1962).J. Amer. Chem. SOC. 84, 1973. Bruice, T. C., Fife, T. H., Bruno, J. J., and Benkovic, P. (1962a). J. Amer. Chem. SOC. 84, 3012. Bruice, T. C., Fife, T. H., Bruno, J. J., and Brandon, N. E. (1962b). Biochemistry 1, 7. Bruice, T. C., and Lapinski, R. (1958). J. Amer. Chem. SOC. 80, 2265. Bruice, T. C., and Marquardt, F. H. (1962). J. Amer. Chem. SOC. 84, 365. Bruice, T. C., and Pandit, U. K. (1960a). J. Amer. Chem. SOC. 82, 3386. Bruice, T. C., and Pandit, U. K. (1960b). J. Amer. Chem. SOC.82, 5858. Bruice, T. C., andpiszkiewicz, D. (1967).j. Amer. Chem. SOC. 89, 3568. Bruice, T. C., and Schmir, G. L. (1957). J. Amer. Chem. SOC. 79, 1663. Bruice, T. C., and Sturtevant, J. M., (1959). J. Amer. Chem. SOC. 81, 2860. Bruice, T. C., and Tanner, D. W. (1965). J. Org. Chem. 30, 1668. Bruice, T. C., and Turner, A. (1970).J. Amer. Chem. SOC. 92, 3422. Bruice, T. C., and York, L. (1961).J. Amer. Chem. SOC. 83, 1382. Buckingham, D. A., Davis, C. E., Foster, D. M., and Sargeson, A. M. (1970b). J. Amer. Chem. SOC. 92, 5571. Buckingham, D. A., Dekkers, J., Sargeson, A. M., and Wein, M. (1972).J. Amer. Chem. SOC. 94,4032. Buckingham, D. A, Foster, D. M., and Sargeson, A. M. (1970a).J. Amer. Chem. SOC. 92, 5701. Buckingham, D. A., Foster, D. M., and Sargeson, A. M. ( 1 9 7 0 ~ )J. . Amer. Chem. SOC. 92, 6151. Buckingham, D. A., Marzilly, L. G., and Sargeson, A. M. (1967a). J. Amer. Chem. SOC. 89, 2772. Buckingham, D. A., Collman, J. P., and Happer, D. A. R. (1967b). J. Amer. Chem. SOC. 89, 1082. Buckingham, D. A., Marzilly, L. G., and Sargeson, A. M. ( 1 9 6 7 ~ ) J. . Amer. Chem. SOC. 89,4539. Bull. H. G., Koehler, K., Pletcher, T. C., Ortiz, J. J., and Cordes, E. H. (1971). J. Amer. Chem. SOC. 93, 3002. Bunnett, J. F., and Hauser, C. F. (1965).J. Amer. Chem. SOC. 87, 2214. Butler, A. R., and Gold, V. (1960). Proc. Chem. SOC. 15. Butler, A. R., and Gold, V. (1961). J. Chem. SOC. 2305. Canfield, R. C. (1963). 1. Biol. Chem. 238, 2699. Caplow, M., and Jencks, W. P. (1962). Biochemistry 1, 883. Caplow, M. (1969). J. Amer. Chem. SOC. 91, 3639. Capon, B. (1963). Tetrahedron Lett. 911. Capon, B. (1969). Chem. Rev. 69,407. Capon, B., McDowell, S, T., and Raftery, W. V. (1973).J.C.S. Perkin 11, 1118. Capon, B., andsmith, M. C. (1965). Chem. Comm. 523. Capon, B., Smith, M. C., Anderson, E., Dahm, R. H., andSankey, G. H. (1969). J . Chem. SOC. 1038. Capon, B., and Thacker, D. (1965).]. Amer. Chem. SOC.87, 4199. Chance, B. (1943).J. Biol. Chem. 151, 553. Chance, B. (1947). Acto. Chem. Scand. 1, 236. Chanley, J. D., and Feageson, E. (1955).J. Amer. Chem. SOC. 77, 4002. Chipman, D. M., and Sharon, N. (1969). Science 165,454. Conley, H . L., Jr., and Martin, R. B. (1965). J. Phys. Chem. 69, 2914, 2923. Cordes, E. H. (1967). Progr. Phys. Org. Chem. 4, 1. Cramer, F., Saenger, W., and Spatz, H-Ch (1967).J. Amer. Chem. SOC. 89, 14. Craze, G., and Kirby, A. J. (1974). J. Chem. SOC. ( B ) 61. Cunningham, B. A., and Schmir, G. L. (1966).J. Amer. Chem. SOC. 88, 551.

PHYSICAL. ORGANIC MODEL SYSTEMS: ENZYMATIC CATALYSIS

119

Cunningham, B. A., andSchmir, G. L. (1967). J. Amer. Chem. SOC. 89, 917. Dahlquist, F. W., and Raftery, M. A. (1968). Biochemistry 9, 3277. Dahlquist, F. W., Rand-Meir, T., and Raftery, M. A. (1969). Biochemistry 8 , 4214. Dickerson, T. A., and Geis, S. M. (1969). “Structure and Action of Proteins”. Harper-Row, New York. Dunn, B., and Bruice, T. C., (1970).J. Amer. Chem. SOC.92, 2410. Dunn, B., and Bruice, T. C., (1971).J. Amer. Chem. SOC. 93, 5725. Dyer, E., Glaudemans, C. P. J., Koch, M. J., and Marchessault, R. H. (1962). J. Chem. SOC. 3361. Eastham, A. M., Blackall, E. L., and Latremouille (1955). J. Amer. Chem. SOC. 77, 2182. Eberson, L. (1964). A c t n Chem. Scand 18, 2015. Edwards, L. J. (1952). Trans Faraday SOC. 48, 696. Eigen, M. (1964). Angew Chem. Ink E d 3, 1. Eliel, E. L. (1962). ‘Stereochemistry of Carbon Compounds”, McGraw-Hill, New York, p. 196. Eriksson, S. O., and Holst, C. (1966). Acta Chem. Scand. 20, 1892. Eyring, H., Lumry, R., and Spikes, J. D. (1954). In “The Mechanism of Enzyme Action”, W. D. McElroy and B. Glass, Eds., Johns Hopkins Press, Baltimore, p. 123. Fedor, L. R.. and Bruice, T. C . (1964).J. Amer. Chem. SOC. 86,4117 Felton, S. M:, and Bruice, T. C. -(1968).-Chem.Comm. 907. Felton, S. M., and Bruice, T. C. (1969). J. Amer. Chem. SOC. 91, 6721. Fersht, A. R. (1971). J. Amer. Chem. SOC. 93, 3504. Fersht, A. R. (1972). J. Amer. Chem. SOC.94, 293. Fersht, A. R., andKirby, A. J. (1967). J. Amer. Chem. SOC. 89, 4857, 4953. Fersht, A. R., and Kirby, A. J. (1968a).J. Amer. Chem. SOC.90, 5818. Fersht, A. R., and Kirby, A. J. (1968b). J. Amer. Chem. SOC.90, 5833. Fersht, A. R., and Raquena, Y. (1971). J. Amer. Chem. SOC. 93, 7079. Fife, T. H. (1965). J. Amer. C h e m SOC. 87, 271. Fife, T. H. (1967).J. Amer. C h e m SOC.89, 3228. Fife, T. H. (1972). Acc. C h e m Res. 5 , 264. Fife, T. H., and Anderson, E. (1970). J. Amer. Chem. SOC. 92, 5464. Fife, T. H., and Anderson, E. (1971a). J. Amer. Chem. SOC. 93, 6610. Fife, T. H., and Anderson, E. (1971b). J. Org. Chem. 36, 2357. Fife, T. H., and Benjamin, B. (1973). J. Amer. Chem. SOC. 95, 2059. Fife, T. H., and Benjamin, B. (1974). Chem. C o m m u n 525. Fife, T. H., and Brod, L. H. (1970).J. Amer. Chem. SOC. 92, 1681. Fife, T. H., and Brod, L. H., (1968).J. Org. Chem. 33, 4136. Fife, T. H., and De, N. C. (1974).J. Amer. Chem. SOC. 96, 6158. Fife, T. H., and Hagopian, L. (1966). J. Org. Chem. 31, 1773. Fife, T. H., and Hutchins, J. E. C. (1972).J. Amer. Chem. SOC. 94, 2837. Fife, T. H., Hutchins, J. E. C., and McMahon, D. M. (1972).J. Amer. Chem. SOC. 94, 1316. Fife, T. H., and Jao, L. K. (1965). J. Org. Chem. 30, 1492. Fife, T. H., and Jao, L. K. (1968). J. Amer. Chem. SOL. 90, 4081. Fife, T. H., and McMahon, D. M. (1970). J. Org. Chem. 35, 3699. Fife, T. H., and Milstien, J. B. (1967). Biochemistry 6, 2901. Frost, L. N., andHegarty, A. F. (1973). Chem. Comm. 82. Gaetjens, E., and Morawetz, H. (1960). J. Amer. Chem. SOC. 82, 5328. Garrett, E. R. (1957).J. Amer. Chem. SOC. 79, 3401, 5206. Gibson, K. D., and Scheraga, H. (1967). PYOC.Nut. A c a d S c i U.S. 58,420.

120

T. H. FIFE

Giudici, T. A., and Bruice, T. C. (1970). Chem. Comm. 690. Goitein, R., and Bruice, T. C. (1972). J. Phys. Chem. 76, 432. Gold, V., Oakenfull, D. G. and Riley, T. (1968). I. Chem. SOC. ( B ) 515. Hartsuck, J. A., and Lipscomb, W. N. (1971). I n “The Enzymes,” 3rd Edition, Vol. 3, P. D. Boyer, Ed., Academic Press, New York. Hegarty, A. F., and Frost, L. N. (1972). Chem. Comm. 500. Henbest, H. B., and Lovell, B. J. (1957). J. Chem. SOC.1965. Henderson, R. (1970). J. Mol. Biol. 54, 341. Higuchi, T., Eberson, L, and Herd, A. K. (1966). J. Amer. Chem. SOC.88, 3805. Higuchi, T., Takechi, H., Pitman, 1. H., and Fung, H. L. (1971).]. Amer. Chem. SOC. 93, 539. Hofmann, K., and Bergmann, M. (1940). J. Biol. Chem. 134,225. Hoppe, J. I., and Prue,J. E. (1957).J. C h e m SOC. 1775. Hubbard, C. D., Kirsch, J. F. (1972). Biochemistry 11, 2483. Hunkapiller, M. W., Smallcombe, S. H., Whitaker, D. R., and Richards, J. H. (1973). Biochemistry 12, 4732. Hutchins, J. E. C., and Fife, T. H. (1972).J. Amer. Chem. SOC.94, 8848. Hutchins, J. E. C., and Fife, T. H. (1973a).J. Amer. Chem. SOC. 95, 2282. Hutchins, J. E. C., and Fife, T. H. (1973b)..]. Amer. Chem. SOC.95, 3786. Hybl, A.,-Rundle, R. E., and Williams, D.-E. (1965). J. Amer. C h e m SOC. 87, 2779. Inagami, T., York, S. S., and Patchornik, A. (1965). J. Amer. Chem. SOC.87, 126. Inward, P. W., andJencks, W. P. (1965). J. Biol. Chem. 240, 1986. Jencks, W. P. (1963). Ann Rev. Biochem 32, 639. Jencks, W. P. (1969). “Catalysis in Chemistry and Enzymology”, McGraw-Hill, New York, N.Y. Jencks, W. P. (1973). Chem. Rev. 73, 705. Jencks, W. P. (1972).J. Amer. Chem. SOC. 94, 4731. Jencks, W. P., andCarriuolo, J. (1960). J. Amer. Chem. SOC. 82, 1778. Jencks, W. P., and Carriuolo, J. (1961). J. Amer. Chem. SOC. 83, 1743. Jencks, W. P., and Gilchrist, M. (1962).J. Amer. Chem. SOC. 84, 2910. Jencks, W. P., and Gilchrist, M. (1968).]. Amer. Chem. SOC. 90, 2622. Jollhs, P. (1964). Angew. Chem. Int. Ed. 3, 28. Johnson, L. N., and Phillips, D. C., (1965). Nature 206, 761. Johnson, S. L. (1967). Adv. Phys. Org. C h e m 5, 237. Kaiser, E. T. and Kaiser, B. L. (1972). Acc. Chem. Res. 5, 219. Kang, E. P., a n d s t o r m , C. B. (1972). Biochem. Biophys. Res. Commun. 49, 621. Kirby, A. J., and Fersht, A. R. (1971). In “Progress in Bioorganic Chemistry,” E. T. Kaiser and F. J. Kkzdy, Ed. Wiley, New York, Vol. 1. p. 1. Kirsch, J. F., and Jencks, W. P. (1964). /. Amer. Chem. SOC. 86, 833, 837. Koshland, D. E., Jr. (1958). Proc. Nut. A c a d SOC.,U.S. 44, 98. Koshland, D. E., Jr. (1962). J. Theoret. Biol. 2, 75. Kroll, H. (1952).]. Amer. Chem. SOC. 74, 2036. Kupchan, S. M., Johnson, W. S., and Rajagopalan, S. (1959). Tetrahedron 7, 47. Kupchan, S. M., and Narayanan, C. R. (1959).J. Amer. Chem. SOC. 81, 1913. Kupchan, S. M., Eriksen, S. P., and Friedman, M. (1962). J. Amer. Chem. SOC. 84, 4159. Kupchan, S. M., Eriksen, S. P., and Friedman, M. (1966a). J. Amer. Chem. SOC. 88, 343. Kupchan, S . M., Eriksen, S. P., and Liang, Y. T. (1966b). J. Amer. Chem. SOC. 88, 347. Kupchan, S. M., and Johnson, W. S. (1956).J. Amer. Chem. SOC. 78, 3864.

PHYSICAL ORGANIC MODEL SYSTEMS: ENZYMATIC CATALYSIS

121

Kupchan, S. M., Slade, P., and Young, R. J. (1960). Tetrahedron Lett. 24, 22. Lipscomb, W. N. (1970). Acc. Chem. Res. 3, 81. Lipscomb, W. N., Hartsuck, J. A, Reeke, G. N., Jr., Ouiocho, F. A., Bethge, P. H., Ludwig, M. L., Steitz, T. A., Muirhead, H., and Coppola, J. C. (1968). Brookhaven Symp. BioL 21, 24. Lipscomb, W. N., Reeke, G. N., Jr., Ouiocho, F. A.. and Bethge, P. H. (1970). Phil. Trans. Roy. SOC.( B ) 251, 177. Lowe, G., Sheppkd, G.,'Sinnott, M. L., and Williams, A. (1967). Biochem. J. 104, 893. Lowry, T. M., and Faulkner, I. J. (1925). J. Chem. SOC. 2883. Lowry, T. M., and Faulkner, I. J. (1927). Z . Physik Chem. Leiprig 130, 125. Lukas, E. C., Caplow, M., and Bush, K. J. (1973).J. Amer. Chem. SOC. 95, 2670. Lumry, R., and Rajender, S. (1971).J. Phys. C h e m 75, 1387. Martell, A. E. (1968). Pure AppL Chem. 17, 129. Martinek, K., Dorovska, V. N., Varfolomeyev, S. D., and Berezin, I. V. (1972). Biochim Biophys. Acta 271, 80. Maugh, T., and Bruice, T. C. (1971).J. Amer. Chem. SOC. 93, 3237. Menger, F. M., and Smith, J. H. (1969). J. Amer. Chem. SOC. 91, 5346. Meriwether, L., and Westheimer, F. H. (1956).J. Amer. Chem. SOC.78, 5119. Milstien, J. B., and Fife, T. H. (1968). J. Amer. Chem. SOC. 90, 2164. Milstien, J. B., and Fife, T. H. (1969).Biochemistry 8, 623. Milstien, S., and Cohen, L. (1972).J. Amer. Chem. SOC. 94, 9158. Morawetz, H., and Oreskes, I. (1958).J. Amer. Chem. SOC. 80, 2591. Nath, R. L., and Rydon, H. N. (1954). Biochem. J. 57, 1. Oakenfull, D. G., and Jencks, W. P. (1971). J. Amer. Chem. SOC. 93, 178. Oakenfull, D. G., Riley, T., and Gold, V. (1966). Chem. Comm. 385. Oakenfull, D. G., Richardson, D. I., and Usher, D. A., (1967). J. Amer. Chem. SOC.89,5491. Okuyama, T., and Schmir, G. L. (1972). J. Amer. Chem. SOC.94, 8805. O'Leary, M. H., and Kluetz, M. D., (1972).J. Amer. Chem. SOC. 94, 3585. Osawa, T., and Nakasawa, Y . (1966). Biochim. Biophys. Acta 130, 56. Page, M., and Jencks, W. P. (1971). Proc. Nut. Acad Sci, 1J.S. 68, 1678. Page, M. I., andJencks, W. P. (1972).]. Amer. Chem. SOC. 94, 8818. Parsons, S. M., and Raftery, M. A. (1972). Biochemistry 11, 1623. Phillipp, M., Pollack, R. M., and Bender, M. L. (1973). Proc. Nut. Acad. Sci U . S . 70, 517. Phillips, D. C. (1969). Sci Amer. 215, 78. Piszkiewicz, D., and Bruice, T. C. (1967). J. Amer. Chem. SOC. 89, 6237. Piszkiewicz, D., and Bruice, T. C. (1968a). J. Amer. Chem. SOC. 90, 5844. Piszkiewicz, D., and Bruice, T. C. (196813). J. Amer. Chem. SOC. 90, 2156. Pocker, Y. (1960). Chem. Znd (London) 968. Polgar, L., and Bender, M. L. (1969). Proc. Nut. Acad. Sci, U S . 64, 1335. Raftery, M. A., and Rand-Meir, T. (1968). Biochemistry 7, 3281. Riordan, J. F., and Muszynska, G. (1974). Biochem. Biophys. Res. Commun. 57, 447. Robertus, J. D., Kraut, J., Alden, R. A., and Birktoft, J. J. (1972). Biochemistry 11, 4293. Rogers, G. A., and Bruice, T. C. (1974). J. Amer. Chem. SOC. 96, 2463. Rony, P. R. (1968).J. Amer. Chem. SOC. 90, 2824. Rupley, J. A. (1967). Proc. Roy. SOC. B167, 416. Rupley, J. A., and Gates, V. (1967). Proc. Nut. A c a d Sci U.S. 57, 496. Rupley, J. A., Gates, V., and Bilbrey, R. (1968). J. Amer. Chem. SOC. 90, 5633. Salomaa, P. (1957). Acta Chem. Scand 11, 132, 141, 239.

122

T. H. FIFE

Salomaa, P. (1964). Suomen Kemistilehti B37, 86. Salomaa, P., and Laiho, S. (1963). Acta Chem. Scand 17, 103. Secemski, I. I., and Lienhard, G. E. (1971).J. Amer. Chem. SOC. 93, 3549. Sigman, D. S., Wahl, G. M., andcreighton, D. J. (1972). Biochemistry 11, 2236. Smidsrod, O.,Haug, A., and Larson, B. (1966). Actn Chem. Scund 20, 1026. Snell, R. L., Kwok, W. K., and Kim, Y. (1967).J. Amer. Chem. SOC.89, 6728. Snoke, J. E., Schwert, G. W., and Neurath, H. (1948). I. BioL Chem. 175, 7. Speck, J. C., Rynbrandt, D. J., and Kochevar,‘I. H. (1965). J. Amer. Chem SOC. 87, 4979. Storm, D. R., and Koshland, D. E., Jr. (1970). Proc. Nut. Acad Sci, U.S. 66, 445. Storm, D. R., and Koshland, D. E., Jr. (1972a). J. Amer. Chem. SOC. 94, 5805. Storm, D. R., and Koshland, D. E., Jr. (1972b).J. Amer. Chem. SOC. 94, 5815. Swain, C. G., and Brown, J. F. (1952). J. Amer. Chem. SOC. 74, 2534. Taft, R. W., Jr. (1956). In “Steric Effects in Organic Chemistry”, M. S. Newman, Ed., Wiley, New York, N.Y., p. 556. Thanassi, J. W., and Bruice, T. C. (1966). J. Amer. Chem. SOC. 88, 747. Usher, D. A., Richardson, D. I., Jr., and Oakenfull, D. G. (1970). J. Amer. Chern. SOC.92,4699. Van Etten, R. L., Clowes, G. A., Sebastian, J. F., and Bender, M. L. (1967a).J. Amer. Chem. SOC. 89, 3242. Van Etten, R. L., Clowes, G. A., Sebastian, J. F., and Bender, M. L. (1967b).J. Amer. Chem SOC. 89, 3253. Vernon, C. A., (1967). Proc. Roy. SOC. ( B ) 167, 389. Wang, J. H., andParker, L. (1967). Proc. Nut. Acad Sci, US.58, 2451. Weeks, D.P., andcrane, J. P. (1973). J. Org. Chem. 38, 3375. Weeks, D. P., Grodski, A., and Fanucci, R (1968). J. Amer. Chem. SOC.90, 4958. Westheimer, F. H. (1962). Adu. Enzymol. 24, 441. Williams, A. (1970). Biochemistry 9, 3383. Winstein, S.,Smith, S., and Darwish, D. (1959). J. Amer. Chem. SOC. 81, 5511. Wright, C. S., Alden, R. A., and Kraut, J. (1969). Nutwe 221, 235. Zachau, H. G., and Karau, W. (1960). Ber. 93, 1830. Zeffren, E., and Hall, P. L. (1973). “The Study of Enzyme Mechanisms”, John Wiley and Sons, New York, N.Y., p. 127.

Charge Density-NMR Chemical Shift Correlations in Organic Ions

D. G . FARNUM Department of Chemistry, Michigan State University, East Lansing, Michigan, U.S.A. 1. 2.

3.

4.

.

Introduction Theory General . Pnmr and Cnmr screening parameters summary . Empirical Observations . . General Basic correlations . Selected problems Conclusions . . References

.

. . . .

.

. . . . . .

123 126 126 127 135 135 135 136 148 172 173

1. INTRODUCTION Five years ago a brief review focused on the applications of nuclear magnetic resonance (nmr) as a method for determining charge density in carbonium ions and pointed out some of the precautions required (Fraenkel and Famum, 1968). Since then, proton nmr (pnmr), which was emphasized in that review, has continued to attract primary attention as a probe into the structure and charge density of organic cations and anions (Olah and Schleyer, 1968,1970,1972, 1973; 0 t h e t al., 1972; Takahashi e t al., 1973; van

124

D. G. FARNUM

der Kooij et al., 1972). Recently, its supremacy as a charge-density probe has been challenged by carbon- 13 nmr (cnmr), whose potential was noted in the early review, but whose practical applications awaited the development of more nearly routine instrumentation (Stothers, 1972). AIthough there are many excellent general reviews of the principles and applications of both pnmr (Becker, 1969; Bovey, 1969; Jackman and Sternhell, 1969; Paudler, 1971; Dixon, 1972) and cnmr (Levy and Nelson, 1972; Stothers, 1972) which treat, in part, their use in determining charge densities, there does not seem to be a comprehensive review which concentrates on a comparison of the two methods, both in theory and in practice, as charge density probes. It is because of this need, particularly at a time when moderately priced, routine cnmr instrumentation promises to swell the already perplexing flood of data appearing in the literature, that this review is presented. The common division into sections on theoretical concepts and empirical observations is followed, but pnmr and cnmr are treated jointly where feasible in each of these sections. The section on theory is unlikely to satisfy those theorists who would like a critical presentation and analysis of conflicting theories. That task, important as it is, is well outside the competence of this reviewer. The presentation here is directed to practising organic chemists who could use a qualitative description of some key concepts emerging from the theoretical treatments to help them avoid possible pitfalls in data interpretation, and prompt them to design models to test some of the ideas. The section on empirical observations is designed to flesh out the theoretical bones as well as provide some data which raise some puzzling problems. The data presented, therefore, do not even approach an exhaustive survey, but are chosen in accordance with the bias of the reviewer t o illustrate points felt to be of particular significance. A less selective approach would obscure the purpose. No attempt has been made in this review to treat other nuclei, since they have not enjoyed the general attention that pnmr and cnmr have. Nonetheless, both theoretical and experimental effort has been spent on fluorine-19 (Adcock et al., 1973; Timberlake et aZ., 1971), and, more recently, oxygen-17 and nitrogen-14 (Mateescu, 1973) nmr as charge density probes. The following conventions and assumptions will be used in reporting data: (1) The practice of reporting both proton and carbon-13 chemical shifts in parts per million (p.p.m.) downfield

CHARGE DENSITY-NMR CHEMICAL SHIFT CORRELATIONS

125

from tetramethylsilane (TMS) will be followed, with a positive sign indicating values downfield. Carbon-13 values are rounded to the nearest p.p.m. (2) Where the reference used experimentally was other than TMS, it is designated in parentheses, along with its chemical shift value used to convert the data to TMS. In some cases, authors have determined chemical shifts using TMS and then converted them to other standards (e.g. CS,) for reporting. In such cases the chemical shift value used by the authors for that standard is given in parentheses and used to convert the data back to TMS, even though it may differ from the accepted value. For example, a chemical shift reported in this review as “214 (ext. TMS, conv. to CS, = 194-6)” means that the authors determined the chemical shift with external TMS as standard, converted it to CS, using the value 194.6 as the chemical shift of CS2 relative to TMS, and reported it relative to CS,. The reported value has been reconverted to 214 relative to external TMS using the authors conversion factor of 194.6. Proton data are reported relative to internal TMS unless otherwise indicated. Carbon-13 data are generally reported relative to external TMS, although the original work does not always state clearly whether the standard is internal or external. The difference is not likely to be significant for the comparisons made here. (3) Chemical shift differences will be reported with a negative sign indicating a downfield shift and a positive sign signifying an upfield shift. Although this may seem inconsistent with convention ( l ) ,it is in accord with most current practice. (4) Solvents used are reported in parentheses in most cases where it is clear from the original work, although in some cases of carbonium ion spectra the solvent is not indicated and may be assumed to be one of several different strong acids used interchangeably. Concentrations are not reported since they are usually not known. The effect of solvent and concentration on chemical shifts will surely be more significant for proton spectra than carbon-13 spectra, and the attempt made in this review to compare data obtained under different conditions of soIvent and concentration (and temperature as well) suffers from the scarcity of systematic studies of these effects on ion spectra. However it seems unlikely that the carbon-13 spectra will be affected too drastically by changes among solvents of comparable acidity, and, even for proton spectra, we can still make use of the trends that emerge. The related phenomenon of ion association has also generally been ignored in this review except in those few cases of carbanions where its influence has been demonstrated. Again, the effect is probably more

126

D. G. FARNUM

significant for protons than carbon-13, and the general feeling seems to be that it is more significant for carbanions than for carbonium ions in powerful acids.

2. THEORY General The development of a general theory of nmr c..emical shifts as a function of electronic structure has been the product of a number of theoretical minds, and has been reviewed by Memory (1968). As is usually the case, the rigorous theory is unmanageable for any compounds of interest, and the manageable theories involve so many approximations that they lose their rigor. Nonetheless, several approximate theories have been developed which correlate chemical shifts (both proton and carbon-13) remarkably well in closely related compounds (e.g., aromatic hydrocarbons, Lazeretti and Taddei, 1971). Furthermore, as stated several times by Grant (Pugmire et al., 1973) the value of approximate theories probably lies less hi their ability to predict chemical shifts accurateIy than in their ability to point out important trends and identify structural features of dominant importance in determining chemical shifts. The familiar separation of the screening constant for an atom A, uA, into a summation of screening constants from several conceptually identifiable factors, as in (l),provides a convenient approximation for elaboration.

Here, ui A is a positive term resulting from diamagnetic shielding by electrons “localized” on atom A, u t A is a term (usually negative) which corrects for any deviation from spherical symmetry of the electrons’ localized on atom A, u A B represents the effect on atom A (shielding or deshielding) of circulation of electrons localized on other atoms, and represents the effect on atom A (shield~

’

~

p

~

~

’

~

~

.

The up term includes excited state mixing and can be separated into both positive and negative terms (Pugmire and Grant, 1968a) corresponding to shielding and deshielding effects as noted in the next section.

CHARGE DENSITY-NMR CHEMICAL SHIFT CORRELATIONS

127

ing or deshielding) of circulation of electrons in delocalized R systems (generally “ring current” effects). The dependence of each of these terms on structure, and a comparison of the contribution each makes to the total screening constant uA for protons and carbon-13 will be explored in the next section.

Pnmr and Cnmr Screening Parameters The “other atom” terms uA and uA~deloc. These terms are treated first since, although they do contribute to uA, their contribution is independent of the nature of the atom A or the charge on atom A (Fraenkel and Farnum, 1968; Stothers, 1972). Their magnitude is therefore the same for both protons and carbon13, and is unchanged by the charge on these atoms.* Since cnmr screening constants are very large compared to u A B and uAide’oc., these “other atom” terms generally contribute 10% or less to u A and not more than 2 p.p.m. to the chemical shift, and are usually ignored in cnmr analysis. For protons, however these terms often dominate uA and the chemical shifts, and accurate corrections for them must be found and applied before chemical shifts can be used t o probe other structural features such as charge density (Fraenkel and Farnum, 1968). An analysis of the quantitative effect on atom A of electron circulations elsewhere in the molecule usually requires an accurate knowledge of the geometry of all important conformations of the molecule, as well as accurate mathematical or empirical models for the screening contribution. With the many assumptions that are necessary, it is usually difficult to make a convincing analysis (Farnum, 1967). The influence of u A B and uA9de’oc. is probably the single most important factor that obscures the effect of charge density on proton chemical shifts. It is no wonder that cnmr was greeted with such enthusiasm as a way to circumvent this problem. The local diamagnetic term, of A This term is comparable in magnitude to u A B and u ~ ~for ~ both protons and carbon-13. Before considering the effect of charge

*

This is not strictly true if the ring current contribution to r ~ ~ includes ~ ~ atom ~ A’ in the delocalized ring. However, the perturbation of an asymmetric charge distribution on the ring current generally seems to be ignored.

~

~

~

’

*

~

D. G. FARNUM

128

on ~2A , it is important to distinguish the ways in which pnmr and cnmr are used to probe charge distribution. Thus, for protons, the chemical shift is interpreted in terms of the charge on the attached carbon, while for carbon-13, the chemical shift is interpreted in terms of the charge directly on the carbon whose chemical shift is being determined. This difference has an important consequence which complicates the interpretation of proton chemical shifts. In fact, the electron density at the proton, which determines u t A , is rarely obtained directly. Rather, the electron density at the attached carbon is calculated by some procedure, and its effect on the electron density at the proton by polarization of the C-H bond is then calculated. In addition to polarization of the C-H bonding electrons by the component of the field along the C-H bond axis, E z , Buckingham (1960) has pointed out that there is also a deshielding effect on u which is proportional to E 2 [equation (211

-

A u = - aEZ - bE2

(2)

Although the E 2 correction is usually negligible, it can assume importance with high fields (i.e., high charges), and Musher (1962) has found that correction of the raw data for some aromatic cations and anions by the E 2 term improves the now well known 10 p.p.m. per unit charge correlation of Fraenkel et al. (1960) (see Section 3). He has also suggested that charged atoms elsewhere in the molecule would polarize the C-H bond, and that corrections from these terms may not be negligible. When both E Z and E 2 terms for all charged carbon atoms are considered, the pnmr spectra of phenyl carbonium ions is adequately correlated by HMO charge densities, while, when the additional terms are ignored, SCF-MO charge densities provide a better fit (Farnum, 1967). In spite of their demonstrated importance, Musher’s corrections seem to have been ignored, probably because they introduce another mathematical complication into the analysis which increases one’s scepticism and resistance. Conversion of the Musher equations into frequency units (Schweizer et al. 1964), as in equation (3), makes them easier to use. As given here, the units of Au are p.p.m., a negative sign is deshielding, q i is the total charge at the carbon in question in units of one electron (a positive charge has a positive sign), and cos Bi and R i (in A) are defined as in Figure 1 (see

CHARGE DENSITY-NMR CHEMICAL SHIFT CORRELATIONS

129

Figure 1. Definition of the terms used in equation (3).

Jackman and Sternhell, 1969, pp. 67-9).3

Aa=

-

1 3 . 3 ( 7 qi

"i)

-

17.0

(2%

R:)2

If one ignores the charge on all but the attached carbon atom. cos B i becomes 1, and the equation then R i becomes -1.09 further reduces to equation (4).Routine application of equation (4)

a,

A~=-11-2q-l2*0q~

(4)

to anions and cations could easily be done, and would allow an estimate of the importance of the q 2 term. Musher has cautioned, however, that use of the bond terminus (the carbonium ion carbon) as the vector terminus violates an assumption in the derivation of equation (3). Thus it is not clear that the theory applies in this simple form. Perhaps dramatization of the cos 0 term by preparation of a system in which a positive charge resulted in shielding of a proton of appropriate geometry as in [ l ] or [2] would call attention to its importance. In each of these geometries the polarization of the C-H bond should result in an increase in the electron density around the proton, an increase in a$ *, and an upfield shift of about 1.3 p.p.m. at a distance of 3 A.4 In Musher's original equations, the vector Ri was measured from the midpoint of the carbon-hydrogen bond to the midpoint of the carbon-carbon bonds. However, subsequent users of his equations seem to have measured the vector from atom to atom. The latter calculation is much easier, and the error inkoduced seems to be small for most molecular geometries encountered, provided the correct values for the constants are used. The value for the first constant is determined empirically to fit the observed linear dependence of u on q for the aromatic cations and anions. Differences in the magnitude of this constant in the literature reflect differences in the accepted proportionality constant and in the assumed geometrical parameters. The value given here (13'3) differs from the Schweizer value (12*5), and was calculated assuming a carbon-hydrogen bond length of 1.09 A and a proportionality constant of 11.2 [see equation (4)] as determined by Mushcr after correction for the E a term It is within the theoretically expected range of -10-16 (Buckingham, 1960; Musher, 1962). The second constant (17.0) is a theoretical estimate (Musher, 1962; Schweizer et aL 1964). If the field vector is taken at the midpoint of the carbon-hydrogen bond the predicted shift is smaller, but still greater than 0-5 p.p.m.

130

D. G. FARNUM

A clear demonstration of the importance of the E 2 term might be achieved in a system where cos 0 was 0, thus cancelling the E Z term. The remaining E 2 term is independent of the s i p of the charge, so that two systems with geometries and charges as in [3] and [4] should give the sume deshielding effect on the indicated proton relative to system [ 5 ] , other factors being equal. For R = 3 A the predicted effect is about 0.2 p.p.m. which is easily detectable.

9

131

[41

[51

By contrast, the influence of charge on udAA for cnmr shifts is determined by direct calculation of the charge on carbon atom “A”. Hence, no correction for the E 2 term or electron polarization by other charged atoms is necessary. The accuracy of the calculated charge density depends only on how faithfully the mathematical model chosen reflects reality. Grant (Pugmire et al. 1973) has emphasized that the effect of charge on u s A is twofold: (1) an increase in electron density at atom A increases the diamagnetic current density and results in an upfield shift, and (2) an increase in electron density increases the effective radius of the orbitals (by , and shielding the nuclear charge), decreases the magnitude of results in a partially compensating downfield shift. Clearly if u$* were the dominant parameter determining the effect of charge on chemical shifts for both protons and carbon-13, cnmr shifts could be more easily and reliably correlated with charge for density than could pnmr shifts. Unfortunately, although protons is negligible (Memory, 1968) because of the high energy of therefore accounts for the full 10 the proton excited states, and p.p.m. per electron (on carbon) charge dependence found for proton

utA

atA

utA

CHARGE DENSITY-NMR CHEMICAL S H F T CORRELATIONS

131

chemical shifts, u t A for carbon-13 can be quite large, and 02, accounts for only some 10% of the charge dependence found for carbon-13 (of the order of 13-18 p.p.m. per electron according to Pugmire et al. (1969)). We must therefore turn to an analysis of the factors affecting ut A if we are to understand cnmr shifts.

The local “paramagnetic” term, up”A This term in the total screening constant corrects for the symmetrical dectron distribution used to calculate the diamagnetic term us A by mixing excited electronic states with the ground states. The resultant electron distribution is a closer approximation to that of the molecule, and the resultant screening constant is a more accurate value. The term enters the total screening expression as though there were a “paramagnetic” circulation of electrons superimposed on the diamagnetic circulation, although this is simply an artifact of the calculation. Like the diamagnetic term, the paramagnetic term shows a direct dependence on electron density (i.e., if there are no electrons in an orbital, there is no u, term for that orbital, and therefore no correction, required), and an indirect dependence as a result of an increase in average orbital radius, r , with an increase in orbital electron density (actually a l / r 3 term for up). In addition there is an inverse dependence on the energy difference, AE, between the ground and excited states for promotion of the electron in question. Obviously, the l/AE term simply says the larger the energy gap between ground and excited states, the less distortion in electron distribution by excited state mixing. Clearly the l/AE dependence accounts for the fact that a t A is negligible in pnmr where electronic excited states require promotion of an electron from a 1s to a 2s or 2p orbital. The term is dominant, however, in larger atoms where excited states are much lower lying. In fact, it is as a “correction” to u t A for quite misleading to speak of carbon-13 screening constants, since changes in up are an order of magnitude greater than those in O d . The dependence of these changes on various structural parameters has been analysed by Grant (Pugmire e t al., 1968, 1973) and a conceptual paraphrase of some of his conclusions follows. Following and elaborating on the approach of Karplus and Das (1961) Grant accepts a separation of up for carbon-13 into a term ubl ) for electrons in orbitals centred entirely on the carbon atom in for electrons in orbitals centred both on question, and a term that atom as well as other atoms in the molecule (ie., “bonding”

utA,

utA

062)

132

D. G. FARNUM

orbitals). He then treats the dependence of u$l ) on orbital charge in some detail. Following the practice of many others in the field, it was assumed that the change in AE with charge could be ignored, and an average value for AE of 10 eV (Pople, 1962), was taken for computational simplification. A number of other assumptions were made at the sacrifice of quantitative accuracy in order to obtain a very useful qualitative picture of the dependence of u$') on q', the orbital electron charge. The direct dependence of of,') on q1 is graphed by the solid line in Figure 2, which is the sum of a deshielding term linear in q' and a shielding q t 2 / 2 term (dashed lines). From this treatment there emerges the surprising prediction that, for a carbon atom denuded of valence electrons (charge = +4), the addition of electron density results in an increase in the deshielding term u ; ' ) , which opposes the shielding term u d . This deshielding increases until q' = 1 (charge = 0), at which point it reverses in direction and additional increase in electron density results in increased shielding. This curious behaviour makes some sense in that, with n o electrons in the orbital there is n o fJd term, and therefore n o "correction" term, up, necessary, while with a completely occupied orbital, the effects of the q' and q f 2 terms exactly cancel (to a first approximation) and u$')goes to zero again. It is important that, if the direct dependence on q f illustrated in Figure 2 were the only effect of q' on uf,'), there would be a zero dependence of of,') on q f in the region q' = 1, the very region most investigated for charge-density-chemical shift correlations, and the region covered by the early empirical correlation of Spiesecke and Schneider (1961). However, the expansion of the orbital with increasing orbital electrolJl density results in an indirect inverse dependence on q' which in the case of u $ ' ) , shows up as a l / r 3 factor. This indirect dependence is represented by Grant by inclusion of the effective nuclear charge parameter $ of Karplus and Pople (1963). This term enters a factor t 3 , which increases the magnitude of 0;') as [, the effective nuclear charge, increases (i.e., as the electrons are held closer t o the nucleus, thereby affecting the field at the nucleus more'). The resultant dependence of u ; ' ) on q f is 3 '2 illustrated by the solid line in Figure 3 as the sum of the q /2 and -E3qf terms (dashed lines). Thus, the slope of the correlation in

as

It must be emphasized that E also enters into the expression for U d , where it results in increased shielding with increasing effective nuclear charge. However, it only enters as a first-order term, so its effect is much less significant than in the already dominant up term, where it enters as a cubic term.

CHARGE DENSITY-NMR CHEMICAL SHIFT CORRELATIONS

133

the region q' = 1 (formal charge = 0) is seen t o be almost entirely the result of the .t3 factor, which results in increased shielding with increasing electron density. Although some other approaches give the

/

-2.01 0

I

I .o

\ '

2.0

91 Figure 2. Direct dependence of the one atom paramagnetic screening constant, oc), on the orbital electron density, q ' , for carbon-13. Modified from Pugmire et al. (1968).

same result (e.g.,'Emsley, 1968), Tokuhiro and Fraenkel (1969) conclude that the t3 dependence accounts for only about one-half the downfield shift in N-heterocycles (see also page 169). In summary, then, the above treatment predicts that, if one starts with a neutral carbon atom in a molecule, and adds electron density to it, the contribution of ubl) to the overall screening constant should decrease (i.e., become less deshielding) in accordance with the t3 dependence. To the extent that oil) is important in the total screening constant, the chemical shift should move upfield. If, however, electron density is removed from the carbon in question, the contribution of to the screening constant should increase (become more deshielding) up to a point. Then, as further electron

06')

134

D. G. FARNUM

density is removed, 0;') should begin to decrease and, to the extent that ug' ) dominates the total screening constant, shielding should actually increase. This startling suggestion has yet to be examined experimentally, but some evidence bearing upon its validity will be explored in Section 3 on empirical observations. I

1

TWO-ELECTRON TERM

\

\,'

I

*.,

.*-

__----.

I

I'

ONE- ELECTRON TERM

3 Figure 3. Total dependence of the one atom paramagnetic screening constant, 0;). on the orbital electron density, q', for carbon-13. Copied with permission from Pugmire et al. (1968).

062)

The term is not treated as explicitly, but its magnitude seems to be primarily dependent on the bond order between the carbon in question and its bonded atoms. An increase in the bond order increases the magnitude of the negative term u p ) and results in deshielding, while a decrease in bonding decreases and results in shielding.

DY

Higher order terms Grant (Pugmire et al., 1973) has suggested that third and fourth order perturbations on the electron distributions, though significant

CHARGE DENSITY-NMR CHEMICAL SHIFT CORRELATIONS

135

in magnitude, are of opposite sign and probably cancel, so that their net effect is not likely t o refine the total screening constant significantly.

Summary The above interpretation of the factors indicated by current theory to be important in determining proton and carbon-13 chemical shifts does not offer much encouragement for the optimistic statement that “[carbon-13 chemical shifts] should provide a more reliable index of charge than the hydrogen shifts” (Fraenkel and Farnum, 1968, p. 251). It is true that the “other atom” terms are a more serious perturbation on proton than on carbon-13 chemical shifts, and are difficult to evaluate. However it is also apparent that carbon-13 shifts are determined by a number of terms (Td, oil), and o i 2 ) whose charge dependence may well be in opposite directions (and variable), and whose relative magnitudes will be very difficult t o assess. A graphical analysis of u y ) as in Figure 2 and Figure 3 for u;’) would be exceedingly valuable, as would a graphical comparison of ud, o i l ) and ui2) as a function of q. However, this has yet to be done. Furthermore, it has been suggested that the assumption that AE is unchanged introduces an unacceptable error into the estimates of up (Jones et al., 1969). In fact, Pugmire e t al. (1969, 1973) now suggest that u- and n-electron densities, bond orders, and AE are all important in determining carbon-1 3 chemical shifts. Clearly, reliable interpretation of carbon-13 chemical shifts in terms of charge densities will require empirical verification of the above ideas and the discovery of reliable correlations. It seems likely that, in cases where magnetic anisotropy of other atoms is not significant, proton shifts will be useful, while in other cases carbon-13 may be better. It is t o illuminate these problems that we now turn t o a consideration of empirical correlations of proton and carbon-13 chemical shifts with charge densities.

3. EMPIRICAL OBSERVATIONS General That ‘H and I3C chemical shifts correlate linearly with charge density in carbonium ions and carbanions has been an operating

136

D. G. FARNUM

assumption for many of us, even though we acknowledge, in principle, the gross oversimplification in that assumption. Indeed, the Fraenkel, and Spiesecke and Schneider expressions, Ao = -10 Aq for pnmr, and A U = -160 Aq for cnmr, have become a part of our grammar. In Section 2 we explored the theoretical bases for these expressions, and found that, although they seem to be sound for small charge densities, we can expect problems at high charge densities such as those found in ions. In this section we will explore the empirical bases for these correlations, attempt to discover other empirical correlations from data available in the literature, and discuss selected problems. The practice of reporting chemical shifts for both protons and carbon-13 in p.p.m. downfield from TMS will be followed.

Basic Correlations Although many workers recognized the possibility of a linear correlation of chemical shift with charge density, the credit for discovery of the values of the constants and clear expression of the relationships usually goes to Fraenkel e t al. (1960), Spiesecke and Schneider (1961), Lauterbur (1961), and MacLean and Mackor (1961). The familiar plot from the Spiesecke and Schneider work is reproduced in Figure 4. The charge densities6 of tropylium ion, benzene, and cyclopentadienide per carbon atom are plotted against the chemical shifts. N o corrections are made for ring current effects, hybridization changes, or E 2 terms. The slopes of the lines in the plots are close to 10 and 160, and these values, since they are good round numbers, have stuck, even though different workers have found somewhat different values. Since the plots cover a narrow range of charge ( - 0 2 to + 0 14),7 and consist of only three points, it would clearly be very desirable t o extend the range and add additional points before using the correlation with confidence, and particularly before making any gross extrapolations. Some attempts to add more points, as well as improve the fit, will be described below. Charge densities are determined by symmetry, and therefore represent total charge density at the carbon in question (neglecting dispersion of charge to the hydrogens), as required for the theoretical analyses described earlier. In some cases in the literature, calculated n-charge densities are used instead of total charge densities. In the sequel the term “charge density” will refer to total charge density unless otherwise qualified. The point for cyclo-octatetraene dianion was not used in determining the slope because of uncertainty about the magnitude of its ring current and extent of dissociation.

137

CHARGE DENSITY-NMR CHEMICAL SHIFT CORRELATIONS

Figure4. Dependence of proton and carbon-13 chemical shifts on charge density. Copied with permission from Spiesecke and Schneider (1961).

The pnmr correlation In order to explore the significance of the E 2 term in determining Aa, Musher (1962) calculated its effect for the C7H: and C5H; ions and found that the correlated points provided a near perfect fit with a correlation line of slope 11.2. Fraenkel et al. (1960) introduced a correction for the variation of ring current with ring size and found it to be small. In Table 1, the raw data for the pnmr shifts for these ions as well as some other monocyclic aromatic ions are given, along TABLE 1 Chemical Shifts and Corrections for Monocyclic Aromatic Ions

Ion

6H

C3H; C,H? C6H6 CgHG CsH3 CsHg

10.80' 9.18a 7*27a 6.8od 5-44' 5.6ga

Ring currenta correction E 2 correctionb Corrected

+la18 -0.07 0 -1.17 +O*ll -1.16

-1-33 -0.24 0 -0.15 -0.48 -0-75

6 H 6 13C

10.65 8-87

7-27 5.48 5-07 3-77

177e 155f 129f logd 102f 85f

~

a Schaefer and Schneider, 1963 ( C & j in CH3CN, 7.27). b tiom equation (4).

Breslow et al., 1967; Farnum et al., 1967 (TMA in HS03F, 3.13). h t z and Garratt, 1963; LaLancette and Benson, 1963 ( C & j in THF, 7-27). Olah and Mateescu, 1970 (ext. TMS, conv. to CS2= 194.6). f Spiesecke and Schneider, 1961 (CH30C02CH3 ext., converted to C6H6 = 128.7).

138

D. G. FARNUM

with the calculated E Z corrections [from equation (4)], ring current corrections (using a point-dipole approximation; Fraenkel et al. 1960; Schaefer and Schneider, 1963; Katz and Garratt, 1964), and corrected values[i.e., the values the chemical shifts would have if the linear first term of equation (4) were the only factor determining

1

3

.

0

0

1

12.0-

11010.09.0-

5 8.07.060-

5.01

4.0

td.1

3.01 td4 td.3 +d.2 do CHARGE

-6.1 -6.2 -d3 4 4

Figure 5 . Dependence of corrected proton chemical shifts on charge density. See Table 1 for data and references.

them]. The corrected values are plotted against charge density in Figure 5. The effect of the various corrections can be seen by comparing the uncorrected, partially corrected, and fully corrected points. The line is drawn through the three corrected original points. Their chemical shifts are considered to be the most reliable since the ring current corrections and the assumptions of complete dissociation and absence of hybridization effects for C9 H i , CBH i , and C3H< are of questionable validity.' It is gratifying that the additional points fall near the line, but the deviations are well outside experimental error, and do little to improve the reliability of the Although the data for C 3 H l were determined in fluorosulfuric acid, while the other shifts were determined in THF, solvent effects o n the aromatic cations are usually only a few tenths p.p.m in contrast to those o n the anions (Grutzner et al., 1972).

CHARGE DENSITY-NMR CHEMICAL SHIFT CORRELATIONS

139

original correlation.' Furthermore, if the effect of hybridization on C3H:, which is probably more than 1 p.p.m.,l0 is considered, the corrected point for C3H: moves upfield and becomes quite bad. Although the correlation of Figure 5 has proved quite useful (see Jackman and Sternhell, 1969, p. 68 for a number of references) it should be realized that neither the theory nor the empirical correlation has been tested outside the narrow range of charge covered by C,H; (-0.2) and C7H: (+0.14). In fact, the chemical shift of the 2-proton of the 2-propyl cation (dimethylcarbonium ion) (13.5) is not satisfactorily correlated by either equation (4)(which predicts -29) or by extrapolation of the line of Figure 5 (which predicts -17) if the carbon atom is assumed to have a single positive charge, and the absence of the ring current effect is taken into account. There are many factors that could account for the discrepancy; including such things as (1) a reduction of the estimated charge at C-2 by dispersion to the methyl groups [e.g., a charge on C-2 of about +0*4 would fit equation (4),while +0.65 would fit Figure 11; (2) incomplete dissociation of the cation [Bacon and Gillespie have established complex formation between methyl or ethyl fluoride and antimony pentafluoride (1971), and ion-pair complexation is a common occurrence with anions (Grutzner et ul., 1972)] ; (3) erroneous theory; (4)all of these. Clearly much more information is necessary if an assessment of the relative importance of these and other factors is to be obtained.

The cnmr correlation The original Spiesecke and Schneider correlation of carbon-13 chemical shifts with calculated charge densities (Figure 4) has been extended by Olah and Mateescu (1970) as replotted in Figure 6. As noted on page 127, correction of the raw data for ring current effects and E 2 terms is unnecessary for carbon-13 shifts. Therefore one would expect an excellent correlation, provided hybridization effects did not interfere. The calculation of charge density is not as reliable for some of the points on the Olah plot, since for triphenylcyclopropenium cation and tetraphenylcyclobutadiene dication it requires assumption of an approximate molecular orbital model which, in If we wanted to be cavalier, we could claim that the deviations for C9HG and C& result fro? the overcoaection for the ring current effect. If so, then C9H; is exhibiting 50%. and C B H 40% ~ of the expected ring current effect. l o Compare the chemical shift of cyclopropene vinyl hydrogens (7-0, Wiberg and Nist, 1961) with normal alkenes (-6.0).

D. G. FARNUM

140

fact, gives .rr-charge, not total charge, and for tetramethylcyclobutadiene dication it requires the assumption that no charge is dispersed on to the methy1 groups. It is questionable to use the raw value for cyclopropenium ion, since hybridization effects are likely to interfere here (compare cyclopropane, 6' C = -2-6, and cyclo-

220 200 -

I80 I60 -

-no140 -

a

120 -

100 -

80 -

60I-

I

l

l

I

I

I

I

I

I

I

I

+0.6 4 5 +0.4 +0.3+0.2 +0.1 0.0 -01 -0.2 -0.3-0.4

CHARGE

'

Figure 6. Dependence of carbon-13 chemical shifts on charge density. Replotted k o ~ Olah and Mateescu (1970) with permission. The point marked by a cross represents C3H3 corrected for hybridization (see text).

hexane, 6l 3 C = 27.8; Stothers, 1972, p. 270). If the former points are ignored as being unreliable, and the C 3 H i point is corrected by 30 p.p.m. to approximate the hybridization effect, then the data are moderately well correlated by the dashed line of slope 190 p.p.m./unit charge. The solid line in Figure 6 is Olah's original correlation line of slope 160 p.p.m./unit charge-identical with the slope of the original Spiesecke and Schneider correlation.' The dashed line is not suggested as a necessarily better correlation, but does indicate that the proportionality constant varies widely, dependent upon the assumptions one chooses to accept. As has been noted by Spiesecke and Schneider, the scatter of the points is well outside expected experimental error. As in the case of the pnmr

''

Olah h'aa suggested a value of 180 p.p.m./unit charge based on, aa yet, unpublished arguments (Olahet al., 1 9 7 2 ~ ) .

CHARGE DENSITY-NMR CHEMICAL SHIFT CORRELATIONS

141

plot above, the value of 280 p.p.m. obtained by extrapolation of the slope 160 line t o unit positive charge does not correctly predict the carbon-13 shift for the dimethyl- or trimethyl-carbonium ions (320 and 330 respectively), although use of the slope 190 line gives a value of 320. Whichever line is chosen, it must be clearly emphasized 120

I

I

I

I

I

I

110-

100 90 I m

-

-

Figure 7. Correlationof corrected proton chemical shifts with carbon-13 chemical shifts See Table 1.

that the data give no evidence for the upfield shift at high positive charge values suggested by Grant’s theoretical treatment (see page 134). Once again, however, the range of charge (+0.5 to -0-25), though greater than that for the pnmr correlation, may not be great enough. As a final counterpoint to this section on these basic correlations, the “corrected” proton shifts are plotted against the carbon-13 shifts in Figure 7. The excellence of the correlation indicates that errors causing scatter in the original plots have largely cancelled. The slope of 13.4 should be a good indication of the ratio of the charge dependence of carbon-13 chemical shifts to the first order charge dependence of proton chemical shifts over the range covered by the data.’

**

A similar plot of proton shifts us. carbon shifts for para-substituted benzenes, where charges differences are much smaller, gives a line of slope 17-2 {Stothers. 1972, p. 199).

142

D. G . FARNUM

Correlation of methyl group chemical shifts The first attempt to correlate proton chemical shifts of methyl groups in carbonium ions with charge on the adjacent carbon atom seems to have been made by MacLean and Mackor (1961). They examined the pnmr spectra of a number of cyclohexadienyl cations (benzenium ions) [6] where R was CH3, and found that chemical shifts for protons directly attached to positive carbons were linearly

[GI

related to those of methyl groups attached in the same position. The correlation slope was about 2-8. Using approximate calculated charge densities, they determined a proportionality constant of 13.4 for the charge dependence for the chemical shift of protons attached to positive carbon, and a value of 4.7 for that of methyls. These values can be considered as only very approximate, however, since they use a very approximate model for the charge density calculation, and are not corrected for either the E 2 term or the influence of other charged atoms suggested by Musher. Indeed, Musher (1962) reanalyzed MacLean and Mackor’s data including these corrections, and found a value 11-0 for the proton correlation, in excellent agreement with his corrected empirical value of 11-2. In addition, he found the corrected methyl shifts to be in agreement with theory. Using Musher’s equations, the empirical value of 11.2 for the linear part of the proton dependence, and reasonable values for bond lengths and angles, one can calculate the expected charge dependence of the proton shifts of methyl groups attached to carbonium ion centres. Relevant distances and angles are given in Figure 8. Thus, using equation (3):

Substituting the appropriate values from Figure 8, and ignoring all contributions from remote charges (the error thus introduced is rather small) gives equation ( 5 ) :

A0 = - 2-33 q

- 0.94 q 2

(5)

CHARGE DENSITY-NMR CHEMICAL SHIFT CORRELATIONS

143

Figure 8. Parameters used to calculate the charge dependence of the proton chemical shift for methyl groups o n carbonium ion centres.

The value 2.33 for the linear dependence is rather far from the uncorrected value of 4.7 from MacLean and Mackor's data. However, it can be shown to be in better agreement with the more reliable empirical correlations following. Data now available in the literature permit further attempts to determine the charge dependence of the chemical shifts of methyl groups attached to carbonium ion centres. Thus, a SpieseckeSchneider type plot for a few systems is given in Figure 9 and the relevant data are tabulated in Table 2. The slope of 3-2 is closer to theory than that from MacLean and Mackor's data. However, the calculation of charge still assumes no charge dispersion on to the methyls, which must be wrong. Also, the number of points is minimal, and, even though the fit is excellent, the correlation cannot be taken too seriously. Once again, the dimethylcarbonium ion shows less than the expected downfield shift-in this case exhibiting about 80% of the expected value. An alternative approach to determination of the charge dependence of the methyl proton shifts is illustrated in Figure 10

CHARGE Figure 9. Charge dependence of proton chemical shifts of methyl groups on ions. See Table 2 for data and references. o Raw data; A corrected for ring current; 0 corrected for ring current and E2.

144

D. G. FARNUM

TABLE 2 Chemical Shifts and Corrections for Some Methyl-Substituted Ions and Molecules

Compound (CH3)IdH

Ring current Charge 6H(CH3) correction

E Z correction

+1-0

5-04a

0

-0.94

4.10

+0.33

2-95'

-0.2

-0.1 1

2-64

2-23d

-0.6

Corrected

6H

CH3

0

1.63

0

CH3 CH3 a

Olah and White, 1969 (ext. TMS in SbFS-SOz). Olah e t al., 1969 (ext. TMS in SbFJ-SOz). Closs et al., 1968 (TMS in CH3CN). Emanuel and Randall, 1969 (TMS in CDCI,).

-

Jz

3.6-

A

3.4-

I

I

I

I

1

I

I

210 220 230 240 250 260 270 280 Figure 10. Correlation of proton chemical shifts of methyl groups in carbonium ions with carbon-13 shifts of the adjacent trigonal carbon atoms. See Table 3 for data and references. o Aryldimethylcarbonium ions; 0 cycloalkenyl cations; A phenylmethylcarbonium ions.

145

CHARGE DENSITY-NMR CHEMICAL SHIFT CORRELATIONS

(data used for Figure 10 and 11 are in Table 3). Here, the methyl proton shifts for a number of closely related carbonium ions are plotted against the carbon-13 shifts of the carbonium ion centres. If both are linearly related to charge, then the slope of the resultant straight line should give the ratio of the dependences. Two systems TABLE 3 Proton and Carbon-13 Chemical Shifts for Methyl Groups Attached to Carbonium Ion Centres Compound

Substituents

X-C6H46(CH3)2a

X

= p-OCH3 X = p-CH3 X = P-F X=H X = m-F X = p-CH3

X-C6H46HCH3

x =p X X X X

- 0 ~ ~ ~ 6

= 2, 4, 6(CH3)3 = p-CH3

=H = P-CF3

R1 = H, R2 = CH3C R , = R2 = CH3C

6HCH3

3.12 3.45 3-38 3.48 3.61 3-71 2.98 3.02 3.6 1 3.73 3-93

220 244 249 256 262 270 200 213 223 235 248

3.36 3.14

247 229

43 34

3.61

262

49

3-45

255

47

R2

Rl

o-,,, a

C

Data from Olah et al., 1972a (ext. TMS, 13C conv. to CS2 = 1946). Data from Olha et al., 1971 (ext. TMS, I3C conv. to CS2 = 1946). Data from Olah and Liang, 1972 (ext. TMS, 13C conv. to CS2 = 1946).

studied by Olah (Olah and Liang, 1972; and Olah e t al., 1972a) were chosen-the methyl substituted cycloalkenyl cations, represented by solid dots, and the substituted aryldimethylcarbonium ions, represented by open circles. The correlation to a line of slope 0-0146 is excellent for all but two of the points. If the Spiesecke-Schneider value of 160 p.p.m. is used for the linear charge dependence of the carbon-13 shifts, then the calculated value for the methyl shifts is

146

D. G. FARNUM

(0.0146) x (160) % 2-3, in excellent agreement with the theoretical vaIue of 2-33. The agreement, of course, depends on the value chosen for the carbon-13 dependence, but even an outside value of 200 still gives 2-9, which is not far from theory. It seems likely that the original proportionality constant of 4-8 suggested by MacLean and

I” 0 50 c 0

0 n 40

60

30 200

220

8 l3c

240 at

260

C+

280

Figure 11. Correlation of carbon-13 chemical shifts of methyl groups in ca~- mium ions with carbon-13 shifts of the adjacent trigonal carbon atoms. See Table 3 for data and references.

Mackor is too large and the actual value may be closer to 3 or less. Attenuation of the effect of the positive charge by a factor near 3 upon insertion of a carbon atom seems reasonable. The correlation should not be taken too seriously, however, since the values for the phenylmethylcarbonium ions (Olah et al., 1971), represented by triangles in Figure 9, which should give a displaced but nearly parallel line, give a rather poor correlation. A more thorough study is needed. A similar approach could be used to calculate the charge dependence of the carbon-13 shifts for methyl groups attached to carbonium ion centres, but only limited data are available. In Figure 11 the carbon-13 shifts for methyl groups in cycloalkenyl cations are plotted against the shifts of the attached carbonium ion carbons as determined by Olah and Liang (1972). The correlation line for the limited number of points has a slope near 0-4, suggesting a value of (0.4) x (160) % 60 for the charge dependence of the carbon-13 shift for methyls attached to carbonium ion centres. Again the attenuation factor of 2.5 seems reasonable. It would be very valuable to have a reliable assessment of both the linearity and

147

CHARGE DENSITY-NMR CHEMICAL SHIFT CORRELATIONS

magnitude of this correlation. It seems quite possible that carbon-13 shifts of methyl groups attached to carbonium ion centres will be quite sensitive to charge, while relatively insensitive to other factors such as diamagnetic anisotropy, hybridization, and bond order which render both proton and carbonium carbon-13 shifts ambiguous.

Other correlations and conclusions Although a number of other attempts have been made to correlate carbon-13 and proton chemical shifts, they do not extend either the range or precision of those already described, and are mentioned in

300 -

-

Figure 12. Correlation of carbon-I3 and proton chemical shifts for trigonal carbon atoms in a number of cations, anions,and neutral molecules.

other reviews quoted above. In order to give an overview of the current possibilities, however, a scatter plot of proton chemical shifts us. carbon-13 chemical shifts for trigonal carbon atoms in a number of cations, anions, and neutral molecules is given in Figure 12. No attempt has been made to correct the raw data for the host of factors which might interfere, although protons and carbons next to heteroatoms have been excluded, since . they are well known to be anomalous, as will be discussed later. If both proton and carbon-13 shifts were precisely proportional to charge with proportionality constants of 10 and 160 respectively, all points should fall on the line of slope 16 drawn through the data. Simplistic as the approach

148

D. G. FARNUM

is, nonetheless the graph dramatizes the following points: (1) There is a general correlation whose slope vanes from about 10 to about 30; (2) the scatter of the points is far too great to allow interpretation of the data in a given case without a careful assessment of the factors which might render the interpretation ambiguous; (3) there is no evidence for a reversal of slope at high positive charge density as would be expected if the up term dominated the carbon-13 shifts throughout and behaved according to theory; (4) many more points are needed in the carbanion region (below 6H-7) and for highly charged carbonium ions (the region above 6H-10). Surely refinement and extension of all of the correlations described above would be valuable. Systematic studies are desperately needed in which spectra of compounds chosen to minimize changes in extraneous variables are determined under carefully controlled conditions of solvent, concentration and temperature.

Selected Problems In this section we shall examine some problems in which cnmr and pnmr spectroscopy have been used to try to answer questions about charge distribution. The selection is, of course, biased, but an attempt has been made to choose examples which one might have expected to be simple, but which turn out, on further analysis, to be quite puzzling.

The triphenylmethyl cation and anion Surely one would expect a system as hoary with age as the triphenylmethyl cation to be well understood by now! Yet despite the effort of many groups to use pnmr and cnmr techniques to determine the detailed charge distribution in both the cation and anion, theory and experimeni are still far from agreement. Perhaps the lack of congruence simply illustrates the principle that the more we know, the more we want to know. Whatever the reason, the fact is a sobering one. The difficulty in determining charge distribution in the triphenylcarbonium ion by pnmr has been discussed in an earlier review (Fraenkel and Farnum, 1968), and only the salient conclusions will be noted here. In short, there are too many approximate models with undetermined variables that intervene between the raw data (the proton shifts) and the rigorous theoretical determination of charge distribution. One must assume an approxi-

CHARGE DENSITY-NMR CHEMICAL SHIFT CORRELATIONS

149

mate model to describe the magnetic anisotropy of the benzene ring, an approximate model for the effect of remote charges on proton shifts, and a more or less arbitrary value for the twist angle of the propeller-shaped ion before one can calculate the “experimental” charge density for comparison with the “theoretical” charge density determined by some approximate theoretical model at the chosen twist angle. Needless to say, the results are not very satisfying, although getting to them provides an interesting exercise in equation juggling and computer programming. TABLE 4 Chemical Shifts for the Triphenylmethyl Cation and Anion _ _ _ _ _ _ _ _ _ _

Position

~

1(4

2

Cation 1Ha I3cb

7-68 2 10.9

Anion

139.9

143.3

7.27

1HC

13cd

3(0)

4(m) 7-86 130.3

6.49

5(P)

8-24 143.1 5.93

92

Farnum, 1967 (TMA in HS03F, 3.13). Ray et. al., 1971 (ext. CS2 in HS03F, 192-8). Sandel and Freedman, 1963 (TMS in THF). Waack et. al., 1966 (ext. CS2 in THF, 192.8).

The pnmr shifts for the anion have also been determined (Sandel and Freedman, 1963; Grutzner et al., 1972) and are compared with those for the cation in Table 4. Although Grutzner e t al. clearly demonstrated a large dependence of the chemical shifts of a number of anions on solvent, counter-ion, and temperature (in contrast to many cations) they also demonstrated that lithium triphenylmethide was reasonably well behaved, particularly in oxygenated solvents. Hence, the shifts reported for the lithium salt in THF are most probably those for the “free” ion. Sandel and Freedman noted that the pnmr spectra of the cation and anion, reproduced in Figure 13, did not show the mirror-image relationship “expected on theoretical grounds.” However, the mirror-image relationship is expected only if charge density at the position in question is linearly related to the proton-shift and is the only factor which determines that shift, sicce it is the charge distribution which theory predicts should have a

150

D. G. FARNUM

Figure 13. Pnmr Spectra for the triphenylmethyl cation and anion. Copied with perkssion horn Farnum (1967), and Sandel and Freedman (1963).

mirror-image relationship. From the earlier comments on the importance of ring-current effects and E 2 terms in determining the observed proton shifts it should be apparent that a mirror-image

Figure 14. Schematic presentation of correction of the proton chemical shifts for the triphenylmethyl cation and anion.

CHARGE DENSITY-NMR CHEMICAL SHIFT CORRELATIONS

151

relationship in the spectra would have been a cause for great concern

over the validity of the theory. Figure 14 is a schematic presentation of an attempt to correct for some of the more important sources of error, largely as an exercise in data juggling to illustrate the ambiguity of interpretation. The chemical shifts determined for the ortho-, metu-, and para-protons for the cation and anion are corrected for the ring current effect of the two other benzene rings assuming a twist angle of 45” about the bond between the phenyl groups and the trigonal carbon in a symmetrical, propeller-shaped cation and anion (Farnum, 1967). These corrected values are then corrected further for the large E 2 term felt by the ortho-protons

Figure 15. Parameters used to calculate E2 corrections for Figure 14.

from the proximate rather high charge on the central carbon. The bond lengths and angles, and distance parameters used for these calculations are shown in Figure 15. The E 2 correction is determined for three different charges ( 4 ) at C-1 using the second term of equation (3), since the charge at C-1 is not known with certainty. It can be seen that for q = +035 the chemical shifts show a nearly mirror image relationship with the anion showing a wider spread. Of course, other corrections have been ignored, and it is not necessary to assume that the cation and anion have the same twist angle or charge at C-1. With so many variables to choose from, it seems likely that the data could be “corrected” to fit almost any theory one would like to choose. The ambiguity of the interpretation of the pnmr spectrum motivated Ray et ul. (1971 and earlier papers) to determine and analyse the cnmr spectrum of the triphenylcarbonium ion (as well as some of its derivatives). The chemical shifts they obtained are given in Table 4, and a comparison of the “empirical” charge densities with those calculated by the HMO and CNDO techniques

152

D. G. FARNUM

TABLE 5 Comparison of Empirical Charge Densities with those Calculated by the HMO and CNDO Methods for Triphenylcarbonium Ion (twist angle = 30')' Position

HMO

1 (Central) 2 3 4 5

0.41 7 0.000 0.065 0~000 0.065 1.00

Total q, a

CNDO

0-249

-0.0 13

0.071 0-015 0.08 1 0.969

Empirical

0.267 -0.048 0-094

0.020 0-107 1.13

Data from Ray e t al. (1971).

for a twist angle of 30° is presented in Table 5. The empirical charge density changes, Aqi, for all carbon atoms except C-1 were obtained using the relationship

Aqi = -ASi/160 where ASi is the chemical shift difference between triphenylcarbinol and the cation for C-i. For the central carbonium ion carbon atom C-1, the charge density q 1 was calculated assuming that the charge on the central carbon atom of trimethylcarbonium ion was +1, and that the charge difference between it and the central carbon atom of triphenylcarbonium ion was proportional t o the chemical shift difference (330-6 ) with the usual constant of 1/160. Thus:

q1 = 1 - (33O-61)/160

It is not at all certain that q 1 can be determined in this way (see page 140), but if it is, the total charge on the cation summed over all positions, q, = +1*13. This is in better agreement with the expected value of +1-0 than the value of +1-48 obtained if q 1 is determined from Figure 6 to be +O-52. The agreement of the empirical charge with the CNDO charge is quite good, at least relatively, as seen by the correlation line of Figure 16 (slope = 1.14, correlation coefficient = 0-985).' Of special interest is the prediction of negative charge density at C-2, which seems to be borne out by the empirical charge densities. This conclusion must be taken with caution, however, since the central carbon atom, attached t o C-2, is undergoing a change l 3 Indeed, if the proportionality constant is taken to be 180, as suggested by Olah et aL (1972), instead of 160, the slope of the line of Figure 16= 1.01.

CHARGE DENSITY-NMR CHEMICAL SHIFT CORRELATIONS

153

from tetrahedral to trigonal on conversion from the carbinol to the carbonium ion. Therefore the carbinol may not be a good model for the chemical shift at C-2.In fact, such a structural difference can be accompanied by an upfield shift of 6-8 p.p.m. For example, the alkyl-substituted carbon of the aromatic ring of cyclohexylbenzene

3

Figure 16. Plot of empirical charge densities us. CNDO charge densities (0 = 30") for triphenylcarbonium ion. Copied with permission from Ray et al. (1971).

resonates at 148, while C-1 in biphenyl resonates at 142; the corresponding carbon resonances of isopropylbenzene and (Ymethylstyrene appear at 149 and 141 p.p.m. respectively (Stothers, 1972, p. 90, 97; Dhami and Stothers, 1965).Correction for that upfield shift accounts for most of the 9 p.p.m. upfield shift observed at C-2 between the carbinol and carbonium ion. Indeed, the shift at C-2 in the carbonium ion (140)is within experimental error of the shift of the similarly constituted atom in amethylstyrene, suggesting it may have n o charge at all. Thus, although the work of Ray et al. (1971)clearly is the best analysis available of the charge distribution in the triphenylcarbonium ion as determined by NMR spectroscopy, it is not free from the ambiguity which plagues attempts t o draw accurate quantitative conclusions using models of uncertain accuracy. Nonetheless, one would hope that some group would find the time and motivation to complete an equally thorough, careful investigation of the triphenylcarbanion for comparison. l4

Work in progress along these lines by Grutzner (1973) should appear soon.

154

n r.. F A R N T I M

The inductive effect of methyl groups The downfield shift of the central carbon resonance in trimethylcarbonium ion (330 p.p.m.) as compared with dimethylcarbonium TABLE 6 Comparison of Chemical Shifts of Methyl Substituted and Unsubstituted Carbonium Ions Case

Rl

1.

2.

R2

Position

R1 R 2 d H

R l R26CH3

A6

320 62 5-04

330 48 4.35

-10

254 35 3-34

281 32,41 2.70, 3-18

-27 +3, -6 +-64, +.16

+14(+13f) + -69

3.

C6H5

235c 3.73c 8-7c

256d 3*48e 8-63e

-21 +-25 + -07

4.

C6H5

227 150

-20

C6H5

200 152 8.38

247 146 230 149 8-32

200c

220d

5.

6.

2.9gc 223c 3.6 lC 234c 248c 3.93c

7. 8. 9.

3-12d 244d 3.45d 24gd 27od 3.7 Id

+4 -30 +3 +so6 -20 -*14 -21 +-16 -15 -2 2 +.22

a Olah

and White, 1969. Olah et al., 1970. (ext. TMS, "C conv. to CSz = 194.6). Olah et al., 1971. Olah et al., 1972. Farnum, 1967 (TMA in HSOJF= 3-13). f Olah and Westerman, 1973 (ext. TMS, conv. to CS2 = 1946).

ion (320 p.p.m.) has been taken as evidence in support of extended Huckel theory calculations which place more positive charge on the central carbon of the tertiary ion than on that of the secondary ion (Olah and White, 1969). Thus, one might conclude that the methyl

CHARGE DENSITY-NMR CHEMICAL SHIFT CORRELATIONS

155

TABLE 7 Chemical Shifts of Some Methyl Substituted and Unsubstituted Allylic Cations4 Case

Position

1.

Unsubstituted cation

0

1

.-.I

3

2

21gb 138’ 8-32

G1, C-3 c-2 H- 2

Methyl substituted cation

0

A6

2

229 139 8.04

-10 -1 +028

2. 1 2

c-1,c - 3 c-2

236’ 147’

249 157

-13 -10

c-1, c-5 c-2, c - 4 c-3 H-2, H-4

182 140 202 8.49

197 136 197 7-70

-15 +4 +5

3.

+ o 79

4.

3

c-1, c-5 (2-2, c - 4 c-3 H-2, H-4

172 139 161 8-20

175 139 186 7-66

-3 0 -25 +054

Gl, C-5 (2-2, c-4 c-3 H-2, H-4

175 139 186 8-05

186 137 176 7.66

-11 +2 +10 +039

5.

156

D. G. FARNUM

TABLE 7-continued -

Case

Position

Unsubstituted cation

0

6.

1

.-4

2

c -2 c-3 H- 2 H- 3 7.

n3 2

c-1 c-3 CH3 cH3

A6

2

24 7 141 206

-28 -3 +13 +0m24 +052

8.08

9.73

/o, 2

24 7 141 206 43 3.36

c-2

~~

3

219‘ 138’ 219’ 8.32 10-25

G1

~

Methvl substituted cation

+18 +2 -23 +10 +022

229 139 229 33 3-14

8. l

a

3

J

32 3

2

c-1 c-2 c-3 H- 2 H-3 9.

262 148 219 8-36 10.60

-26

262 148 219 3.61

249 157 249 2.8 7

+13 -9 -3.0 + u 74

2

2

207 152 207 8-34 10.18

255 141 21 1 8.05 9-58

236’ 147’ 236’ 8.65 11-26

-1 +17 +0.29

m3 q

+0*66

2

c -1 c-2 c-3

W a -1 10.

0, o3

1

c -1 c -2 c-3 H- 2 H- 3

.-4

-48 +11 -4 +-29 +0m60

“Data from Olah c t af., 1972b, Olah and Liang, 1972, and Olah and Porter, 1971 (cxt. TMS. I3C conv. to CS2 = 194-6). The data for the cyclohcxenyl and cyclopcntcnyl cations were transposed in the original article.

CHARGE DENSITY-NMR CHEMICAL SHIFT CORRELATIONS

157

group is electron-withdrawing by comparison with hydrogen. The fact that that conclusion does violence t o our intuition is not sufficient reason for rejection, especially since our intuition is based on indirect evidence from rate and equilibrium comparisons rather than knowledge of actual charge distribution. Nonetheless, it is sufficient reason t o ask that the data be compelling. The data collected in Tables 6 and 7 show unequivocally that the usual effect of a methyl group on the carbon-13 shift of a carbonium ion centre is t o deshield it and cause a downfield shift in the neighbourhood of 20 p.p.m. (bold face entries in last column). The effect is often larger than the ca. 9 p.p.m. deshielding observed upon substitution of methyl for hydrogen in alkanes, and very much larger than that observed upon substitution of methyl for hydrogen at neutral trigonal centres. For example, ethylene, 2-butene, and 2,3dimethyl-2-butene all show resonance at 123 k 1 p.p.m., while the aromatic carbons of benzene and hexamethylbenzene appear at 129 and 133 p.p.m. respectively (Stothers, 1972, Chapter 3). The effect is not restricted to cations, as is shown by a comparison of diphenylmethyllithium ( a - C = 7 9 p.p.m., Waack et a/., 1966) and 1,ldiphenylhexyl lithium (a-C= 83 p.p.m., McKeever et al., 1971). Nevertheless, t o conclude that the downfield shift reflects an increase in positive charge at the carbonium ion centre results in some serious inconsistencies in the data. For example, the carbon-13 shift of a methyl group attached to a carbonium ion centre moves upfield as much as 10 p.p.m. when the carbonium ion is substituted further by a methyl group (Table 6, cases 1 and 2, Table 7, case 7). The proton shift of such methyl groups similarly moves upfield (Table 6, cases 1, 2, 3, 6, 7, 9; Table 7, cases 7 and 9) with only one exception (Table 6, case 6). Both of these results indicate that the positive charge at the carbon to which the methyl group is attached has become less upon further methyl substitution, a conclusion difficult to reconcile with an electron withdrawing methyl group. The upfield shift of both the para-carbon (Table 6, cases 4, 5) and para-hydrogen (Table 6, cases 3, 5) in phenylcarbonium ions upon further substitution of methyl for hydrogen at the carbonium ion centre is also indicative of less positive charge. Finally, the upfield shift of the remaining uncharged carbon atoms and vinyl protons in most of the cases of Table 7 suggests that positive charge is being dispersed by the methyl groups leaving less charge on the allylic system. Since our earlier analysis concluded that there was neither theoretical nor empirical justification for a linear relationship of

158

D. G. FARNUM

positive slope between charge density and carbon-13 chemical shift at high positive charge density, it would seem very premature to interpret the downfield shift which accompanies methyl substitution on carbonium ions in terms of increased positive charge density, particularly in the face of the considerable evidence to the contrary.

Charge delocalization b y phenyl and cyclopropyl Olah and Westerman (1973) have determined the carbon-13 spectra of a number of phenyl and cyclopropyl carbonium ions in order to assess the reIative abilities of phenyl, cyclopropyl, and methyl to delocalize charge. In Table 8 their results are presented in such a way as to simplify the assessment of their conclusion that TABLE 8

Comparison of Chemical Shifts of Methyl, Cyclopropyl, and Phenyl Carbonium Ionsa

1.

2. 3. 4. 5. 6. 7.

8. 9.

320 62 5.04 330 48 4-35 281 32,41 254 276 235 8.7 255 36 157 8.63 247 146 230 149 8.32

254 35 3.34 28 1 41 3-18,2*70 276 39 255 272 227 150 247 24 146 262 137 236 146,151

235 3-73 255 36 3-57 247 24 226 262 200 152 8.38 230 32 149 8-32 236 151,146 21 1 143 8.24

Data from Olah and Westerman, 1973 (ext. TMS, I3C conv. to CS2 = 194.6).

CHARGE DENSITY-NMR CHEMICAL SHIFT CORRELATIONS

159

phenyl is better than cyclopropyl, which, in turn is better than methyl at delocalizing positive charge. If we accept the thesis that, for these closely related ions, the carbonium centre carbon-13 shift will parallel the positive charge at that centre, then the carbon-13 resonance should move progressively upfield from left to right within each group of three cations in the table, if Olah and Westerman are right. With only few exceptions (cases 4, 8, and 9), this is so. Even the exceptions interchange the cyclopropyl and methyl cations, but leave the phenyl substituted cation at highest field. The chemical shifts of the methyl carbons also are in general accord with the order with a few exceptions, while the protons show more discrepancies.

Figure 1 7. Possible geometry for the cyclopropyldiphenylcarboniurnion.

The proton shifts, of course, will be more affected by neighbounnggroup anisotropies. One might be concerned over the order of the para-carbon chemical shifts in cases 6, 7, and 8, which suggest that there is less positive charge at the para-position when the benzylic cation is substituted by cyclopropyl than by phenyl. However, the charge at the para-position will parallel the charge at the alphaposition only if all the ions have the same degree of coplanarity. Because of the different size of methyl, cyclopropyl, and phenyl, this may well not be the case (cfi Farnum, 1967). The cyclopropyldiphenylcarbonium ion of case 8 provides an interesting illustration of this paint. The phenyl groups in this ion are not equivalent, as evidenced by the quite different ring carbon resonances. For one ring, the para-carbon resonates at 151 p.p.m., for the other at 146 p.p.m. The difference of 5 p.p.m. is much too large to be accounted for by a long-range anisotropic effect of the cyclopropane ring. Furthermore, the difference is similar at the ortho-positions (4-3 p.p.m.) and much less at the meta-positions (1.5 p.p.m.) which seems quite inconsistent with an explanation based on magnetic anisotropy. The phenomenon finds a ready explanation, however, in different degrees of coplanarity of the two phenyl groups as illustrated in Figure 17. Thus, the more hindered, more

160

D. G. FARNUM

twisted phenyl syn to the cyclopropyl group would not delocalize charge as effectively as the more nearly planar anti-phenyl. The difference in charge delocalization on to the phenyls would be manifest in a difference in the chemical shifts of the ring carbons, which should be greatest at the ortho- and para-positions, as observed. It seems likely that a similar distortion from coplanarity is involved in the anomalous para-carbon resonances of c a s ~ s6 , 7, and 8. Our confidence, then, in the conclusion that phenyl is more effective at delocalizing positive charge than 'cyclopropyl parallels our confidence in the charge density chemical shift correlation for these compounds. Caution is again warranted, particularly since we have already seen that conversion of an adjacent carbon atom from tetrahedral to trigonal can cause a significant upfield shift (page 153). Just such a change takes place upon replacement of cyclopropyl by phenyl. The effect of heteroatoms Substitution of heteroatoms on carbonium ion centres often has a profound and unexpected effect on the carbon-13 chemical shift which is hardly designed to bolster our confidence in the simplicity of the charge density-chemical shift correlation. A few examples will illustrate the point (in the following structures carbon-13 shifts are given in regular type, proton shifts in italics). In example (a) (Stothers and Lauterbur, 1964) the carbonyl carbon-13 resonance moves upfield when placed next to an electronCH3. CO .CH3

t

t

198

205

CH3COi 4

I

182

.

CH3. CO .CO CH3

CHjC02H

+

2.101 177

CH3CO;Hz 4

3.18

1

196

(4

CHARGE DENSITY-NMR CHEMICAL SHIFT CORRELATIONS

7.95

7-69

161

9.53

withdrawing carbonyl group. In example (b) (Gray et al., 1969; Olah and White, 1969) the carboxyl carbon resonances of both acetate anion and the protonated cation are downfield from that of acetic acid. In (c) (Pugmire and Grant, 1968a) the alpha-carbon of pyridine moves upfield on N-protonation, while the other carbons move downfield as expected. In (d) (Pugmire and Grant, 1968b) all carbons of imidazole move upfield with increasing positive charge on the system as one goes through the series anion, neutral molecule, cation. Proton shifts, where reported, are much more in accord with expectation. Much effort has been expended, particularly by Grant, to unravel the complex interplay of factors which produce these phenomena. Data abound, many collected by Olah. Yet our understanding of the data still leaves much to be desired. In this section some data available on the oxygen substituted systems will be presented, followed by a summary of the nitrogen heterocycles. Oxygen substitution Olah has determined the pnmr and cnmr spectra of a number of protonated carbonyl compounds in fluorosulphuric acid-antimony pentafluoride-sulphur dioxide solution (Olah and White, 1968; Olah et al. 1972). The carbon-13 data are given in Tables 9 and 10. The proton data are not included in Table 9, since the anisotropy of the oxygen function obscures any interpretation. A correlation of some

TABLE 9

Carbon-13 Chemical Shifts of, Hydroxycarbonium Ions (in FS03H-SbF5-SO2 Solution at -78 ) and Their Uncharged Precursorsu (CH3)3C13 + (CH C ~ ~ O H + CH,3d?3(OH)f 0 3 (OH); CH3OC l 3 (OH); HC13 (OH) f a

341 250 193 166 164 178

(CH3)2C13=CH2 (CH3)2C13=0 CH3C1302CH3 (CH3O)2Cl3O (CH30)2C130 HC1302CH3

Data from Olah and White, 1968 (ext. TMS,conv. to CS2 = 194.6).

140 207 172 159 159 162

162

D. G. FARNUM

W

0.50.4 O3

3;O

I

I

I

I

300

250

200

150

~TMS

Figure 18. Correlation of carbon-13 chemical shifts with charge for some oxygcnated cations. Copied from Olah and White (1968). with permission.

of the data with a-electron charge density as calculated by a simple HuckeI treatment was done, and this is reproduced in Figure 18, The correlation, although quite satisfactory, is subject to a number of qualifications: more points are needed; the charge on the central carbon atom of the t-butyl cation was assumed to be unity; the mathematical model used to calculate charge on the other systems is a very approximate one, and the adjustable parameters are somewhat arbitrarily chosen. The slope of 306 p.p.m. per unit charge may, therefore, be quite misleading. It is certainly much larger than the approximately 160 p.p.m. value found for most other systems, and may reflect an unusual sensitivity of the carbon-13 resonance of oxygenated cations to positive charge density. Olah’s data for protonated a,P-unsaturated ketones are given in Table 10. In the abstract of the article it is concluded that the generaIIy Iarger downfield shift of the p-carbon resonance (C-3) than the oxygenated carbon (C-1) upon protonation indicates more positive charge at the 0-carbon. However, that conclusion is modified in the text of the article because of uncertainties introduced by the oxygen function. If, indeed, the charge dependence of the carbon-13 chemical shift at an oxygenated carbon is 300 p.p.m./unit charge as determined above, then one could conclude at least that there was more buzld-up of charge at C-3 than C-1, since it would then be true that the calculated change in charge at C-3 (Aq3 in the last

TABLE 10

I onb

6C1

6a

214 (-19) 214 (-17) 208 (-15) 207 (-10) 228 (-29) 226 (-27) 222 (-23) 224 (-24) 218 (-20) 233 (-23) 226 (-27) 214 (-15)

6b 6c

6d 7a 7b 7c 7de 7ee 8

9a 9b

6C2

6C3 179 (-41) 176 (-39) 203 (-49) 200 (-48) 160 (-30) 161 (-21) 183 (-38) 188 (-48) 184 (-29) 184 (-39) 193 (-42) 185 (-60)

Chemical Shifts for Protonated

CH3

(Y,

D-Unsaturated Aldehydes and Ketonesac

HIR'l CH3 H I

7.60 (-1.20) 2-48 (-0.68)

R

2

2-40 (-0.70) 3.40 (-1.17) 3-40 (- 1~ 0 7 ) 3.24 (-1.10) 3 28 (-1.10) 314 (-1.13)

7.40 (-1.52) 2.50 (-0.59)

7.20 (-1.20) 7.42 (- 1.22) 7.30 (- 1*40)

7.16 (-1.28)

3.00 (-0.90)

2.80 (-0.90)

2.38 (-0.58)

3

8.50 (-2-1 0) 812 (-1-7 7) 9.30 (-2.30)

8.90 (-1.22) 7.74 (-1.54) 8-56 (-2.52)

7-24 (-1.24)

R

H

8.30 (-1.90) 8-00 (-1.90)

7-80 (- 1*60)

r

1CH3

891 (-2.11)

2 84 (- 1.08)

1

CH3

H'

Aq jd 023

2.88 (-0.90)

9.70 (-0.20) 9.60 (-0-10) 9.50 (0) 9.35 (+0*15)

8.11 (-1.87) 8.89 (-2.76) 7.98 (-1-20)

0.22

0.27 0.27

'83

0.17

6

0.12

ii15

0.21

2.73 (-0.82) 3 00 (-0.99)

$

0-16

n

g.

8

9 60 (-1 -96)

0.22

yE

9 00 (- 1*96)

0.23

3

290 (-0.8 7)

0-33

Data from Olah et al., 197% (ext. TMS,I3C conv. to CS2 = 194.6). For structures see text. Proton shifts are italicized. Numbers in parentheses are shift differences from unyrotonated parents. A negative sign means a downfield shift on protonation. Determined using 180 p.p.m./unit charge. The proton data seem to be transposed in the original.

c1 Q)

w

164

D. G. FARNUM

I

R3

PI R , = R ~ = H

f

0

+\H

PI ~ , R ~ = R ~ = R ’ = H b, R l = H; Rz = R’ = CH3

column of Table 10) would be greater than that at C- 1 in every case. However, that conclusion is not consistent with some of the proton data. Thus, protonation t o give [7c], [7d], and [7e] results in a downfield shift of the a-methylproton resonance of 1.10, 1.10, and 1.13 p.p.m. respectively, while the o-methyl resonances move downfield by only 0 9 0 , 0.82 and 1-08and 0.99. These shifts suggest more buildup of positive charge at C-1 than C-3. In any event it does not seem likely that the magnitude of the total charge at C-1 is less than that at (2-3, since the chemical shift at C-1 is displaced some 30 p.p.m. to lower field in most cases. One can conclude with confidence that C-2 is much less bothered by the whole business, but beyond that, the data are really not too revealing about the charge distribution in these ions. As an incidental point, ions [gal and [9b] provide one more exception to the general rule that methyl substitution results in a downfield shift of the carbonium carbon-13 resonance.

Nitrogen substitution The dramatic effect of nitrogen substitution on carbon-13 chemical shifts was illustrated at the beginning of this section by the fact that the alpha-carbon resonance of pyridine moves upfield upon protonation at nitrogen. A further demonstration that nitrogen substitution does still more extensive violence to charge densitychemical shift correlations is seen in Figures 19 and 20. Thus, in

CHARGE DENSITY-NMR CHEMICAL SHIFT CORRELATIONS

160 -

165

0 0 0

O 0 0 0 0

e

I50-

o

0

0

0 0

0

0

a p?

0

0

0

140 -

0

e

b 0

0 .

0

0

. I

0

b

b

120

0

o

'

'

In)

8.0

' ' '

'

9.0

' ' ' '

I

' '

10.0

6H Figure 19. Correlation of proton and carbon-I3 chemical shifts for a number of heterocycles and their anions and cations. o Positions alpha to N; other positions.

Figure 19 the carbon-13 chemical shifts for a number of nitrogen heterocycles and their anions and cations are plotted against the proton shifts at the same positions. The scatter is excessive, and is not restricted to positions alpha to nitrogen (open circles). In Figure

5 t

0 0

0

'

0

0

0 0

'

b

-'"t -15 -2.0

-1.0

0.0

1.0

ASH

Figure20. Correlation of the change in proton and carbon-13 chemical shifts for a number of heterocyclic systems upon protonation. o Positions alpha to N; other positions.

D. G. FARNUM

166

20 an attempt has been made to correct for the anisotropy of the nitrogen and its lone pair by plotting the change in chemical shift which occurs at a given position upon addition of a proton to a nitrogen atom in a heterocycle or its anion. The scatter is not improved, and positions alpha to nitrogen (open circles) are, again, no worse than any others. The failure to show a correlation must mean that either the proton, or carbon shifts, or both are not linearly related to charge in these systems. Of course, the upfield shift of the alpha-carbon on protonation of pyridine tells us that alphacarbon-13 shifts are anomalous, but the proton shifts are also puzzling. Upon protonation, the alpha-proton moves downfield 1-08 p.p.m., the beta, 1.71 p.p.m., and the gamma 1.75 p.p.m. The increasing downfield shift further from the site of protonation, and the large downfield shift at the beta-position are contrary to intuitive estimates of charge distribution in the ion. Hence, it seems quite possible that neither cnmr nor pnmr shifts will correlate linearly with charge in the nitrogen heterocycles. The influence of the paramagnetic term from the nitrogen lone pair on the screening constant for the alpha-proton in pyridine was suggested by Gil and Murrell (1964). Fraenkel and coworkers (1968) related the earlier suggestion to their discovery of anomalous proton shifts in phenyl-lithium and phenylmagnesium bromide, where the protons ortho to the carbon-metal bond showed large downfield shifts contrary to expectations based on charge polarization. Shortly thereafter Jones et al. (1969) reported similar anomalous downfield shifts in carbon-13 spectra of phenyl-lithium and phenylmagnesium bromide, and again emphasized the similar behaviour between these organometallics and pyridine on protonation. The data are indicated

171.9 (+ 43.2)

(+ -76) 8-02

(- -24) 7.02 (- *30) 6.96

141.6 (+ 12.9) 126.3 (- 2.4)

(+ ‘38) 7-64

61 (+ 113)

H -52

(- 1.08) 8 . 6 8 0 150.4 (+ 7.8) (- 1-75) 7.62

124-1 (- 5 - 1 ) 123.7 (- 24-8)

126.0 (- 2.7) 124.7 (- 4.0)

125.6 (- 3.1)

(- 1.71) 7-20

140.1 (+ 11-4)

(- -24) 7.02

9 ; ;

@142-6

9.3 7

129.2 148-5

CHARGE DENSITY-NMR CHEMICAL SHIFT CORRELATIONS

167

on the structures following. Carbon-13 shifts are in bold type, proton shifts are in italics, and the change in shift downfield (negative values) or upfield (positive values) upon protonation is given in parentheses. Of the several factors which affect up, the bond order and AE were considered to be the two likely contenders to account for the upfield shift of both carbon-13 and protons ortho to the carbonmetal bond, and for the upfield shift of carbon-13 and the less-thanexpected downfield shift of protons ortho to the nitrogen lone pair upon protonation. Although both factors operate in the right direction in pyridine (C=N becomes less of a double bond on protonation, the bond order decreases, and up decreases; AE becomes larger upon protonation because the lone pair is lost, and op decreases), the bond order in the metallobenzenes should actually increase on protonation, thereby increasing up, and causing a downfield shift. Since the charge polarization difference should also cause a downfield shift, the only factor left to cause an upfield shift is the change in AE. It was inferred that this was also the dominant effect in pyridine, especially since calculations suggested that the bond order change could account for only some 20% of the upfield shift upon protonation of pyridine. Conceptually, it is difficult to see why the effect of a change in AE might change sign for ortho-, rnetu-, and paru-positions of the ring. Using a different model for the paramagnetic term-the McConnell point/dipole model (1957) -Fraenkel et al. (1968) have calculated up for the ortho- and metu-protons of phenyl-lithium, assuming that the entire shift at the remote para-position results from a change in up. The significant feature of the McConnell model for these purposes is its angular dependence as expressed in equation (6) and Figure 2 1. AX AuP = y (1 - 3 COS*O)

3R

Figure 21. Definition of parameters for equation (6).

168

D. G. FARNUM

One can, therefore, calculate a value for Ax assuming all the downfield shift at the para-position on protonation of phenyl-lithium results from the change in Au, (Aup = -0-30). From the calculated A x one can then calculate paramagnetic shifts at the ortho- and para-positions of +0*85 and -0.22, respectively, which are in good agreement with the observed values of +0.76 and -0-24. Agreement for phenylmagnesium bromide is less satisfactory but of the correct sign. It is surprising that no one has yet applied this simple and satisfying approach t o the carbon-13 shifts of these systems. TABLE 11 Calculated and Observed Carbon-13 Chemical Shifts in Azines, Jaffk and BenC CNDO-MO's, Average Excitation Energy (AEE) and Different Excitation Energy (DEE) Approximationsa

Compound

Position

CJ

obsd

AEE

DEE

Pyridine

150.4 124.1 136.1

141.7 125.5 133-1

149-0 126.7 135.2

Pyrazine

145.8

Pyridazine

152.6 127.7

137.8 140.4 130.0

145'0 149.7 129.0

Pyrimidine

159.2 157-6 122.6

155.8 145.3 123-1

167'4 157.2 129.5

'Data from Tokuhiro and Fraenkel, 1969 (C&

= 128.7).

Unfortunately, it cannot be applied to pyridine securely because of the indeterminate effect of charge delocalization on the para chemical shift. Nonetheless the sign and magnitude of the effects in phenyl-lithium would go a long way towards removing the discrepancies in the pnmr spectra of pyridine and pyridinium ion. The quantitative effect of changes in AE on the cnmr shifts for several heterocycles was explored by Fraenkel (Tokuhiro and Fraenkel, 1969). They concluded that the experiment was critically dependent on the MO approach chosen t o calculate the charge densities and A E ' s and found that the Ben&-JaffCCNDO wave functions, gave the best correlation. The best calculated and observed chemical shifts are given in Table 11, and plotted in Figure 22. The scatter from the perfect correlation line is quite acceptable, but the approach

CHARGE DENSITY-NMR CHEMICAL SHIFT CORRELATIONS

169

has not been applied to protonated heterocyclic cations yet. Their analysis of the charge dependence suggested that the l / r 3 term (Grant’s t3 term) accounted for only about one-half of the factor of 160 in the expression Ao = 160 q. The remainder was due to terms which were less uniformly linear in q. These conclusions are contrary to those of Emsley (1968). Later work by Adam et al. (1969) on the cnmr spectra of heterocyclic anions, neutral molecules and cations was in general agreement that changes in AE were dominant in accounting for these

8 13C OBSERVED Figure 22. Correlation of observed and calculated carbon-13 chemical shifts for a number of heterocycles. For data and references see Table 11. o Using AEE approximation; 0 DEE approximation.

carbon-13 shifts. However, the quantitative agreement with protonation shifts was less than satisfactory as shown in Table 12. The sign of the shift is correctly predicted in every case for carbons not alpha to nitrogen, and in nine of the fifteen cases for carbons alpha to nitrogen. However, the magnitudes of the shifts, and particularly the downfield shifts, are most often grossly underestimated. The theoretical treatment seems to accommodate the upfield shifts alpha to nitrogen by a displacement of all shifts upfield. The authors’ calculations suggest that an increase of AE by 20% (2 eV) results in an upfield shift of approximately 50 p.p.m. Grant has come to the position that both bond order effects and changes in AE must be considered in order to understand chemical shifts in heterocycles (Pugmire et al., 1968- 1973). Some of his data are also collected in Table 12. In Column B, carbon shift differences

170

D. G. FARNUM

TABLE 12 Calculated and Observed Carbon Shifts on Protonation of a Number of Heterocyclic Anions, Neutral Molecules, and Cations Carbon Shift on Protonation Calcd. Obsd.'

A0

Bb

+1-1 -2.0 -1.0

+0*6 -0.8 -1.9

Pyrazine

+7.8 -5.0 -12'4 +2*6

+4*7

Pyrazine cation

-1.0

+1-4 +2*4

Species Pyridine

Pyridazine Pyrimidine

Pyrimidine cation

Imidazole anion Imidazole Pyrrole anion Pyrazole anion Pyrazole

Positiond

+ l a 1

-10-1 +7-2 -1.3 -2.5(-3*Oe) +la2 -0.2 -3.6(-3.1e) +8.5(8*ge) +4.4 +1*6 +2*2 +8-5 -1.6 +4*0 -2.0 -0.5 -3.5

+0*5 +0.8 -0.3 -2.9 +la3 -08 +1.0 -0.2 -0.4 +1.3 +2-1 +I-3 -0.4 +1-4 -0.4 -0.6 -0.5 +2-0 -0.5

+Om4 -2.1 +4.9 -0.3 -1.2 +11-2 +1.4 -0.9 -0.4 -2.8 -7.4 -2.2 +1*7 -9.2 -8-3 -7.4 -1 1.9 -3.1

cc

+44 +2'3 -1.6 +3-1 +7-2 -4.3 -2-4 -2.6 -7.4 +2*0

Data from Adam et. al., 1969. Data from Pugmire and Grant,1968, ignoring changes in AE. Data from Pugmire and Grant, 1968, corrected for changes in AE. Positions alpha to nitrogen are in bold-face type. Values from Pugmire and Grant, 1968 (neat or sat. aq. solution).

are calculated ignoring changes in AE, but including the bond order term." Again, the sign is correctly predicted for all carbons not alpha to nitrogen, and for eight of the fifteen cases alpha to nitrogen. In column C, the shifts are corrected for changes in AE among the l S The differences between the models used by Grant and Adam to calculate chemical shifts are discussed in some detail by Adam et al. (1969). Among other things, Grant ignores ud. Although Ud is occasionally of significant magnitude, changes in Ud on protonation are not (Adam et al., 1969), so that the error thereby introduced is negligible.

CHARGE DENSITY-NMR CHEMICAL SHIFT CORRELATIONS

171

heterocycles by inclusion of a factor f which corrects the 10 eV average excitation energy to a value more in accord with the relative value calculated from the molecular orbital model used (CNDOSCF). The correction has, unfortunately, only been applied t o the five-membered heterocycles, where it corrects the sign for three positions alpha to nitrogen, changes the sign of one beta-carbon in the wrong direction, and improves the overall quantitative fit. Grant

Figure 23. Correlation of observed and calculated changes in carbon-13 shifts for a number of heterocyclic systems upon protonation. For data and references see Table 12. o Data from Adam et al. (1969) (Column A); 0 Data from Pugmire and Grant (1968) ignoring changer in aE (Column B). x Data from Pugmire and Grant (1968) corrected for changes in AE (Column C).

notes that the f factor deviates from unity the most for those systems which have two lone pairs, and the least for those which are devoid of lone pairs, as expected if the n+r* transition is important in determining AE. His calculations suggest that a change in AE of 0-15 eV (1.5%) results in a change in chemical shift of about 5 p.p.m. An important point emerges from comparison of columns B and C of Table 12: namely, the AE correction results in about the same upfield shift at positions alpha to nitrogen and at other positions in the molecule. It, theretore, cannot account for the anomalous behauiour of the alpha-carbon shift as compared with the other positions. The status of these several treatments is perhaps best illustrated in Figure 23 where the calculated protonation shifts are plotted against

172

D. G. FARNUM

the observed shifts. The open circles are from Adam’s approach, and show excessive scatter around the perfect correlation line. In general, the range of calculated values is much too narrow, in accord with the earlier suggestion that the approach does not sufficiently distinguish between alpha positions and others. The solid circles are from Grant’s data in column B of Table 12, and include bond order effects, but are not corrected for differences in AE. Although the range of values is much improved, the scatter is still excessive. The crosses are from Grant’s data including corrections for differences in AE (Column C of Table 12). There is definitely an improvement in the correlation, but scatter is still unacceptable and more points are needed. It is quite apparent that, because of the complexity of these systems and the required approximations in the mathematical models used to describe them, one is inhibited from drawing conclusions by an immobilizing tangle of qualifications. Nonetheless, the ideas that emerge from the tangle deserve the careful attention of creative minds that may find less equivocal ways of testing them.

4. CONCLUSIONS

Both the theoretical and empirical studies discussed above emphasize the need to consider several factors in addition to linear charge polarization when attempting to interpret either proton or carbon-13 chemical shifts in terms of charge distribution. Prime among these are (1) the effect of neighbouring group anisotropy on proton shifts, (2) the E 2 term in proton shifts at high positive or negative charge densities, (3) the paramagnetic term in proton shifts in systems with lone pair eleckons, (4) a possible reversal of slope for the carbon-13 shift-charge density correlation at high positive charge density, ( 5 ) the importance of bond order terms in determining carbon-13 shifts, (6) the importance of average excitation energies (AE)in determining carbon-13 shifts, particularly in systems with lone pairs. Attempts to find reliable quantitative empirical assessments of each of these factors in the literature were thwarted either by the narrow range of charge over which charge density could be determined with confidence, or the complexity of the systems studied, which required emasculating assumptions in the mathemajical models used for calculation of the electronic properties of the molecules and the chemical shifts. Confident correction of raw

CHARGE DENSITY-NMR CHEMICAL SHIFT CORRELATIONS

173

chemical shift data for these factors, in all but the simplest cases, covered adequately by reliable empirical correlations, must await more thorough empirical studies in which systems are chosen carefully t o isolate some of the effects. Although we can certainly echo Olah and Westerman’s statement ‘‘I C nmr shifts, if used with proper consideration of all factors involved, are a very powerful tool in studying the structure of carbocations, including the trend of charge distribution” (1973), there is still n o general agreement on what constitutes “proper consideration of all factors”, indeed, on what the factors are, nor on the quantitative contribution each makes to a given case. Until such agreement emerges, the powerful tool implied in that statement, as well as in the earlier review by Fraenkel and Farnum (1968), remains largely illusory for hosts of substances of interest.

Acknowledgement The author is indebted t o Dr. Anthony D. Wolf for assistance in the initial literature research and to the faculty and staff of the Chemistry Department at the University of California, San Diego, for their hospitality and stimulation during my sabbatical leave there. I would especially like t o give credit t o Dr. Charles E. Perrin, whose incisive, forthright questions laid bare a number of weaknesses in the theoretical discussion. Any remaining weaknesses are there not because Dr. Perrin failed to discover them, but because I could not do justice t o these points.

REFERENCES Adam, W., Grimison, A., and Rodriguez, G. (1969). J. Chem. Phys. 50,645. Adcock, W., Dewar, M. J. S., and Gupta, B:D. (1973).J. Amer. Chem. SOC. 95, 7353. Bacon, J., and Gillespie, R. J. (1971). J. Amer. Chem. SOC. 93, 6914. Becker, E. D. (1969). “High Resolution Nuclear Magnetic Resonance.” Academic Press, New York. Bovey, F. A. (1969). “Nuclear Magnetic Resonance Spectroscopy”. Academic Press, New York. Breslow, R., Groves, J. T., and Ryan, G. (1967). J. Amer. Chem. SOC. 89, 5048. Buckingham, A. D. (1960). Can. J. Chem. 38,300. Closs, G. L., Boll, W. A., Heyn, H., and Dev, V. (1968).J. Amer. Chem. SOC. 90, 173. Dhami, K. S., and Stothers, J. B. (1965). Con. J. Chem. 43, 510. Dixon, W. T. (1972). “Theory and interpretation of Magnetic Resonance Spectroscopy.” Plenum Publishing Co., New York.

174

D. G. FARNUM

Emanuel, R. V., and Randall, E. W. (1969).J. Chem. SOC.( A ) ,3002. Emsley, J. W. (1968). J. Chem SOC. ( A ) ,1387. Farnum, D. G. (1967).f. Amer. Chem. SOC.89, 2970. Farnum, D. G., Mehta, G., and Silberman, R. G. (1967).J. Amer. Chem. SOC. 89, 5048. Fraenkel, G., Carter, R. E., McLachlan, A. D., and Richards, J. H. (1960). J. Amer. Chem. SOC.82, 5846. Fraenkel, G., Adams, D. G., and Dean, R. R. (1968).]. Phys. Chem. 72,944. Fraenkel, G., Dayagi, S., and Kobayashi, S. (1968).]. Phys. Chem. 72, 953. Fraenkel, G., and Farnum, D. G. (1968). In “Carbonium Ions”, G. A. Olah and P. von R. Schleyer, Eds. Vol. 1. Wiley, Interscience, New York. Gil, V. M. S., and Murrell, J. N. (1964). Trans. Faraday SOC. 60, 248. Gray, G. A., Ellis, P. D., Traficante, D. D., and Maciel, G. E. (1969). J. Magnetic Res. 1, 41. Grutzner, J. B., Lawlor, J. M., and Jackman, L. M. (1972).J. Amer. Chem. SOC. 94, 2306. Grutzner, J. B. (1973). Paper No. 34, First Fall Organic Chemistry Conference of the ACS, Cape Cod, and private communication. Jackman, L. M., and Sternhell, S. (1969). “Applications of Nuclear Magnetic Resonance Spectroscopy in Organic Chemistry.” Pergamon Press, New York. Jones, A. J., Grant, D. M., Russel, J. G., and Fraenkel, G. (1969). 1.Phys. Chem. 73, 1624. Karplus, M., and Das, T. P. (1961). J. Chem. Phys. 34, 1683. Karplus, M., and Pople, J. A. (1963).J. Chem. Phys. 38,2803. Katz, T. J., and Garratt, P. J. (1963). J. Amer. Chem. SOC.85, 2852. Katz, T. J., and Garratt, P. J. (1964).J. Amer. Chem. SOC.86, 5194. LaLancette, E. A., and Benson, R. E. (1963).J. Amer. Chem. SOC.85,2853. Lauterbur, P. C. (1961). Tetrahedron Letters 274. Lazzeretti, P., andTaddei, F. (1971). Org. Magn. Res. 3, 283. Levy, G. C., and Nelson, G. L. (1972). “Carbon-13 Nuclear Magnetic Resonance for Organic Chemists.” Wiley-Interscience, New York. McConnell. H. (1957).J. Chem. Phys. 27, 226. McKeever, L. D., and Waack, R. (1971).J. Organometal. Chem. 28, 145. MacLean, C., and Mackor, E. L. (1961). Mol. Phys. 4 , 2 4 1 . Mateescu, G. D. (1973). Abstracts of 165th ACS Meeting, Dallas. Memory, J. D. (1968). “Quantum Theory of Magnetic Resonance Parameters.” McGraw-Hill, New York. Musher, J. L. (1962).J. Chem. Phys. 37, 34. Olah, G. A., and Schleyer, P. von R., Eds. (1968). “Carbonium Ions.” Vol. I. Wiley-Interscience, New York. OIah, G. A., and White, A. M. (1968)./. Amer. Chem. SOC.90, 1884. Olah, G. A., and White, A. M. (1969).J. Amer. Chem. SOC. 91,5801. Olah, G. A., Bollinger, J. M., and White, A.M. (1969).J. Amer. Chem. SOC. 91, 3667. Olah, G. A., and Mateescu, G. D. (1970).J. Amer. Chem. SOC. 92, 1430. Olah, G. A., and Schleyer, P. von R., Eds. (1970). “Carbonium Ions.” Vol. 11. Wiley-Interscience, New York. Olah, G. A., Kelly, D. P., Jeuell, C. L., and Porter, R. D. (1970). J. Amer. Chem. SOC.92, 2544. Olah, G . A., and Porter, R. D. (1971).]. Amer. Chem. SOC.93,6877. Olah, G. A., Porter, R. D., and Kelly, D. P. (1971).]. Amer. Chem. SOC. 93, 464. Olah, G . A., andLiang, G. (1972).J. Amer. Chem. SOC.94, 6434.

CHARGE DENSITY-NMR CHEMICAL SHIFT CORRELATIONS

175

Olah, G. A., and Schleyer, P. von R., Eds. (1972). “Carbonium Ions.” Vol. 111. Wiley-Interscience, New York. Olah, G. A., Porter, R. D., Jeuell, C. L., and White, A. M. (1972a). J. Amer. C h e m SOC. 94,2044. Olah, G. A., Liang, G., andMo, Y. K. (1972b). J. Amer. Chem. SOC. 94,3544. Olah, G. A., Halpern, Y., Mo, Y. K., and Liang, G. (1972c).J. Amer. Chem. SOC. 94,3554. Olah, G. A., and Schleyer, P. von R., Eds. (1973). “Carbonium Ions.” Vol. IV. Wiley-Interscience, New York. Olah, G. A., and Westerman, P. W. (1973).]. Amer. Chem. SOC. 95, 7530. Oth, J. F. M., Baumann, H., Giles, J. M., and Schroeder, G. (1972). J. Amer. Chem. SOC. 94,3498. Paudler, W. W. (1971). “Nuclear Magnetic Resonance.” Allyn and Bacon, Boston. Pople, J. A. (1962).]. Chem. Phys. 37, 5360. Pugmire, R. J., and Grant, D. M. (1968a).J. Amer. Chem. SOC. 90, 697. Pugmire, R. J., and Grant, D. M. (1968b).J. Amer. Chem. SOC. 90,4332. Pugmire, R. J., Grant, D. M., Robins, M. J., and Robins, R. K. (1969).J. Amer. Chem. SOC. 91, 6381. Pugmire, R. J., and Grant, D. M. (1971).J. Amer. Chem. SOC. 93, 1880. Pugmire, R. J., Robins, M. J., Grant, D. M., and Robins, R. K. (1971).]. Amer. Chem. SOC. 93,1887. Pugmire, R. J., Grant, D. M., Townsend, L. B., and Robins, R. K. (1973). 1 . Amer. Chem. SOC. 95,2791. Ray, G. J., Kurland, R. J., and Colter, A. K. (1971). Tetrahedron 27, 735. Sandel, V., and Freedman, H. H. (1963).J. Amer. Chem. SOC. 85,2328. Schaefer, T., and Schneider, W. G. (1963). Can. J. Chem. 41, 966. Schweizer, M. P., Chan, S. I., Helmkamp, G. K., and Ts’o,P. 0. P. (1964). J. Amer. Chem. SOC. 86,696. Spiesecke, H., and Schneider, W. G. (1961). J. Chem. Phys. 35, 722. Spiesecke, H., and Schneider, W. G. (1961). Tetrahedron Letters 468. Stothers, J. B., and Lauterbur, P. (1964). Can. J. Chem. 42, 1563. Stothers, J. B. (1972). “Carbon-13 Nuclear Magnetic Resonance Spectroscopy.” Academic Press, New York. Strong, A. B., Ikenberry, D., and Grant, D. M. (1973).]. Magn. Res. 9, 145. Takahashi, K., Konishi, K., Ushio, M., Takaki, M., and Asami, R. (1973). J. Organometal. Chem. 50, 1. Timberlake, J. W., Thompson, J. A., and Taft, R. W. (1971). J. Amer. Chem. SOC. 93, 274. Tokuhiro, T., and Fraenkel, G. (1969).J. Amer. Chem. SOC. 91, 5005. van der Kooij, J., Velthorst, N. H., and MacLean, C. (1972). Chem. Phys. Lett. 12, 596. Waack, R., Doran, M. A., Baker, E. B., and Olah,G. A. (1966).J. Amer. Chem. SOC. 88, 1272. Wiberg, K. B., and Nist, B. J. (1961).J. Amer. Chem. SOC. 83, 1226.

This Page Intentionally Left Blank

The Norbornyl Cation: A Reappraisal of

its Structure Under Stable Ion Conditions' G. M. KRAMER Corporate Research Laboratories, Exxon Research and Engineering Co., Linden, New Jersey 07036, U.S.A.

1. 2. 3. 4. 5.

6. 7. 8. 9. 10. 11.

.

Introduction Possible Structures of the Norbornyl Cation Solvolytic Background Deamination of 2-Norbornylamine Isotope Effects Theoretical Status The Search for a Protonated Cyclopropyl Ring The Control System ESCA I3c-nmr. 'H-nmr Raman Spectra Related Ions . Summary References

.

. .

.

.

.

. .

.

. .

. . . . . . . . . . . . . . .

177 179 180 188 190 192 194 196 199 202 211 215 218 221 222

1. INTRODUCTION Determining the structure of the norbornyl cation has piqued the imagination and challenged the resources of chemists for a lengthy period. The problem apparently was initiated by the suggestion of Editor's footnote: A forthcoming volume of the series is expected to contain an article by C. A. Olah, J. E. Nordander, and P. v. R. Schleyer on the norbornyl cation, its structure and relevance to chemistry.

178

G. M. KRAhfER

Nevell, de Salas and Wilson (1939) that the rearrangement of camphene hydrochloride [ 11 into isobornyl chloride [ 21 involved the

formation of a mesomeric or non-classical carbonium ion [3]. The rearrangement which had earlier been studied by Meerwein and van

[31

Emster (1922) could also be explained by the alternative proposal that two rapidly equilibrating classical ions [4] and [5] were the actual reaction intermediates.

Ions [3], [4] and [ 5 ] were the forerunners of an extensive collection of cation structures which have been embroiIed in the non-classical-classical controversy. Solvolysis studies have played a major role in providing the information which has fuelled the arguments. The solvolysis data have by and large been obtained in systems where the pertinent intermediates or transition states have but a fleeting existence. In the norbornyl system such information was originally interpreted as providing the strongest case for the existence of a non-classical species. More recently the major arguments in the norbornyl system have been challenged by H. C. Brown’s systematic studies which appear to indicate that many results can be explained by consideration of changes in strain and of steric hindrance occurring in the ionization or product-forming steps where classical ions represent the structure of intermediates. Brown’s results tend to negate nonclassical theory as it applies to the structure of the ion reached during solvolysis of norbornyl derivatives and, with few exceptions, there seems to be no compelling reason to require that a bridged ion is formed during solvolysis

THE NORBORNYL CATION

179

(Brown, 1972). On the other hand, there are data which some authors have interpreted as indicating varied degrees of participation or bridging of the bond between C-1 and C-6 during the solvolysis of exo-2 norbornyl compounds. This survey will briefly review the major features of the solvolysis studies but is primarily concerned with appraising a substantial body of information stemming from studies of the norbornyl cation in strong acid media. When a norbornyl derivative is dissolved in SbF,-HS03F or a large number of related acids, an apparently long-lived ion is formed. The solutions have been subjected to many spectral studies in efforts to elucidate its structure under “stable” ion conditions. These include the use of ESCA (C-1s x-ray electron spectroscopy), C-nmr, H-nmr and Raman spectroscopy. Under these conditions the ion undergoes intramolecular rearrangements which have been studied by nmr, and, if it forms a stable nonclassical structure, it must exhibit a protonated cyclopropyl ring. The formation and proton exchange behaviour of protonated cyclopropanes during the rearrangement of alkyl cations in similar solutions has recently been studied and these techniques have also been applied t o the norbornyl problem. In addition to these probes, theoretical calculations are being made in efforts to decide upon the inherently most stable structure of the ion. These will be discussed along with the conclusions which have been derived from the study of isotope effects in solvolysis and substituent effects in “stabilizing” media.

2. POSSIBLE STRUCTURES OF THE NORBORNYL CATION The non-classical-classical debate centres on the question of the relative energy of various structures of the ion. This energy must reflect the bond lengths, angles, charge distributions and intramolecular interactions present in the structure. Before proceeding further, it is advisable to consider the major characteristics of the classical and non-classical structures of the norbornyl ion. 7

[61 Classical

[ 71

[81 Possible non-classicalstructures

[91

180

G. M. KRAMER

The classical structure [ 61 is geometrically differentiated from the non-cIassica1 possibilities by virtue of C-6 not being symmetrically positioned between C-1 and C-2. The positive charge density is considered to be largest at C-2 but the actual charge density at C-2 is unknown. The charge is certainly distributed throughout the structure but knowledge of the distribution must await a complete SCF a6 initio calculation. Non-classical structures contain a bridging C-6 atom that requires the existence of an elongated cyclopropyl ring. One can consider the structure as arising by the sharing of u-electrons in the 1-6 bond between the 1, 2 and 6 atoms. The possible participation of these electrons as evidenced by enhanced solvolytic rates is a major factor in kinetic discussions. The electron distribution was at one time considered to lead to the spreading of positive charge to the 1 and 6 position (Winstein and Trifan, 1949, 1952). To the extent that charge is delocalized to C-6 structures [ 71 and [8] acquire varying weight as resonance forms of the comer protonated cyclopropyl ring. Both structures contain a carbon atom forming five bonds and whether a carbon atom can form a stable structure with this many covalent bonds is a matter of concern, the existence of species such as CH: notwithstanding. The non-classical ion may exist as edge or face protonated nortricyclene, represented by structure [9], as an alternative. This species is not a resonance structure of [7] or [8]. If it is a discrete intermediate, one might expect to be able to detect the displaced proton via a suitably designed exchange experiment. This review now considers the evidence which may allow US to choose between the various structures of the ion.

3. SOLVOLYTIC BACKGROUND Rates and product distributions obtained during the solvolysis of secondary and tertiary cationic precursors provide the data for the norbornyl controversy. One of the reasons the debate has not been ended is the complexity of the solvolysis mechanism. Rate measurements should provide information on the energy difference between the ground and transition states of the reaction. Their analysis, however, is often clouded by a series of questions. The first question concerns the reaction mechanism. Solvolysis of

THE NORBORNYL CATION

181

most of the substrates in the debats are often interpreted in terms of a pure S N l mechanism in which the rate-determining step is the ionization. An immediate problem is caused by the possible existence of an SN2 pathway, the participation of solvent in the ionization of secondary substrates being an important but often a relatively uncertain factor. In the limiting case of no solvent participation there are still difficulties in interpreting rate processes. Thus the importance of tight and separated ion pairs and the possibility of internal return complicate both rate interpretations and the analysis of product distributions (Winstein and Trifan, 1952; Winstein and Schreiber, 1952). Before turning to the principal arguments and lines of evidence it is interesting to consider the use to which rate data is often applied. Substituents at different carbon atoms around the norbornyl skeleton usually result in a change of the solvolysis rate. Assuming that such substitution affects a rate determining sN1 ionization, the data very ofteh are found to lend themselves to two interpretations which have different conclusions but yet are difficult to resolve. The dilemma stems from the fact that substitution inevitably introduces both an electronic and a steric change which are important in both the ground and transition states of the reaction. The assessment of the relative importance of the electronic and steric factors is often critical in drawing conclusions as to the importance of bridging or 5-participation in the process and therein lies the basis for debate. Three major arguments have been presented as evidence of bridging or o-participation during solvolysis of norbornyl compounds. (a) Unusually fast rates of solvolysis of exo-derivatives, camphene hydrochloride reacting 6000 times faster than t-butyl chloride (Brown, Hughes, Ingold and Smith, 1951). (b) High exolendo rate ratios as illustrated by a value of 350 for the acetolysis of the 2-norbornyl brosylates (Winstein and Trifan, 1949 and 1952). (c) The highly selective formation of exo-products upon acetolysis of exo- and endo-compounds, even with 7,7-dimethyl substituents (Winstein and Trifan, 1949 and 1952). The relative rates of solvolysis of camphene hydrochloride and t-butyl chloride indicate that the free energy of activation is 5 kcal mole-' higher with the latter compound. This might be attributed either to non-classical stabilization of the camphenyl transition state

182

G. M. KRAMER

or to the relief of 5 kcal of strain energy upon ionization of camphene hydrochloride. After measuring the relative rates of ethanolysis of compounds [ 11, [ 101, [ 111, [ 121 and [ 131 (Brown and Chloupek, 1963), given below, Brown has indicated that it is quite possible that the latter explanation is correct.

13,600

[ 11

355 [121

60 ~ 3 1

In all of these compounds solvolysis will lead t o a tertiary ion. The series [ 101, [ 131 , [ 111 clearly indicates the strain argument, and one may note that the difference in rates between [ 11 and [ 111 corresponds to an energy difference of only 1-1kcal mole-'. The data do not prove that non-classical stabilization of the transition state in [ 11 and [ 121 is not partly responsible for the rate differences but rather suggests that relief of strain could account for the results. Other factors, particularly differential solvation of the ground state and transition state and the possibility that solvolysis may not be of a limiting type but involve reaction with solvent, may also play a role but are difficult to evaluate. In any case the rate of solvolysis of exo-compounds does not appear to be unusually rapid when viewed in this light. The second factor used to support the non-classical position is the high exo/endo rate ratio usually found in solvolysis. The high ratios have been taken as evidence for delocalization of the u electrons of the bond between C-1 and C-6 into the rear of a developing p-orbital at C-2 as an exo-substituent ionizes. Similar delocalization during the ionization of an endo-substituent is not anticipated because of poor orbital alignment. From the classical point of view it is proposed that the high rate ratio stems from a normal solvolysis rate of the em-compound

183

THE NORBORNYL CATION

accompanied by a slow solvolysis of endo-compounds because of steric interaction of the leaving group and the endo-6 hydrogen. Several series of experiments may be noted which tend t o substantiate both views. The first are relative rates for [ 141 - [ 171 obtained by Schleyer and coworkers and given below on the effects of substituents at the 6-position in augmenting solvolysis of 2-exo-norbornyl tosylates (Schleyer, 1972; Schleyer et al., 1965; Strang and Schleyer, 1968).

1.00

0.5

~ 4 1

~

5

1

A

CH30

OTs

The methyl, dimethyl and methoxy substituents would be expected t o facilitate participation of the 1-6 electrons during ionization of the tosylates by lowering the positive charge at C-6. Thus, if bridging were important, [ 171 should solvolyse considerably faster than [ 141 . In fact, all of the substituted compounds react more slowly leading t o the apparent conclusion that o-participation is not manifested in these solvolyses. Perhaps it might be argued that the rate-determining step in these reactions followed ionization to a tight ion pair but the results are not easily reconciled with the non-classical hypothesis. Nevertheless, it is possible t o offer an alternative explanation which might reconcile the apparent lack of a substituent effect in [ 151 , [ 161 and [ 171 with bridging in the norbornyl ion itself. This is accomplished by proposing that the substituents interfere with the formation of a bridged ion by introducing steric interactions. The argument is essentially the reverse of that used by the classical school to explain the low rate of solvolysis of endo-compounds. Thus, the lower rate of solvolysis of [15] as opposed t o [14] might be rationalized by noting that formation of the bridged ion in [15] introduces a non-bonded interaction between the methyl group and

184

G. M. KRAMER

the hydrogen at C-1 as these groups acquire a conformation with decreasing exo-character. Donaldson (1958) appears to have been one of the earliest to advance this argument as an explanation of the low rate of solvolysis of the 6,6-dimethyl tosylate [ 161. The steric interaction is presumed to be sufficiently strong to negate the enhancement which might have been expected on the basis of the electromeric properties of the alkyl groups. Both preceding arguments appear to have sufficient merit to render a positive assessment of ionic structure or the extent of participation from exolendo rate ratios rather dubious, and, though we favour the electromeric argument, the possibility of steric hindrance to bridged ion formation is certainly important. Sargent (1972) has reviewed a number of related studies (Corey and Glass, 1967; Baker and Hudec, 1967) in which bridging appears to be impeded by locking remote positions of the norbornyl esters into fixed configurations (see Table 1). This has been interpreted as introducing a strain barrier to the bridging process because of sizeable energy differences between the ions containing the exo and TABLE 1 Relative Rate of Acetolysis at 25' 1

W

O

T

S

ATs

643

2-3 OTS

X = H, Y = OTs, 0.67 X = OTs, Y = H, 5.70

(a) Corey and Glass, 1967; (b) Calculated from Schleyer et al., 1965.

T H E NORBORNYL CATION

185

endo fused rings which should be partly reflected in the bridged species. As a result, the exolendo rate ratio is markedly reduced from that in the parent compounds and is taken as indirect evidence of bridging in the latter. This interpretation has not been universally accepted, one reason being that the product ratios do not decline in like manner to the rate ratios. Another factor is that, even if the rate ratios are taken to indicate the existence of some participation of the electrons in the bond between C-1 and C-6 in the solvolysis of exo-2-norbornyl tosylate, this is far removed from proving the existence of the symmetrical and hence fully delocalized ion at the transition state or as an intermediate in the resction. Hence, if one concludes that the exo- and endo-derivatives of the restricted systems are solvolysing t o essentially unbridged or classical ions, the extent of bridging to be deduced from high exolendo rate ratios found with less restricted systems is unclear. On the other hand, it is difficult to explain why the exo- and endo-derivatives should solvolyse at such similar rates from the classical view. Model compounds do not indicate marked changes in leaving-group interactions between the parent and restricted systems .and convincing arguments to explain the apparently reduced exoreactivity without invoking inhibition of partial (T participation are not obvious. Another series of experiments was also aimed at finding out if the exolendo ratios might reflect steric inhibition t o the leaving group. To do this, a series of tertiary methyl p-nitrobenzoates of three bicyclic systems with increased crowding were prepared and the rates of solvolysis in 80% acetone determined, (Brown et al. 1967). The exolendo ratios in 1181, { 191 and [ Z O ] and their epimers systematically increased as crowding increased exhibiting a good correlation. Whilst a wide review of solvolysis results is beyond the scope of this review, brief mention must be made of the Foote-Schleyer

02

CH, OPNB OPNB

exo --

endo

- 17

[I81

885

4300

186

G. M. KRAMER

correlation (Schleyer, 1964; Foote, 1964; Brown et al., 1966). In principle, this correlation was proposed as a means of calculating the rate of solvolysis of exo- or endo-compounds in the absence of either steric effects or a-participation, assuming that the solvolysis of model compounds was of a limiting type. Deviations from the calculated values would then reflect these factors. Thus, if calculation for an endo-compound agreed with the experimental rate, steric 4factors would not be important, whereas if calculations for exo-compounds agreed with experiment, a-participation must be minimal. The correlation generally predicted markedly higher rates than were observed when it was applied to crowded endo-tosylates (Brown, 1972).The natural conclusion is that there appears to be an increase in non-bonded strain on going from the ground state to the transition state. The utility of the correlation is however unclear and the assumptions upon which it was based are currently being reconsidered (Fry et al., 1970). Even without this support, the classical interpretation of the exolendo rate ratios appear as plausible as the nonclassical in view of the fluid nature of the arguments which can be made with respect to participation and interaction of substituents at C-6. The third factor, namely that there is an extremely high selectivity to exo-products during the acetolysis of most norbornyl derivatives loses much of its force when Goering and Schewene's reaction coordinate diagram for the acetolysis of exo- and endo-norbornyl acetates is considered. Although the data were originally thought to favour the non-classical model (Goering and Schewene, 1965), cogent arguments have been presented to eliminate a unique interpretation of the data (Brown, 1972). The main point is that the product ratio is controlled by two transition states, one involving the reaction of norbomyl ion with a nucleophile leading to the exoacetate and the other for the same reactants yielding the endoacetate. The product ratio itself does not provide enough information to characterize the norbornyl ion. The situation may perhaps be made clearer by noting that high exolendo product ratios are also found in the reaction of tertiary norbornyl derivatives where Uparticipation is generally believed to be absent, and hence the product ratio must be a poor barometer of structure. The free-energy diagram for the solvolysis of 2p-anisylcamphenilyl p-nitrobenzoate in 80% acetone, which parallels that for the acetolysis of norbornyl tosylate, illustrates Brown's arguments (Figures 1 and 2). One notes that a difference of 5 kcal in the free energy of activation corresponds to a product ratio of 5000/1.

THE NORBORNYL CATION

187

Figure 1. Free-energy diagram for the acetolysis of exo- and endo-norbornyl tosylate (Brown, 1972).

AG-5.2 I --

AG'

20.1

-I? LGC 26.4

Figure 2. Freeenergy diagram for the solvolysis of the 2-p-anisylcamphenilyl p nitrobenzoates;80%aqueous acetone at 25" (Brown, 1972).

188

G. M. KRAMER

On balance, the solvolysis results just described indicate that there is little reason to require the norbornyl ion reached under solvolytic conditions to be nonclassical. Before we turn to other aspects of the problem, it is interesting briefly to consider the results of deamination studies in the norbornyl system.

Deamination of 2-Norbornylamine A remarkably thorough study of the deamination of optically active exo- and endo-2-norbornylamines in acetic acid has been conducted by two groups of investigators, Berson and Remanick (1964) and Corey and coworkers (1963). Rather similar results were found, those of Berson and Remanick being listed in Table 2. Both the exo- and endo-reactants gave rise to a closely similar distribution of products. The major difference between the reactions being the retention of much more optical activity in the products of the endo-isomer. Both reactants yielded more than 95% exoderivatives and it is important to note that the exo-acetate obtained from the exo-amine retained 11% optical purity. TABLE 2 Deamination of Optically Active 2-Norbornylamines in Acetic Acid ~~

From exo-isomer

ProducP

Yield, %

Total acetate Total alcohol Total nitrate Total exo Total endo

87 10 2 f 0.5 97 ? 1 1.85 f 0.45

Acetate, % endo exo

98

Alcohol, % endo exo

Total products

'

2 f 0-5

1

99

Optical purity, %

From endo-isomer

Yield, %

Optical purity, %

80 16 4 95 48 -

11+2

4 7 + 0.5 95.3 5.3 0-5 94.7

85 ?r 12 18 ? 0.6

_+

13

-

21

Product distributions are based on the total recovered product (Berson and Remanick, 1964).

THE NORBORNYL CATION

189

The extent of retention of optical purity differentiates Corey’s results from those of Berson and Remanick. Corey found similar optical purity in the products of both epimers, from which it was concluded that a common ion was an intermediate in both reactions. Since optical purity could not be retained, if the symmetrical nonclassical ion was an intermediate or transition state in the reaction, he concluded that deamination led to the classical cation from both the exo- and endo-reactants. It is probably due to a difference in experimental procedure, that Berson and Remanick found substantially more retention in products obtained from the endo-isomer than from the exo. The optically active compounds again must be attributed to the intervention of classical ions. The varying extent of racemization has been a subject for considerable debate. The classical view is straightforward, namely that the products are all the results of kineticalIy controlled captures following initial formation of the classical ion. Racemization is viewed as a result of the Wagner-Meerwein rearrangement competing with solvent or anion capture of the ion. The increased yield of endo-acetate and retention obtained from the endo-epimer is primarily due to the reaction of tight ion pairs. The greater degree of racemization obtained from the exo-epimer is attributed to the lack of steric inhibition of the IVagner-Meerwein shift because of the remote location of the leaving group. Sargent (1972) has summarized the more flexible arguments from the non-classical view. Thus Berson and Remanick’s data are taken to indicate that all products in both reactions cannot arise from a common intermediate. One of the intermediates from either the exoor endo-isomer is assumed to be the classical ion. The nature of the other intermediate which leads to racemization is uncertain but the data could all be rationalized if it were the bridged ion. As we understand the argument, the data could be explained equally well on the preceding basis, namely that the extent of Wagner-Meerwein rearrangement occurring before the intermediate is trapped is identical with the proposed extent of bridged-ion formation. There is considerable uncertainty as to how to translate the results of deamination studies t o solvolysis experiments. Deamination normally proceeds with a substantially lower activation energy and is expected to produce a hotter or freer ion because of the greater exothermicity of the reaction. Whether this should aid or hinder bridging, ion pairing or solvent participation is not any clearer than

190

G. M. KRAMER

the distinction between this and the cold carbonium ions which are formed in solvolysis. The firmest conclusion that can be drawn with regard to the structural problem is that both the exo- and endo-amines yield a substantial amount of cIassical norbornyl ions during deamination. If stabilization by non-classical bridging is in competition with classical ion formation, one might deduce from the optical purity of the exo-acetate produced from the exo-amine that it amounts to less than 1-4kcalmole-’. On the other hand all the data could be reconciled without invoking the non-classical species provided that a small barrier is proposed for the Wagner-Meerwein shift.

Isotope Effects An excellent review of the use of secondary a and (3 isotope effects in the norbornyl system has recently been written by Scheppele (1972). We shall briefly summarize a few of the factors and conclusions arrived at because of their relevance to the interpretation of other solvolysis data, but the interested reader should turn to that article and the original sources quoted for a more extensive discussion. The kinetic isotope effect has its origin in force constant changes occurring at a n isotopically substituted position as the reactant is converted into an activated complex. Hence it provides information about the transition state in the solvolysis reaction, but not necessarily about the structure of possible intermediates. This limits the utility of information drawn from isotope studies in resolving the structure of ions under “stabilizing” conditions. A further limitation stems from a detailed examination of the possible reaction coordinate in an SN1 reaction. Consider a scheme where substrate A may solvolyse to form an intimate ion pair, B, that dissociates and then reacts with solvent. Shiner et al. (1969), Shiner and Dowd (1971) and Shiner and Fisher (1971) have suggested that maximum a-effects will be found where there is a rate-determining conversion of the intimate ion pair into a solvent-separated ion pair, and Murr and Donnelly (1970) have shown how partitioning of the ion pair, B + A and C, results in an isotope effect in the diphenylmethyl RX A

k1 k-1

R+XB

kl G

k -2

R+IIXC

k3

HOS

ROS

191

THE NORBORNYL CATION

benzoate system. Thus not only may the isotope effect be attributed to force constant changes in the initial ionization process but a substantial contribution can be made by the partitioning of the ion pairs. Realization of the latter process complicates the interpretation not only of isotope effect data but of the solvolysis rate information previously discussed. Let us now consider Scheppele's discussion of the solvolysis of 2-d-exo- and -endo-bicyclo [ 2.2.11 hept-2-yl p-nitrobenzoates and p-toluenesulfonates. A maximum 01 isotope effect of 1-22 is suggested for k H / i z D for the limiting solvolysis of simple secondary sulphonates (solvolyses where the rate is not affected by the nucleophilicity of the medium). The endo-esters yield a similar value which might be taken to imply the absence of steric hindrance in their reaction. However, whereas the maximum isotope effect for the simple sulphonates is believed to arise partly from an activated complex in the conversion of the internal ion pair into a solventseparated ion pair, B + C , the effect in the endo-norbornyl compounds is suggested to arise during the transition state involved in the formation of the intimate ion pair, A + B. The latter is deduced from the fact that polarimetric and titrimetric rate constants are the same in acetic acid, aqueous acetone and ethanol (Winstein and Trifan, 1949 and 1952). If this is the case, the endo-effect of 1-20 should be TABLE 3 &-Isotope Effects in the Solvolysis of Exo- and Endo-Norbornyl Compounds

Compound

4 X

X

Solvent

t("c)

kH/kDa

OBS OBs OBS OTS

CH3C02H 80% aq. C2HsOH C H J C O ~ HCH3C02K , CH3C02H

24.85 25 25

1.11 1.124 1.118 1.20

OBS OBs OBS OBs Br

CH3C02H Aq. dioxane HCOzH 80% aq. C2H50H Aq. C2H50H

50

1.203' 1.20' 1.22' 1-193 1-28

a Per atom of D. b Per 0.94 atom of D.

Lee and Wong, 1964. Mum and Conkling, 1970. Humski et al., 1970. f Lee and Wong, 1965. Schaefcr et al.. 1967.

50.20 25 25 60.21

Ref. C

d d e c

f f e

t?

192

G. M. KRAMER

compared with a normal a-effect for compounds solvolysing via a limiting mechanism with a negligible partitioning isotope effect. Just what this value is is not clear. However, Scheppele indicates that it should be considerably lower than 1-20and hence that the endo-data might be construed as being high, this therefore being evidence for steric hindrance in the ionization process. The reason for the large value in ethanolysis of the bromide is not understood and no interpretation is offered. These views are not universally held, one reason being that Schleyer has presented evidence that the solvolysis of simple tosylates and endo-norbornyl esters may not occur via a limiting mechanism (Fry et al., 1970 and accompanying papers). Just as intriguing are the studies on the exo-compounds. Two sets of data are available, those having a low value of 1-12 and the higher result of Humski, MalojEiC, BorEiC and Sunko, 1.20. The former value can be rationalized with a non-classical transition state in the A-B process but does not seem to require it. The larger value appears to be consistent with the presence of a normal ion-pair like transition state (B-C) involving the classical ion and is difficult to rationalize with the presence of a bridged ion. Isotope effect measurements have been conducted with labelling at the 6- and 3-positions as well as C-2. These results appear to be no more conclusive than those briefly discussed; Scheppele’s conclusion concerning y effects at C6 in solvolysis of exo-2-norbomyl brosylate being that they are (a) probably inconsistent with classical theory, but (b) not clearly consistent with non-classical theory. Although not resolving the non-classical issue, the studies indicate many difficulties which obscure the interpretation of any rate data. It has been suggested that, regardless of the transition state structure during solvolysis, the norbornyl ion might have yet another structure under the stabilizing influence of strong acid media. Before considering those systems it is appropriate to discuss the current theoretical status as provided by molecular orbital calculations.

4. THEORETICAL STATUS

A rigorous all-electron, non-empirical ab initio calculation ought to be able to define the bond lengths, angles and charge distribution and hence the structure of an unsolvated norbornyl cation. The most

193

THE NORBORNYL CATION

extensive attempt to resolve the problem in this manner to date appears to have been made by Goetz and Allen (1971). However, principally because of the size of the ion, we are far from having a complete potential energy surface for C7H:1. Goetz and Allen used Pople’s STO-3G and 4-31G basis sets which are representative of small and medium sets to obtain points on the potential energy surface of the ion (Hehre et al., 1969; Ditchfield et al., 1969). These sets had proved useful in computing the potential surface for the C3H; ion (Radom e t al., 1971) and it was assumed they would be valuable in describing the C-1: C-2: C-6 portion of the norbornyl ion.

The STO-3G set was employed in an initial geometry search and later supplemented by selected calculations with the 4-31G basis. According to Goetz and Allen, “The STO-3G basis appears to be deficient in its representations of the carbon atoms, leading to incorrect energy comparisons between geometries differing in their non-bonded C-C distances. The 4-31G basis set is a considerable improvement over the STO-3G basis set for energy comparisons due to its splitting of the valence orbitals into independent short- and long-range components. However, the 4-31G basis is less adequate than a split Hartree- Fock A 0 with polarization functions and therefore it must be recognized that our basis sets are definitely more limited than those which have frequently been employed for small polyatomics. ” Recognizing these limitations, we may summarize the results of the calculations (see Figure 3). The STO-3G calculation of the classical ion was made by fixing the bond length between C-1 and C-6at 1-59 the C-1:C-2bond length at 1-47 A and minimizing the energy by varying the C-1:C-2:C-6 bond angle. A minimum energy of -268.05150 Hartrees was found with a 96” bond angle.

a,

11540

1.539

Figure 3. Approximate geometries of the norbornyl ion derived from the STO-3G basis set (Goetz and Allen, 1971).

194

G. M. KRAMER

For the non-classical ion an energy minimum was deduced with a bond length of 1-40 A between C-1 and C-2 and 1.837 A between C-1 and C-6. The energy was -268.04377 Hartrees or about 4.3 kcal mole-' higher than the classical structure. The energy difference is expected to be reduced and perhaps eliminated when the calculations are repeated utilizing the 4-3 1G basis set because of Pople's experience with the C3H: ion. The non-classical form should be preferentially stabilized by the improved representation of the atomic orbitals, which should enable calculations to give a better account of the long bonds and three-centre ring. The interesting result of these considerations is that the inherent energy difference between the principal participants in the debate is likely to be very small. The potential energy surface may well be nearly flat with less than a kilocalorie separating the classical and non-classical ions. When the ion is placed in a solvent, specific solvation of one or the other structure may well become a stabilizing driving force. We now turn to a consideration of structural information that has been acquired by studies in strong acid media where the norbornyl ion is believed to exist as a relatively long lived intermediate.

5 . THE SEARCH FOR A PROTONATED CYCLOPROPYL RING

The presence of a protonated cyclopropyl ring distinguishes the two norbornyl ions and one might hope t o characterize a stabilized ion by its proton exchange behaviour. To do this one would have to investigate the exchange behaviour of protonated alkylcyclopropanes and tertiary and secondary cations in an acid sufficiently strong to stabilize all species for some time. Protonated cyclopropane has been considered to be a real intermediate since the work of Aboderin and Baird (1964), and protonated alkylcyclopropanes have also recently come to be considered intermediates as opposed to transition states in many rearrangements occurring in strong acid media where SbF, is the Lewis acid. Thus Brouwer and Oelderik (1968a) suggested that protonated methylcyclopropane is an intermediate in the isomerization of sec-butyl1-13C to sec-butyl-2-13C cations in the HF-SbF, system, and Saunders e t al. (1968) implicated this species in the rearrangement of

THE NORBORNYL CATION

195

sec-butyl cations in SbFs -SO2ClF. Similar species have become the subject of recent review articles (Collins, 1969; Saunders et al., 1973). The question of whether protonated cyclopropanes are intermediates during the rearrangement of alkyl ions in SbF, -HS03F or whether they should be considered as transition states in a process leading to the formation of a solvated proton and an alkylcyclopropane is central t o the interpretation of proton exchange reactions in this medium. Brouwer (196813) has provided a thermodynamic argument indicating that the protonated species are intermediates, and we are inclined to accept this result. The following discussion, which is based on this assumption, enables one to distinguish between the edge and corner protonated cyclopropanes which may be formed from alkyl systems in SbFS-HS03 F. It also enables one t o draw a conclusion as to the most reasonable structure of the solvated norbornyl ion. If the assumption is invalid, structural distinctions within the alkyl ions cannot be drawn but one can still provide an estimate of the structure of the norbornyl cation. Tracer studies have shown that during the rearrangements of butyl, amyl and hexyl cations in SbF, -HS03 F an intermediate often forms which exchanges one proton with the acid (Kramer, 1970, 1973). Because control experiments showed that the parent tertiary or secondary ions would not exchange unless formed in the back reaction of the same intermediate leading t o products, and because multiple exchange did not occur, the intermediate is considered t o be an edge-protonated cyclopropane [ 211 rather than the corner protonated compound [ 221. A more detailed discussion of these studies is

presented on page 196.

HA

H

H

H H

The norbornyl system has been subjected to an analogous series of experiments with the results shown in Table 4. When exo-2-chloronorbornane was solvolysed in the presence of a hydride donor good yields of norbornane were obtained. The product had not acquired any protons from the acid and thus an edge-protonated intermediate had not formed before reaction with the hydride donor occurred.

G. M. KRAMER

196

TABLE 4 Trapping the Norbornyl Ion by Hydride Transfer from Methylcyclopentane in 2M SbFS/H(T)S03F. ([MCP]/[norbornyl ion] = 40) Procedure

% HzO

A A A B B

2 5 5 5 5

t("C) -50 -50 -78 -50 -78

Norbornane, %

60 100 100 90 50

Protons exchanged

0.0 1 0.01 Trace 0.08 0.08

A-The ion was trapped by solvolysing in the presence of methylcyclopentane. B-Methylcyclopentane was added 10 minutes after ionization.

Since in the alkyl series a corner-protonated intermediate is less stable than an edge-protonated one, the data also suggest that the former is not formed during the reaction. The norbornyl system is remarkable in that proton exchange is extremely slow, even when the ion is left in the acid for extended times before being quenched. During this period at -5OOC it undergoes many intramolecular rearrangements without forming a species that will deprotonate. This behaviour is clearly consistent with the presence of classical ions equilibrating by Wagner-Meerwein and 2,6 hydride shifts. It is inconsistent with the presence of either an edge- or corner-protonated species. In order t o rationalize the data from a non-classical point of view one would have to propose that the relative stability of the edge- and corner-protonated species is different in norbornyl than in alkyl ions and that they are far apart in energy or that only the latter is an intermediate on the energy surface. There are no obvious reasons to believe that this is the case so that the exchange studies favour the classical ion.

The Control System The behaviour of alkyl cations in SbF, solutions was used to establish the basic information needed to interpret the exchange studies with the norbornyl ion. The studies were conducted with butyl, amyl and hexyl ions but only the results with the C4 system are discussed here.

THE NORBORNYL CATION

197

Exchange studies were carried out by solvolysing a series of butyl chlorides in ZM solutions of antimony pentafluoride in fluorosulfonic acid at -50" and -78". The acid contained tracer levels of T,O and small amounts of water to provide sufficient nucleophiles t o catalyse proton exchange reactions with some of the intermediates formed in the butyl system (Kramer, 1970, 1973). TABLE 5 Kinetically Controlled Product Distributions in 2M SbF5 / H ( T ) S 0 3 F Reactant

.

t("C) H 2 0 , % Conv.'

i-c4H10, %

Hexb n-C4H10, % Hex

Me3 C C1

-50

c

100

100

0.2

MezCH.CH2C1

-50

2

18

87


MeCH2.CHCIMe -50

2

90

-

-78

5

MeCHz CHClMe -78

5

Me2CH.CH2Cl

.

88 71

-

13 100

<0.1

12 100

0.9


a The conversions are approximations, based on the detection of unconverted reactants in the products. Material balances are not well known. b Hex is the average number of protons exchanged in the trapped isomer. Commercial HS03F was used.

Solvolysis was in some cases carried out in the presence of an excess of methylcyclopentane to trap intermediate ions as fast as possible via hydride transfer. In other experiments the hydride donor was added ten minutes after solvolysis. The results of a series of exchange experiments on the butyl system are shown in Table 5. Solvolysis of t-butyl chloride in the presence of methylcyclopentane yields a t-butyl ion which is quantitatkely converted to isobutane before exchanging hardly any protons with the acid. Secondary butyl chloride yields n-butane which has not exchanged a proton with the acid if 2% water is present. The same reaction with 5% water in the acid yields partly exchanged butane at -78O. The reaction of isobutyl chloride at -50' with 2% water present leads t o a mixture of 13% n-butane which has exchanged 0.9 protons and 87% isobutane which essentially has not exchanged any. With 5% water, unexchanged isobutane is again recovered, but the n-butane has exchanged nearly two protons with the acid, even though the reaction wLs run at -78".

198

G. M. KRAMER

The behaviour of the butyl system provides important information on the nature of the intermediate formed during the rearrangement of the isobutyl to the 2-butyl cation. Thus, from the observation that isobutyl chloride yields n-butane which has exchanged one proton with the acid, while the solvolysis of 2-butyl chloride in the same acid (2% H, 0),yields unexchanged n-butane one might deduce that an intermediate was formed during the former’s solvolysis which exchanged one proton with the acid before it converted to a secondary butyl ion. A reasonable mechanism is shown in Scheme 1.

Scheme 1

There it is proposed that isobutyl chloride first solvolyses to a “primary” ion. This ion undergoes two reactions, mainly a 1,2 hydride shift to form a t-butyl ion, which ultimately yields unexchanged isobutane, and to a lesser extent cyclization to form a protonated methylcyclopropyl species [ 231 . This species exchanges one proton with the acid and then undeigoes a cleavage reaction yielding the 2-butyl ion. When these reactions were carried out with the acid containing more water and hence more nucleophiles, slightly more exchange was found in the n-butane recovered from each of these alkyl halides because the nucleophiles increased the rate of exchange of the 2-butyl ion. Protonated methylcyclopropane [ 231 might be a corner-, edge-, or face-protonated compound, but, for our purpose, it is sufficient to neglect the distinction between the last two. On the basis of the exchange information, we can make a choice between the comer and edge species. Thus, if the intermediate is edge-protonated, it would be reasonable to expect the “edge” proton t o exchange with the acid before it added to a carbon atom and induced the cleavage of a C-C bond. In this case, the product would have acquired one proton from the acid.

THE NORBORNYL CATION

199

If the intermediate were corner-protonated, however, and undergoing exchange with the acid, it would be most reasonable to expect all the protons at the comer t o exchange with the acid before ring cleavage occurred. Thus, corner protonated methylcyclopropane would be expected to exchange three protons with the acid if it exchanged any. The fact that only one proton exchanges therefore indicates that [23] is an edge protonated compound. The behaviour of the butyl system is typical of the pentyl system and hexyl system in all of which exchange studies indicate that edge-protonated cyclopropanes form during many ionic rearrangements in this acid (again with the assumption that a protonated cyclopropane is a bona fide intermediate) (Kramer, 1970). The combined studies on the norbornyl and butyl ions lead to the strong suggestion that norbomyl is classical. This conclusion would also be reached if the protonated cyclopropanes in the alkyl series were only regarded as transition states in the exchange reaction. In that case one would expect the non-classical norbornyl intermediate to be a participant in deprotonation reactions when placed in the same media where the alkyl species could not be stabilized. The absence of exchange again is evidence against the presence of a protonated cyclopropyl ring in the norbornyl ion and indicates that the solvated species is classical. This result is contrary to what has been deduced from other studies of the norbomyl ion in strong acid media. Accordingly, we now turn to a reappraisal of the available ESCA, H-nmr, C-nmr, kinetic, Raman, quenching and isotope studies that have been brought to bear on the structural problem.

6. ESCA The C-1s x-ray electron spectroscopic spectra of the norbornyl ion has recently been reported (Olah et al., 1972). In principle, ESCA appears to provide an ideal method to decide upon the ion’s structure, as one would expect a classical ion to exhibit two peaks with an area ratio of 6: 1 and the non-classical ion to show two peaks with a 5:2 ratio. In preliminary experiments, it was reported that the t-butyl and cyclopentyl cations exhibited ESCA spectra wherein the peak representing the cationic centre was shifted about 4 eV to the high

G. M. KRAMER

200

binding energy side of a single band representing the other carbon atoms, the shift being somewhat larger in the secondary than the tertiary system. The norbornyl spectrum shows a much smaller shift, estimated as 1-7eV, and this, together with the fact that the area ratio was reported as 5:2, suggested that the non-classical ion was being observed. This decision, however, appears to be based on a miscalculation. A closer inspection of the published spectra leads to the conclusion that the area ratio is much closer to 6:l than 5:2 and therefore it cannot represent the non-classical ion. Thus, if the spectrum is that of a norbornyl cation it must be that of the classical ion and would therefore be consistent with the implications of the exchange study just described. The area ratio has been measured in two ways: first, after enlarging the published spectra and then resolving the peaks by a geometrical construction, and second, with the aid of a Du Pont 310 curve resolver using Gaussian peaks. By construction, we obtain a peak ratio slightly larger than 6/1.With the curve resolver we find 82 k 1% of the area in the large band and 18 k 1% in the other. This ratio is considerably closer to 6/1 than 2.5/1. Another reason for believing that ESCA is revealing the classical ion is obtained if one considers the band shape to be expected for each ion. Thus, even though the peaks may be poorly resolved, the experimental shape is closely related to that corresponding to the superposition of 3 bands of areas 4:2:1 (in order of increasing binding energy), as is expected for the classical ion. The non-classical ion might also fit the spectrum, but it would be more fortuitous because here one would expect the overlap of four bands with relative areas of 2:2:1:2 at increasing binding energies. C

4:2:1

2:2:1:2

While the area ratio apparently eliminates the non-classical ion, one must ask whether the small binding energy shift observed can be accounted for if the ion is classical. This might be possible because the shift is dependent on a large number of factors which differ

THE NORBORNYL CATION

201

markedly between the classical norbomyl ion and the model US$ (cyclopentyl). However, recent theoretical calculations by Allen an Goetz (1974)’indicate that the expected shift is larger than observec and the spectra therefore appear to be incompatible with both ions. The qualitative arguments which might have rationalized the spectra and the classical ion are given below. Thus, the ESCA shift is normally assumed to be influenced by the electrostatic environment through which the electron ejected from the C-1s orbital is passing. The environment about the Cz atom in the classical norbomyl ion is very different from its counterpart in the cyclopentyl ion for several reasons. First, the extent of rehybridization in the C-2-H bond should be different because of strain considerations. Second, neighbouring group interactions are quite different, as can be seen by observing Dreiding models. In the classical norbornyl ion, the 1-6 bond and vacant p-orbital are well disposed for some C-C hyperconjugative interaction, whereas G H - p orbital interactions predominate in cyclopentyl. The norbomyl ion is also prone t o exhibit a charge-induced dipole interaction between C2 and the C1-C6 bond which is unmatched in the cyclopentyl ion. If these classical interactions occur, and if C-C hyperconjugation is more important than C-H, as theoretical models currently indicate, it is clear that the charge density and electrostatic situation in the respective ions are riot comparable. The differences which exist all appear t o be operating to reduce the positive charge density at C2 in the norbomyl ion. Since it is believed that the charge density at such a site is already low in classical ions such as t-butyl, cyclopentyl, and isopropyl, a further reduction due t o the unique interactions in the norbomyl ion might be expected to result in a large percentage change, as is apparently observed by ESCA. Thus, both the area ratio and the binding energy shifts might be rationalized with the classical ion by qualitative arguments. As the Allen and Goetz study apparently negates these arguments it leaves us with the possibility that the “norbornyl” spectrum is that of a derivative or mixture of derivatives rather than that of a stable cation. At this time many derivatives like alkylated solvent species and dialkylhalonium ions seem possible. The proposition that a mixture of species is actually present has in our view been reinforced by a recent report which presents a second “norbornyl” ESCA spectrum, Olah et al. (1973). Although it is claimed that both spectra are identical, a careful comparison indicates that they are different, the most noticeable change being the apparent growth of a new high

202

G. M. KRAMER

binding energy band on the shoulder of what was supposedly the cationic band. If the area of this band is deleted the remaining spectra again yields a 6: 1 area ratio. Of more importance is the fact that there now exist two different spectra for the "same" species. This coupled with the fact that the spectra have the wrong shape for the bridged ion, makes it clear that ESCA has not resolved the structural problems.

7. 13C-NMR Proton and carbon 13 nmr studies and Raman spectral investigations have provided important information about the norbornyl ion. These have been offered as proof of the non-classical structure. The principal arguments are that the Raman spectrum proves the existence of a cyclopropyl ring, that the 13C chemical shifts are incompatible with the existence of a pair of equilibrating classical ions, that the H-nmr shows that a corner-protonated rather than an edge-protonated nortricyclene exists and that several quenching results are consistent with the non-classical ion. This interpretation of the data is plausible but not the only one possible. Indeed, the postulate of equilibrating ions would again appear to be at least as consistent with these observations as the non-classical structure. C information. The C spectrum of Let us first consider the the norbornyl ion was obtained in SbF5-S0, at -70". The spectral data for it and reference compounds are listed in Table 6. Of particular interest is the shift of C-2, 101.8 p.p.m. in the norbornyl cation at -70°, and the separation of this peak into components at +70 and +173 p.p.m. at -154". Table 7 contains C shift information for a number of other compounds that have been used to analyse the norbornyl spectra. Included among these are data for the isopropyl ion which has been used as the principal model for the secondary ion, and for the 1,2-dimethylnorbomyl cation which has been reported to be an example of an equilibrating classical cation, the Wagner-Meerwein shift being too fast to be stopped at -140". In the table it may also be noticed that the t-butyl cationic centre experiences a shift of about 11 p.p.m. on being transferred from neat SbFs to SbFs -SO,. The observed I3C chemical shifts of carbon atoms in the norbornyl ion are said t o rule out the presence of equilibrating ions. In arriving

203

THE NORBORNYL CATION

TABLE 6 13C and 'H Nuclear Magnetic Resonance Chemical Shifts and Coupling Constants in the Norbornyl Cation, Cyclopentyl Cation, Nortricyclene, and Norbornanee

F&

3

Solvent

&(13c)a

&('HIb

ICHC

SbFs-SO2 -70"

C-2 + 101.8 C-3 + 162.5 C-4 + 156.1

5-01 1-86 2-82

53.3 140.2' 153

S ~ F S - S O ~ C I F - S O ~ FC-1 ~ + 173 -154" c-2+ 70

3-05 6.59

d d

c-2 + 95-4

4.48

C-2 + 183.6 C-3 + 161.8 C-4 + 165.2

0-99 1.19 1.89

C-2 + 161.2 c-7 + 155.4 C-1 + 157.6

1.30 2.10 2.18

+--.

. I '

t

2

SbFs-SO2ClF -70"

cc14 35O

28.5

175 133 148

*

(130 2) (130 140 f 2)

2

a In parts per million from

13cs2.

b In parts per million (6)from internal TMS.

In Hertz, Was not obtainable in natural carbon abundance samples. Olah et al., 1964, 1970, 1971.

at this conclusion, it has been assumed that the 3 C shift reflects charge density at carbon atoms in the framework and that the charged carbon of the isopropyl cation is a model for the ionic centre in the classical structure. Inspection of the 1 3 C shift in Tables 6 and 7, however, raises immediate doubt of the quantitative relation between charge density and chemical shift.2 Thus, whereas an uncharged carbon appears near +165 p.p.m. (us. CS2) and the tertiary carbon of the t-butyl cation at -135.4 p.p.m. the cationic centre of the isopropyl ion is found at A more wide-ranging survey of chemical shift-charge density correlations in organic ions is given in the preceding chapter by D. G. Farnum (Editor).

G. M. KRAMER

204

TABLE 7 13C-nmr Chemical Shifts in Alkyl Cations, Norbornyl Cations and Aryl CationsO

I on

613C

Solvent

Temp., OC

+

C’

CH3

c-1

-139.2

c-1 c-2 c-1

-135.4 +146*3 -146.5

SbFS-SO2CIF SbFS-SO2 SbF5, neatb

-20

c-1 c-2

-125.0 +132.8

Sb Fs-SO2 C1F

-20

c-1

-86.8

SOz-Sb F 5

-60

c-1

-61.1

c-1 c-2

+118-0 -76.1

FS03 H-SbF 5-SO2

-80

c-1 c-2

+130-8 -65.8

FS03H

-35

C-1, C-2

+26

-70

C-G

+147

-140

-60

+30

I

CH3

@ C’, ( CH3 I)

Olah et al., 1970. Olah et al., 1964. Olah etal., 1971.

-60

THE NORBORNYL CATION

205

-125-0 p.p.m. This implies that the positive charge density at the secondary site is lower than it is at the tertiary carbon in the t-butyl ion, a notion in qualitative discord with the fact that tertiary alkyl ions are more stable than secondary ions. This is usually attributed to the ability of alkyl groups to permit the flow of electrons toward the cationic site, thus reducing its positive charge and simultaneously spreading some charge density throughout the rest of the ion. Whether this flow is due to an inductive release through the sigma bonds or to a manifestation of hyperconjugation, the result is the same; less positive charge density is to be expected at a tertiary than at a secondary cationic centre. Accordingly, if the C shift quantitatively reflects charge density, one would expect the isopropyl centre to be significantly downfield of the tertiary in t-butyl. The fact that it is not suggests that the C shift in this ion cannot provide a general model for secondary cations and its use to predict the shifts of the classical norbomyl ion is highly dubious. This serious limitation of the utility of 3C-nmr has been recognized (Olah et al., 1970), but dismissed by recourse to the results of extended Huckel M.O. calculations by Hoffmann (1964). These calculations indicated that the charge on the cationic centre of t-butyl is, in fact, more positive than in the isopropyl ion. Hoffmann noted that his results are at odds with the widely held assumption that a methyl group is a better electron donor than hydrogen but showed that the correct order of carbonium ion stability could be predicted even if a methyl group is electron withdrawing with respect to hydrogen. It is difficult to assess Hoffmann’s conclusions, particularly in view of more recent M.O. calculations dealing with systems similar to his but sometimes with varied agreement. For example, in considering the relative stability of various structures for the ethyl cation, Hoffmann concluded that the essentially classical CH,CHt ion is slightly more stable than a proton-bridged ethylene and contains the charge distribution shown in Figure 4a. Pfeiffer and Jewett (1970), however, have made ab initio calculations on the ethyl cation and report the charge distributions in Figure 4b for the “most” stable ethyl ion. Their calculations agree with Hoffmann’s in predicting that the classical ethyl structure is more stable than a bridged structure, but their calculated charge distribution is entirely different. Other semi-empirical and ab initio calculations also lead to

’

’

206

G. M. KRAMER

divergent conclusions about the charge distribution in the ethyl cation. Thus, an NDDO treatment (Sustmann et al., 1969) leads to a charge of +0-294 at the carbon atom of the CH2 group, while a recent ab initio calculation finds a charge of +0-249 at the same centre (Radom e t al., 1972). These are to be compared with the value of -0-339 found by Pfeiffer and Jewett. Radom, Pople and

Figure 4. Calculated charge densities in the ethyl cation. (a) Hoffmann, 1964; (b) Pfeiffer and Jewett, 1970; (c) Sustmann et al., 1969.

Schleyer (1972) found that their calculations also support the notion that the tertiary carbon in the t-butyl cation is more positively charged than the secondary carbon in the isopropyl iorr; but this conclusion and the enormous discrepancy with Pfeiffer and Jewett are not readily understood. Another problem concerning the use of I 3 C shifts in model ions to estimate peak positions in unknowns stems from the fact that there is a solvent dependence on the shift of the t-butyl ion, C1 moving from -146.5 in neat SbF, to -135.4 in SbF, -SO2 CIF. This shift might be due to a change in shielding caused by bulk solvent properties or it might be due to rapid reversible alkylation of the solvent by the cation. If the latter is a factor, it would be more important in affecting shifts of secondary ions. To date, little information bearing on this point is available. Still another barrier to employing the 3 C shifts to resolve the norbornyl problem stems from recent studies of H. C. Brown and Peters (1973). These authors have measured the relative rates of solvolysis of 2-methyl-, 2-t-butyl-, 2-phenyl- and 2-cyclopropyl2-propyl p-nitrobenzoate in 80% aqueous acetone to gather information about the electron-releasing ability of these groups. Upon comparing their results with the 3 C shifts of the carbonium carbon of the same cation, they found no correlation. Although it might be

207

THE NORBORNYL CATION

argued that factors stabilizing a transition state during solvolysis may differ from those stabilizing the ion under “stable” ion conditions, equilibria between triphenylmethyl or tricyclopropylmethyl cations and their carbinols in H2SO4 determined by Den0 and coworkers (1965) appear to correlate well with the implications of Brown’s solvolysis studies. Therefore, the quantitative use of C shifts determined for one ion as a model for another might be questionable. In spite of these difficulties, the shifts of the isopropyl ion have been used to “prove” that the equilibrating classical ion could not give rise to either the -150’ or -70’ 3 C spectrum actually observed. Thus, taking -125 p.p.m. for the shift of the secondary site and +158 p.p.m. for that of the bridgehead carbon leads to an estimate of +16.5 p.p.m. for the equilibrating carbons at -150’. This is substantially less than the +70 shift observed. A similar shortcoming is found upon analysing the -70’ data where three equilibrating carbons lead to an estimated shift of 50 p.p.m. in the classical ion that is much less than the 101.8 p.p.m. shift actually found. Now, although the classical model does not appear to fit these estimates, it should be noted that non-classical ion would lead to exactly the same poor estimate, as long as the same amount of positive charge is distributed between C-1and C-2in both models. In fact, n o estimate of the C shifts for the non-classicah structure have been advanced; but, if they were and if the chemical shift is linearly proportional to charge density which has been assumed, then structure [ 81 with half a unit charge at each of the bridged C-atoms of the cyclopropyl ring must lead to the same deficiency as equilibrating ions [4]and [ 5 ] . Hence, one must conclude that either C shift is significantly influenced by factors other than charge, the or that there is substantial leakage of charge out of the C-1:C-2: C-6 portion of the ion, irrespective of its structure. One might also note that structure [ 7 J which had originally been proposed as a model for the non-classical ion is incompatible with the C shifts, C-6offering no indication of the presence of a sizeable positive charge. The difficulty of using isopropyl as a model for the norbornyl cation can be illustrated by attempting to use the model for other equilibrating secondary ions and t-butyl for equilibrating tertiary cations on a somewhat different basis from what has thus far been employed. The experimental change in C shifts at the cationic site and neighbouring carbon atoms upon converting a model hydrocarbon to a carbonium ion is taken as typical of the shifts t o be found in equilibrating ions. The shifts on converting isobutane t o the

’

’

’

G. M. KRAMER

208

t-butyl ion and propane to the isopropyl ion are chosen as models for tertiary and secondary systems, viz., Model Systems

c-c-c

+

c-6-c

C

C

I

I

c-c-c+c-c-c +

AI3C -303, +C -45, CH3

-304, +C -23, CH3

The shift at a neighbouring group is assumed to be constant irrespec-

I

tive of whether the group is CH,, CH, , CH or -C-.

I

In order to make the calculation, the 13C shift of the parent hydrocarbon must be known or estimated. This process has been applied to a series of compounds with the results listed in Table 8. The tabulation shows that the t-butyI system is a reasonable model for some equilibrating ions. It fails badly, however, when applied t o the norbornyl compounds. The isopropyl system is a poor model for sec-butyl and cyclopentyl ions and is a very poor model for the norbornyl cation. The failure of the models to provide reasonable estimates of the shifts in the tertiary norbornyl cations which are undergoing either Wagner-Meerwein shifts or hydride migration makes it clear that the experimental shifts in the secondary system cannot be used as structural proofs. Rather they should be regarded as fascinating results t o be rationalized in terms of the structure, whatever it may be. C chemical shift information cannot be For these reasons the used unambiguously to define the structure of the ion. The data are important in indicating that positive charge is somehow highly spread out about the ion, but a multitude of factors probably lead t o the observed shifts. Cheney and Grant (1967) indicate that the magnetic shielding responsible for the observed shift is composed of three terms; dd),a diamagnetic term induced by the magnetic field at the atom being observed, d p ) , a paramagnetic term originating in the intrinsic angular momentum of non-spherical orbitals centred on the atom, and d"),a term including screening contributions from other atoms in the molecule. The major factor is likely to be the paramagnetic term, the magnitude of which will reflect variations in local electronic structure, polarization and non-bonded interactions. Without accounting for this term one is unable to make theoretical

209

THE NORBORNYL CATION

TABLE 8 Use of Models for Predicting Shifts of Equilibrating Ions Equilibrating Ion 6, PPm

Difference (Obs-Calc) x NO. of equivalent C+centres

Parent 6I3CR

Calc

C-2 156

-7.5

-3.4

41x 2

C-2 153 C-3 150

-12.0

-11.5'

05x 2

C-1 139

-24.5

-24

05x 2

C-1 147 C-2 153

-13.5

26h

39-5 x 2

C-2 169

-5.0

21.2'

26 x 2

C-1 167

88

95.41

7x5

C-1 156 C-2 162

-15

7@

85x 2

Obs

%b

Model:

A

+ --t

\

b

0' c

'Shifts us. CS2, p.p.m.

Estimated (?2p.p.m.) by the proceduie of Savitsky and Namikawa, 1964. Pekhk et al., 1970. 'Paul and Grant, 1963. Burke and Lauterbur, 1964.

f Olah et aL, 1974. Olah and White, 1967; Olah et al., 1970; Olah et al., 1973. Olah and Liang, 1974. f Olah and White, 1969b. J Olah and White, 1969a

\

G. M. KRAMER

210

estimates of C shifts in alkanes and it seems likely that this is also the source of difficulty in utilizing the isopropyl or cyclopentyl ions as models for the cation. In Table 6, J(' 3C-H) spin coupling constants are reported for the norbornyl ion. It has been noted (Olah et al., 1970) that the distinction between the classical and non-classical ion cannot be TABLE 9 Estimated Spin Coupling Constants, J( 13C1,2,6-H) in the INorbornyl Cation

H H

J?

=

H

H

H

J(13C-H)

42

3(175) 12

J(13C-H) = -= 43.9 Hz

H

& H

2(130) + l(140) + l(169) = 47-4Hz 12

J ( I3C-H) =

2(125) + 2(175) = 50.0 Hz 12

(Experimental value: J(13C-H) = 53.3 Hz)

made on the basis of these values, but it is still interesting to consider the data. In Table 9, estimates of the average coupling constants for I(' 3C1,2,6-H) in the classical ion and both the edge- and cornerprotonated non-classical ions are given. In making these estimates, the following models were used. For the classical ion, coupling in the CH2 group, 130 Hz, the bridgehead C-H, 140 Hz, and secondary C-H, 169 Hz, were estimated from data on norbomane and the isopropyl cation. For edge-protonated nortricyclene, it was assumed that three protons are coupled to the cyclopropane ring as in the parent molecule and the fourth proton is not coupled at all. For corner-protonated nortricyclene, it was assumed that the coupling in the bridging CH2 group is 125 Hz and 175 Hz in the C-H group, as in nortricyclene. Most of these assignments have previously been made by Olah and White (1969).

THE NORBORNYL CATION

211

It is clear from the data that the spin coupling constants for the three models are all in reasonable agreement with observation, as already indicated. That there is any agreement whatsoever, is perhaps just as surprising as the fact that all models are reasonable. This is because theory indicates that spin coupling arises as a result of a number of complex factors, an important term being derived from transitions between the singlet ground and excited triplet state of the molecule being studied. Since little is known of the excited triplet state of the norbornyl cation or the reference molecules and ions, the agreement indicated in Table 9 is perhaps fortuitous.

8. 'H-NMR The proton nmr behaviour of the norbomyl ion provides a wealth of information which, however, also appears t o be of limited value in the structural problem. The cation has been observed by 'H-nmr in many solutions containing SbF, as well as in GaBr,-SOz (Jensep and Beck, 1966). At -80" in SbF5-SO2 it exhibits 3 peaks, 61.86 (area 6), 2.82 (area 1) and 5-01 (area 4). The assignments and experimental coupling constants are shown in Figure 5 (Olah et al., 1970). The 5-01 p.p.m. peak indicates equivalence of the four HA protons, which is caused by rapid 6,1,2-hydride shifts. At lower temperatures, the spectrum of the HA protons changes. Between -130 and -154' this peak separates into two equal but broad components at 63-05 and 6.59. The methylene spectrum at 1-86 p.p.m. also separates and exhibits a broad high-field shoulder at 1-70. These shifts are all measured relative t o the Hc band at 2-82 p.p.m., which is assumed t o remain constant. Above -80', the spectrum of the norbornyl ion coalesces as a result of the increasing rate of the 3,2-hydnde shift. This, coupled

Figure 5. PMR characteristics of the norbomyl ion.

212

G. M. KRAMER

with lY2,6-hydrideshifts and the Wagner- Meerwein rearrangement, equilibrates all protons and results in a singlet at ambient temperatures. A multi-site shape analysis of the data indicates that the 3,2and 6,1,2-hydride shifts has the following rate constants (Saunders et al., 1964): k3,2 = k1.2.6

101 2.3

e- 10 , 8 0 O / RT

- 1012.1 e-5900/RT -

These spectral changes are consistent with either (a) freezing out a stable non-classical ion, or (b) the presence of the classical ion in which the 6,2-hydride shift has been slowed down, but the WagnerMeerwein rearrangement is still exceedingly rapid. If the ion is non-classical, the spectrum indicates that it must be a comerprotonated nortricyclene rather than the edge-protonated ion because the latter would have split into a 1:2: 1 pattern. The proposal that the -154' spectra is consistent with the non-classical ion suffers in our view from results with the 1,2dimethylnorbornyl cation (Olah et al., 1971). This cation was found

t o be a classical equilibrating ion by a series of C, H and Raman studies similar to those used t o characterize the norbomyl ion. In this case, the Raman spectra provided n o evidence of a cyclopropyl ring (still to be discussed with respect t o norbornyl), and C-1 and C-2 C shift at -140' in fair agreement with that exhibited an average estimated by assuming the coalescence of the tertiary and bridgehead carbons of the frozen 2-methylnorbornyl cation. It was therefore concluded that the Wagner-Meerwein rearrangement between tertiary ionic centres was occurring too fast t o be frozen out at -140'C. Since tertiary ion [24] in its classical form must be more stable than the classical secondary norbornyl cation [6] and since the barrier t o the Wagner-Meerwein rearrangement is negligible in [24] it is difficult t o argue that the Wagner-Meerwein

THE NORBORNYL CATION

213

rearrangement in [6] could have anything but a still lower barrier. Hence, if norbornyl is classical the Wagner- Meerwein rearrangement ought to be occumng at -154O. Sorensen and Ranganayakulu (1972) have supported the view that rapid Wagner- Meerwein rearrangement is occurring in the 1,2dimethyl-2-norbornyl cation by an independent nmr study of the 3or 7-dimethyl substituted analogue of this cation. In that case, averaged chemical shifts were obtained which were temperaturedependent, indicating that equilibration between the non-equivalent structures of the ion was occurring a t temperatures as low as -1 22.5'. The temperature dependence of the H-nmr spectra of the norbomyl ion has been subjected t o a line shape analysis t o provide an activation energy of 10-8 kcal mole-' for the 3,2-hydride shift and 5.9 kcal mole-' for the 6,1,2-hydride shift. It has been argued that the energy required for the 3,2 shift is substantially higher than a 2-4 kcal mole-' estimate for 1,2-hydride shifts around the cyclopentyl cation, the 7-9 kcal mole-' difference being attributed to stabilization of the norbornyl system by the formation of the comer protonated non-classical structure (Olah 8 al., 1969). From the classical point of view, a reasonable case can be made for the slow 3,2-hydride shift being at least partly due to poor alignment of the p-orbital of the secondary ion and the sp3 orbital of the exo-C-H bond involved in the hydride transfer. The fact that orbital orientation can play a significant role in 1,2-hydride and alkyl migrations is well established, although a quantitative relationship between the activation energy for migration and the dihedral angle between the orbitals is not available (Brouwer and Hogeveen, 1970; Schleyer et al., 1970; and Majerski et al., 1970). A striking manifestation of this effect is found in secondary adamantyl cations which change into tertiary ions by intermolecular hydride transfer rather than by intramolecular 1,2-shifts because of orbital alignment. This effect is likely t o be most important in constrained or inflexible systems such as would be expected in the classical norbornyl ion. Of particular importance are recent results on the energetics of the exo-2,3-hydride shift in the 2,3-dimethyl-2-norbornyl and the 1,2,3,4-tetramethyl-2-norbornylcations. Line broadening studies reveal a free energy of activation of 6-6 and 7.3 kcal mole-' for this process in the respective ions (Huang e t al., 1973;Jones et al., 1974). Thus the exo-2,3-shift in the tertiary norbornyl ions has been

214

G. M. KRAMER

found to be much slower than the analogous reaction in alkyl cations and favoured by 3 to 4 kcal mole-' over the reaction in the parent norbornyl ion. The latter difference might be rationalized by noting that a steric interaction exists between the endo-3-methyl group and the endo-5-hydrogen atom will tend to rotate C-3 and help orient the exo-3-hydrogen for the hydride shift in the tertiary ions. Since the interaction is missing in the norbornyl cation, a higher activation energy would be expected. Another result of the kinetic analysis of the 3,2- and 1,2,6-hydride shifts which tends to support the classical view of the ion is the close similarity in the A-factors in the Arrhenius equation. A is 10' for the 3,Z-shift and 10' 2 * 7 for the 1,2,6-equilibration process. Both factors are roughly normal and might be expected for rearrangements where the rate-determining step is a relatively slow hydride transfer between carbon atoms in the classical ion. On the other hand, if the non-classical ion is a stable intermediate, the transition state for the 3,Z hydride shift requires a substantial reorganization, including the cleavage of the cyclopropyl ring, and, by analogy with unimolecular gas phase processes, a much higher pre-exponential factor might be expected. [In the cyclopropanepropylene reaction log A is 15.17 (Chambers and Kistiakowsky, 1954).] Contrary to expectation, the observed pre-exponential for the 3,2-shift is actually a little lower than for the 1,2,6-equilibration process. In summary, the temperature dependence of the proton nmr spectra does not provide a sufficient basis t o decide upon the norbornyl cation structure. The slow 3,Z-hydride shift might be rationalized by either a non-classical or classical interpretation, but both the activation energies and the pre-exponential factors might be influenced by solvation. If this is the case the arguments based on the pmr observations become still less persuasive. Before we turn to other investigations of the stabilized ion it is interesting briefly to compare the rates of hydride shifts found in the strong acid systems with rates obtained under solvolytic conditions. Collins and Lietzke (1967) and Berson et al. (1967) have deduced the rates of 3,2- and 6,Z-hydrogen migration relative to the rate of solvent capture for norbomyl ions in acetic acid, formic acid and other common solvents from elaborate C scrambling studies due to Roberts and Lee (1951) and Roberts et al., (1954) and tritium labelling studies of Lee and Lam (1966). From the ratio of these

*.'

'

THE NORBORNYL CATION

215

rates one finds that the 6,2-shift is roughly 10 to 100 times faster than the 3,2-shift in acetic acid or formic acid. On the other hand, use of the kinetic expressions for the nmr data indicates that the 6,2-shift is about 8000 times faster than the 3,2 under stable ion conditions at 25°C. The discrepancy between the rates of the intramolecular shifts points strongly t o the presence of controlling solvation effects which, as indicated, restrict the conclusions to be gained from these results.

9. RAMAN SPECTRA Next, one may consider the implications of laser-Raman studies of the norbornyl ion. Raman spectra of the norbornyl cation at -70" to -80" were reported in the F S 0 3H-SbF, -SOz system (Olah et al., 1968). The spectra of it and related compounds are given in Table 10. The norbornyl cation and substituted nortricyclenes were shown to have five relatively weak bands in the C-H stretching region (-3000 cm-'). Only three were present in norbornane and its derivatives, and these were generally at lower frequency than those in the ion. In the C-C stretching region, the norbornyl ion and nortricyclene exhibit very strong peaks at 952 and 951 cm-' and similar spectra between 700 and 1000 cm-' . The alkylnorbornyl cations, 2-methyl and 2-ethyl, on the other hand, show substantial similarities to norbomane in this region. These similarities gave rise to the original suggestion that the ion had a geometry closely related to nortricyclene, C3" symmetry, and was probably an "edge-protonated" nortricyclene, even though such a species no longer retained this symmetry. In view of subsequent C studies, the Raman data have also been deemed compatible with a corner-protonated nortricyclene which now has C, symmetry. It is thus evident that the Raman data may be viewed as being consistent with either an edge- or comer-protonated nortricyclene and one may ask if it might also be consistent with the presence of a classical ion of C1-symmetry. In our view, it is extremely difficult to use the Raman data to dismiss any structure of the norbornyl ion. The main reason for this is that a normal-coordinate analysis of the corner-protonated

TABLE 10 Main Raman Spectral Lines (cm-I) of the Norbornyl and 2-Methyl- and 2-Ethylnorbornyl Cations and Vibrational Data of Nortricyclene and NorbornaneaVb

C-H Region (3100-2700 c m - l ) Norbornyl cation 3-Br-Nortricy clene Nortricy clene 2-Me-Norbornyl cation 2-Et-Norbornyl cation Norcamphor Norbornane 2-C1-Norbornane 2-Br-Norbornane 1-Mc-Nortricyclene 2-Mc-NorbornaneC 2-Et-NorbornaneC a

3110 3087 3080

3030 3006 2990

’

3069

2994

3010 2975 2945 2984 2980 2964 2964 2976 2978 2967 2952 2947

.

2947 2955 2915 2960 2965 2920 2936 2933 2933 2930 Source Source

Olah et aL, 1968. s = strong; m = medium; w = weak; p = polarized; vs = very strong. Aleksanyan and Sterin, 1957.

2860 2875 2867 2855 2873 2882 2873 2878 2878 2871 2869 2869

C-C Region (1000-250 cm-l)

911 (s) 920 (s) 938 (s) 920 (s, P) 925 (s, P) 925 (s) 905 (s, wp) 923 (s) 923 (m)

972 (vs, p) 959 (vs, p) 951 ( v s , ~ ) 880 (s, P) 884 (s, P) 880 (m) 871 (s, P) 879 (s, P) 878 (s) 857 (vs, wp) 873 (m) 878 (m)

796 732 783 770 765 786 753 764 762 794 719 696

(w) (m) (w) (m) (m) (m)

(m) (m) (m) (m) (m)

(m)

345 294

m P

21 7

THE NORBORNYL CATION

TABLE 11 Allowed Vibrations in the Norbornyl Cation

I on

Symm. type

Skeletal modes

H modes

PI

A' A" A

9

17 16 33 5

[GI Observed

6 15 3-4

Total

Select. rules

26 IRandRaman 22 IRandRaman 48 IR and Raman

nortricyclene and classical structures leads to the conclusion that there are many more bands allowed in both the C-C and C-H regions than are actually observed. The result of this analysis is shown in Table 11. Because of the discrepancy between the number of allowed and observed bands, it would appear that no firm conclusions can be drawn. Attention has also been focused upon the 3 110 cm-' vibration as evidence for the presence of a cyclopropyl ring by analogy with the 3080 cm-' band in nortricyclene. This, however, appears to be a tenuous argument, because the protonation of the cyclopropyl ring might be expected to remove electron density from the ring system and associated C-H bonds, thus weakening and shifting them to TABLE 1 2 Trapping the Norbornyl Cation Acid

Ion source

SbFs-SO2

7-Chloronorbornanea

-20 to

SbF5 -SO2

1Ghloronorbornanea

SbFs-FS03H-SO2

Quench

Major products

H2O

7-Norbornanol + 2-exonorbornanol

-20to -50

H2O

1-Norbornanol + 2-exonorbomanol

exo-2-Chloronorbornaneb

Low

Pyridine

Nortricyclene

SbF5-FS03H-S02

exo-2-Chloronorbornaneb

Low

(Me0 - ) MeOH

Nortricyclene and exo-2norbomyl methyl ether

SbFs-FS03H

exo-2-Chloronorbornane

-50 to -80

a Schleyer et al., 1964.

Olih et al., 1968.

t("C)

-50

Norbornane

218

G. M. KRAh4ER

lower frequencies. Rehybridization of the C-2-H bond, sp3 to sp2 , would, however, be expected to cause an opposite shift, the net result being somewhat in doubt. Alternatively, it is not possible easily to estimate the stretching frequency of the C-H bond at the charged carbon or elsewhere in ion [ 6 ] .The Raman spectrum is thus considered to provide contributory but inconclusive information regarding the structure of the norbornyl ion.

Quenching studies Also inconclusive are the results of the quenching experiments shown in Table 12. Thus, reaction of the ion with pyridine has been reported to yield nortricyclene, while water and methanol yield mainly 2-exo-norbornyl derivatives. If the ion is classical, the production of alcohols and methyl ethers is regarded as due to condensation of it and water, methanol, or methoxide. The production of nortricyclene would then be regarded as due to a “deep seated” reaction of pyridine with ion [2] leading to a transition state resembling pyridine, a proton and nortricyclene. If the ion is non-classical, nortricyclene may be viewed as the expected deprotonation product, while the hydrolysis products arise from a “deep seated” rearrangement to a transition state resembling the classical ion and the nucleophile.

10. RELATED IONS The solvolysis of camphenyl hydrochloride sparked off much of the research that has followed to elucidate the structure of the parent system. What can be said of the nature of the substituted ions themselves in the light of these studies? Again, solvolysis provides an insight into the nature of a transition state during the reaction. For a copious summary of these results the excellent review of Sargent (1972) should be consulted. By and large these studies seem to point to a classical structure for most of the derivatives but we will confine our remarks to studies of a few monoor di-substituted 2-norbornyl ions under “stable” ion conditions. The ‘H-nmr parameters for the 2-methyl, 1-2-dimethyl and 1,2dimethoxy-2-norbornyl cations are listed in Table 13. The 2-methyl derivative was found to form an ion in which the exo- and

219

THE NORBORNYL CATION

endo-6-hydrogens were distinguishable. While this strongly indicates that the ion is classical, it has been argued from C shifts that some o-delocalization is also present (Olah et al., 1971). The classical structure of this ion has been attributed to the fact that rearrangement must be an uphill process since it involves the conversion of a tertiary to a secondary ion. TABLE 13 Pmr Parameters of Substituted 2-Norbornyl Cations

6

Ion

6

A

B

A

B

3

4

5

6

7

H

CHja

464

3.00

2.70

2.70

1.47

1-71

CH3

CH3a8C

2.43

2.43

2.42

2.86

2.42

CH3O

CH30b

1.35

4.66

2-84

2.84

1.52 t(6), 42) 1-1-2.4

3.28~ 1-09n 2.86 1.1-2.4

1.1-2.4

~~

~~~

Olahetal., 1971. Nickon and Liu, 1969. x = exo, n = endo, 6 = chemical shift in p.p.m. from TMS, t = triplet, d = doublet, ( ) = coupling constant in Hz.

McFarland, 1969.

The 1,2-dimethoxynorbornyl ion studied by Nickon and Liu (1969) has also been deduced to be classical on the basis of temperature-dependent changes in the nmr spectrum. Whether the behaviour of methoxy derivatives is much related to that of alkyl derivatives is, however, not clear because much of the positive charge is expected to be delocalized to the oxygen atom of the substituent. This might be expected to provide an unusual stabilization to the classicd structure. No such stabilization is present in the 1,Z-dimethyI-Z-norbornyl cation and the proton spectrum does not permit a choice because the exo- and endo-6-protons exhibit a common peak. The methyl groups also afford a single non-resolvable band and the spectra might be compatible with either the bridged ion or two rapidly equilibrating

220

G. M. KRAMER

structures. This dilemma appears to be resolved by examination of the 1 3 C spectra (Olah et al., 1971). The 2-methylnorbornyl ion is used as a model for the dimethyl cation. It was previously shown that C-1 and C-2 had shifts of +118 and -76 p.p.m. in the frozen ion (Olah et al., 1969). These give +21 p.p.m. for the expected averaged shift of C-1 and C-2 under equilibrating conditions. Although one might ask if it is any more legitimate to extrapolate the shifts of tertiary systems than of secondary ones, the close comparison of the TABLE 14 Cmr Parameters of Substituted Norbornyl Cations'

Ion

Chemical shiftb

A

B

C-I

C-2

H CH3

CH3 CH3

118

-76

~~~~~

C-1,C-2 C-3,C-7

C-4

C-5

C-6

CH,

153

170

141 147

176

21 26

179

~

'Olah et af., 1971. Shifts are in p.p.m. relative to external I3CSz.

model with the dimethylnorbornyl ion makes it reasonable to conclude that the 26 p.p.m. shift found with the latter is evidence for the classical nature of this species. We have already noted that Sorensen and Ranganayakulu (1972) have also found that 3- and 7dimethyl-substituted derivatives of the 1,2-dimethyl-2-norbornyl ion yield the n.m.r. spectra of equilibrating ions rather than of a bridged species. The conclusion that these species are classical is also strongly buttressed by ESCA studies of the 2-methylnorbornyl ion (Olah et al., 1972). The C-1s electron spectrum of this cation exhibited two fairly well resolved bands with an area ratio of 711 and a 3.7 eV binding energy shift. The conditions of obtaining this spectrum are not clear from the paper but it is probable that the temperature was low enough to preclude any reactions or rearrangements. This is important because, if for some reason more than one equally long-lived species were present, the spectrum would represent a weighted average of the mixture.

THE NORBORNYL CATION

221

One thus comes t o the conclusion that under stable ion conditions the mono- and dialkyl-substituted norbomyl ions are open or classical ions. It is not difficult to rationalize the structure of the monoalkyl ion on the basis that bridging would be expected to be unfavourable because it leads the system towards an inherently unstable secondary ion. The reason for a classical preference of the ly2-dimethyl-2-norbomyl ion is less obvious from the non-classical view because electron release from 2-methyl groups might have been expected to augment the stability of a bridged structure. Perhaps it could be argued that the 1,2-dimethyl substituents introduce a stenc barrier to bridging. Such a proposal has been offered by Sargent (1972) to rationalize the finding that the solvolysis of 1,Z-diaryl-2-norbornyl precursors leads to classical ions (Schleyer et al., 1963). The argument presented is that bridging is primarily due to relief of steric strain present in the norbornyl ring, and consequently any structural modification which might reintroduce strain would impede the process. The argument is plausible when 1,Z-diaryl substituents are present because of non-bonded repulsions between the aromatic systems. On the other hand, inspection o f models of the 12-dimethyl-2norbornyl ion fails to reveal a similar barrier and it is doubtful if a steric argument permits a reasonable explanation of the data. Consequently, we believe that the data imply that o-participation is not only not strong enough to stabilize the symmetrically substituted ion relative to the classical tertiary species, but it would most probably be also too weak to stabilize the parent system.

11. SUMMARY The structure of the norbornyl ion under stable-ion conditions appears to be as elusive as that of the transition state reached during solvolysis. Theory suggests that the intrinsic energies of the classical and non-classical forms of the ion are extremely close and hence the structure to be found in solution will be governed by solvation effects. Inspection of strong acid solutions in which the ion may be stabilized by spectral and exchange probes fail, in our opinion, to establish the presence of a protonated cyclopropyl ring and therefore do not support the presence of the non-classical ion. The proton

222

G. U KRAMER

exchange studies favour a classical representation of the ion; the exchange information containing the inherent limitation that the behaviour of the norbornyl ion can be inferred from that of rearranging alkyl ions. Proton and 3C-nmr, ESCA, and Raman studies provide a wealth of information which unfortunately is not subject to a unique interpretation. The main conclusion to be drawn therefore is that the structure of the solvent stabilized cation is still unproven. Gas phase estimates of the heat of formation of the norbornyl cation imply a rather marked stability of the structure relative to other secondary ions (Kaplan et al., 1970). When combined with other estimates of the heat of formation of the t-butyl cation, however, these data suggest that hydride transfer from isobutane to the norbornyl ion will be endothermic by 6 to 15 kcal mole-’. This is contrary to experience in the liquid phase behaviour of the ion, and the author’s conclusion that their observation of enhanced stability is evidence of stabilization by bridging deserves further scrutiny.

REFERENCES Abodenn, A, and Baird, R L. (1964). J. Amer. Chem. SOC. 86, 2300. Aleksanyan, V. T., and Sterin, K. E. (1957). Fiz. Sb. L’vov. Cos. Univ. 1, 59;cf. Chem. Abs. 1959,53, 21158a. Allen, L. C., and Goetz, D. W. personal communication (1974). Baker, R., and Hudec, J- (1967). Chem. Comm. 929. Berson, J. A., Bergman, R. G., Hammons, J. H., and McRowe, A. W. (1967).]. Amer. Chem. SOC. 89, 2581. Berson, J. A, and Remanick, A. (1964). J. Amer. Chem. SOC. 86, 1794. Brouwer, D. M., and Oelderik, J. M. (1968a). Pet. Chem. Preprints 13, 184. Brouwer, D. M. (1968b). Rec. Truv. Chim. 87, 1435. Brouwer, D. M., and Hogeveen, H. (1970). Rec. Truv. Chim. 89,211. Brown, F., Hughes, E. D., Ingold, C. K., and Smith, J. F. (1951). Nature 168, 65. Brown, H. C., andchloupek, F. J. (1963).J. Amer. Chem. SOC. 85, 2322. Brown, H. C., Rothberg, I., Schleyer, P. von R., Donaldson, M. M., and Harper, J. J. (1966).Proc. Nut. Acud S c i , US.56, 1653. Brown, H. C., Hammer, W. J., Kawakami, J. H., Rothberg, I., and Jagt, D. L. van der (1967).J. Amer. Chem. SOC. 89, 6381. Brown, H. C. (1972). “Boranes in Organic Chemistry”, Chapt. IX, X,XI Cornell University Press. Brown, H. C., and Peters, E. (1973). J. Amer. Chem. SOC. 95, 2400. Burke, J. J., and Lauterbur, P. C. (1964).]. Amer. Chem. SOC. 86, 1870. Chambers, T. C., and Kistiakowsky, G. B. (1954). J. Amer. Chem. SOC. 56, 399. Cheney, B. V., and Grant, D. M. (1967). J. Amer. Chem. SOC. 89, 5319.

THE NORBORNYL CATION

223

Collins, C. J. (1967). Chem. Rev. 69, 543. Collins, C. J., and Lietzke, M. H. (1967). J. Amer. Chem. SOC. 89, 6565. Corey, E. J., andGIass, R. S. (1967). J. Amer. Chem. SOC. 89, 2600. Corey, E. J., Casanova, J., Vatakencherry, P. A., and Winter, R. (1963). J. Amer. Chem. SOC. 85, 169. Deno, N. C., Richey, H. C., Liu, J. S., Lincoln, D. N., and Turner, J. 0. (1965). J. Amer. C h e m SOC. 87, 4533. Ditchfield, R., Hehre, W. J., andPople, J. A. (1969). J. Chem. Phys. 51, 2657. Donaldson, M. M. ( 1958). Ph.D. Dissertation, Princeton University. Foote, C. S. (1964). J. Amer. Chem. SOC. 86, 1853. Fry, J. L., Lancelot, C. J., Lam, L. K. M., Harris, J. M., Bingham, R. C., Raber, D. J., Hall, R. E., and Schleyer, P. von R. (1970). J. Amer. Chem. SOC. 92, 2538. Goering, H. L., andSchewene, C. B. (1965).J. Amer. Chem. SOC. 87, 3516. Goering, H. L., andHumski, K. (1968). J. Amer. Chem. SOC. 90, 6213. Goetz, D. W., and Allen, L. C. (1971). XXIII International Congress of Pure and Applied Chem. Vol. 1, Boston. Hehre, W. J., Stewart, R. F., and Pople, J. A. (1969). J. Chem. Phys. 51, 2657. Hoffmann, R. (1964). J. Chem. Phys. 40, 2480. Huang, L., Ranganayakulu, K., and Sorensen, T. S. (1973). J. Amer. Chem. SOC. 95, 1936. Humski, K., MalajEiC, R., BorTiC, S., and Sunko, D. E. (1970). J. Amer. Chem. SOC. 92, 6534 and 6537, footnote 17. Jensen, F. R., and Beck, B. H. (1966). Tetrahedron Lett. 36, 4287. Jones, A. J., Huang, L., Haseltine, R., and Sorensen, T. S. (1974). Personal Communication. Kaplan, F., Cross, P., and Prinstein, R. (1970). J. Amer. Chem. SOC. 92, 1446. Kramer, G. M. (1970).J. Amer. Chem. SOC. 92, 4344. Kramer, G. M. (1973). Unpublished results. Lee, C. C., and Wang, E. W. C. (1964). J. Amer. Chem. SOC. 86, 2752. Lee, C. C., and Wang, E. W. C. (1965). Can. J. Chem. 43, 2254. Lee, C. C., and Lam, L. K. M. (1966). J. Amer. Chem. SOC. 88, 2831, 5355. Majerski, A., Schleyer, P. von R., and Wolf, A. (1970). J. Amer. Chem. SOC. 92, 5731. McFarland, J. T. (1969). Ph.D. Thesis, Calif. Institute of Technology, Pasadena, California Meerwein, H., and van Emster, K. (1922). Chem. Ber. 55, 2500. Murr, B. L., andconkling, J. A. (1970a).J. Amer. Chem. SOC. 92, 3462. Murr, B. L., and Donnelly, M. F. (1970b). J. Amer. Chem. SOC. 92, 6886. Nevell, T. P., de Salas, E., and Wilson, C. L. (1939). J. Chem. SOC. 1188. Nickon, A., andLiu, Y. (1969). J. Amer. Chem. SOC. 91, 6861. Olah, G. A., Baker, E. B., andComisarow, M. B. (1964). J. Amer. Chem. SOC. 86, 1265. Olah, G. A., Commeyras, A., and Lui, C. Y. (1968). J. Amer. Chem. SOC. 90, 3882. Olah, G. A., and White, A. M. (1969). J. Amer. Chem. SOC. 91, 3954. Olah, G. A., and White, A. M. (1969b). J. Amer. Chem. SOC. 91, 5801. Olah, G. A., DeMember, J. R., Lui, C. Y., and White, A. M. (1969). J. Amer. Chem. SOC. 91, 3958. Olah, G. A., White, A. M., DeMember, J. R., Commeyras, A., and Lui, C. Y . (1970). J. Amer. Chem. SOC. 92, 4627. Olah, G. A., DeMember, J. R., Lui, C. Y., and Porter, R. D. (1971). J. Amer. Chem. SOC. 93, 1442.

224

G. M. KRAMER

Olah, G. A., Mateescu, G. D., and Riemenschneider, J. L. (1972).J. Amer. Chem. Soc; 94, 2529. Olah, G. A., Liang, G., Mateescu, G. D., and Riemenschneider, J. L. (1973). J. Amer. Chem. SOC.95, 8698. Paul, E. G., andGrant, D. M. (1963).]. Amer. Chem. SOC.85, 1701. Pekhk, T. I., Lippmaa, E. T., Belikova, N. A., and Plate, A. F. (1970). Dokl. Akad. Nauk SSSR 195, 885. Pfeiffer, G . V., and Jewett, J. G. (1970).J. Amer. Chem. SOC. 92, 2143. Radom, L., Pople, J. A., Buss, V., and Schleyer, P. von R. (1971). J. Amer. Chem. SOC.93, 1813. Radom, L., Pople, J. A., and Schleyer, P. von R. (1972). J. Amer. Chem. SOC. 94, 5935. Roberts, J. D., and Lee, C. C. (1951). J. Amer. Chem. SOC. 73, 5009. Roberts, J. D., Lee, C. C., and Saunders, W. H. (1954). J. Amer. Chem. SOC. 76, 4501. Sargent, G . D. (1972). Chapter 24 In “Carbonium Ions”, G. A. Olah and P. von R. Schleyer, Eds. Wiley-Interscience, New York, p. 1099. Saunders, M. L., Schleyer, P. von R., and OIah, G. A. (1964). J. Amer. Chem. SOC. 86, 5680. Saunders, M. L., Hagen, E. L., and Rosenfeld, J. (1968). J. Amer. Chem. SOC. 90, 6882. Saunders, M. L., Vogel, P., Hagen, E. L., and Rosenfeld, J. (1973). Accounts of Chem. Res. 6, 53. Savitsky, G. B., and Namikawa, K. (1969). J. Phys. Chem. 68, 1956. Schaeffer, J. P., Dagani, M. J., and Weinberg, D. S. (1967).]. Amer. Chem SOC. 89, 6938. Scheppelle, S. E. (1972). Chem Rev. 72, 511. Schleyer, P. von R. (1964). J. Amer. Chem. SOC.86, 1854, 1856. Schleyer, P. von R., Watts, W. E., Fort, Jr., R. C., Comisarow, M. B., and Olah, G. A. (1964).J. Amer. Chem. SOC.86, 5679. Schleyer, P. von R., Donaldson, M. M., and Watts, W. E. (1965).J. Amer. Chem. SOC.87, 375. Schleyer, P. von R., Lam, L. K. M., Raber, D. J., Fry, J. L., McKervey, M. A., Alford, J. R., Cuddy, B. D., Keizer, V. G., Geluk, H. W., and Schlatmann, J. L. M. A. (1970).]. Amer. Chem. SOC.92, 5246. Schleyer, P. von R. to H. C. Brown (1972). Private communication reported in H. C. Brown, “Boranes in Organic Chem.”, Cornell Univ. Press. Shiner, Jr., V. J., Dowd, W., Fisher, R. D., Hartshorn, S. R., Kessick, M. A., Milakofsky, L., and Rapp, M. W. (1969).J. Amer. Chem. S O C . 91, 4838. Shiner, Jr., V. J., and Dowd, W. (1971a).J. Amer. Chem. SOC. 93, 1029. Shiner, Jr., V. J., and Fisher, R. D. (1971b).J. Amer. Chem. SOC.93, 2553. Sorensen, T. S., and Ranganayakulu, K. (1972). Tetrahedron Lett. 24, 2447. Stang, P. J., and Schleyer, P. von R. (1968). Abst. of the 155th Meeting of the Am. Chem. SOC.(San Francisco), p. 192, corrected value. Sustmann, R., Williams, J. E., Dewar, M. J. S., Allen, L. C., and Schleyer, P. von R. (1969). J. Amer. Chem. SOC.91, 5350 (1969). Winstein, S., and Schreiber, K. C. (1952).J. Amer. Chem. SOC. 74, 2165. Winstein, S., and Trifan, D. S. (1949). J. Amer. Chem. SOC. 71, 2953. Winstein, S., and Trifan, D. S. (1952).J. Amer. Chem. SOC.74, 1147, 1154.

Nucleophilic Aromatic Photosubstitution J. CORNELISSE, G. P. DE GUNST, E. HAVINGA Department of Organic Chemistry, Gorlaeus Laboratories, Leiden University, The Netherlands. 1.

.

Introduction . Review of Main Results 1956-1966. . Evaluation and Reorientation. Some complementary Novel Data and Reactions . . The Excited State From Which the Nucleophilic Aromatic Photosubstitution Starts. Kinetics . Introductory Remarks . . Recent Results; Kinetics Orientation Rules in Nucleophilic Aromatic Photosubstitution Meta- Activation . . Ortho-para Activation by Methoxy-(alkoxy-) Substituents &Reactivity in Polycyclic Aromatic Compounds Specific Influences of the Leaving Group, the Nucleophile and the medium. Investigations on Intermediates . Introduction Experimental Details . Results and Discussion Conclusion Epilogue References .

.

2.

3.

. .

.

4.

5.

.

. .

. .

. .

. . . . . . . .

. . .

225 225 233 235 236 236 238 245 245 246 249 250 253 253 254 255 260 261 264

1. INTRODUCTION Review of Main Results 1956-1 966 Almost 20 years ago it was discovered that nitrophenyl phosphates and sulphates upon ultraviolet irradiation of their solutions in water

226

J. CORNELISSE, G. P. DE GUNST, E. HAVINGA

show a substitution reaction to yield the corresponding nitrophenol and phosphate (sulphate) (Havinga, de Jongh and Dorst, 1956). At that time most photoreactions of organic molecules investigated were of a homolytic nature leading to free radical chemistry, and no well-studied heterolytic photoreactions were known. It was therefore a stimulating surprise to find these aromatic photosubstitution reactions that in their overall pattern are of heterolytic nature. Another peculiar feature that sparked further study was that in these seemingly nucleophilic photosubstitutions it was the metanitro-compounds rather than the ortho- and para-isomers, that showed the most efficient and clean reaction. This behaviour strongly contrasted with what might have been expected on the basis of the classical orientation rules for thermal reactions. I t was soon found that not only esters but also nitrophenyl ethers are capable of photosubstitutions (Havinga and de Jongh, 1962; Zimmerman and Somasekhara, 1963). When their solutions in alkaline medium are irradiated the alkoxy-group is replaced by hydroxyl. Photoinduced replacement of a nitro-substituent was reported at an early stage (Gold and Rochester, 1960, 1964; Johnson and Rees, 1964; Letsinger and Ramsay, 1964; de Vries and Havinga, 1965). Besides variation of the leaving group the change to other nucleophiles also proved feasible. IIlumination of (rneta-) nitrophenyl esters and ethers in the presence of amines resulted in the formation of (meta-)nitroanilines ( 1 ) (Havinga, 1962; Kronenberg et al., 1967).

NO2

Gradually it was recognized that nucleophilic aromatic photosubstitution is a fairly general reaction (Havinga et al., 1967; Havinga and Kronenberg, 1968). It can be realized also with polycyclic and heterocyclic aromatic systems. Various solvents (water, alcohols,

NUCLEOPHILIC AROMATIC PHOTOSUBSTITUTION

227

acetonitrile, etc.) may be used. As a result of contributions from several laboratories, a whole range of leaving groups and substituting nucleophiles is now known as exemplified in the following lists: Leaving groups: phosphate, sulphate, OR-, NO;, I - , Br-, C1-, F-, SO;-, SO,-, Nz, H.

Reacting nucleophiles:

HzO, ROH, primary, secondary and tertiary mines, pyridines, OH-, OR-, CN-, CNO-, H-, CH;, COO-, NO;.

In many of the cases studied a nitro-group is present as a substituent in the aromatic reactant and one gets the impression that this is favourable to the reaction. On the other hand, quite a few examples are known where no nitro-group plays a role, e.g., in the reactions of anisoles (Barltrop et al., 1967; Nilsson, 1971; Lok and Havinga, 19 73), in the photocyanation of aromatic hydrocarbons (Vink et al., 1972a), and in the photosubstitution of- aromatic ketones (Letsinger and Colb, 1972). It should be emphasized that the wide scope of nucleophilic aromatic photosubstitution does not imply that it will work indiscriminately with any combination of aromatic compound and nucleophile. On the contrary, there are pronounced selectivities. The general picture now arising shows a field with certainly as much variability and diversification as chemists, in the course of growing experience, have learned to appreciate in the area of classical (thermal) aromatic substitution. It is one of the aims of this article to contribute to a description and understanding of the various reaction paths and mechanisms of nucleophilic aromatic photosubstitution, hopefully to the extent that valuable predictions on the outcome of the reaction in novel systems will become feasible. Curiously, the earliest examples of aromatic photosubstitution found, the photohydrolysis and photoamination of (meta-) nitrophenyl esters and ethers, already show many of the aspects, the principles and the fundamental questions characteristic of the whole field. The photohydrolyses are often very clean reactions with good quantum yields. The absorption spectra of the system taken during the reaction show isosbestic points up to high degrees of conversion (Figure 1). The reaction may follow a different course with different nucleophiles. Experiments with O-labelling showed that mnitrophenyl phosphate upon illumination undergoes hydrolytic

228

J. CORNELISSE, G.

P. DE GUNST, E. HAVINGA

attack by water at the phosphorus atom, whereas with hydroxide ion and methylamine genuine aromatic substitution through attack at the ring carbon atom takes place (de Jongh and Havinga, 1966). The reaction is first order in light intensity and zero order in the light absorbing aromatic reaction partner (as long as its concentration is high enough completely to absorb the light).

0.4

0

I

300

500

400

h innm Figure 1. Isosbestic points; photohydrolysis of 3,5-dinitroanisole (0-01MNaOH; 313

4.

The quantum yield of formation of substitution product increases with increasing concentration of the nucleophile at relatively low concentrations. At higher concentrations the value of the quantum yield levels off to reach a constant value (Figure 2). These kinetic data suggest a pathway in which the nitrophenyl ester or ether, brought into an excited state by absorption of a light quantum, reacts in a bimolecular process with the nucleophile or returns to the ground state of the original molecule. At high nucleophile concentrations every excited molecule has one or more encounters with the nucleophilic reaction partner and the

NUCLEOPHILIC AROMATIC PHOTOSUBSTITUTION

229

“spontaneous” deactivation gets no chance to occur. The fact that under such conditions the quantum yield does not become unity but as a rule reaches values of 0.1-0.6 is then most easily described by the assumption that effective encounters between the excited Q,,,,, =0,42

Q

I

.loz

I

30-

20

i

04 0

4

8

12

16

“H31

20

(%I

___.c

Figure 2. Quantum yield as a function of nucleophile concentration (m-nitroanisole in

NH3 /CH30H).

aromatic species and the nucleophile yield substitution product in a certain percentage of cases only. The rest lead t o deactivation without overall chemical reaction.

l N*

product Figure 3. Schematic representation of pathways in a simple nucleophilic photosubstitution.

This picture, schematically represented in Figure 3, predicts linear “Stern-Volmer” plots for l/@us. l/nucleophile concentration (1/“1):

23 0

J. CORNELISSE, G. P. DE GUNST, E. HAVINGA

This has been experimentally verified in a great number of examples (Figure 4) and excellent linear relationships have been found. It is only at high concentrations that deviations occur, and then the value of CP becomes constant at lower concentrations than strictly corresponds with equation (2). Here it may become manifest that k in equations (2) and (3) is not a real kinetic constant, owing to the lack of a sufficient number of collisions before an encounter and subsequent reaction between the partners occurs. This should mean that the requirements for the statistics underlying equation (2) are no longer met.

04

0

20

40

60

a0

100

120

140

160

180

Figure 4. Linear relationships of l / O and 1/[N] for the reaction of Figure 2.

Concentrations of the nucleophile needed for good efficiency (quantum yield) are generally of the order of 1 0 - 4 - 1 0 - ’ ~ . Quenchers have to be present in the same concentration range to be effective. Assuming the reactions to be diffusion controlled, this suggests that the excited aromatic species which reacts is short-lived (10-7-10-10sec.). We shall discuss in Section 2 the question whether we are dealing with a singlet or a triplet state. The quantum yields are the same, independent of whether the irradiation takes place with light absorbed in the first or the second r + a * absorption band (exceptions were found later, e.g., in the case of the azulenes (Section 2), and in the photo-amination of nitrobenzene, (van Vliet et al., 1969). At long wavelength [A > 330 nm (for the photohydrolysis of m-nitroanisole)] a small but significant decrease in is observed, tentatively ascribed to absorption of part of the light by a hidden n + r* band. The obvious conclusion is that the reacting species is the aromatic molecule in its lowest excited

NUCLEOPHILIC AROMATIC PHOTOSUBSTITUTION

231

n + n* singlet state or in a triplet state originating from this lowest singlet. It was recognized a t an early stage that the fact that m-nitrophenyl phosphate and m-nitroanisole give a much cleaner and more efficient reaction than their ortho- and para-isomers, in itself, is not sufficient t o conclude that the rate of reactions of the excited meta-compound with the nucleophile is higher. Such a more efficient photoreaction might also be explained by a longer lifetime of the excited metacompound, by a different intersystem crossing efficiency, etc. However, with compounds such as nitroveratrole, where the lifetime is excluded as a discriminating factor between the meta- and para-position, the results indeed indicate a higher reactivity at the meta-position (Stratenus, 1966). These systems at the same time furnish nice examples of a complementarity between the reactions in the ground state (ortho-para-activation) and in the excited state (meta-activation with respect to the nitro-group). The conception of OCH, I

(4)

@OCH3

NO2

CH3

HO NO*

232

J. CORNELISSE, G. P. DE GUNST, E. HAVINGA

these nucleophilic aromatic photosubstitutions as bimolecular reactions of the nucleophile with the excited aromatic compound is supported by semiquantitative considerations and calculations of the charge densities (Fernindez Alonso, 1951; Grinter and Heilbronner, 1962; Havinga and de Jongh, 1962; Zimmerman and Sandel, 1963;

Figure 5. Calculated charge densities of 4-nitrocatechol and nitrohydroquinone in SO and Sl (Mulder, 1966).

Mulder, 1965/1966) (Figure 5 ) . From these it follows that in the first excited singlet state (and probably also in the triplet) the position meta with respect to a nitro-group is relatively more positive than the ortho- and para-positions. A nucleophilic attack at the rneta-position thus seems rational. An equivalent description in many respects consists of a primary transfer of an electron from a nucleophile molecule close to the meta-position of the excited aromatic compound and subsequent bond formation between the partners of the radical pair so formed.

233

NUCLEOPHILIC AROMATIC PHOTOSUBSTITUTION

Eualuation and Reorientation Although many questions remained t o be answered, it Iooked some years ago as if a consistent picture of the nucleophilic aromatic photosubstitutions was emerging from the experimental data and from charge density calculations. This picture would present an understanding of the photosubstitution process somewhat equivalent t o what had been reached in the course of time for the well-known thermal aromatic substitution reactions. These optimistic views received a first blow by the discovery of another category of nucleophilic aromatic photosubstitutions occurring in liquid ammonia as a solvent and nucleophile. In this medium the nitroanisoles still show the pattern that had become familiar, i.e., photosubstitution of OCH, by N H 2 with preference for reaction at the position meta with respect to the nitro-group. However, nitrobenzene, dinitrobenzenes, and nitrohalogenobenzenes @NO2

hv liq. NH3

H2N@N0~

c1

&NO2

hv

A

liq. NH3

+

(7)

‘@NO2

NH2

give substitution at the ortho- and para-positions [equations (6) and (7)J . Remarkably, we are dealing here with genuine photoreactionsthe reactions would not proceed at all in the dark-but the substitution pattern is the one t o be expected from molecules in the ground-state: ortho9ara-activation with respect to the nitro-group. The overall picture certainly became n o simpler when quite a few examples were encountered where the nature of the nucleophile is decisive for the position at which the photosubstitution occurs. Just one recent example, taken from the work of Lammers (1974) is given’ in equation (8). To account for this influence of the nucleophile, the formation of an excited complex between aromatic compound and nucleophile has been suggested as the primary intermediate (de Vries, 1970). Now, by constructing a scheme for the reaction pathway consisting of a combination of photoadditions and consecutive thermal

234

J. CORNELISSE, G. P. DE GUNST, E. HAVINCA

OCH3 I

I

NO2

I

CN

reactions, it remains possible formally t o describe the different phenomena in the various nucleophilic aromatic photosubstitutions. But it was felt at this stage that in order to arrive at a real understanding and a rationalization of a quantitative nature, allowing detailed predictions to be made, a penetrating study using a variety of physical and chemical methods should be undertaken. Such a should aim at answering the following questions: What is the nature of the excited state from which the reactions start (T + A * , n + A * , singlet or triplet) and how is this state reached (spectroscopy, sensitization and quenching experiments)? What are the orientation rules than can be extracted from the data obtained with various aromatic compounds and nucleophiles? What is the reaction pathway and what are the intermediates on the way to relaxation and product formation (flash methods, esr, C.I.D.N.P.)? What is the influence of the detailed molecular structure of aromatic and nucleophilic partners, of the nature of the reaction medium and of the temperature? Is there any correlation with charge densities, localization energies etc.? What are the mechanisms of the consecutive reaction steps?

235

NUCLEOPHILIC AROMATIC PHOTOSUBSTITUTION

In the following sections of this article we shall deal with investigations concerning the questions enumerated above. In order to obtain an effective basis, it seems appropriate here to complement the review of data from older studies given in this introducti0.n by a selection of some pertinent results and novel reactions from recent investigations.

Some Complementary Novel Data and Reactions With respect to the classes of aromatic compounds studied, it should be mentioned that bi- and tricyclic compounds (derivatives of naphthalene, azulene, biphenyl, phenanthrene, etc.), the reactions of which had incidentally been touched in the past, have received considerable attention in recent years. It is found that generally the polycyclic aromatic compounds undergo nucleophilic aromatic photosubstitution more easily and cleanly than benzene derivatives. Here, even the unsubstituted hydrocarbons prove capable of undergoing light-induced substitution (Vink e t al., 1972a; 1972b). As a second, interesting feature may be reported the fact that in the photoreactions the influence of a substituent in one ring extends t o the positions in the other (adjacent) ring much more strongly than with thermal (ground state) reactions. Characteristically for such reactions an “extended meta-activation” is found (Beyersbergen van Henegouwen and Havinga, 1970; Pirkinyi, 1972). We shall come back t o these reactions as well as to the photosubstitutions of heterocyclic compounds later in this article (Sections 2 and 3). r

n

O

C

H

3

hv, OH-

moCH3 hv, OH-

O2N

OCH3

O2N

OH

Amongst the nucleophiles most used in the study of nucleophilic aromatic photosubstitution, the cyanide ion has proved of special interest (Letsinger and McCain, 1966, 1969; Letsinger and Hautala, 1969; Lok et al., 1970). In many instances it gives efficient substitution, promoted by presence of oxygen (air). With some

236

J. CORNELISSE, G. P. DE GUNST, E. HAVINGA

nitrobenzene derivatives, presumably those where the lowest triplet has much n - + a * character, cyanide ions effect photoreduction of the nitro- to the nitroso-group (Vink et af., 1971; Petersen and Letsinger, 1971).

In the range of leaving groups, fluorine has been recognized as a valuable substituent that has practically no heavy atom effect and that in many cases is smoothly replaced under the influence of irradiation (Brasem e t al., 19 72). With Z-fluoro-4-nitroanisoleit even proved capable of efficient substitution by the weak nucleophile water, a reaction that has not been equalled by any other substituent.* Curiously, the photosubstitution of fluorine by cyanide is generally less efficient than that by other nucleophiles.

2. THE EXCITED STATES FROM WHICH THE NUCLEOPHILIC AROMATIC PHOTOSUBSTITUTION STARTS. KINETICS Introductory Remarks As has been indicated in Section 1 , the investigations on nucleophilic aromatic photosubstitution reactions during the first ten years brought a wealth of circumstantial data and semi-quantitative results allowing speculation concerning the reaction pathway on the one hand and making evident the need for more quantitative and mechanistic studies on the other hand. That the excited state from which the actual reaction begins would have a + a* rather than n + a* character was no difficult guess. Semi-quantitative considerations and simple calculations of electron distribution already make it dear that upon n +a* 2 The photohydrolysis of the nitrophenyl phosphates in neutral medium consists of an attack of water at the phosphorus atom. There are indications that with 3,5-dinitroanisole photosubstitution by water might occur under certain conditions (de Gunst, 1971).

NUCLEOPHIUC AROMATIC PHOTOSUBSTITUTION

23 7

absorption of nitro-aromatics negative charge moves from the nitrogroup into the ring, making this latter unattractive to nucleophilic attack. Since P.P.P. (Pariser, Parr, Pople) calculations for the first A + K* singlet (and also for the first A + A* triplet in the few cases where calculations were performed) indicate a high electron density at the position metu with respect to a nitro-substituent, the balance is in favour of a K -+ A* state as a candidate for undergoing a nucleophilic attack. Moreover, no choice is left in the case of photosubstitutions of the aromatic hydrocarbons themselves where n -+ 7 ~ *transitions do not occur (Vink et ul., 1972a). The question whether we are generally dealing with singlet or triplet reactions was in the beginning considered somewhat in favour of the singlet, for which an estimated lifetime of lo-'- lo-' O sec. (see Section 1) seemed normal. The lack of influence of oxygen (air) pointed in the same direction. However, the question became fully open again when evidence accrued that nitrobenzenes have high rates of S T intersystem crossing and, consistently, very short triplet lifetimes (Hurley and Testa, 1968). Fortunately, some 5 years ago the problem was ripe for systematic tackling by sensitization and quenching methods as well as by laser photolysis. At the time a range of sensitizers and quenchers of well-known behaviour had become available, and much experience in this field had been gained. An early investigation concerned the photoreactions of pnitroanisole. Letsinger and Steller ( 1969) elegantly demonstrated not only that its photosubstitution by hydroxide ion and by pyridine could be sensitized (by benzophenone), but also that the product distribution of the photohydrolysis (p-nitrophenollp-methoxyphenol 1:4) was the same whether it was the result of direct irradiation or of a sensitized reaction. The first systems with meta-activation deliberately investigated by sensitization (and quenching) experiments were m-nitroanisole in liquid ammonia (van Vliet et ul., 1970) and 1-methoxy-6-nitronaphthalene in alkaline medium (Beyersbergen van Henegouwen and Havinga, 1970). In these two cases indications of a singlet reaction were found. With m-nitroanisole in liquid ammonia the benzophenonesensitized reaction yields inter aliu 2-methoxy-4-nitroaniline as a product and no m-nitroaniline, which is formed in very high yield upon direct irradiation in liquid ammonia as well as in NH, /CH3OH. In the latter instance l/@varies linearly with 1/[NH3], suggesting that the reaction is either singlet or triplet but not of a mixed type. --f

-

238

J. CORNELISSE, G. P. DE GUNST, E. HAVINCA

One is thus led to conclude that we are dealing with a photosubstitution starting from a singlet. The photohydrolysis of 1methoxy-6-nitronaphthalene can be sensitized (benzophenone), although this procedure yields more side-products than the direct irradiation. Quenching by triplet quenchers (sorbic acid, tetramethyldiazetine dioxide) cannot be effected. The same is true for the phototo give 6-nitro-N-methyl-asubstitution by methylamine naphthylamine, indicating a singlet mechanism for this reaction also (Lammers, 1974).

R ecen t Results; Kinetics General aspects In recent years a series of representative nucleophilic aromatic photosubstitution reactions has been systematically investigated by product analysis, by sensitization and quenching, by flash photolysis and in a few cases also by nanosecond laser photolysis (see Section 4). It appears by these criteria that many of the classical examples such as the photohydrolysis of m-nitroanisole (den Heyer et al., 1973) and 3,5-dinitroanisole (de Gunst and Havinga, 1973), as well as the recently studied photohydrolysis and photoamination of l-fluoro-3-nitronaphthaleneand the photocyanation of 2-nitrothiophene are triplet reactions. The lifetimes of the triplets are short (in water 10-6-10-9 sec. and of course still shorter in the presence of strong nucleophiles and/or quenchers). The rates of reaction are very high, varying from ca. one hundredth to ca. one third of the rate of triplet quenching. Examples (a) In order to explore how far a scheme of the type indicated in Section 1 would prove adequate consistently to describe quantitative results obtained from photosubstitution reactions of a chosen aromatic compound with various nucleophiles and quenchers in a wide concentration range, experiments were performed on solutions of l-fluoro-3-nitronaphthalenewith hydroxide ion and methylamine as nucleophiles and 3,3,4,4-tetramethyldiazetine dioxide (TMDD) as a quencher (Tamminga and Lammers, 1973). From the orientation rules now known, it may be expected that the fluorine atom at position 1 of naphthalene and meta with respect to the nitro-

NUCLEOPHILIC AROMATIC PHOTOSUBSTITUTION

239

substituent will be capable of smooth and efficient nucleophilic photosubstitution (see Section 3). It gives very clean reactions and good quantum yields indeed as shown in equation (12). The reactions could be sensitized by benzophenone (triplet energy of

For N = OH-, For N = CH3NH2,

@=

0.35 at [OH-]

= 0.1~

= 0.25 at [CH3NH2] = 0 3 M

1-fluoro-3-nitronaphthalene = 57 kcal mole- from phosphorescence measurements at -190°C in tetrahydrofuran). Quenching could be effected by TMDD. Kinetic data indicate a reaction starting from one excited state (I/@ is linear with [Q] and @ goes to zero with increasing [Q] ). Evidently, we are dealing with a pure triplet state reaction. For description and terminology we use the following simple scheme (Figure 6).

proaucr Figure 6. Scheme for photoreaction of a triplet species with a nucleophile (N) in the presence of a quencher (Q).

The scheme of Figure 6 leads to:

240

J. CORNELISSE, G. P. DE GUNST, E. HAVJNGA

The following linear equations were deduced from the experimental data: : l/@ = 2-4 + 0*057/[OH-] For N = O H For N = CH, NH2 For [ N ] = 0-1 M OH-

Q = TMDD

I

: 1/@= 3.6 : l/@ =

+ 0.046/[CH3NH2 ]

2-8 + 1300 [TMDD]

Assuming by analogy that @isc = 0-8 [the value of 2-nitronaphthalene (Rusakowicz and Testa, 1971)], one obtains:

For N = CH3NH2 : k 2

- 2k1

Izd = 0.037 k,

The values were checked by ilieasurement at [OH-] = 0-01 M . While k [CH3NH2 ] k [OH-], the values for k 2 differ by a factor of two. Adopting the value 10' for k , , one has k l = k,/100 lo8 and

-

-

-

0-046k, 4.6 X l o 6 . The triplet lifetime is thus found to be -2 x lo-' s, a value of the expected order of magnitude.

kd =

24

0

r

1

Figure 7. Relation between nitrothiophene by cyanide ion.

2

I

3

1 / 0 and l/(CN-]

I

L

5

1

6

'/[CN-].1031 mold' in the photosubstitution of 2-

24 1

NUCLEOPHILIC AROMATIC PHOTOSUBSTITUTION

For the sake of comparison a similar analysis was performed on the photosubstitutions of the analogous methoxy-compound (1methoxy-3-nitronaphthalene)with the same nucleophiles (OH- and CH3NH2) (Lammers, 1974). The quantum yields of these reactions (at comparable nucleophile concentration) are lower than those for the reactions of the fluoro-compound. This is found to be by no means indicative of a low reactivity of the excited species. On the contrary, making a reasonable guess for @ i s c , the value of which lies between 0-1 and 0.7, and assuming it to be 0.3 one finds k l 3 5-10' and kz * l o 9 mole/l sec, i.e. rather higher than for the fluorine analogue. The lower value of the quantum yields of photosubstitution of the methoxy compound appears to be caused by the predominating influence of the short lifetime of its triplet state: T = lo-' sec. The quantitative data obtained explicitly prove that it is essential in the discussion and the interpretative use of quantum yields to break these down on the basis of appropriate measurements to vdues of intersystem crossing efficiencies, life-times and reaction rate constants. (b) As a second example (Groen and Havinga, 1974) the photocyanation of 2-nitrothiophene will be briefly dealt with. This compound upon irradiation in water or methanol undergoes smooth substitution of the nitro-group by nucleophiles (CN-, CH30-) as shown in equation (15). In Figure 7 the linear relationship between

l/@ and l/[CN-] is demonstrated. The reaction can be triplet sensitized (acetone) and is quenched by sodium sorbate. The results of the measurements are described by: l/
kl +k2-

- 2.4 ;

~

1 x kd - 6.73 x @isc

k2

kl

1

-x

!% = 3.04

Qisc

kl

Assuming Q i s c to be not far from unity, and k, to be ca. 10'

kl

2:

1/3 k,

2:

0.3 x 10" 1 mol-' s - l ; k z k d 21

2 x lo6 s-'

gives:

0.4 x 10'' I mol-' s - l ;

242

J. CORNELISSE, G. P. DE GUNST, E. HAVINGA

Consequently, the triplet lifetime of 2-nitrothiophene in water is cu. 5.10-' s and cu. lo-' s in 0.01 M KCN. Owing to the extremely fast reaction with cyanide ion, the triplet lifetime, which in water is reasonably long, appears to be considerably reduced in the presence of even moderate cyanide concentrations. Unimo lecular kine tics Already, with one of the first nucleophilic aromatic photosubstitutions encountered, curious behaviour was found when studying the rate of reaction as a function of pH. rn-Nitrophenyl sulphate shows no increase in the quantum yield of photohydrolysis with increase of hydroxide ion concentration up to values as high as 0-1 M . This behaviour, also found with one other compound ( 5 chloro-3-nitroanisole), is in clear contradistinction to what is TABLE 1 Quantum Yields of Photohydrolysis of some Halogenopyridines at Different pH Values: ax,Disappearance of Hdogenopyridine; @OH, Formation of Hydroxypyridine Compound

pH

13

2-C1-pyridine

@X

2-Br-pyridine

4

0.11

0.11

@OH

0.07

0.07

@x

0.47 0.32

0.45 0.31

0.40 0.29

0.3 0-14

0-2 0.15

0.2 0.18

0.16

0.14

@OH

0.13

0.14

@OH

0-2

0.2

@OH

2-I-pyridine

ax

3-Br-pyridine

@X

@OH

4-Br-pyridine

7

0.12

2.4

2.1

1.2

0.30

0.20

0.03

0.0

observed with the majority of the compounds investigated. However, in the photoreactions of heterocyclic compounds studied recently, quite a few more cases of comparable behaviour were encountered (van der Stegen, 1972). Table 1 summarizes the data for photohydrolysis of the halogenopyridines. It is seen from Table 1 that the series of halogenopyridines show quantum yields of photohydrolysis independent of pH in the neutral and alkaline region and decreasing only at pH-values where the pyridine nitrogen becomes protonated to a considerable percentage. A clue to an understanding of what may go on in these systems was offered during a recent extension of the investigations to incorporate 5-membered heterocyclics (Groen and Havinga, 1974).

NUCLEOPHILIC AROMATIC PHOTOSUBSTITUTION

243

The photocyanation of 2-nitrothiophene in KCN solution in water (see page 241) shows not only a normal increase of @ with increasing [CN-] but also the quantum yield of its consumption equals that of the formation of 2-cyanothiophene. With 2-nitrofuran under similar conditions the @ value for formation of 2-cyanofuran again increases in the usual way with increase of [CN-] , levelling off at a limiting 0.1 M . However, the quantum yield of value of 0-51 at [CN-] disappearance of the nitrofuran is 0.51, independent of the cyanide concentration. Evidently, there is a second photoreaction (with water) that “complements” the cyanation. Indeed ybutenolide could be found as a product of irradiation of 2-nitrofuran in water.

Equation (16) offers a picture that is able to describe the phenomena, if it is assumed that the formation of X is quantum yield-determining and that its reaction with water and other nucleophiles to form “photoproducts” is very fast. Of the various guesses one can make about the nature of X two hypotheses deserve closer investigation. The first is that X might be a triplet and the quantum yields are determined by intersystem crossing only. One of the objections t o this assumption, at least in the case of the hdogenopyridines, is that the quantum yields of product formation seem to be lower than is expected from known values for intersystem crossing (for pyridine itself ca. 0.36, Cundall et al., 1964). A second hypothesis is that X may be a positive ion intermediate in a substitution of an A S N 1 type. For the furan derivatives, the formation of such a cation might be favoured by assistance through electron donation from the ring oxygen. The hypothesis seems more difficult, e.g., for pyridine derivatives, unless it is assumed that the excited species which reacts has some n + K* character. Whichever this may be, the old “anomalous” photohydrolysis, where no increase of rate was found upon increase of hydroxide ion concentration, now seems consistently to fall into a wider category of reactions where the maximum quantum yield of conversion is reached merely by reaction with water.

244

J. CORNELISSE. G. P. DE GUNST, L HAVINGA

Azulene derivatives With the reactions discussed so far, variation of the wavelength of irradiation has only a minor influence on the quantum yield of nucleophilic aromatic photosubstitution (Section 1). Nitroazulenes are found to be interesting exceptions (Lok et al., 1973). TABLE 2 Quantum Yields of the Photoreaction of 1-Nitroazulene with Methoxide Ion (1 M ) in Methanol at Various Wavelengths Wavelength (in nm) 546 405 366 313 254

Quantum yield

< 0.0010 0.0020 0*0020 0.0034 0.004

In Table 2 are presented the quantum yields of the substitution of 1-nitroazulene by methoxide ion (1 M , methanol) upon irradiation at different wavelengths. Although the values may have been influenced somewhat by complex formation of the reaction partners in the ground state, it can be easily shown that this does not change the picture essentially. For the reaction of 1-cyano-3-nitroazulene, quantum yields can be corrected for the influence of complex formation on the basis of spectroscopic data, yielding the values 9.10-4 at 3 13 nm and 8.10-’ at 405 nm (0.01 M CH,ONa in methanol). The substitutions are triplet reactions and the variation in quantum yield may reflect the low intersystem-crossing from the first singlet and a direct intersystem crossing from higher excited singlet states to the triplet manifold. It seems significant that the fluorescence of 1-nitroazulene is also wavelength-dependent (Dhingra and Poole, 1968). For the quenching of the photomethoxylation of 1-nitroazulene (1 M CH, ONa in methanol) by TMDD the following linear relationship holds: l/@= 625 + 2-73 x l o 4 [Q]. This leads to a value of ca. 20 ns for the lifetime of the reacting nitroazulene triplet in the absence of quencher.

NUCLEOPHILIC AROMATIC PHOTOSUBSTITUTION

245

3. ORIENTATION RULES IN NUCLEOPHILIC AROMATIC PHOTOSUBSTITUTION

Meta-activation More than anything else it was the curious meta-activation found in the photosubstitution of nitrophenyl esters and ethers by various nucleophles that stimulated closer investigation of this class of reactions (Section 1). In order to avoid misunderstanding, it should be stressed that the concept of activation at the meta-position &th respect to the nitro-group does not exclude reactions taking place at other positions in the excited molecule. It merely means that, other things being equal, substitution is preferred at the meta-position over the ortho- and para-positions. A few examples of this meta-activation have been given in Section 1 and many more are to be found in a recent compilation (Cornelisse and Havinga, 1975). The study of the photosubstitution of nitromethoxynaphthalenes leads to an extension of this rule, showing that, e.g., in the reactions of 5 - and 6-nitro-2,3-dimethoxynaphthalene, the influence of the nitro-group extends to the methoxy-substituted ring much more than in ground state reactions and effects a clear meta-activation here also. In a recent systematic study on a complete series of fluoro-1-nitronaphthalenes reacting with hydroxide and methoxide ions, the photosubstitution is again found to proceed more cleanly and smoothly as soon as there is a meta-relationship between the fluoro- and the nitro-substituent (Lammers, 1974). The meta-activation thus defined remained the characteristic orientation rule for almost 10 years. It finds a rationalization in charge density calculations for the lowest (a + a*) singlet and, in the few cases where calculations have been made, also for the lowest a + a* triplet. The nucleophile is supposed to show a higher rate of attack at the positions of relatively high positive charge. Through this picture, part of the influence of the medium can also be understood from its capacity to make an excited state with a + a* character the one of lowest energy. It is a a +a* rather than n + a* or a charge transfer state from which nucleophilic substitution can easily take its start. This line of thought, based on ideas and work of Porter and Suppan (1965; Porter, 1967) on the photochemistry of aromatic

246

J. CORNELISSE. G. P. DE GUNST, E. HAVINGA

carbonyl compounds and on data of Hurley and Testa (1968) was elaborated by Letsinger and Hautala (1969). Letsinger and McCain (1969) direct attention to the fact that the leaving tendencies of the substituents in the aromatic compound should play a role in the determination of the position where substitution will occur. As has been pointed out in Section 1, the simple picture in which the charge distribution of the excited aromatic compound determines the position of substitution, although containing elements of truth, cannot adequately cover the newer experimental findings and rules which we now will summarize. Ortho-para activation by methoxy-(alkoxy-) substituents The idea of orthopara activation by alkoxy-substituents in nucleophilic aromatic photosubstitution has been explicitly formulated very recently (Lok and Havinga, 1974). Looking back, a few older findings can be described by this new orientation rule. We mention the easy replacement of the halogeno-substituent in 2halogeno-4-nitroanisoles (Nijhoff and Havinga, 1965; Nijhoff, 196 7; Brasem et al., 1972) and the substitution of the nitro-group of p-nitroanisole (Letsinger et al., 1965; de Vries and Havinga, 1965). The remarkably clean photosubstitution of the methoxy substituent of m-nitroanisole may be ascribed not only to meta-activation, but also to the fact that in this compound the nitro-group is not activated

*

AH3

C ,H3

-0.276 -0.015

(J0.052+0.154

0

+0.021 -O.O9I

10.109

-0.055

SO

+0.054

b

-0.03 1 -0.019 0-0.099

s2

Sl

C , H3

0 +0.234

C ,H3

0 +0.329

b b -0.046 -0.031 -0.019 -0.075

Tl

+0.199 0.167

0+0.171

50.186

T2

Figure 8. Charge distribution of anisolc in the ground state and in singlet and triplet excited states (see Lodder and Havinga, 1972).

247

NUCLEOPHILIC AROMATIC PHOTOSUBSTITUTION

by methoxy. Recent data that can be covered by the rule are the formation of 0- and p-cyanoanisole from anisole (Nilsson, 1971) and the easy photocyanation at the 5-position of 6-methoxyquinoline (Letsinger, 1972) as well as the formation of 1-amino-2-methoxyanthraquinone upon irradiation of 2-methoxy-anthraquinone in ammonia (Griffith and Hawkins, 1973). However, in contradistinction to the meta-activation by the nitrogroup, there is no attractive simple rationalization of the orthoactivation by the methoxy group. One may formulate in general terms a complementarity of rules for thermal and photochemical processes. However, such a generalization could possibly predict ortho-para orientation but not activation. First of all, careful experimental testing of the hypothesis seems required. We present a selection of very recent results. , 1. Illumination of 1-methoxynaphthalene in media containing cyanide ion gives a good yield of 1-cyano-4-methoxynaphthalene (and very little 1-cyanonaphthalene) whereas photocyanation of naphthalene itself proceeds with only small yields (Lok and Havinga, 1974). 2. The three dimethoxybenzenes as well as the three trimethoxybenzenes undergo smooth photocyanation, yielding in all cases the products to be expected from ortho-para orientation (Figure 9) (den Heyer, 1973). It is clear from these examples that there is not only orthopara orientation but also activation. If several positions in the same molecule are competing in the substitution reaction, factors such as

OCH3

C H 3 0Lf O g OC CHH3 3

L@0CH3 OCH3

OCH3

OCH3

I

C H 3 0 q O C H 3

OCH3

Figure 9. Positions of preferential attack in photocyanation.

248

J. CORNELISSE, G. P. DE GUNST, E. HAVINGA

intersystem crossing efficiency and lifetime of the reacting excited species cannot operate as discriminating factors. 3. Whereas in examples 1 and 2 above it may be argued that the direction and ease of reaction is (partly) due to “merging resonance stabilization” during product formation, this cannot be a positive factor in cases where OH- or OR- function as the nucleophile. Valuable examples are seen in the reaction of some o- and phalogenoanisoles. o-Fluoroanisole and especially p-chloroanisole give smooth photosubstitution upon illumination in the presence of various nucleophiles (Brasem and Lammers, 1972; den Heyer, 1973). OCH3 I

CN

Cl

6 OCH3

4. Particularly striking is the reaction of 5-nitro-1,3-dimethoxybenzene (Wiegerink, 1973). Equation (18) demonstrates that the ortho-para-orientation by the two methoxy-groups may successfully compete with the meta-activation by the nitro-substituent.

Summarizing the experimental data, we may say that ortho-para activation by methoxy-substituents seems at the moment to be as well established as the meta-activation by a nitro-group. In some of the cases (photocyanation) this ortho-para activation may be understood on the basis of the resonance stabilization during product

249

NUCLEOPHILIC AROMATIC PHOTOSUBSTITUTION

formation. For the photosubstitution of methoxybenzene derivatives by nucleophiles such as OH- and OCH; a simple rationalization does not seem obvious.

&-Reactivity in Polycyclic Aromatic Compounds A third orientation rule presents itself from a survey of the nucleophilic photosubstitu tions of bi- and tricyclic aromatic compounds (Vink et al., 1972a; Vink et al., 1972b; Lok e t al., 1973).

&-

&

t Figure 10. Positions of preferential photosubstitution by CN‘ (in t-BuOH/H,O).

Biphenyl, naphthalene, azulene, phenanthrene, etc. show preferential “a-reactivity” in photocyanation (Figure 10). This preference of photoreaction with a nucleophile at position 1 of azulene and naphthalene (4and 2 in biphenyl, 9 in phenanthrene) is also evident upon considering the products from the reactions of derivatives of these hydrocarbons (Lok, 1972). In many other cases besides those represented in Figure 10 and equations (19) and (ZO), the “a-reactivity” can be recognized as a major orientation rule.

+

NO2

azoxynitronaphthalene (19)

250

J. CORNELISSE, G. P. DE GUNST, E. HAVINGA

For azulene and its derivatives, the preference for the 1-position can be rationalized by considering the charge densities calculated for the various electronic states (Table 3, Pariser, 1954).

TABLE 3 Calculated Charge Densities for the Ground State (So), the Lowest Excited Singlet State ( S l ) , and the Lowest Triplet State (T1) of Azulene (Pariser, 1954)

Atom no.

SO

Sl

T1

123 2 4,8 5, 7 6

-0.096 +0*021 +0*122 -0.049 +Om052

+Om145 -0.1 18 -0.080 +0-072 -0.108

+0.138 -0.126 -0.067 +0-092 +0.128

From Table 3 it can be understood that in the ground state position l(3) will be favoured in electrophilic substitution, whereas in the first excited singlet or triplet this same position shows the largest positive charge facilitating attack by nucleophiles. In agreement with this line of thought, in 1,3-di-t-butylazulene the photocyanation has been found to occur at the position 5(7), which is the next best from consideration of charge distribution (Lok et al., 1973). However gratifying this rationalization may seem to be in the case of azulene derivatives, it can of course not be used with naphthalene, biphenyl, etc. Here we have to confine ourselves to the statement that nucleophilic photosubstitution preferentially occurs at the positions that are known to be generally the most reactive in thermal reactions also.

Specific influences of the leaving group, the nucleophile and the medium Although strictly speaking the influences of the specific properties of the attacking nucleophile, the leaving group and their interaction during the photosubstitution process may not be classified as “orientation rules”, the interplay of these factors and the rules evidently determines the course and outcome of the reactions in a given system. It seems appropriate therefore to give in this section a

NUCLEOPHILIC AROMATIC PHOTOSUBSTITUTION

25 1

survey of some regularities observed in this intriguing field, where rationalization is still fragmentary and of a tentative nature. Of course one should not expect the relative leaving tendencies of substituents to be the same in the ground state and in an excited state of an aromatic compound. Nor should this be the case for the nucleophilicities of the partners reacting in the one instance in a thermal process governed by a free energy of activation and in the other in an extremely fast reaction where a description by way of ideas derived from transition state theory would not seem to be applicable. If one realizes furthermore that the reactions occur in solution and that the specific properties of substituents are dependent on the nature of the medium (solvent), it is clear that there is such wide variability and delicate balance between a series of factors that a knowledge of facts and chemical intuition will certainly remain a valuable asset as it still is in ground state chemistry. Notwithstanding the complexity to be anticipated of the phenomena, one or two valuable generalizations can already be deduced from the rapidly growing body of experimental results. In the following, we shall confine ourselves to some indications and remarks based mainly on our own experience. These have mostly to be of a qualitative nature since systematic quantitative measurements that allow calculation of rate constants have been started only recently (Section 2).

1. An important factor that somehow s e e m to determine the course of many aromatic photosubstitutions and that may form the background of some of the selectivities in the interaction between a given nucleophile and the substituents present in the excited aromatic reaction partner can be described as “merging (resonance) stabilization” during product formation. This line of reasoning, formulated already on the basis of early results (Havinga et al., 1967; Havinga and Kronenberg, 1968), has proved applicable t o quite a few of the recent results also. It is clear that the ortho-para-directing influence of alkoxy groups on a substituting cyan0 group can be rationalized on the basis of the existence of a favourable interaction between the electron-donating ether group and the electronwithdrawing cyano-substituent. On the other hand, cyanide ion as a substituting reagent will meet a barrier to settling down at the positions ortho or para with respect t o a nitro-group, in contrast t o an attacking hydroxide or methoxide ion. A difference may be expected between tertiary amines that become electron-withdrawing

252

J. CORNELISSE, G. P. DE GUNST, E. HAVINGA

upon being bound t o the ring carbon atom and primary and secondary amines that eventually may show up as electron donating groups.

\

CN-

, CN

2. Groups that generally function well as leaving groups in nucleophilic aromatic photosubstitution are the methoxy(alkoxy)group and the nitro-group. There is less experience with esters such as phosphates and sulphates but the data available confim the expectation that these also are readily photosubstituted by various nucleophiles. Of the halogens, the fluoro-substituent is a favourite in photosubstitution (cf. Section 2). It shows a curious selectivity in that as a rule it is not easily substituted by cyanide ion, a nucleophile that functions so well with other leaving groups. Chlorine and bromine react under favourable conditions (activation by other substituent, appropriate nucleophile). Iodine may also be photosubstituted by nucleophiles but is easily induced to enter into homolytic reaction pathways. One has to bear in mind that the heavier substituents (iodine, bromine and even chlorine) increase the rate of intersystem crossing which, depending on the conditions, increases or decreases the quantum yield. Hydrogen can be substituted by cyanide ion (22) and less easily by hydroxide ion. The leaving tendency of hydrogen is enhanced by oxygen (and probably other oxidants). Even so, the replacement of a nitro-group (23) is generally preferred over that of hydrogen. CN

JJ - (0 0

hu CN-

&

NUCLEOPHILIC AROMATIC PHOTOSUBSTITUTION

253

3. Although nucleophilicities will be different in photosubstitution as compared t o thermal substitution and although one should also remain aware of the different influence of the solvent, the efficiencies of nucleophiles in aromatic photosubstitution largely follow expectations based on chemical insight and experience of ground state chemistry. Water and methanol react in specially favourable systems only. Amines are much more effective as became clear at the beginning of the investigations by the possibility of photoamination in amine/ water and amine/methanol mixtures. Interestingly, nitroanisoles, like other aromatic nitro-compounds, when irradiated in pure diethylamine give only reduction products (BarItrop and Bunce, 1968). Efficient nucleophiles are OH- and OCH;. The system NaOCH3 /CH30H in particular often shows clean reactions with good yields of methoxy-aromatic compounds. The cyanide ion is another valuable nucleophile that will give efficient substitution with many partners. It has the additional possibility of functioning as a deoxygenating agent with n + K* excited nitroaromatic compounds, yielding nitroso-compounds and cyanate (Petersen and Letsinger, 1971; Vink et al., 1971). A similarly interesting nucleophile is the hydride ion (borohydrides etc.) that may also function as a substituting or as a reducing agent (Petersen and Letsinger, 1971). For synthetic purposes the cyanate ion is a promising reagent leading to amines in water and t o carbamates in alcohol (Hartsuiker et al., 1971). Alkyl-lithium reagents may also function in nucleophilic photosubstitutions with aromatic partners containing suitable leaving groups (Shapiro and Tomer, 1968).

4. INVESTIGATIONS ON INTERMEDIATES

Introduction Clearly one of the methods of elucidation of the reaction pathway of nucleophilic aromatic photosubstitution consists of the study of

254

J. CORNELISSE, G. P. DE GUNST, E. HAVINGA

intermediates occurring during the conversion of the starting material to reaction products. In the search for short-lived intermediate species in photochemical processes flash photolysis presents itself as a logical tool. This is true in particular for aromatic photosubstitutions where the sequence of events following the exciting flash can easily be followed by monitoring the ultraviolet absorption spectra. After it had been found in some exploratory investigations that short-lived species could be observed upon flashing solutions of nitroanisoles and nucleophiles (Cornelisse and Havinga, 1966) a more detailed investigation was started. In the following we deal with the results (de Gunst, 1971; de Gunst and Havinga, 1973) obtained with conventional flash photolysis as well as with laser photolysis. A few remarks will be made concerning electron spin resonance results in connection with the flash work. A scheme of a possible reaction pathway will be offered.

Experimental Details In the flash photolysis experiments nitrogen-filled discharge lamps were used. With photoflashes of 1500- 2000 Joule, duration about 20 ps, solutions of aromatic compound ( M ) and nucleophilic reagent (0.5- 0-005M ) in acetonitrile- water (1:1) were excited. Absorption spectra were recorded with spectroflashes of 50 Joule, duration about 15 p s , via a spectrograph on photographic plates. Time-dependent absorption changes, measured with a continuous light source, were displayed on an oscilloscope. In the laser photolysis experiments the aromatic compound (4.10-4 M ) and the nucleophile (0.04 M ) in acetonitrile- water (1:1) were irradiated with the frequency doubled pulse (100 mJ, 6 ns, 347 nm) of a ruby laser. Only time-dependent absorption changes were measured (double pulsed xenon flash lamp with 10 ps continuous output as light source); absorption spectra were constructed from these measurements at 12 or 25 nm intervals. The electron spin resonance spectra were run in nitrogen-saturated solutions of aromatic compound (ca. l o e 3 M ) and nucleophile (0.05-0-1 M ) in the solvent(s) indicated. Irradiation in the cavity was effected with a high pressure mercury arc. Electrolysis was performed with the platinum cathode in the cavity, tetraethylammonium perchlorate as electrolyte and electric currents of 10- 250 PA.

NUCLEOPHILIC AROMATIC PHOTOSUBSTITUTION

255

Results and Discussion Flash excitation of solutions of 3,5-dinitroanisole in the presence of nucleophiles produces a short-lived species with an absorption maximum at 550-570 nm (Figure 11, curve a). Decay of this species follows first order kinetics. The lifetime varies rather strongly with the nature of the nucleophile, from 40 ms with hydroxide ion t o 0.1 ms with thiocyanate (Table 4).The absorption spectrum is independent of the reagent used. TABLE 4 Lifetime

(7) of

the Absorption at 550-570 nm in the Presence of Various Nucleophiles in Acetonitrile-Water 1: 1 .

~~

Nucleophile

r/ms

OH-

40

H2 0 '

12 8 2 1.5 1 0.5 0-2 0.1 0.3

'In a mixture of acctonitriie-water ( 1 :9). b In methanol.

In mixed solvents containing more than 50% water the 550-570 nm species can be generated without a nucleophile added. Evidently water plays a nucleophilic role. Decay of the species in these water-rich media in the presence of nucleophilic reagents is n o longer first order. The behaviour of other aromatic nitrocompounds (e.g. m nitroanisole, 3,5-dimethoxynitrobenzene and 4-nitrobiphenyl) follows the same pattern: the same short-lived absorption is produced upon exciting an aromatic compound in the presence of a variety of nucleophilic agents, whereas the lifetime of the species formed depends on the nature of the reagent. Two hypotheses concerning the nature of the 550-570 nm species have been tested: the species might be either a complex between the excited aromatic compound and the nucleophile of such a structure

256

J. CORNELISSE, G. P. DE GUNST, E. HAVINGA

that it can yield starting materials and photosubstitution products, or it might be a radical anion formed by electron transfer from the nucleophile to the aromatic compound, a process which is favoured by the enhanced electrophilicity of the aromatic nitro-compound in its excited state. To test the first hypothesis, solutions of 3,5-dinitroanisole and hydroxide ions were flashed and the absorption spectra at different time intervals after excitation were compared. The absorption (X,, = 400-410 nm) that remains after all time-dependent absorptions have decayed can be shown to be due to 3,5-dinitrophenolate anion, the photosubstitution product of 3,5-dinitroanisole with hydroxide ion. When the absorption band of the 550-570 nm species is subtracted from the spectrum of the solution immediately after the flash, there remains an absorption at 400-410 nm, which can also be ascribed to 3,5-dinitrophenolate anion. The quantity of this photoproduct does not increase during the decay of the 550-570 nm species. Therefore the 550- 570 nm species cannot be intermediate in the aromatic photosubstitution reaction of 3,5-dinitroanisole with hydroxide ion to yield 3,5-dinitrophenolate. Repetition of the experiment with a variety of nucleophiles on this and other aromatic compounds yielded invariably the same result: nucleophilic aromatic photosubstitution is, in all cases studied, completed within the flash duration (about 2 0 ~ s of ) our classical flash apparatus. Experiments along the second line consisted of irradiating nitrogen saturated solutions of 3,5-dinitroanisole and nucleophiles in the cavity of an electron spin resonance spectrometer. It was found that radicals with, surprisingly, only one nitrogen coupling constant were formed. The e.s.r. spectrum hardly changed on variation of the nucleophile. The absence of the second nitrogen splitting appears to be a rather general phenomenon for dinitrocompounds in protic solvents (Ward, 1960; Ayscough et al., 1963; Gough and Cross, 1968 and de Gunst, 1971). There is no difference in the ability of photogeneration of radicals between systems that show photosubstitution and systems that do not, for example amines with 3,5-dinitroanisole and m-dinitrobenzene, respectively. This indicates that the formation of radicals from excited aromatic nitro-compounds in the presence of nucleophiles has no direct relation with the photosubstitution reaction. When radicals are produced electrochemically in the same solvents and under similar conditions again only one nitrogen splitting is observed in the e.s.r. spectrum. The change to a non-protic solvent,

NUCLEOPHILIC AROMATIC PHOTOSUBSTITUTION

25 7

such as acetonitrile, gives the expected radical anion of dinitrocompounds, with both nitrogen coupling constants apparent. Some illustrative data on the e.s.r. spectra are collected in Table 5 . TABLE 5 Electron Spin Resonance Coupling Constants (Gauss) of Radicals of Some Aromatic Nitrocompounds, Produced Electrochemically and Photochemically, in Protic and Aprotic Media Aromatic compound Solvent 3,5-dinitroanisole 3,5-dinitroanisole 3,5-dinitroanisole m-dinitrobenzene m-dinitrobenzene m-dinitrobenzene m-nitroanisole m-nitroanisole m-nitroanisole

Method aN (1) QN (2) QH-Z "H-4

CH~CN-HZOPhoto CH~CN-HZ0 Electro CH3CN Electro CH3CN-HZO Photo CH3CN-Hz 0 Electro CH3CN Electro CH~CN-HZOPhoto CH3CN-Hz0 Electro CH3 CN Electro

122 13.4 441 12.5 13.5 4.56 13.3 13.2 102

0 041

441 0 04 456 -

-

"H-5 OH-6

3.56 3.56 - 3.08 3.25 3.25 - 3.25 441 2.63 - 1441 3.37 3.37 1.08 3.37 3.27 3.27 1.00 3.27 3-064.13 1-064.13 3.31 3.31 1.09 3.56 3-403-401.10 3.40 3-543-541.06 3.81

We come to the conclusion that in the radical anions of the dinitrocompounds one of the nitro-groups is excluded from contributing to the spin density dishibution, most probably by fast protonation by the solvent. In this way a reduction process is started. That such reactions may play a significant role in the consecutive processes of 3,5-dinitroanisole radical anion is substantiated by the upon electroformation of 3,3'-dimethoxy-5,5'-dinitroazoxybenzene lysis of 3,5-dinitroanisole in acetonitrile-water ( 1 : 1). In order t o measure the absorption spectra, the radical anions were generated electrochemically in the optical path of a spectrophotometer. The absorption spectrum of 3,5-dinitroanisole radical anion (Figure 11, curve c) is very similar to that of the 550-570 nm species produced photochemically. So we believe this species to be the radical anion formed by electron transfer from the nucleophile t o the excited 3,5-dinitroanisole and decaying by interaction with its surroundings including the nucleophile radical cation. The behaviour described seems to be rather general for aromatic nitro-compounds since it is observed with a series of these compounds with various nucleophilic reagents. The results with the classical discharge flash were positive in that a short-lived species, the radical anion, could be traced and identified,

J. CORNELISSE, G. P. DE GUNST, E. HAVINGA

258

E

480

400

640

560

.

nm

A,

Figure 11. Absorption spectrum of 3,5-dinitroanisole radical anion, produced from 3.5-dinitroanisole by: (a) irradiation in akaline acetonitrile-water with a 20 ps discharge flash; (b) irradiation in alkaline acetonitrile-water with a frequency doubled ruby laser pulse (347 nm, 6 ns); (c) electrolysis in acetonitrile in the presence of tetraethyl ammonium pcrchlorate.

A

1

II

400

I I

II

1 1

II

I I

I I

500

I I

II

600

A,

nm

Figure 12. Absorption spectra of a solution of 3,5dinitroanisole and hydroxide ion in acetonitrile-water (1:l) at different time intervals after irradiation with a laser pulse. 0 = 0, 1 = 200,2 = 400,s = 600 and 4 = 900 ns.

NUCLEOPHILIC AROMATIC PHOTOSUBSTITUTION

259

but not of much help to our main problem, since this species is found not to be an intermediate, at least not in the photosubstitutions under discussion. Laser photolysis offers new possibilities, allowing the study of much faster processes. Once again the system 3,5-dinitroanisole and hydroxide ion was irradiated. The constructed absorption spectra (Figure 12) at different time intervals after pulse irradiation indicate two decaying absorptions with different lifetimes, probably representing two different species: 1. 412 nm species, A, 2. 475 nm species, h,,,

--

412 nm, lifetime 500 ns 475 nm, lifetime 12 ns.

-

During the decay of the 4 12 nm species an absorption at 5 12- 550 nm builds up. An isosbestic point is observed at ca. 535 nm (Figure 12). The optical density of the solution continues to increase until about 1.5 ps after the exciting pulse and then remains constant (at 51 2- 550 nm) up to the maximum time of measurement (10 ps). We think this last absorption to be due to the radical anion of 3,5-dinitroanisoIe (Figure 11, curve b) formed from the 412 nm species. After 1.5 ps there is no further decay of the 412 nm absorption. The optical density of the solution at that wavelength is higher than before the pulse, due to the formation of the photosubstitution = 400- 41 0 nm). Alproduct (3,5-dinitrophenolate anion; A, though the formation of the substitution product from the 412 nm species can thus not be observed directly, we propose this pathway as the most plausible. In the absence of nucleophile, neither the 412 nm species nor the formation of the radical anion, nor that of the photosubstitution product is found. It is concluded therefore that the 412 nm species results from some kind of interaction between the (excited) aromatic compound and the nucleophilic reagent. The character of this (6 aromatic compound-nucleophile-complex” is as yet unknown. However, in our present view, the nature of the complex has t o allow for the formation of both the radical anion and the photosubstitution product(s). An attractive possibility for this complex remains the o-complex, in formal analogy with the Meisenheimer complexes in the thermal nucleophilic reactions with aromatic compounds. An “exciplex” forms another possibility. The second short-lived species, with A,, = 475 nm, does not depend for its occurrence on the presence of a nucleophile. The

260

J. CORNELISSE, G. P. DE GUNST, E. HAVINGA

lifetime varies with increasing nucleophile concentration: T is about 1 2 ns (vide supra) in the presence of 0.04 M NaOH; T is ca. 5 5 ns in the absence of hydroxide ions. Apparently, the hydroxide ion plays a quenching role with respect to the 475 nm species. It is known (Section 2) that in this system we are dealing with a substitution starting from a triplet state. From experiments with triplet quenchers, [OH-] = 0.025 M , it was found by applying Stern-Volmer kinetics that k,. T = 350 (de Gunst and Havinga, 1973). Assuming the quenching to be diffusion controlled (kd8f = 1.3 x 10’ 1 mol-’ s - l in the solvent used) the lifetime of the triplet is calculated to be T = 27 ns. This number agrees very well with the measured lifetimes of the 475 nm species. We therefore consider that the 475 nm species corresponds with the triplet state of 3,5-dinitroanisole. Owing to the small difference between the pulse duration and the triplet lifetime in solutions containing the nucleophile, direct evidence could not be obtained for the formation of the “aromatic compound-nucleophile-complex” from the triplet state. However, we have n o indications against this route and for the moment we wish to adopt it since it gives a simple picture consistent with all experimental data.

Conclusion At the end of this section we present a reaction s’cheme (Figure 13) which accounts for the experimental evidence obtained so far.

Figure 13. Reaction scheme for nucleophilic aromatic photosubstitution (the values in parentheses refer to the photohydrolysis of 3,5-dinitroanisole in alkaline medium).

NUCLEOPHILIC AROMATIC PHOTOSUBSTITUTION

26 1

The excited singlet state reached upon excitation ( 1 ) decays to the ground state by internal conversion (2) or to the triplet state by intersystem crossing (3). It has been shown in Section 2 that we deal, at least in many photosubstitutions of aromatic nitro-compounds, with a a-a* triplet as the reactive excited state. Interaction with the nucleophile leading to an “aromatic compound-nucleophilecomplex” ( 5 ) competes with the decay of the triplet to the ground state (4). This complex has several reaction paths at its disposal: (I) dissociation into starting compounds (6); (11) dissociation into aromatic radical anion and nucleophile radical cation (7); and (111) loss of a substituent to form substitution product(s) (8). Path I represents a “chemical quenching” of the excited state. Path 11, probably only important with nitroaromatic compounds, leads via the radical anion to reduction products (10) (a process that is supposed here to start from a a-r* state). The radical anion may return to ground state aromatic compound by loss of an electron (9),forming another way of chemical quenching. The photosubstitution product(s) emerge from path I11 (8), whereby the ratio of the rate constants of processes ( 6 ) and (8) largely determines the quantum yield of the nucleophilic aromatic photosubstitution.

5. EPILOGUE

In this review we have concentrated on giving a picture of the main lines and some recent developments in the field of nucleophilic aromatic photosubstitution. A more complete enumeration of the many reactions discovered will be presented elsewhere (Cornelisse and Havinga, 1975). We thought it preferable not to include in the discussion the electrophilic aromatic photosubstitution since well-studied examples are still scarce (proton-deuteron and proton-triton exchange in acidic media; protodeboronation). From our experience we have the feeling that many electrophiles are very efficient quenchers and that moreover it is not easy to choose systems where the concentrations of the potentially reacting electrophiles can be made high enough to react efficiently with the short-lived excited species. In Section 1, after a short survey of older investigations from various laboratories, we formulated a number of fundamental questions and suggested lines of investigation to get more insight into the pathways and the intermediates of nucleophilic aromatic photo-

262

J. CORNELISSE, C. P. DE GUNST, E. HAVINGA

substitution, the mechanism of the reaction steps and the factors governing product composition and quantum yields. In Sections 2, 3 and 4 the results of recent studies in these directions are presented. Let us try to summarize the main novel developments compared with the situation in the field as described in the last survey some six years ago (Havinga and Kronenberg, 1968). 1. Perhaps the most striking new result is that, in all the various reactions investigated so far by flash photolysis, the end products of the photosubstitution are formed within a period of 10-6s or less. Free radical anions are formed in some of the systems; they have lifetimes of the order of 10-3-10-’ s and they d o not contribute significantly t o substitution product formation. Evidently in order to trace intermediates of the substitution reaction we have to resort to still faster methods (laser photolysis, Section 4). 2. In some of the reactions investigated (sensitization and quenching) the excited species which reacts seems to be the lowest n + T * singlet, in many other cases the reaction is found to start from a (T + T * ) triplet. Also in the latter cases the reacting species is very short lived; lifetimes of the triplets measured were.10-6-10-7 s in the absence of quenchers and nucleophiles. 3. Quantitative data on rates of reaction have been obtained for some of the triplet reactions. Assuming triplet quenching to be approximately diffusion controlled, the rate constants for the reactions between excited species and nucleophile are lo-’ O 1 mole-’ s - I . The data show that in comparing and interpreting quantum yields-even in the case of related systems-one should proceed to determine separately rate constants as well as intersystem crossing efficiencies and lifetimes of the reacting excited species. 4. Indications are increasing that the first “chemical” reaction step consists of the formation of a complex between the excited aromatic compound and the nucleophile (exciplex, excited .rr or u complex?). From this complex the reaction products are formed via one o r more subsequent reaction steps. Such a pathway makes it possible to rationalize inter ufiu the large influence of the nucleophile on the position where it will eventually become attached through o-bonding at the ring carbon atom. The results and interpretations summarized in points 1-4 have been represented diagrammatically in Figure 13. This forms in a sense a concrete formulation of the speculations presented earlier (Havinga and Kronenberg, 1968). I t makes clear that the quantum yields and the product composition are determined by a combination of

NUCLEOPHILIC AROMATIC PHOTOSUBSTITUTION

263

photo-processes and thermal reactions and that in general they cannot simply be derived from the properties of the excited aromatic compound only. In this connection we wish t o mention preliminary results (van Riel and Lodder, 1973) of a study in progress on systems where substituting and leaving groups are identical. The rates of exchange of the methoxy-groups (labelled with “0) of the three isomeric nitromisoles in NaOCH, /CH, OH are found t o be practically the same. In terms of the scheme of Figure 13 this may indicate that the meta-activation observed with other reactions has its origin mainly in the second part of the process, possibly in the balance between the splitting off of the original substituent and the newly attached nucleophile. 5. Some four principles and orientation rules are now rather well established: meta-activation by the NOz -group; ortho-para-activation by alkoxy-substituents; “a-reactivity” in polycyclic aromatic compounds; merging (resonance) stabilization during product formation. As in ground state chemistry, a judicious handling of these rules often enables qualitative prediction of the results of reactions in novel systems. In this connection it seems in order to formulate one broad generalization. In many cases nucleophilic aromatic photosubstitution occurs at the positions at which electrophilic substitution takes place in the ground state. However, this generalization has not only its important exceptions (photoamination of nitrobenzenes in liquid ammonia) but it also brings the danger of contenting oneself with a general statement that in the various instances may rest on different mechanistic factors. We have to leave open of course the possibility that here again this is a consequence of a general complementarity between ground state reactions and photoreactions, as was recognized earlier in a striking form, e.g., with the cyclization reactions of conjugated systems such as dienes and trienes. Coming back to the orientation rules and to the selectivities in the reactions of certain leaving groups and nucleophiles, one might wonder how these selectivities are to be reconciled with the old adage that specificity of reaction decreases with increasing reactivity of a reagent. This may constitute n o real difficulty if one realizes that generalizations of ground state chemistry should not be extended to photochemical reactions, where lines of thinking related t o

264

J. CORNEUSSE, G. P. DE GUNST, E. HAVINGA

transition state theory may prove misleading. Furthermore we should keep in mind that the orientation rules until now mostly reflect the results of product analysis. What may seem a considerable difference on the basis of product isolation can be the effect of relatively small differences in rates of reaction. Finally the picture which has been developed, in which nucleophilic aromatic photosubstitution proceeds through a sequence of steps of different nature, naturally leaves room to explain the selectivities and variability of the overall reaction. The future hopefully will bring the precise and quantitative elucidation of the pathways and mechanistic aspects of a few deliberately chosen model systems. A beginning has been described for the photohydrolysis of 3,5-dinitroanisole. It should be possible to complement this by studying systems in which the step of product formation can be directly followed. On the other hand the multiplicity of orientation rules found and the growing list of novel reactions with often surprising aspects guarantees that nucleophilic aromatic photosubstitution remains a field that will inspire chemists to rewarding and enjoyable exploration for a long time in the future.

REFERENCES Ayscough, P. B., Sargent, F. P., and Wilson, R. (1963).]. Chem. Soc. 5418. Barltrop, J. A., Bunce, N. J., and Thomson, A. (1967).J. Chem. Soc. (C)1142. Barltrop, J. A., and Bunce, N. J. (1968).]. Chem. Soc. (C),1467. Beyersbergen van Henegouwen, G. M. J., and Havinga, E. (1970).Rec. Trav.

Chim. 89,907.

Brasem, P., and Lammers, J. G. (1972).Unpublished results, Leiden University. Brasem, P., Lammers, J. G., Lugtenburg, J., Cornelisse, J., and Havinga, E. (1972).Tetrahedron Letters 685. Cornelisse, J., and Havinga, E. (1966).Tetrahedron Letters 1609. Cornelisse, J., and Havinga, E. (1975).Chem. Rev. to be published. Cundall, R. B., Fletcher, F. J., and Milne, D. G. (1964).Truns. Faraday Soc. 60,

1146.

de Gunst, G. P. (1971). Thesis, Leiden. de Gunst, G. P., and Havinga, E. (1973).Tetrahedron 29, 2167. de Jongh, R. O., and Havinga, E. (1966).Rec. Trav. Chim. 85, 275. de Vries, S., and Havinga, E. (1965).Rec. Trav. Chim. 84,601. de Vries, S. (1970).Thesis, Leiden. den Heyer, J., Spee, T., de Gunst, G. P., and Cornelisse, J. (1973).Tetruhedron Letters 1261. den Heyer, J. (1973).Unpublished results, Leiden University. Dhingra, R. C., and Poole, J. A. (1968).Chem Phys. Letters 2, 108. Fernindez Alonso, J. T. (1951).Compt. Rend. 223,403. Gold, V . , and Rochester, C. H. (1960).Proc. Chem. Soc. 403. Gold, V., and Rochester, C. H. (1964).J. Chem. Soc. 1687-1735.

NUCLEOPHILIC AROMATIC PHOTOSUBSTITUTION

265

Gough, T. E., and Cross, J. M. (1968). J. Chem. Soc. ( B ) 798. Griffith, J., and Hawkins, C. (1973). J.C.S. Chem. Comm. 11 1. Grinter, R., and Heilbronner, E. (1962). Helv. Chim. Acta 45, 2496. Groen, M. B., and Havinga, E. (1974). Mot. Photochem. 6, 9. Hartsuiker, J., de Vries, S., Cornelisse, J., and Havinga, E. (1971). Rec. Trau. Chim. 90, 61 1. Havinga, E., de Jongh, R. 0.. and Dorst, W. (1956), Rec. Trau. Chim. 75, 378. Havinga, E., and de Jongh, R. 0. (1962). Bull. Soc. Chim. Belg. 71, 803. Havinga, E. (1967). Proc. 13th Conf. Chemistry, Brussels 1965; Interscience. Havinga, E., de Jongh, R. O., and Kronenberg, M. E. (1967). Helv. Chim. Acta 50, 2550. Havinga, E., and Kronenberg, M. E. (1968). Pure Appl. Chem. 16, 137. Havinga, E. (1962). Chimica 16, 145. Hurley, R., andTesta, A. C. (1968). J. Amer. Chem. Soc. 90, 1949. Johnson, R. M., and Rees, C. W. (1964). Proc. Chem. Soc. 213. Kronenberg, M. E., van der Heyden, A., and Havinga, E. (1967). Rec. Trau. Chim. 86, 254. Lammers, J. G. (19 74). Thesis, Leiden. Letsinger, R. L., and Ramsay, 0. B. (1964). j . Amer. Chem. Soc. 86, 1447. Letsinger, R. L., Ramsay, 0. B., and McCain, J. H. (1965). J. Amer. Chem. Soc. 87, 2945. Letsinger, R. L., and McCain, J. H. (1966). J. Amer. Chem. Soc. 88, 2884. Letsinger, R. L., and Steller, K. E. (1969). Tetrahedron Letters 1401. Letsinger, R. L., and Hautala, R. R. (1969). Tetrahedron Letters 4205. Letsinger, R. L., and McCain, J. H. (1969). J. Amer. Chem. Soc. 91, 6425. Letsinger, R. L. (1972). Quoted in P. Courtot, “Elements de Photochirnie AvancCe”, p. 343, Hermann, Paris. Letsinger, R. L., andColb. A. (1972). J. Amer. Chem. Soc. 94, 3665. Lodder, G., and Havinga, E. (1972). Tetrahedron 28, 5583. Lok, C. M., Lugtenburg, J. Cornelisse, J., and Havinga, E. (1970). Tetrahedron Letters 4701. Lok, C. M. (1972). Thesis, Leiden. Lok, C. M., den Boer, M. E., Cornelisse, J., and Havinga, E. (1973). Tetrahedron 29, 867. Lok, C. M., andHavinga, E. (1974). Koninkl. Nederl. Akad. Wet. Proc. B 77, 15. Mulder, J. J. C. (1965/1966). Quoted in Havinga, 1967 and Havinga and Kronenberg, 1968. Nijhoff, D. F., and Havinga, E. (1965). Tetrahedron Letters 4199. Nijhoff, D. F. (1967). Thesis, Leiden. Nilsson, S. (1971). Thesis, Lund. Pariser, R. (1954). J. Chem. Phys. 25, 1112. Piirkinyi, C. (1972). Personal communication. Petersen, W. C., and Letsinger, R. L. (1971). Tetrahedron Letters 2197. Porter, G., and Suppan, P. (1965). Trans. Faraday Soc. 61, 1664. Porter, G. (1967). Proc. 13th Conf. Chemistry, Brussels 1965; Interscience. Rusakowicz, R., and Testa, A. C. (1971). Spectrochim. Acta 27A, 787. Shapiro, R. H., and Tomer, K. (1968). Chem. Comm. 460. Stratenus, J. L. (1966). Thesis, Leiden. Tamminga, J. J., and Lammers, J. G. (1973). Unpublished results, Leiden University. van der Stegen, G. H. D., Poziomek, E. J., Kronenberg, M. E., and Havinga, E. (1966). Tetrahedron Letters 6371. van der Stegen, G. H.D. (1972). Thesis, Leiden.

266

J. CORNELISSE, G. P. DE GUNST, E. HAVINGA

van Riel, H. C. H. A., and Lodder, G. (1973). Unpublished results, Leiden University. van Vliet, A., Kronenberg, M. E., and Havinga, E. (1966). Tetrahedron Letters

5975.

van Vliet, A., Cornelisse, J., and Havinga, E. (1969).Rec. Trav. Chim. 88, 1339. van Vliet, A., Kronenberg. M. E., Cornelisse, J., and Havinga, E. (1970). Tetrahedron 26, 1061. Vink, J. A. J., Cornelisse, J., and Havinga, E. (1971).Rec. Trav. Chim. 90,

1333.

Vink, J. A. J., Lok, C. M., Cornelisse, J., and Havinga, E. (1972a).J.C.S. Chem. Comm. 710. Vink, J. A. J., Verheydt, P. L., Cornelisse, J., and Havinga, E. (1972b). Tetrahedron 28, 5081. Ward, R. L. (1960).J.C h e m Phys. 32,410. Wiegerink, F. J. (1973).Unpublished results, Leiden University. Zimmerman, H. E., and Sandel, V. R. (1963).j.Amer. Chem. SOC. 85,915. Zimmerman, H. E., and Somasekhara, S. (1963).J. Amer. Chem. SOC. 85,922.

Alternative Protonation Sites in Ambident Conjugated Systems

M. LILER School of Chemistry, The University, Newcastle upon Tyne, England

1.

2. 3.

.

Introduction Methods of Investigation Spectroscopic Methods . Correlation Analysis Cation Stability and Solvation Intrinsic Cation Stability Cation Solvation Tautomerism Kinetic vs. Thermodynamic Stability Protonation Sites in Conjugated Molecules N/N Alternatives N/O (or N/S) Alternatives . N/C Alternatives O/O (or O / S ) Alternatives O/C (or S/C) Alternatives . N/C/O (or N/C/S) Alternatives . References

. .

. .

4.

. .

. . . .

.

. .

. .

.

.

. . .

. . . . . . .

.

.

267 270 270 277 287 287 291 296 298 300 301 328 351 363 370 378 382

1 . INTRODUCTION Over the past 10-15 years there has been a considerable increase of interest in the question of protonation sites in organic molecules, which contain two or three possible protonation sites in conjugation.

268

M. LILER

This is mainly due t o the advent and wider availability of instrumentation for nuclear magnetic resonance spectroscopy, which offers a means of identifying various types of protonation either by direct observation of the spectra of the protonated species or by the effect on the spectra of proton exchange phenomena. The stimulus provided by the nmr studies has led to greater activity in the application to the same question of other spectroscopic methods, which have been available and have successfully been used for some time. The oldest method of deciding about protonation sites, by consideration of acid-base strengths, has also seen useful new developments in the more sophisticated approach of correlation analysis. Molecular orbital calculations, which have become easier with more extensive use of computers, have proved useful in corroborating conclusions about protonation sites reached by other methods and in predicting unknown protonation sites. In fact, experiment and theory may be said t o have interacted in this area t o produce greater knowledge and understanding of the whole question of protonation sites. The question of protonation sites is one of the basic questions in the behaviour of complex organic molecules in solution, since protonated molecules are intermediates in synthetic organic chemistry, and the knowledge of protonation sites is important for the theory of reaction mechanisms of acid-catalysed reactions. It is also of fundamental importance for structural theory in general, since it is intimately connected with the concepts of mesomerism, electron density and bond polarization. In conjugated molecules one or other of the possible protonation sites may be more or less favoured by solvation effects and for this reason sites of protonation are often solvent dependent. In some instances, similar stability of two possible cations results in tautomeric equilibria and these too may be solvent dependent. Just as solute-solvent interactions have an effect on the relative stability of two possible cations formed from a conjugated molecule, so in solid salts stability relationships depend on the mode of packing of ions, which determines interactions with the nearest neighbours. Therefore the types of cation observed in solid salts are not necessarily the most stable ones in solution. The present review article will be primarily concerned with the structure of cations in solution. It will deal with the methods of structural investigation, with the nature of the acidic solvents used and their effect on cation stability, and with the results obtained so

ALTERNATIVE PROTONATION SITES

269

far for numerous conjugated molecules in various media. Some implications of these results for the foxmulation of reaction mechanisms will also be discussed. The most general formulation of the systems under discussion may be given by [ 13 or 121 in which there is a conjugative (mesomeric)

,E Z

\Y(-)

-

+

z<..-

X-

Y(-)

[21

interaction between X and Y, and Z may be a more or less complex, usually carbon containing$ructure, inclusive of rings. The group X is the source of the electron pair and the structure -Z = Y may be viewed as an electron sink. X and Y may both be N or 0, or N and 0 (or S), or N and C, or 0 (or S) and C, and depending on the valency of these elements, may carry substituents or may be linked through two or more other atoms to form a ring. In one of the cations formed by proton addition [3] conjugation is eliminated, and in the other [4]it is preserved. The structure of Z

H [41

may be such as to allow further spreading of charge in the cation

t41A number of methods used to decide which type of cation is formed rely upon spectroscopic comparisons of protonated molecules with the methylation or alkylation products in general ([3] and [4],with Me or R instead of H). In alkylation products there is no mobility of the group R and the structure of the ion is thus fixed. There is no possibility of mobile tautomeric equilibria. The properties of the ?r-electron system, which determines nmr coupling constants and ultraviolet/visible absorption (and, not quite

270

M. LILER

so cleanly, also the vibrational frequencies), may thus be compared with those in the conjugate acids in order to identify the structure of the cation as [ 3 ] or [ 4 ] . This article will not be concerned with orders of protonation in molecules with several basic centres, each possessing a lone electron pair independently of the others, but only with preferences in the protonation of two mutually conjugated centres in terms of the definition [ 11 or [ 21.

2. METHODS OF INVESTIGATION Spectroscopic Methods Nuclear magrtetic resonance spectra Nmr spectroscopic methods of determination of sites, of protonation may be divided into two types, depending on the kind of phenomena that are observed upon protonatim. The first type is concerned with the spectral changes (chemical shifts, coupling constants) in the skeleton of the molecule Z caused by protonation; the second is concerned with spectral changes at the protonation sites, X or Y (proton exchange phenomena, coupling to the captured proton, and the observation of the resonance of the captured proton itself). Spectral changes in the skeleton are observed when the base is changed into the conjugate acid, i.e., they are clearly observable at acidities of the medium which are some two logarithmic units of acidity greater than that required for half-protonation of the base (299% protonation). In the protonation equilibrium of a base B:

B + SH; + BH' + SH

(1)

where SHZ is the lyonium ion of the solvent SH, the chemical shifts of the protons in B and the coupling constants of the protons within B (other than those involving exchanging protons or the captured proton) are weighted averages of the values in B and BH' and their values may be used to determine the extents of protonation. The changes in coupling constants arise from changes in bond orders in Z (or Z-X and Z-Y) and sometimes in molecular geometry (e.g. free or restricted rotation around bonds, such as Z-X). The coupling

ALTERNATIVE PROTONATION SITES

27 1

constants in the cations may be used for structural assignment only by comparison of the values observed with those in model systems of known structure (either related conjugate acids of known unambiguous structure or methylation products with Me fixed to X or Y). Comparisons of chemical shifts are also possible, but must be made with more caution, especially when methylation products are used, because the chemical environment of protons close to the methyl groups will be different from that in the conjugate acids (i.e., only chemical shifts of the protons sufficiently far removed from the protonation and methylation sites may safely be compared). The spectral phenomena accompanying proton exchange and proton fixation are an even more important source of information on protonation sites. In the limit they include the observation of the resonance of the captured proton itself and couplings of other protons to it. This becomes possible only when the captured proton spends a sufficiently long time at its site of attachment. The acidities required for this may be some ten logarithmic units of acidity higher than the acidity at half-protonation, i.e., the complete nmr spectrum of the stable cation is not so readily observable as are the limiting chemical shifts and coupling constants of the skeleton of the conjugate acid. This is because the time scale of nmr observations is long compared with individual acts of proton exchange. The theory of rate phenomena in nmr spectroscopy has been fully worked out (see e.g. Pople et al., 1959, Connor and Loewenstein, 1963, or Reeves, 1965) and only the most important conclusions referring to proton exchange will be mentioned briefly here. In any protonation equilibrium (1) there are two limiting conditions: one at low acidity, where the protonation is very slow and the base exists free, and one at high acidity, where the deprotonation of the conjugate acid is very slow. Under such limiting conditions separate resonances of the base and the solvent or of the conjugate acid and the acidic solvent are observed, unaffected by proton exchange. The full spectrum of B or of BH' (inclusive of the resonance of the captured proton) is expected to be observable. Analogous situations then arise when the rates of protonation or deprotonation increase by varying the acidity of the solutions. For example, for the protonation the rate expression (2) may be written, where T is the mean lifetime of a proton on B between exchanges.

M. LILER

27 2

The rate is seen to depend on acid concentration. The values of T may be obtained from the appearance of the spectra, from collapsing chemical shifts or spin-spin multiplets (bearing in mind that T,x,h and rSspin are usually related through a statistical factor) (Saika, 1960), using the diagrams produced by Gutowsky and Holm (1956) or the corresponding approximate formulae, or total line-shape analysis. Similarly, where the spectra of the cations are concerned,

0

I

2

3

4

5

6

7

e

9 p.p.m.

10

Figure 1. Nmr spectra of the methylammonium ion (pKa = 10.5) at various pH-value% showing the effects of exchange (recorded at 60 MHz and 33°C). After Cmnwald et af., 1957.

the rates of proton abstraction increase with decreasing acidity of the solutions, and the observed resonances again show collapse of couplings to the captured proton and possibly some other changes as well. A good example of the latter is the spectrum of the methylammonium ion at various pH values (Figure l), first reported by Grunwald et al. (1957). It is important to note that the fully developed spectrum of the cation is observed only at an acidity some 1O'O times greater than the acidity at half-protonation. With increasing rates of exchange the width of the solvent line increases, because this now becomes an averaged resonance of all the exchanging protons. At fast rates of exchange no separate resonance of the captured proton (BH') is observable, and the rate of exchange may be obtained from the broadening of the solvent line.

ALTERNATIVE PROTONATION SITES

273

When secondary or tertiary amino groups are involved in conjugated systems (X = NHR or N R 2 ) , restricted rotation results at sufficiently low temperature, and there may be in general two isomeric forms present in equilibrium (3). Protonation on N then

binds the electron pair involved in conjugation, and the Z-N bond becomes a single bond with free rotation around it. This means that there is averaging of the resonances of the groups R in the two non-equivalent positions and only one resonance line for these groups is observed. In addition, in secondary amino groups, there is a collapse of couplings of the NH proton t o the protons in the group R owing t o NH exchange. The appearance of the -NR or -NR2 resonance under the conditions of rapid exchange affords n o information about the extent of protonation (only chemical shift measurements afford this), and the full spectrum o f the Nprotonated cation may be expected to be observed only at acidities some 10' times greater than the acidity at half-protonation. The rates of N-protonation are obtainable from the collapse of the NH-R coupling or from the collapse of the -NR2 chemical shift with increasing acid concentration. Protonation at the atom conjugated to the nitrogen (Y) does not destroy the restricted rotation around the Z-N bond and may indeed increase the rotational barrier. The rates of protonation on Y may in general also be obtained from couplings to any protons in Z, but the rates of protonation on a carbon already carrying a proton may also be obtained by the indirect method of deuterium exchange. The integral of the CH resonance decreases as deuterium is substituted in the exchange process, and also any couplings to the CH proton collapse with deuterium substitution. A conjugated system which involves three kinds of basic centre (-NH2, carbonyl 0 and ethylenic CH) is found in 0-aminoalkenones [ 5 ] . The methods described above, applied t o such molecules H

274

M. LILER

(Kramer, 1966), have shown that in general the rates of protonation follow the order ko >> kN 3- b - c .In view of this sequence of protonation and proton exchange rates, the resonances of the captured proton are observed most easily when protonation occurs at carbon, and with greater difficulty when it occurs at N or 0 atoms. This means that in order to observe the resonance of the captured proton, higher acidities of the medium and lower temperatures are needed for basic centres of comparable basicity on a nitrogen or an oxygen atom than on a carbon atom. There is a further difficulty in the observation of a captured proton on a nitrogen atom, namely the broadening of the resonance due to quadrupole coupling to the 4 N nucleus ( I = 1 ) (see Figure 1). This may be eliminated either by double irradiation techniques or by substitution of N for N in N nucleus has a spin of f ,which means that the the molecule. The resonances of the protons attached to it will be split into doublets, with coupling constants characteristic of the state of hybridization of the nitrogen orbitals. It is important to emphasize that the presence of two or three basic centres of comparable basicity in a conjugated molecule may lead to protonation processes at all of them. These processes may occur at various rates, but if all the rates are fast, nmr spectra d o not enable us to say what are the relative amounts of variously protonated species present at equilibrium. Increasing the acidity of the medium and lowering the temperature in order to observe the resonances of the captured proton may have the effect of shifting the tautomeric equilibria in favour of one or other of the protonation sites. Information on the position of tautomeric equilibria in protonation processes is thus not obtainable from nmr spectra under the conditions of rapid proton exchange. In connection with the question of stabilization of the cation and the possibility of observation of its fully developed spectrum, it is of interest to mention observations on the spectra of variously substituted methyl- and dimethyl-anilinium ions in trifluoroacetic acid (Giger and Simon, 1970). These studies have shown that only bases with pK,-values greater than 4- 5 give sufficiently stable cations for the observation of couplings to the captured proton in that medium. The H , value of trifluoroacetic acid has been reported as -3-03 (Hyman and Garber, 1959), i.e., an acidity higher by some 7-8 logarithmic units is here needed for sufficient cation stabilization. The Ho values of some other commonly used strongly acidic media are given in Table 1.

'

ALTERNATIVE PROTONATION SITES

275

TABLE 1 Ho-values for Some Strongly Acidic Media Commonly Used in Protonation Studies by Nmr Spectroscopy Medium

HO

Reference

CF JCOOH 100%HF 100%H2S04 Pure HS03F HSOj F-SbFS (equimolar)

-3.03 -9.9 7 -1 1-94 -15.07

Hyman and Garber Hyman and Garber Ryabova et al. (1966) Gillespie and Peel (1971)

-17.5 to -18

Olah et al. (1970d)

Superacid solvent systems have recently been reviewed by Gillespie and Peel (1971), who report Ho data for a number of superacid mixtures.

Ultraviolet and visible spectra The light absorption act is fast compared with proton exchange rates and therefore the ultraviolet spectrum of a system in protonation equilibrium (1) is the superposition of the spectra of the base and the conjugate acid, each contributing according to its concentration. So the observed molar extinction coefficient is given by (4), where n! is the fraction of the base converted into the

conjugate acid. Similarly, when the base is converted into more than one kind of conjugate acid, the observed absorption is the sum of contributions of all the species present. The unravelling of the spectrum may be complicated in such instances, and introduction of substituents which shift the tautomeric equilibrium may be helpful in drawing conclusions. The spectrum of the cation expected for protonation at one or other site of protonation may be obtained by synthesizing the two methylated products, with methyl groups alternatively attached t o each of the two sites. Weak hyperconjugation, of which methyl groups are capable, does not usually disturb the system of conjugated bonds any more than does a proton. The effect of NHS groups is also similar to that of methyl groups. Thus the spectra of two alternative cations may be obtained using model compounds in similar or identical solvents to those in

276

M. LILER

which protonation is studied, and hence conclusions about the nature of the protonated species drawn. The same principle can be extended to second and third protonations of polybasic compounds.

Infrwed and Raman spectra Uncertainties that beset the infrared method of structural determination in general, namely, ambiguities of assignment of absorption frequencies to particular elements of structure, are enhanced in protonation studies by the fact that many frequencies may be shifted by proton addition. Furthermore, the method can be applied primarily t o solid salts and to solutions in aprotic solvents. Studies in aqueous solutions, even when possible, are not generally useful owing to the broad absorption band of the solvent itself, which covers the spectral region above 3000 cm-' , in which the -NH: and the -OH group frequencies in many protonated compounds may be expected to appear. Also, conclusions reached from spectra of solid salts do not necessarily apply to cation structure in solution, especially when protonation at two basic centres of similar basicity is involved, because second order stabilization effects (crystal structure, packing) may favour one type of cation in the solid, whereas the other form may be more stable in certain types of solvent. Probably the most obvious evidence that protonation in solution and formation of solid salts may involve different interactions is the fact that, as well as salts in the molar ratio of acid to base of 1:1, solids in molar ratios 1:2 and 2:l are often obtainable. Interactions in some of these solids may be of a hydrogen bonding rather than a proton transfer type, and, if proton transfer is present, there may be association between BH' and A-, depending on the structure of A (e.g., there is usually association with C1-, but not with ClO, or SbCl,). However, if examination is restricted to salts of the 1: 1 type and the uncertainties in assignments are reduced by isotopic substitution, infrared studies of salts can in many instances provide information on protonation sites which is in agreement with conclusions reached from other evidence in solution. Raman scattering is a method that has been used to a small extent for solutions so far, but is likely to gain in popularity with the introduction of laser light sources, because it is capable of quantitative estimation of the species involved. A limitation is the necessity to use rather concentrated solutions, which can be a source of complications when solution equilibria are involved, but this

ALTERNATIVE PROTONATION SITES

277

necessity may become less compelling with increasing intensity of available light sources. The application of this method t o solid salts has been fairly common, but again the relevance of such studies t o phenomena in solution is somewhat limited, as has been pointed out above.

Correlation Analysis Examination of acid-base strengths is one of the oldest methods of structural investigation of molecules, which possess acidic or basic properties (see e.g., Brown et al. 1955). Its first applications to questions of alternative protonation sites lie at the very root of the theory of resonance and date from the end of the 19th century, when data on acid-base strengths became more readily aviilable following the development of the theory of electrolyte solutions. Amongst the oldest and classical examples of this are the enhanced basicity of guanidine (Ostwald, 1886), as compared with other amino bases, and the enhanced basicity of y-pyrone (Collie and Tickle, 1899) as compared with other oxygen compounds; both of these examples were eventually interpreted in terms of cation structures of low energy due t o mesomerism or resonance ( [ 6 ] and [ 7 ] ,

respectively). Alternative cation structures would have led to much lower basicity, by analogy with other simple nitrogen and oxygen bases. Relationships of acid-base strengths are in fact reIationships between free energies of reaction in terms of the thermodynamic

2 78

M. LILER

equation ( 5 ) , where AGO is the standard free energy change in going from the reactants in their standard states t o the products in their AGO = - R T In K

(5)

standard states, K is the equilibrium constant, R the universal gas constant, and T the absolute temperature. For acid-base reactions involving neutral bases, it is usually preferable to write the reaction as ( 6 ) , i.e., as the ionization of the conjugate acid, rather than the

BH++ HS + B + H2S+

(6)

reverse, equation (1). pK, , the negative logarithm of the equilibrium constant defined in (7), may then be introduced into equation ( 5 ) ,

yielding (8). AGO = 2.303 RTpK, In view of equation (8), any relationships observed between pK, -values may be termed free energy relationships. More specifically, linear free energy relationships are observed between AGOvalues of various analogously substituted acid-base systems, and these are often useful in determining protonation sites. Also, insofar as a series of compounds, which protonate on the same kind of atom and yield cations of similar structure, would be expected to undergo protonation to extents which are determined by the electron densities on this atom, correlations may be expected between the AGO-values and spectroscopic properties of molecules, which also depend upon these electron densities. The method of arriving at conclusions about structure and reactivity from such relationships has in recent times become known as correlation analysis. Correlation analysis depends on quantitative information (pK, -values, substituent constants, spectroscopic parameters) and is deductive in character. An outline of its methods and examples of some relationships, which will be useful in subsequent discussions, will now be given under three headings.

Some general pK,-relationships One of the fundamental questions in considerations of acid-base strengths in solution of molecules involving mesomerism is the rela-

279

ALTERNATIVE PROTONATION SITES

tive importance of this effect compared with inductive effects affecting the acidic or basic centre. The inductive effect, understood as a bond polarization effect, is always present whenever an acidic or basic centre is linked to a group which also enters into quantum mechanical resonance with it. For example, when the acid strength of alcohols (pK, = 15-16)is compared with that of phenols (pK, = lo), then part of the enhancement of acidity may be ascribed to an inductive electron withdrawal by the phenyl group (since sp2 hybridized carbon atoms are more electronegative than sp3 hybridized) [ 8a] and part to quantum mechanical resonance [ 8b], which

6 6.--6-6 \

-

\

(2 structures) [8bl

[gal

delocalizes the negative charge in the anion. Quantum mechanical resonance leads to an enhancement of the negative charge at the ortho and para positions of the benzene ring, which thus become potentially competing protonation sites with the oxygen. There is no evidence that any protonation at these sites actually occurs and phenols are known to contain the -OH group; the mesomeric effect [8b] is not sufficiently important and the negative charge, which tends to reside on the most electronegative atom, does so predominantly in the phenolate anion. The relatively small enhancement of acidity of benzoic acid (pK, = 4.18)as compared with acetic acid, (pK, = 4-75)suggests that the inductive [gal and mesomeric effects [9b] in this system act in opposition and largely compensate each’ other (Brown et al., 1955). Changes in electron density at the acidic or basic centre inevitably lead to differences in solvation, and these are especially important in the charged species of the conjugate acid-base pairs. In most instances stabilization of ions by solvation O& /;o

b-o+-b

o\

,o-

-wc/0-

(2 structures) t9bl

-o\

/o-

+

280

M. LILER

may constitute only a minor factor in determining acid-base strength, but it becomes a factor which must be reckoned with when close comparisons are made of acid-base strengths in similar systems. This is especially true when solvation occurs through hydrogen bonding. In the two systems so far discussed it is impossible to obtain a quantitative idea of the relative importance of the inductive and resonance effects because it is impossible t o achieve the operation of one of the effects without the other. When nitrogen is the basic centre, this becomes possible by steric fixation of the nitrogen lone pair orbital in the plane of the benzene ring, which virtually eliminates its overlap with the a-electron orbital of the ring carbon and hence also the mesomerism. So the enhanced acidity of the anilinium ion (pK, = 4-62) as compared with methylammonium (pK, = 10.67) has been shown (Wepster, 1952) t o be half inductive and half mesomeric in origin by a consideration of the following systems ( [ 101 - [ 121 ):

Quinuclidine [ 101 is a typical tertiary amine, in benzoquinuclidine [ l l ] only the inductive effect of the benzene ring is operative (mesomerism is sterically inhibited) and finally in dimethylaniline [12] both effects are present. All three systems yield tertiary ammonium ions (i.e. ions with comparable hydrogen bonding solvation) and the pK, -differences thus suggest approximately equal inductive and mesomeric effects upon electron density on the nitrogen. In [ 121, the benzene ring acts as an electron sink with respect to the amino group, which is the source of the electron pair. The total reduction in the basicity of the amino group in aniline as compared with aliphatic amines has at one time (Pauling, 1939, p. 207) been ascribed to the mesomeric interaction, as in [ 121, but the above analysis of the pK,-values of compounds [ 101 -[ 121 shows that this view is unjustified. In general, the relative importance of electron withdrawal by inductive and mesomeric effects is not known, although some

ALTERNATIVE PROTONATION SITES

28 1

estimates, like the one above, are available in favourable cases for systems containing nitrogen as the electron source. Indirect information about the importance of the mesomeric effect, which makes an electron pair less available at the electron source (X in [ 11) and more available at the electron sink (Y in [ l ] ) , is obtainable from experimental determinations of protonation sites. There resides the primary interest in such studies. The findings may be compared with theoretical calculations by molecular orbital methods (page 288), and more often than not the theory and the experiment reinforce each other.

Substituent effects-the Hammett equation Both steric and electronic effects of substituents upon experimental pK, -values provide information about protonation sites. In view of the small steric requirement of the proton, steric effects in protonation arise most often from steric inhibition of solvation of cations, which results in a reduction of their stability (pK, smaller). More important and more widely applicable criteria of protonation sites are available in the effects of more distant substituents, in particular metu and para in benzene derivatives, which are as a rule purely electronic in origin. Exceptionally, they may be complicated by steric effects (e.g., in polysubstituted derivatives). The most widely useful approach is that of Hammett (1940),who suggested that the effects of substituents on the ionization constants of benzoic acid may be taken as a measure of their effectiveness in other systems involving other reaction centres and in reactions other than acid-base equilibria. He thus defined substituent constants, u, by the equation

KO, KH

u = log--

where KO, is the ionization constant of the substituted benzoic acid and K& is that of the parent acid. Depending upon the position of the substituent, meta or para, for each substituent X, there are two constants, u, and u p . Hammett assumed that in reactions involving different reaction centres the effect of the substituents will be proportional to that upon the ionization constants of benzoic acid, so that for equilibrium constants, K , the relationship (10) would log Kx - - pa KH

282

M. LILER

hold, where p is a constant independent of the substituents, but characteristic of the reaction (reaction constant). Linear relationships between the logarithms of the equilibrium constants for a reaction and u (i.e., the logarithms of the ionization constants of benzoic acid) are, in fact, in view of equation ( 5 ) , also linear relationships between the standard free energies of the two types of reaction. The Hammett equation, of course, also applies to rate constants, but in questions concerning the sites of protonation, it is only the acid ionization constants (pK, -values) that are of interest. Good adherence to the Hammett equation has been found in many systems, especially for meta substituents. This is because meta substituents interact with the benzene ring both inductively and mesomerically, thus altering the electron density at the site of attachment of the reaction centre, but do not enter into direct conjugation with the reaction centre itself. The nature of the interaction is thus determined by the benzene ring and the substituent only. The effect on the electron density within the ring is transmitted to the reaction centre purely inductively. Para substituents, on the other hand, do enter into direct conjugation with the reaction centre via the benzene ring (cross-conjugation or through resonance) and the interaction is, therefore, dependent upon the nature of the reaction centre as well. A single set of up constants has thus proved inadequate to correlate all types of reactions, and significant deviations from the Hammett equation have been found in two types of system in particular. These will be discussed here, drawing examples from acid-base reactions only. A para-substituent may stabilize mesomerically either the conjugate acid of an acid-base pair rather more than it stabilizes benzoic acid, or it may stabiIize the conjugate base rather more than it stabilizes the benzoate anion. The first situation is found in carboxonium ions [ 131, where the delocalization of the positive charge on to a mesomerically electron-donating substituent stabilizes the cation. A similar resonance in the benzoic acid molecule [14] involves a separation of charge and affects the binding of the proton

ALTERNATIVE PROTONATION SITES

283

less directly. The effect of X on the pK,-value of benzoic acid is therefore less pronounced, which means that the up constant derived from this system will not be adequate in Hammett plots of pK,-values for aromatic ketones, for example. These systems require more negative substituent constants for mesomerically electron-donating substituents than Hammett’s up-constants. Such constants have been evaluated from studies of electrophilic aromatic substitution (Stock and Brown, 1963) and are called oi-constants. Conversely, if a para substituent stabilizes the conjugate base of an acid-base pair rather more than it stabilizes the benzoate ion, more positive substituent constants are required to achieve linearity in Hammett plots. Examples of this are acid dissociations of phenols and anilinium ions, where mesomerically electron-withdrawing substituents (Y = -NO,, - E N ) are more effective in enhancing acid strength than they are in benzoic acid, because charge delocalization of the type [ 151 is not possible in the benzoate anion.

Y

0 Y-

A criterion of an unknown site of protonation may then be obtained in some instances by considering which type of substituent constant leads to a better correlation of log K , values with substituent constants. The type of constant used constitutes evidence for one type of cation rather than another. In cases of possible tautomerism, it is the dominant cation that determines the type of the correlation. Substituent constants devised for substituted benzenes are often useful for substituents in meta-type (1, 3) positions in other ring systems and good correlations provide corroborating evidence for

284

M. LILER

postulated protonation sites. In some non-cyclic conjugated systems correlations with 0; constants may prove useful. The type of interaction measured by 0, constants is rather more widespread than their original definition implies and extends to other systems in which substituents are carried on an sp2 -hybridized carbon. An example of this is given in the next Section, and further examples are discussed on pages 303 and 306.

C o v e la tio ns with spectroscopic properties It is the stretching frequencies of bonds that may be expected to show correlations with electron densities, because the resonance contribution of the polar form of the bond will determine t o some extent the strength of the bond. When bonds involving hydrogen are concerned the contribution of the polar form of [ 161 may also be

expected to be reflected in the acidity of the compound. Hence correlations may arise between stretching frequencies of such bonds and standard free energies of ionization (pK,-values). Such correlations with stretching frequencies of the O-H bond have in fact been observed for carboxylic acids (Goulden, 1954) and phenols (Bavin and Canady, 1957; Canady, 1960), but it has been found that compounds of structurally different types fall on different lines. For double bonds of the type [ 1 7 1 , the contribution of the polar form \

,Z=Y

-

\,z +

-Y-

~171

determines on the one hand the strength of the bond, and, on the other, the availability of the negative charge on Y and hence its basicity. The carbonyl group being a group of this type, it is not surprising that correlations of basic strength of carbonyl compounds with carbonyl stretching frequencies have been found, in the first instance for variation with ring substitution in substituted acetophenones (Stewart and Yates, 1958) and benzoic acids (Stewart and Yates, 1960). More general correlations of the basicity of carbonyl compounds of the type R-CO-X (X being H, CH, and a variety of

ALTERNATIVE PROTONATION SITES

285

groups capable of mesomeric electron donation) and the carbonyl frequency in both the aliphatic and the aromatic series have also been observed (Liler, 1967; Liler, 1971c, p. 134). These correlations are very good for all carbonyl bases with the exception of amides (the one for aromatic compounds is shown in Figure 2). This fact 12

-

//

-PK,

0 -

p

2 I

qcOcm-'

N", I

I

I

Figure 2. l h e correlation of PKBH +-values of aromatic carbonyl compounds RCOX with carbonyl frequencies.

assumes special significance with regard t o the question of the site of protonation of amides, especially in conjunction with some observations on the carbonyl frequency itself. is a massThe carbonyl frequency in compounds R-CO-X insensitive vibration (Fuson et al., 1954) and is known to be determined by the nature of both the group R (conjugation leading to lower carbonyl frequencies) and the group X. The interactions of the groups X with the carbonyl group are analogous to those with the benzene ring. This follows from the fact that carbonyl frequencies in compounds R-CO-X ( R = alkyl or aryl) are correlated with the urn-constants of groups X (Figure 3). I t was stated in the preceding Section that groups X interact with the benzene ring inductively and mesomerically [ 181 , the relative importance of the two effects being different for different groups X. It is the combination of these two effects that is reflected in the changes in electron

286

M. LILER

I800

-

-

1750-

0

0

U

P

1700

-

1650

-0.4

- 0.2

0

0.2

0.4

'Jm

Figure 3. The correlation of carbonyl frequencies of compounds RCOX with Omconstants of groups X: A, acetyl compounds; B, benzoyl compounds (Liler, 1967; reproduced with permission from Liler, 1971c, p. 90).

density within the ring, of which um-constants are a measure. A correlation of carbonyl frequencies with u, -constants of groups X thus represents evidence that the same combination of inductive and mesomeric effects determines the polarity of the carbonyl group [ 193 (Liler, 1967). It can be seen from Figure 3 that the NH2 group

obeys this correlation and amides are thus not exceptional. Their deviation from the correlation of pK, -values vs. carbonyl frequency (Figure 2) is thus a strong indication that under the conditions of pK,-determination, protonation does not occur on the carbonyl oxygen, but that the N-protonated form greatly predominates. Oxygen basicity of amides may be predicted from these correlations with reasonable approximation. The question is further discussed in Section 4, page 333. This is an example of the usefulness of correlations involving spectroscopic properties in gaining information about protonation sites.

ALTERNATIVE PROTONATION SITES

287

3. CATION STABILITY AND SOLVATION This section will be concerned with several aspects of cation stability, especially with relation to the solvent medium. Molecular orbital calculations offer a method of estimating the relative basicity of protonation sites in a conjugated molecule, i.e., the relative stability of the two (or three) kinds of cation. This method involves certain assumptions, including a choice of parameters for the heteroatoms, which is up to a point arbitrary. Its predictive value is circumscribed further by the fact that the calculations refer to structures in the gas phase. When such structures are placed into a solvent medium, interactions occur which may modify significantly the relative stability picture that emerges from molecular orbital calculations. The nature of the medium determines to a certain degree the kind of cation that is experimentally observed. In cases of closely similar stability of two kinds of cation, tautomeric mixtures may be present in some media, while one or other tautomer may be favoured in others. The rates of protonation, which involve polarizability effects, are also medium dependent because the protonating agents in various media can be more or less polarizing. Faster rates of protonation at one site rather than another may lead to transient formation of the less stable kind of cation, which in time converts to the more stable cation at equilibrium. All these questions require more detaiIed discussion.

In trinsic Cation Stability The intrinsic stability of molecules and ions is determined by the bond energies of the o- and R-bond frameworks, and, in conjugated systems, by the resonance or delocalization energy of the R electron system. Resonance between several valence bond structures leads t o changes in bond lengths and hence in the energies of the o-bond framework. The resonance stabilization and the o-bond energies are thus not separable. The assumption is commonly made that the compression energy involved in changing the bond lengths from the values in model compounds to those found in the conjugated system may be included in the value of the resonance energy (Wiberg, 1964). The resonance energy is thus an ill-defined concept. Even in the most favourable cases, such as benzene, resonance energy is not

288

M. LILER

obtainable from purely theoretical considerations and resort must be made to thermochemically determined delocalization energy (from heats of combustion and hydrogenation) in order t o estimate the value of the resonance integral. But it must be remembered that in obtaining resonance energies from thermochemical data assumptions are also involved (Pauling, 1939). Such an assumption is the constancy of bond energies which are used in the calculation of the theoretical bond energy of a molecule of given structure. This assumption is useful only as a first approximation and therefore many thermochemical resonance energies are n o more than rough estimates. Resonance energy may either be lost or retained upon proton addition to a conjugated molecular framework. The loss of resonance energy occurs when one of the electron pairs which participates in the conjugated system is localized by proton addition, so that the remaining 'IT-electron system involves one less atom and two less electrons (Streitwieser, 1961). The resulting change in the A-electron energy is called the localization energy and is given by (1 l ) , where

A& = (&)BH+ - & ) B

(11)

E , stands for the R-electron energy, and the subscripts BH' and B refer to the protonated and the unprotonated system, respectively. Since the unprotonated system is more resonance stabilized (lower in energy) than the protonated system, the localization energy is a positive quantity (A& > 0). Molecular orbital theory enables calculations of this energy to be made for aromatic hydrocarbons, assuming that the resonance integral remains unchanged by the positive charge. The theoretical aspects of protonation of hydrocarbons have been reviewed by Perkampus (1966) and will not be discussed here. There is, however, an important relationship between the localization energy in such systems and the basicity, which is applicable t o other conjugated systems also. The thermodynamic relationship between the standard free energy of protonation and the basicity constant K b is

AGO = - RT In Kb

(12)

On the other hand AGO

= M - TAS"

where and ASo are the standard enthalpy and entropy changes for the reaction, respectively, and T is the absolute temperature.

ALTERNATIVE PROTONATION SITES

289

Now, in a series of closely similar protonation reactions, ASo may be assumed to be invariant, and then variations in AGO will reflect primarily variations in AHO. In condensed phases and for cations of closely similar structure, these will be practically equal to the changes in internal energy, i.e.

A(AG") 2 A(A@) Z AE,

(14)

Hence follows a linear relationship between pKb for aromatic hydrocarbons and the localization energies (the changes in the o-bond energies are the same in the protonation of any aromatic hydrocarbon). The first calculations of this kind are due to Gold and Tye (1952) and the correlations have been demonstrated by Dallinga et al. (1957) and by Mackor et al. (1958). From such theoretical calculations on aromatic hydrocarbons, it is possible to predict the preferred site of protonation, i.e., the most stable proton addition complex. Analogous theoretical calculations on conjugated systems involving heteroatoms are more difficult, because more numerous assumptions about the unknown Coulomb integrals and resonance integrals are needed. Conclusions about the relative basicity of two conjugated sites obtained from such calculations are therefore necessarily more tentative. Caution is also needed in linking up the idea of resonance energy with the changes in acid-base strengths of various sites in a molecule. For example, Pauling (1939, p. 207) suggested that the total reduction in the basicity of aniline relative to an aliphatic primary amine may be ascribed to the loss of the resonance energy of interaction of the amino-group with the benzene ring [ 201 upon proton addition, H

[201

(2 structures)

and calculated hence a resonance energy of 8.4 kcal mole-' (35.2 kJ mol-'). As has been pointed out on page 280, however, only about one half of the reduction in the basicity of aniline as compared with an aliphatic amine may be ascribed to resonance, the other half

M. LILER

290

being due to an inductive electron withdrawal. In benzoquinuclidine [ 111 the nitrogen lone pair is localized, and hence the change in the free energy of protonation of N,N-dimethylaniline [ 121 as compared with benzoquinuclidine [ 111 is approximately equal to this localization energy in view of the approximation (14). The estimated

A(AG") = 2.303 RTApK,

=

3-7 kcal mole-' (15.6 kJ mol-')

(15)

empirical resonance energy (additional to the resonance energy of the benzene ring) of 6 kcal mole-' (25 kJ mol-I) (Pauling, 1939, p. 139) is in moderately good agreement with this figure. The N-protonated cation possesses only the resonance energy of the benzene ring, which is considerable (39 f 3 kcal mole-' ) (Pauling, 1939, p. 129; Aston, 1955, p. 560). According to the above resonance structures [ 201, the alternative protonation sites in aniline are the ortho and thepara carbon atoms of the ring. Localization energies for these sites have not been estimated, but are certainly greater than that for N-protonation, since in the formation of a-complexes the great stabilization of the benzene ring is partially destroyed. The para-carbon localization energy of aniline is, however, certainly less than that of benzene. A pK, -value for para-carbon protonation of aniline of approximately -1 1 has recently been estimated by Ridd (private communication) from pK,-values of aromatic hydrocarbons, using the Hammett equation. When additional stabilization by other amino-groups is present, C-protonation becomes possible in aqueous solution (see page 356). Under aqueous conditions, the N-protonated cation owes part of its stability to hydrogen bonding solvation, which may contribute as much as 4 5 kcal mole-' (ca. 190 kJ mol-') of stabilization energy to the primary anilinium ion, in view of its localized charge (Sheraga, 1968). The C-protonated cation would certainly be less stabilized by solvation, since the positive charge is shared between the nitrogen atom and the ring [21] and C-H bonds are known to have poor hydrogen bonding ability. H

I

:N-H

ALTERNATIVE PROTONATION SITES

29 1

An alternative theoretical approach for finding the relative basicity of possible protonation sites in a molecule is the calculation of electrostatic potentials that these sites create in the space around them. The molecular wave function enables the calculation of the electrostatic potential V ( r ) for various points in a molecule which would attract a point charge q at a distance r , this being a basic model for proton addition to a base. The coulombic interaction energy q V is then a first approximation to the true energy of interaction. However, several effects remain ignored here. This treatment takes no account of the polarization of the molecule by the charge, nor of the possible changes in the geometry at the site of protonation upon proton addition. The inclusion of these effects (termed dynamic) has recently been attempted in some calculations on formamide (Mkly and Pullman, 1972), but even these more refined approaches ignore solvation.

Cation Solvation Most chemists’ ideas about the basic strength of common basic centres in organic molecules are shaped by the knowledge about the pK,-values of common organic bases in aqueous solution, i.e., in a medium which has quite exceptional solvating properties, especially for cations. For inorganic cations, this solvation may be understood in terms of charge-dipole interactions, but, for organic cations formed by proton addition to neutral bases, there is also hydrogen bonding of the cation to the solvent. This type of interaction has first been invoked in order to explain the anomaly in the pKa-values of aliphatic amines (Trotman-Dickenson, 1949). NH:

CH3NH;

(CH3)zNH:

(CH3 )3NH+

PKa

9-25

10.66

10-73

9.75

pK,(stat. con.)

9.85

11.14

11-03

9-75

After allowing for a statistical correction for the number of protons that may be lost from the cation, the order of acid strengths (per NH-proton) of these cations remains unchanged. The initial trend of decreasing acidity with methyl substitution is reversed in the trimethylammonium ion, because with a decreasing number of protons in the cation, its hydration stabilization by hydrogen

292

M. LILER

bonding to the solvent ([ 221 us. [ 231 ) decreases. Evidence for this has also been obtained from nmr studies on anilinium salts (Fraenkel

and Kim, 1966). This type of interaction is sometimes called chemical hydration, because of its stoichiometry (BH+-*OH2) and of the exchange forces involved in the quantum mechanical formulation of the hydrogen bond. Studies of protonation of mesomeric molecules, in which the common basic sites (N,O) show reduced basic strength and give rise t o alternative protonation sites in the molecule, often require the use of more acidic media. The most obvious extension of the dilute aqueous acid solutions (pH 2 1 ) are concentrated acids (No< l), which in the limit becomes anhydrous acids. Such strongly acidic solvents (100% H2 SO4, HS03 F, H S 0 3 F-SbF,-S02, HF, HF-BF3) have been extensively used in recent years, especially in protonation studies of very weakly basic molecules by nmr methods. This has been necessary in order to slow down the rates of proton exchange, as discussed on page 274. In a progression from dilute aqueous acids t o anhydrous acids, the solvating properties of the medium change, and it is no exaggeration to say that they change drastically. This change in solvating properties of the medium lies at the root of the theory of acidity functions (see, e.g., Rochester, 1970; Liler, 1 9 7 1 ~ ) . In a protonation equilibrium such as (16), water molecules are involved in solvating all the dissolved species, but especially the

B(H2O)j + H'(H2O)j

BH'(H2O)k

i=

+ ( i + j - k)H2O

(16)

hydronium ion and the conjugate acid BH+ (which may contain several hydrogen bonding protons). In view of the strong solvation of the proton (H,O:, j = 4), some water molecules are released in practically all protonation reactions ( h = i + j - k > 0). The equilibrium constant for the above reaction may be written as:

ALTERNATIVE PROTONATION SITES

293

all the species being assumed hydrated, y's being the activity coefficients on the molarity scale and a, being the activity of water. Since the acidity function H , is operationally defined by (18) and

the scale is anchored to the standard state of infinite dilution in water (a, = 1)' then it follows that K , = 1/K and pK, = log K . Hence Ho=logK-10g---ogW+

PI

I-

h

YBH+

YBYH+

+ l o g a, ~

CH

(19)

If the simplifying assumption is made that Y H * S Y B H + (since the hydrated proton and the BH+ ion are both large) and since activity coefficients of neutral bases do not vary much with acid concentration, then the major factor which determines the variation of the acidity function with acid concentration is the activity of water (a, vanes from 1 in infinitely dilute acid t o --oo in strictly anhydrous acids). These were essentially the assumptions of Bascombe and Bell (1957) and of Wyatt (1957), who showed that the acidity functions of a number of strong acids at equal molality of the proton are, in fact, a common function of the activity of water. For a long time the basic postulate made by Hammett and Deyrup (1932)' that y BH+/yB is the same for all neutral bases was n o t questioned and the hydration of the base and the conjugate acid BH' were not thought to be very important (e.g., Bascombe and Bell assumed that i = k = 0 in their treatment of the acidity function). The importance of the hydration of the conjugate acid BH' particularly became clear when tertiary anilinium ions were used as indicators instead of primary aniIinium ions (which were the chief indicators of Hammett and Deyrup), and generated an acidity function (H:') more negative than the original Hammett acidity function Ho (= HA) (Arnett and Mach, 1964). This could arise only if the Hammett activity coefficient postulate, that Y B H + / Y B is independent of the nature of the base, breaks down in this instance. This is what would be expected if chemical hydration of substituted ammonium ions plays a part in their stabilization in solution, because the more highly hydrated primary anilinium ions should become progressively more destabilized with decreasing availability of water in more concentrated acid solutions than the tertiary cations (yBp$" >ytge)ri? with increasing acid concentration). Using a generalized definition of an

294

M. LILER

acidity function (20), it can be seen that such a relationship between the activity coefficients of primary and tertiary anilinium ions leads ~H+YB Hi = - log -

Y BH+

to Mo > H t ' . The more complete definition of H o given above (19) shows that differences in cation hydration lead to differences in acidity functions also owing to different h-values ( h is greater for tertiary anilines than for primary, because k = 1 instead of 3, the differences in i being probably less important). Since the activity coefficients of cations, Y B H + , are not measurable quantities and h cannot be estimated with certainty, it is in general impossible to separate these two effects. The reality of differences in the variation of Y B H + with acid concentration for cations of different structure has been demonstrated by measurements of activity coefficients of a number of cations relative to tetraethylammonium ion (TEA') as a standard (Boyd, 1963). This was achieved by measuring the solubilities of salts of the type BH'PCP-, where PCP- is the pentacyanopropenide anion which does not become protonated in quite strong acid. Some results of such measurements are shown in Figure 4. Acidity functions differing from Ho have been found to arise from protonation of several other types of neutral base, such as indoles (Hinman and Lang, 1964) and amides (Yates et af., 1964). Acidity function studies have thus provided further evidence for solvation by hydrogen bonding of conjugate acids of neutral bases in aqueous media. A corollary of this, which is of importance in the present context, is that in conjugated systems with two alternative protonation sites, one may yield a cation which is stabilized more by hydration than the other in dilute acid. This cation will become progressively more destabilized than the other with increasing acid concentration, with a possible tautomerization into the other cation. Of the two cations [3] and [4], the one that possesses the more localized charge [3] will be expected to form stronger (and perhaps also more numerous) hydrogen bonds with the solvent. In other words, cations with localized charges [3], which are stabilized by solvation in the highly polar aqueous acids, may be dominant over the resonance stabilized forms [4], but the latter may emerge as dominant under anhydrous conditions. Some examples of this will be discussed in Section 4. Cations in other polar solvents (DMSO, alcohols, DMF, TFA, etc.), which are poorer hydrogen bond acceptors than water, will generally

ALTERNATIVE PROTONATION SITES

295

be less stabilized by solvation. This will favour, relatively, forms with more delocalized charges. Such solvents are very often used in protonation studies, firstly, because they are usually better solvents for organic compounds than aqueous media, and secondly, because proton exchange in them is slower than in aqueous solutions, which means that cations may be more readily observable by nmr methods.

Figure 4. Activity coefficients of cations relative to the tetraethylammonium cation in sulphuric acid-water mixtures. A-anilinium ions: 1 , unsubstituted; 2, p-Cl; 3, p-NOz; 4, m-NOz. B-benzamide cation. C-a typical carbonium ion, @-MeOC6H4)3C+. A and C, after Boyd, 1963; B, after Sweeting and Yates, 1966.

Also such solvents are easier to handle than the superacids (HSO, F, HSO, F-SbF, -SO,, HF-BF, , 100% H2SO, etc.). Finally, protonation studies are often attempted in non-polar solvents with additions of the strongest carboxylic acids (e.g. CHC1, or CH,C12 with CF3COOH or CC13COOH). The content of the polar component must clearly be sufficient for protonation, but mixtures with only small additions of these acids may have too low dielectric constants for the formation of free cations of whatever kind. Rather, ion pairs may result or indeed only hydrogen-bonded complexes.

296

M. LILER

This discussion, which emphasizes the difference between aqueous acids and other less solvating and less polar solvents as protonating media, has led to the conclusion that the protonation site is not solely determined by intrinsic factors and is not medium invariant: the more polar media generate cations with more localized charges more easily than the less polar media. The same conclusion would follow from the application of the principle of hard and soft acids and bases (Pearson, 1963). The hydronium ion, being positively charged, is a “hard” acid and would react with the hardest basic centre in a molecule, which would in most cases be the nitrogen (or oxygen), even though more or less conjugated to a carbon containing structure. Anhydrous acids or uncharged acids in polar solvents are “soft” and would protonate softer sites in conjugated molecules (which could be a carbon or an oxygen, which have acquired partial negative charge by conjugation with a nitrogen, oxygen or sulphur atom, and are located in a polarizable system of bonds). There is here an apparent analogy with alkylation or acylation of systems with alternative nucleophilic sites, the important difference being that in protonation, fast proton exchange ensures that the final product is always the thermodynamically stable form (or mixture of forms) under a given set of conditions. In alkylation and acylation reactions kinetically as well as thermodynamically stable products are obtained depending on the reagent and conditions. Such reactions have been discussed in detail by Gompper (1964).

Tautomerism In any kind of conjugated system with two protonation sites [ 11 there is the possibility of both cations [3] and [4]being formed in varying amounts, depending on conditions. As has been pointed out in the preceding section, solvation effects may play a dominant role here in tipping the balance from resonance-stabilized forms to solvation-stabilized forms. Thus under a given set of conditions we may have protonation at both sites as shown in equations (21) and

(22)-

t 21

[31

ALTERNATIVE PROTONATION SITES

297

Here HzS+is the protonating agent assumed to be the lyonium ion of the solvent SH, and KH and K; are the ionization constants of the two types of conjugate acid [3] and [4] in the solvent HS, as defined in (23) and (24).

and

The activity coefficients are referred t o some chosen standard state, in which they are taken as unity (e.g., in aqueous solution to the state of infinite dilution). The tautomeric equilibrium constant is then given by (25).

If tautomerism occurs in dilute aqueous acid, then Kk and K; will be the thermodynamic acid ionization constants and (26) will hold; thus pKT = pK: - p K i = ApK,

(26)

the position of the tautomeric equilibrium will be determined by the difference between the pK,-values of the two forms of cation (Mason, 1958). As pointed out on page 274, under conditions of rapid proton exchange (such as would exist in dilute aqueous acid), the position of the tautomeric equilibrium is not obtainable from nmr spectra, but it is obtainable from uv and visible spectra, provided that the two forms of cation differ sufficiently in their absorption and provided that both forms are present in amounts detectable by this method. Usually no less than 3% of one form can be detected with certainty, which means that tautomeric quotients of 0.03- 30 at most can be estimated (ApK, = k1.5). Thus observable tautomerism will arise when the basicities of the two protonation sites are fairly closely matched.

M. LILER

298

When some other factor also enters into the tautomerization equilibrium, such as a variable medium, then tautomerism may be observable over a certain range of medium compositions. Examples of this kind of behaviour may be found in concentrated aqueous acids, where varying degrees of hydration of the two types of cation can lead to a dependence of the tautomeric equilibrium upon the activity of water, as discussed in the preceding section. This happens in the tautomerization of amide and phenol (or anisole) cations and further discussion of this phenomenon will be deferred until later (see pages 333 and 372).

Kinetic us. Thermodynamic Stability As already mentioned on page 296, owing t o rapid proton exchange between the solvent medium and the two types of cation, it is always the thermodynamically stable cation (or cation mixture) that is observed in the protonation of a conjugated system. However, it is known from the work of Eigen and his collaborators (Eigen, 1964) that the rates of protonation at the two sites in a conjugated system, such as an enolate ion, may differ widely. The measurements in aqueous acid solution have shown that protonation which is associated with a minimum of electronic shift and which occurs at a site already hydrogen bonded to the aqueous solvent occurs at diffusion-controlled rates for bases which become protonated in the aqueous pH range. So for example, ammonia is protonated according to the equation (27) with a rate constant

NH3 (aq.) + H 3 0 +

kf

NH: (aq.) + H 2 0

kr

kf =

k H + = 4 - 3 x 10' M s-' , which corresponds to the diffusion rate constant. Since the equilibrium constant for this reaction in

dilute aqueous solution is given by (28) and the forward and reverse rates must be equal at equilibrium, then, from (29), it follows that

299

ALTERNATIVE PROTONATION SITES

the deprotonation rate constant is k, = 5.75 x x 4.3 x 10'O = 24.7 s A 1 .The reverse reaction is thus relatively slow and is activation controlled (since the proton goes from a more basic t o a less basic site, i.e. across a high potential energy barrier). In the protonation of conjugated systems, the rates of protonation at the two possible sites are different. So for example, the mesomeric anion of acetylacetone, [24], is protonated by the hydronium ion at

~ 4 1

the enol oxygen at the diffusion controlled rate (kH+ = 3.10 x 10'' M s - ' ) and at the carbon at a slower rate ( k H + = 1-2 x 10' M - l s - l ). This is because the oxygen centre is a better hydrogen bond acceptor than the partially negatively charged carbon and also because the protonation on oxygen involves less reorganization in the electronic structure of the anion (Eigen, 1964). Similar considerations apply t o neutral conjugated molecules. Much less is known about rates of protonation in nonaqueous media, but some of the general principles that apply to aqueous solutions probably hold in other media also. Thus, proton transfers t o the most electronegative atoms and to the best hydrogen bonding sites are probably faster than proton transfers to carbon. The finding mentioned earlier for enaminocarbonyl compounds, that k, S k N S k, (Kramer, 1966), is in line with these principles. This means that if there are two basic sites of comparable basicity in a molecule, the proton transfer t o oxygen or nitrogen will be faster than to carbon, and consequently relatively higher concentrations of the N- or 0-protonated cations may be formed initially. This may not be readily observable if protonation at all sites is fast relative t o the time-scale of the method of observation. An example has, however, been reported (Alais et al., 1971) of a slow tautomerization of an N-protonated cation to a more stable C-protonated cation, which was observed by nmr and which is worth mentioning because it shows an interesting medium effect on the relative rates. The protonation of N,N-dimethylisobutylenamine [25] can lead to two types of cation: the N-protonated

-'

Me\ C=CH-N, Me/

/Me Me

Me, +,Me ,C=CH-N, Me I Me

Me, ,CH-CH=N Me

+

,Me 'Me

300

M. LILER

enammonium ion [26] and the C-protonated immonium ion [27]. In 70% perchloric acid, in which hydronium ions are the protonating agents, the N-protonated cation is formed to the extent of 90% and the C-protonated cation to the extent of only 7% initially, but there is a slow transformation of the first into the second at room temperature. Carboxylic acids (acetic, trifluoroacetic) in DMSO and CDC13 lead directly to the C-protonated product, whereas hydrochloric acid in ether gives the N-protonated cation, which changes rapidly into the C-protonated cation at room temperature (Opitz and Griesinger, 1963). The C-protonated cation is thus the thermodynamically stable form in all these solvents, but the rates of C-protonation relative to N-protonation vary. This has been explained (Alais et a/., 1971) by the principle of hard and soft acids and bases and the situation is analogous to the alkylation and acylation of ambident nucleophiles, which was mentioned on page 296. Owing to its ionic character, the hydronium ion is “hard” and will attack faster that centre in the molecule which has the highest negative charge. Undissociated carboxylic acids in nonpolar and polar non-aqueous solvents are “soft” acids and will attack the “soft” partially negatively charged carbon at a relatively faster rate.

4. PROTONATION SITES IN CONJUGATED MOLECULES The results that have been obtained by the application of the methods discussed in Section 2 to various conjugated systems will be discussed in this Section, following a classification according to the identity of atoms X and Y in the general formula of the conjugated system [ I ] . Nitrogen being the most common basic centre, systems containing two nitrogen alternatives will be discussed first. Amongst them are some well-known systems, but as there has been some recent work on practically all of them (usually cation characterization by nmr), these will be discussed also, stating briefly the established position. This will be followed by N/O (or N/S) and N/C alternatives. Amongst this range of compounds are some very weak bases, which become protonated only in moderately concentrated acid. Molecules with two oxygen alternatives and O/C alternatives are usually still weaker bases. Overall, the basicities of the compounds to be discussed span a pK,-range of some 22 units, from pK, = 13.6 for guanidine t o pK, of ca. -8.5 for benzoic acid. Any pK,-values

301

ALTERNATIVE PROTONATION SITES

mentioned, for which no literature source has been indicated, come from the compilation and supplement by Perrin (1965 and 1972) and refer to 25"C, unless otherwise stated. The literature has been covered up to about mid-1973.

N / N Alternatives Systems containing two or more conjugated nitrogen atoms are numerous. It is convenient to subdivide them into open-chain and cyclic systems. As open-chain systems we shall regard those in which the two nitrogens are conjugated, possibly even through an aromatic ring, but are not part of a ring themselves. In cyclic systems one or both of the nitrogens may be part of a ring.

Open-chain systems The enhanced basicities of acetamidine [28] and amidines with conjugated double bonds linking the two nitrogen atoms [29] are H,N--C=NH

HRN--C(=CH--CH),=NR

I

I

CH3

(R = H or alkyl)

[281

~ 9 1

classical examples of the effect of resonance stabilization of symmetrical cations [30] and [31] upon basic strength (Schwarzenbach [H2NXJ;3NH2]+

[

HRNzCz(CHxH)n"-NRH HI

~301

[311

I

+

and Lutz, 1940). The resonance stabilization is not greatly dependent upon the chain length, since amidine has a pK, = 12-41 and glutacon-diethylimid ([31], n = 2, R = Et) has a pK,-value of 11.90. Information on the importance of resonance in the neutral base [ 321

M. LILER

302

has recently been sought by studying hindered rotation around the C-NMe, bond in N'-t-butyl-N,N-dimethylformamidine [ 331 by nmr (Harris and Wellman, 1968). In pyridine and toluene at room temperature there is rapid rotation around that bond, but at low temperature (below -15OC) the N-methyl singlet broadens and splits into a doublet (at -54'C). The activation energy for rotation, which has been obtained from the spectra by total line shape analysis [AG' = 1 2 f 0.8 kcal mole-' (50 f 3.3 kJ mol-')I , is lower than that in amides (see page 330). This activation energy is a reflection of the a-bond energy of the C-NMe, bond, since the a-bond is broken in rotation, but it is not its precise measure, because part of the activation barrier to rotation is due t o the repulsion of groups carried on the atoms forming the bond. The smaller value found in amidines than in amides may be ascribed to the smaller electron-withdrawing capacity of the imino-nitrogen as compared with the carbonyl oxygen . The structure of amidinium cations has been confirmed by infrared spectra of their salts (Mecke and Kutzelnigg, 1960; Grivas and Taurins, 1959). Since amidines are strongly basic, the NHprotons become sufficiently stable in dilute mineral acids to allow their observation by nmr (Neuman and Hammond, 1963). Thus acetamidinium NH-proton resonances, which are not observable at pH 5-6, become observable as a broad singIet in more acidic solutions. Also in the spectrum of N,N'-dimethylacetamidinium chloride in water there are two singlet N-methyl resonances (owing to methyl groups in two non-equivalent positions, as in [34] ), but in 14% sulphuric acid two doublet resonances (J = 5 Hz) arise, owing to splitting of the methyl resonances by NH-protons. The relaxation times of the two doublets are different. In > 70% sulphuric acid these doublets collapse to broad singlets and in 90% acid to a single singlet. These changes are due to the second protonation, which leads to cations [35], involving free rotation around the C-N bonds. The very high acidity required for the onset of this second protonation ( H , < - 6 ) is also a measure of the considerable reasonance I

pN-H CH,-CI;

I+

CI-

CHI
N-H +

ALTERNATIVE PROTONATION SITES

303

stabilization of the monoprotonated cation. The mechanisms of proton exchange of these cations have been discussed by Neuman and Hammond (1963). The alternative possibility of protonation of amidines on the amino-nitrogen depends on its basicity, after its lone pair of electrons has been localized. This localization energy is not known and neither is the inductive effect of the residue attached to the amino-nitrogen in [ 361 , but an amino-nitrogen attached t o an sp2 -hybridized carbon R-C<..

NH NH2

carrying an electronegative substituent would be expected to be less basic than ammonia, so that there is at Ieast a factor of lo3 in the basicity in favour of the resonance stabilized cation. It is perhaps worth mentioning that substitution of other groups for the C-methyl group in amidinium ions has an effect on their pK,-values, which is best correlated with Hammett’s u,,, -constants ( p = -11.98), and the same is true of C-substituted N-phenylamidinium ions ( p = -12.08) (Charton, 1965). This is in line with the u -constants are a measure of the suggestion made on page 286, that , inductive and mesomeric interactions of groups with’a carbon in the sp2 state of hybridization. The even greater basicity of guanidine (pK, = 13.6) as compared with amidines is due to an even greater resonance stabilization of the cation, since three equivalent resonance structures enter into resonance [6]. The base itself resonates between three nonequivalent structures [37], and has according to Pauling (1939,

p. 213) a resonance energy of 6-8 kcal mole-’ (25 - 33 kJ mol-’ ) less than the cation. The effect of N-alkyl substituents on the basicity is minimal (Angyal and Warburton, 1951), which shows the dominant role of resonance stabilization and the lesser effects of inductive electron release and cation solvation, probably owing to the delocalized positive charge. The resonance stabilization is such

U LILER

304

that further protonation is very difficult. Guanidine becomes diprotonated to the extent of only 27% in a 0.08 M solution of guanidinium perchlorate in 99.9% sulphuric acid ( H , = -11.43) (Williams and Hardy, 1953). Guanidine is diprotonated in the superacid system HS03F-SbFs (Olah and White, 1968b). At -80°C the NH2-groups give two non-equivalent NH-resonances, whereas the NHg-group appears as a singlet at lower field. In negatively substituted guanidines [38], i.e., guanidines with an electron attracting substituent, there is little doubt that the substituent is carried on the imino-nitrogen, since the protons would X

.N/x

I

:N:-

tend to move away from a negatively substituted site. Imino-nitrogen protonation leads here to the resonance stabilized cation, but, since the effect of the substituent on the resonance energy is difficult to assess, the question arises whether the relative basicity of the amino-nitrogens, which are less directly affected by the substituent, may not increase sufficiently to make them dominant protonation centres. Kessler and Leibfritz (1969) have shown that this is not so in the protonation of 1,1,3,3-tetramethy1-2-phenylguanidine in CDC13/CF3COOH. In the nmr spectrum at low temperature (-60°C) the two NMe2-groups (a) become non-equivalent and (b) split into two signals. This can only be explained by the assumption of hindered rotation around both C-NMe, bonds and around the C-NHPh bond in the cation [39]. The NMe2-group next to the CH3 I

CH3 1391

phenyl ring appears at higher field. The effect of para-substituents in the ring on the rotational barriers is rather small, smaller than in the corresponding guanidine bases, but negative substituents increase the rates of rotation.

ALTERNATIVE PROTONATION SITES

305

There has been some controversy over the site of protonation of nitroguanidine in recent years. Bonner and Lockhart (1958) studied the protonation of several nitroguanidines in sulphuric acid-water mixtures by ultraviolet spectroscopy and concluded that imino nitrogen protonation occurred. They first of all observed that the strong absorption band at 260- 270 nm [ E = (1 2- 17) x l o 3 ] , shown by all nitroguanidines, is quite different from the broad band at 225 nm ( E = 5900)in the spectrum of nitramide, which is consistent with the view that nitroguanidines do not exist in the nitramino form but in the nitrimino form [38] (X= NO,). The characteristic absorption of nitramide appears in the spectra of nitroguanidines in strong acid solution and reaches maximum intensity in 40-50% sulphuric acid, due to complete conversion of the nitroguanidines into their conjugate acids. Plots of the logarithms of the ionization ratios us. Ho yielded pK,-values ranging from -0.93 for nitroguanidine and -0.86 for 1-methyl-2-nitroguanidine to -1.20 for the 1,l-dimethyl derivative. The weaker basicity of the latter was ascribed to steric hindrance involved in achieving coplanarity in the resonance stabilized cation. This conclusion of imino-nitrogen protonation in nitroguanidines has been challenged more recently by Price et al. (1967),who studied the nmr spectra of 1,1,3,3-tetramethyl-Z-nitroguanidinein strong acid solution and failed to find evidence of restricted rotation in 37.5% hydrochloric acid down to -35°C. They also observed a splitting of the methyl signal by the captured proton (J = 3 Hz at -20°C to -3O"C), which they ascribe to a proton on the aminonitrogen, as in [40].This interpretation is not entirely convincing

CH3

because (a) a coupling constant of only 3 Hz is considerably smaller than vicinal coupling constants in similar situations in other compounds (which are always > 4 Hz and usually about 5 Hz), and (b) if the proton has a sufficient life-time on one of the NMe,-groups for this coupling to be observable, one would expect the resonance of

M. LILER

306

the other NMez-group to appear as a separate signal at higher field. This has not bken reported. It is known, on the other hand, that in strongly conjugated systems coupling may occur across more bonds than in singly bonded structures and that fixed geometry also tends to enhance it. It therefore appears possible that the observed coupling is a coupling across five bonds to a proton on the imino-nitrogen. A temperature of -35°C may not be sufficiently low to observe restricted rotation in the resonance stabilized nitroguanidinium ion, especially as it has been shown that the rates of rotation are increased by negative para-substituents in 111,3,3-tetramethyl-2-phenylguanidiniumions (Kessler and Leibfritz, 1969). Since the question of the dominant site of protonation of nitroguanidine does not seem to have been resolved by nmr observations, it is worth pointing out that information about this is also obtainable from a correlation of pK,-values of negatively substituted guanidines with known parameters of the substituents. Such correlations were already attempted by Charton (1965), who used uI- and urn-constants of the substituents and claimed a better correlation with uI-constants. Since in negatively substituted guanidines [ 381 the substituents are carried on an $'-hybridized nitrogen atom, as they are in the nitroso compounds, X-N=O, where substituent effects are known to be governed by the same factors as in carbonyl compounds (Bellamy and Williams, 1957), a correlation with urn-

-

15

PKll

10

-

5 -

0-

-5

I

L

0

+0 . 5

I

+LO

Figure 5. Substituent effects on the basicity of guanidine.

ALTERNATIVE PROTONATION SITES

307

constants appears theoretically better justified (see page 286). Such a correlation is shown in Figure 5. The pK,-data used are given in Table 2 and the om-constants were taken from the compilation by Pal'm (1961). There is some uncertainty in the om-valuesfor the Ph group, so two points are shown, and the pK,-value of guanidine, 13.60 (Angyal and Warburton, 1951), has been statistically corrected for six equivalent protons present in the cation. It can be seen that, TABLE 2 The pK,-values of Some Substituted Guanidines

Substi tuen t Phenyl Acetyl Carbamoyl Methoxy Cyano Nitro

10.77 8-26 8.00 7-51 -0.4, -0.3 -0.93

Temp. ("C)

Ref.

25 20 25 23 25 25

Q

b C

d e, f t?

'Davis and Eldcrfield (1932). Albert et al. ( 1 948). Hirt and Schmitt (1958). Borck and Clarke (1938). Hirt et aL (1961). f Takimoto (1 964). 8 Bonner and Lockhart (1958).

apart from a somewhat larger deviation of the Me0 substituent, the correlation is reasonably good ( p = -20.6). There is thus no evidence here of a change in the protonation site in the nitro-substituted derivative. Deductions from the very small differences in pK, observed for N,N-dialkyl derivatives, such as have been attempted by Price et al. (1970), are notoriously uncertain and cannot provide firm evidence for a particular site of protonation. The protonation site of biguanides [41] has also been somewhat controversial. The suggestion that the terminal (N') nitrogen is the

I

R*

I

R4

308

M. LILER

preferred protonation site (Shapiro et al., 1959) has recently been disproved by nmr studies of the hydrochloride of N' -n-butylbiguanide in DMSO (Wellman and Harris, 1967), which show that the a-CH, -group of the n-butyl substituent is coupled to one proton only. It is suggested that the first proton goes on to the imino-nitrogen, N2. The second protonation occurs in 3- 18 N sulphuric acid, as shown by the disappearance of the ultraviolet absorption of the monocation at 236 nm ( E = 14,800). Exchange of the NH-proton is too rapid in 18 N sulphuric acid for the observation of NH-proton resonances, but in concentrated sulphuric acid a new low field resonance (7 = 0.82) was ascribed to the second captured proton on N 3 . Thus the suggested structures of the monocations and dications are [42] and [43]. Some broadening of the NH,-signals by exchange

with concentrated sulphuric acid could be accounted for by a third protonation on N2 or N4 which destroys the symmetrical resonance system of the dication (Wellman and Harris, 1967). The first and second pK,-values of biguanide are 11.51 and 2-93 (Das Sanna, 1952) and the structure of the neutral base may be tautomeric, with one of the protons on N 3 . There remain in this Section to be considered p-aminobenzylideneaniline and p-aminoazobenzene and their derivatives. In these molecules two nitrogen basic centres are conjugated through a benzene ring, and the two systems are isoconjugate. The protonation of benzylideneaniline and its p- and p'-dimethylamino derivatives ([44] and [45]) (Schiff bases) was studied by

ALTERNATIVE PROTONATION SITES

309

ultraviolet spectroscopy by Reeves and Smith (1963). These compounds hydrolyse rapidly in highly acid solution and therefore a flow technique had to be used to obtain reliable spectra of the conjugate acids. The protonation of benzylideneaniline itself and several p-trimethylammonium derivatives leads t o bathochromic shifts of the ultraviolet absorption in the 305- 332 nm spectral region. This shift was taken to be characteristic of the azomethine nitrogen protonation, and, since a similar shift was found for the p-dimethylamino-derivative, it was concluded that the conjugate acid of [44] exists almost entirely in the anilium form. Low concentrations of the ammonium form could not be detected, however, since the absorption of this species is characterized by a shoulder, rather than a peak. p-Dimethylaminobenzylideneanilineand its conjugate acid have absorption coefficients several times larger than those of other derivatives, and the stability of the conjugate acid towards hydrolysis is greater. The spectral changes associated with the protonation of p'dimethylamino-derivatives [45] contrast sharply with those described above. There is a hypsochromic shift of the maximum absorption and the absorption spectrum of the conjugate acid of benzylidenep'-dimethylaminoaniline is closely similar t o that of benzylideneaniline free base, which leads t o the conclusion that the monocation has the ammonium structure. The conjugate acid of benzy1idenep'-dimethylaminoaniline shows, however, an additional peak at 455 nm, which was ascribed t o the anilium form, but the position of the tautomeric equilibrium could not be determined without the knowledge of the extinction coefficients of the two forms. pK,-values of only a few of these compounds were estimated from these measurements on flowing solutions. So, for example, p-trimethylammonium-benzylidene-p-toluidinehas a pK,-value of 1.85, and for variable p'-substituents a p-constant of 1.5 has been calculated. The p'-dimethylamino group may be expected to have pK, in excess of 5 . Before discussing the protonation sites of p-aminoazobenzene and its derivatives, it should be mentioned that there has been some controversy over the structure of the protonated azo-group itself. In their study of the protonation of the derivatives of azobenzene, Yeh and Jaffk (1959a) suggested that the protonated azo-group has a bridge structure (46) with the proton attached equally t o the two nitrogens. This view implies a cis configuration of the cation, whether it comes from cis- or trans-azobenzene. That cations from

M. LILER

3 10

cis- and trans-azobenzene give distinct ultraviolet spectra was demonstrated by Gerson et al. (1960), and this result means that the classical structure [47], with the proton attached to one of the two

[461

[471

nitrogens, is the correct one. When substituents are present in one or in both rings, depending on their nature, one of the two nitrogens becomes the preferred site of protonation. Correlations of pK, -values of disubstituted derivatives with o- and o+-constants (Yeh and Jaffi, 1959b) may be accounted for in these terms, as has been discussed by Liler (1971, p. 102). Some recent detailed studies of the effect of ortho- substituents on the pK, -values of azobenzenes confirm the view that an NH o-bond is formed in the cation (Hoefnagel et af., 1969; Haselbach, 1970). The protonation behaviour of a number of azobenzenes containing ionic solubilizing substituents has also been studied recently (Reeves, 1966; Reeves and Kaiser, 1969). The introduction of a p-amino-substituent enhances the basicity of azobenzene very considerably. The value of pK, changes from -2-90 (Yeh and JaffC, 1959a) to +F82 (Bascombe and Bell, 1959). The controversy over the site of protonation of p-aminoazobenzene has been resolved in favour of a tautomeric mixture of the ammonium [48] and azonium [49] forms on the basis of an analysis of

1481

[491

ultraviolet and visible spectra. This evidence has been reviewed by Lewis (1960). Absorption bands at approximately 320 nm were ascribed to the ammonium ion because they lie in the same spectral region as the K band of azobenzene itself. The azonium ions, which are resonance stabilized, as shown in [ 5 0 ] , absorb at longer wavelength (500-540 nm). Substituents cause an inverse variation of the @ l - N O h Mq

[501

@-~=N-@-NMe,

31 1

ALTERNATIVE PROTONATION SITES

two bands, since they stabilize preferentially one form or the other. The tautomeric equilibrium constant for p-dimethylaminoazobenzene, KT = [azonium] /[ammonium], has been estimated by Yeh and Jaffk ( 1 9 5 9 ~ and ) shown to increase from 1.4 in 6.6% sulphuric acid to 2.8 in 24.2% sulphuric acid, i.e., with increasing acid concentration the azonium form becomes progressively more favoured. This is a medium effect, such as has been discussed on page 294. The results of Tanizaki et al. (1966) on ultravioletlvisible spectra of p-aminoazobenzene in 30% aqueous ethanol- sulphuric acid mixtures show apparently a complete shift of the tautomeric equilibrium towards the azonium form with increasing acid concentration, which is then followed by diprotonation. TABLE 3 Medium Effect on the Tautomeric Equilibrium Constant of the Cations of p - Amino- and p-Dimethylamino-derivatives of Azobenzene (Korolev and Titova, 1 9 7 1 )

KT = [azonium] /[ammonium] Derivative P-NH2 p-NMez

Nitromethane

Acetonitrile

Water

23

4 9.5

0.25 2

8

The monoprotonation of p - [ (p-dimethylamino)phenyI]-azobenzenesulphonic acid (methyl orange) also gives a mixture of the azonium ion and the ammonium ion, in which the former predominates according t o Reeves (1966). Bolton et al. (1973). however, using linear free energy relationships, conclude that the ammonium form is dominant, and question the general acceptance of the tautomeric model for the protonation of p-aminoazobenzene and its derivatives. The tautomerism of p-amino- and p-dimethylamino-azobenzene cations has also been studied recently in nitromethane and acetonitrile (Korolev and Titova, 1971), and it has been shown that azonium forms are much more favoured in these solvents than in aqueous acid. The results are shown in Table 3. The estimates are based on the assumption that the extinction coefficient of the ammonium ion [48] may be taken t o be equal to that of p-trimethylammoniumazobenzene and that of the azonium ion to be equal t o that of the conjugate acid of 4-amino-3,5-di-t-butylazobenzene (Yeh and Jaffk, 1959c), which is assumed t o exist only in

312

M. LILER

the azonium form because of steric hindrance to amino-nitrogen pro tonation. Recent studies of the protonation of p-triphenylphosphazoazobenzene [ 51 J have shown that the absorptions of the two tautomeric

cations, the phosphazo- and the azo-protonated forms, occur in the same spectral regions as those of p-dimethylaminoazobenzenes, i.e., at about 350- 360 nm and at about 500- 520 nm, respectively (Zhmurova et al., 1971). However, protonation in dilute aqueous alcoholic acid occurs predominantly on the phosphazo-group. In more concentrated acid, 3 - 6 ~hydrochloric, weak absorption bands due to the azo-protonated form also appear, but hydrolysis becomes too rapid t o allow the determination of tautomeric ratios. It has been estimated that p-triphenylphosphazoazobenzene has a pK, -value in water of 8-36, i.e., that it is some 5.5 pK-units more basic than p-aminoazobenzene. Negative substituents in the para-position of the azobenzene part of the molecule increase the proportion of the azonium form of the cation, which leads to coloured salts (Zhmurova et al., 1972). In the phenylazoazulene [ 521 , the more basic of the two nitrogens is the nitrogen in the 0-position to the azulene ring (Gerson and Heilbronner, 1959), as predicted from molecular orbital calculations. In p-dimethylaminophenylazoazulene [ 531 , the same nitrogen re.

4 mains the protonation centre, as could be deduced from a consideration of substituent effects on the ultraviolet spectrum of phenylazoazulene (Gerson and Heilbronner, 1959). This may be

ALTERNATIVE PROTONATION SITES

313

explained by the fact that unsubstituted phenylazoazulene is several pK-units more basic than azobenzene, an effect which outweighs the tendency of the p-amino-group to enhance the basicity of the N(a) atom. Rather the basicity of the N(P) atom is enhanced. There was no evidence of amino-nitrogen protonation. By contrast, the mdimethylamino-isomer is protonated mostly on the amino-nitrogen, although a measurable amount of the azonium cation is also present (Gerson, 1963). Finally, it remains t o be mentioned at this point that the protonation of Schiff base derivatives of p-aminoazobenzene has also been studied (Ricketts and Cho, 1961). Benzy1idene-p’-phenylazoaniline [ 541 and its p-dimethylamino-derivativein mild acid solution show absorption bands at both 320 nm and 500 nm. These bands are

identical with those of p-dialkylaminoazobenzenes and suggest that the monocations exist as tautomeric mixtures of ions protonated on the azomethine nitrogen and on the azo-nitrogen. The tautomeric ratios have not been determined. These compounds are more stable to hydrolysis than the ordinary Schiff bases mentioned earlier. Other cases of amino-imino conjugation through more extended systems of conjugated bonds involving rings are also known and lead to resonance stabilized cations. Examples are Z-amino-l,4-naphthoquinone imines, which become protonated in strong acid at the imino-nitrogen, giving vinylogs of amidinium salts (Bullock et al., 1969). Much more complex systems are substituted 5-arylamino-9dimethylamino-9H-benzo [ a ] phenoxazines (derivatives of meldol blue), which protonate on the 5-arylamino-group (Stransky and Stuzka, 1968). This was concluded from the relatively high p-value (1.64 in 50% ethanol) in a correlation of pK,-values with o-constants of the substituents in the phenyl ring.

Ring systems with an exocyclic nitrogen atom This section will deal with ring systems containing only one of the two conjugated nitrogen atoms as part of a ring, the second usually being within an exocyclic amino-group. Systems containing both

314

M. LILER

conjugated nitrogen atoms as part of a ring or rings are discussed in the next section. The simplest systems, and most closely related t o the open chain systems discussed in the preceding section, are cyclic amidines, such as 2-iminopymolidine ([55], n = 3), 2-iminopiperidine ([55], n = 4) and their N-alkyl derivatives. Infrared spectra of their hydrochloride

H

I551

salts show an absorption at 1635-1700 cm-’, ascribed to the partially double C-N bonds in the cation [56] (Moriconi and Cevasco, 1968). This is entirely as expected from the behaviour of open chain aliphatic amidines. Conjugated systems analogous to the amidinium ion arise in the protonation of 2-aminopyridines, which are known t o be protonated on the ring nitrogen, giving the resonance stabilized ion [ 571 (Albert

et al., 1948). Both pyridine and an amino-group carried on an aromatic ring are weakly basic (pK, = 5-2 and ca. 4.6, respectively), whereas 2-aminopyridine has a pK,-value of 6.71, owing to this resonance stabilization. Analogous, but much more effective, resonance stabilization is found in 4-aminopyridinium cation [ 581 , +

H

ALTERNATIVE PROTONATION SITES

315

which has a pK,-value of 9-11, indicating an enhancement of pyridine basicity by some 4 pK-units. These effects are discussed more fully by Albert (1968, p. 74). All other y-amino-derivatives of polycyclic analogues of pyridine (quinoline, acridine) show a similar effect. 3-Aminopyridine is much less basic (pK, = 6-03) and it also protonates first at the ring nitrogen atom (Mason, 1960; Albert, 1960; Barlin, 1964). The second protonation of these compounds, studied more extensively recently for a number of amino- and aminonitro-pyridines (Bellobono and Favini, 1971), occurs only with difficulty when the amino-group is in the 2- or 4-position (pK,-values for the second protonation are -7.65 and -6.60 for 2- and 4-aminopyridine, respectively, but only -1-26 for 3-aminopyridine). The first and second protonation of N,N-dimethylaminopyridines have also been studied (Forsythe et al., 1972). The protonation sequence is the same as for aminopyridines, as expected. The basic strength is enhanced in the first protonation, as has also been found by Albert et al. (1948) and by Cruege et al. (1970), owing to the inductive stabilization of the monocation. The second protonation follows the H:’ acidity function and the second pK,-values are all more negative than for the primary amino-derivatives (-10.21, -2-1 and -9.28 for the 2-, 3- and 4-dimethylamino derivative, respectively). This is the usual behaviour of tertiary us. primary amines and is believed to be due to reduced hydrogen bonding stabilization of the tertiary as compared with the primary protonated amino-groups (TrotmanDickenson, 1949). Quantum mechanical calculations on the electronic structure of aminopyridines and their mono- and di-cations have recently been reported (Konishi et al., 1970). Similar studies have been extended t o diamino-derivatives of pyridine and t o amino-derivatives of more complex heterocycles (Bellobono and Favini, 197 1). In diaminopyridines, enhancements of basicity comparable t o those in aminopyridines are observed whenever the amino-groups are present in the 2- and/or 4- (and/or 6-) position, and preferred protonation sites are therefore not difficult to decide. The second protonation of these compounds is made easier by the extra amino-group present and occurs in largely aqueous acid. The third protonation, however, now requires H o values <--lo, depending on the relative positions of the two amino-groups, the last one protonated being always the group 01 or y to the protonated ring nitrogen. While monoamino-derivatives of acridine all protonate on the ring

M. LILER

316

nitrogen, as shown by studies of ultraviolet-visible spectra (Craig and Short, 1945; Tumbull, 1945), the 4,5-diamino-derivatives protonate on the primary amino-groups first, since the dication has an ultraviolet spectrum identical to that of acridine (Craig, 1946). A similar conclusion has recently been reached for the 4-hydroxy-5-amino derivative (Corsini and Billo, 1970). The two +M groups thus oppose each other's base-strengthening effect on the ring nitrogen atom. Aminopynmidines have been studied by Albert et al. (1948), and, by relying on the pK,-values at 20" of aminopyridines, there is no doubt that the 4-amino-derivative is protonated on the ring nitrogen N-1 (pK, = 5.69, as compared with pK, = 1.23 for unsubstituted pyrimidine). The 2-amino-derivative is a weaker base ( p K , = 3.45) and the 5-amino-derivative is weaker still (pK, = 2.60). An interesting ultraviolet spectral correlation has recently been pointed out (Albert and Taguchi, 1973) between compounds containing the 2-aminopyrimidinium and 2-pyrimidone structures ([59] and [60] ). \N

\N

\N \&NH2 HI

\NA4H2 HI

A0 - \;Ao-

N N

c----f

HI

\N

HI

1591

[GO1

This is due t o a formal similarity between the guanidinium and urea types of resonance. The second pK,-value of 2-aminopyrimidine is reported as -4.1 (Bean et al., 1967). The monocation [61] and the dication [62] of 2-aminopyrimidine have been fully characterized by nmr (Wagner and von Philipsborn, 1970). The monocation [ S l ] is present in trifluoroacetic acid, but the two NH2-protons remain equivalent even at low

I

H

H

(2 structures) [621

temperature owing t o rapid proton exchange. In fluorosulphuric acid, the symmetrical spectrum of the dication [62] is observed. Earlier pK,-measurements on substituted pyridazines have recently been extended by Cookson and Cheeseman (1972), who

317

ALTERNATIVE PROTONATION SITES

found that, for a series of substituted dimethylamino-pyridazines [63], substituent effects of the groups R are correlated with Me2N

$1 R ~ 3 1

c'. -

Me2N

R

Me2N'

QH R

~ 4 1

u,-constants by the equation pK, = 5.14 - 6.14 urn. This is consistent with the protonation of the ring nitrogen atom ortho to the dimethylamino-group, which is meta to the substituent. The enhancement of basicity of the parent compound by resonance in the cation [ 641 is substantial (pyridazine has pK, = 2-24.) The basic ionization constants of a large number of 2,4diaminopyrimidines and condensed pyrimidine derivatives have recently been measured and earlier values verified (Roth and Strelitz, 1969). The first protonation occurs at the N-1 ring atom (pK, = 7.40) and the effect of 5-substituents is correlated quite well with u,-constants for the +M substituents ( p = 4.85). The -M substituents (-NO2, -CN) show an enhanced base-weakening effect, which may be accounted for in terms of the direct conjugation shown in [65]. This provides additional evidence for protonation at N-1.

A further ring nitrogen atom in the 5-position in 2,4-diamino1,3,5-triazine leads t o a considerable reduction in basic strength (pK, = 3.91) and the site of protonation remains N-1 (or the equivalent N-5) (Roth and Strelitz, 1969). The effect of 6-substituents on pK, correlates with Hammett's urn-constants, and it is claimed that calculations of a-electron densities and nmr spectra indicate protonation at N-3, rather than at N-1 as in the unsubstituted compound (Marimoto, 1966). In 2,4,6-triaminopyrimidine there are two equivalent ring nitrogen atoms (N-1 and N-3), both in a-positions t o an amino-group and

3 18

M. LILER

hence preferred protonation sites (Roth and Strelitz, 1970). The ionization constant K 1 must therefore be statistically corrected by a factor of 2 for comparison with 2,4-diaminopyrimidines. The value of pK, (6.72) so obtained shows that the third amino-group has a net base-weakening effect, which may be accounted for by the resonance [66], which is known to be important in 2- and 4aminopyridine (see Albert, 1968, p. 75). The second protonation

occurs in moderately strong acid on the second ring nitrogen, and there are differences in the pK,-values obtained in sulphuric acid (1.7) and in hydrochloric acid (1-31),probably owing to solvation effects (Roth and Strelitz, 1970). On the basis of similarities in ultraviolet spectra of the cations of melamine [ 671 and ammeline [ 681 and the dianion of cyanuric acid [69], it was concluded that amidinium-type resonances exist in the

melamine cations (Hirt and Schmitt, 1958; Takimoto, 1964). The second and third protonations of melamine, the first and second protonation of ammeline, and the protonation of amelide [70] all lead to further amidinium or guanidinium type resonances. Prototropic equilibria of such molecules have recently been reviewed by Safta (1969).

ALTERNATIVE PROTONATION SITES

319

A bicyclic system, 5-azaindole7 may be mentioned here because it owes its relatively high basic strength (pK, = 8.3) t o a resonance stabilization of the cation [71] which is analogous to that of 4-aminopyridinium ion (Adler and Albert, 1960). Nmr spectra of the cations of 5-azaindole and its 1-phenyl derivatives have recently been reported (Dvoryantseva et al., 1973). Amino-substitution in more complex condensed heterocyclic rings, when in the 2- or 4-position relative to the ring nitrogen, leads t o protonation at that site. So, for example, 2- and 4-aminopteridins ([72] and [ 7 3 ] ) (pK,-values 4-26 and 3.54, respectively) are

[721

[731

protonated on N-1 (Albert e t al., 1951). Here there is n o exaltation of basic strength by resonance (pK, for pteridine is 4-12), which is so characteristic of the simpler ring systems discussed above, but which is known to decrease with the introduction of further nitrogen atoms into the rings (Albert et al., 1948). Recent quantum mechanical calculations confirm protonation on N-1 (Kapoor, 1973). Amidinium type resonance determines protonation sites also in some more complex systems, such as cytosine [74] and pterin [75]. Cytosine is protonated on the ring nitrogen N-1 (Katritzky and Waring, 1963) and calculations on the electronic structure of the cation have been published (Denis and Gilbert, 1968). Pterin is

Z-amin0-4-0~0-3,4-dihydropteridin. From the virtual identity of the ultraviolet spectra of the cations of this compound, [ 7 6 ] , and its 1-and 3-methyl derivatives, it was concluded that protonation occurs at N-1 (Pfleiderer e t al., 1960). These cations and many substituted derivatives have more recently been fully characterized by nmr

320

M. LILER

spectroscopy (Dieffenbacher and von Philipsborn, 1969). In trifluoroacetic acid monocations are present, but in fluorosulphuric acid a second proton is added on t o N-8. The chemical shifts of the NH2-protons of the monocation are linearly reiated to the pK,values for the first protonation.

Ring systems with two (or more) ring nitrogen atoms In this section systems containing two or more conjugated ring nitrogens will be discussed, including some carrying aminosubstituents. There is thus n o sharp division between the systems discussed in this section and the preceding one. A sharp division would have been difficult t o achieve without fragmentation of material. The simplest systems containing two conjugated ring nitrogens are 2-imidazolines [ 7 71 , which are cyclic amidines and give cations [ 781.

The basicity of imidazoline ( [ 7 7 ] , R' = R2 = H) has been estimated from a correlation of pK, -values of several 2-substituted imidazolines with u,,, -constants, which have been found t o correlate the basicities of open-chain amidines (see page 303). The correlation for imidazolines is pK, = 10.45 - 9.15 urn (Elguero et al., 1969b). The pK,-value of 10.45 shows that imidazolines are strong bases, although less strong than acetamidine (pK, = 12-41). The resonance stabilization is presumably less effective in the five-membered ring owing t o steric strain. The difference in the p-value compared with open-chain amidines ( p = 11.98) may have the same origin. The basicity of 2-phenylimidazoline is much enhanced by 1-methyl substitution (pK, changes from 9.90 t o 10.90). While the +I effect of the methyl group certainly plays a part, it appears that the unsubstituted base may be stabilized by hydrogen bonding of the NH-proton to water molecules, which is enhanced by the partial positive charge on N-1. Such hydrogen bonding is not possible in the 1-methyl derivative, hence the unusual increase in basicity, which is greater than for similar N-methylation of primary amines or anilines (Elguero et al., 1969b).

32 1

ALTERNATIVE PROTONATION SITES

1,2-Diaryl-2-imidazolines are also protonated on N-3, since otherwise their pK,-values could have been expected t o be close t o those of the corresponding arylamines (Harnsberger and Riebsomer, 1964). The pK,-value of the parent compound is high (9.26) and the substituent effects are correlated with a--constants of the substituents for aryl groups in position 1 (Fernindez et al., 1973). This is due t o an aniline type resonance in the base [79] involving N-1, which affects electron densities at N-3 indirectly.

(2 structures) [791

A higher analogue of 2-imidazolines, which has recently been examined by nmr (Lloyd et al., 1973) is 2,3-dihydro-lH-l,4diazepine [80], which protonates on the imino-nitrogen to give cation [81]. Nmr spectra show that this has a half-chair shape, which inverts rapidly at room temperature.

(3 H

In imidazole [82], one of the nitrogen lone pairs (N-1) forms part of the aromatic sextet, and so protonation on N-3 is expected. The fact that imidazole is a stronger base than pyridine (pK, = 6.99 us. 5-23) and not weaker, as might have been expected owing t o the

electron withdrawing effect of N-1, has long been ascribed to resonance stabilization of the symmetrical cation [ 831 . This may also be regarded as an amidinium type resonance. The fact that benzimidazole (pK, = 5.53) is much less basic than diphenylacetamidine (pK, = 8.30) is a consequence of steric strain in the

322

M. LILER

five-membered ring as compared with the open-chain system (Schwarzenbach and Lutz, 1940). Nmr spectra of imidazolium ions have been studied more recently (Mannschreck et al., 1963). In concentrated sulphuric acid the exchange frequency of NH-protons is sufficiently reduced for the observation of NH-CH couplings. The spectra prove that imidazole and benzimidazole are definitely protonated on N-3. The tautomerism of 4- (or 5-) substituted imidazoles has also been studied by nmr (Staab and Mannschreck, 1963). The protonation of 2-d-4(or 5)-bromoimidazole [ 841 in concentrated D2S04and of 1,2-d2-4(or5)-bromoimidazole [ 851 in

y.g”’

Conc. D2S04

D

H

H

Conc. HzS04

D ~

’ D

5

1

WI

concentrated H, SO4 yield different cations ( [ 861 and [ 871 ) owing to slow proton exchange in these media (the half-life of NH protons is about 30 minutes at 35°C and 0.3 M ) . This proves that the dominant tautomer in the base is the one with the proton further removed from the electron-withdrawing substituent, i.e., the neutral base is 4-bromoimidazole. Protonation shifts, obtained from nmr spectra in deuterochloroform and trifluoroacetic acid, for 1-methylimidazoles also suggest that the cations show amidinium type resonance, because the signal of the proton on C-2 shifts downfield much more than those of the protons on C-4 and C-5 (Barlin and Batterham, 1967). Amidinium type resonance in benzimidazoles activates the proton on C-2, which becomes susceptible to base-catalysed exchange (Elvidge et al., 1973). The effects of protonation on electron densities in the imidazole ring have been calculated by extended Huckel theory (Adam et al., 1967). The protonation of 1-methyl-l,2,4-triazole [ 881 leads t o a cation stabilized by amidinium type resonance [ 891 only if the proton adds on t o N-4. Chemical shifts caused by protonation confirm this,

ALTERNATIVE PROTONATION SITES

323

because the downfield shift of the proton on C-5 is much greater than that on C-3 (Barlin and Batterham, 1967). The crystal structure of 3,4,5-triamino-172,4-triazole hydrobromide has recently been reported (Seccombe and Kennard, 1973) and shows that protonation occurs on N-2. The structure of the cation is best described by the resonance [go] , although there is some spreading of charge on to N-4 as well.

+

~901

The rather high basic strength of 3,5-diazaindole [91] (pK, = 6.1) is not due to the protonation of the imidazole ring (for benzimidazole pK, = 5-53), but rather of the pyridine ring. Cation [92] is formed and this is stabilized by the same kind of resonance as

1911

5-azaindole [ 711 , mentioned in the preceding section (Albert, 1968, p. 200). Purine [93] is analogously protonated on N-1 (Coburn et al., 1965). With a pK,-value of 2.30, it has an enhanced basicity compared with that of pyrimidine (pK, = 1.23). In the nmr spectrum in trifluoroacetic acid, the resonance of the captured proton is not observable owing t o exchange, but it can be observed in NH2 I

H

H

324

M. LILER

pure fluorosulphuric acid at -55"C, together with the resonances of two more NH-protons (at N-7 and N-9), all at very low field (Wagner and von Philipsborn, 1971). An additional proton at N-3 is observed in 1: 1 HS03 F-SbF, diluted with SO2. Thus the sequence of proton additions on purine is N-1, N-7 and N-3. The first protonation on N- 1 has also been confirmed for 8-methylthiopurine (Reichman et al., 1973). Adenine (6-aminopurine, [ 941 ) has been shown to be protonated on N-1 in the solid hydrochloride by x-ray measurements of electron densities in the crystals (Cochran, 1951). Similarly, adenosine 3'-phosphate dihydrate has a zwitterion structure, in which one of the phosphate protons is on the N-1 nitrogen of the adenine part of the molecule (Sundaralingam, 1966). The pK,-value of adenine (4.12) shows that the basicity of the purine ring is enhanced by the 6-amino-substituent owing to the amidinium type resonance, which arises on N-1 protonation. In trifluoroacetic acid- sulphur dioxide mixtures at -55°C restricted rotation of the amino-group of the amidinium system is observable, which gives rise to two nonequivalent protons in the nmr spectrum (Wagner and von Philipsborn, 1971). The chemical shifts show that the 1,7diprotonated cation is present in that medium. In fluorosulphuric acid at -55°C there are altogether four individual NH-resonances (in addition to NH2), corresponding t o protons in positions 1, 3, 7, and 9, i.e., the trication is present. The presence of the cation protonated on N-1 cannot account for the fluorescence of aqueous acidic adenine solutions (pH = 2), since the 1-methyl derivative does not fluoresce under the same conditions (Borresen, 1967). It has therefore been suggested that other tautomeric forms of the cation are also present, the fluorescent tautomer probably being protonated on the amino-group with another proton on N-7. Quantum mechanical calculations (Veillard and Pullman, 1963) indicate similar proton affinity for N-1 and N-3, and a lesser one for N-7. There are numerous calculations in the literature on the electronic structure of adenine (see Boyd, 1972, and references quoted therein) and a recent one on N-7-H and N-9-H tautomers protonated on N-1 (Jordan and Sostman, 1972). The N-9-H form is preferred according to both MIND0 and CND0/2 calculations. Guanine (2-aminopurine-6-one, [ 951 ), in which N-1 already carries a proton, is protonated at the N-3 position (pK, = 2-95), since the ultraviolet spectrum of the cation [96] is closely similar to that of

ALTERNATIVE PROTONATION SITES

325

purine-2,6-dione [ 971 . The 9-methyl derivative, however, is protonated on N-7 (imidazole ring), since its ultraviolet spectrum is

similar to that of 2-amino-7,9-dimethyl-6-oxopurinium cation (Pfleiderer, 1961). Nmr spectra in anhydrous trifluoroacetic and fluorosulphuric acids, studied by Wagner and von Philipsborn (1971), show the protonation sequence for guanine to be N-3, N-7, 0. The last protonation is only achieved in FSO3H-SbF5-S02. Barriers to rotation about the exocyclic C-N bonds, due to resonance interactions with the purine rings, have recently been calculated by the CNDO/2 method for all the nucleotide bases (cytosine, adenine, guanine) (Rao and Rao, 1973). Some other ten n-electron N-heterocycles, but with a nitrogen atom shared by two rings, also present two or more conjugated nitrogen protonation sites. One such system is imidazo[ 1,2-a] pyridine (l-azaindolizine, [98]), which may be viewed as an

[981

[991

analogue of 2-aminopyridine. In fact, all the protons on the sixmembered ring are deshielded at least to the same extent as the protons in 2-aminopyridine and the molecule is planar (Paudler and Blewitt, 1966). An andysis of the chemical shifts in the methiodide and the hydrochloride leads to the conclusion that both protonation and quaternization occur on N-1. In azaindolizines protonation generally takes place on the non-bridgehead nitrogen atom, irrespective of its position in the molecule (Armarego, 1965). Protonation on N-4 would destroy the aromaticity of the ring and indolizine itself is protonated on C-3 (see page 360). The introduction of further ring nitrogen atoms into [ 98 J usually does not change the site of protonation, which remains N-1 (Armarego, 1965). The condensed ring system [99], however, adds

326

M. LILER

the proton on t o N-2, and is a much weaker base (pK, = 0.42). Some further such systems have been studied by Paudler and Helmick (1968), who have shown that they all protonate either on N-1 or on N-3. Pyrazole( 2-azapyrrole, [ 1001 ) is also an aromatic ring, in which the N-1 atom donates its electron pair to the aromatic sextet.

Alternatively, conjugation may be thought t o occur as in [ l o l l , which would tend t o enhance the basicity of N-2. There is little evidence of this, however, since pyrazole has a pK,-value (2.52) very similar to that of pyridazine (2.3). Both are much less basic than pyridine (pK, = 5-23), which may be attributed to the effect of the neighbouring nitrogen. Pyrrole is a very weak base (and in fact C-protonated, see page 357), so that it is not surprising that pyrazole is protonated on N-2 t o give the resonance stabilized cation [ 1021. For pyrazoles substituted in positions 3 and 5, Staab and Mannschreck (1962) have shown by nmr spectra in concentrated sulphuric acid that each substituent is coupled t o one NH-proton, and Elguero et af. (1967b) have shown that the cation has a plane of symmetry. Elguero et af. (1968) have estimated that the cation [lo21 is dominant by a factor of at least lo3 over the alternative N-1 protonated cation, whose formation would require a considerable localization energy of the nitrogen electron pair involved in the aromatic resonance. Since cation [lo21 is the dominant form, substituent effects on pK,-values can be discussed entirely in terms of that structure (Elguero et al., 1968). A number of variously substituted pyrazolium ions have been fully characterized by nmr . orbital calculations (Adam et af., (Elguero et al., 1 9 6 9 ~ )Molecular 1967) show very good correlations between total electron densities and the chemical shifts (protons and I3C) for azoles and their conjugate acids, thus indirectly proving structure [ 1021 . In 2-pyrazolines [ 1031 the conjugation involves N-1 and C-3 atoms. 2-Pyrazolines may be regarded as cyclic hydrazones, which have the advantage over non-cyclic products of being stable to hydrolysis. It has been shown by Elguero and Jacquier (1965) that protonation occurs at N-1, giving cation [ 1041, while forms with the proton at N-2 and C-3 [ 1051 may exist in amounts of less than 1-5%.

ALTERNATIVE PROTONATION SITES

H'

327

'R'

An sp3 nitrogen is usually regarded as more basic than s p 2 , but this does not mean that one can predict which nitrogen atom is the more basic in 2-pyrazolines because the state of hybridization of N-1 is not pure sp3 owing to its conjugation with the C-N double bond. 2-Pyrazoline (pK, = 4-62) is some 3 pK-units less basic than pyrazolidine (estimated pK, = 7.8, the same as for diethyl-1,Zhydrazine). A reduction of about 2 pK-units may be ascribed to conjugation with the C-N double bond, and the rest to the effect of the sp2 nitrogen (Elguero et al., 1969a). The effects of substituents in position 3 on pK,-values are best correlated with a,-constants ( p = 6.14). This may be understood in terms of an analogy between the p-n conjugation in pyrazolines and that in anilines. The C=N bond (like the C=C bond) transmits substituent effects better than the aromatic ring ( p = 2.8 for aniline). The substituents in position 5 cause a reduction of basicity by steric hindrance to solvation of the cation (Elguero et a/., 1969a). There is thus little doubt that N-1 protonated cations are dominant in solution. Indirect evidence for the presence of other forms (N-2 and C-3 protonated) is obtainable, however, from a consideration of the mechanisms of some acid catalysed reactions of 2-pyrazolines (prototropy, cis-trans isomerisations, hydrogen exchange) (Elguero et a[., 1970). The ambiguity of infrared criteria for the identification of protonation sites in imidazoles and 2-pyrazolines has been discussed by Elguero et af. (1967a), who have also provided new information on the infrared spectra of the hydrochloride of 3,5,5-trimethyl-2pyrazoline in chloroform solution. There is a shift of the stretching vibration of the C=N bond from 1624 cm-' in the base to 1649 cm-' in the hydrochloride salt in chloroform, i.e., in the N-1 protonated cation. This is analogous to similar shifts of the carbonyl frequency in some amide salts (see page 338). Closely similar to 2-pyrazolines are 1,4,5,6-tetrahydropyridazines [106]). They have been shown to undergo both protonation and quaternization on N-l (Aubagnac et al., 1970). Small amounts of the N-2 protonated form may be present, but nmr spectra show no evidence of protonation on C-3.

M. LILER

328

I

3-Pyrazolines [ 1071 also offer three possible protonation sites, the two nitrogens and C-4. They are in fact cyclic enehydrazines and their protonation will be discussed on page 354.

N / O ( o r N / S ) Alternatives Open-chain systems The simplest conjugated system involving nitrogen and oxygen atoms in an open chain is the amide group [108]. The site of R-CHO ‘NH~

-

R-C

yo%H*

protonation of amides has been the subject of a prolonged controversy, and the problem has been examined from more angles than most. This has only recently resulted in an answer, namely that both N- and 0-protonation are possible depending o n conditions. There is a vast difference in the basicity between ammonia (pK, = 9.25) and a typical carbonyl compound, such as acetone (pK, = -7.7) which is half-protonated only in 80% sulphuric acid (see Liler 1971c, p. 124). The resonance interaction [ 1081 decreases the basicity of the nitrogen and enhances the basicity of the oxygen atom. The resonance energy of the amide group has been estimated by Pauling (1939, pp. 207-8) t o be 21 kcal mole-’ (88 kJ mol-’ ). On this basis it was claimed that the basicity of the nitrogen atom was negligible, and that amides therefore could not form salts by adding a proton on to the amino-group. This resonance energy is a gross overestimate, since an analysis of the pK,-values of amides which are now known leads t o a much smaller figure. It has been pointed out on page 286 that amides obey the correlation of carwith om -constants of bony1 frequencies of compounds R-CO-X

329

ALTERNATIVE PROTONATION SITES

groups X (Figure 3), and that therefore the relative importance of the inductive and mesomeric effects of the C=O group on the electron density of the NH2-group should be the same as that of the benzene ring, since phenyl and acyl groups are electron sinks of an analogous character. In aniline (see page 280) the inductive effect of the ring is responsible for about one half of the reduction in the basicity of the amino-nitrogen atom, the mesomeric delocalization of electrons being responsible for the other half. Similar relationships may be expected in the amide group and are in fact found (Liler, 1971c, p. 107):

LH3

pK, W

a

[1091 10.58

[ 1101 5.33 5.25

[1111

-0.36 5.69

2,2-Dimethylquinuclidone-6 [ 1101 was synthesized by Pracejus (1959) as an amide involving n o mesomerism and its pK,-value was determined. Pracejus e t al. (1965) later assumed that the -M effect of the CO-group reduces the basicity of the NH2-group by at least a further 4 pK-units (since their estimates of the pK,-values of several amides lay in the range 0 t o + l ) , and, using the approximate relationship discussed on page 289, estimated a lower limit for the resonance energy of the amide group as 5.4 kcalmole-' (23 kJ mol-') (the implicit assumption being that the protonation of amides occurs on the nitrogen). A reliable value for the pK,-value o f dimethylacetamide in aqueous sulphuric acid, obtained more recently from chemical shift measurements (Haake et al., 1967; Liler, 1969), is shown in the above scheme. It can be seen that within some 0-4 pK-units, the reduction in basicity due t o mesomerism is equal t o that due to the inductive effect. The analogy with the aniline scheme (see page 280) is very close, and bearing in mind the analogous character of the phenyl and the carbonyl groups as electron sinks (as discussed on page 285), it indicates that dimethylacetamide is N-protonated under the conditions of pK, -determination (halfprotonation occurs in ca. 9% sulphuric acid). Using again the approach of Section 3, the reduction in basicity of 5-69 pK-units is found to correspond t o a resonance energy of 7.8 kcalmole-' (33 kJ mol-I). This is lower than the activation energy for rotation around the C-N bond in dimethylacetamide, which was found to be

M. LILER

330

Ea = 11.6 k 0.8 kcal mole-’ (48.0 3-3 kJ mol-’) or AGg = 20.1 kcal mole- (88 kJ mol-’ ) (Rogers and Woodbrey, 1962), because in rotation around this bond, in addition t o the locdizatisn energy of the nitrogen lone pair, energy to overcome non-bonded interactions of groups carried on the carbon and nitrogen atoms must also be supplied. The resonance energy of 7.8 kcal mole-’ (33 kJ mol-I) for the amide group is about twice as large as the extra resonance energy in aniline (see page 290).

0

5

10

15

p KO&! OH) Figure 6. A linear free energy relationship, showing the effect of substitution upon the acidity of -OH and -NHI groups.

There is also a relationship between pK,-values of ethanol, phenol and acetic acid on the one hand, and ethylamine, aniline and acetamide on the other (Liler, 1971c, p. log), which may be presented as a linear free energy relationship (Figure 6 ) by regarding Et, Ph and CH,CO as substituents (R). This excellent relationship leaves no doubt that the acidic group in the R-NH; series remains unchanged, i.e., that acetylammonium ion is formed in the protonation of acetamide in aqueous acid (half-protonation in ca. 19% sulphuric acid). Analogies between groups -0- fi =O and -NH, * = N H i in spectroscopic properties of a number of mesomeric systems have been pointed out on page 316. They extend also to aniline and phenolate anion (‘Jones, 1945) and t o benzamide and benzoate anion (1, a x 225 nm and 224 nm, respectively, with similar cm a values). The relationship in Figure 6 shows that there is also a parallel

33 1

ALTERNATIVE PROTONATION SITES

between -0- and -NH, in the free energies of protonation of these simple mesomeric systems. The electronic similarity of the two series in Figure 6 is also reflected in the slope of the plot, which is very close t o unity (1-05). Further evidence for nitrogen as the site of protonation of amides in largely aqueous acid comes from studies of substituent effects on the pK,-value of benzamide (Edward et al., 1960; Yates and Stevens, 1965). It has been observed that the pK,-values are correlated with u-constants of the substituents, rather than with u+. This means that the structure of the dominant form of the cation is N-protonated as in [ 1121 , because resonance interactions of para-substituents in this kind of cation are similar to those in benzoic acid. The p-value (0.92)

[1121

~

3

1

is also closely similar to p = 1 for benzoic acid. In the O-protonated cation [ 1131, the resonance interactions of the para-substituents should be similar to those in other carbonyl protonated compounds, which all show correlations with u+-constants of the substituents (see e.g., Liler, 1971c, pp. 134-136). The expected p-value for carbonyl protonation would be about 3.4, as may be estimated from a correlation of p-values with carbonyl frequencies of several carbonyl protonated series (Liler, 1965). This is much larger than the observed value. These correlations thus show unambiguously that the protonation of amides in aqueous acid occurs on the nitrogen atom. This conclusion is in apparent conflict with nmr evidence showing that only O-protonated cations are observable in concentrated sulphuric acid, 72% perchloric acid (Fraenkel and Niemann, 1958; Fraenkel and Franconi, 1960) and pure fluorosulphuric acid (Gillespie and Birchall, 1963). Doubts about the interpretation of these spectra (Spinner, 1960a) have been dispelled by Fraenkel et al. (1961). The resonance of the proton on the oxygen was also recorded (Gillespie and Birchall, 1963) and consequently O-protonation of amides became generally accepted as a universal phenomenon. This was despite the fact that nmr spectra of amides in aqueous acid show all the features associated with N-protonation, namely rapid exchange of NH-protons with the solvent, with the consequent loss of coupling t o the N-methyl group in N-methyl amides [ 1141, and free rotation

M. LILER

332 R ,

H 8 ;

CZN, Me

8-

t 1141

[1151

around the C-N bond, with the merging of the resonances of cis and trans N-methyl groups in N-methyl and N,N-dimethyl amides [ 1 151 (Berger et al., 1959; Liler, 1969; Liler, 1971b). The assumption of N-protonation is essential in order to account for these phenomena, but it was thought that the concentration of the N-protonated form was negligible, since no N-protonated cation was detectable at high acid concentrations when proton exchange was slowed down (Berger et al., 1959). The rates of protonation, obtainable from the collapse of NH-NMe couplings and the onset of free rotation (and therefore referring to N-protonation) were determined by several authors (Berger et al., 1959; Takeda and Stejskal, 1960; Bunton et al., 1959). The rate constants obtained show the process to be a slow one as compared with the protonation of typical amines (including aniline) which is diffusion controlled (Eigen, 1964). For example, the rate = constant for the protonation of N-methylacetamide is k 400 M - ~s-’. The slowness of the process may be accounted for in terms of the electronic reorganization that is necessary in the N-protonation of the conjugated system [Equation (30)]. Bunton et

al. (1959) pointed out that the knowledge of the reverse rate constant ( k - l ) was necessary in order to calculate the concentration of the N-protonated form, because the two rate constants enable the calculation of the equilibrium constant for the protonation where K , is the equilibrium (30) from the relationship K , = k acid ionization constant of the cation. This rate constant is not obtainable from nmr spectra. It cannot be assumed to be equal to the diffusion rate constant (as has been done by Martin, 1972), because processes involving solvent molecules cannot be assumed to occur with diffusion-limited rate constants. However, the experimental K,-values of amides have been shown by the above pK,-relationships to refer to N-protonation, and therefore k - can be estimated from the known K,- and k - values. Taking pK, for N-methylacetamide as -0-72 ( K , = 5-25) (Liler, 1969), the value of k - is 2100 s - l . The reverse reaction is thus also slow and activation controlled, as would

,

ALTERNATIVE PROTONATION SITES

333

be expected for proton transfers between two sites of comparable basicity (pK, for the hydronium ion may be taken t o be -1.74). In a study of acid catalysed proton exchange in aqueous urea solutions, much faster rates have been found (k = 9 x 1O6 M s-l) (Vold et al., 1970). This was ascribed to N-protonation. A higher rate is expected for a stronger base [pK, for urea is 0.8 or 1.6, according to Arnett (1963) and 0.18, according to Parry et al. ( 1 9 6 9 ) l . The exchange of NH-protons of thiourea is acid-catalysed about 6 0 times as efficiently as that of N-methylacetamide (Vold and Correa, 1970). The oxygen basicity of amides has so far been estimated reliably only for benzamide (Liler, 1971a) from the correlation of pK,-values of a number of carbonyl protonated compounds with the stretching frequency of the carbonyl bond (Figure 2). Benzamide deviates from that relationship, not unexpectedly in view of the free energy relationships discussed above which show that in aqueous acid the proton adds on to the nitrogen. The estimated pK,-value for carbonyl protonation is -5.8. Since the measured pK,-values for N-protonation range between -2-16 (Edward et al., 1960) and -1.74 (Yates and Stevens, 1965), and bearing in mind medium effects on the spectra which are invariably reported in the literature on ultraviolet studies of the protonation of amides, an estimate of -2.0 for the pK, -value of the benzamide nitrogen seems reasonable. This leads to an estimate of the tautomeric equilibrium constant

,

where KT is given by (32), y N and yo being the activity coefficients

K

[PhCONH:]

YN

- [PhC(OH)NH:]yo

of the two forms. Thus in highly dilute aqueous acid, the Nprotonated form of benzamide is favoured by a factor of almost lo4 over the O-protonated form. A similar, but less reliable, estimate for acetamide is ca. l o 6 (Liler, 1971a). It now remains to explain why this tautomeric equilibrium shifts towards dominant O-protonation in concentrated and anhydrous acid. The reasons for the changeover can best be understood in terms of the developments in the theory of acidity functions, which have been outlined on page 293. Cation solvation must be taken into account in considering the tautomeric equilibrium (33). Ph-CO-NH;

. p H 2 0 + Ph--C(OH)NG. q H 2 0 + rH20

(33)

334

M. LILER

Strong solvation of the cations formed in dilute acid is apparent both from the fact that amides generate an acidity function of their own ( H A ) , which is less negative than the Hammett acidity function Ho (Yates et al., 1964), and from the measurements of relative activity coefficients (yBH+/yT A + ) for benzamide cations (Sweeting and Yates, 1966) (see Figure 4). These relative activity coefficients show sharply rising values, more so than for anilinium ions, which means that benzamide cations become strongly destabilized at higher acid concentrations (with decreasing water activity), and this in turn means that they are strongly stabilized by hydration in aqueous acid. Strong hydration is expected for N-protonated cations of benzamide because there are three hydrogen bonding sites t o the three NHprotons and one to the carbonyl oxygen. Strong hydrogen bonding is favoured by the localized positive charge on the nitrogen, and there may be as much as 46 kcal mole-' (190 kJ mol-') of solvation energy for the NH:-group alone, with further contributions from carbonyl hydration (see Sheraga, 1968, p. 140). The 0-protonated form retains the relatively small resonance energy, but has only three hydrogen bonding sites and the charge is delocalized, so that weaker hydrogen bonds would be formed. Therefore, in terms of primary hydration, an estimate of r = 1 in equation (33) is reasonable (Liler, 1971a), and the tautomeric equilibrium constant in concentrated acids should include the activity of water [Equation (34)]. The KT =

[ PhCONHf as.] Y N [PhC(OH)NHl aq.] yoa,

(34)

activities of water in sulphuric acid-water mixtures show a sharp decrease from 65% acid (a, 0.1) t o 100% acid (a, 2 (Giauque et al., 1960). This brings about a shift in the tautomeric equilibrium (33), so that in anhydrous acid the 0-protonated form becomes dominant. 0-Protonated cations of amides in concentrated and anhydrous acids are now well characterized by nmr spectroscopy. 0-Protonated cations of N,N-dimethyl amides are most easily observed, even in 72% perchloric acid which has a water activity of about because for tertiary amides the N-protonated forms is relatively less stabilized by hydration (Liler, 1972a). 0-Protonated cations of N-alkyl amides show considerable exchange of NH-protons with the solvent in 72% perchloric acid owing to the intervention of the N-protonated form. For primary amides (acetamide), however, 0protonated cations are not observable in that solvent (Liler, 1972b),

ALTERNATIVE PROTONATION SITES

335

because relatively high concentrations of the N-protonated cations are present, which cause rapid NH-exchange. In 100% sulphuric acid, pure fluorosulphuric acid and HSO, F-SbF, -SO2 solutions, 0protonated cations of all amides are stable on the nmr time scale, but the OH-proton resonance is observable only in the latter two media at low temperature. For primary amides, the resonances of the two non-equivalent NH-protons are more readily observed in Nenriched samples, because NH-resonances in ordinary amides are broadened by quadrupole relaxation of the N-nucleus. NH-protons then show coupling to the , N nucleus ( l = f )with large coupling constants ( J > 90 Hz), but the geminal coupling is not resolvable (Liler, 1972b). The use of N-benzamide has a particular advantage, because, owing to the large splitting by the N nucleus, it obviates to a large extent the overlap of the NH resonances with the aromatic region (Liler, 1974). Cis- and trans-isomers of the 0-protonated cations ([ 1161 and [ 1171 ) of a number of N-alkylamides have also

’

’

trans

’

’

R\ /R’ /C-=N HO’ + ‘H cis

been observed (Liler, 1972a). There is a general preference for the trans-isomer, but in N-alkylformamides (R = H), the proportion of the cis-form increases with increasing bulk of the group R‘. The configuration of the OH-proton has also been determined for several primary amides at low temperature (-8OoC, in 1:l HS03F-SbF, diluted with SO2) and found to be trans to the NH2-group (Brouwer and van Doom, 1971). Urea and substituted ureas have also been shown to be 0protonated in anhydrous media, with a second protonation ensuing on one of the N atoms in pure fluorosulphuric acid and in “magic acid” (Birchall and Gillespie, 1963; Olah and White, 1968). Both 0- and N-protonated cations of N-acylaziridines have been observed by nmr spectroscopy, the first in HSO,F-SbF5-SO2 at -60°C, and the second as the hexafluoroantimonate in liquid SO2 at -6OOC (Olah and Szilagyi, 1969). Presumably pure liquid sulphur dioxide is a better stabilizing medium for N-protonated cations than the “magic acid” mixture. At room temperature the OH-proton of the 0-protonated cations exchanges rapidly with the solvent in all strongly acidic solvents. The

336

M. LILER

NH-protons are brought into rapid exchange with the solvent medium by adding water t o 100%sulphuric acid, for example, owing t o partial conversion of the O-protonated cation into the Nprotonated cation (Liler, 1972b). Under such conditions of rapid proton exchange, it is impossible to estimate the position of the tautomeric equilibrium from the appearance of the amide group resonance. There may be differences in the chemical shifts of the groups attached t o the amide group in the N- and O-protonated cations, but these appear to be small. Information on the position of the tautomeric equilibrium is obtainable from ultraviolet spectra. The literature had in fact contained evidence of tautomeric change for amide cations in concentrated acids for some time. Hantzsch (1931) demonstrated that the ultraviolet spectrum of benzamide in 100% sulphuric acid was virtually identical with that of the cation of ethylbenzimidate in the same medium, as expected for the O-protonated cation. With the nmr evidence from the spectra of "N-benzamide (Liler, 1974), the presence of the O-protonated cation of benzamide in 100% sulphuric acid is thus well established. On the other hand, Edward et al., (1960) pointed out that the spectrum of benzamide in 60% sulphuric acid, in which the amide was taken to be fully protonated, was closely similar t o that of unprotonated acetophenone in 50% sulphuric acid. This was as expected for the N-protonated cation, since replacement of a methyl group by NHS has little effect on the ultraviolet spectra of many aromatic compounds (Jones, 1943; Doub and Vandenbelt, 1947). An unexplained medium effect on the spectrum of benzamide in >60% sulphuric acid had also been reported (Edward and Meacock, 1957). A closer examination of these spectra showed that there was a shift of the maximum absorption from 245-5 nm ( c m a x= 12.6 x l o 3 ) in 60% sulphuric acid t o 252-5 nm (emax= 13.8 x l o 3 ) in 100% sulphuric acid and that an isosbestic point is present at 247.5 nm, thus showing that the spectral changes are due to a shifting equilibrium (Liler, 1 9 7 2 ~ )It. is possible to estimate the tautomeric quotient from the ultraviolet spectra and, using known values of the activity of water and KT = lo3" [Equation (31)], t o calculate the ratio of activity coefficients of the two forms ( y N / y O ) from Equation (32). The result obtained shows that y o increases rather less steeply with increasing acid concentration than y N , as expected for different degrees of hydration stabilization of the two kinds of cation (Liler, 1974).

ALTERNATIVE PROTONATION SITES

337

Large medium effects on the ultraviolet spectra of the protonated form of a number of other amides have been reaffirmed in some recent protonation studies (Farlow and Moodie, 1970; Congdon and Edward, 1972; Barnett and O’Connor, 1973). These undoubtedly conceal evidence of tautomeric change. Protonation studies of aliphatic amides by ultraviolet spectroscopy are more difficult because the main absorption bands occur in the far ultraviolet region. It is interesting that in the first study of this kind, in the spectral range 200-210 nm, Goldfarb et al. (1955) were unable to account for the spectral changes in terms of a single protonation equilibrium, but had to postulate two equilibria owing to spectral changes at more negative values of H , , the second of which they were unable to assign with certainty. A new study has recently been reported on dimethylacetamide, purporting to prove dominant O-protonation in dilute acid (Benderly and Rosenheck, 1972), although the spectral shift of the maximum absorption at 196 nm with increasing acid concentration is towards shorter wavelengths as expected for N-protonation. Maximum absorption at 191 nm in 30% sulphuric acid corresponds t o 95% protonated amide (pK, = 0.36) (Liler, 1969), and further changes are observed at still higher acid concentrations, thus suggesting tautomerization. Thioamides have likewise been shown by nmr studies to protonate on the sulphur in pure fluorosulphuric acid (Birchall and Gillespie, 1963). Thioacetamide has been studied by ultraviolet spectroscopy in aqueous hydrochloric acid by Hosoya et al. (1960), who interpreted the hypsochromic shift of the absorption maximum (from 260 nm in water to 235 nm in 6M HCI and to 237 nm in 1 2 HCI) ~ as being due to N-protonation, but observed that the crossing of absorption curves was imperfect for solutions of higher acid concentrations. The absorption maximum of thioacetamide in 96% sulphuric acid, in which it is undoubtedly largely S-protonated, has also been found at 234 nm (Janssen, 1961), but with log E , a x = 4.16, which is much higher than that found in 1 2 HCI. ~ Thus there is evidence here of spectral change in concentrated acids other than that due to initial protonation. The basicity of thioacetamide is similar to that of acetamide, and hence suggests N-protonation in dilute acid. The interpretation of the infrared and Raman spectra of solid salts of amides has been the subject of much disagreement. One must first distinguish solid complexes of amides with carboxylic acids from salts with strong mineral acids. In the first, the interaction is

338

M. L E E R

undoubtedly of the hydrogen bonded type, as some recent x-ray studies of 2 : l complexes of benzamide with succinic acid and of furamide with oxalic acid show (Huang et al., 1973). Solid amide salts with mineral acids may be of the 1 : l or 2 : l type (e.g. hydrochlorides). Some studies of 1 : l salts have led to firm conclusions in favour of N-protonation. These are: Raman studies of acetamide hydrochloride (Kahovec and Knollmiiller, 1941), of Nmethylacetamide hydrochloride (Bonner et al., 1966), and of formamide in concentrated hydrochloric acid (Smith and Robinson, 1957); infrared and Raman studies of acetamide and urea hydrochlorides as solids and in concentrated hydrochloric acid (Spinner, 1959); an infrared study of N-ethylacetamide in carbon tetrachloride containing dry hydrogen chloride (Cannon, 1955), and an infrared study of urea nitrate (Davis and Hopkins, 1957). It may be noted that in these studies mostly primary and secondary aliphatic amides were used and most often the acid was hydrochloric. The conclusions rest largely on the shift to higher frequencies of the characteristic strong band ascribed to the stretching vibration of the C=O bond (a shift to higher frequencies is expected for a strongly electronwithdrawing substituent such as -NH:, in view of the correlation in Figure 3). Doubts have been cast on these conclusions by Janssen (1961) and Cook (1964), who studied infrared spectra of the salts of dimethylacetamide. The strong C=O band in the solid amide appears at 1639 cm-'. In the solid hydrochloride a band appears at 1693 cm-', which was assigned to the C=N bond in the 0-protonated cation (Janssen, 1961), but could equally be due to the C=O vibration in the N-protonated cation. Rarnan spectra of dimethylacetamide (2 M ) in conc. hydrochloric acid show the same intense band at 1693 cm-' (de Loze et al., 1972). In view of the high concentrations of both the amide and the acid, it is probable that the amide is present in these solutions largely as the 0-protonated cation, so that the assignment of this band to the C=N frequency seems likely, especially for a tertiary amide. Infrared and Rarnan spectra are thus capable of interpretation in terms of both N- and 0-protonation, depending to some extent on the amide (primary, secondary or tertiary), .and cannot provide unambiguous evidence for the site of protonation. The infrared findings of Gompper and Altreuther (1959) on benzamide salts, claimed to indicate 0-protonation in perchlorates, are equally inconclusive. The carbonyl frequency at 1656 cm-' in solid benzamide (in KBr) appears in the perchlorate in almost the

ALTERNATIVE PROTONATION SITES

339

same position, but there is a shift t o higher frequency in Nethylbenzamide and its salt (1633 cm-' t o 1677 cm-'). 0Methylated N,N-dimethylbenzamide perchlorate absorbs at 1639 cm- in dichloromethane, whereas the N-ethylated salt absorbs at much higher frequencies (1724 cm-' and 1783 cm-'), which led to the conclusion that 0-protonation occurred in benzamide perchlorate, since it showed n o strong bands in this high frequency region. Some observations on benzamide hydrogen sulphate are of interest in this connection (Liler, 1974, and unpublished results). This salt has been found to crystallize from sulphuric acid-water mixtures, of less than 75% acid content, when attempts are made to dissolve 1 mole of amide per litre (>75% acid dissolves this amount readily). Thus the salt crystallizes out of solvent mixtures which contain predominantly the N-protonated cation, and it therefore presumably contains such a cation. The crystals (m.p. 12OoC) are also obtainable by mixing benzamide and 100% sulphuric acid in the molar ratio 1 : l . Their infrared spectrum (in KBr) shows a strong absorption at 1657 cm-' (like the perchlorate, i.e., n o shift relative to the free solid amide), but other bands associated with the primary amide group, namely, the strong band ascribed t o NH-deformation vibrations at 1618 cm-' and the strong C-N stretching vibration at 1400 cm-' , disappear. Instead, a medium strong broad absorption appears between 1505 and 1525 cm-' , in a region where vibrations of the -NHg group are observed (Cross, 1964, p. 61). (The bisulphate ion shows only a weak absorption at 1630 cm-' over this spectral range.) These facts are consistent with the N-protonated cation structure. The absence of a shift of the carbonyl vibration may be due to enhanced conjugation with the ring in the cation [ 1181,

'

which would counter any shift of the frequency t o higher values produced by the -I effect of the -NH: group. It may also be mentioned at this point that the carbonyl absorption of the cations of 2,2-dimethylquinuclidone-6,which is certainly N-protonated, is found at the very high frequency of 1799 cm-' (Pracejus, 1959). This bicyclic system is strained, however, and strain is known to increase the carbonyl frequency in cyclic ketones (Cross, 1964, p. 59).

340

M. LILER

Several lines of kinetic evidence have also to be considered in connection with the site of protonation of amides. Salts of the type [119] ( R = Me or Ph) have been prepared by Klages and Zange R-G<

0 NEt:

SbCl,

(1957) and have been found t o react with water only very slowly in the cold (a statement to the contrary was made by Katritzky and 196 1). The N-protonated cation of 2,2-dimethylJones, quinuclidone-6 also hydrolyses relatively slowly (half-life, 16 min at 2OoC) (Pracejus, 1959), even though the bicyclic ring possesses eclipsing strain and also some angle strain. Depending on the acid concentration, amides in dilute acid solution exist only fractionally as cations and hydrolysis is slow. It has been suggested by Smith and Yates (1972) that the acetylpyridinium [120] and acetylimidazolium [121] cations may be regarded as models for Nprotonated acetamide and that comparable rates of hydrolysis are t o

be expected for these cations and for protonated acetamide, if it were N-protonated. For the purpose of comparison, the measured rate constants for hydrolysis of acetamide were recalculated using known degrees of protonation of acetamide, t o yield the rate constant for the rate-determining step, namely the nucleophilic attack by water molecules upon the cation. Since the calculated rate constant R (pseudo-first order, for reaction with solvent molecules in aqueous solution) was found t o be three orders of magnitude smaller than that for the hydrolysis of acetylimidazolium ions and some six orders of magnitude smaller than that for acetylpyridinium ions, Smith and Yates (1972) concluded that the concentration of the N-protonated cation must be about l o 6 times less than the total protonated amide, i.e., that the dominant form in dilute acid must be 0-protonated. The argument that acetylammonium ions should have the same reactivity as acetylpyridinium ions (because of the localized charge on the N) is invalidated by consideration of the nature of

34 1

ALTERNATIVE PROTONATION SITES

the rate-determining step in the reaction. This is the formation of the tetrahedral intermediate [ 1221, since it is well known that only

OH 11221

slight oxygen exchange occurs in the acid-catalysed hydrolysis of benzamide (Bender and Ginger, 1955). The rate of formation of the intermediate is determined by the electron density at the site of nucleophilic attack, i.e., at the carbonyl carbon. This will vary with the electron donating ability of the amine part of the molecule. The rates of nucleophilic attack should thus be a function of the pK,-value of the amine, and a plot of the relative rates of Smith and Yates (1972) for acetyl compounds us. the pK,-values of the amines (Figure 7) shows that this is indeed the case. It follows that protonated acetamide in aqueous acid is as reactive towards hydrolysis as may be expected for the N-protonated cation. Hence the rate of hydrolysis is consistent with dominant N-protonation. Another feature of the kinetics of acid-catalysed hydrolysis of amides may be of relevance in connection with the nature of the protonated cation, that is the rate maximum which is found at 3- 6 M acid for all amides. The usual interpretation of this is that increasing

, 3

4

5

6

7

8

9

10

PKa

Figure 7. A linear free energy relationship for rates of hydrolysis of cations CH3CONf*, consistent with N-protonation of acetamide in aqueous acid.

342

M. LILER

extents of protonation are counteracted by decreasing availability of water for the reaction. An alternative suggestion has been made recently, that decreasing rate at higher acid concentrations may be due to decreasing concentrations of the N-protonated form in favour of the 0-protonated form (de Lockerente et al., 1970). This is certainly not the case for primary amides, but may be true for tertiary amides. It has been assumed in all these considerations that 0-protonated cations do not undergo hydrolysis. Many textbooks on reaction mechanism (e.g., Breslow, 1969) assume on the contrary that a tetrahedral intermediate is formed from the 0-protonated cation in the first instance, and then undergoes a rapid prototropic change into intermediate [ 1221. It is accepted that only intermediate [ 1221 is consistent with the virtual absence of oxygen exchange in acidcatalysed hydrolysis (see also Liler, 197 l c , pp. 198- 201). The recognition of dominant N-protonation of amides in dilute acid removes the necessity for assuming two types of intermediate in the hydrolysis of amides. Contrary to some earlier conclusions (Kurland and Wilson, 1957), more recent microwave studies of formamide have shown that the molecule is non-planar, the group C- NH2 forming a shallow pyramid (Costain and Dowdling, 1960). This means that the lone electron pair on the nitrogen is available for protonation to some extent. Numerous molecular orbital calculations on the structure of amides (primarily formamide) have not been very helpful in deciding the site of protonation, firstly, because the conclusions about charge distribution depend so much upon the assumptions made, and, secondly, because solvation effects are not taken into account. As an illustration of the first statement, the paper by Basch e t al. (1967) may be quoted, where calculations on formamide, considering the n-electrons only, yielded fractions of electrons on C, N and 0 as +0.31, +0.19 and -0.50, respectively (+ deficiency, - excess of n-electrons), but when account was taken of the redistribution of o-electrons as well (donation of charge by the three hydrogens), the charge distribution changed considerably, turning N from electron deficient (+0.19) to strongly electron rich (-0.75). T o illustrate the second statement, we may quote the theoretical calculations by Mtly and Pullman (1972) on formamide with particular reference to the approach of a positive charge, i.e., protonation, which take account of polarization and transfer of charge within the molecule, but disregard solvation. The result that the N-protonated form, while

ALTERNATIVE PROTONATION SITES

343

stabilized with respect t o the amide structure, is some 30 kcal mole-' less stable than the O-protonated form, therefore refers to gas phase protonation. Molecular orbital calculations by Hopkinson and Csizmadia (1973) put the same difference at only 6.2 kcal mole-', again ignoring solvation. Protonation studies of methyl and ethyl carbamate [123] and their N-methyl and N,N-dimethyl derivatives (Armstrong and 0

11

Eto4-N\H

,H

(or Me) (or Me)

Moodie, 1968), by measurements of internal chemical shifts in aqueous sulphuric and perchloric acid, have shown that their behaviour is generally similar to that of amides, their basicities being intermediate between those of amides and esters. The ionization ratios follow the H A acidity function only moderately well, and half-protonations occur at HA-values between -3.0 and -3.3. In superacid solutions (HS03F-SbF, -SOz ) these substances are protonated on the carbonyl oxygen (Olah and Calin, 1968), but this does not imply dominant O-protonation in dilute acid any more than it does for amides. It is interesting that for N,N-di-isopropylcarbamate N-protonation has been found in 90-98% sulphuric acid or fluorosulphuric acid at temperatures below 0°C (Armstrong et al., 1968). The broad NH-peak disappears in 80% acid and at higher temperatures owing to accelerated exchange. It has been suggested that, with bulky isopropyl groups, steric hindrance to coplanarity destabilizes the O-protonated form in favour of the N-protonated form, but the same does not apply to N,N-di-isopropyl derivatives of acetamide and benzamide which are O-protonated in anhydrous acids. Olah et al. (1971a) investigated methyl and ethyl N,N-diisopropyl carbamates further in HSO, F-SO, ClF, a solvent system in which temperatures as low as -120°C are attainable. At this temperature carbonyl oxygen was found to be protonated, but when the temperature was raised to -3OoC, rearrangement took place to the thermodynamically stable N-protonated form. The protonation of more complex molecules containing amide and carbamate elements of structure has also been studied in HS03FSbF,-SO, mixtures by Olah and his collaborators. Thus allophanates [124] were found to be protonated in that medium at

344

M. LILER

0

0

II

II

H2 N--C-NH--C?)R ~ 2 4 1

-60°C on both carbonyl oxygen atoms (Olah et al., 1971b). N-alkoxycarbonyl-substituted amino-acids are also carbonyl protonated in that medium (Olah and Brydon, 1970). Some simple peptides and insulin have also been studied (Olah et al., 1970a). An interesting example of an amide which appears to remain N-protonated in pure fluorosulphuric acid at low temperature is acrylamide (Farona et al., 1969). In this molecule, the carbonyl group enters into conjugation [ 1251 with both the amino-group and

~

5

1

the ethylenic double bond. Owing t o this competition, the lone electron pair on the nitrogen atom is more readily available for protonation than in an ordinary amide. Molecular orbital calculations predict smaller localization energy for the lone electron pair on the nitrogen than on the oxygen atom, much more so for acrylamide than for formamide (Farona et al., 1969). In accordance with this theoretical prediction, nmr spectra of acrylamide and N,N-dimethylacrylamide in pure fluorosulphuric acid suggest N-protonation. In particular, the N,N-dimethyl group resonance of N,N-dimethylacrylamide remains a sharp singlet down to -50°C, in contrast to N,N-dimethylacetamide, in which restricted rotation around the C-N bond in the O-protonated cation gives a clearly resolved doublet even at room temperature (Gillespie and Birchall, 1963; Liler, 1972a). While sulphonamides [126] may be considered to be formally similar to amides, the sp3-hybridized sulphur atom has no vacant

p-orbitals for mesomeric interaction with the amino-group. The basicities of aliphatic sulphonamides, determined from nmr chemical shift measurements, fall in the range of -5.5 t o -6.0 on the H , scale

345

ALTERNATIVE PROTONATION SITES

(Laughlin, 1967). The proton has been shown t o add on to the nitrogen in pure fluorosulphuric acid (Birchall and Gillespie, 1963). The large base-weakening effect of the CH, SOz -group upon the amino-nitrogen atom (ca. 16 pK-units) is difficult to explain by a purely inductive effect, however. In fact, protonation deshields the S-CH, group more than the N-CH3 groups, which suggests that there is a certain amount of double bond character in the S-N bond (Laughlin, 1967). The basicities of aromatic sulphonamides have been determined spectrophotometrically (Virtanen and Maikkula, 1968). Benzenesulphonamide behaves as a Hammett base and has a pK,-value of -6.64. The protonation of N,N-disubstituted sulphonamides follows the Hz’ (tertiary aniline) acidity function (Virtanen and Heinamaki, 1969). In the cyclic sulphonamide [ 1271, N-protonation has been confirmed and it has been shown that NH-exchange occurs with retention of configuration (Menger and Mandell, 1967). The conclusion regarding the site of protonation of phosphinamides has been reached rather indirectly, from the effect of protonation on coupling across the P-N bond. Phosphinamides [128] are very labile in acid solution (Haake and Koizumi, 1970),

0 II Ph-P-NH2

Mefp-FJUvle

(or NMe2)

(or N R 2 )

Me

[I281

Me ~

9

1

which suggests that a good leaving group (-NHS) is present in the protonated compound. The ring compound [ 1291 is more stable, and its protonation could be studied by chemical shift measurements (Haake and Koizumi, 1970). The pK,-values of -3-2 and -1.7 for the N-methyl and N,N-dimethyl derivatives show a large effect of N-methyl substitution, thus suggesting N-protonation. The P-N-CH coupling constant decreases between solutions in chloroform and in sulphuric acid, probably due to a decrease in the P-N bond order consequent upon N-protonation. This coupling constant criterion was supported by a study of 0-alkylated salts, but caution is needed if changes in coupling constants are small (Debruin et al., 1971). In nitrosamines [130] an amino-group interacts with a nitrosogroup in a manner akin to its interaction with a carbonyl group,

346

M. LILER

since the vibration frequencies of nitroso-compounds X-NO are linearly correlated with the carbonyl frequencies of compounds R-CO-X (Bellamy and Williams, 1957). The protonation of N,Ndimethyl and N,N-di-isopropyl-nitrosamine has been studied spectrophotometrically in sulphuric acid- water mixtures (Layne et al., 1963). The absorption maximum of the first compound a t 230 nm in water, ascribed to the T - T * transition ( E = 6800), undergoes a hypsochromic shift up t o about 6.5 M acid, while the low intensity absorption at 345 nm ( E = 65) decreases steadily with acid concentration. At still higher acidities there are further shifts in the absorption, and Layne et al. (1963) interpret the results in terms of two equilibria and invoke hydrogen bonding with the solvent in order to explain this. This is a tentative explanation and the general position seems to be similar to that of amides. Nmr spectra of N,N-dimethylnitrosamine in pure fluorosulphuric acid at
Ring systems In ring systems the most important possibility of NfO conjugation is when oxygen is exocyclic. Here belong firstly cyclic aliphatic amides, lactams. The basicities of some lactams and thiolactams have been measured over the past ten years by ultraviolet spectrophotometry in sulphuric acid- water mixtures. pK,-values of -0.94 for pynolidone and of -0-92 for N-methylpyrrolidone (Virtanen and Sodervall, 1967) are closely similar t o the pK,-values of amides, and point to N-protonation in aqueous acid. Thiopyrrolidone with a pK,-value of -2.0 (Edward and Stollar, 1963) is only a little less basic and therefore also probably N-protonated in aqueous acid. Thiopiperidone and thiocaprolactam (pK,-values -1.4 and -1.6 respectively) are of similar basicity (Edward and Stollar, 1963). Measurements using an indicator method (Brkant and Dupin, 1969) gave pK, = -0.17 for N-methylpyrrolidone, i.e. well within the amide pK, -range. Nmr chemical shift measurements in aqueous sulphuric acid gave for E-caprolactam a pK,-value of -0.46 (Farber

ALTERNATIVE PROTONATION SITES

347

and Brieux, 1966). The protonation of e-caprolactam with dry hydrogen chloride in chloroform was found, however, to occur on the oxygen (Ottenheym et al., 1961), but in this solvent the interaction may be predominantly of a hydrogen bonding type. When two carbon atoms of an aromatic six-membered ring are as in 2-pyridone, the replaced by an amide group (-CONH-) aromatic character of the ring is not lost, owing t o the importance of

resonance [131]. With replacement of further pairs, as in uracil [ 1321 and cyanuric acid [ 1331, aromatic character is still preserved

A&

HN$

0

H

(Albert, 1968, pp. 62-3). Protonation on the nitrogen would have to destroy this high stabilization and it does not occur. Oxygen is protonated instead. 4-Pyridone is stabilized by a similar type of resonance [ 1341 and the site of protonation is also oxygen. Evidence

for this comes from ultraviolet and nmr spectra and has already been reviewed (Katritzky and Jones, 1961). In more recent work (Katritzky and Reavill, 1963; Cox and Bothner-By, 1969), the cations of 2-pyridone and derivatives have been fully characterized by nmr in sulphuric acid, deuteriosulphuric acid and deuterium chloride- deuterium oxide mixtures. It was also shown, however, that

M. LILER

348

the cations of 4-pyridone exchange the NH-proton fairly rapidly in 63% deuteriosulphuric acid (van der Haak and de Boer, 1964). The observation of the OH-stretching frequency in the 3500 cm-* region of the infrared spectra of the solid hexachloroantimonates has also been reported (Bell et al., 1963). This band was apparently not discernible in the hydrochloride salts (Spinner, 1960b). The protonation of 2-pyndones follows the H A acidity function (Brignell et a/., 1968). The 1,5-dimethyl derivative has a pK,-value of 0.79, i.e., it is more basic than open-chain amides. 2-Hydroxypyrimidine and 4-hydroxypyrimidine, which exist predominantly as tautomers [135] and [136] (Gronowitz and Hoffman, 1960), become monoprotonated in trifluoroacetic acid on

r2

0

H

[I351

0

'

4,

HN?

~ 3 6 1

nitrogen to give cations [ 1371 and [ 1381, and only the second protonation in fluorosulphuric acid occurs on the oxygen atom (Wagner and von Philipsborn, 1970). 0

0 H J,J H [I371

11381

The first protonation of uracil in trifluoroacetic acid was found by nmr spectroscopy to occur at the oxygen atom in the 4-position, giving a cation with extensive charge delocalization [ 1391, and the second protonation in fluorosulphuric acid occurs at the oxygen in position 2, giving cation [140], in which the two charges are delocalized over the whole molecule (Wagner and von Philipsborn,

ALTERNATIVE PROTONATION SITES

349

1970). Infrared evidence on alkyluracyl salts supports oxygen protonation (Cook, 1966). The pK,-value of the mono-cation is -3.38, according t o Katritzky and Waring (1962), and -2.07, according to Antonovskii et al. (1972). Monoprotonation of barbituric acid [ 1411 and its 1,3-diphenyl derivative in concentrated sulphuric acid occurs at oxygen atoms in

positions 4 or 6 (Bobranski, 1969). The 5,5-diethyl and the 5-ethyl5-phenyl derivatives in concentrated sulphuric acid are monoprotonated at the oxygen atom in position 2 to the extent of 66%, according to cryoscopic and conductometric measurements. In the same medium, 5,5-diallylbarbituric acid is diprotonated, at the oxygen atom in position 2 and at one of the ally1 groups (Bobranski, 1969). The protonation site of pyridazone [ 1421 appears also to be the exocyclic oxygen, since a hypsochromic change occurs in the ultraviolet absorption upon protonation (Cookson and Cheeseman, 1972). The pK,-values of the 6-substituted derivatives follow a correlation with up-constants of the substituents: pK, = 1.122 - 2.920,.

(36)

Pyridazinethiones have also been studied (Albert and Barlin, 1962). When two amide groups in a six-membered aromatic ring are “in opposition” t o each other, as in maleic hydrazide, the 0x0 structure [143] is not so highly stabilized and tautomerism is observed in which the 6-hydroxy-3-pyridazone form [ 1441 predominates (Katritzky and Lagowski, 1963). The pK,-value of this compound is

350

M. LILER

-0.97 and the tautomeric equilibrium appears to be affected by changing acidity (Cookson and Cheeseman, 1972). Similar studies have recently been extended to aminopyridones and thiones and to aminopyrimidones (Barlin and Pfleiderer, 197 1; Barlin, 1972). Protonation of 3- and 5-amino-2-hydroxypyridine and 3,4-diamino-Z-hydroxypyridine occurs first at the 3- (or 5-) aminogroup, but 4- and 6-amino-2-hydroxypyridine and 2- and 3-amino-4hydroxypyridine first protonate at the oxygen atom, because the amino-group is involved in conjugation with the partially positively charged ring nitrogen. Similar results are found for the aminopyridine-2(and 4)-thiones (Barlin, 1972). The site of protonation of amino-substituted pyridazones [ 1451 is not firmly established (Cookson and Cheeseman, 1972). Protonation on the ring nitrogen leading t o a cation stabilized by amidinium-type resonance appears unlikely, since it would be expected t o produce a bathochromic shift in the ultraviolet absorption. The reverse is observed and therefore oxygen protonation is more likely. 0x0-derivatives of pyrazole also contain an amide group replacing two carbon atoms in the aromatic pyrrole ring. Antipyrine [ 1461 is

such a derivative. Its site of protonation has been shown to be oxygen by an infrared study of its salts (Cook, 1963). Oxygen protonation preserves the aromatic sextet and the pK,-value of 1.45 shows enhanced basicity as compared with that of amides. Sydnones, which contain a cyclic ylid type structure [ 1471, have one nitrogen and two possible oxygen protonation sites. Calculations quoted by Coulson (1952) show the largest negative charge on the carbonyl oxygen and protonation at that site has been observed by

ALTERNATIVE PROTONATION SITES

35 1

nmr spectroscopy in H S 0 3 F-SbF, (Olah e t a f . , 1 9 7 0 ~ )The . formulation of a mechanism of acid-catalysed hydrolysis in aqueous acid is, however, easier in terms of the N-2 protonated species (Aziz et at., 1971). While it has been known for some time that pyndine-1-oxides are protonated on oxygen (Gardner and Katritzky, 1957), aminopyridine-1-oxides have attracted more attention recently. Using the Hammett equation, it was concluded that the cation of 3-aminopyridine-1-oxide was an equilibrium mixture of 0-and N-protonated forms in the ratio 8 : l (Jaffk, 1965). A similar conclusion was reached for 3-dimethylaminopyridine-1-oxide (Forsythe et al., 1972). In these compounds there is no conjugation between the amino-nitrogen and oxygen basic centres, but the 2- and 4-amino derivatives may be thought of as involving conjugation [ 1491. This is

supported by their enhanced basicity (pK,-values 2.27 and 3.88, respectively) as compared with that of the 3-derivative (pK, = 1.92) (Forsythe e t al., 1972). They are also predominantly 0-protonated.

N / C Alternatives Systems with alternative protonation sites on nitrogen and carbon are numerous. Open-chain and alicyclic systems may conveniently be treated together. This will then be followed by systems with an exocyclic nitrogen or carbon, and finally by aromatic systems. Open-chain and alicyclic systems Open-chain systems containing conjugated N and C basic centres are found in enamines [ 1501 (sometimes called vinylogous amines) and the isoelectronic hydrazones [ 15 11 . \

0.

,N-C=C,/

I

-

\ +

/N=C-C,

I

- /

35 2

M. LILER

Basicity measurements on open-chain and cyclic enamines have led to apparently conflicting results. While Adams and Mahan (1942), Leonard et a/., (1955) and Leonard and Hauck (1957), studying 6-membered cyclic enamines bearing alkyl groups in the 2-position [152], found them to be more basic than the corresponding saturated amines (e.g., [152], R = Me has a pK,-value of 11.4 as compared with 10-3 for a tertiary amine), Stamhuis et al. (1965), studying open-chain tertiary enamines [ 1531, found them to be less

H3cYCH3 QCH3 R

~521 (R = H or Me)

[I531 (piperidino, pyrrolidino and morpholino compounds)

basic than tertiary saturated amines. The results of Stamhuis et al. (1965) could be rationalized with the help of earlier pK,-measurements on quinuclidine and dehydroquinuclidine (Grob et al., 1957) to indicate a weakening effect of the C=C bond on the nitrogen basicity in the manner already discussed for aniline and amides (see pages 280 and 329):

Q ..

pK, b K a

1095

9.82

1.13

8.4-8.7

-1.2

These pK,-values show that the inductive effect of the C=C bond accounts for about one half of the reduction in the basicity of the enamine nitrogen atom, the other half being due to the mesomeric effect (Stamhuis et al., 1965). The localization energy of the lone pair o n the nitrogen, calculated from the second ApK,, is only about 1.6-1.7 kcal mole-’ (6.7-7-1 kJ mol-I). There are indications from the substituent effects on the stretching frequency of the ,C=CH2 bond that this group is an electron sink of the same kind

-..

ALTERNATIVE PROTONATION SITES

353

as the carbonyl group (Bellamy and Williams, 1957; Liler, 1971, p. 91), as might have been expected from the sp2-state of hybridization of the carbon atom. It is a much less powerful electron sink, however. By contrast, Adams and Mahan (1942), Leonard et al. (1955) and Leonard and Hauck (1957) thought that they were dealing with C-protonated cations, and the latter authors provided infrared evidence for the absence of ammonium-type salts in solid perchIorates of compounds [152]. There was some evidence for the presence of C=C bonds and NH-bonds in solid salts of the same compounds without an alkyl group in the 2-position. It has been suggested more recently (Elguero et al., 1965 and 1966) that these apparently conflicting facts may be explained by assuming that both N- and C-protonation occur, depending on substituents and conditions. C-Protonation appears to be especially favoured in cyclic systems [152], probably because it relieves steric strain to some extent. The protonation of enamines usually follows the course shown in (37), i.e., N-protonation occurs first, followed in some instances by >N-C=C<

I

H+

\ +

;;N--C=C<

I

I

--+ >h=C--C-H I I

(37)

the transformation of the N-protonated form into the C-protonated form (Elguero et al., 1965 and 1966). The site of protonation also appears to depend on acidity. In 6 M HCl the N-protonated form of 1-N-morpholino-l-isobutylene is dominant. Although the N-protonated form of 1-dimethylamino-1-isobutylene'[ 1541 is also dominant in dilute acid, and the C-protonated form [155] becomes

observable only in > 1 * 5 ~HCl, the coupling of the N-methyl protons to the NH-proton is fully developed only in 6 M HCl (doublet, J = 5 Hz). Alkylation also can give either the N-alkylated or the C-alkylated product depending on the substrate and conditions. The effect of the nature of the protonating agent upon the transformation (37) has been investigated in greater detail more

354

M LILER

recently (Alais et al., 1971). Concentrated perchloric acid (70%), in which the protonating agents are hydronium ions, leads at 0°C initially to 90% of the N-protonated product of l-dimethylamino-lisobutylene [154]. This cation changes only very slowly at room temperature into the more stable C-protonated cation [ 1551 in 70% perchloric acid, but the change in ether containing hydrochloric acid is very rapid (Opitz and Griesinger, 1963). This shows that the transformation is not an intramolecular process, but requires basic molecules to transport the proton between the two sites. The greater kinetic stability of the N-protonated form in concentrated perchloric acid is due to the virtual absence of free water molecules, which play this part. The rate of the transformation increases with increasing dilution of the acid (Elguero et al., 1965). Carboxylic acids in non-polar solvents (CDCl, ) or in equimolar mixtures with dimethyl sulphoxide lead directly to very high percentages (90-100%) of the C-protonated form. This may be understood in terms of the principle of “hard” and “soft” acids and bases (Alds et al., 1971) as has already been discussed on page 300. All this leads to the conclusion that the relative stabilities of the N- and C-protonated forms of enamines are not very different and that relatively minor structural differences or differences of medium favour one form over the other. 2-Alkyl substituents especially favour C-protonation (Hinman, 1968). They certainly greatly enhance the basicity of pyrroles which are C-protonated (see page 358). Open-chain enehydrazines [ 1561 are very unstable substances, but their cyclic analogues, 3-pyrazolines [ 1571, have been the subject of

, .. .. ,N-N-C=C< I I

~ 5 6 1

extensive protonation studies by Elguero and his collaborators. There are three possible protonation sites in [157], namely, N-1, N-2 and C-4. Protonation on N-1 would correspond to the first protonation of a substituted hydrazine. The pK,”O-value of ethylhydrazine being 7.99, N-1 protonation will not occur if the adjacent enamine system is more basic. As was pointed out above, N-protonation of enamines leads to pK,-values of the order of 8.5, whereas C-protonation in alicyclic rings leads to pK,-values of >ll. Thus N-2 protonation or

355

ALTERNATIVE PROTONATION SITES

C-4 protonation would be expected t o occur preferentially in 3-pyrazolines. It is in fact found that the thermodynamically stable conjugate acid is protonated on C-4 (Aubagnac et al., 1967a and 1969), but a transitory existence of an N-protonated form could also be demonstrated (Aubagnac et al., 1967b). This was identified as N-1 protonated by measurements of chemical shifts caused by proto[ 1581 . The signal of nation of 1,2,4-tnmethyl-3-phenyl-3-pyrazoline

- PhTMe Ph

Ph

€1

+

Me’

I

MeAN;N, Me

Me

Me’

H

‘N

I

Me

the N-1 methyl group shifts downfield more than that of the N-2 methyl group, owing to the positive charge generated on N-1 in [ 1591. The unstable N-protonated form was observed in a hydrochloric acid solution in water-deuterioacetone, and evolves slowly into the C-4 protonated form [I601 (Aubagnac et al., 1969). The nature and position of substituents at C-3 and C-4 have a profound effect on the competition between N- and C-protonation. In 4,5disubstituted-3-pyrazolines protonation at C-4 occurs stereospecificsalt (Elguero ally t o yield the trans-4,5-disubstituted-P-pyrazolinium et al., 1971). Cyclic analogues of hydrazones [ 1511, 2-pyrazolines, show both protonation and alkylation on N-1, as has already been discussed on page 326. The sp2 nitrogen (which distinguishes these systems from enamines) does not appear t o play any direct part. The protonation of 1,P-diazepines may be mentioned within this section. Contrary to an earlier report, 3,5,7-triphenyl- 1,2-diazepines,

N-N

Ph

1 2/

N-N

Ph

I

Me

[I611

1621

[161] and [162], have been shown by nmr in trifluoroacetic acid, sulphuric acid and CDCl3-HC1 to be N-1 or N-2 protonated (Thomas et al., 1972).

356

M. LILER

Systems with an exocyclic N o r C In amino-substituted benzenes the resonance interaction of the amino-groups with the ring enhances the basicity of the ortho- and para-carbon atoms and the possibility of ring protonation arises. A single amino-group has been estimated t o enhance the basicity of the para-carbon atom to a pK,-value of only -11, as mentioned on page 290. Aniline is consequently N-protonated in aqueous acid. The suggestion that s-triaminobenzene may be C-protonated was first made on the basis of its electronic spectrum in aqueous acid (Kohler and Scheibe, 1956). Mataga (1963) calculated the 7-r-electron structure of the s-triaminobenzenium ion [ 1631 and showed that the

calculated transition energies agree with the observed electronic spectrum. C-protonation of s-triaminobenzene and other polyaminobenzenes has been confirmed more recently by nmr spectroscopy. Yamaoka et al. (1968) observed the signal of the methylene group at 6 -3.3 p.p.m. in acid solution. Ultraviolet spectra, however, also show some N-protonation in aqueous solution at p H z 4 . 3 . Ring protonation is the more important at room temperature, but Nprotonation becomes dominant at lower temperature. In the doubly charged cation two amino-groups are protonated and the spectrum is very similar t o that of aniline (Kohler and Scheibe, 1956). The pK,-value for the triply charged cation has been estimated t o be 0.6, that for the doubly charged cation 2.95, whereas that for the singly charged cation is 5.5, i.e., considerably enhanced compared with aniline (4.60) and rn-phenylenediamine (4.7 7). This suggests a change in the structure of the cations (C-protonation), but the stability of the N-protonated cation is only a little smaller. The C-protonated cation gives rise t o a strong absorption at 360 nm which approaches the absorption region of similarly conjugated aliphatic cations (Kohler and Scheibe, 1956). The shifting C-protonated cation-N-protonated cation-free base

ALTERNATIVE PROTONATION SITES

35 7

equilibrium was studied in more detail by analysing the pH dependence of the ultraviolet absorption intensity, and thermodynamic functions for the two protonations were derived (Yamaoka et al., 1970). The formation of the N-protonated cation is less exothermic than that of the C-protonated cation, so that the equilibrium involving C-protonation is the more dependent on temperature. Thcrc is also some C-protonation for N,N-dimethyl-rn-phenylenediamine, but none for the para-isomer. Quantum mechanical calculations on C-protonated cations of aromatic amines satisfactorily predict cases of C-protonation and the electronic transitions of these ions (Yamaoka, 1970). Amino-groups in the 1-, 3-, and 5-positions greatly stabilize the benzenium ion [ 1631, the effect of amino-groups in other positions being much smaller. An example of an N/C conjugated system involving an exocyclic carbon is the enamine conjugation in cyanine dyes, in which. the 0-carbon atom constitutes the methine bridge ( [ 1641 , n = 0), or part of the bridge ([ 1641, n = 1). Protonation of these cations occurs in

CH-(CH=CH),

I

I

the aqueous pH range (the apparent pK,-value of 1,l'-diethyl2,2'-cyanine ([164], n = 0) is ca. 4 ) and the compounds are monoprotonated in pure trifluoroacetic acid (Feldman et al., 1968). Nmr spectra show that in this solvent the proton is added t o the bridge carbon in cyanines (2,2'-, 2,4'- and 4,4'-compounds, n = 0) and to one of the carbons p to the ring nitrogens in carbocyanincs (corresponding compounds with n = 1). Aromatic ring systems Pyrrole [ 1651 is the simplest aromatic ring, in which the nitrogen may be regarded as being conjugated with the a-and 0-carbon atoms of the ring. The nitrogen lone electron pair is part of the aromatic sextet and N-protonation is not expected. Pyrrole itself is unstable in dilute acid (trimerization occurs), but in more concentrated acid solutions, in which protonation occurs, a spectrophotometric study

358

M. LILER

of very dilute solutions was possible (Chiang and Whipple, 1963). Methyl-substituted pyrroles are more stable. Their protonation behaviour follows the indole acidity function H I rather better than H , .

The pK,-value of pyrrole on this scale is -3-80. The effects of methyl substitution are very large [e.g., according to Abraham et al. (1959), the pK,-value of 2,3,4-trimethylpyrrole is 3-94]. This can only be accounted for in terms of protonation of ring carbon atoms to give cations [166] and [167] (Hinman, 1968). Nitrogen is estimated to carry about 50% of the positive charge in these cations, the rest being carried by the sp2-carbon atoms of the ring (Chiang and Whipple, 1963). Methyl groups on these carbon atoms help to delocalize the charge further. Both a-C- and 0-C-protonation are observed by nmr spectra in concentrated sulphuric acid. The /3/a ratio is independent of the medium, but introduction of an N-phenyl group into 2,5-dimethylpyrrole decreases this ratio from 0-42 to 0.19 (Whipple et al., 1963). When two ortho-methyl groups are introduced into the phenyl ring, P-protonation predominates, probably owing to steric hindrance to a-protonation. With the known a//3 ratios, it has proved possible to analyse the complex ultraviolet spectra as well (Chiang e t al., 1967). Recent theoretical calculations on pyrrole and its protonated species using the MIND0/2 method predict the reactivity sequence CY > /3 > N, in agreement with experiment (Gleghorn, 1972). The protonation of acetyl- and formyl-pyrroles occurs on the carbonyl group (Skylar et al., 1966). So, for example, in the nmr spectrum of 2,4-dimethyl-3-acetyl-pyrrole in concentrated acid the vinyl (C-5) proton resonance is retained, although there is exchange in deuteriated acid (Melent'eva e t al., 1971). The behaviour of carbethoxy-substituted pyrroles is more complicated, however. While 3-carbethoxy derivatives [ 1681 and the 2-carbethoxy derivatives unsubstituted at the 5-position [ 1691 protonate at the a-position of the ring and are not exceptional, the 2-carbethoxy derivatives with a methyl group at the 5-position [170] give evidence in the nmr

359

ALTERNATIVE PROTONATION SITES

COOEt

[168]

COOEt

H

H

[ 1691

(R = H, Me, Et)

Rlr-Me Rp-JMe

N COOEt ‘H [ 1 7 1 ] (R = H, Me or Et)

Me

COOEt H [170] ( R = H, Me, or Et) Me

H’

spectrum in 11- 12-5 M sulphuric acid of two protonated forms, a transitory one and a stable one (Struchkova et al., 1972). The transitory form is characterized by a greater downfield shift of the methylene protons of the carbethoxy group than the stable form, and it is presumed therefore to be carbonyl protonated. The stable protonated form shows a broad peak at 6 4.6-4-8 p.p.m., which appears with time and is assigned to the NH,-group (the 5-methyl group signal simultaneously becomes a triplet). This means that the N-protonated cation [ 1711 is the stable form, and the slowness of its formation is connected with the electronic rearrangement involved in the break-up of the aromatic sextet. Ultraviolet spectra also show a change with time. The phosphorus analogue of pyrrole, phaphole, has a degree of aromatic character, according to molecular orbital calculations and nmr spectra (Brown, 1962; Chuchman et al., 1971). 1-Methylphosphole has a pK,-value of 0-5 (Quin e t al., 1969), much higher than that of pyrrole. It polymerizes rapidly in aqueous acid. The site o f protonation of 1,2,5-triphenylphosphole is phosphorus according to the infrared spectra of some o f its stable salts (Chuchman et al., 1971). Indoles have been shown by ultraviolet and nmr spectroscopy to be protonated predominantly at the 3-position, giving the conjugate acids [172] (Hinman and Whipple, 1962). Hydrogen exchange at

2@ :

t--f

a\ 3

R‘ ~ 7 2 3 (R’, R2,R 3 = H or Me)

R R’

2

360

M. LILER

positions 1 and 2 also occurs through the two less stable C-2 and N-1 protonated conjugate acids. For example, skatole (R' = RZ = H, R3 = Me) in 1 8 ~ deuteriosulphuric acid exchanges the NH-proton within 60 s, and the C-2 proton within 30 minutes. Finite concentrations of these forms are not detectable by nmr (sensitivity about 5%). Indoles become protonated in dilute and moderately concentrated acid. They d o not follow the Hammett acidity function H o , but most of them generate a new acidity function HI (Hinman and Lang, 1964). The basicities of indoles, methyl substituted in the heteroring, range from +Om3 t o -4-5, and are not strictly comparable with those of pyrroles, because in pyrroles both a-C- and 0-C-protonation occur. It is noteworthy, however, that a 2-methyl group enhances the basicity by more than 4 pK,-units as compared with a 3-methyl group (pK,-values 0.28 and -4.55, respectively) owing t o the delocalization of the partial positive charge on carbon-2. Isoindole has been isolated only in the C-alkyl form [173] (Bender and Bonnett, 1966). Nmr spectra show that its protonation

WH Me

[I731

Me

[I741

in acidic solvents occurs cleanly on the a-carbon atom, giving cation [174]. In systems consisting of a pyrrole ring fused with a pyridine ring, hydrogen exchange also occurs in the 3-positions in acid solutions (El-Anani et al., 1973). Similar studies have also been carried out on condensed heterocyclic systems with a bridgehead nitrogen. Indolizine [ 1751 and its 2-methyl-, 1,Z-, 2,6- and 2,8-dimethyl- and 1,2,3-trimethyl derivatives protonate preferentially on C-3 in trifluoroacetic acid, giving cations [176] (Fraser et al., 1962). In general, when position 3 is

ALTERNATIVE PROTONATION SITES

361

unsubstituted, protonation occurs at that site, and, when it is substituted, the protonation site depends on the nature of the substituent and is also affected by substituents at positions 1,2 and 5 (Fraser e t al., 1966). 3-Methylindolizine is 21% 3-protonated and 79% 1-protonated (cation [ 1771 ). In 1,3-di-substituted derivatives, if both substituents are the same, C-3 protonation occurs. Substituents in the 2- and 5-positions of 3-methylindolizine increase the C-3 to C-1 protonation ratio, apparently by relieving steric crowding of adjacent substituents in the plane of the molecule. Protonation in dilute aqueous hydrochloric acid also leads to mixtures of species protonated on C-1 and C-3 (Armarego, 1966). Protonation of indolizine at C-3 is predicted by Hiickel molecular orbital calculations of charge densities (Coulson and Longuet-Higgins, 1947), but not by the calculated localization energies (Galbraith et al., 1961). The sulphur n-electron analogue of indolizine, pyrrolo [ 2,1431 thiazole [ 1781, and some methyl substituted derivatives protonate in

[ 1781

[I791

trifluoroacetic acid exclusively on C-5 (Molloy et al., 1965) giving rise t o a six n-electron system in the sulphur containing ring (cation [ 1791 ). The only exception is 5,6-dimethyIpyrrolo[ 2,141 thiazole, which forms a mixture of cations protonated on C-5 and C-7, in which the former is dominant. The nitrogen n-electron analogue of indolizine, pyrrolo[ 1,2-a] imidazole [ 1801 , shows more complicated protonation behaviour

q-I

than indolizine. In trifluoroacetic acid the protonation sites are C-5 and C-7, and the majority of variously substituted derivatives show variable amounts of both cations, with C-5 protonation predominating (60-90%) (Alekseeva et a/., 1972a). A methyl group on C-5 leads to predominant C-7 protonation, whereas a methyl group

362

M. LILER

on C-7 leads t o exclusive C-5 protonation. The basicity of pyrroloimidazoles is considerably greater than that of indolizine. Quantum mechanical calculations by the Huckel molecular orbital method predict C-5 as the site of protonation of pyrroloimidazole (Alekseeva et al., 1972a). Pyrrolo [ 1,2-a] benzimidazoles [ 181] (R3= H, Me, Ph; R4 = Me, CH,Ph), unsubstituted at the 1- and 3-positions, protonate in trifluoroacetic acid exclusively on C-1 (Alekseeva e t al., 1972b). A methyl substituent in the l-position leads to mixtures of C-1 and C-3 protonated forms, .the relative amounts depending on the presence and nature of substituents at C-3. Without a C-3 substituent, the extent of protonation at the position is 81%, but decreases t o 18% in the 3-methyl- and 3-phenyl derivatives; the basicity of the derivatives increases simultaneously. Alekseeva et al. (1972b) have carried out a comparison of calculated values of localization energies and free valency indices of pyrrolo [ 1,2-a] imidazole, pyrrolo [ 1,2-a]benzimidazole and indolizine. In all these molecules the a-position of the pyrrole ring is calculated to be more reactive than the &position. The free valence indices increase in the order: indolizine < pyrrolobenzimidazole < pyrroloimidazole, which is also the order of increasing basicity. The protonation of a number of azaindolizines occurs at the additional non-bridgehead nitrogen (whether in the 5-membered or in the 6-membered ring). The sole exception is 5-azaindolizine [ 1821, in which protonation at carbon leads t o the establishment of a six a-electron cation in the 6-membered ring ( [ 1831, [ 1841). Several

N”4

[I821

m N’

+

11831

m41

of its alkyl and aryl derivatives have been shown by nmr in trifluoroacetic acid to be preferentially protonated on C-1 or C-3 (Fraser, 1971). Methyl substituents at the 1- or 3-position direct the proton to the opposite site. When both C-1 and C-3 are equally substituted, protonation occurs exclusively at C-3 (Fraser, 1972). The perchlorates of 1,2,3,6-tetramethyl- and 1,3,6-trimethyl-Zphenyl-5-azaindolizine show initially N-5 protonation in trifluoroacetic acid, but equilibration takes place with the formation of cations protonated on C-3. The only other system containing a n-equivalent nitrogen and found to protonate on carbon, is 2H-cyclopenta[d] pyridazine [ 1851

363

ALTERNATIVE PROTONATION SITES

(Anderson and Forkey, 1969). In these systems protonation has been shown by nmr and ultraviolet spectra to occur at the 5- and/or 7-positions, giving cations [ 1861 and [ 1871 (R = Me or Ph). The

[I861

[1851

[1871

ultimate electron pair donor is N-2 and a pyridazinium 6-membered ring results upon protonation. Other fused ring systems in which C-protonation in the fivemembered ring leads to pyridinium or quinolinium aromatic rings, respectively, are 2-phenyl-2-pyrindine (Anderson and Harrison, 1964) and 3-acyl-1,2-dihydro-4H-P-quinindines (Kholodov et al., 1969). The corrole ring system [188], like porphyrin, contains an aromatic 18 n-electron chromophore as shown by its electronic and

GFt, , N

Me \ " H

Me

/ NH /

HN A

Me

/

Me

Et

Et

Me

Me [ 1891

[1881

nmr spectra (Johnson and Kay, 1965). Mono-protonation on N-22 occurs readily without destroying aromaticity. In concentrated sulphuric acid, green solutions are obtained containing the dication. The nmr spectrum of a solution in fluorosulphuric acid indicates that the second protonation occurs on C-5 to give the cation [189] (Broadhurst e t al., 1972).

010 (or 01s) Alternatives It has been known since the turn of the century that amongst oxygen-containing compounds crystalline salts with protonic acids are most readily formed by 2,6-dimethyl-4-pyrone [ 1901 (Collie and

364

M. LILER

Tickle, 1899). Which of the two oxygens was the site of proton addition became clear only when Hantzsch (1919) demonstrated the close optical analogy between the salts and methiodides of dimethylpyrone and pyridinium salts and argued that a benzene-like ring arises in the pyroxonium salts [ 1911. The discussion of the structure

a 0

Me

Me

Me

0 Me

[191] (R = H or Me)

11901

of these salts has made a considerable contribution to the development of the theory of resonance (for a historica1 account see Nenitzescu, 1968). While the readiness with which pyroxonium salts are formed has always indicated their greater basicity as compared with that of other oxygen compounds, few measurements of basicity have been available until recently. The most reliable early determination of the pK,-value of y-pyrone as -0.28 (Rordam, 1915), obtained from measurements of the hydrolysis of the hydrochloride, shows this base to be much stronger than either ethers, for which estimates of Ho values at half-protonation range between -3.2 and -4.3 (Amett and Wu, 1960a; Arnett et af., 1962), or aliphatic ketones (e.g., acetone, which has a pK,-value of -7.7) (Campbell and Edward, 1960; Liler, 1971, p. 124). An enhancement of basicity of this order of magnitude is consistent with aromatic stabilization of the cation [ 191J . Recent spectrophotometric determinations of the basicity of variously substituted 4-pyrones (Tolmachev et al., 1968) show them to be a little more basic than amides (e.g., 2,6-dimethyl-4pyrone, pKa = 0 2 1 ; the 2,6-diphenyl derivative, pK, = -0.21). The 2-pyrones are considerably less basic (the pK, -value of 4,6-dimethyl2-pyrone is -3.78). The fusion of pyrone rings with benzene rings leads to a reduction of basicity (e.g., chromone [ 1921, coumarin [ 1931, and xanthone [ 1941 ). Sulphur and selenium analogues ([ 1921 -[ 1941, X = S or Se) are more basic (Tolmachev et al., 1967; 0

[192] (X= 0)

0

[193] (X= 0)

I1941 ( X = O )

365

ALTERNATIVE PROTONATION SITES

Degani et a/., 1968). A comparison of the basicity of these compounds with those of the corresponding pyridone analogues (2- and 4-quinolones and 9-acridones) is given in Table 4. The order of basicities for these compounds is NH > S > Se > 0. They are all protonated on the carbonyl oxygen. Such is the resonance stabilization of the pyrone cations that they do not accept any further TABLE 4 Comparison of Basicity of Some Heterocyclic Analogues of Chromone, Coumarin and Xanthone

p K , for X = Compound

0

S

Se

-2.00 -2.05

-1.12 -1-20

-1.46

-4.9 7

-4.28 -4.70

NH

Ref.

2-27

a

b

-0.3 1

U

b

0

-4.08 -4.12

-4.03 -3.95

-0.32 -4.3 6

U

b

a Tolmachev et UL (1967). Degani et al. (1968).

protons in concentrated sulphuric acid (Oddo and Scandola, 1910; Wiley and Moyer, 1954). The infrared spectrum of the hydrobromide of 7-pyrone shows that the carbonyl stretching frequency of the free pyrone at 1639 cm-' shifts to lower frequency in the salt (vco = 1488 c m - l ) (Cook, 1961). The protonation of o-phenylene carbonate [ 1951, a further analogue of coumarin in which a C=C group has been replaced by 0,

@oh=o ' 0

- a?*-

366

M. LILER

may also be mentioned here (Balaban, 1969). The aromaticity of the oxygen-containing ring is well established. The charge separation appears to be of considerable importance since the ultraviolet spectrum does not change significantly when the substance is dissolved in concentrated sulphuric acid. The compound is simply protonated in that solvent since it is recovered unchanged when the solution is diluted with water. Its pK,-value is unknown. Crystalline salts (picrate, perchlorate) could not be prepared. A crystalline hexafluoroantimonate of vinylene carbonate is, however, obtainable from a solution in HF-SbF5-S02 (Olah and White, 1968a). Alternative oxygen protonation sites also exist in carboxylic acids, esters and anhydrides, and in carbonic acid and its derivatives. All these compounds are very weak bases and some cations are stable only at low temperature. Some crystalline adducts of carboxylic acids with strong mineral acids have been known for some time, e.g., with sulphuric acid (Kendall and Carpenter, 1914; Tutundiid e t al., 1954). Esters also form adducts, e.g. diethyl oxalate with the complex acid H4Fe(CN), (Baeyer and Villiger, 1901). Benzoic anhydride forms 2:1, 1:l and 1:2 adducts with sulphuric acid (Kendall and Carpenter, 19 14). Structural information on these solids is not available, and whether they should be regarded as true salts o r as hydrogen-bonded complexes is unknown. It has long been known, however, that carboxylic acids are monoprotonated in 100% sulphuric acid (Hantzsch, 1908; Odd0 and Casalino, 1917). The first indication that the site of protonation of carboxylic acids is the carbonyl oxygen came from a study of substituent effects on the basicity of benzoic acid (Stewart and Yates, 1960), because it was found that pK,-values of variously substituted benzoic acids show a correlation with u' substituent constants, and not with Hammett a-constants. This is expected for cations of structure [ 1961, but not OH

[ 1961

[ 1971

for [ 1971, and is found also in the protonation of other carbonyl bases, e.g., acetophenones (Stewart and Yates, 1958). Carbonyl protonation of carboxylic acids and esters has been amply confirmed in recent years by spectroscopic methods, primarily nmr. Thus Birchall and Gillespie (1965) found in the spectrum of protonated aliphatic acids in HS03F-SbF5 two peaks for the protons on the carboxyl group, but only one peak for the proton captured by ethyl

ALTERNATIVE PROTONATION SITES

367

acetate. Since this peak showed n o spin-spin splitting, the proton was assumed to be on the carbonyl oxygen. The non-equivalence of the protons on the carboxyl group has been explored in detail for formic acid in HF-BF, and HF-SbF, mixtures, in which protonated formic acid is more stable than in concentrated sulphuric acid (Hogeveen e t al., 1966 and 1967). Three kinds of OH-proton are present at -67°C and these have been ascribed t o two different species of protonated formic acid, trans [ 1981 and cis [ 1991. Three

trans (77%)

cis (23%)

~ 9 8 1

[I991

isomers of protonated thioformic acid have also been observed (Olah e t al., 1969). The 3C-H coupling constants for the three different types of proton in protonated formic acid (one direct and two indirect) have also been measured (Hogeveen et al., 1968). Nmr information on the protonated forms of other carboxylic acids is also available and has been reviewed recently (Olah et al., 1970d). Generally, the trans-form of the protonated carboxyl group [ 2001 predominates in higher carboxylic acids. This is probably stabilized by hydrogen-bonding.

’

A question that may justifiably be raised here is whether these proofs of exclusive carbonyl protonation of carboxylic acids in concentrated and anhydrous acids necessarily imply the dominance of this form in dilute acid. Evidence that this is not so for amides has been discussed on pages 328 ff. It is possible that the dternative protonated form of carboxylic acids [ 2011 is dominant in aqueous acid, but as the overall extents of protonation are small, it is not detectable by any spectroscopic method. Unlike amides, carboxylic acids become measurably protonated in quite concentrated acid (>60% sulphuric acid), which would tend to favour the formation of the protonated form with a delocalized charge. The form [201]

3 68

M. LILER

could be kinetically important in acid-catalysed esterification and oxygen exchange in dilute acid. Ultraviolet spectra of benzoic acid in sulphuric acid solutions, published by Hosoya and Nagakura (1961), show a considerable medium effect on the spectrum of the unprotonated acid, but a much smaller one in concentrated acid. The former is probably connected with a hydrogen-bonding interaction of benzoic acid with sulphuric acid which is believed to be responsible for a peculiarity in the activity coefficient behaviour of unprotonated benzoic acid in these solutions (see Liler, 1971, pp. 62 and 129). The absence of a pronounced medium effect on the spectra in >85% acid is consistent with dominant carbonyl oxygen protonation. In accordance with this, Raman spectra show the disappearance in concentrated sulphuric acid of the carbonyl stretching vibration at 1650 cm-' (Hosoya and Nagakura, 1961). Molecular orbital calculations on the structure of the carbonyl protonated benzoic acid have also been carried out (Hosoya and Nagakura, 1964). The protonation of esters in superacid solvents (HSO, F-SbF, and HF-BF,) also occurs on the carbonyl oxygen (cation [ 2 0 2 ] ) H ,OR'+ R-C

[2021

80

12031

and many protonated esters undergo cleavage at ordinary and higher temperatures (Olah et al., 1967; Hogeveen, 1967). In aqueous acid, the alternative alkoxy-protonated form [ 2031, in terms of which the mechanism of acid-catalysed hydrolysis of esters (whether A- 1 or A-2) is most often formulated, may be the more important. The protonation of lactones in HSO, F-SbF, -SO2, also on the carbonyl oxygen, leads t o two isomeric species at -60'C for most lactones (Olah and Ku, 1970). For example, there are two C=OH proton signals for 6-butyrolactone, corresponding to isomers [ 2041 and [205] in the ratio 52:48. The relative amounts of the two

isomers differ for various lactones. Exceptionally, only one isomer is observed for protonated 6-valerolactone, a-acetyl-y-butyrolactone, coumarin, dihydrocoumarin and 4-h ydroxycoumarin.

ALTERNATIVE PROTONATION SITES

369

An extensive nmr study of the protonation of acyclic carboxylic acid anhydrides in HS03F-SbF5-S02 solution at -7OOC has proved the formation of carbonyl-protonated anhydrides [ 2061 ,

which undergo rapid intra- and intermolecular proton exchange (Olah et al., 1972). In an excess of the superacid, there is evidence of double carbonyl-protonation, and in large excess cleavage occurs. Carbonic acid and dialkyl carbonates are also carbonyl-protonated in HS03 F-SbF, -SO, at low temperature (Olah and White, 1968). The trihydroxycarbonium ion [207] is stable up to O°C in the absence of SOz. This and analogous ions formed by the protonation

~

7

1

[2081

of dialkylcarbonates [208] have been fully characterized by both proton and C nmr spectroscopy. The protonation of alkyl- and arylsulphonic and sulphinic acids and a k y l sulphonates and sulphinates has also been studied in HS03 F-SbF, (+SO, ClF) solutions (Olah et al., 1970b). At -6OOC RS03H: ions are observed. It is presumed that protonation occurs on the sulphonyl oxygen, but the S-OH proton resonances could not be seen in the nmr spectra, probably because of overlapping with the solvent peak at 6 = 11- 12.8 p.p.m. For protonated methanesulphonic acid, two methyl group signals are observed in the ratio 60:40, probably owing to hindered rotation around the S-OH bond in the cation [209] and [210]. In higher homologues only one protonated isomer is present.

M. LILER

370

O/C (or SIC) Alternatives

Open-chain and alicyclic systems The simplest system which offers the possibility of protonation on oxygen or carbon is carbon monoxide. The carbon protonated cation, the formyl cation [211] , is known to be unstable from

attempts at protonation of formic anhydride in H S 0 3 F-SbF, -SO2, which led only t o decomposition products (protonated formic acid and CO) (Olah et al., 1972). Several other attempted methods of preparation of the formyl cation in superacid media, including direct protonation of carbon monoxide under 50 atm pressure and at -13OoC, have all proved fruitless. The cation is known t o exist in the gas phase, and therefore Olah e t al. (1972) suggest that its instability in solution may be due to a second protonation, on the oxygen, which leads t o the break-up of the resulting species [212] by charge- charge repulsion. The protonation of ketene, CH2 =C=O. in siiperacid solution leads exclusively t o the acetyl cation, CH3 . CO+ (Olah e t al., 1972), but initial O-protonation is not ruled out. Molecular orbital calculations confirm that the cations resulting from protonation of ketene at the 0 carbon atom are more stable than the O-protonated ones, the a-Cprotonated cations being the least stable (Hopkinson, 1973). Alkyl vinyl ethers [213], systems analogous to enamines, are readily hydrolysed by dilute acid to acetaldehyde and the 6-

CH2-CH-OR

[213]

6+

(R = alkyl)

...........

+ CH3SH-R

~ 1 4 1

appropriate alcohol. Therefore, direct evidence on the site of their protonation is lacking, but inferences from the mechanistic interpretation of the kinetics of their acid-catalysed hydrolysis point t o alkoxy carbonium ions [214] as intermediates (Jones and Wood, 1964). The rates of hydrolysis increase in the order t-Bu > i-Pr > Me (Ledwith and Wood, 1966). For a wider range of groups R, a linear correlation exists between the logarithms of the rate constants for hydrolysis and a*-constants of the substituents, which supports the

37 1

ALTERNATIVE PROTONATION SITES

view that substituents enhance the conjugation shown in [213] by an inductive effect on the electron densities (Trifonov e t af., 1969). The 0-protons of ethyl vinyl ether undergo acid-catalysed hydrogen exchange also via cations [214] (Kresge and Chiang, 1969). The protonation sites of alicyclic conjugated ethers, such as [ 2151, [ 2161 and [ 2171, have also been deduced indirectly from an

[2151

[2161

~ 1 7 1

analysis of the products of hydrolysis (Rogers and Sattar, 1964 and 1965). The hydrolysis of dienol ether [ 2 151 gave products suggesting that a-protonation was preferred. The products of hydrolysis of [216] suggest that there is 45% of a-protonation, whereas for [ 2171 there is exclusive y-protonation (Rogers and Sattar, 1965). Methyl substituents in a-or y-positions of [215] favour protonation at the unsubstituted site. The analogue of [217] with an exocyclic CH,group [ 2181 is also y-protonated. Molecular orbital calculations on the open-chain system [219] predict the observed reactivities of the a-and y-sites.

n PI81

~ 1 9 1

The type of conjugation in a-diazoketones [220] which is responsible for the partial double bond character of the C-C bond 0

0

II

R-C-CH=~=N

+--+

II -

+

R-C-CH-N=N

[2201

and the two rotational isomers of diazoacetone (Kaplan and Meloy, 1966), leaves open the possibility of a-carbon or carbonyl protonation. The latter in fact occurs in HSO, F-SbF, -SOz, and a transoid [221] and a cisoid [222] isomer of the cation have been

,c=c, A:

CH3,

HO

H

372

M. LILER

identified by nmr (Allard et al., 1969). The latter predominates in the approximate ratio of 4: 1. Aromatic systems By far the most important systems under this heading are those with an exocyclic oxygen, i.e., phenols and aromatic ethers. They show an extremely wide variety of behaviour and many of the observations are still not fully understood. The question of their site of protonation was the subject of a controversy for a few years, but it is now clear that both the 0- and the C-protonated forms [223], [224] and [225] occur, depending on the conditions (medium, temperature) and the substituents (R').

(R = H or alkyl)

In an attempt to determine the pK,-values of aromatic ethers spectrophotometrically in sulphuric acid- water mixtures, Arnett and Wu (1960b) found the spectra time-dependent due to sulphonation. Nevertheless, measurements at O°C, extrapolated t o zero time, gave sigmoid curves of absorption us. H o , characteristic of protonation phenomena. The estimated pK,-values, ranging from -5.4 t o -7.4 (on the Ho-scale of Paul and Long, 1957), were ascribed t o 0-protonation on the basis of tE,e similarity of the electronic structure and spectra to those of anilines. No simple protonation behaviour could, however, be confirmed in perchloric acid (Yates and Wai, 1965). When it became established by nmr spectroscopy that C-protonated cations were present in liquid hydrogen fluoride (MacLean and Mackor, 1962) and in fluorosulphuric acid at low temperature (Birchall and Gillespie, 1964; Birchall et al., 1964), doubt was cast on the spectrophotometric findings, until Kresge and Hakka (1966) pointed out that it could be expected that the 0-protonated cation of anisole (half-protonated in 7 7% sulphuric acid) would change to the C-protonated cation a t higher acid

ALTERNATIVE PROTONATION SITES

373

concentrations owing to decreasing availability of water for hydration stabilization of oxonium ions [ 2231 . The benzenonium ions [224] and [225] are expected t o be less stabilized by hydration, owing to the delocalized positive charge (see page 294). An estimate of the basicity of the para-carbon atom of anisole suggested half-protonation in 92% sulphuric acid. The ultraviolet spectrum in 100% sulphuric acid, which differs from that in moderately concentrated acid, could not be regarded as confirming this, however, because anisole is rapidly sulphonated in concentrated sulphuric acid and is not therefore recoverable unchanged upon dilution (Ramsey, 1966). Sulphonation is even more rapid in fluorosulphuric acid (Svanholm and Parker, 1972) and occurs also in solutions of sulphuric acid and fluorosulphuric acid in trifluoroacetic acid. Nevertheless, in deuteriosulphuric acid or in CDC1, CF, COOD-D2 SO,, there is exchange of ortho- and para-protons, which indicates that C-protonation does occur, even if it is overruled by sulphonation. The suggestion of Kresge and Hakka (1966) that O-protonation changes t o C-protonation in stronger acid thus remains valid. The difference in the relative basicity of the C- and O-sites of anisole, according to their estimates, is not large, and secondary effects, such as the medium, temperature and substituents, may favour one form over the other. A clear demonstration of the importance of these effects was first produced by Brouwer et al. (1966), who showed that both 0- and C-protonated cations are observable by nmr spectroscopy in the mixed solvent HF-BF,. The ratio of concentrations of carbonium to oxonium ion was found to be strongly temperature dependent. It increases from 1.5 at -80°C to over 5 0 at O°C. This is in contrast with the situation in fluorosulphuric acid, where only the carbonium ions are observed. Numerous further examples of protonation of variously substituted hydroxy- and alkoxy-benzenes in several acid solvent mixtures have been reported more recently and show a bewildering variety of behaviour. Several hydroxy- or alkoxy-groups favour C-protonation. Thus, 1,3,5-trihydroxybenzene and its ethers have been known for some time to be C-protonated at moderate acid concentrations (35-65% perchloric acid) (Kresge et al., 1962). In partial methyl and ethyl ethers of 1,3,5-trihydroxybenzene,there is a preference for protonation para to the hydroxy- rather than the alkoxy-group (Kresge et al., 1971a). 1,3-Dihydroxybenzene and its methyl ethers are also C-protonated in perchloric and sulphuric acid, whereas C-protonation of anisole would be expected only in >70%

374

M. LILER

perchloric acid. Ultraviolet absorption was used to determine the ionization ratios of such compounds, but no single acidity function was satisfactory for the determination of their pK,-values (Kresge et a/., 1971b). The protonation of phenol and anisole in pure fluorosulphuric acid occurs on the para-carbon (Birchall et al., 1964), but the introduction of methyl substituents in the para-position leads t o competing ortho-carbon protonation (Hartshorn et af., 1971). So, for example, in the solutions of 5-methoxyhemimellitene, cations [ 2261 and [227] occur in the ratio 2:5 at low temperature. Similar results Me

li

OMe

were obtained with 3,4-dimethyl-phenol and -anisole. However, 4-methyl-, 4-isopropyl-, 2,4-dimethyl-, 2,4-diisopropyl-, 2,4,6trimethyl- and 2,4,6-triisopropyl-phenols d o not seem to be protonated in fluorosulphuric acid (Bertholon and Perrin, 1972), but some undergo decomposition, e.g. the 4-isopropyl derivative. In trifluoromethanesulphonic acid at its m.p. (-34OC) the protonation of 4-isopropylphenol occurs at the para-carbon to the extent of 30%. Trifluoromethanesulphonic acid is thus a protonating medium equivalent to fluorosulphuric acid, with the advantage that the solutions of phenols in it are more stable. Extensive nmr studies of the protonation of trihydroxybenzenes and their methyl ethers (Olah and Mo, 1972) and mono- and dihydroxybenzenes and their methyl ethers (Olah and Mo, 1973) in four different superacid media have recently been published. In the order of decreasing acidity, the media used were: (I)

SbFS-HF ( l : l , M/M) -SOZClF

(11) SbFs-HSO3F ( l : l ,M/M)--SO,CIF

(111) SbFS-HS03F (1:4, M / M ) -S02ClF

(IV) HSO3F-SO2ClF

Both C- and/or O-protonated cations were found and characterized by proton and 3C nmr spectroscopy. O-Protonation is

375

ALTERNATIVE PROTONATION SITES

generally favoured in weaker acid media, while C-protonation usually occurs in stronger superacids. Thus, 0-protonated phenol is observed in hydrogen fluoride containing small amounts of antimony pentafluoride at -105' t o -8O'C, while C-protonated phenols are found in acids of higher strength and at higher temperature. 0-Protonated ions generally exchange protons with the solvent rapidly, even in these media of very low nucleophilicity. The C-protonated ions show a much lesser tendency t o exchange with the solvent. A few particular results from these studies deserve mention. Thus, while 4-methylphenol (p-cresol) is not protonated in pure fluorosulphuric acid (Bertholon and Perrin, 1972), it is completely Cprotonated at the 2-position in superacid (11). Its methyl ether is only partially protonated on C-2 under identical conditions. CProtonation of p-dimethoxybenzene was not observed even in the strongest superacid medium (I), but rather di-0-protonation. In the least acidic medium (IV), rn-methylanisole was completely Cprotonated, while rn-cresol was only partially 0-protonated (25%) under identical conditions, which suggests that a methoxy-group can stabilize a benzenium ion better than a hydroxy-group. This may be reasonable, but it does not apply to preferred protonation sites in 3,5-dimethoxyphenol, where the protonation para to the hydroxygroup is four times more favourable (Kresge et al., 1971; Olah and Mo, 1973). There are similar inconsistencies in the findings on other phenols and phenol ethers, which defy understanding at present. Diprotonation was observed only in the strongest superacid (I). Apart from the di-0-protonated p-dimethoxybenzene already mentioned, 3,5-dimethylanisoIe gives a di-C-protonated cation [ 2281 H+

Me

Me

H H

and 2,4-dimethylanisole gives the 0- and C-diprotonated cation [229] (Olah and Mo, 1973). Ring protonation has been looked for in some other phenolic systems also. Thus a C-protonated cation is observed at low temperature when 2,4,6-trihydroxybenzoic acid is dissolved in 70% perchloric acid and pure fluorosulphuric acid (Birchall e t al., 1964). A

376

M. LILER

C-protonated cation has also been reported for 2,4,6-trimethylbenzoic acid in 99-103% sulphuric acid, in which it exists in equilibrium with the carbonyl-protonated cation and the corresponding ,benzoyl cation (Beistel and Atkinson, 1969). The latter appears in 100% sulphuric acid, increases in concentration with increasing acid concentration at the expense of the other forms and is the only form observed at concentrations >105% sulphuric acid. 4-Hydroxyazobenzene and 4-hydroxyazobenzene-4’-sulphonic acid become diprotonated in 100% sulphuric acid, the second proton going on the phenolic 0-site, rather than on carbon (Strachan et af., 1969). The alternative structural situation of an exocyclic methylene group in conjugation with a ring oxygen is found in methylene-4pyrans (vinylogous pyrones), in which a 4-pyrone-like resonance [ 2301 occurs. These compounds are little known. Their synthesis

and properties have been the subject of recent studies by Strzelecka (1966). Salts with strong acids (perchloric, picric) are known, and are undoubtedly pyrylium salts [ 23 11. In the acyl derivatives ( [ 23 1] ,

R’ = Ac or PhCO, R = Me or Ph) two cations can arise, [ 2321 and [233], both involving aromatic stabilization of the ring. In the ultraviolet spectrum of the acetyl derivative (with R = Me) there is a bathochromic shift of the maximum absorption in ethanol when perchloric acid is added, which suggests a preference for the more extensively conjugated cation [233] in the salt. The infrared spectrum of the perchlorate shows no band near 1700 cm-’ , which

ALTERNATIVE PROTONATION SITES

377

supports the cation structure [233] (Balaban et af., 1962). For the perchlorate of the benzoyl derivative (R = Ph) Strzelecka and Simalty (1968)assume cation structure [232],however, apparently on mechanistic grounds. The oxygen and sulphur analogues of pyrrole, furan and thiophen, may be expected t o behave similarly t o pyrrole in regard t o protonat i o n This is so for thiophen, but furan is sensitive t o acids and undergoes hydrolysis or polymerization, depending on the acid concentration and conditions (Albert, 1968, p. 263). Unstable cations are probable intermediates in these reactions. Protonated thiophen has been studied in hydrogen fluoride at temperatures below -60" but decomposition occurs at higher temperatures (Hogeveen, 1966). Nmr spectra show exclusive protonation at the &position for thiophen itself and several methyl substituted derivatives, including the 2,5-dimethyl derivative, which gives cation [ 2341.

Me

Me [ 2341

[235] (R = H, Ph, OH)

S-Protonated thiophens are not observed even at -9OOC. The protonation of 2-thienyl carbonyl compounds [235] in HSO, F-SbF, -SOz at -85OC occurs primarily on the carbonyl oxygen, and only the second protonation may occur on the ring (Kaper and de Boer, 1970). Thus the protonated carboxylic acid, with a methyl group in the 5-position, adds a second proton on (2-4. In fused ring systems involving the pyran ring, protonation in the homocyclic rings is favoured by conjugation with the pyran ring t o yield pyrylium salts, but alternative carbon protonation sites may arise. Systems [236] and [237] have been studied by nmr spectroscopy in trifluoroacetic acid and acetic acid with addition of 70%

F%qph \4

Ph ~ 3 6 1

perchloric acid, respectively (Boyd and Ellis, 1966). Of the two possible protonation sites, C-5 or C-7,the first is more likely in 2,4,6-triphenylcyclopenta[b]pyran [ 2361, and C-1 is certainly the protonation site of 2-phenylcyclopenta[b] [ 11 benzopyran [ 2371.

378

M. LILER

In cyclopenta[c]thiapyran [ 2381 there are also two possible protonation sites, C-5 and C-7 (Anderson and Harrison, 1964). A firm assignment of the position of protonation on the basis of the

t 2381

t 2391

~ 4 0 1

nmr spectrum in 98% sulphuric acid was not possible, but a comparison of the ultraviolet spectrum with that of the azulenium ion in the same solvent suggests that cation [240] is favoured over cation [ 2391. It may be mentioned here that 1-formyl- and 1-acetyl-azulene [241] also offer alternative C/O protonation sites, with two carbon alternatives, C-1 and C-3. Nmr spectra show that, in 1-formylazulene,

[241]

(R= H or Me)

~ 4 2 1

2431

carbonyl oxygen protonation occurs (Schulze and Long, 1964). The relative amounts of the planar cations [242] and [243] have been shown to be 1:l for 1-formylazulene, but a methyl group in the peri-position 8 favours configuration [ 2431 with the hydroxy-group in the anti-position to the tropylium ring, owing to steric hindrance (Meuche e t al., 1967). Cations of 1-acetylazulene without substituents in the 2- and 8-positions assume the anti-configuration, but a methyl group in the 8-position leads to C-1 protonated cations, in which steric crowding is reduced (Meuche et al., 1967). Similarly, in 1-nitroazulene the nitro-group is protonated rather than C-3 (Meuche and Heilbronner, 1962), but in l-nitro-4,6,8-trimethylazulene C-1 protonation occurs (Schulze and Long, 1962).

N/C/O (or N / C / S ) Alternatives The simplest systems in which N/C/O alternative sites of protonation occur are a$-unsaturated 0-aminocarbonyl compounds

379

ALTERNATIVE PROTONATION SITES

(aldehydes, ketones, esters). In these systems the conjugation with the nitrogen lone pair of electrons as shown in [244] leads t o P a

>N<=CH-C=O

I

I

-

P a! ,N=C--CH=C--O-

\+

I

I

enhancements of the basicity of the a-carbon and oxygen atoms, but, depending on the extent of electron involvement, nitrogen also remains a possible protonation site. Protonation does in fact occur at all three sites, the rates, as determined by nmr spectroscopy (see page 273), being ko > k~ S k , . C (Kramer, 1966). In open-chain systems nmr indications of predominant 0- (or C-) protonation are the restricted rotation around the CB-N bond and the slowness of NH-exchange. The nmr spectrum of P-dimethylaminoacrolein in aqueous perchloric acid shows that the non-equivalence of the two N-methyl groups is retained (Kramer and Gompper, 1964). 3-BenzyIaminoacrolein is also mainly 0-protonated (Kramer, 1966). Conformational studies on a number of 0-protonated N,Ndimethyl enaminoaldehydes and ketones in trifluoroacetic acid show that for aldehydes the s-truns conformation [ 245a] is preferred, and the s-cis-conformation [ 245b] for the ketones, decreasingly so with H0-CGR

H,c-c,

,R’ N-Me +

N-Me

1

I

Me

Me s-trans [a1

s-cis [2451

[b I

(R = H or alkyl, R’ = H or Me)

decreasing bulk of the group R (Kozerski and Dabrowski, 1972). Alkyl substitution at the a-carbon atom leads to dominant a-Cpro tonation. In more complex enaminocarbonyl systems, the dominant form of the cation depends t o some extent upon the nature of the system, e.g., whether or not part of the chain [244] is involved in a large structure, such as an alicyclic or aromatic ring, and whether or not the carbonyl group takes part in another competing conjugation, as in the ester group. Ultraviolet and infrared spectroscopy are

M. LILER

380

convenient tools for distinguishing the three possible cation structures because of their different spectral characteristics (Dolby et al., 1971). These are summarized in Table 5. Consistent conclusions using both criteria presuppose that the site of protonation is identical in solution and in crystalline salts. For example, an early report of TABLE 5 Spectral Characteristics of N-, C- and 0-protonated Cations of Enaminocarbonyl Compounds System

>~;~--cH=cH--c=o I

Ultraviolet spectrum

Infrared spectrum

Intense absorption around 300 nm

Two bands around 1600 cm-'

Absorption around 220 nm

Absorption above 1650 cm-' (unsat. carbonyl)

Insignificant absorption

Carbonyl band f imminium band around 1670 cm-'

Absorption around 300 nm

No bands between 1650 and 1800 cm-'. One or two strong bands near 1600-'.

(base)

\+

H-N--CH=CH--C=O /

I

(N-protonated) \+ / N=CH--CHz--C=O

I

(C-protonated) $=CH--CH=C--OH

I

(0-protonated)

0-protonation of 5,5-dimethyl-3-(N-pyrrolidyl)-2-cyclohexene-l-one [246],based on the infrared spectrum of its perchlorate (Leonard and Adamcik, 1959),was confirmed by ultraviolet spectra (Alt and Speziale, 1965). The protonation of the bicyclic system 12471, which is certainly N-protonated owing to steric hindrance t o mesomerism, leads t o a cation with a maximum absorption at 212 nm (in ethanolic hydrochloric acid) and three strong infrared absorptions in the solid hydrochloride at 1655 cm-' and 1720 cm-' for the double bonds and at 2430 cm-' for the NH-vibrations (Dolby et al., 1971).

u"' 0

..

Me

I

ALTERNATIVE PROTONATION SITES

38 1

However, 2,3-dihydro-4( lH)-pyridone [ 2481 is also claimed to be N-protonated (Sugiyama et al., 1969) because the absorption maximum of its hydrochloride occurs virtually at the same wavelength (325 nm) as for the base itself (327 nm). According to Table 5 , this would be an indication of 0-protonation. Monocyclic ketones ([ 2491, R = MeCO) with uninhibited mesomerism show O-proto-

OR I

hie

R

[2491 (R = MeCO or M e 0 . CO)

~ 5 0 1 (R = H or Me)

nation (Wenkert e t al., 1968), whereas analogous esters ([249] ; R = Me0 . CO) are C-protonated in solution (Powers, 1965; Wenkert et al., 1968; Dolby e t al., 1971) but N-protonated in the hydrated hydrochloride salts (Dolby et al., 1971). The lactones [250], on the other hand, are 0-protonated in both the crystalIine form and in ethanolic solution (Dolby et al., 1971). This suggests that in these systems (as in enamines themselves) all sites of protonation are of closely similar basicity, and secondary factors, such as steric, additional conjugation, or the medium, then determine the actual site of protonation. Still more complicated enaminocarbonyl compounds, with the involvement of yet another heteroatom (0 or S), are heterocyclic y&;Et

R

2

R

ygkt

~ 5 1 1 (n = I or 2, X = 0 or S )

x*>.-.

~ 5 2 1

enaminoesters [ 2511, studied by Wamhoff (1970). These have been shown by nmr spectroscopy in trifluoroacetic acid t o protonate on C-3, giving an immonium ion structure [252]. The analogous

Me

NH2

382

M. LILER

aromatic systems, furans and thiophens [ 253 J , however, protonate at the 5-position, i.e., cy t o the heteroatom, but the amino-group also contributes to the stabilization of the cation [ 254 J . Additional cation stabilization by another heteroatom is also available in 2-acylmethylene azolines [ 255 J . The protonated forms

of these compounds exist in acid solution as keto-enol equilibrium mixtures [256], indicating that both C- and 0-protonation occur OH

I

H=C-R

(Ciurdaru et al., 1970). Increasing acid concentration displaces the equilibrium towards the keto-form. These tautomeric equilibria support the view that in enaminoaldehyde, -ketone and -ester systems generally the relative basicities of the C- and 0-protonation sites are comparable, whereas Nprotonation is encountered only when steric hindrance to mesomerism is present and, exceptionally, in some solid salts. REFERENCES Abraham, R. J., Bullock, E., and Mitra, S. S. (1959). Can. J. Chem. 3 7 , 1859. Adam, W., Grimison, A., and Rodriguez, G. (1967). Tetrahedron 23, 2513. Adams, R., and Mahan, J. E. (1942). J. Amer. Chem. SOC. 64, 2588. Adler, T. K., and Albert, A. (1960).J. Chem. SOC. 1794. Alak, L., Michelot, R., and Tchoubar, B. (1971). Compt. Rend. C 273, 261. Albert, A. (1960). J. C h e m SOC. 1020.

ALTERNATIVE PROTONATION SITES

38 3

Albert, A. (1968). “Heterocyclic Chemistry.” Second edn. The Athlone Press, University of London. Albert, A., and Barlin, G. B. (1962). J. Chem. SOC. 3129. Albert, A., andTaguchi H. (1973). J.C.S. Perkin 11, 1101. Albert, A., Goldacre, R., and Phillips, J. (1948). J. Chem. SOC. 2240. Albert, A., Brown, D. J., and Cheeseman, C. (1951). /. Chem. SOC.474. Alekseeva, L. M., Dvoryantseva, G. G., Persianova, I. V., Sheinker, Yu. N., Druzhinina, A. A., and Kochergin, P. M. (1972a). Khim. Geterotsikl. Soedin. 492. Alekseeva, L. M., Dvoryantseva, G. G., Persianova, I. V., Sheinker, Yu. N., Palei, R. M., and Kochergin, P. M. (197213). Khim. Geterotsikl. Soedin. 1132. Allard, M . , Levisalles, J., and Sommer, J. M. (1969). Chem. Comm. 1515. Alt, G . H., and Speziale, A. J. (1965). J. Org. Chem. 30, 1407. Anderson, A. C., and Forkey, D. M. (1969).J. Amer. Chem. SOC. 91, 924. Anderson, A. G., and Harrison, W. F. (1964). J. Amer. Chem. SOC. 86, 708. Angyal, S. J., and Warburton, W. K. (1951).J. Chem. S O C . 2492. Antonovskii, V. L., Gukovskaya, A. S., Prokop’eva, T. M., and Sukhorukov, B. I. (1972). Dokl. Akad. Nauk. SSSR 205, 461. Armarego, W. L. F. (1965).J. Chem. SOC. 2778. Armarego, W. L. F. (1966).]. Chem. SOC.( B ) 191. Armstrong, V. C., and Moodie, R. B. (1968).]. Chem. SOC. ( B ) 275. Armstrong, V. C., Farlow, D. W., and Moodie, R. B. (1968). Chem. Comm. 1362. Amett, E. M. (1963). Prop. Phys. Otg. Chem. 1,223. Amett, E. M., andMach, C . W. (1964).J. Amer. Chem. SOC. 86, 2671. Amett, E. M., and Wu, C. Y. (1960a).J. Amer. Chem. SOC.82, 4999. Arnett, E. M., and Wu, C. Y. (1960b). J. Amer. Chem. S O C . 82, 5660. Amett, E. M., Wu, C. Y., Anderson, J. N., and Bushick, R. D. (1962)./. Amer. Chem. SOC. 84, 1674. Aston, J. G. (1955). In “Determination of Organic Structures by Physical Methods” (E. A. Braude and F. C. Nachod, eds.) p. 560. Academic Press, New York. Aubagnac, J.-L., Elguero, J., and Jacquier, R. (1967a). Bull. SOC. chim. France 3316. Aubagnac, J.-L., Elguero, J., Jacquier, R., and Tizani, D. (1967b). Tetrahedron Lett. 3709. Aubagnac, J.-L., Elguero, J., and Jacquier, R. (1969). Bull. SOC. chim. France 3292. Aubagnac, J.-L., Elguero, J., and Jacquier, R. (1970). Compt. Rend. C 270, 1829. Aziz, S., Buglass, A. J., and Tillett, J. C. (1971). J. Chem. SOC. ( B ) 1912. Baeyer, A., and Villiger, V. (1901). Ber. 34, 2679. Balaban, A. T. (1969). Rev. Roum. Chim. 14, 1323. Balaban, A. T., Frangopol, P. T., Katritzky, A. R., and Nenitzescu, C. D. (1962). J. Chem. S O C . 3889. Barlin, C. B. (1964).J. Chem. SOC. 2150. Barlin,G. B. (1972).J.C.S. Perkin 11 1459. Barlin, G. B., and Batterham, T. Y. (1967).J. Chem. SOC.( B ) 516. Barlin, G. B., and Pfleiderer, W. (1971).]. Chem. SOC. ( B ) 1425. Barnett, J. W., andO’Connor, C. J. (1973).J.C.S. Perkin I f 1331. Basch, H., Robin, M. B., and Kuebler, N. A. (1967). J. Chem. Phys. 47, 1201. Bascombe, K. N., and Bell, R. P. (1957). Discuss. Faraday SOC. 24, 158. Bascombe, K. N., and Bell, R. P. (1959).J. Chem. SOC. 1096.

384

M. LILER

Bavin, P. M. G., and Canady, W. J. (1957). Can. J. Chem. 35, 1555. Bean, G. P., Brignell, P. J., Johnson, C. D., Katritzky, A. R., Ridgewell, B. J., Tarhan, H. O., and White, A. M. (1967).J. Chem. SOC. ( B ) 1222. Beistel, D. W., and Atkinson, E. R., Jr. (1969). J. Mol. Spectroscopy 29, 244. Bell, C. L., Shoffner, J., and Bauer, L. (1963). Chem. and Ind. 1353. Bellamy, L. J., and Williams, R. L. (1957). J. Chem. SOC. 863. Bellobono, I. R., and Favini, G. (1971). J. Chem. SOC. ( B ) 2034. Bender, C. O., and Bennett, R. (1966). Chem. Commun. 198. Bender, M. L., and Ginger, R. G. (1955).J. Amer. Chem. SOC. 7 7 , 348. Benderly, H., and Rosenheck, K. (1972). Chem. Comm. 179. Berger, A., Loewenstein, A., and Meiboom, S. (1959).J. Amer. Chem. SOC. 81, 62. Bertholon, G., and Perrin, R. (1972). Compt. Rend. C 275, 645. Birchall, T., and Gillespie, R. J. (1963). Can. J. Chem. 41, 2642. Birchall, T., and Gillespie, R. J. (1964). Can. J. Chem. 42, 502. Birchall, T., Bourns, A. N., Gillespie, R. J., and Smith, P. J. (1964). Can. I. Chem. 42, 1433. Bobranski, B. (1969). Rocz. C h e m 43, 1971; C.A. 72, 78269a. Bolton, P. D., Ellis, J., Fleming, K. A., and Lantzke, I. R. (1973). Austral. J. Chem. 26, 1005. Bonner, 0. D., Bunzl, K. W., and Woolsey, G. B. (1966). Spectrochirn. Acta 22, 11 25. Bonner, T. G., and Lockhart, J. C. (1958).J. Chem. SOC. 3858. Borek, E., and Clarke, H. T. (1938).J. Biol. Chem. 125, 483. Borresen, H. C. (1967). Acta Chem. Scand. 21, 2463. Boyd, D. B. (1972). J, Amer. Chem. SOC. 94, 64. Boyd, G. V., and Ellis, A. W. (1966).J. Chem. SOC. ( B ) 349. Boyd, R. H. (1963). J. Amer. Chem. SOC. 85, 1555. BrCant, M., and Dupin, M. (1969). Compt. Rend. C 269, 306. Breslow, R. (1969). “Organic Reaction Mechanisms.” Second edn., p. 198. Benjamin, New York. Brignell, P. J., Katritzky, A. R., and Tarhan, H. 0. (1968). J. Chem. SOC. ( B ) 1477. Broadhurst, M. J., Grigg, R., Shelton, G., and Johnson, A. W. (1972). J.C.S. Perkin I 143. Brouwer, D. M., and van Doorn, J. A. (1971). Tetrahedron Lett. 3339. Brouwer, D. M., Mackor, E. L., and MacLean, C. (1966). Rec. Trau. Chim. PUYS-BQS85, 109, 114. Brown, D. A. (1962).J. Chem. S O C . 929. Brown, H. C., McDaniel, D. H., and Hafliger, 0. (1955). I n “Determination of Organic Structures by Physical Methods” (E. A. Braude and F. C. Nachod, eds.), pp. 567-662. Academic Press, New York. Bullock, F. J., Tweedie, J. F., and McRitchie, D. D. (1969). J. Chem. SOC. ( C ) 1799. Bunton, C. A., Figgis, B. N., and Nayak, B. (1959). Roc. 4th Intern. Meeting Mol. Spectroscopy, Bologna 3, 1209 (Published 1962); C.A. 61, 1740e. Campbell, H. J., and Edward, J. T. (1960). Con. J. Chem. 38, 2109. Canady, W. J. (1960). Can. J. Chem. 38,1018. Cannon, C. G. (1955).Microchim. Acto 562. Charton, M. (1965).J. Org. Chem. 30, 969. Chiang, Y., and Whipple, E. B. (1963).]. Amer. Chem. SOC. 85, 2763. Chiang, Y., Hinman, R. L., Theodoropoulos, S., and Whipple, E. B. (1967). Tetrahedron 23, 745. Chuchman, R., Holah, D. G., Hughes, A. N., and Hui, B. C. (1971). J. Heterocyclic Chem. 8, 877.

ALTERNATIVE PROTONATION SITES

385

Ciurdaru, G., Farcasan, M., and Denes, V. I. (1970). Rev. Roum. Chim. 15, 1975. Cobum, W. C., Thorpe, M. C., Montgomery, J. A., and Hewson, K. (1965). J. Org. Chein. 30, 1110, 1114. Cochran, W. (1951). Acta Cryst. 4, 81. Collie, J. N. and Tickle, T. (1899).]. Chem. S O C . 75, 710. Congdon, W. I., and Edward, J. T. (1972).J. Amer. Chem. S O C . 94, 6099. Cook, D. (1961). Can.]. Chem. 39,1184. Cook, D. (1963). Can. J. Chem. 41, 2794. Cook, D. (1964). Can. J. Chem. 42, 2721. Cook, D. (1966). Can.J. Chem. 44, 335. Cookson, R. F., and Cheeseman, G. W. H. (1972). J.C.S.PerkinII 392. Corsini, A., and Billo, E. J. (1970). J. Inorg. Nucl. Chem. 32, 1241. Costain, C. C., and Dowdling, J. M. (1960).J. Chem. Phys. 32, 158. Coulson, C. A. (1952). “Valence”, p. 324. Oxford University Press. Coulson, C. A., and Longuet-Higgins, H. C. (1947). Trans. Faraday S O C . 43, 87. Cox, R. H., and Bothner-By, A. A. (1969). J. Phys. Chem. 73, 2465. Craig, D. P. (1946).]. Chem. S O C . 534. Craig, D. P., and Short, L. N. (1945). J. Chem. S O C . 419. Cross, A. D. (1964). “An Introduction to Practical Infra-red Spectroscopy.” Butterworths, London. Crucge, F., Girault, G., Coustal, S., Lascombe, J., and Rumpf, P. (1970). Bull. SOC. chim. France 3889. Dallinga, G., Verrijn Stuart, A. A., Smit, P. J., and Mackor, E. L. (1957). Z. Electrochem. 61, 1019. Das Sarma, B. (1952).J. Indian Chem. S O C . 29, 217. Davis, M., and Hopkins, L. (1957). Trans. Faraday SOC.53, 1563. Davis, T. L., and Elderfield, R. C. (1932).J. Amer. Chem. SOC.54, 1499. Debruin, K. E., Padilla, A. G., and Johnson, D. M. (1971). Tetrahedron Lett. 4279. Degani, I., Fochi, R., and Spunta, G. (1968). Boll. S c i Fac. Chim. Ind., Bologna 26, 3, 31. de Lockerente, S . R., Nagy, 0. B., and Bruylants, A. (1970). Org. Magnetic Resonance 2, 179. de Loze, C., Combelas, P., Bacelon, P., and Garrigou-Lagrange, C. (1972). /. Chim. Phys. 69, 397. Denis, A., and Gilbert, M. (1968). Theor. Chim. Ac t a 11, 31. Dieffenbacher, A., and von Philipsborn, W. (1969). Helv. Chim. Acta 52, 743. Dolby, J., Hasselgren, K.-H., Nilsson, J. L. G., and Elander, M. (1971). Acta Pharm. Suecica 8, 97. Doub, L., and Vandenbelt, J. M. (1947). J. Amer. Chem. SOC. 69, 2714. Dvoryantseva, G. G., Ul’yanova, T. N., Sheinker, Yu. N., andyakhontov, L. N. (1973). Khim. Geterotsikl. Soedin. 6, 767. Edward, J. T., and Meacock, S. C. R. (1957).J. Chem. SOC. 2000. Edward, J. T., and Stollar, H. (1963). Can. J. Chem. 41, 721. Edward, J. T., Chang, H. S., Yates, K., and Stewart, R. (1960). Can. /. Chem. 38, 1518. Eigen, M. (1964). Angew. Chem. Internat. Ed. 3, 1. El-Anani, A., Clementi, S., and Katritzky, A. R. (1973).].C.S. Perkin I I 1072. Elguero, J., and Jacquier, R. (1965). Tetrahedron Lett. 11 75. Elguero, J., Jacquier, R., and Tarrago, G. (1965). Tetrahedron Lett. 4719. Elguero, J., Jacquier, R., and Tarrago, G. (1966). Tetrahedron Lett. 11 12. Elguero, J., Gil, R., and Jacquier, R. (1967a). Spectrochim. Acta 23A, 383.

386

M. LILER

Elguero, J., Gonzalez, E., and Jacquier, R. (1967b). Bull. SOC. c h i m Frunce 2998. Elguero, J., Gonzalez, E., and Jacquier, R. (1968). Bull. SOC. chim. Frunce 5009. Elguero, J., Gonzalez, E., and Jacquier, R. (1969a). Bull. SOC. chim. France 2054. Elguero, J., Gonzalez, E., Imbach, J.-L., and Jacquier, R. (1969b). Bull. SOC. chim. France 4075. Elguero, J., Jacquier, R., and Tizan6, D. ( 1 9 6 9 ~ )Bull. . SOC. chim. France 1687. Elguero, J., Jacquier, R., and Marzin, C . (1970). Tetrahedron Lett. 3099. Elguero, J., Jacquier, R., and Tizank, D. (1971). Tetruhedron 27, 123. Elvidge, J. A., Jones, J. R., O’Brien, C., Evans, E. A., and Turner, J. C. (1973). J.C.S. Perkin II432. Farber, S. Y.,and Brieux, J. A. (1966). Chem. and Ind. 599. Farlow, D. W., and Moodie, R. B. (1970). J. Chem. SOC. (If) 334. Farona, M. F., Ayers, W. T., Ramsey, B. G., and Grasselli, G. (1969). Inorg. Chim. Acta 3, 503. Feldman, L. H., Herz, A. H., and Regan, T. H. (1968). J. Phys. Chem. 72,2008. Ferndndez, B., Perillo, I., and Lamdan, S. ( 19 73). J. C.S. Perkin 11 1 37 1. Forsythe, P., Frampton, R., Johnson, C. D., andKatritzky, A. R. (1972). J.C.S. Perkin 11671. Fraenkel, G., and Franconi, C . (1960). J. Amer. Chem. SOC. 82, 4478. Fraenkel, G., and Kim, J. P. (1966). J , Amer. Chem. SOC. 88, 4203. Fraenkel, G., and Niemann, C. (1958). Proc. Nut. Acud. Sci. U.S.A. 44,688. Fraenkel, G., Loewenstein, A., and Meiboom, S. (1961). J. Phys. Chem. 65, 700. Fraser, M. (1971). J. Org. C h e m 36, 3087. Fraser, M. (1972). J. Org. Chem. 37, 3027. Fraser, M., Malera, A., Molloy, B. B., and Reid, D. H. (1962). J. Chem. SOC. 3288. Fraser, M., McKenzie, S., and Reid, D. H. (1966). J. Chem. SOC. ( B ) 44. Fuson, N., Josien, M.-L., and Shelton, E. M. (1954). J. Amer. Chem. SOC. 76, 2526. Galbraith, A., Small, T., Barnes, R. A., and Broekelheide, V. (1961). J. Amer. Chem. SOC. 83,453. Gardner, J. N., and Katritzky, A. R. (1957). J. Chem. SOC. 4375. Gerson, F. (1963). Helv. Chim. Actu 46, 1109. Gerson, F., and Heilbronner, E. (1959). Helv. Chim. Acta 42, 1877. Gerson, F., Heilbronner, E., van Veen, A., and Webster, B. M. (1960). Helv. Chim. Actu 43, 1889. Giauque, W. F., Hornung, E. W., Kunzler, J. E., and Rubin, T. R. (1960). J. Amer. Chem. SOC. 82, 62. Giger, W.,and Simon, W. (1970). Helv. Chim. Acta 53, 1609. Gillespie, R. J., and Birchall, T. (1963). Cun. J. Chem. 41, 148. Gillespie, R. J., and Peel, T. E. (1971). Advun. Phys. Org. Chem. 9, 1. Gleghorn, J. T. (1972). J.C.S. Perkin II479. Gold, V., andTye, F. L. (1952). J. Chem. SOC. 2184. Goldfarb, A. R., Mele, A., and Gutstein, N. (1955). J. Amer. Chem. SOC. 77, 6194. Gompper, R. (1964). Angew. Chem. 76,412. Gompper, R., and Altreuther, P. (1959). 2. unalyt. Chem. 170, 205. Goulden, J. D. S. (1954). Spectrochim. Acta 6 , 129. Grivas, J. C., andTaurins, A. (1959). Can. J. Chem. 37, 1260. Grob, C. A., Kaiser, A., and Renk, E. (1957). Chem. and Ind. 598. Gronowitz, S., and Hoffman, R. A. (1960). Arkiu Kemi 16,459.

ALTERNATIVE PROTONATION SITES

38 7

Grunwald, E., Loewenstein, A,, and Meiboom, S. (1957). J. Chem. Phys. 27, 630. Gutowsky, H. S., and Holm, J. C. (1956).]. Chem. Phys. 25, 1228. Haake, P., and Koizumi, T. (1970). Tetrahedron Lett. 4845. Haake, P., Cook, R. D., and Hurst, G. H. (1967).J. Amer. Chem. Soc. 89, 2650. Hammett, L. P. (1940). “Physical Organic Chemistry.” McGrawiHill, New York and London. Hammett, L. P., and Deyrup, A. J. (1932).J. Amer. Chem. Soc. 54, 2721. Hantzsch, A. (1908). 2. Phys. Chem. 61, 257. Hantzsch, A. (1919). Ber. 52, 1535, 1544. Harnsberger, B. G., and Riebsomer,‘ J. L. (1964). J. Heterocyclic Chem. 1 , 188. Harris, D. L., and Wellman, K. M. (1968). Tetrahedron Lett. 5225. Hartshom, M. P., Richards, K. E., Vaughan, J., and Wright, G. J. (1971). J. Chem. Soc. ( B ) 1624. Haselbach, E. (1970). Hefu. Chim. Acta 53, 1526. Hinman, R. L. (1968). Tetrahedron 24, 185. Hinman, R. L., and Lang, J. (1964). J. Amer. Chem. Soc. 86, 3796. Hinman, R. L., and Whipple, E. B. (1962).J. Amer. Chem. Soc. 84, 2534. Hirt, R. C., andSchmitt, R. G. (1958). Spectrochim. Acta 12, 127. Hirt, R. C., Schmitt, R. G., Strauss, H. L., and Koren, J. G. (1961). J. Chem. Eng. Data 6,610. Hoefnagel, M. A,, van Veen, A., and Webster, B. M. (1969). Rec. Trau. Chim. PUYS-BUS 88, 562. Hogeveen, H. (1966). Rec. TTQU.Chim. Pays-Bas 85, 1072. Hogeveen, H. (1967). Rec. Trau. Chim. Pays-Bas 86, 816. Hogeveen, H., Bickel, A. F., Hilbers, C. W., Mackor, E. L., and MacLean, C. (1966). Chem. Comm. 898. Hogeveen, H., Bickel, A. F., Hilbers, C. W., Mackor, E. L., and MacLean, C. (1967). Rec. Trau. Chim. Pays-Bas 86, 687. Hogeveen, H., Mackor, E. L., Ros, P., and Schachtschneider, J. H. (1968). Rec. Trav. Chim. Pays-Bas 87, 1057. Hopkinson, A. C. (1973).J.C.S. Perkin I I 795. Hopkinson, A. C., and Csizmadia, I. G. (1973). Can. J. Chem. 51, 1432. Hosoya, H., and Nagakura, S. (1961). Spectrochim. Acta 17, 324. Hosoya, H., and Nagakura, S. (1964). Bull. Chem. Soc. japan 37, 1500. Hosoya, H., Tanaka, J., and Nagakura, S. (1960). Buff. Chem. Soc. Japan 33, 850. Huang, C. M., Leiserowitz, L., and Schmidt, G. M. J. (1973). J.C.S. Perkin I1 503. Hyman, H. H., and Garber, R. A. (1959).J. Amer. Chem. Soc. 81, 1847. Jaffi, H. H. (1965). J. Amer. Chem. Soc. 77, 4445. Janssen, M. J. (1961). Spectrochim. Acta 17,475. Johnson, A. W., and Kay, I. T. (1965).J. Chem. Soc. 1620. Jones, D. M., and Wood, N. F. (1964).J. Chem. Soc. 5400. Jones, R. N. (1943). Chem. Rev. 32, 14. Jones, R. N. (1945).J. Amer. Chem. Soc. 67,2127. Jordan, F., and Sostman, H. D. (19721.1. Amer. Chem. Soc. 94, 7898. Kohovec, L., and Knollmuller, K. (1941). 2. Phys. Chem. 51B, 49. Kaper, L., and de Boer, Th. J. (1970). Spectrochim. Acta 26A, 2125. Kaplan, F., and Meloy, C. K. (1966). J. Amer. Chem. Soc. 88, 950. Kapoor, K. L. (1973). Int. J. Quantum Chem. 7, 27. Katritzky, A. R., and Jones, R. A. Y . (1961). Chem. and Ind. 722. Katritzky, A. R., and Lagowski, J. M. (1963). Adu. Heterocycl. Chem. 1, 366.

388

M. LILER

Katritzky, A. R., and Reavill, R. E. (1963). J. Chem. SOC.753. Katritzky, A. R., and Waring, A. J. (1962).]. Chem. SOC. 1540. Katritzky, A. R., and Waring, A. J. (1963). J. Chem. SOC.3046. Kendall, J., and Carpenter, C. D. (1914).J. Amer. Chem. SOC.36, 2498. Kessler, H., and Leibfritz, D. (1969). Tetrahedron 25, 5127. Kholodov, L. E., Tishchenkova, I. F., Persianova, I. V., and Yashunskii. V. G. (1969). Reukts. Sposob. Org. Soedin. 6, 1000. Klages, F., and Zange, E. (1957). Ann. Chem. 607, 35. Kohler, H., and Scheibe. A. (1956). 2. Anorg. Atlgem. Chem. 285, 221. Konishi; H., Kato, H., and Yonezava, T. (1970). Theor. Chim. Acta 19, 71. Korolev, B. A., andTitova, S. P. (1971). Zhur. Org. Khim. 7 , 1188. Kozerski, L., and Dabrowski, J. (1972). Org. Magnetic Resonance 4, 253. Kramer, H. E. A. (1966).Ann. Chem. 696, 15. Kramer, H. E. A., and Gompper, R. (1964). 2 . Phys. Chem. (Frankfurt)43, 349. Kresge. J. A., and Chiang, Y. (1969).J. Amer. Chem. SOC. 91, 1025. Kresge, A. J., and Hakka, L. E. (1966). J. Amer. Chem. SOC. 8 8 , 3868. Kresge, A. J., Barry, G. W., Charles, K. R., and Chiang, Y. (1962). J. Amer. C h e m S O C . 84,4343. Kresge, A. J., Chiang, Y., and Hakka, L. E. (1971a). J. Amer. Chem. SOC. 93, 6167. Kresge, A. J., Chen, H. J., Hakka, L. E., and Kouba, J. E. (1971b).J. Amer. Chem. SOC. 93,6174. Kwland, R. J., and Wilson, E. B., Jr. (1957). J. Chem. Phys. 27, 585. Laughlin, R. G. (1967). J. Amer. C h e m SOC.89, 4268. Layne, W. S., Jaffk, H. H., and Zimmer, H. (1963). J. Amer. Chem. SOC.85, 1816. Ledwith, A., and Wood, H. J. (1966). J. Chem. SOC.( B ) 753. Leonard, N. J., and Adamcik, J. A. (1959). J. Amer. Chem. SOC. 81, 595. Leonard, N. J., and Hauck, F. P., Jr. (1957). J. Amer. Chem. SOC. 79, 5279. Leonard, N. J., Hay, A. S., Fulmer, R. W., and Gash, V. W. (1955). J. Amer. Chem. SOC. 77,439. Lewis, G. E. (1960). Tetrahedron 10, 129. Liler, M. (1965). Chem. Comm. 244. Liler, M. (1967). Spectrochim. Acta 23A, 139. Liler, M. (1969). J. Chem. SOC.( B ) 385. Liler, M. (1971a). Chem. C o m m 115. Liler, M. (1971b). J. Chem. SOC.( B ) 334. Liler. M. f 1971cl. "Reaction Mechanisms in Sulphunc Acid." Academic Press, London and New York. Liler, M. (1972a). J.C.S. Perkin IZ 720. Liler, M. (1972b) J.C.S. Perkin I1 816. Liler, M. ( 1 9 7 2 ~ )C. h e m Comm. 527. Liler, M. (1974). J.C.S. Perkin I1 71. Lloyd, D., Mackie, R. K., McNab, H., and Marshall, D. R. (1973).J.C.S. Perkin II 1729. Loewenstein, A., and Connor, T. M. (1963). Ber. Bunsenges. Phys. Chem. 67, 280. Mackor, E. L., Hofstra, A., and van der Wads, J. H. (1958). Trans. Faraday SOC. 54, 66. Maclean, C., and Mackor, E. L. (1962). Discuss. Faraday SOC. 34, 165. Mannschreck, A., Seitz, W., and Staab, H. A. (1963). Ber. Bunsenges. Phys. Chem. 67,470.

ALTERNATIVE PROTONATION SITES

389

Marimoto, G. (1966). Nippon Kagaku Zasshi 87, 790. Martin, R. B. (1972). Chem. Comm. 793. Mason, S. F. (1958). J. Chem. SOC. 674. Mason, S. F. (1960). J. Chem. Soc. 219. Mataga, N. (1963). Bull Chem. Soc. Japan 36, 1109. Mecke, R., and Kutzelnigg, W. (1960). Spectrochim. Acta 16, 1216. Melent’eva, T. A., Filippova, T. M., Kazanskaya, L. V., Kustanovich, 1. M., and Berezovsku, V. M. (1971). J. Gen. Chem. USSR 41, 175. MCly, B., and Pullman, A. (1972). Compt. Rend. C 274, 1371. Menger, F. M., and Mandell, L. (1967). J. Amer. Chem. SOC. 89,4424. Meuche, D., and Heilbronner, E. (1962). Hefv. Chim. Acta 45, 1965. Meuche, D., Dreyer, D., Hafner, K., and Heilbronner, E. (1967). Helv. Chim. Acta 5 0 , 1178. Molloy, B., Reid, D., and McKenzie, S. (1965).J. Chem. SOC. 4368. Moriconi. E. J., and Cevasco, A. A. (1968).J. Org. Chem. 33, 2109. Nenitzescu, C. D. (1968). In “Carbonium Ions” (G. A. Olah and P. von R. Schleyer, eds.) Vol. I, p. 23. Interscience, New York. Neuman, R. C., and Hammond, G. S. (1963). J. Phys. Chem. 67, 1659. Oddo, G., and Casalino, A. (1917). Garz. Chim. Ital. 47,11, 200,232. Oddo, G., and Scandola, E. (1910). Carz. Chim. Ztnf. 40,II, 163. Olah, G. A., and Brydon, D. L. (1970). J. Org. Chem. 35, 313. Olah, G. A., and Calin, M. (1968).J. Amer. Chem. SOC. 90,401. Olah, G. A., and Ku, A. T. (1970). J. Org. Chem. 35, 3916. Olah, G. A., andMo, Y. K. (1972). J. Amer. Chem. SOL. 94, 5341. Olah, G. A., andMo, Y. K. (1973). J. Org. Chem. 38,353. Olah, G. A., and Szilagyi, P. J. (1969). J. Amer. Chem. SOC. 91, 2949. Olah,G. A., and White, A. M. (1968a). J. Amer. Chem. SOC. 90, 1884. Olah, G. A., and White, A. M. (1968b). J. Amer. Chem. SOC. 90, 6087. Olah, G. A., O’Brien, D. H., and White, M. (1967). J. Amer. Chem. SOC. 89, 5 694. Olah, G. A., Ku, A., and White, A. M. (1969)./. Org. Chem. 34, 1827. Olah, G. A., Brydon, D. L., and Porter, R. D. (1970a).J. Org. Chem. 35, 317. Olah, G. A., Ku, A. T., and Olah, J. A. (1970b). /. Org. Chem. 35, 3908. Olah, G. A., Kelly, D. P., andSwein, N. ( 1 9 7 0 ~ )J.. Amer. Chem. SOC.92, 3133. Olah, G. A., White, A. M., and O’Brien, D. H. (1970d). Chem. Rev. 70, 561. Olah, G. A., White, A. M., and Ku, A. T. (1971a). J. Org. Chem. 36, 3585. Olah, G. A., Ku,A. T., and Olah, J. A. (1971b). J. Org. Chem. 36, 3582. Olah, G. A., Dunne, K., Mo, Y. K., andszilagyi, P. (1972). J. Amer. Chem. Soc. 94, 4200. Opitz, G., and Griesinger, A. (1963).Ann. Chem. 665, 101. Ostwald, W. (1886).J. Prakt. Chem. [2], 33, 367. Ottenheym, J. H., van Raayen, W., Smidt, J., Groenewege, M. P., and VeerKamp, Th. A. (1961). Rec. Trav. Chim.Pays-Bas 80, 121 1. Pal’m, V. A. (1961). Russ. Chem. Rev. 30, No. 9,471. Parry, E. P., Hem, D. H., and Burr, J. G. (1969). Biochim. Biophys. Acta 182, 570. Paudler, W. W., and Blewitt, H. L. (1966). J. Org. Chem. 31, 1295. Paudler, W. W., and Helmick, L. S. (1968). J. Heterocycf. Chem. 5, 691. Paul, M. A., and Long, F. A. (1957). Chem. Rev. 57, 1. Pauling, L. (1939). “The Nature of the Chemical Bond.” Oxford University Press, London. Pearson, R. G. (1963). J. Amer. Chem. SOC.85, 3533. Perkampus, H.-H. (1966). Adv. Phys. Org. Chem. 4, 195.

390

M. LILER

Pernn, D. D. (1965). “Dissociation Constants of Organic Bases in Aqueous Solution.” Butterworths, London. Pernn, D. D. (1972). “Dissociation Constants of Organic Bases in Aqueous Solution: Supplement 1972.” Butterworths, London. Pfleiderer, W., Liedek, E., Lohrmann, R., and Rukwied, M. (1960). Chem. Ber. 93, 2015. Pople, J. A., Schneider, W. G., and Bernstein, H. J. (1959). “High-Resolution Nuclear Magnetic Resonance”, p. 100 ff. McGraw-Hill, New York. Powers, J. C. (1965). J. Org. Chem. 30, 2534. Pracejus, H. (1959). Chem. Ber. 92,988. Pracejus, H., Kehlen, M., Kehlen, H., and Matschiner, H. (1965). Tetrahedron 21, 2257. Price, E., Barefoot, R. D., Tompa, A. S., and Lowe, J. U., Jr. (1967). J. Phys. Chem. 71, 1608. Price, E., Person, L. S., Teklu, Y. D., and Tompa, A. S. (1970).J. Phys. Chem. 74, 3826. Quin, L. D., Bryson, J. G., and Moreland, C. G. (1969).J. Amer. Chem. SOC. 91, 3308. Ramsey, B. G. (1966).J. Amer. Chem. SOC. 88, 5358. Rao, K. G., and Rao, C. N. R. (1973).J.C.S. Perkin II889. Reeves, L. W. (1965). Adv. Phys. Org. C h e m 3, 187. Reeves, R. L. (1966). J. Amer. Chem. SOC. 88, 2240. Reeves, R. L. and Kaiser, R. S. (1969). J . Phys. Chem. 73, 2279. Reeves, R. L., and Smith, W. F. (1963). J. Amer. Chem. SOC. 85, 724. Reichman, U., Bergmann, F., Lichtenberg, D., and Neiman, Z. (1973). J.C.S. Perkin 1 7 9 3 . Ricketts, J. A., and Cho, C. S. (1961).J. Org. Chem. 26, 2125. Rochester, C. H. (1970). “Acidity Functions.” Academic Press, New York and London. Rogers, M. T., and Woodbrey, J. C. (1962). J. Phys. Chem. 66, 540. Rogers, N. A. J., and Sattar, A. (1964). Tetrahedron Lett. 1311. Rogers, N. A. J., and Sattar, A. (1964). Tetrahedron Lett. 1471. Rordam, H. N. K. (1915).J. Amer. Chem. SOC. 37,557. Roth, B., and Strelitz, J. Z. (1969).]. Org. Chem. 34, 821. Roth, B., and Strelitz, J. Z. (1970). J. Org. Chem. 35, 2696. Ryabova, R. S., Medvetskaya, I. M: and Vinnik, M. I. (1966). Zhur. Fir. Khim: 40, 339. Safta, M. (1969). Studii Cercet. Chim. 17, 107. Saika, A. (1960). J. Amer. Chem. SOC.82, 3540. Schulze, J., and Long, F. A. (1962). Proc. Chem. SOC. 364. Schulze, J., and Long, F. A. (1964). j . Amer. Chem. SOC. 86, 322. Schwarzenbach, G., and Lutz, K. (1940). Helv. Chim. Acta 23, 1162. Seccombe, R. C., and Kennard, C. H. L. (1973). J.C.S. Perkin 111. Shapiro, S. L., Parrino, V. A., and Freedman, I. (1969). J. Amer. Chem. SOC. 81, 2220. Scheraga, H. A. (1968). Adv. Phys. Org. Chem. 6 , 140. Sklyar, Y. E., Evstigneeva, R. D., and Preobrazhenskii, N. A. (1966). Khim. Geterotsikl. Soedin. 216. Smith, C. H., and Robinson, S. D. (1957). J. Amer. Chem. SOC. 79, 1349. Smith, C. R., and Yates, K. (1972). Can. J. Chem. 50, 771. Spinner, E. (1959). Spectrochim Acta 15, 95. Spinner, E. (1960a). J. Phys. Chem. 64, 275. Spinner, E. (1960b).J. Chem. SOC. 1226.

ALTERNATIVE PROTONATION SITES

39 1

Staab, H. A,, and Mannschreck, A. (1962).Tetrahedron Lett. 913. Stamhuis, E.J., Maas, W., and Wynberg, H. (1965).J. Org. Chem. 30,2160. Stewart, R., and Yates, K. (1958).J. Amer. Chem. SOC. 80, 6355. Stewart, R.,and Yates, K. (1960).J. Amer. Chem. SOC.82,4059. Stock, L. M., and Brown, H. C. (1963).Adv. Phys. Org. Chem. 1, 35. Strachan, W. M. J., Dolenko, A., and Buncel, E. (1969).Can. J . Chem. 47,

3631.

Stransky, Z., and Stuzka, V. (1968).Coil. Czech. Chem. Commun. 33,1802. Streitwieser, A., Jr. (1961).“Molecular Orbital Theory for Organic Chemists”, p. 335.J. Wiley and Sons, New York. Struchkova, M. I., Dvoryantseva, G. G., Kostyuchenko, N. P., Sheinker, Y. N., Sklyar, Yu. E., and Evstigneeva, R. P. (1972).Khim. Ceterotsikl. Soedin. 336. Strzelecka, H. (1966).Ann. Chim. [14],1, 201. Strzelecka, H., and Simalty, M. (1968). Bull. SOC. Chim. France 832. Sugiyama, N., Yamamoto, M., and Kashima, Ch. (1969).Bull. Chem. SOC. Japan

42,2690.

Sundaralingam, M. (1966).Acta Cryst. 21,495. Svanholm, U.,and Parker, V. D. (1972).J.C.S. Perkin II 962. Sweeting, L. M., and Yates, K. (1966).Can. J. Chem. 44,2395. Takeda, M.,and Stejskal, E. 0. (1960).]. Amer. Chem. SOC. 8 2 , 25. Takimoto, M. (1964).Nippon Kagaku Zasshi 85, 159. Tanizaki, Y., Kobayashi, T., and Hoshi, T. (1966).Bull. Chem. SOC.Japan 39,

558.

Thomas, M. T., Snieckus, V., and Klingsberg, E. (1972).Chem. Comm. 504. Tolmachev, A. I., Shulezhko, L. M., and Kisilenko, A. A. (1967).Zhur. Obshch. Khim. 37, 367. Tolmachev, A. I., Shulezhko, L. M., and Kisilenko, A. A. (1968).Zhur. Obshch. Khim. 38,118. Trifonov, B. A., Yemelyanov, I. S., Yaselman, M. E., Atavin, A. S., Prokop’ev, B. V., Gusarov, A. V., Vanyukhin, G. N., and Ovchinikova, M. M. (1969). Reakts. Sposob. Org. Soedin. 6 , 934. Trotman-Dickenson, A. F. (1949).J. Chem. SOC.1293. Tumbull, N. H. (1945).J.Chem. SOC. 441. Tutundiic, P. S., Liler, M., and Kosanovic, D. (1954).Bull. SOC. Chim. Belgrade,

19, 549.

van der Haak, P. J.,and d e Boer,Th. J. (1964).Rec. Trav. Chim. Pays-Bas 83,186. Veillard, A., and Pullman, B. (1963).J. Theoret. B i d . 4,37. Virtanen, P. I. O., and Heinamaki, K. (1969).Suom. Kemistilehti B 42, 142. Virtanen, P. I. O., and Maikkula, M. (1968).Tetrahedron Lett. 4855. Virtanen, P. I. O., and Sodentall, T. (1967).Suom. Kemistilehti B 40,337. Vold, R. L.,and Correa, A. (1970).J. Phys. Chem. 74,2674. Vold, R. L., Daniel, E. S., and Chan, S. 0. (1970).J. Amer. Chem. SOC.92,

6771.

Wagner, R., and von Philipsbom, W. (1970).Helv. Chim. Acta 53,299. Wagner, R., a n d v o n Philipsbom, W. (1971).Helv. Chim. Acta 54,1543. Wamhoff, H. (1970).Tetrahedron 26,3849. Wepster, B. M. (1952).Rec. Trav. Chim. Pays-Bas 71,1159,1171. Wellman, K. M., and Harris, D. L. (1967).Chem. C o m m 568. Wenkert, E., Dave, K. G., Haglid, F., Lewis, R. G., Oishi, T., Stevens, R. V., and Terashima, M. (1968).J.Org. Chem. 33, 747. Whipple, E. B., Chiang, Y., and Hinman, R. L. (1963).]. Amer. Chem. SOC. 85,26. Wiberg, K. (1964).“Physical Organic Chemistry”, p. 74.J. Wiley and Sons, New York.

392

M. LILER

Wiley, R. H., andMoyer, A. N. (1954).J. Amer. Chem. SOC.76,5706. Williams, G.,and Hardy, M. L. (1953).J. Chem. SOC. 2560. Wyatt, P. A. H. (1957).Discuss. Faroday SOC. 24, 163. Yamaoka, T. (1970).Bull. Chem. SOC. Japan 43,3086. Yamaoka, T.,Hosoya, H., and Nagakura, S. (1968).Tetrahedron 24, 6203. Yamaoka, T.,Hosoya, H., and Nagakura, S. (1970).Tetrahedron 26,4125. Yates, K., and Stevens, J. B. (1965).Can. J. Chem. 43,529. Yates, K., and Wai, H. (1965).Can. J. Chem. 43, 2131. Yates, K., Stevens, J. B., and Katritzky, A. R. (1964).Con. J. Chem. 42, 1957. Yeh, S.-J., and Jafft, H. H. (1959a).J.Amer. Chem. SOC. 81,3274. Yeh, S.-J.,and Jafft,H. H. (1959b).J. Amer. Chem. SOC.81,3279. Yeh, S.-J., and Jafft, H. H. (1959~). J. Amer. Chem. SOC. 81,3283. Zhmurova, I. N., Yurchenko, R. I., Kukhar, V. P., Petrashenko, A. A., and Kirsanov, A. V. (1971).Zhur. Org. Khim. 7 , 1027. Zhmurova, I. N., Yurchenko, R. I., Kukhar, V. P., Zolotareva, L. A., and Kirsanov, A. V. (1972).Zhur. Obshch. Khim. 42, 1954.

Author Index Numbers in italics refer to the pages on which references are 1i.sted at the end o f each article.

A

B

Aboderin, A., 194, 222 Abraham, R. J., 358,382 Adam, W., 169, 170, 171, 173, 322, 326, 382 Adamcik, J. A.. 380, 388 Adams, D. G., 166, 174 Adams, R., 352,353,382 Adcock, W.,124, 173 Adler, T. K.,319,382 Alais, L., 299, 300, 354,382 Albert, A., 307, 314, 315, 316, 318, 319, 323,347,349, 377,382, 383 Alden, R. A., 56,121, 122 Aldersley, M. F., 77, 116 Aleksanyan, V. T., 216, 222 Alekseeva, L. M., 361, 362, 383 Alexander, M. D., 68, 69, 116 Alford, J. R. 213, 223 Allard, M., 372, 383 Allen, L. C., 193, 201, 206, 222. 223, 224 Alt, G. H., 380,383 Altreuther, P., 338,386 Anderson, A. G., 363, 378,383 Anderson, B. M., 32, 116 Anderson, E., 18, 22, 23, 29, 85, 86, 87, 90, 92.93.96, 100, 114, 116. 118, 119 Anderson, J. N., 364,383 Angyal, S. J., 303, 307,383 Antonovskii, V. L., 349, 383 Armarego, W. L. F., 325, 361, 383 Armstrong, V. C. 343,383 Arnett, E. M., 293, 333, 364, 372, 383 Asami, R., 123, 175 Aston, J. G., 290,383 Atavin, A. S., 371, 391 Atkinson. E. R. Jr., 376,384 Atkinson, R. F., 88, 109, 113, 116 Aubagnac, J. L., 327,355,383 Ayers, W. T., 344, 386 Ayscough, P. B., 256,264, Aziz, S., 351, 383

Bacelon, P., 338, 385 Bacon, J., 139, 173 Baeyer, A., 366,383 Baird, R. L., 194, 222 Baker, E. B., 149, 157, 175, 203, 204, 223 Baker, R., 184, 222 Balaban, A. T., 366, 377, 383 Banerjee, S. K., 82, 116 Barefoot, K. o., 305, 390 Barlin, G . B.. 315, 322, 323, 349, 350, 383 Barltrop, .J. A., 227, 253, 264 Barnes, R. A., 361,386 Barnett, J. W., 337,383 Barry, G. W., 373,388 Basch, H.. 342, 383 Bascombc, K. N., 293,383 Batterham, T. Y., 322, 323,383 Bauer, L., 348,384 Baumann, H., 125, 177 Bavin, P. M. G., 284, 384 Bean, G . P., 3 16,384 Beck, B. H., 21 1, 222 Becker, E. D., 124, 173 Beistel, D. W., 376, 384 Belikova, N. A., 209,224 Belke, C. J., 43, 52, 116 Bell, C. L.. 348, 384 Bell, R. P., 293, 383 Bcllamy, I>.J. 306, 346, 353, 384 Bellobono, 1. R., 31 5, 384 Bender, C. O., 360,384 Bender, At. I.., 2, 3, 4, 23, 24, 29, 30, 32, 34, 36, 37, 38, 39, 44, 46, 49, 54, 58, 60, 61, 62, 66, 67, 75, 76,92, 116, 117, 121, 122,341,384 Benderly, H., 337,384 Benjamin, B.. 49, 119 Benkovic, P., 40, 58, 118 Benkovic, S. J., 2, 3, 4, 6, 7, 14, 17, 29, 30, 31. 36, 37, 45, 67, 117 Bennett, R., 360,384

393

394

AUTHOR INDEX

Benson, R. E., 137,174 Berezin, I. V., 63, 121 Berezovskii, V. M., 358,389 Berger, A., 332,384 Bergman, R. G., 214,222 Bergman, F., 324,390 Bcrgman, M., 64, 1 2 0 Bernstein, H. J., 271,390 Berson, J. A., 188, 214, 222 Bertholm, G., 374,375,384 Bethge, P. H., 64, 121 Beyersbergen van Henegouwen, G. M. J., 235,237,264 Bickel, A. F., 367,387 Bilbrey, R., 82, 121 Billo, E. J., 316, 385 Bingham, R. C., 186, 192,223 Birchall, T., 331, 335, 337, 344, 345, 366, 372,374,375,384, 386 Birktoft, J. J., 37, 39, 56,117, 121 Blackall, E. L., 20, 117 Blackburn, G. M., 33, 117 Blake,C. C. F., 28, 81,117 Blewitt, H. L., 325,389 Blow, D. M., 37, 39, 117 Bobranski, B., 349,384 Bockelheide, V., 361,386 Boll, W. A., 146, 175 Bollinger, J. M., 144, 174 Bolton. P. D., 31 1, 384 Bonner, 0. D., 338,384 Bonner, T. G., 305, 307,384 Borcid, S., 191, 223 Borek, E., 307,384 BBrresen, H. C., 324, 384 Bothner-By, A. A., 347,385 Bourns, A. N., 372,374,375,384 Bovey, F. A., 124, 173 Boyd, D. B., 324,384 Boyd, G. V., 377,384 Boyd, R. H., 294,295,384 Boyer, P. D., 4, 117 Bradbury, W. A., 13, 14, 17, 62, 117 Bradshaw, R. A., 64, 1 1 7 Brandon, N. E., 39,118 Brasem, P., 236, 246, 248,264 Brass, H. J., 60,117 Bdant, M., 346,384 Breslow, R., 68, 70, 71, 80, 117, 137, 173, 342,384 Brieux, J. A., 374,386 Brignell, P.J., 316, 348,384 Broadhurst, M. J., 363,384 Brod, L. H., 85, 86, 103, 119 Bromilow, R. H., 23, 117 Bronsted, J. N., 83, 11 7 Brouw*er, D. M.. 194, 195, 213, 222, 335, 373,384

Brown, A., 11, 110, 117 Brown, D. A., 359,384 Brown, D. J., 319,383 Brown, F., 181, 222 Brown, H. C., 178, 179, 182, 185, 186, 187, 206, 222, 224, 277, 279, 283,384, 391. Brown, J. F., 20, 122 Bruice,T. C., 2, 3, 4, 6, 7, 8,9, 11, 13, 14, 15, 17, 18, 21, 29, 30, 31, 32, 36, 37, 39, 40, 42, 45, 46, 49, 51, 55, 56, 58, 62, 67, 75, 79, 86, 88, 89, 90, 94, 97, 100, 105, 106, 109, 110, 113, 114, 116, 117, 118, 119, 120, 121, 122 Bruno, J. J., 39, 40, 58,118 Bruylants, A., 342,385 Brydon, D. L., 344,389 Bryson, J. G.. 359,390 Buckingham, A. D., 128, 129,173 Buckingham, D. A., 68,69, 70, 118 Buglass, A. J., 351, 383 Bull, H. G., 85, 118 Bullock, E., 358,382 Bullock, F. J., 319,384 Bunce, N. J., 227, 253, 264 Buncel, E., 376,391 Bunnett, J. F., 14,118 Bunton, C. A., 332,384 Bunzl, K. W., 338,384 Burke, J. J., 209, 222 Burr, J. G., 333, 389 Busch, D. H., 33, 68, 69, 116 Bush, K. J., 33, 121 Bushick, R. D., 364,383 Buss, V., 193, 224 Butler, A. R., 65, 119

C Calin, M., 343,389 Campbell, H.J., 364,384 Canady, W. J., 284,384 Canfield, R. C., 81, 118 Cannon, C. C., 338,384 Caplow, M., 33, 34, 36, 118, 121 Capon, B., 51, 83, 86, 90, 100, 102, 104, 108,116, 118 Carpenter, C. D., 366, 388 Carrigou-Lagrange, C., 338, 385 Carriuolo, J., 31, 4 0 , 1 2 0 Carter, R. E., 128, 136, 138, 174 Casalino, A., 336, 389, Casanova, J., 188, 223 Cevasco, A. A., 314,389 Chambers, T. C., 214,222 Chan, S . I., 128, 129, 175 Chan, S . O., 333,391 Chance, B., 5,118

395

AUTHOR INDEX Chang, H. S., 331, 333, 336,385 Chanley, J. D., 23, I18 Charles, K. R., 373,388 Charton, M., 303,306,384 Cheeseman, G. W. H., 316, 319, 349, 350, 383, 385 Chen, H. J., 374,388 Cheney, B. V., 208,222 Chiang, Y.,358, 371, 373,384, 388, 391 Chipman, D. J., 71, 83, 104, 117, 118 Chloupek, F., 75, 76, 1I7 Chloupek, F. J., 182, 222 Cho, C. S., 313,390 Chuchman, R., 3 5 9 , 3 8 5 Ciurdaru, G., 382, 385 Clarke, H. T., 3 0 7 , 3 8 4 Clementi, S., 360, 385 Closs, G. L., 146, 1 7 5 Clowes, G. A., 58, 1 2 2 Coburn, W. C., 323,385 Cochran, W., 324,385 Cohen, L., 8 , 121 Colb, A., 227,265 Collie,.J. N., 363, 364,385 Collins, C. J., 195, 214, 223 Colter,A. K., 149, 151, 1 5 3 , 1 7 5 Combelas, P., 338, 385 Comisarow, M. B., 205, 206, 217, 221, 224 Commeyras, A., 205, 206, 207, 210, 215, 216,217,223 Congdon, W. I., 337,385 Conkling, J. A., 191,223 Conley, H. I.. Jr., 67, 118 Connor T. M., 27 1 , 3 8 8 Cook, D., 338, 349, 350, 365,385 Cook, R. D., 329,387 Cookson, R. F., 316, 349, 350, 385 Coppola, J. C., 64, I21 Cordes, E. H.,4, 32,84, 85,116, 118 Corey, E. J., 184, 188, 223 Cornelisse, J., 227, 230, 235, 236, 237, 238, 244, 245, 246, 249, 250, 253, 261, 264, 265, 266 Correa, A., 333, 391 Corsini, A., 316,385 Costain, C. C., 342, 385 Coulson, C. A., 350, 361, 385 Coustal, S., 315,385 Cox, R. H., 347,385 Craig, D. P., 3 1 6 , 3 8 5 Cramer, F., 59, 118 Crane, J. P., 110, 122 Craze, G.-A., 92, 118 Creighton, D. J., 72,122 Cross, A. D., 339, 385 Cross, J. M., 256, 264 Crucge, F., 315,385 Csizmadia, I. G., 343,387

Cuddy, B. D.. 213,223 Cundall, R. B., 243,264 Cunningham, B. A., 21, 33, 119 D

Dabrowski, J., 3 7 9 , 3 8 8 Dagani, M. J . , 191, 224 Dahlquist, F. W.,82, 1 1 9 Dahm, T. H., 9 0 , 1 0 0 , 1 1 8 Dallinga, G., 289, 385 Daniel, E. S., 333,391 Darwish, D., 20, 122 Das, T. P., 1 3 1 , 1 7 4 Das Sarma, B., 308,385 Dave, K. G., 3 8 1 , 3 9 1 Davis, C. E., 69, 118 Davis, M., 338, 385 Davis, T. L., 307,385 Dayagi, S., 166, I 7 4 De, N. C., 105, 109,119 Dean, R. R., 166,174 de Boer, Th. J., 348, 377,387, 391 Debruin, K. E., 345,385 Degani, I., 365,385 de Gunst, G. P., 236, 238, 254, 256, 260, 264 de Jongh, R. O., 226, 228, 232, 251, 264, 265 Dekkers, J., 69, 118 de Lockerente, S. R., 342, 385 de Loze, C., 338.385 DcMcrnber, J. R., 205, 206, 207, 210, 212, 213,218,219,220,223 den Boer, M. E., 2 4 4 , 2 4 9 , 2 5 0 , 2 6 5 Denes, V. I., 382,385 dcn Heyer, J., 238, 247, 248, 264 Denis, A., 319, 385 Deno, N. C., 2 0 7 , 2 2 3 de Salas, E., 1 7 8 , 2 2 3 Dev, V., 146,175 de Vries, S., 226, 233, 246, 253, 264, 265 Dewar, M. J. S., 124, 173, 206, 224 Deyrup, A. J., 293,387 Dhami, K. S., 153, 173 Dhingra, R. C., 244, 264 Dickcrson, T. A., 4, I19 Dieffenbacher, A., 3 2 0 , 3 8 5 Ditchfield, R., 193, 223 Dixon, W. 'l'. 124, , 173 Dolby, J., 380, 381,385 Dolcnko, A., 376, 391 Donaldson, M. M., 184, 186, 222, 223 Donnclly, M. F., 190, 223 Doran, M. A., 149, 1 5 7 , 1 7 5 Dorovska, V. N., 63, I 2 1 Dorst, W., 226, 265 Doub, L., 336,385

396

AUTHOR lNDEX

Dowd, W., 190,224 Dowdling, J. M., 342,385 Dreyer, D., 378,389 Druzhinina, A. A., 361, 362,383 Dunn,B.,4,91,97, 114,117, 1 1 9 Dunne, K., 369, 370,389 Dupin, M., 346, 384 Dvoryantseva, G. G., 319, 359, 361, 362, 383, 385, 391 Dyer, E., 8 7 , 1 1 9 E Eastham, A. M., 20, 117, 119 Eberson, L., 13, 19, 75, 119, 120 Edward, J. T., 331, 333, 336, 337, 346, 364,384, 385 Edwards, L. J., 73,119 Eigen, M., 43,119, 298, 299, 332,385 El-Anani, A., 360,385 Elander, M., 380,381,385 Elderfield, R. C., 307,385 Elguero, J., 320, 326, 327, 353, 354, 355, 383, 385, 386 Eliel, E. L. 102,119 Ellis, A. W., 377, 384 Ellis, J., 311, 384 Ellis, P. D., 161, 174 Elvidge, J. A., 322,386 Emanuel, R. V., 144,174 Emsley, J. W., 133, 174 Ericsson, L. H.,64, 117 Eriksen, S. P., 25, 26, 1 2 0 Eriksson, S. O., 21, 119 Evans, E. A., 322,386 Evstigneeva, R. P., 358, 359, 390, 391 Eyring, H.,28, 119

F Fairweather, R., 68,117 Fanucci, R., 108,122 Farber, S. Y. 347,386 Farcasan, M., 382,385 Farlow, D. W., 337, 343,383, 386 Farnum, D. G., 123, 127, 128, 135, 137, 148, 149, 150, 151,154,159,167, 173, 174 Farona, M. F., 344, 386 Faulkner, I. J., 20, 121 Favini, G., 315,384 Feageson, E., 23, 118 Fedor, L. R., 1 7 , 1 1 9 Feldman, L. H., 357,386 Felton, S. M., 17, 32, 119 Fernindez Alonso, J. T., 232,264 Fernindez, B., 321,386

Fersht, A. R., 6, 13, 17, 33,60, 73, 74, 7 7 , 79,119,120 Fife, T. H., 2, 4, 18, 22, 23, 29, 32, 37, 39, 40, 44, 47, 49, 55, 56, 58, 61, 63, 85, 86, 87, 92, 93, 96, 100, 103, 104, 105, 108, 109, 111, 113, 114,116, 118, 119, 120 Figgis, B. N., 332,384 Filippova, T. M., 358,389 Fisher, R. D., 190,224 Fleming, K. A., 31 1 , 3 8 4 Fletcher, F. J., 243, 264 Fochi, R., 364,385 Foote, C. S., 186,223 Forkey, D. M., 363,383 Forsythe, P., 315, 351,386 Fort, R. C. Jr., 217,221,224 Foster, D. M., 69, 70, 118 Fraenkel, G., 123, 127, 128, 133, 135, 136, 137, 138, 148, 166, 167, 173, 174, 175, 292,33 1,386 Frampton, R., 3 15,35 1,386 Franconi, C., 33 1,386 Frangopol, P. T., 377,383 Fraser, M., 360, 361, 362,386 Freedman, H. H., 149, 150,175 Freedman, I., 308,390 Friedman, M., 25, 120 Frost, L. N., 4 6 , 1 1 9 Fry, J. L., 186, 192, 213, 223, 224 Fulmer, R. W., 352, 353,388 Fung, H. L., 22,120 Fuson, N., 285,386 G

Gaetjens, E., 17, 36, 51, 79,119 Galbraith, A., 361,386 Garber, R. A., 274, 275,387 Gardner, J. N., 351,386 Garratt, P. J., 137, 138, 174 Garrett, E. R., 73, 1 1 9 Gash, V. W., 352, 353,388 Gates, V., 82, 114, 121 Geis, S. M., 4, 119 Geluk, H.W., 2 13,223 Gerson, F., 310, 312, 313,386 Giauque, W. F., 334,386 Gibson, K. D., 46,119 Giger, W., 274, 386 Gil, R., 327,385 Gil, V. M. S., 166, 174 Gilbert, M., 319,385 Gilchrist, M., 40, 120 Giles, J. M., 125, 177 Gillespie, R. J., 139, 173, 275, 331, 335, 337, 344, 345, 366, 372, 374, 375,384, 386

AUTHOR lNDEX Ginger, R. G., 341, 384 Girdult, G . , 315,385 Giudici, T. A., 89, 1 2 0 Glass, R. S., 184,223 Glasson, W. A., 39, 117 Glaudemans, C. P. J., 87, 119 Gleghorn, J. T., 358,386 Glowds, G. A., 58, 122 Goering, H. L., 186, 223 Goetz, D. W., 193, 201, 222, 223 Goitein, R., 46, 51, 1 2 0 Gold, V., 65, 79, 118, 120, 121, 226, 264, 289,386 Goldacre, R., 307, 314, 315, 316, 319,383 Goldfarb, A. R., 337,386 Gompper, R., 296,338,379,386, 388 Gonzalez, E., 320, 326, 327,386 Cough, T. E., 256,264 Goulden, J. D. S., 284,386 Grant, D. M., 126, 131, 133, 134, 135, 161, 166, 169, 170, 171, 174, 175, 209,224 Grasselli, G., 344, 386 Gray, G. A., 161,174 Griesinger, A., 300, 3 5 4 , 3 8 9 Griffith, J., 247, 264 Grigg, R., 363,384 Grimison, A., 169, 170, 17 1, 173, 322, 326, 382 Grinter, R., 232,264 Grivas, J. C., 302,386 Grob, C. A., 352,386 Grodski, A., 108, 122 Groen, M. B., 2 4 1 , 2 4 2 , 2 6 4 Groenewege, M. P., 347,389 Gronowitz, S., 348,386 Groves, J. T., 137,173 Grunwald, E., 272,387 Grutzner, J. B., 138, 139, 149, 1 5 3 , 1 7 4 Gukovskaya, A. S., 349,383 Gunter, C. R., 3, 29,30, 3 2 , 6 1 , 6 2 , 1 1 7 Gupta, B. D., 124,173 Gurd, R. M., 67, 121 Gusarov, A. V., 37 1,391 Gutowsky, H. S., 272,387 Gutstein, N., 337,386

H Haake, P., 329, 345,387 Hafliger, O., 277, 279, 384 Hafner, K., 378,389 Hagen, E. L., 194, 195,224 Haglid, F., 381,391 Hagopian, L., 103, 108, 11 1 , 1 1 9 Hakka, L. E., 372, 373, 374, 388 Hall, P. L., 20, 122 Hall, R. E., 186, 192,223 Halpern, Y.,140, 152, 161, 163,175

397

Hamilton, G . A., 3 2 , 1 1 7 Hammer, W. J., 185,222 Hammett, L. P., 281, 293,387 Hammond, G. S., 3 0 2 , 3 0 3 , 3 8 9 Hammons, J. H., 214,222 Hantzsch, A., 336, 364, 366,387 Haroy, M. L., 304,392 Harnsbcrger, B. G., 321, 387 Harper, J. J., 186, 222 Harris, D. L., 302, 308,387, 391 Harris,D.O., 11,117 Harris, J. M., 186, 192,223 Harrison, W. F., 363, 3 7 8 , 3 8 3 Hartlcy, B. S., 37, 39, 1 1 7 Hartshorn, M. P., 374,387 Hartshorn, S. R., 190, 224 Hartsuck, J. A., 28,64,120, 121 Hartsuiker, J., 253, 265 Haselbach, E., 310,387 Haseltine, R., 213, 223 Hasselgren, K. H., 380, 381, 385 Hauck, F. P. Jr., 352, 353,388 Haug, A., 101, 122 Hauser, C. F., 14, 118 Hautala, R. R., 235, 246, 265 Havinga, E., 226, 227, 228, 230, 232, 235, 236, 237, 238, 241, 242, 244, 245, 246, 249, 250, 251, 253, 254, 260, 261, 262, 264,265,266 Hawkins, C., 247, 264 Hay, A. S., 352, 353,388 Heck, H. d'A., 34, I1 7 Hegarty, A. F., 4 6 , 1 1 9 Hehre, W. J., 193,223 Heilbronner, E., 232, 264, 310, 312, 378, 386, 389 Heinamaki, K., 345,391 Helmick, I,. S., 326, 389 Helmkamp, G. K., 128, 129,175 Henbest, H. B., 56, 1 2 0 Henderson, R., 56, 120 Herd, A. K., 13, 1 9 , 1 2 0 Hrrn, D. H., 333,389 Herz, A. H., 357,386 Iiewson, K., 323,385 Heyn, H., 146, 175 Higuchi,'I., 13, 18, 2 2 , 1 2 0 Hibcrs, C. W., 367,387 Hinman, R. L., 294, 354, 358, 359, 360, 384, 387.391 Hirt, R. C., 307, 318,387 Hoefnagel, M. A., 310,387 Hoffman, R. A., 348,386 Hoffmann, R., 207,222 Hofmann, K., 6 4 , 1 2 0 Hofstrd, A., 289, 388 Hogeveen, H., 213, 222, 367, 368,377,387 Holah, D. G., 359,385

398

AUTHOR INDEX

Holm, J. C., 272,387 Holst, C., 21, 119 Homer, R. B., 4 4 , 1 1 7 Hopkins, L., 338, 385 Hopkinson, A. C., 343, 370,387 Hoppe, J. I., 6 8 , 1 2 0 Hornung, E. W., 334,386 Hoshi, T., 31 1 , 3 9 1 Hosoya, H., 337, 356, 357, 368, 387, 392 Huang, C. M., 338,387 Huang, L., 213,223 Hubbard, C. D., 3 4 , 3 6 , 6 1 , 1 2 0 Hudec, J., 184,222 Hughes, A. N., 359,385 Hughes, E. D., 181,222 Hui, B. C., 359,385 Humski, K., 191,223 Hunkapiller, M. W., 38, 1 2 0 Hurley, R., 237,246,265 Hurst, G. H., 329,387 Hutchins, J. E. C., 37, 39, 44, 47, 55, 61, 119,120 Hybl, A., 58, 1 2 0 Hyman, H. H., 274,215,387 I

Ikenberry, D., 1 7 5 Imbach, J. L., 320,386 Inagami, T., 3 4 , 1 2 0 Ingold, C. K., 181, 222 Inward, P. W., 3 2 , 1 2 0

J Jackman, L. M., 124, 129, 138, 139, 149, 174 Jacquier, R., 320, 326, 327, 353, 354,355, 383, 386 Jagt, D. L. van der, 185,222 Jaff6, H. H., 309, 310, 311, 346, 351,387, 388, 392 Janssen, M. J., 337, 3 3 8 , 3 8 7 Jao, I,. K., 85,87, 103, 104, 108,119 Jencks, W. P., 2 , 4 , 6 , 19, 23, 28, 29, 30, 31, 32, 33, 36, 40, 41, 65, 89, 116, 117, 118,120,121 Jensen, F. R., 2 1 1, 222 Jeuell, C. L., 145, 154, 174, 1 7 5 Jewett, J. G., 205, 206, 224 Johnson, A. W., 363,384, 3 8 7 Johnson, C. D., 315, 316, 351,384, 386 Johnson, D. M., 345,385 Johnson, L. N., 28, 8 1 , 1 2 0 Johnson, R. M., 226,265 Johnson, S. L., 4, 120 Johnson, W. S., 56, 1 2 0 Jolles, P., 81, 1 2 0

Jones, A. J., 135, 166,174, 213, 223 Jones, D. M., 370,387 Jones, J. R., 322,386 Jones, R. A. Y., 340, 347,387 Jones, R. N., 330, 336,387 Jordan, F., 3 2 4 , 3 8 7 Josien, M.-L., 285,386 K Kahovec, L., 338,387 Kaiser, A., 352,386 Kaiser, B. L., 4, 64, 80, 1 2 0 Kaiser, E. T., 4, 64, 80, 1 2 0 Kaiser, R. S., 310, 390 Kang, E. P.. 6 5 , 1 2 0 Kaper, L., 377,387 Kaplan, F., 37 1 , 3 8 7 Kapoor, K. L., 319,387 Karau, W.,56, 1 2 2 Karplus, M., 131, 132, 1 7 4 Kashima, Ch., 381,391 Kato, H., 315, 388 Katritzky, A. R., 294, 315, 316, 319, 334, 340, 347, 348, 349, 351, 360, 377,383, 384, 385,386,387, 388, 392 Katz,T. J., 137, 138,174 Kawakami, J. H., 185,222 Kay, I. T., 3 6 3 , 3 8 7 Kazanskaya, L. V., 358,389 Keana, J. J., 68,117 Kehlen, H., 329,390 Kehlen, M., 329,390 Keizer, V. G., 213,223 Kelly, D. P., 145, 146, 154, 174. 351, 389 Kendall, J., 366,388 Kennard, C. H. L., 323,390 Kessick, M. A., 190, 224 Kessler, H., 304, 306,388 KPzdy, F. J., 3, 29, 30, 32, 34, 54, 61, 62, 117 Kholodov, L. E., 363,388 Kim, J. P., 292,386 Kim, Y., 24,122 Kirby, A. J., 6, 13, 17, 23, 73, 74, 77, 79, 92,116,117,118,119,120 Kirsanov, A. V., 312,392 Kirsch, J. F., SO, 34, 36, 61, 120 Kisilenko, A. A., 364, 365,391 Kistiakowsky, G. B., 214,222 Klages, F., 340,388 Kleinfelter, D. C., 221,223 Klingsberg, E., 355,391 Kluetz, M. D., 34,121 Knollmiiller, K., 338,387 Kobayashi, S., 166,174 Kobayashi, T., 31 1 , 3 9 1 Koch, M. J. 8 7 , 1 1 9 Kochergin, P. M., 361,362, 383

AUTHOR INDEX

399

Layne, W. S., 346,388 Kochevar, I. H., 108,122 Koehler, K., 85,118 Lazzeretti, P., 126, 174 Koenig, D. F., 28, 81,117 Ledwith, A., 370,388 Kohler, H., 356,388 Lee, C. C., 191, 214, 215,223, 224 Leibfritz, D., 304, 306, 388 Koizumi, T., 345,387 Koltun, W. L., 67,121 Leisrowitz, L., 338,387 Konishi, H., 315,388 Leonard, N. J., 352,353, 380,388 Konishi, K., 123, 175 Letsinger, R. L., 226, 227, 235, 236, 237, Koren, J. G., 307,387 246, 247,253,265 Korolev, B. A., 31 1,388 Levisalles, J., 372, 383 Kosanovic, D., 366, 391 Levy, G. C., 124,174 Koshland, D. E. Jr., 8, 9 , 10, 11, 12, 28, Lewis, G. E., 310,388 120,122 Lewis, R. G., 381,391 Kostyuchenko, N. P., 359,391 Liang, G., 145, 146, 156, 163, 174, 1 7 5 Kouba, J. E., 374,388 201,224 Kozerski, L., 378,388 Liang, Y. T., 26,120 Kramer,C. M., 195, 197, 199,223 Lichtenberg, D., 324,390 Kramer, H. E. A., 274,299,379,388 Liedek, E., 319,390 Kraut, J., 56, 121, 122 Lienhard, C. E., 90, 122 Kregar, I., 8 2 , 1 1 6 Lietzke, M. H., 214,223 Kresge, A. J., 371, 372, 373, 374, 375, 388 Liler, M., 285, 286, 292, 310, 328, 329, Kroll, H., 66,120 330, 331, 332, 333, 334, 335, 336, 337, Kronenberg, M. E., 226, 237, 251, 262, 339, 342, 344, 353, 364, 366, 368,388, 265, 266 391 Ku, A. T., 343,344,367,368,369,389 Lincoln, D. N., 207, 223 Kuebler, N. A., 342,383 Lippmaa, E. T., 209,224 Kukhar, V. P., 312,392 Lipscomb, W. N., 28, 64, 120, 121 Kunzler, J. E., 334,386 Liu, J. S., 207, 223 Kupchan, S. M., 25,26,56, 120.121 Liu, Y., 219,223 Kurland, R. J., 149, 151, 153, 175, 342, Lloyd, D., 321,388 388 Lockhart, J. C., 305, 307,384 Kustanovich, I. M., 358,389 Lodder, G., 246, 263, 265 Kutzelnigg, W., 302,389 Loewenstein, A., 271, 272, 331, 332, 384, Kwok, W. K., 24,122 386, 387, 388 Lohrmann, R., 319,390 Lok, C. M., 227, 235, 237, 244, 246, 249, L 250,265, 266 Long, F. A., 327, 378,389, 390 Lagowski, J. M., 349,387 Longuet-Higgins, H. C., 36 1 , 38.5 Laiho, G. E., 108,122 Lovcll, B. J., 56, 120 LaLancette, E. A., 137,174 Lowe, G., 28, 81, 121 Lam, L. K. M., 186, 192, 213, 215, 223, L o w e , J . V. Jr., 305, 390 224 Lowry, T. M., 20, 121 Lamdan, S., 321,386 Lucas, E. C., 33, 121 Lammers, J. G., 233, 236, 238, 245, 246, Ludwig, M. L., 64, 121 248,264, 265 Lugtenburg, J,, 236, 246, 264, 265 Lancast-er, P. W., 77, 116 Lui, C. Y., 205, 206, 207, 210, 212, 213, Lancelot, C. J., 186, 192,223 215,216, 217,218,219, 220,223 Lang, J., 294,360,387 Lumry, R., 28, 63,119, 121 Lantzke, I. R., 31 1,384 Lutz, K., 301, 322, 390 Lapinski, R., 39,118 Larson, B., 101, 122 Lascombe, J., 315,385 M Latremouille, G. A., 20, 119 Laughlin, R. G., 345,388 Maas, W., 352,391 Lauterbur, P. C., 136, 160, 174, 175, 209, McAllister, C., 70, 80, I 1 7 222 McCain, J. H., 235, 246, 265 Lawlor, J. M., 23, 24, 92, 117, 138, 149, McConnell, H., 169,176 174 McDanicl, D. H., 277, 279,384

400

AUTHOR INDEX

McDowell, S. T., 51, 118 McFarland, J. T., 219,223 Mach, C. W.,293,383 Maciel, G. E., 161, 174 McKeever, L. D., 157, 174 McKenzie, S., 361,386, 389 McKervey, M. A., 213,223 kiackie, R. K., 321,388 llackor, E. L., 136, 142, 174, 289, 367, 372, 373,384, 385,387,388 McLachlan, A. D., 128, 136, 137, 138,174 McLean, C., 123, 124, 136, 142, 174, 175, 367,372,373,384,387,388 McMahon, D. M. 3 2 , 3 7 , 1 1 9 McNab, H., 321,388 McRitchie, D. D., 3 13, 384 McRowe, A. W.,214,222 Mahan, J. E., 352, 353,382 Maikkula, M., 345, 360, 391 Mair, C. A., 28,81, 117 Majerski, A., 213, 223 Malajcid, R., 191, 223 Malera, A., 360,386 Mandell, I,., 345, 389 Mannschreck, A., 322,326,388, 390 Marchessault, R. H., 87, 119 Marimoto, G., 31 7,389 Marquardt, F. H., 42, 118 Marshall, D. R., 321,388 Martell, A. E., 66, 121 Martin, R. B., 67, 118, 332,389 Martinek, K., 63, 121 Marzilly, L. C., 68, 118 Marzin, C., 327,386 Mason, S. F., 297, 315,389 Mataga, N., 356,389 Matccscu, G. D., 124, 137, 139, 140, 174 199,201,209, 220,224 Matschiner, H., 329,390 Maugh,T., 21, 79,121 Meacock, S. C. R., 336, 385 Mecke, R., 302,389 Medvetskaya, I. M., 275,390 Meerwein, H., 178,223 Mehta, G., 137,174 Memory, J. D., 126, 130, 174 Meiboom, S., 272, 331, 332, 384, 386. 387 Mele, A., 337,386 Melent’eva, T. A., 358,389 Meloy, C. K., 371,387 MPly, B., 291, 342,389 Menger, F. M., 25,121, 345,389 Meriwether, L., 68, 121 Meuche, D., 378,389 Michelot, R., 299, 300, 354,382 Milakofsky, L., 190,224 Milne, D. G., 243, 264 Milstien, J. B., 32, 63, 119, 121

Milstien, S., 8, 121 Mitra, S. S., 358,382 Mo, Y. K., 140, 152, 156, 161, 163, 175, 369,370,374,375,389 Molloy, B. B., 360, 361,386, 389 Montgomery, J. A., 323,385 Moodie, R. B., 337,343,383, 386 Morawetz, H., 17, 36, 51, 78, 79, 119, 121 Moreland, C. C., 359, 390 Moriconi, E. J., 314,389 Moyer, A. N., 365,392 Muirhead, H., 64,121 Mulder, J. J. C., 232, 265 Mum, 8.L., 190, 191,223 Murrell, J. N., 166, 174 Musher, J. L., 128, 129, 137, 142, 174 Muszynska, G., 8 0 , 1 2 1 N Nagakura, S., 337, 356, 357, 368,387, 392 Nagy, 0. B., 342,385 Nakamura, K., 34, 36, 11 7 Nakasawa, Y.,121 Narayanan, C. R., 56, 120 Nath, R. L., 9 7 , 1 2 1 Nayak, B., 332,384 Neiman, Z., 324,390 Nelson, G. L., 124, 174 Nenitzescu, C. D., 364, 377,383, 389 Ncuman, R. C., 302,303,389 Neurath, H., 64, 117, 122 Nevell, T. P., 178, 223 Neveu, M. C., 75, 76,117 Nickon, A., 219,223 Niemann, C., 331,386 Nijhoff, D. F., 246,265 Nilsson, J. L. G., 380, 381, 385 Nilsson, S., 227, 247, 265 Nist, B. J., 139,175 North, A. C.T., 28,81,117

0 Oakenfull, D. G., 26, 41, 79, 120. 121, 122 O’Brien, C., 322,386 O’Brien, D. H., 275, 367, 368,389 O’Connor, C. J., 337,383 Oddo, G. 365,366,389 Oelderik, J. M., 194, 222 Oishi, T., 38 1,391 Okuyama, T., 4 3 , 4 4 , 1 2 1 Olah,G. A., 123, 137, 139, 140, 144, 145, 146, 149, 152, 154, 156, 157, 158, 162, 163, 173, 174, 175, 199,201,203,204, 205, 209, 210, 211, 212,213, 215, 216, 217, 219, 220, 221,223, 224, 275,304, 335, 343, 344, 351,366, 367,369,370, 374,375,389

401

AUTHOR INDEX Olah, J. A., 344,368,369,389 O'Leary, M. H., 34,121 Opitz, C., 300, 354,389 Oreskes, I., 78, 121 Ortiz, J. J., 85, 118 Osawa, T., 121 Ostwald, W., 277,389 Oth, J. F. M., 123,175 Ottenheym, J. H., 347,389 Ouicho, F. A., 64,121 Ovchinikova, M. M., 371,391 Overman, L. E., 7 1, 117

P Padilla, A. G., 345,385 Page, M. I., 6, 19, 121 Palei, R. M., 362,383 Pal'm, V. A., 307,389 Pandit, U. K., 8, 9, 18,36, 49, 51, 118 Pariser, R., 250, 265 Pkkinyi, C., 235,265 Parker, L., 60, 122 Parker, V. D., 373,391 Parrino, V. A., 308,390 Parry, E. P., 333,389 Parsons, S. M., 82,121 Patchornik, A., 34,120, 121 Paudler, W. W., 124,175,325,326,389 Paul, E. G., 209,224 Paul, M. A., 372,389 Pauling, L., 280, 288, 289, 290, 328,389 Paziomek, E. J., 265 Pearson, R. G., 296, 307,389 Peel, T. E., 275,386 Pekhk, T. I., 209,224 Perillo, I., 321, 386 Perkampus, H.-H., 288,389 Perrin, D. D., 301,390 Perrin, R., 374, 375,384 Persianova, I. V., 361, 362, 363, 383, 388 Person, L. S., 307, 390 Peters, E., 206,222 Petersen, W. C., 236, 253, 265 Petrashenko, A. A., 312,392 Pfeiffer, G. V., 205, 206, 224 Pfleiderer, W.,319, 325, 350,383, 390 Philipp, M., 34, 60,121 Phillips, D. C., 28,81, 82, 117, 120, 121 Phillips, J., 307, 314, 315, 316, 319,383 Piszkiewicz, D., 86, 95, 100, 105, 106,118. 121 Pitman, I. H., 22,120 Plate, A. F., 209,224 Pletcher, T. C., 85, 118 Pocker, Y.,20,121 Polgar, L., 38, 121 Pollack, R. M., 3 4 , 6 0 , 1 2 1

Poole, J. A., 244,264 Pople, J. A., 132, 174, 175, 193, 206,223, 224 271, 3 9 0 Porter, G., 345, 265 Porter, R. D., 145, 146, 154, 156,174, 175, 203, 204, 212, 219, 220, 223, 344,389 Powers, J. C., 381,390 Pracejus, H., 329, 339, 340,390 Preobrazhenskii, N. A., 358,390 Price, E., 305, 307,390 Prokop'ev, B. V., 371,391 Prokop'eva, T. M., 349,383 Rue, J. E., 68,120 Pugmire, R. J., 126, 130, 131, 133, 134, 135, 161, 169, 170, 171, 175 Pullman, A., 291,342,389 Pullman, B., 324,391

Q Quin, L. D., 359,390

R Raher, D. J., 186, 192,213,223, 224 Radom, L., 193,206,224 Raftery,M. A., 51, 81, 82, 107, 118. 1 2 1 Rajagopalan, S., 56,120 Rajender, S., 63, 121 Ramsey, B. G., 344, 373,386, 3 9 0 Ramsey, 0. B., 226,246,265 Randal:, E. W., 144,174 Rand-Meier,T.. 81, 82, 107, 119. 121 Ranganayakulu, K., 213, 220,223, 224 Rao, C. N. R., 325,390 Rao. K. G.. 325.390 Rapp, M. W., 190,224 Raquena, Y., 33,119 Ray,G. J., 149, 151, 152, 153, 175 Reavill, R. E., 347,388 Reeke, G. N. Jr., 64,121 Rees, C. W., 226, 265 Reeves, L. W. 271,390 Reeves, R. L. 309, 310, 31 1 , 3 9 0 Regan, T. Xi.,357,386 Reichman, U., 324,390 Reid, D. H., 360, 361,386, 389 Remanick, A., 188,222 Renk, E., 352,386 Richards, J. H., 38, 120, 128. 136, 137, 138, 174 Richards, K. E., 374,387 Richardson, D. I. Jr., 26, 121, 122 Ric'hey, H. G., 207,223 Ricketts, J. A., 312,390 Ridgewell, B. J., 316,384 Riebsomer, J. L., 321,387

402

AUTHOR INDEX

Riemcnschneider, J. L., 199, 201, 220, 224 Riley, T., 79,120, 121 Riordan, J. F., 80,121 Roberts, J. D., 214,224 Robertus, J. D., 56,121 Robin, M. B., 342,383 Robins, M. J., 135, 169,175 Robins, R. K., 126, 130, 131, 134, 135, 169,175 Robinson, S. D., 338,390 Rochester, C. H., 226, 264, 292,390 Rodriguez, G., 169, 170, 171, 173, 322, 326,382 Rogers, G. A., 121 Rogers, M. T., 330,390 Rogers, N. A. J., 37 1 , 3 9 0 Rony, P. R., 20,121 Rordam, H. N. K., 364,390 Ros, P., 367, 387 Rosenfeld, J., 194, 195,224 Rosenheck, K.,337,384 Roth, B., 317, 318,390 Rothberg, I., 185, 186, 222 Rubin, T. R., 334,386 Rukwied, M., 319,390 Rumpf, P., 3 15,385 Rundle, R. E., 58, 120 Rupley, J. A., 82, 114,116, 121 Rusakowicz, R., 240,265 Russel, J. G . , 135, 166, 174 Ryabova, R. S., 275,390 Ryan, G., 139,175 Rydon, H. N., 9 7 , 1 2 1 Rynbrandt, D. J., 108,122 S

Saenger, W.,59, 118 Safta, M., 318,390 Saika, A., 272,390 Salomaa, P., 108, 121, 122 Sandel, V. R., 149, 150,175,232,266 Sankey, G. H., 90, 100, 118 Sargent, F. P., 256,264 Sargent,G. D., 184, 189,218,221,224 Sargeson, A. M., 68,69, 70, 118 Sarma, V., 28,81, 117 Sattar, A., 37 1 , 3 9 0 Saunders, M. L., 194, 195, 212,224 Saunders, W. H., 214,224 Scandola, E., 365,389 Schachtschneider, J. H., 367,387 Schaefer, T., 137, 138,175 Schaeffer, J. P., 191,224 Scheibe, A., 356,388 Scheppele, S. E., 190,224 Scheraga, H. A., 46,119, 290,390 Schewene, C. B., 186,223

Schlatmann, J. L. M. A., 213, 223 Schleyer, P. von R., 123, 174, 175, 183, 184, 186, 192, 193,206, 212,213, 217, 221,222, 223,224 Schmidt, G. M.J., 338,387 Schmir, G. L., 21, 30, 33, 43,44,118, 119, 121 Schmitt, R. G., 307, 318,387 Schneider, W. G., 132, 136, 137, 138, 175, 271,390 Schreiber, K. C., 181, 224 Schroeder, G., 125, 1 7 7 Schulze, J., 378,390 Schwarzenbach, G . , 301, 322,390 Schweizer, M. P., 128, 129, 175 Schwert, G. W., 64,122 Sebastian, J. F., 58,122 Seccombe, R. C., 323,390 Secemski, I. I., 90,122 Seitz, W.,322,388 Shafer, J. A., 43, 52, 116 Shapiro, R. H., 253,265 Shapiro, S. L., 308,390 Sharon, N., 83, 104,118 Sheinker, Yu, N., 319, 359, 361, 362, 385, 391 Shelton, E. M., 285, 386 Shelton, G., 363,384 Sheppard,G., 2 8 , 8 1 , 1 2 1 Shiner, V. J. Jr., 190,224 Shoffner, J., 348,384 Short, L. N., 316,385 Shuezhko, L. M., 364, 365,391 Sigman, D. S., 72, 122 Silberman, R. G., 137, 174 Simalty, M., 377,391 Simon, W., 274,386 Sinnott, M. L., 28, 81, 121 Sklyar, Y.E., 358, 359,390, 391 Slade, P., 56, 121 Small, T., 36 1, 386 Smallcombe, S. H., 38,120 Smidrod, O., 101, 122 Smidt, J., 347,389 Smit, P. J., 289,385 Smith, C. H., 3 9 0 Smith, C. R., 338, 340, 341,390 Smith, J. F., 181,222 Smith, J. H., 25,121 Smith, M. C., 83,90, 100,118 Smith, P. J., 372,374, 375,384 Smith, S., 20, 122 Smith, W. F., 309,390 Snell, R. L., 24, 122 Snieckus, V., 355,391 Snoke, J. E.,64,122 Sodervall, T., 346,391 Somasekhara, S., 226,266 S o m m a , J. M., 372,383

403

AUTHOR INDEX Sorensen, T. S., 213, 220,223, 224 Sostman, H.D., 324,387 Spatz, H.-Ch., 59, 118 Speck, J. C., 108,122 Spel, T., 238,264 Speziale, A. J., 380,383 Spiesecke, H., 132, 136, 137, 175 Spikes, J. D., 28,119 Spinner, E., 33 1, 348,390 Spunta, G., 364,385 Staab, H. A., 322,326,388,390 Stamhuis, E. J., 352.391 Stang, P. J., 183,224 Steitz, T. A., 64,121 Stejskal, E. O., 332,391 Steller, K. E., 237, 265 Sterin, K. E., 216, 222 Sternhell, S., 124, 129, 139,174 Stevens, J. B., 294, 33 1, 333, 334,392 Stevens, R. V., 381,391 Stewart, R., 284, 331, 333, 336, 366, 385, 391 Stewart, R. F., 193,223 Stock, L. M., 283,391 Stollar, H.,346,385 Storm, C. B., 65,120 Storm, D. R.,8, 10, 11, 12,122 Stothers, J. B., 124, 127, 140, 141, 153, 157,160,173,175 Strachan, W. M. J., 376, 391 Stransky, 2 , 3 13,391 Stratenna, J. L., 231, 265 Strauss, H. L., 307,387 Streitwieser, A. Jr., 288,391 Strelitz, J. Z., 317, 318,390 Strong, A. B., 175 Struchkova, M. I., 359,391 Strzelecka, H.,376, 377,391 Sturtevant, J. M., 30, 31,36,118 Stuzka, V., 313,391 Su, S. C. K., 43,52,116 Sucio, N., 351,389 Sugiyama, N., 381, 391 Sukhorukov, B. I., 349,383 Sundaralingham, M., 324,391 Sunko, D. E., 191,223 Suppan, P., 245,265 Sustmann, R., 206,224 Svanholm, U., 373,391 Swain, C. G., 20,122 Sweeting, L. M., 295, 334,391 Szilagyi, P. J., 335, 369,370,389

T Taddei, F., 126,174 Taft, R. W. Jr., 57,122, 124,175

Taguchi, H., 316,383 Takahashi, K., 123,175 Takaki, M., 123,175 Takechi, H., 22,120 Takeda, M., 332,391 Takimoto, M., 307, 318,391 Tamminga, J. J., 238,265 Tanaka, J., 337,387 Tanizaki, Y ., 3 1 1,391 Tanner, D. W., 55,118 Tarhan, H.O., 3 16, 348,384 Tarrago, G., 353, 354,385 Taurins, A., 302,386 Tchoubar, B., 299, 300, 354,382 Teklu, Y. D., 307,390 Terashima, M., 381, 391 Testa, A. C., 237,240, 246,265 Thacker, D., 104, 108, 118 Thanassi, J. W., 17, 75, 122 Theodoropoulos, S., 358,384 Thomas, M. T., 355,391 Thompson, J. A., 124,175 Thornson, A., 227,264 Thorpe, M. C., 323,385 Tickle, T., 363, 364,385 Tillett, J. G., 351,383 Timberlake, J.-W., 124, 175 Tischenkova, I. F., 363,388 Titova, S. P., 3 1 1,388 Tizani, D., 326, 355,383, 386 Tokuhiro, T., 133, 168,175 Tolmachev, A. I., 364, 365,391 Tomer, K., 253, 265 Tompa, A. S., 305,307,390 Townsend, L. B., 126, 130, 131, 134, 135, 169,175 Traficante, D. D., 161,174 Trifan, D. S., 180, 181, 191, 224 Trifonov, B. A., 37 1,391 Trotman-Dickenson, A. F., 291, 315,391 Tso’o, P. 0. P., 128, 129, 175 Turk, V.,82,116 Turnbull, N. H., 316,391 Turner, A,, 45, 46, 79, 118 Turner, J. C., 322,386 Turner, J. O., 207,223 Turnquest, B. W., 30, 66, 67.11 7 Tutundiic, P. S., 366, 391 Tweedie, J. F., 313, 384 Tye, F. L., 289,386

U

Ul’yanova, T. N., 319,385 Usher. D. A.. 26.121. 122 Ushio; M., 123, 175

404

AUTHOR INDEX

v Vandenbelt, J. M., 336,385 van der Haak, P. J., 384,391 van der Heyden, A., 226, 265 vanderKooij, J., 123, 124, 175 van der Stegen, G. H. D., 242,265 van der Waals, J. H., 289,388 van Doorn, J. A., 335,384 van Emster, K., 178, 223 Van Etten, R. L., 58, 5 9 , 1 2 2 van Raayen, W., 347,389 Van Riel, H. C. H. A., 263,265 van Veen, A., 310,386, 387 van Vliet, A., 230, 237, 266 Vanyukhin, G. N., 371,391 Varfolomeyev, S. D., 63,121 Vatakencherry, P. A., 188,223 Vaughan, J., 374,387 Veerkamp, Th., A., 347,389 Veillaro, A., 324, 391 Velthurst, N. H., 123, 124, 175 Verheydt, P. L., 235,249,266 Vernon, C. A., 82, 122 Verrijn Stuart, A. A., 289,385 Villiger, V., 366, 383 Vink, J. A. J., 227, 235, 236, 237, 249, 253.266 Vinnik,.M. I., 275,390 Virtanen, P. I. O., 345,391 Vogel, P., 195, 224 Vold, R. L., 333,391 von Philipsborn, W., 316, 320. 324, 325, 348,349,385,391

W Waack, R., 149, 157,174, 1 7 5 Wagner, R., 316, 324, 325, 348, 349,391 Wahl, G. M., 72, 122 Wai, H., 372 392 Walsh, K. A., 64, 117 Wamhoff, H., 381,391 Wang, J. H., 60, 122 Warburton, W. K., 303, 307,383 Ward, R. L., 256,266 Waring, A. J., 319, 349,388 Watts, W. E., 185, 186, 217, 221, 224 Wedlar, F. C., 37,117 Weeks, D. P., 108, 110, 122 Wein, M., 69,118 Weinberg, D. S., 191,224 Wellman, K. M., 302, 308,387, 391 Wenkert, E., 381,391 Wepstrr, B. M., 280, 310,386, 391 Westerman, P. W., 154, 158, 173, 175 Westheimer, F. H., 4, 6, 68, 121, 122 Whipple, E. B., 358,359,384, 387, 391 Whitaker, D. R., 38,120

White, A. M., 144, 145, 154, 161, 162, 174, 175, 203, 204, 205, 209, 210, 211, 220, 223, 224, 275, 304, 316, 335, 343, 366, 367,369,384, 389 White, M., 368,389 Wiberg, K. B., 139,175, 287,391 Wiegerink, F. J., 248, 266 Wiley, R. H., 365,392 Williams, A., 28, 60, 61, 81,121, 1 2 2 Williams, D. E., 58, 1 2 0 Williams, G., 353, 392 Williams, J. E., 206, 224 Williams, R. L., 304,306, 346,384 Wilson, C. L., 178, 223 Wilson, E. B. Jr., 342, 388 Wilson, N. K., 170, 1 7 7 Wilson, R., 256,264 Winstein, S., 20, 122, 180, 181, 191, 224 Winter, R., 188,223 Wold, A., 213,223 Wong, E. W. C., 191,223 Wood, H. J., 370,388 Wood, N. F., 370,387 Woodbrey, J. C., 330,390 Woolsey, G. B., 338,384 Wright. C. S.. 3. 122 Wright; G.J.1 3$4,387 Wu, C. Y., 364, 372,383 Wyatt, P. A. H., 293, 392 Wynberg, H., 352, 391 Wynne-Jones, W. F. K., 83, 117

Y Yakhontov, L. N., 319,385 Yamamoto, M., 381, 391 Yamaoka, T., 356,357,392 Yaselman, M. E., 371,391 Yashunskii, V. G., 363,388 Yates, K., 284, 294, 295, 331, 333, 334, 336, 340, 341, 366, 372,385, 390, 391, 392 Yeh, S. J., 309, 310, 31 1 , 3 9 2 Yemelyanov, I. S., 37 1 , 3 9 1 Yonezava, T., 315,388 York, L., 40,118 York, S. S., 34, 120, 121 Young, R. J., 5 6 , 1 2 1 Yurchenko, R. I., 312,392 2

Zachau, H. F., 56, 122 Zange, E., 340,388 Zeffren, E., 20, 122 Zerner, B., 54, 117 Zhmurova, I. N., 312,392 Zimmer, H., 346, 388 Zimmerman, H. E., 226, 232,266 Zolotareva, L. A., 312,392

Cumulative Index to Authors

Anbat, M., 7, 11 5 Bell, R. P., 4, 1 Bennett, J. E., 8, 1 Bentley, T. W.,8, 151 Bethell, D., 7, 153; 10, 53 Brand, J. C. D., 1, 365 Brinkman, M. R., 10, 5 3 Brown, H. C., 1, 35 Cabell-Whiting, P. W., 10, 129 Cacace, F., 8, 79 Carter, R. E., 10, 1 Collins, C. J., 2, 1 Cornelisse, J., 11, 225 Crampton, M. R., 7, 21 1 de Gunst, G. P., 11, 225 Farnuin, D. C.,11, 123 Fendler, E. .J., 8, 271 Fendler, J. H., 8, 2 7 1 Ferguson, ti., 1, 203 Fields, E. K., 6, 1 Fife, T. H., 11, 1 Fleischmann, M., 10, 155 Frey, H. M., 4, 147 Gilbert, B. C., 5, 53 Gillespie, R. J., 9, I Gold, V., 7, 259 Greenwood, H. H., 4, 73 Havinga, E., 11, 225 Hogeveen, H., 10, 29, 129 Johnson, S. L., 5 , 2 3 7 Johnstone, R. A. W.,8, 151 Kohnstam, G., 5, 121 Krarner, G. M., 11, 177 Kreevoy, M. M:, 6 , 6 3 Liler, M., 11, 26 7 Long, F. A., 1, 1 Maccoll, A., 3, 91 McWeeny, R., 4, 73 Melander, L., 10, 1 Mile, B., 8 , 1 Miller, S. I., 6, 185 Modena, G., 9, 185 405

406 More O'Ferrall, R. A., 5, 331 Norman, R. 0. C., 5, 53 Olah, G. A., 4 , 3 0 5 Parker, A. J., 5, 173 Peel, T. E., 9, 1 Perkampus, H. H., 4, 195 Pittmann, C. U., Jr., 4 , 305 Pletcher, D., 10, 155 Ramirez, F., 9, 25 Rappoport, Z., 7, 1 Reevts, L. W., 3, 187 Robertson, J. M., 1, 203 Samuel, D., 3, 123 Schaleger, L. L., 1, 1 Scheraga, H. A., 6, 103 Shatenshtein, A. I., 1 , 156 Silver, B. L., 3, 123 Simonyi, M., 9, 127 Stock, L. M., 1, 35 Symons, M. C. R., 1, 284 Thomas, A., 8, 1 Tonellato, U., 9, 185 Tudos, F., 9, 127 Turner, D. W., 4, 31 Ugi, I., 9, 25 Ward, B., 8, 1 Whalley, E., 2 , 9 3 Williams, J. M., Jr., 6, 63 Williamson, D. G., 1, 365 Wolf, A. P., 2, 201 Zollinger, H., 2, 163 Zuman, P., 5, 1

INDEX

Cumulative Index of Titles

Abstraction, hydrogen atom, from 0 - H bonds, 9, 127 Acid solutions, strong, spectroscopic observation of alkylcarbonium ions in, 4, 305 Acids, reactions of aliphatic diazo compounds with, 5, 331 Activation, entropies of, and mechanisms of reactions in solution, 1, 1 Activation, heat capacities of, and their uses in mechanistic studies, 5, 12 1 Activation, volumes of, use for determining reaction mechanisms, 2,93 Aliphatic diazo compounds, reactions with acids, 5, 331 Alkylcarbonium ions, spectroscopic observation in strong acid solutions, 4, 305 Ambident conjugated systems, alternative protonation sites in, 11, 267 Ammonia, liquid, isotope exchange reactions of organic compounds in, 1, 156 Aromatic photosubstitution, nucleophilic, 11, 225 Aromatic substitution, a quantitative treatment of directive effects in, 1, 35 Aromatic substitution reactions, hydrogen isotope effects in, 2, 163 Aromatic systems, planar and non-planar, 1, 203 Arynes, mechanisms of formation and reactions at high temperatures, 6, 1 A - S E ~reactions, developments in the study of, 6, 63 Base catalysis, general, of ester hydrolysis and related reactions, 5, 237 Basicity of unsaturated compounds, 4, 195 Bimolecular substitution reactions in protic and dipolar aprotic solvents, 5, 173 Carbene chemistry, structure and mechanism in, 7, 163 Carbon atoms, energetic, reactions with organic compounds, 3,201 Carbon monoxide, reactivity of carbonium ions towards, 10, 29 Carbonium ions (alkyl), spectroscopic observation in strong acid solutions, 4, 305 Carbonium ions, gaseous, from the decay of tritiated molecules, 8, 79 Carbonium ions, photochemistry of, 10, 129 Carbonium ions, reactivity towards carbon monoxide, 10, 29 Carbonyl compounds, reversible hydration of, 4, 1 Catalysis, enzymatic, physical organic model systems and the problem of, 11, 1 Catalysis, general base and nucleophilic, of ester hydrolysis and related reactions, 5, 237 Catalysis, micellar, in organic reactions: kinetic and mechanistic implications, 8, 27 1 Cations, vinyl, 9, 185 Charge density-N.M.R. chemical shift correlations in organic ions, 11, 125 Chemically induced dynamic nuclear spin polarization and its applications, 10, 53 CIDNP and its applications, 10, 53 Conformations of polypeptides, calculations of, 6, 103 Conjugated molecules, reactivity indices in, 4, 73 Diazo compounds, aliphatic reactions with acids, 5, 331 Dipolar aprotic and protic solvents, rates of bimolecular substitution reactions in, 5, 173 Directive effects in aromatic substitution, a quantitative treatment of, 1, 35 40 7

40 8

CUMULATIVE INDEX

Electrode processes, physical parameters for the control of, 10, 155 Electron spin resonance, identification of organic free radicals by, 1, 284 Electron spin resonance studies of short-lived organic radicals, 5, 5 3 Electronically excited molecules, structure of, 1, 3 6 5 Energetic tritium and carbon atoms, reactions of, with organic compounds, 2, 201 Entropies of activation and mechanisms of reactions in solution, 1, 1 Enzymatic catalysis, physical organic model systems and the problem of, 1 1 , 1 Equilibrium constants, N.M.R. measurements of, as a function of temperatures, 3, 187 Ester hydrolysis, general base and nucleophilic catalysis, 5, 237 Exchange reactions, hydrogen isotope, of organic compounds in liquid ammonia, 1 , 156 Exchange reactions, oxygen isotope, of organic compounds, 3 , 1 2 3 Excited molecules, structure of electronically, 1, 365 Free radicals, identification by electron spin resonance, 1, 284 Free radicals and their reactions at low temperature using a rotating cryostat, study of, 8, 1 Gaseous carbonium ions from the decay of tritiated molecules, 8, 79 Gas-phase heterolysis, 3, 9 1 Gas-phase pyrolysis of small-ring hydrocarbons, 4, 147 General base and nucleophilic catalysis of ester hydrolysis and related reactions, 5 , 237

HzO-DzO Mixtures, protolytic processes in, 7, 259 Heat capacities of activation and their uses in mechanistic studies, 5, 121 Heterolysis, gas-phase, 3, 9 1 Hydrated electrons reactions of, with organic compounds, 7, 115 Hydration, reversible, of carbonyl compounds, 4, 1 Hydrocarbons, small-ring, gas-phase pyrolysis of, 4, 147 Hydrogen atom abstraction from 0 - H bonds, 9,127 Hydrogen isotope effects in aromatic substitution reactions, 2 , 163 Hydrogen isotope exchange reactions of organic compounds in liquid ammonia, 1 , 1 5 6 Hydrolysis, ester, and related reactions, general base and nucleophilic catalysis of, 5, 237 Ionization potentials, 4, 31 Ions, organic, charge density-N.M.R. chemical shift correlations, 11, 1 2 5 Isomerization, permutational, of pentavalent phosphorus compounds, 9, 25 Isotope effects, steric, experiments on the nature of, 10, 1 Isotope effects, hydrogen, in aromatic substitution reactions, 2 , 1 6 3 lsotope exchange reactions, hydrogen, of organic compounds in liquid ammonia, 1, 150 lsotope exchange reactions, oxygen, of organic compounds, 3, 1 2 3 Isotope and organic reaction mechanisms, 2, 1 Kinetics, reaction, polarography and, 5 , 1 Mass spectrometry, mechanism and structure in: a comparison with other chemical processes, 8, 152 Mechanism and structure in carbene chemistry, 7, 1 5 3 Mechanism and structure in mass spectrometry: A comparison with other chemical processes, 8 , 152 Mechanisms, organic reaction, isotopes and, 2, 1 Mechanisms, reaction, use of volumes of activation for determining, 2 , 9 3 Mechanisms of formation and reaction of arynes at high temperatures, 6, 1 Mechanisms of reactions in solution, entropies of activation and, 1 , 1 Mechanistic studies, heat capacities of activation and their uses in, 5, 121 Meisenheimer complexes, 7, 2 1 1 Micellar catalysis in organic reactions: kinetic and mechanistic implications, 8, 27 1 N.M.R. chemical shift-charge density correlations, 11, 125 N.M.R. measurements of reaction velocities and equilibrium constants as a function of temperature, 3, 187

CUMULATIVE INDEX

409

Non-planar and planar aromatic systems, 1, 203 Norbornyl cation: reappraisal of structure, 1 1 , 179 Nuclear magnetic resonance, see N.M.K. Nucleophilic aromatic photosubstitution, 1 1 , 225 Nucleophilic catalysis of hydrolysis and related reactions, 4, 237 Nucleophilic vinylic substitution, 7, 1

0 - H bonds, hydrogen atom abstraction from, 9, 127 Oxygen isotope exchange reactions of organic compounds, 3, 123 Permutational isomerization of pentavalent phosphorus compounds, 9 , 25 Phosphorus compounds, pentavalent, turnstile rearrangement and pseudorotation in permutational isomerization, 9 , 2 5 Photochemistry of carbonium ions, 10, 129 Photosubstitution, nucleophilic aromatic, 1 1 , 225 Planar and non-planar aromatic systems, 1 , 203 Polarizability, molecular refractivity and, 3, 1 Polarog~aphyand reaction kinetics, 5 , 1 Polypeptides, calculations of conformations of, 6, 103 Protic and dipolar aprotic solvents, rates of bimolecular substitution reactions in, 5, 173 Protolytic processes in H20-DzO mixtures, 7, 259 Protonation sites in ambident conjugated systems, 1 I , 267 Pseudorotation in isomerization of pentavalent phosphorous compounds, 9, 25 Pyrolysis, gas-phase, of small-ring hydrocarbons, 4, 147 Radicals, organic free, identification by electron spin resonance, 1 , 284 Radicals, short-lived organic, electron spin resonance studies of, 5 , 53 Reaction kinetics, polarography and, 5, 1 Reactions mechanisms, use of volumes of activation for determining, 2 , 9 3 Reaction mechanisms in solution, entropies of activation and, 1 , 1 Reaction velocities and equilibrium constants, N.M.R. measurements o f , as a function of temperature, 3, 187 Reactions of hydrated electrons with organic compounds, 7, 115 Reactivity indices in conjugated molecules, 4, 73 Refractivity, molecular, and polarizability, 3, 1 Resonance, electron-spin, i d e n t i h a t i o n of organic free radicals by, 1 , 284 Resonance, electron-spin, studies of short-lived organic radicals, 5, 63 Short-lived organic radicals, electron-spin resonance studies of, 5 , 53 Small-ring hydrocarbons, gas-phase pyrolysis of, 4, 147 Solution, reactions in, entropies of activation and mechanisms, 1 , 1 Solvents, protic and dipolar aprotic, rates of bimolecular substitution reactions in, 5, 173 Spectroscopic observation of alkylcarbonium ions in strong acid solutions, 4, 305 Steric isotope effects, experiments on the nature of, 10, 1 Stereoselection in elementary steps of organic reactions, 6, 185 Structure and mechanism in carbene chemistry, 7, 153 Structure of electronically excited molecules, I , 365 Study of free radicals and their reactions at low temperatures using a rotating cryostat, 8, 1 Substitution, aromatic, a quantitative treatment of directive effects in, 1, 35 Substitution reactions, bimolccular, in protic and dipolar aprotic solvents, 5 , 173 Substitution reactions, aromatic, hydrogen isotope effects in, 2, 163 Superacid systems, 9, 1 Temperature, N.M.R. measurements of reaction velocities and equilibrium constants as a function of, 3, 187 Tritiated molecules, gaseous carbonium ions from the decay of, 8, 79 Tritium atoms, energetic, reactions with organic compounds, 2, 201 Turnstile rearrangement in isomerization of pentavalent phosphorus compounds, 9, 25

410

CUMULATIVE INDEX

Unsaturated compounds, basicity of, 4, 195 Vinyl cations, 9, 185 Volumes of activation, use of, for determining reaction mechanisms, 2,93