Solvents and Solvent E¤ects in Organic Chemistry docx

Following the same layout as in the second edition, all topics retained have beenbrought up to date, with smaller and larger changes and additions on nearly every page.Two Sections 4.4.7

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Prof Dr Christian Reichardt

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6 2003 WILEY-VCH Verlag GmbH & Co KGaA, Weinheim

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All rights reserved (including those of translation in other languages) No part of this book may be reproduced in any form – by photoprinting, microfilm, or any other means – nor transmitted or translated into machine language without written permission from the publishers Registered names, trademarks, etc used in this book, even when not specifically marked as such, are not to be considered unprotected by law.

Composition: Asco Typesetters, Hong Kong Printing: Strauss O¤setdruck GmbH, Mo¨rlenbach Bookbinding: J Scha¨¤er GmbH & Co KG, Gru¨nstadt

Printed in the Federal Republic of Germany.

and in memory of my parents

Preface to the Third Edition

Meeting the demand for the second edition of this book, which is – despite a reprint in

1990 – no longer available, and considering the progress that has been made during thelast decade in the study of solvent e¤ects in experimental and theoretical organic chem-istry, this improved third edition is presented to the interested reader

Following the same layout as in the second edition, all topics retained have beenbrought up to date, with smaller and larger changes and additions on nearly every page.Two Sections (4.4.7 and 5.5.13) are completely new, dealing with solvent e¤ects onhost/guest complexation equilibria and reactions in biphasic solvent systems and neo-teric solvents, respectively More than 900 new references have been added, giving pre-ference to review articles, and many older ones have been deleted New references eitherreplace older ones or are added to the end of the respective reference list of each chapter.The references cover the literature up to the end of 2001

From the vast number of published papers dealing with solvent e¤ects in all areas

of organic chemistry, only some illustrative examples from the didactic and systematicpoint of view could be selected This book is not a monograph covering all relevantliterature in this field of research The author, responsible for this subjective selec-tion, apologizes in advance to all chemists whose valuable work on solvent e¤ects isnot mentioned in this book However, using the reviews cited, the reader will find easyaccess to the full range of papers published in a certain field of research on solvente¤ects

Great progress has been made during the last decade in theoretical treatments ofsolvent e¤ects by various quantum-chemical methods and computational strategies.When indicated, relevant references are given to the respective solution reactions orabsorptions However, a critical evaluation of all the theoretical models and methodsused to calculate the di¤erential solvation of educts, activated complexes, products,ground and excited states, is outside the expertise of the present author Thus, a book onall kinds of theoretical calculations of solvent influences on chemical reactions andphysical absorptions has still to be written by someone else

Consistent use of the nomenclature,a) symbols,b) terms,c) and SI unitsd) mended by the IUPAC commissions has also been made in this third edition

recom-For comments and valuable suggestions I have to thank many colleagues, in ticular Prof E M Kosower, Tel Aviv/Israel, Prof R G Makitra, Lviv/Ukraine, Prof

par-N O Mchedlov-Petrossyan, Kharkiv/Ukraine, and Prof K Mo¨ckel, Mu¨hlhausen/Germany For their assistance in drawing formulae, preparing the indices, and provid-ing me with di‰cult to obtain literature, I thank Mr G Scha¨fer (technician), Mrs S.Schellenberg (secretary), and Mrs B Becht-Schro¨der (librarian), all at the Department

a) G J Leigh, H A Favre, and W V Metanomski: Principles of Chemical Nomenclature – A Guide to IUPAC Recommendations, Blackwell Science Publications, London, 1998.

b) I Mills, T Cvitas, K Homann, N Kallay, and K Kuchitsu: Quantities, Units and Symbols in

c) P Mu¨ller: Glossary of Terms used in Physical Organic Chemistry – IUPAC Recommendations

1994, Pure Appl Chem 66, 1077 (1994).

of Chemistry, Philipps University, Marburg/Germany Special thanks are due to thesta¤ of Wiley-VCH Verlag GmbH, Weinheim/Germany, particularly to Dr ElkeWestermann, for their fine work in turning the manuscript into the final book Lastly,

my biggest debt is to my wife Maria, not only for her assistance in the preparation of themanuscript, but also for her constant encouragement and support during the writing ofthis book

Preface to the Second Edition

The response to the first English edition of this book, published in 1979, has been bothgratifying and encouraging Its mixed character, lying between that of a monograph and

a textbook, has obviously made it attractive to both the industrial and academic chemist

as well as the advanced student of chemistry

During the last eight years the study of solvent e¤ects on both chemical tions and absorption spectra has made much progress, and numerous interesting andfascinating examples have been described in the literature In particular, the study ofionic reactions in the gas phase – now possible due to new experimental techniques –has allowed direct comparisons between gas-phase and solution reactions This has led

reac-to a greater understanding of solution reactions Consequently, Chapters 4 and 5 havebeen enlarged to include a description of ionic gas-phase reactions compared to theirsolution counterparts

The number of well-studied solvent-dependent processes, i.e reactions andabsorptions in solution, has increased greatly since 1979 Only a representative selection

of the more instructive, recently studied examples could be included in this secondedition

The search for empirical parameters of solvent polarity and their applications

in multiparameter equations has recently been intensified, thus making it necessary torewrite large parts of Chapter 7

Special attention has been given to the chemical and physical properties oforganic solvents commonly used in daily laboratory work Therefore, all AppendixTables have been improved; some have been completely replaced by new ones A newwell-referenced table on solvent-drying has been added (Table A-3) Chapter 3 has beenenlarged, in particular by the inclusion of solvent classifications using multivariate sta-tistical methods (Section 3.5) All these amendments justify the change in the title of thebook to Solvents and Solvent E¤ects in Organic Chemistry

The references have been up-dated to cover literature appearing up to the firstpart of 1987 New references were added to the end of the respective reference list ofeach chapter from the first edition

Consistent use of the nomenclature, symbols, terms, and SI units recommended

by the IUPAC commissions has also been made in the second edition.*)

I am very indebted to many colleagues for corrections, comments, and valuablesuggestions Especially helpful suggestions came from Professors H.-D Fo¨rsterling,Marburg, J Shorter, Hull/England, and R I Zalewski, Poznan´/Poland, to whom I amvery grateful For critical reading of the whole manuscript and the improvement of myEnglish I again thank Dr Edeline Wentrup-Byrne, now living in Brisbane/Australia

Dr P.-V Rinze, Marburg, and his son Lars helped me with the author index Finally,

I would like to thank my wife Maria for her sympathetic assistance during the tion of this edition and for her help with the indices

* Cf Pure Appl Chem 51, 1 (1979); ibid 53, 753 (1981); ibid 55, 1281 (1983); ibid 57, 105 (1985).

The organic chemist usually works with compounds which possess labile covalentbonds and are relatively involatile, thereby often rendering the gas-phase unsuitable as areaction medium Of the thousands of reactions known to occur in solution only fewhave been studied in the gas-phase, even though a description of reaction mechanisms ismuch simpler for the gas-phase The frequent necessity of carrying out reactions in thepresence of a more or less inert solvent results in two main obstacles: The reactiondepends on a larger number of parameters than in the gas-phase Consequently, theexperimental results can often be only qualitatively interpreted because the state ofaggregation in the liquid phase has so far been insu‰ciently studied On the other hand,the fact that the interaction forces in solution are much stronger and more varied than inthe gas-phase, permits to a¤ect the properties and reactivities of the solute in manifoldmodes.

Thus, whenever a chemist wishes to carry out a chemical reaction he not only has

to take into consideration the right reaction partners, the proper reaction vessels, andthe appropriate reaction temperature One of the most important features for the success

of the planned reaction is the selection of a suitable solvent Since solvent e¤ects onchemical reactivity have been known for more than a century, most chemists are nowfamiliar with the fact that solvents may have a strong influence on reaction rates andequilibria Today, there are about three hundred common solvents available, nothing tosay of the infinite number of solvent mixtures Hence the chemist needs, in addition tohis intuition, some general rules and guiding-principles for this often di‰cult choice.The present book is based on an earlier paperback ‘‘Lo¨sungsmittele¤ekte in derorganischen Chemie’’ [1], which, though following the same layout, has been completelyrewritten, greatly expanded, and brought up to date The book is directed both towardthe industrial and academic chemist and particularly the advanced student of chemistry,who on the one hand needs objective criteria for the proper choice of solvent but on theother hand wishes to draw conclusions about reaction mechanisms from the observedsolvent e¤ects

A knowledge of the physico-chemical principles of solvent e¤ects is required forproper bench-work Therefore, a description of the intermolecular interactions betweendissolved molecules and solvent is presented first, followed by a classification of solventsderived therefrom Then follows a detailed description of the influence of solvents onchemical equilibria, reaction rates, and spectral properties of solutes Finally, empiricalparameters of solvent polarity are given, and in an appendix guidelines to the everydaychoice of solvents are given in a series of Tables and Figures

The number of solvent systems and their associated solvent e¤ects examined is

so enormous that a complete description of all aspects would fill several volumes Forexample, in Chemical Abstracts, volume 85 (1976), approximately eleven articles perweek were quoted in which the words ‘‘Solvent e¤ects on ’’ appeared in the title Inthe present book only a few important and relatively well-defined areas of generalimportance have been selected The book has been written from the point of view ofpractical use for the organic chemist rather than from a completely theoretical one

In the selection of the literature more recent reviews were taken into accountmainly Original papers were cited in particular from the didactic point of view rather

than priority, importance or completeness This book, therefore, does not only have thecharacter of a monograph but also to some extent that of a textbook In order to helpthe reader in his use of the literature cited, complete titles of the review articles quotedare given The literature up until December 1977 has been considered together with afew papers from 1978 The use of symbols follows the recommendations of the SymbolsCommittee of the Royal Society, London, 1971 [2].

I am very grateful to Professor Karl Dimroth, Marburg, who first stimulated myinterest in solvent e¤ects in organic chemistry I am indebted to Professors W H Pirkle,Urbana/Illinois, D Seebach, Zu¨rich/Switzerland, J Shorter, Hull/England, and numer-ous other colleagues for helpful advice and information Thanks are also due to theauthors and publishers of copyrighted materials reproduced with their permission(cf Figure and Table credits on page 495) For the careful translation and improvement

of the English manuscript I thank Dr Edeline Wentrup-Byrne, Marburg Without theassistance and patience of my wife Maria, this book would not have been written

References

Weinheim 1973;

E¤ets de solvant en chimie organique (translation of the first-mentioned title into French, by

I Tkatchenko), Flammarion, Paris 1971;

Rastvoriteli v organicheskoi khimii (translation of the first-mentioned title into Russian, by E R Zakhsa), Izdatel’stvo Khimiya, Leningrad 1973.

[2] Quantities, Units, and Symbols, issued by The Symbols Committee of the Royal Society, don, in 1971.

Lon-Preface to the First Edition

XII

1 Introduction 1

2 Solute-Solvent Interactions 5

2.1 Solutions 5

2.2 Intermolecular Forces 10

2.2.1 Ion-Dipole Forces 10

2.2.2 Dipole-Dipole Forces 11

2.2.3 Dipole-Induced Dipole Forces 13

2.2.4 Instantaneous Dipole-Induced Dipole Forces 13

2.2.5 Hydrogen Bonding 15

2.2.6 Electron-Pair Donor/Electron-Pair Acceptor Interactions (EPD/EPA Interactions) 19

2.2.7 Solvophobic Interactions 27

2.3 Solvation 30

2.4 Selective Solvation 38

2.5 Micellar Solvation (Solubilization) 42

2.6 Ionization and Dissociation 46

3 Classification of Solvents 57

3.1 Classification of Solvents according to Chemical Constitution 57

3.2 Classification of Solvents using Physical Constants 62

3.3 Classification of Solvents in Terms of Acid-Base Behaviour 73

3.3.1 Brønsted-Lowry Theory of Acids and Bases 73

3.3.2 Lewis Theory of Acids and Bases 79

3.4 Classification of Solvents in Terms of Specific Solute/Solvent Interactions 82

3.5 Classification of Solvents using Multivariate Statistical Methods 84

4 Solvent E¤ects on the Position of Homogeneous Chemical Equilibria 93

4.1 General Remarks 93

4.2 Solvent E¤ects on Acid/Base Equilibria 95

4.2.1 Brønsted Acids and Bases in Solution 95

4.2.2 Gas-Phase Acidities and Basicities 99

4.3 Solvent E¤ects on Tautomeric Equilibria 106

4.3.1 Solvent E¤ects on Keto/Enol Equilibria 106

4.3.2 Solvent E¤ects on other Tautomeric Equilibria 113

4.4 Solvent E¤ects on other Equilibria 121

4.4.1 Solvent E¤ects on Brønsted Acid/Base Equilibria 121

4.4.2 Solvent E¤ects on Lewis Acid/Base Equilibria 123

4.4.3 Solvent E¤ects on Conformational Equilibria 126

4.4.4 Solvent E¤ects on cis/trans or E/Z Isomerization Equilibria 132

4.4.5 Solvent E¤ects on Valence Isomerization Equilibria 135

4.4.6 Solvent E¤ects on Electron-Transfer Equilibria 137

4.4.7 Solvent E¤ects on Host/Guest Complexation Equilibria 139

5 Solvent E¤ects on the Rates of Homogeneous Chemical Reactions 147

5.1 General Remarks 147

5.2 Gas-Phase Reactivities 155

5.3 Qualitative Theory of Solvent E¤ects on Reaction Rates 162

5.3.1 The Hughes–Ingold Rules 163

5.3.2 Solvent E¤ects on Dipolar Transition State Reactions 173

5.3.3 Solvent E¤ects on Isopolar Transition State Reactions 187

5.3.4 Solvent E¤ects on Free-Radical Transition State Reactions 199

5.3.5 Limitations of the Hughes–Ingold Rules 215

5.4 Quantitative Theories of Solvent E¤ects on Reaction Rates 218

5.4.1 General Remarks 218

5.4.2 Reactions between Neutral, Apolar Molecules 219

5.4.3 Reactions between Neutral, Dipolar Molecules 225

5.4.4 Reactions between Neutral Molecules and Ions 233

5.4.5 Reactions between Ions 234

5.5 Specific Solvation E¤ects on Reaction Rates 237

5.5.1 Influence of Specific Anion Solvation on the Rates of SNand other Reactions 238

5.5.2 Protic and Dipolar Aprotic Solvent E¤ects on the Rates of SN Reactions 243

5.5.3 Quantitative Separation of Protic and Dipolar Aprotic Solvent E¤ects for Reaction Rates by Means of Solvent-Transfer Activity Coe‰cients 254 5.5.4 Acceleration of Base-Catalysed Reactions in Dipolar Aprotic Solvents 259 5.5.5 Influence of Specific Cation Solvation on Rates of SNReactions 262

5.5.6 Solvent Influence on the Reactivity of Ambident Anions 269

5.5.7 Solvent E¤ects on Mechanisms and Stereochemistry of Organic Reactions 273

5.5.8 Influence of Micellar and Solvophobic Interactions on Reaction Rates and Mechanisms 292

5.5.9 Liquid Crystals as Reaction Media 298

5.5.10 Solvent Cage E¤ects 303

5.5.11 External Pressure and Solvent E¤ects on Reaction Rates 308

5.5.12 Solvent Isotope E¤ects 315

5.5.13 Reactions in Biphasic Solvent Systems and in Neoteric Solvents 317

6 Solvent E¤ects on the Absorption Spectra of Organic Compounds 329

6.1 General Remarks 329

6.2 Solvent E¤ects on UV/Vis Spectra 330

6.2.1 Solvatochromic Compounds 330

6.2.2 Theory of Solvent E¤ects on UV/Vis Absorption Spectra 340

6.2.3 Specific Solvent E¤ects on UV/Vis Absorption Spectra 348

6.2.4 Solvent E¤ects on Fluorescence Spectra 352

6.2.5 Solvent E¤ects on ORD and CD Spectra 359

6.3 Solvent E¤ects on Infrared Spectra 363

6.4 Solvent E¤ects on Electron Spin Resonance Spectra 369

Contents

XIV

6.5 Solvent E¤ects on Nuclear Magnetic Resonance Spectra 375

6.5.1 Nonspecific Solvent E¤ects on NMR Chemical Shifts 375

6.5.2 Specific Solvent E¤ects on NMR Chemical Shifts 381

6.5.3 Solvent E¤ects on Spin-Spin Coupling Constants 387

7 Empirical Parameters of Solvent Polarity 389

7.1 Linear Gibbs Energy Relationships 389

7.2 Empirical Parameters of Solvent Polarity from Equilibrium Measurements 396

7.3 Empirical Parameters of Solvent Polarity from Kinetic Measurements 402 7.4 Empirical Parameters of Solvent Polarity from Spectroscopic Measurements 411

7.5 Empirical Parameters of Solvent Polarity from other Measurements 443

7.6 Interrelation and Application of Solvent Polarity Parameters 445

7.7 Multiparameter Approaches 452

Appendix 471

A Properties, Purification, and Use of Organic Solvents 471

A.1 Physical Properties 471

A.2 Purification of Organic Solvents 471

A.3 Spectroscopic Solvents 479

A.4 Solvents as Reaction Media 488

A.5 Solvents for Recrystallization 488

A.6 Solvents for Extraction and Partitioning (Distribution) 490

A.7 Solvents for Adsorption Chromatography 492

A.8 Solvents for Acid/Base Titrations in Non-Aqueous Media 496

A.9 Solvents for Electrochemistry 496

A.10 Toxicity of Organic Solvents 500

References 509

Figure and Table Credits 581

Subject Index 583

Author Index 599

List of Abbreviations

Abbreviations and Recommended Values of Some Fundamental Constants andNumbersa,b)

[¼ 1=ðm0
c0 Þ; m0¼ permeability ofvacuum]

gas (at t¼ 0C and p¼ 100 kPa)

22.711 L
mol1

Abbreviations and Symbols for Unitsa,b)

(millilitre mL; 106m3)

volume

a) I Mills, T Cvitasˇ, K Homann, N Kallay, and K Kuchitsu: Quantities, Units and Symbols in

C degrees centigrade (Celsius) temperature

dm3 cubic decimetre (litre L; 103m3) volume

(magnetic field)

percent (%) part per hundred (102) dimensionless fraction

a electric polarizability of a molecule,

polarizability volume

C2
m2
J1or 4pe0
cm3

hydrogen-bond donor acidity (Taftand Kamlet)

basicity (Palm and Koppel)

BMeOD IR based empirical parameter of

solvent Lewis basicity (Palm andKoppel)

c) P Mu¨ller: Glossary of Terms used in Physical Organic Chemistry – IUPAC Recommendations

1994 Pure Appl Chem 66, 1077 (1994).

BPhOH IR based empirical parameter of

solvent Lewis basicity (Koppel andPaju; Makitra)

tendency or ‘basity’ (Swain)

hydrogen-bond acceptor basicity(Taft and Kamlet)

density) of a solvent

MPa (¼ 106Pa)

ci; cðiÞ molar concentration of solute i mol
L1

CA; CB Lewis acidity and Lewis basicity

parameter (Drago)

DHA molar bond-dissociation energy of the

bond between H and A

kJ
mol1

basicity, based on a 1,3-dipolarcycloaddition reaction (Nagai et al.)

[¼ DH(DaaSbCl5)]

kcal
mol1

(Marcus)

d; dH Hildebrand’s solubility parameter MPa1 =2

(Taft and Kamlet)

acidity (Palm and Koppel)

EA; EB Lewis acidity and Lewis basicity

parameter (Drago)

EN

parameter, based on the n! pabsorption of an aminyloxide radical(Mukerjee; Wrona)

EK empirical solvent polarity parameter,

based on the d! pabsorption of amolybdenum complex (Walther)

kcal
mol1

EMLCT empirical solvent polarity parameter,

based on the d! p absorption of atungsten complex (Lees)

ET molar electronic transition energy,

molar electronic excitation energy

kJ
mol1or kcal
mol1

ETð30Þ empirical solvent polarity parameter,

based on the intramolecular CTabsorption of a pyridinium-N-phenolate betaine dye (Dimroth andReichardt)

T empirical solvent polarity parameter,

based on the n! pabsorption of anS-oxide (Walter)

kcal
mol1

er relative permittivity (¼e=e0)

(‘‘dielectric constant’’)

based on the n! pabsorption ofketones (Dubois)

parameter (Schleyer and Allerhand)

DG standard molar Gibbs energy change kJ
mol1

kJ
mol1

DH0 standard molar enthalpy of activation kJ
mol1

vapourization

kJ
mol1

HBD hydrogen-bond donor

HOMO highest occupied molecular orbital

(Kova´ts)

(n¼ 1) and bimolecular (n ¼ 2)reactions

(L
mol1)n1
s1

k0 rate constant in a reference solvent or

in the gas phase for monomolecular(n¼ 1) and bimolecular reactions(n¼ 2)

(L
mol1)n1
s1

constant of unsubstituted substrates

(L
mol1)n 1
s1with

n¼ 1 or 2

K; Kc equilibrium constant (concentration

basis; v¼ stoichiometric number)

(mol
L1)Sv

Ka; Kb acid and base ionization constants (mol
L1)Sv

Kauto autoionization ion product,

Ko=w 1-octanol/water partition constant

(Hansch and Leo)

LUMO lowest unoccupied molecular orbital

parameter of substituents (Menger)

mi standard chemical potential of solute i kJ
mol1

myi standard chemical potential of solute i

at infinite dilution

kJ
mol1

n; nD refractive index (at sodium D line)

(¼ c0=c)

nucleophilicity (Winstein andGrunwald)

(nucleophileþ solvent)-systems(Ritchie)

n frequency in the gas phase or in an

inert reference solvent

Hz, s1

based on a Diels-Alder reaction(Berson)

bar (¼ 105Pa)

(Palm and Koppel)

based on 19F NMR measurements(Taft)

based on the p! p emission ofpyrene (Winnik)

Po =w 1-octanol/water partition coe‰cient

(Hansch and Leo)

(abbreviation of potentia hydrogenii

or puissance d’hydroge`ne (So¨rensen1909)

polarizability parameter, based

on the p! pabsorption ofsubstituted aromatics (Taft andKamlet)

pazo empirical solvent dipolarity/

polarizability parameter, based on the

p! pabsorption of azomerocyanine dyes (Buncel)

r; rA Hammett reaction resp absorption

constants

based on the Z-values (Brownstein)

tri-n-propylamine with iodomethane(Drougard and Decroocq)

DS0 standard molar entropy of activation J
K1
mol1

(Abraham)

hydrogen-bond donor acidity(Catala´n)

hydrogen-bond acceptor basicity(Catala´n)

dipolarity/polarizability, based on the

p! pabsorption of substituted nitrofluorenes (Catala´n)

Vm; Vm; i molar volume (of i) cm3
mol1

xi; xðiÞ mole fraction of iðxi¼ ni=PnÞ

based on an SE2 reaction (Gielen andNasielski)

wR; wB empirical solvent polarity parameters,

based on the p! pabsorption ofmerocyanine dyes (Brooker)

kcal
mol1

OyXS;WyXS solvent-transfer activity coe‰cient of

a solute X from a reference solvent(O) or water (W) to anothersolvent (S)

ionizing power, based on t-butylchloride solvolysis (Winstein andGrunwald)

YOTs empirical parameter of solvent

ionizing power, based on 2-adamantyltosylate solvolysis (Schleyer andBentley)

and Koppel)

negative for anions

based on the intermolecular CTabsorption of a substitutedpyridinium iodide (Kosower)

kcal
mol1

‘‘Agite, Auditores ornatissimi, transeamus alacres ad aliud negotii! quum enim sicsatis excusserimus ea quatuor Instrumenta artis, et naturae, quae modo relinquimus,videamus quintum genus horum, quod ipsi Chemiae fere proprium censetur, cui certeChemistae principem locum prae omnibus assignant, in quo se jactant, serioque tri-umphant, cui artis suae, prae aliis omnibus e¤ectus mirificos adscribunt Atque illudquidem Menstruum vocaverunt.’’*)

Hermannus Boerhaave (1668–1738)

De menstruis dictis in chemia, in:Elementa Chemiae (1733) [1, 2]

1 Introduction

The development of our knowledge of solutions reflects to some extent the development

of chemistry itself [3] Of all known substances, water was the first to be considered as asolvent As far back as the time of the Greek philosophers there was speculation aboutthe nature of solution and dissolution The Greek alchemists considered all chemicallyactive liquids under the name ‘‘Divine water’’ In this context the word ‘‘water’’ wasused to designate everything liquid or dissolved

The alchemist’s search for a universal solvent, the so-called ‘‘Alkahest’’ or struum universale’’, as it was called by Paracelsus (1493–1541), indicates the impor-tance given to solvents and the process of dissolution Although the eager search ofthe chemists of the 15th to 18th centuries did not in fact lead to the discovery of any

‘‘Men-‘‘Alkahest’’, the numerous experiments performed led to the uncovering of new solvents,new reactions, and new compounds**) From these experiences arose the earliest chem-ical rule that ‘‘like dissolves like’’ (similia similibus solvuntur) However, at that time,the words solution and dissolution comprised all operations leading to a liquid productand it was still a long way to the conceptual distinction between the physical dissolution

of a salt or of sugar in water, and the chemical change of a substrate by dissolution, forexample, of a metal in an acid Thus, in the so-called chemiatry period (iatrochemistryperiod), it was believed that the nature of a substance was fundamentally lost upon dis-solution Van Helmont (1577–1644) was the first to strongly oppose this contention Heclaimed that the dissolved substance had not disappeared, but was present in the solu-tion, although in aqueous form, and could be recovered [4] Nevertheless, the dissolution

* ‘‘Well then, my dear listeners, let us proceed with fervor to another problem! Having su‰ciently analyzed in this manner the four resources of science and nature, which we are about to leave (i.e fire, water, air, and earth) we must consider a fifth element which can almost be considered the most essential part of chemistry itself, which chemists boastfully, no doubt with reason, prefer above all others, and because of which they triumphantly celebrate, and to which they attribute above all others the marvellous e¤ects of their science And this they call the solvent (menstruum).’’

** Even if the once famous scholar J B Van Helmont (1577–1644) claimed to have prepared this

‘‘Alkahest’’ in a phial, together with the adherents of the alkahest theory he was ridiculed by his contemporaries who asked in which vessel he has stored this universal solvent.

Solvents and Solvent Effects in Organic Chemistry, Third Edition Christian Reichardt

Copyright 8 2003 WILEY-VCH Verlag GmbH & Co KGaA, Weinheim

of a substance in a solvent remained a rather mysterious process The famous Russianpolymath Lomonosov (1711–1765) wrote in 1747: ‘‘Talking about the process of disso-lution, it is generally said that all solvents penetrate into the pores of the body to bedissolved and gradually remove its particles However, concerning the question of whatforces cause this process of removal, there does not exist any somehow reasonableexplanation, unless one arbitrarily attributes to the solvents sharp wedges, hooks or,who knows, any other kind of tools’’ [27].

The further development of modern solution theory is connected with three sons, namely the French researcher Raoult (1830–1901) [28], the Dutch physical chemistvan’t Ho¤ (1852–1911) [5], and the Swedish scientist Arrhenius (1859–1927) [6] Raoultsystematically studied the e¤ects of dissolved nonionic substances on the freezing andboiling point of liquids and noticed in 1886 that changing the solute/solvent ratio pro-duces precise proportional changes in the physical properties of solutions The observa-tion that the vapour pressure of solvent above a solution is proportional to the molefraction of solvent in the solution is today known as Raoult’s law [28]

per-The di‰culty in explaining the e¤ects of inorganic solutes on the physical erties of solutions led in 1884 to Arrhenius’ theory of incomplete and complete dissoci-ation of ionic solutes (electrolytes, ionophores) into cations and anions in solution,which was only very reluctantly accepted by his contemporaries Arrhenius derived hisdissociation theory from comparison of the results obtained by measurements of elec-troconductivity and osmotic pressure of dilute electrolyte solutions [6]

prop-The application of laws holding for gases to solutions by replacing pressure byosmotic pressure was extensively studied by van’t Ho¤, who made osmotic pressuremeasurements another important physicochemical method in studies of solutions [5].The integration of these three basic developments established the foundations ofmodern solution theory and the first Nobel prizes in chemistry were awarded to van’tHo¤ (in 1901) and Arrhenius (in 1903) for their work on osmotic pressure and electro-lytic dissociation, respectively

The further development of solution chemistry is connected with the pioneeringwork of Ostwald (1853–1932), Nernst (1864–1941), Lewis (1875–1946), Debye (1884–1966), E Hu¨ckel (1896–1980), and Bjerrum (1879–1958) More detailed reviews on thedevelopment of modern solution chemistry can be found in references [29, 30]

The influence of solvents on the rates of chemical reactions [7, 8] was first noted

by Berthelot and Pe´an de Saint-Gilles in 1862 in connection with their studies on theesterification of acetic acid with ethanol: ‘‘ l’e´the´rification est entrave´e et ralentie parl’emploi des dissolvants neutres e´trangers a` la re´action’’ [9]*) After thorough studies onthe reaction of trialkylamines with haloalkanes, Menschutkin in 1890 concluded that areaction cannot be separated from the medium in which it is performed [10] In a letter

to Prof Louis Henry he wrote in 1890: ‘‘Or, l’expe´rience montre que ces dissolvantsexercent sur la vitesse de combinaison une influence conside´rable Si nous repre´sentonspar 1 la constante de vitesse de la re´action pre´cite´e dans l’hexane C6H14, cette constantepour la meˆme combinaison dans CH3aaCOaaC6H5, toutes choses e´gales d’ailleurs sera847.7 La di¤e´rence est e´norme, mais, dans ce cas, elle n’atteint pas encore le maxi-

* ‘‘ the esterification is disturbed and decelerated on addition of neutral solvents not belonging

to the reaction’’ [9].

mum Vous voyez que les dissolvants, soi-disant indi¤e´rents ne sont pas inertes; ilsmodifient profonde´ment l’acte de la combinaison chimique Cet e´nonce´ est riche enconse´quences pour la the´orie chimique des dissolutions’’ [26]*) Menschutkin also dis-covered that, in reactions between liquids, one of the reaction partners may constitute anunfavourable solvent Thus, in the preparation of acetanilide, it is not without impor-tance whether aniline is added to an excess of acetic acid, or vice versa, since aniline inthis case is an unfavorable reaction medium Menschutkin related the influence of sol-vents primarily to their chemical, not their physical properties.

The influence of solvents on chemical equilibria was discovered in 1896,simultaneously with the discovery of keto-enol tautomerism**) in 1,3-dicarbonyl com-pounds (Claisen [14]: acetyldibenzoylmethane and tribenzoylmethane; Wislicenus [15]:methyl and ethyl formylphenylacetate; Knorr [16]: ethyl dibenzoylsuccinate andethyl diacetylsuccinate) and the nitro-isonitro tautomerism of primary and secondarynitro compounds (Hantzsch [17]: phenylnitromethane) Thus, Claisen wrote: ‘‘Es gibtVerbindungen, welche sowohl in der Form aaC(OH)bbCaaaaCOaa wie in der FormaaCOaaCaaHaaCOaa zu bestehen vermo¨gen; von der Natur der angelagerten Reste, vonder Temperatur, bei den gelo¨sten Substanzen auch von der Art des Lo¨sungsmittels ha¨ngt

es ab, welche von den beiden Formen die besta¨ndigere ist’’ [14]***) The study of theketo-enol equilibrium of ethyl formylphenylacetate in eight solvents led Wislicenus tothe conclusion that the keto form predominates in alcoholic solution, the enol form inchloroform or benzene He stated that the final ratio in which the two tautomeric formscoexist must depend on the nature of the solvent and on its dissociating power, whereby

he suggested that the dielectric constants were a possible measure of this ‘‘power’’.Stobbe was the first to review these results [18] He divided the solvents into two groupsaccording to their ability to isomerize tautomeric compounds His classification reflects,

to some extent, the modern division into protic and aprotic solvents The e¤ect of vent on constitutional and tautomeric isomerization equilibria was later studied in detail

sol-* ‘‘Now, experience shows that solvents exert considerable influence on reaction rates If we

enor-mous, but in this case it has not even reached its maximum So you see that solvents, in spite of appearing at first to be indi¤erent, are by no means inert; they can greatly influence the course of chemical reactions This statement is full of consequences for the chemical theory of dissolutions’’ [26].

** The first observation of a tautomeric equilibrium was made in 1884 by Zincke at Marburg [11].

He observed that, surprisingly, the reaction of 1,4-naphthoquinone with phenylhydrazine gives the same product as that obtained from the coupling reaction of 1-naphthol with benzenediazonium salts This phenomenon, that the substrate can react either as phenylhydrazone or as a hydroxyazo compound, depending on the reaction circumstances, was called Ortsisomerie by Zincke [11] Later

on, the name tautomerism, with a di¤erent meaning however from that accepted today, was introduced by Laar [12] For a description of the development of the concept of tautomerism, see Ingold [13].

dissolved compounds, also on the nature of the solvent, which of the two forms will be the more stable’’ [14].

by Dimroth [19] (using triazole derivatives, e.g 1,2,3-triazole) and Meyer [20] (using ethyl acetoacetate).

5-amino-4-methoxycarbonyl-1-phenyl-It has long been known that UV/Vis absorption spectra may be influenced bythe phase (gas or liquid) and that the solvent can bring about a change in the position,intensity, and shape of the absorption band*) Hantzsch later termed this phenomenonsolvatochromism**) [22] The search for a relationship between solvent e¤ect and sol-vent property led Kundt in 1878 to propose the rule, later named after him, thatincreasing dispersion (i.e increasing index of refraction) is related to a shift of theabsorption maximum towards longer wavelength [23] This he established on the basis

of UV/Vis absorption spectra of six dyestu¤s, namely chlorophyll, fuchsin, anilinegreen, cyanine, quinizarin, and egg yolk in twelve di¤erent solvents The – albeit limited– validity of Kundt’s rule, e.g found in the cases of 4-hydroxyazobenzene [24] and ace-tone [25], led to the realization that the e¤ect of solvent on dissolved molecules is a result

of electrical fields These fields in turn originate from the dipolar properties of the ecules in question [25] The similarities in the relationships between solvent e¤ects onreaction rate, equilibrium position, and absorption spectra has been related to the gen-eral solvating ability of the solvent in a fundamental paper by Scheibe et al [25].More recently, research on solvents and solutions has again become a topic ofinterest because many of the solvents commonly used in laboratories and in the chemicalindustry are considered as unsafe for reasons of environmental protection On the list ofdamaging chemicals, solvents rank highly because they are often used in huge amountsand because they are volatile liquids that are di‰cult to contain Therefore, the intro-duction of cleaner technologies has become a major concern throughout both academiaand industry [31–34] This includes the development of environmentally benign newsolvents, sometimes called neoteric solvents (neoteric¼ recent, new, modern), constitut-ing a class of novel solvents with desirable, less hazardous, new properties [35, 36] Theterm neoteric solvents covers supercritical fluids, ionic liquids, and also perfluorohydro-carbons (as used in fluorous biphasic systems) Table A-14 in Chapter A.10 (Appendix)collects some recommendations for the substitution of hazardous solvents, together withthe relevant literature references

mol-For the development of a sustainable chemistry based on clean technologies, thebest solvent would be no solvent at all For this reason, considerable e¤orts haverecently been made to design reactions that proceed under solvent-free conditions, usingmodern techniques such as reactions on solid mineral supports (alumina, silica, clays),solid-state reactions without any solvent, support, or catalyst between neat reactants,solid-liquid phase-transfer catalysed and microwave-activated reactions, as well as gas-phase reactions [37–42] However, not all organic reactions can be carried out in theabsence of a solvent; some organic reactions even proceed explosively in the solid state!Therefore, solvents will still be useful in mediating and moderating chemical reactionsand this book on solvent e¤ects will certainly not become superfluous in the foreseeablefuture

* A survey of older works of solvent e¤ects on UV/Vis absorption spectra has been given by Sheppard [21].

** It should be noted that the now generally accepted meaning of the term solvatochromism di¤ers from that introduced by Hantzsch (cf Section 6.2).

2 Solute-Solvent Interactions

2.1 Solutions

In a limited sense solutions are homogeneous liquid phases consisting of more than onesubstance in variable ratios, when for convenience one of the substances, which is calledthe solvent and may itself be a mixture, is treated di¤erently from the other substances,which are called solutes [1] Normally, the component which is in excess is called thesolvent and the minor component(s) is the solute When the sum of the mole fractions ofthe solutes is small compared to unity, the solution is called a dilute solution*) A solu-tion of solute substances in a solvent is treated as an ideal dilute solution when the soluteactivity coe‰cients g are close to unity (g¼ 1) [1, 171] Solute/solvent mixtures A þ Bthat obey Raoult’s law over the entire composition range from pure A to pure B arecalled ideal solutions According to Raoult, the ratio of the partial pressure of compo-nent AðpAÞ to its vapour pressure as a pure liquid (p

A) is equal to the mole fraction of

AðxAÞ in the liquid mixture, i.e xA¼ pA=p

A Many mixtures obey Raoult’s law verywell, particularly when the components have a similar molecular structure (e.g benzeneand toluene)

A solvent should not be considered a macroscopic continuum characterized only

by physical constants such as density, dielectric constant, index of refraction etc., but as

a discontinuum which consists of individual, mutually interacting solvent molecules.According to the extent of these interactions, there are solvents with a pronouncedinternal structure (e.g water) and others in which the interaction between the solventmolecules is small (e.g hydrocarbons) The interactions between species in solvents (and

in solutions) are at once too strong to be treated by the laws of the kinetic theory ofgases, yet too weak to be treated by the laws of solid-state physics Thus, the solvent isneither an indi¤erent medium in which the dissolved material di¤uses in order to dis-tribute itself evenly and randomly, nor does it possess an ordered structure resembling acrystal lattice Nevertheless, the long-distance ordering in a crystal corresponds some-what to the local ordering in a liquid Thus, neither of the two possible models – the gasand crystal models – can be applied to solutions without limitation There is such a widegulf between the two models in terms of conceivable and experimentally establishedvariants, that it is too di‰cult to develop a generally valid model for liquids Due to thecomplexity of the interactions, the structure of liquids – in contrast to that of gases andsolids – is the least-known of all aggregation states Therefore, the experimental andtheoretical examination of the structure of liquids is among the most di‰cult tasks ofphysical chemistry [2–7, 172–174]

Any theory of the liquid state has to explain – among others – the following facts:Except for water, the molar volume of a liquid is roughly 10% greater than that of thecorresponding solid According to X-ray di¤raction studies, a short-range order of sol-vent molecules persists in the liquid state and the nearest neighbour distances are almostthe same as those in the solid The solvent molecules are not moving freely, as in the

* The superscript y attached to the symbol for a property of a solution denotes the property of an infinitely dilute solution.

Solvents and Solvent Effects in Organic Chemistry, Third Edition Christian Reichardt

Copyright 8 2003 WILEY-VCH Verlag GmbH & Co KGaA, Weinheim

gaseous state, but instead move in the potential field of their neighbours The potentialenergy of a liquid is higher than that of its solid by about 10% Therefore, the heat offusion is roughly 10% of the heat of sublimation Each solvent molecule has an envi-ronment very much like that of a solid, but some of the nearest neighbours are replaced

by holes Roughly one neighbour molecule in ten is missing

Even for the most important solvent – water – the investigation of its inner finestructure is still the subject of current research [8–15, 15a]*) Numerous di¤erent models,e.g the ‘‘flickering cluster model’’ of Franck and Wen [16], were developed to describethe structure of water However, all these models prove themselves untenable for acomplete description of the physico-chemical properties of water and an interpretation

of its anomalies [304] Fig 2-1 should make clear the complexity of the inner structure

of water

Liquid water consists both of bound ordered regions of a regular lattice, andregions in which the water molecules are hydrogen-bonded in a random array; it is per-meated by monomeric water and interspersed with random holes, lattice vacancies, andcages There are chains and small polymers as well as bound, free, and trapped watermolecules [9, 176] The currently accepted view of the structure of liquid water treats it

as a dynamic three-dimensional hydrogen-bonded network, without a significant ber of non-bonded water molecules, that retains several of the structural characteristics

num-of ice (i.e tetrahedral molecular packing with each water molecule hydrogen-bonded

to four nearest neighbours), although the strict tetrahedrality is lost [176] Its dynamicbehaviour resembles that of most other liquids, with short rotational and translationalcorrelation times of the order of 0.1 to 10 ps, indicating high hydrogen-bond exchangerates [176, 305]

In principle, other hydrogen-bonded solvents should possess similar complicatedstructures [306] However, whereas water has been thoroughly studied [17, 176, 307], theinner structures of other solvents are still less well known [172, 177–179] By way ofexample, the intermolecular structure of acetone is determined mainly by steric inter-actions between the methyl groups and, unexpectedly, only to a small extent by dipole/dipole forces [308], whereas the inner structure of dimethyl sulfoxide is dictated bystrong dipole/dipole interactions [309]

From the idea that the solvent only provides an indi¤erent reaction medium,comes the Ruggli-Ziegler dilution principle, long known to the organic chemist Accord-ing to this principle, in the case of cyclization reactions, the desired intramolecularreaction will be favoured over the undesired intermolecular reaction by high dilutionwith an inert solvent [18, 310]

The assumption of forces of interaction between solvent and solute led, on theother hand, to the century-old principle that ‘‘like dissolves like’’ (similia similibus sol-vuntur), where the word ‘‘like’’ should not be too narrowly interpreted In many cases,the presence of similar functional groups in the molecules su‰ces When a chemical

* The amusing story of ‘‘polywater,’’ which excited the scientific community for a few years during the late 1960’s and early 1970’s, has been reviewed by Franks [175] It turned out that polywater was not a new and more stable form of pure water, but merely dirty water The strange properties

of polywater were due to high concentrations of siliceous material dissolved from quartz capillaries

in which it was produced.

Fig 2-1 Two-dimensional schematic diagram of the three-dimensional structure of liquid water [9].

similarity is present, the solution of the two components will usually have a structuresimilar to that of the pure materials (e.g alcohol-water mixtures [19]) This rule ofthumb has only limited validity, however, since there are many examples of solutions ofchemically dissimilar compounds For example, methanol and benzene, water and N,N-dimethylformamide, aniline and diethyl ether, and polystyrene and chloroform, are allcompletely miscible at room temperature On the other hand, insolubility can occur inspite of similarity of the two partners Thus, polyvinyl alcohol does not dissolve inethanol, acetyl cellulose is insoluble in ethyl acetate, and polyacrylonitrile is insoluble inacrylonitrile [20] Between these two extremes there is a whole range of possibilitieswhere the two materials dissolve each other to a limited extent The system water/diethylether is such an example Pure diethyl ether dissolves water to the extent of 15 mg/g at

25C, whereas water dissolves diethyl ether to the extent of 60 mg/g When one of thetwo solvents is in large excess a homogeneous solution is obtained Two phases occurwhen the ratio is beyond the limits of solubility A more recent example of a rea‰rma-tion of the old ‘‘like dissolves like’’ rule is the di¤erential solubility of fullerene (C60),consisting of a three-dimensional delocalized 60p-electron system, in solvents such asmethanol (s¼ 0:01 mg/mL) and 1-chloronaphthalene (s ¼ 50 mg/mL) [311]

However, rather than the ‘‘like dissolves like’’ rule, it is the intermolecular action between solvent and solute molecules that determines the mutual solubility Acompound A dissolves in a solvent B only when the intermolecular forces of attraction

inter-KAA and KBB for the pure compounds can be overcome by the forces KAB in solution[21]

The sum of the interaction forces between the molecules of solvent and solute can

be related to the so-called polarity*) of A and B Denoting compounds with large actions A   A or B    B, respectively, as polar, and those with small interactions asnonpolar, four cases allowing a qualitative prediction of solubility can be distinguished(Table 2-1)

inter-An experimental verification of these simple considerations is given by the bility data in Table 2-2

solu-Table 2-1 Solubility and polarity [22].

Solubility of

A in B

a) Not much change for solute or solvent.

b) Di‰cult to break up B    B.

c) Di‰cult to break up A    A.

* For a more detailed definition of solvent polarity, see Sections 3.2 and 7.1.

The solubilities of ethane and methane are higher in nonpolar methane, whereas the opposite is true for chloromethane and dimethyl ether A survey

tetrachloro-of the reciprocal miscibility tetrachloro-of some representative examples tetrachloro-of organic solvents is given

in Fig 2-2

Solubility is commonly defined as the concentration of dissolved solute in a vent in equilibrium with undissolved solute at a specified temperature and pressure For

sol-a deeper sol-and more detsol-ailed understsol-anding of the diverse rules determining the solubility

of organic compounds in various solvents, see references [312–316]

The solubility parameter d of Hildebrand [4, 24], as defined in Eq (2-1), can often

be used in estimating the solubility of non-electrolytes in organic solvents

Table 2-2 Solubilities of methane, ethane, chloromethane, and dimethyl ether in

tetrachloro-methane (nonpolar solvent) and acetone (polar solvent) [22].

miscibility; little miscibility; without line: immiscible.

respectively d is a solvent property which measures the work necessary to separate thesolvent molecules (i.e disruption and reorganization of solvent/solvent interactions) tocreate a suitably sized cavity, large enough to accommodate the solute Accordingly,highly ordered self-associated solvents exhibit relatively large d values (d¼ 0 for the gasphase) As a rule, it has been found that a good solvent for a certain non-electrolyte has

a d value close to that of the solute [20, 24, 25]; cf Table 3-3 in Section 3.2 for a tion of d values Often a mixture of two solvents, one having a d value higher and theother having a d value lower than that of the solute is a better solvent than each of thetwo solvents separately [24]; cf also Section 3.2

collec-A nice example demonstrating mutual insolubility due to di¤erent d values hasbeen described by Hildebrand [180], and was later improved [181] A system of eightnon-miscible liquid layers was constructed The eight layers in order of increasing den-sities are para‰n oil, silicon oil, water, aniline, perfluoro(dimethylcyclohexane), whitephosphorus, gallium, and mercury This system is stable indefinitely at 45C; this tem-perature is required to melt the gallium and phosphorus [181] A simplified, less harmfulversion with five colourless liquid phases consists of mineral (para‰n) oil, methyl siliconoil, water, benzyl alcohol, and perfluoro(N-ethylpiperidine) (or another perfluoro-organic liquid), in increasing order of density [317] Addition of an organic-soluble dyecan colour some of the five layers

2.2 Intermolecular Forces [26, 27, 182–184]

Intermolecular forces are those which can occur between closed-shell molecules [26, 27].These are also called van der Waals forces, since van der Waals recognized them as thereason for the non-ideal behaviour of real gases Intermolecular forces are usually clas-sified into two distinct categories The first category comprises the so-called directional,induction, and dispersion forces, which are non-specific and cannot be completely satu-rated ( just as Coulomb forces between ions cannot) The second group consists ofhydrogen-bonding forces, and charge-transfer or electron-pair donor–acceptor forces.The latter group are specific, directional forces, which can be saturated and lead to stoi-chiometric molecular compounds For the sake of completeness, in the following theCoulomb forces between ions and electrically neutral molecules (with permanent dipolemoments) will be considered first, even though they do not belong to the intermolecularforces in the narrower sense

2.2.1 Ion-Dipole Forces [28, 185]

Electrically neutral molecules with an unsymmetrical charge distribution possess a manent dipole moment m If the magnitude of the two equal and opposite charges of thismolecular dipole is denoted by q, and the distance of separation l, the dipole moment isgiven by m¼ q  l When placed in the electric field resulting from an ion, the dipole willorient itself so that the attractive end (the end with charge opposite to that of the ion)will be directed toward the ion, and the other repulsive end directed away The potentialenergy of an ion-dipole interaction is given by

is positioned next to the ion in such a way that the ion and the separated charges ofthe dipole are linearly arranged ( or ) Equation (2-2) gives thefree energy for the interaction of an ionic charge z e and a so-called ‘point-dipole’(for which l¼ 0) in vacuum For typical interatomic spacings (rA300–400 pm), theion-dipole interaction is much stronger than the thermal energy k T at 300 K Forthe monovalent sodium cation (z¼ þ1, radius ¼ 95 pm) near a water molecule(radius A 140 pm; m¼ 5:9  1030 Cm), the maximum interaction energy calculated by

Eq (2-2) amounts to U¼ 39k  T or 96 kJ  mol1at 300 K [26b].

Only molecules possessing a permanent dipole moment should be called dipolarmolecules Apart from a few hydrocarbons (n-hexane, cyclohexane, and benzene) andsome symmetrical compounds (carbon disulfide, tetrachloromethane, and tetra-chloroethene) all common organic solvents possess a permanent dipole moment ofbetween 0 and 18 1030 Cm (i.e Coulombmeter) Among the solvents listed in theAppendix, Table A-1, hexamethylphosphoric triamide is the one with the highest dipolemoment (m¼ 18:48  1030 Cm), followed by propylene carbonate (m¼ 16:7  1030Cm), and sulfolane (m¼ 16:05  1030 Cm) The largest dipole moments amongst fluidsare exhibited by zwitterionic compounds such as the sydnones (i.e 3-alkyl-1,2,3-oxadiazolium-5-olates) For example, 4-ethyl-3-(1-propyl)sydnone, a high-boiling liquid(tbp¼ 155C/3 Torr) with a large relative permittivity (e

r¼ 64:6 at 25C), has a dipolemoment of m¼ 35:7  1030 Cm (¼10.7 D) [318] The peculiar physical properties ofsuch room temperature liquid sydnones make them to good nonaqueous dipolar sol-vents for many ionophores (electrolytes)

Ion-dipole forces are important for solutions of ionic compounds in dipolar vents, where solvated species such as Na(OH2)lm and Cl(H2O)mn (for solutions of NaCl

sol-in H2O) exist In the case of some metal ions, these solvated species can be su‰cientlystable to be considered as discrete species, such as [Co(NH3)6]3lor Ag(CH3CN)l2 4.For a comprehensive review on ion/solvent interactions, see reference [241]

2.2.2 Dipole-Dipole Forces [29]

Directional forces depend on the electrostatic interaction between molecules possessing

a permanent dipole moment m due to their unsymmetrical charge distribution Whentwo dipolar molecules are optimally oriented with respect to one another at a distance r

as shown in Fig 2-3a, then the force of attraction is proportional to 1=r3 An alternativearrangement is the anti-parallel arrangement of the two dipoles as shown in Fig 2-3b

* It should be noted that Eqs (2-2) to (2-6) are valid only for gases; an exact application to tions is not possible Furthermore, Eqs (2-2) to (2-6) are restricted to cases with r g l.

Unless the dipole molecules are very voluminous, the second arrangement is themore stable one The two situations exist only when the attractive energy is larger thanthe thermal energies Therefore, the thermal energy will normally prevent the dipolesfrom optimal orientation If all possible orientations were equally probable, that is, thedipoles correspond to freely rotating molecules, then attraction and repulsion wouldcompensate each other The fact that dipole orientations leading to attraction are sta-tistically favored leads to a net attraction, which is strongly temperature dependent,according to Eq (2-3) (kB¼ Boltzmann constant; T ¼ absolute temperature) [29].

to Eq (2-3), for pairs of dipolar molecules with m¼ 3:3  1030 Cm (¼1 D), at a ration of 500 pm, the average interaction energy is about 0.07 kJ  mol1 at 25 C.This is clearly smaller than the average molar kinetic energy of 3/2 k T ¼ 3:7

sepa-kJ mol1at the same temperature [26d].

Among other interaction forces, these dipole-dipole interactions are mainlyresponsible for the association of dipolar organic solvents such as dimethyl sulfoxide [30]

or N,N-dimethylformamide [31]

It should be mentioned that dipoles represent only one possibility for the chargearrays in electric multipoles (n-poles) n-Poles with an array of point charges with ann-pole moment (but no lower moment) are n-polar Thus, a monopole (n¼ 1) is a pointcharge and a monopole moment represents an overall charge (e.g of an ion Naþ or

Cl) A dipole (n¼ 2; e.g H2O, H3CaaCOaaCH3) is an array of partial charges with

no monopole moment (i.e no charge) A quadrupolar molecule (n¼ 4; e.g CO2, C6H6)has neither a net charge nor a dipole moment, and an octupolar molecule (n¼ 8; e.g

CH4, CCl4) has neither charge nor a dipole or quadrupole moment In addition todipole/dipole interactions, in solution there can also exist such higher intermolecularmultipole/multipole interactions Therefore, to some degree, octupolar tetrachloro-methane is also a kind of polar solvent However, the intermolecular interaction energyrapidly falls o¤ at higher orders of the multipole [26d] The anomalous behaviour of the

Fig 2-3 (a) ‘‘Head-to-tail’’ arrangement of two dipole molecules; (b) Antiparallel arrangement of two dipole molecules.

chair-configured, non-dipolar solvent 1,4-dioxane, which often behaves like a polar vent even though its relative permittivity is low (er¼ 2:2), is caused by its large nonidealquadrupolar charge distribution [411].

sol-2.2.3 Dipole-Induced Dipole Forces [32]

The electric dipole of a molecule possessing a permanent dipole moment m can induce

a dipole moment in a neighbouring molecule This induced moment always lies in thedirection of the inducing dipole Thus, attraction always exists between the two partners,which is independent of temperature The induced dipole moment*) will be bigger thelarger the polarizability a of the apolar molecule experiencing the induction of the per-manent dipole The net dipole/induced dipole energy of interaction for two di¤erentmolecules, each possessing a permanent dipole moment m1and m2and polarizabilities a1and a2, often referred to as the induction or Debye interaction [32], is given by Eq (2-4)

Similarly, a charged particle such as an ion introduced into the neighbourhood of

an uncharged, apolar molecule will distort the electron cloud of this molecule in thesame way The polarization of the neutral molecule will depend upon its inherentpolarizability a, and on the polarizing field a¤orded by the ion with charge z e Theenergy of such an interaction is given by Eq (2-5)

Uion-induced dipole¼  1

ð4p  e0Þ2z22 e r24 a ð2-5ÞThe importance of both of these interactions is limited to situations such as solutions ofdipolar or ionic compounds in nonpolar solvents

2.2.4 Instantaneous Dipole-Induced Dipole Forces [33, 34, 186]

Even in atoms and molecules possessing no permanent dipole moment, the continuouselectronic movement results, at any instant, in a small dipole moment m, which canfluctuatingly polarize the electron system of the neighbouring atoms or molecules Thiscoupling causes the electronic movements to be synchronized in such a way that amutual attraction results The energy of such so-called dispersion or London [33] inter-

elec-tric polarizability of the molecule; E elecelec-tric field strength).

actions can be expressed as

Udispersion¼  1

Dispersion forces are extremely short-range in action (depending on 1=r6!)

Dispersion forces are universal for all atoms and molecules; they alone areresponsible for the aggregation of molecules which possess neither free charges norelectric dipole moments Due to the greater polarizability of p-electrons, especiallystrong dispersion forces exist between molecules with conjugated p-electron systems (e.g.aromatic hydrocarbons) For many other dipole molecules with high polarizability aswell, the major part of the cohesion is due to dispersion forces For example, the calcu-lated cohesion energy of liquid 2-butanone at 40C consists of 8% orientational energy,14% inductional energy, and 78% dispersion energy [35] Two molecules with

a¼ 3  1030 m3, I ¼ 20  1019 J, and r¼ 3  1010 m have an interaction potential of

11.3 kJ/mol (2.7 kcal/mol) [35a] These values of a, I, and the average intermoleculardistance r correspond to those for liquid HCl It is instructive to compare the magnitude

of these dispersion forces with that of the dipole-dipole interactions For two dipoles,both with dipole moments of 3:3  1030 Cm (1.0 D), separated by a distance of

r¼ 3  1010 m and oriented as in Fig 2-3a, the interaction energy is only5.3 kJ/mol(1.1 kcal/mol) [35a] Thus, for HCl and most other compounds, the dispersion forcesare considerably stronger than the dipole-dipole forces of nearest neighbour distance inthe liquid state However, at larger distances the dispersion energy falls o¤ rapidly

As a result of the a2term in Eq (2-6b), dispersion forces increase rapidly with themolecular volume and the number of polarizable electrons The polarizability a is con-nected with the molar refraction and the index of refraction, according to the equation

of Lorenz-Lorentz Therefore, solvents with a large index of refraction, and hence largeoptical polarizability, should be capable of enjoying particularly strong dispersionforces As indicated in Table A-1 (Appendix), all aromatic compounds possess relativelyhigh indices of refraction, e.g quinoline (n¼ 1:6273), iodobenzene (n ¼ 1:6200), aniline(n¼ 1:5863), and diphenyl ether (n ¼ 1:5763); of all organic solvents, carbon disulfide(n¼ 1:6275) and diiodomethane (n ¼ 1:738) have the highest indices of refraction.Solvents with high polarizability are often good solvators for anions which alsopossess high polarizability This is due to the fact that the dispersional interactionsbetween the solvents and the large, polarizable anions like Im3 , Im, SCNmor the picrateanion are significantly larger than for the smaller anions like Fm, HOm, or R2Nm[36].Perfluorohydrocarbons have unusually low boiling points because tightly held electrons

in fluorine have only a small polarizability

2.2.5 Hydrogen Bonding [37–46, 187–190, 306]

Liquids possessing hydroxy groups or other groups with a hydrogen atom bound to anelectronegative atom X are strongly associated and have abnormal boiling points Thisobservation led to the contention that particular intermolecular forces apply here Theseare designated as hydrogen bridges, or hydrogen bonds, characterized by a coordinativedivalency of the hydrogen atom involved A general definition of the hydrogen bond is:when a covalently bound hydrogen atom forms a second bond to another atom, thesecond bond is referred to as a hydrogen bond [38]

The concept of hydrogen bonding was introduced in 1919 by Huggins [37].The first definitive paper on hydrogen bonding – applied to the association of watermolecules – was published in 1920 by Latimer and Rodebush [191] All three wereworking in the Laboratory of G N Lewis, University of California, Berkeley/USA

A hydrogen bond is formed by the interaction between the partners RaaXaaHand :YaaR0according to Eq (2-7).

ð2-7Þ

RaaXaaH is the proton donor and :YaaR0 makes available an electron pairfor the bridging bond Thus, hydrogen bonding can be regarded as a preliminarystep in a Brønsted acid-base reaction which would lead to a dipolar reaction product

RaaXm   HaaYlaaR0 X and Y are atoms of higher electronegativity than hydrogen(e.g C, N, P, O, S, F, Cl, Br, I) Both inter- and intramolecular hydrogen bonding arepossible, the latter when X and Y belong to the same molecule

The most important electron pair donors (i.e hydrogen bond acceptors) are theoxygen atoms in alcohols, ethers, and carbonyl compounds, as well as nitrogen atoms inamines and N-heterocycles Hydroxy-, amino-, carboxyl-, and amide groups are themost important proton donor groups Strong hydrogen bonds are formed by the pairs

OaaH    O, OaaH    N, and NaaH    O, weaker ones by NaaH    N, and theweakest by Cl2CaaH    O and Cl2CaaH    N The p-electron systems of aromaticcompounds, alkenes, and alkynes can also act as weak hydrogen bond acceptors [189].When two or more molecules of the same type associate, so-called homo-intermolecular hydrogen bonds are formed (Fig 2-4) The association of di¤erent mole-cules (e.g RaaOaaH    NR3) results in hetero-intermolecular hydrogen bonds Thedesignations homo- and heteromolecular [192] as well as homo- and heteroconjugatedhydrogen bond are also in use A remarkable example of a competitive solvent-dependent equilibrium between homo- and hetero-intermolecular hydrogen-bond asso-ciated species has been found in solutions of 4-hydroxyacetophenone and 2-(2-hexyloxyethoxy)ethanol [319]

Hydrogen bonds can be either intermolecular or intramolecular Both types ofhydrogen bonds are found in solutions of 2-nitrophenol, depending on the Lewis basic-ity of the solvent [298] The intramolecularly hydrogen-bonded form exists in non-hydrogen-bonding solvents (e.g cyclohexane, tetrachloromethane) 2-Nitrophenol breaksits intramolecular hydrogen bond to form an intermolecular one in electron-pair donor(EPD) solvents (e.g anisole, HMPT)

Circular hydrogen bonds have been found in the hexahydrate of a-cyclodextrin(cyclohexaamylose) [193] Hydration water molecules and hydroxy groups of the ma-cromolecule cooperate to form a network-like pattern with circular OaaH    O hydro-gen bonds If the OaaH    O hydrogen bonds run in the same direction, the circle iscalled homodromic Circles with the two counter-running chains are called antidromic,and circles with more randomly oriented chains are designated heterodromic [193]; cf.Fig 2-4a Such circular hydrogen bonds can be of importance with respect to the innermolecular structure of water and alcohols (cf also Fig 2-1).

The question of the exact geometry of hydrogen bonds (distances, angles, pair directionality) has been reviewed [194]

lone-The bond dissociation enthalpy for normal hydrogen bonds is ca 13 42 kJ/mol(3 10 kcal/mol)*) For comparison, covalent single bonds have dissociation enthalpies

of 210 420 kJ/mol (50 100 kcal/mol) Thus, hydrogen bonds are approx ten timesweaker than covalent single bonds, but also approx ten times stronger than the non-

Fig 2-4 Homo-intermolecular hydrogen bonds in alcohols, carboxylic acids, and amides (the hydrogen bonds are denoted by dotted lines).

Fig 2-4a Three types of circular hydrogen bonds: (a) homodromic, (b) antidromic, and (c) dromic hydrogen bonds [193].

hetero-* Bond dissociation enthalpies outside these limits are, however, known Examples of weak,

proton-acceptor and the acidity of the proton-donor molecule Compounds with very strong drogen bonds have been reviewed [320].

hy-specific intermolecular interaction forces The question as to whether or not a hydrogenbond is stronger than the equivalent deuterium bond is addressed in reference [321]: theD-bond seems to be somewhat stronger than the H-bond in the case of neutral hydro-gen-bonded complexes, but the reverse is true for charged complexes.

Hydrogen bonds are characterized by the following structural and spectroscopicfeatures [39]: (a) the distances between the neighbouring atoms involved in the hydrogenbond [X and Y in Eq (2-7)] are considerably smaller than the sum of their van derWaals radii; (b) the XaaH bond length is increased and hydrogen bond formationcauses its IR stretching mode to be shifted towards lower frequencies (for exceptions seereference [190]); (c) the dipolarity of the XaaH bond increases on hydrogen bond for-mation, leading to a larger dipole moment of the complex than expected from vectorialaddition of its dipolar components RaaXaaH and YaaR0; (d) due to the reduced elec-tron density at H-atoms involved in hydrogen bonds, they are deshielded, resulting insubstantial downfield shifts of their1H NMR signals; (e) in hetero-molecular hydrogenbonds, a shift of the Brønsted acid/base equilibrium RaaXaaH    YaaR0S Raa

Xm   HaaYlaaR0to the right-hand side with increasing solvent polarity is found (cf.Section 4.4.1 and references [195, 322] for impressive examples)

Up until now there has been no general agreement as to the best description of thenature of the forces in the hydrogen bond [42–46] The hydrogen bond can be described

as a dipole-dipole or resonance interaction Since hydrogen bonding occurs only whenthe hydrogen is bound to an electronegative atom, the first assumption concerning thenature of the hydrogen bond was that it consists of a dipole-dipole interaction such as

RaaXdmaaHdl   YdmaaR0 This viewpoint is supported by the fact that the strongesthydrogen bonds are formed in pairs in which the hydrogen is bonded to the most elec-tronegative elements (e.g FaaH    Fm, DH ¼ 155 kJ/mol) The greater strength ofthe hydrogen bond compared with non-specific dipole-dipole interactions is due to themuch smaller size of the hydrogen atom relative to any other atom, which allows it toapproach another dipole more closely This simple dipole model accounts for the usuallinear geometry of the hydrogen bond, because a linear arrangement maximizes theattractive forces and minimizes the repulsion

However, there are reasons to believe that more is involved in hydrogen bondingthan simply an exaggerated dipole-dipole interaction The shortness of hydrogen bondsindicates considerable overlap of van der Waals radii and this should lead to repulsiveforces unless otherwise compensated Also, the existence of symmetrical hydrogen bonds

of the type Fdm   H    Fdm cannot be explained in terms of the electrostatic model.When the XaaY distance is su‰ciently short, an overlap of the orbitals of the XaaHbond and the electron pair of :Y can lead to a covalent interaction According to Eq.(2-8), this situation can be described by two contributing ‘‘protomeric’’ structures, whichdi¤er only in the position of the proton*)

ð2-8Þ

* The term ‘‘protomeric structure’’ was obviously introduced in analogy to the well-known

‘‘mesomeric structures’’, which are used to describe the electronic ground state of aromatic pounds such as benzene in terms of a resonance hybrid [323].

The approximate quantum mechanical description of proton states by linearcombination of these protomeric structures has been called protomerism (symbol p) [323,324] It seems to be applicable to hydrogen bond systems in which a proton transfer mayoccur between two potential minima of equal depth [323, 324].

Solvents containing proton-donor groups are designated protic solvents [36] orHBD solvents [196]; solvents containing proton-acceptor groups are called HBA sol-vents [196] The abbreviations HBD (hydrogen-bond donor) and HBA (hydrogen-bondacceptor) refer to donation and acceptance of the proton, and not to the electron pairinvolved in hydrogen bonding

Solvents without proton-donor groups have been designated aprotic solvents [36].However, this term is rather misleading, since, for example, solvents commonly referred

to as dipolar aprotic (e.g CH3SOCH3, CH3CN, CH3NO2) are in fact not aprotic Inreactions where strong bases are employed, their protic character can be recognized.Therefore, the term aprotic solvents should be replaced by nonhydroxylic or better still

by non-HBD solvents [197]

Typical protic or HBD solvents are water, ammonia, alcohols, carboxylic acids,and primary amides Typical HBA solvents are amines, ethers, ketones, and sulfoxides.Amphiprotic solvents can act both as HBD and as HBA solvents simultaneously (e.g.water, alcohols, amides; cf Fig 2-4)

In type-A hydrogen bonding, the solute acts as a HBA-base and the solvent as aHBD-acid; in type-B hydrogen bonding, the roles are reversed [196]

Hydrogen bonding is responsible for the strong, temperature-dependent self- andhetero-association of amphiprotic solvents (e.g water, alcohols, amides)

The molecular structure of binary HBD/HBA solvent mixtures is largely mined by intermolecular hydrogen bonding between the two components, which usuallyleads to pronounced deviations from ideal solution behaviour [306, 325–327] Repre-sentative examples are trichloromethane/acetone [326] and trichloromethane/dimethylsulfoxide mixtures [327], which readily form hydrogen-bonded 1:1 and 2:1 complexes,respectively, with distinct changes in their physical properties as a consequence

deter-Hydrogen bonding plays a particularly important role in the interactions betweenanions and HBD solvents Hence, HBD solvents are good anion solvators Due to thesmall size of the hydrogen atom, small anions like Fm, Clm, or HOm are more e¤ec-tively solvated by such solvents than the larger ones, e.g Im3 , Im, SCNm, or the picrateion [36] This is also one of the reasons why the Gibbs energy of hydration,DGsolv, ofthe halide ions decreases in the series Fm> Clm

> Brm

> Im[49]

Hydrogen bonding is of paramount importance for the stabilization and the shape

of large biological molecules in living organisms (e.g cellulose, proteins, nucleic acids).For instance, the anaesthetic properties of some halogen-containing solvents such aschloroform, halothane (CF3aaCHClBr), and methoxyflurane (CH3OaaCF2aaCHCl2)have been connected with their ability to hinder the formation of biologically importanthydrogen bonds This is shown in the following equilibrium [300]:

Halohydrocarbon solvents containing an acidic CaaH bond shift this equilibrium infavour of free or less associated species, thus perturbing the ion channels which deter-mine the permeability of neuron membranes to Kl/Nal ions in the nervous system.Hydrogen bonds play a decisive role in determining the structure and dimension of theseion channels, on which this permeability depends [300].

Hydrogen-bonding also seems to be the molecular basis of sweetness All sweetcompounds seemingly have a H-bond donor and a H-bond acceptor ca 250 400 pmapart, which can form hydrogen bonds with a complementary pair on the sweet receptor

in the tastebuds of the tongue [328]

The e¤ectiveness of solvents (and solutes) as hydrogen-bond donors and/oracceptors has been studied experimentally using suitable reference compounds, com-prising representative HBDs or HBAs, in order to construct quantitative scales of sol-vent (and solute) hydrogen-bond acidity and hydrogen-bond basicity, respectively Forreviews on their construction and application to physicochemical and biochemical pro-cesses, see references [329–334] as well as Chapter 7 Scales of hydrogen-bond acidityand basicity have mostly been set up using complex formation constants, as determined

in inert solvents [329–332] For example, the strength of HBAs has been measured fromthe Gibbs energy changeDGHB for the formation of 1:1 hydrogen-bonded complexesbetween all kinds of HBAs (bases) and the reference HBD 4-fluorophenol in tetra-chloromethane at 25C [331, 332] Other attempts to construct scales of HBD and HBAstrengths, e.g the a and b scale of Taft and Kamlet [333, 334], are described in Chapter

7 Not unexpectedly, the pKHB scales derived in this way do not correspond to thecommon pKa and pKbscales, i.e to the normal acidity or basicity constants

2.2.6 Electron-Pair Donor/Electron-Pair Acceptor Interactions (EPD/EPA Interactions)[50–59, 59a, 59b]

When tetrachloromethane solutions of yellow chloranil and colourless benzene are mixed, an intensely red solution is formed (lmax¼ 517 nm [50]) This is due

hexamethyl-to the formation of a complex between the two components, and is only one example

of a large number of so-called electron-pair donor/electron-pair acceptor complexes(EPD/EPA complexes)*) It is generally accepted that the characteristic long-wavelength absorptions of these EPD/EPA complexes are associated with an electrontransfer from the donor to the acceptor molecule Mulliken termed these absorptions