Christian Reichardt anddl.booktolearn.com/ebooks2/science/chemistry/...edition called Solvent E¤ects in Organic Chemistry appeared in 1979, followed by a second English edition in

Christian Reichardt and

Thomas Welton

Solvents and Solvent E¤ects in

Organic Chemistry

Solvents and Solvent Effects in Organic Chemistry, Fourth Edition. Edited by Christian Reichardt and Thomas WeltonCopyright 8 2011 WILEY-VCH Verlag GmbH & Co. KGaA, WeinheimISBN: 978-3-527-32473-6

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Christian Reichardt and Thomas Welton

Solvents and Solvent E¤ects in

Organic Chemistry

Fourth, Updated and Enlarged Edition

The Authors

Prof. Dr. Christian Reichardt

Fachbereich Chemie

der Philipps-Universität

Hans-Meerwein-Strasse

35032 Marburg, Germany

Prof. Dr. Thomas Welton

Department of Chemistry

Imperial College London

South Kensington Campus

London SW7 2AZ

United Kingdom

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Previous Editions

1st Edition 1979

2nd Edition 1988

1st Reprint 1990

3rd Edition 2003

1st Reprint 2004

2nd Reprint 2005

4th Edition 2011

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statements, data, illustrations, procedural

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To Mariaand in memory of my parents

C. R.

To Mikeand my parents

T. W.

Preface to the Fourth Edition

About 40 years ago, in 1969, a German paperback entitled Lösungsmittele¤ekte in derorganischen Chemie, written by the first author, was published by Verlag Chemie inWeinheim. Based on this paperback and its second edition in 1973, an enlarged Englishedition called Solvent E¤ects in Organic Chemistry appeared in 1979, followed by asecond English edition in 1988 with the now enlarged title Solvents and Solvent E¤ects inOrganic Chemistry. A first and second reprint in 2004 and 2005 of the third, updatedand enlarged English edition of this book, published in 2003, demonstrate the continu-ing common interest in the study of solvent e¤ects on chemical reactions and physicalprocesses. This and the progress that has been made in recent years in this field of re-search encouraged us to present now to the interested reader a fourth, again updatedand enlarged, edition of this book. This was only possible because a junior authorhelped the meanwhile retired senior author with the preparation of the manuscript forthis new edition, particularly in writing the new Chapter 8.

This new chapter deals with the relationship between solvents and green chemis-try, the classification of solvents by their environmental impact, and the replacement oftraditional by non-traditional solvents for chemical reactions.

During the seven years after publication of the third edition in 2003, the numberof solvent-dependent processes studied has increased to such an extent (particularly inthe field of ionic liquids) that only a careful selection of instructive and representativeexamples could be additionally included in this fourth edition. The literature has beencovered up to 2009, partly up to 2010. New references have been added at the end of thereference list of each chapter.

Consistent use of the nomenclaturea), symbolsb), termsc) and SI unitsd) recom-mended by the respective IUPAC Commissions has again been made in this fourthedition.

For useful comments and valuable suggestions we thank many colleagues, inparticular Prof. Dr. N. O. Mchedlov-Petrossyan, Kharkov/Ukraine, Dr. T. Rager,Basel/Switzerland, and Prof. Dr. G. N. Papatheodorou, Rio/Greece. For their assis-tance in providing us with di‰cult to obtain literature and in preparing the final manu-script, C. R. thanks Mrs. B. Becht-Schröder (librarian) and Mr. G. Schäfer (technician)of the Department of Chemistry, Marburg, and also Mrs. Maria Reichardt, Marburg,


a) G. J. Leigh, H. A. Favre, and W. V. Metanomski: Principles of Chemical Nomenclature –A Guide to IUPAC Recommendations, Blackwell, Oxford, 1998; R. Panico, W. H. Powell, andJ.-C. Richer: A Guide to IUPAC Nomenclature of Organic Compounds – Recommendations 1993,Blackwell, Oxford, 1993.b) E. R. Cohen, T. Cvitaš, J. G. Frey, B. Holmström, K. Kuchitsu, R. Marquardt, I. Mills,F. Pavese, M. Quack, J. Stohner, H. L. Strauss, M. Takami, and A. J. Thor: Quantities, Units andSymbols in Physical Chemistry (IUPAC 2007), 3rd ed., Royal Society of Chemistry, Cambridge,2007.c) P. Müller: Glossary of Terms Used in Physical Organic Chemistry – IUPAC Recommendations1994, Pure Appl. Chem. 66, 1077 (1994).d) G. H. Aylward and T. J. V. Findlay: SI Chemical Data, 6th ed., Wiley, Milton/Australia, 2008;see also Bureau International des Poids et Mesures (BIPM): Le Système International d’Unités(SI), 8th ed., STEDI Media, Paris, 2006.

for her continuous support of this project. T.W. thanks the final-year Imperial CollegeChemistry students and Green Chemistry Master students for helpful discussions.

We both express our thanks to the sta¤ of Wiley-VCH Verlag GmbH, Weinheim,particularly to Dr. Elke Maase and Dr. Stefanie Volk, for their help and excellent workin turning the manuscript into this final book.

Marburg (Lahn) Christian ReichardtLondon Thomas WeltonSummer 2010

VIII Preface to the Fourth Edition

Preface to the Third Edition

Meeting the demand for the second edition of this book, which is – despite a reprint in1990 – 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 areasof 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) recom-mended by the IUPAC commissions has also been made in this third edition.

For comments and valuable suggestions I have to thank many colleagues, in par-ticular Prof. E. M. Kosower, Tel Aviv/Israel, Prof. R. G. Makitra, Lviv/Ukraine, Prof.N. O. Mchedlov-Petrossyan, Kharkiv/Ukraine, and Prof. K. Möckel, Mühlhausen/Germany. For their assistance in drawing formulae, preparing the indices, and provid-ing me with di‰cult to obtain literature, I thank Mr. G. Schäfer (technician), Mrs. S.Schellenberg (secretary), and Mrs. B. Becht-Schröder (librarian), all at the Department


a) G. J. Leigh, H. A. Favre, and W. V. Metanomski: Principles of Chemical Nomenclature – AGuide to IUPAC Recommendations, Blackwell Science Publications, London, 1998.b) I. Mills, T. Cvitaš, K. Homann, N. Kallay, and K. Kuchitsu: Quantities, Units and Symbols inPhysical Chemistry, 2nd ed., Blackwell Science Publications, London, 1993.c) P. Müller: Glossary of Terms used in Physical Organic Chemistry – IUPAC Recommendations1994, Pure Appl. Chem. 66, 1077 (1994).d) G. H. Aylward and T. J. V. Tristan: SI Chemical Data, 4th ed., Wiley, Chichester, 1999;Datensammlung Chemie in SI-Einheiten, 3rd ed., Wiley-VCH, Weinheim/Germany, 1999.

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.

Marburg (Lahn), Spring 2002 Christian Reichardt

X Preface to the Third Edition

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 anda textbook, has obviously made it attractive to both the industrial and academic chemistas well as the advanced student of chemistry.

During the last eight years the study of solvent e¤ects on both chemical reac-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 ledto 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 selectionof the more instructive, recently studied examples could be included in this secondedition.

The search for empirical parameters of solvent polarity and their applicationsin 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 recommendedby 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. Försterling,Marburg, J. Shorter, Hull/England, and R. I. Zalewski, Poznań/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 prepara-tion of this edition and for her help with the indices.

Marburg (Lahn), Spring 1988 Christian Reichardt


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

Preface to the First Edition

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 hasto take into consideration the right reaction partners, the proper reaction vessels, andthe appropriate reaction temperature. One of the most important features for the successof 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 ‘‘Lösungsmittele¤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 isso 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, Zürich/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 improvementof 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.

Marburg (Lahn), Summer 1978 Christian Reichardt

References

[1] C. Reichardt: Lösungsmittele¤ekte in der organischen Chemie. 2nd edition. Verlag Chemie,Weinheim 1973;E¤ets de solvant en chimie organique (translation of the first-mentioned title into French, byI. 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, Lon-don, in 1971.

XIV Preface to the First Edition

Contents

1 Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1

2 Solute-Solvent Interactions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 7

2.1 Solutions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 72.2 Intermolecular Forces . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 122.2.1 Ion-Dipole Forces . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 132.2.2 Dipole-Dipole Forces. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 142.2.3 Dipole-Induced Dipole Forces . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 152.2.4 Instantaneous Dipole-Induced Dipole Forces . . . . . . . . . . . . . . . . . . . . . . . . . 162.2.5 Hydrogen Bonding . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 172.2.6 Electron-Pair Donor/Electron-Pair Acceptor Interactions (EPD/EPA

Interactions) . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 232.2.7 Solvophobic Interactions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 312.3 Solvation . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 342.4 Preferential Solvation. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 432.5 Micellar Solvation (Solubilization) . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 482.6 Ionization and Dissociation . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 52

3 Classification of Solvents . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 65

3.1 Classification of Solvents according to Chemical Constitution . . . . . . . . . 653.2 Classification of Solvents using Physical Constants . . . . . . . . . . . . . . . . . . . . 753.3 Classification of Solvents in Terms of Acid-Base Behaviour. . . . . . . . . . . . 883.3.1 Brønsted-Lowry Theory of Acids and Bases . . . . . . . . . . . . . . . . . . . . . . . . . . 883.3.2 Lewis Theory of Acids and Bases . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 933.4 Classification of Solvents in Terms of Specific Solute/Solvent

Interactions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 963.5 Classification of Solvents using Multivariate Statistical Methods . . . . . . . 99

4 Solvent E¤ects on the Position of Homogeneous Chemical Equilibria . . . . 107

4.1 General Remarks . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1074.2 Solvent E¤ects on Acid/Base Equilibria . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1094.2.1 Brønsted Acids and Bases in Solution . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1094.2.2 Gas-Phase Acidities and Basicities. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1144.3 Solvent E¤ects on Tautomeric Equilibria . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1214.3.1 Solvent E¤ects on Keto/Enol Equilibria . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1214.3.2 Solvent E¤ects on Other Tautomeric Equilibria . . . . . . . . . . . . . . . . . . . . . . . 1284.4 Solvent E¤ects on Other Equilibria. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1364.4.1 Solvent E¤ects on Brønsted Acid/Base Equilibria . . . . . . . . . . . . . . . . . . . . . 1364.4.2 Solvent E¤ects on Lewis Acid/Base Equilibria . . . . . . . . . . . . . . . . . . . . . . . . 1384.4.3 Solvent E¤ects on Conformational Equilibria . . . . . . . . . . . . . . . . . . . . . . . . . 1424.4.4 Solvent E¤ects on cis/trans or E/Z Isomerization Equilibria . . . . . . . . . . . 1484.4.5 Solvent E¤ects on Valence Isomerization Equilibria . . . . . . . . . . . . . . . . . . . 1504.4.6 Solvent E¤ects on Electron-Transfer Equilibria . . . . . . . . . . . . . . . . . . . . . . . 1534.4.7 Solvent E¤ects on Host/Guest Complexation Equilibria . . . . . . . . . . . . . . . 156


5 Solvent E¤ects on the Rates of Homogeneous Chemical Reactions. . . . . . 165

5.1 General Remarks . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1655.2 Gas-Phase Reactivities . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1735.3 Qualitative Theory of Solvent E¤ects on Reaction Rates. . . . . . . . . . . . . . 1805.3.1 The Hughes–Ingold Rules . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1815.3.2 Solvent E¤ects on Dipolar Transition State Reactions . . . . . . . . . . . . . . . . 1925.3.3 Solvent E¤ects on Isopolar Transition State Reactions. . . . . . . . . . . . . . . . 2065.3.4 Solvent E¤ects on Free-Radical Transition State Reactions . . . . . . . . . . . 2205.3.5 Limitations of the Hughes–Ingold Rules . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2355.4 Quantitative Theories of Solvent E¤ects on Reaction Rates . . . . . . . . . . . 2395.4.1 General Remarks . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2395.4.2 Reactions between Neutral, Apolar Molecules . . . . . . . . . . . . . . . . . . . . . . . 2405.4.3 Reactions between Neutral, Dipolar Molecules. . . . . . . . . . . . . . . . . . . . . . . 2465.4.4 Reactions between Neutral Molecules and Ions . . . . . . . . . . . . . . . . . . . . . . 2545.4.5 Reactions between Ions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2555.5 Specific Solvation E¤ects on Reaction Rates . . . . . . . . . . . . . . . . . . . . . . . . . 2595.5.1 Influence of Specific Anion Solvation on the Rates of SN and other

Reactions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2595.5.2 Protic and Dipolar Aprotic Solvent E¤ects on the Rates of SN

Reactions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2655.5.3 Quantitative Separation of Protic and Dipolar Aprotic Solvent E¤ects

for Reaction Rates by Means of Solvent-Transfer Activity Coe‰cients 2775.5.4 Acceleration of Base-Catalysed Reactions in Dipolar Aprotic Solvents 2825.5.5 Influence of Specific Cation Solvation on the Rates of SN Reactions. . . 2855.5.6 Solvent Influence on the Reactivity of Ambident Anions. . . . . . . . . . . . . . 2925.5.7 Solvent E¤ects on Mechanisms and Stereochemistry of Organic

Reactions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2985.5.8 Influence of Micellar and Solvophobic Interactions on Reaction Rates

and Mechanisms . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3175.5.9 Liquid Crystals as Reaction Media . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3265.5.10 Solvent Cage E¤ects . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3315.5.11 External Pressure and Solvent E¤ects on Reaction Rates . . . . . . . . . . . . . 3365.5.12 Solvent Isotope E¤ects . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3435.5.13 Reactions in Biphasic Solvent Systems and in Neoteric Solvents. . . . . . . 345

6 Solvent E¤ects on the Absorption Spectra of Organic Compounds . . . . . . 359

6.1 General Remarks . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3596.2 Solvent E¤ects on UV/Vis Spectra . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3606.2.1 Solvatochromic Compounds . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3606.2.2 Theory of Solvent E¤ects on UV/Vis Absorption Spectra . . . . . . . . . . . . . 3716.2.3 Specific Solvent E¤ects on UV/Vis Absorption Spectra . . . . . . . . . . . . . . . 3806.2.4 Solvent E¤ects on Fluorescence Spectra. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3846.2.5 Solvent E¤ects on ORD and CD Spectra . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3936.3 Solvent E¤ects on Infrared Spectra . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3976.4 Solvent E¤ects on Electron Spin Resonance Spectra . . . . . . . . . . . . . . . . . . 4036.5 Solvent E¤ects on Nuclear Magnetic Resonance Spectra. . . . . . . . . . . . . . 410

XVI Contents

6.5.1 Nonspecific Solvent E¤ects on NMR Chemical Shifts . . . . . . . . . . . . . . . . . 4106.5.2 Specific Solvent E¤ects on NMR Chemical Shifts . . . . . . . . . . . . . . . . . . . . . 4176.5.3 Solvent E¤ects on Spin-Spin Coupling Constants . . . . . . . . . . . . . . . . . . . . . 422

7 Empirical Parameters of Solvent Polarity . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 425

7.1 Linear Gibbs Energy Relationships. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 4257.2 Empirical Parameters of Solvent Polarity from Equilibrium

Measurements. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 4327.3 Empirical Parameters of Solvent Polarity from Kinetic Measurements . 4387.4 Empirical Parameters of Solvent Polarity from Spectroscopic

Measurements. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 4487.5 Empirical Parameters of Solvent Polarity from Other Measurements . . . 4817.6 Interrelation and Application of Solvent Polarity Parameters . . . . . . . . . . 4837.7 Multiparameter Approaches . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 490

8 Solvents and Green Chemistry . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 509

8.1 Green Chemistry . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 5098.2 Reduction of Solvent Use. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 5118.3 Green Solvent Selection . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 5138.4 Non-Traditional Solvents . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 5148.4.1 Water . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 5148.4.2 Supercritical Carbon Dioxide (sc-CO2) . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 5298.4.3 Ionic Liquids. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 5348.4.4 Polyethylene Glycols (PEGs). . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 5438.4.5 Biomass-Derived Solvents . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 5448.5 Outlook . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 548

Appendix . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 549

A. Properties, Purification, and Use of Organic Solvents. . . . . . . . . . . . . . . . . . 549A.1 Physical Properties . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 549A.2 Purification of Organic Solvents. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 556A.3 Spectroscopic Solvents. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 557A.4 Solvents as Reaction Media. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 562A.5 Solvents for Recrystallization . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 563A.6 Solvents for Extraction and Partitioning (Distribution) . . . . . . . . . . . . . . . . 570A.7 Solvents for Adsorption Chromatography . . . . . . . . . . . . . . . . . . . . . . . . . . . . 572A.8 Solvents for Acid/Base Titrations in Non-Aqueous Media . . . . . . . . . . . . . 574A.9 Solvents for Electrochemistry . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 578A.10 Toxicity of Organic Solvents . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 578

References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 587

Figure and Table Credits . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 675

Subject Index . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 677

XVIIContents

List of Abbreviations

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

NA Avogadro constant 6:0221 � 1023 mol�1c0 speed of light in vacuum 2:9979 � 108 m � s�1e0 permittivity of vacuum

[¼ 1=ðm0 � c02Þ; m0 ¼ permeability ofvacuum]

8:8542 � 10�12 F � m�1

e elementary charge 1:6022 � 10�19 Ch Planck constant 6:6261 � 10�34 J � sR molar gas constant 8.3145 J � K�1 � mol�1

(or 0.08206L � atm � K�1 � mol�1)

kB Boltzmann constant (¼ R=NA) 1:3807 � 10�23 J � K�1Vm standard molar volume of an ideal

gas (at t ¼ 0 �C and p ¼ 100 kPa)22.711 L � mol�1

T0 zero of the Celsius scale 273.15 K

p ratio of the circumference to thediameter of a circle

3.1416

e exponential number and base ofnatural logarithms (ln)

2.7183

ln 10 natural logarithm of ten (ln x ¼ ln10 � lg x; lg ¼ decadic logarithm)

2.303

Abbreviations and Symbols for Unitsa,b)

bar bar (¼ 105 Pa ¼ 105 N � m�2) pressurecg/g centigram/gram weight percent

cL/L, cl/l centilitre/litre volume percent

cmol/mol centimol/mol mole percent

cm centimetre (10�2 m) length


a) E. R. Cohen, T. Cvitaš, J. G. Frey, B. Holmström, K. Kuchitsu, R. Marquardt, I. Mills,F. Pavese, M. Quack, J. Stohner, H. L. Strauss, M. Takami, and A. J. Thor: Quantities, Units andSymbols in Physical Chemistry (IUPAC 2007), 3rd ed., Royal Society of Chemistry, Cambridge,2007.b) G. H. Aylward and T. J. V. Findlay: SI Chemical Data, 6th ed., Wiley, Milton/Australia, 2008;see also Bureau International des Poids et Mesures (BIPM): Le Système International d’Unités(SI), 8th ed., STEDI, Paris, 2006.

cm3 cubic centimetre(millilitre mL; 10�6 m3)

volume

Q coulomb electric charge�C degrees Celsius temperaturedm3 cubic decimetre (litre L; 10�3 m3) volumeJ joule energy

kJ kilojoule (103 J) energy

K kelvin temperature

L, l litre (1 dm3; 10�3 m3) volumem metre length

min minute time

mol mole amount of substance

MPa megapascal (106 Pa) pressure

mT millitesla (10�3 T) magnetic flux density(magnetic field)

nm nanometre (10�9 m) lengthPa pascal (1 N � m�2 ¼ 10�5 bar) pressurepercent (%) part per hundred (10�2) dimensionless fractionppm part per million (10�6) dimensionless fractions second time

Abbreviations and Symbols for Propertiesc)

ai activity of solute i

að1HÞ ESR hyperfine coupling constant(coupling with 1H)

Hz or mT (¼ 10�3 T)

Aj the solvent’s anion-solvating tendencyor ‘acity’ (Swain)

AN solvent acceptor number, based on31P NMR chemical shift of Et3PO(Gutmann and Meyer)

a electric polarizability of a moleculeor polarizability volume

C2 � m2 � J�1 or 4pe0 � cm3

a empirical parameter of solventhydrogen-bond donor acidity (Taftand Kamlet)

B empirical parameter of solvent Lewisbasicity (Palm and Koppel)

c) P. Müller: Glossary of Terms used in Physical Organic Chemistry – IUPAC Recommendations1994. Pure Appl. Chem. 66, 1077 (1994).

XX List of Abbreviations

BMeOD IR based empirical parameter ofsolvent Lewis basicity (Palm andKoppel)

BPhOH IR based empirical parameter ofsolvent Lewis basicity (Koppel andPaju; Makitra)

Bj the solvent’s cation-solvatingtendency or ‘basity’ (Swain)

b empirical parameter of solventhydrogen-bond acceptor basicity(Taft and Kamlet)

c cohesive pressure (cohesive energydensity) of a solvent

MPa (¼ 106 Pa)

ci; cðiÞ molar concentration of solute i mol � L�1CA;CB Lewis acidity and Lewis basicity

parameter (Drago)

cmc critical micellisation concentration mol � L�1DHA molar bond-dissociation energy of the

bond between H and AkJ � mol�1

Dp empirical parameter of solvent Lewisbasicity, based on a 1,3-dipolarcycloaddition reaction (Nagai et al.)

DN solvent donor number (Gutmann)[¼ �DH(DaaSbCl5)]

kcal � mol�1

DNN normalized solvent donor number(Marcus)

d; dH Hildebrand’s solubility parameter MPa1=2

d chemical shift of NMR signals ppm

d solvent polarizability correction term(Taft and Kamlet)

E energy, molar energy kJ � mol�1E electric field strength V � m�1E enol constant (K. H. Meyer)

E empirical parameter of solvent Lewisacidity (Palm and Koppel)

EA;Ea Arrhenius activation energy kJ � mol�1EA;EB Lewis acidity and Lewis basicity

parameter (Drago)

EA electron a‰nity kJ � mol�1ENB empirical solvent Lewis basicity

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

XXIList of Abbreviations

EK empirical solvent polarity parameter,based on the d ! p� absorption of amolybdenum complex (Walther)

kcal � mol�1

E �MLCT empirical solvent polarity parameter,based on the d ! p� absorption of atungsten complex (Lees)

ET molar electronic transition energy,molar electronic excitation energy

kJ � mol�1 or kcal � mol�1

ETð30Þ empirical solvent polarity parameter,based on the intramolecular CTabsorption of a pyridinium-N-phenolate betaine dye (Dimroth andReichardt)

kcal � mol�1

ENT normalized ETð30Þ solvent polarityparameter (Reichardt)

E SOT empirical solvent polarity parameter,based on the n ! p� absorption of anS-oxide (Walter)

kcal � mol�1

EPA electron-pair acceptor

EPD electron-pair donor

er relative permittivity (¼ e=e0)(‘‘dielectric constant’’)

1

F empirical solvent polarity parameter,based on the n ! p� absorption ofketones (Dubois)

G IR based empirical solvent polarityparameter (Schleyer and Allerhand)

DG� standard molar Gibbs energy change kJ � mol�1DG0 standard molar Gibbs energy of

activationkJ � mol�1

DG�solv standard molar Gibbs energy ofsolvation

kJ � mol�1

DG�hydr standard molar Gibbs energy ofhydration

kJ � mol�1

DG�t ðX;O!SÞ,DG�t ðX;W!SÞ

standard molar Gibbs energy oftransfer of solute X from a referencesolvent (O) or water (W) to anothersolvent (S)

kJ � mol�1

gi activity coe‰cient of solute i

DH � standard molar enthalpy change kJ � mol�1DH0 standard molar enthalpy of activation kJ � mol�1DHv molar enthalpy (heat) of vaporization kJ � mol�1

XXII List of Abbreviations

H0 acidity function (Hammett)

HBA hydrogen-bond acceptor

HBD hydrogen-bond donor

HOMO highest occupied molecular orbital

Ei; I ; IP ionization energy kJ � mol�1I, Ic ionic strength (concentration basis)

(¼ 1/2 � Pci � zi2)mol � L�1

I gas-chromatographic retention index(Kováts)

J NMR spin-spin coupling constant Hz

k rate constant for monomolecular(n ¼ 1) and bimolecular (n ¼ 2)reactions

(L � mol�1)n�1 � s�1

k0 rate constant in a reference solvent orin the gas phase for monomolecular(n ¼ 1) and bimolecular reactions(n ¼ 2)

(L � mol�1)n�1 � s�1

k0 in Hammett equations the rateconstant of unsubstituted substrates

(L � mol�1)n�1 � s�1 withn ¼ 1 or 2

K ;Kc equilibrium constant (concentrationbasis; v ¼ stoichiometric number)

(mol � L�1)Sv

Ka;Kb acid and base ionization constants (mol � L�1)SvKauto autoionization ion product,

autoprotolysis constantmol2 � L�2

KAssoc;KDissoc,Kion;KT

equilibrium constants of association,dissociation, ionization, resp.tautomerization reactions

(mol � L�1)Sv

Ko=w 1-octanol/water partition constant(Hansch and Leo)

KB kauri-butanol number

L desmotropic constant (K. H. Meyer)

LUMO lowest unoccupied molecular orbital

l wavelength nm (¼ 10�9 m)m mass of a particle g

Mr relative molecular mass of a substance(‘‘molecular weight’’)

M miscibility number (Godfrey)

MH microscopic hydrophobicityparameter of substituents (Menger)

m empirical solvent softness parameter(Marcus)

XXIIIList of Abbreviations

m permanent electric dipole moment ofa molecule

C � m (or D)

mind induced electric dipole moment of amolecule

C � m (or D)

m�i standard chemical potential of solute i kJ � mol�1myi standard chemical potential of solute i

at infinite dilutionkJ � mol�1

n; nD refractive index (at sodium D line)(¼ c0=c)

N empirical parameter of solventnucleophilicity (Winstein andGrunwald)

Nþ nucleophilicity parameter for(nucleophile þ solvent)-systems(Ritchie)

n frequency Hz, s�1

n� frequency in the gas phase or in aninert reference solvent

Hz, s�1

~nn wavenumber (¼ 1=l) cm�1W empirical solvent polarity parameter,

based on a Diels-Alder reaction(Berson)

p pressure Pa (¼ 1N � m�2),bar (¼ 105 Pa)

P measure of solvent polarizability(Palm and Koppel)

P empirical solvent polarity parameter,based on 19F NMR measurements(Taft)

PA proton a‰nity kJ � mol�1Py empirical solvent polarity parameter,

based on the p� ! p emission ofpyrene (Winnik)

Po=w 1-octanol/water partition coe‰cient(Hansch and Leo)

pH �lg c(H3Oþ), or more precisely�lg a(H3Oþ)abbreviation of potentia or pondushydrogenii, power of hydrogen, orpuissance d’hydrogène (Sørensen1909). The pH scale ranges usuallyfrom 1 to 14, but is open-ended,allowing for pH values below 0 orabove 14!

XXIV List of Abbreviations

pK �lg Kp internal pressure of a solvent MPa (¼ 106 Pa)p� empirical solvent dipolarity/

polarizability parameter, basedon the p ! p� absorption ofsubstituted aromatics (Taft andKamlet)

p�azo empirical solvent dipolarity/polarizability parameter, based on thep ! p� absorption of azomerocyanine dyes (Buncel)

px hydrophobicity parameter ofsubstituent X in H5C6-X (Hansch)

r radius of sphere representing an ionor a cavity

cm (¼ 10�2 m)

r distance between centres of two ionsor molecules

cm (¼ 10�2 m)

r density (mass divided by volume) g � cm�3r; rA Hammett reaction or absorption

constants

S generalized for solvent

S empirical solvent polarity parameter,based on the Z-values (Brownstein)

S lg k2 for the Menschutkin reaction oftri-n-propylamine with iodomethane(Drougard and Decroocq)

DS� standard molar entropy change J � K�1 � mol�1DS0 standard molar entropy of activation J � K�1 � mol�1Sp solvophobic power of a solvent

(Abraham)

SA empirical parameter of solventhydrogen-bond donor acidity(Catalán)

SB empirical parameter of solventhydrogen-bond acceptor basicity(Catalán)

SPP empirical parameter of solventdipolarity/polarizability, based on thep ! p� absorption of substituted 7-nitrofluorenes (Catalán)

s Hammett substituent constant

s NMR screening constant

XXVList of Abbreviations

t Celsius temperature �CT thermodynamic temperature K

tmp melting point�C

tbp boiling point�C

U internal energy kJ

DUv molar energy of vaporization kJ � mol�1Vm;Vm; i molar volume (of i) cm

3 � mol�1DV0 molar volume of activation cm3 � mol�1xi; xðiÞ mole fraction of i ðxi ¼ ni=

PnÞ

amount-of-substance fraction1

X empirical solvent polarity parameter,based on an SE2 reaction (Gielen andNasielski)

wR; wB empirical solvent polarity parameters,based on the p ! p� absorption ofmerocyanine dyes (Brooker)

kcal � mol�1

OySX;WySX solvent-transfer activity coe‰cient for

the transfer of a solute X from areference solvent (O) or water (W) toanother solvent (S)

Y empirical parameter of solventionizing power, based on t-butylchloride solvolysis (Winstein andGrunwald)

YOTs empirical parameter of solventionizing power, based on 2-adamantyltosylate solvolysis (Schleyer andBentley)

Y measure of solvent polarization (Palmand Koppel)

zi charge number of an ion i(zi ¼Qi/e)

positive for cations,negative for anions

Z empirical solvent polarity parameter,based on the intermolecular CTabsorption of a substitutedpyridinium iodide (Kosower)

kcal � mol�1

XXVI List of Abbreviations

‘‘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 certe

Chemistae principem locum prae omnibus assignant, in quo se jactant, serioque tri-

umphant, cui artis suae, prae aliis omnibus e¤ectus mirificos adscribunt. Atque illud

quidem 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 developmentof 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 Greek philosopher Thales ofMiletus (ca. 640–546 bc) asserted that water is the origin out of which everything aroseand into everything resolved itself.

From these ancient times, a familiar and today often cited quotation of the fa-mous Greek philosopher Aristotle (384–322 bc) was handed down, which reads in LatinCorpora non agunt nisi fluida (or liquida) seu soluta, and was translated into English as‘‘Compounds do not react unless fluid or if dissolved’’ [43]. However, according toHedvall [44], this seems to be a misinterpretation of the original text given in Greek asTá ńgrá miktá málista ton somáton (Ta hygra mikta malista ton somaton), which isprobably taken from Aristotle’s work De Generatione et Corruptione [45]. According toHedvall, this statement should be better read as „ . . . it is chiefly the liquid substanceswhich react’’ [44] or „ . . . for instance liquids are the type of bodies most liable to mix-ing’’ [45c]. In this somewhat softened version, Aristotle’s statement is obviously lessdistinct and didactic. With respect of the many solid/solid reactions known today, it isquite understandable that solid-state chemists were not very happy with the commonfirst version of Aristotle’s statement [43, 44].

The alchemist’s search for a universal solvent, the so-called ‘‘Alkahest’’ or ‘‘Men-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 of


* ‘‘Well then, my dear listeners, let us proceed with fervor to another problem! Having su‰cientlyanalyzed 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 themost essential part of chemistry itself, which chemists boastfully, no doubt with reason, preferabove all others, and because of which they triumphantly celebrate, and to which they attributeabove all others the marvellous e¤ects of their science. And this they call the solvent (menstruum).’’

the chemists of the 15th to 18th centuries did not in fact lead to the discovery of any‘‘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 dissolutionof 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 dissolutionof 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 per-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].

The di‰culty in explaining the e¤ects of inorganic solutes on the physical prop-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].

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 the osmotic pressure and thetheory of electrolytic dissociation in dilute solutions, respectively.

The further development of solution chemistry is connected with the pioneeringwork of Ostwald (1853–1932), Nernst (1864–1941), Lewis (1875–1946), Debye (1884–

* 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 hiscontemporaries who asked in which vessel he has stored this universal solvent.

2 1 Introduction

1966), E. Hückel (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 notedby Berthelot and Péan de Saint-Gilles in 1862 in connection with their studies on theesterification of acetic acid with ethanol: ‘‘. . . l’éthérification est entravée et ralentie parl’emploi des dissolvants neutres étrangers à la réaction’’ [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 letterto Prof. Louis Henry he wrote in 1890: ‘‘Or, l’expérience montre que ces dissolvantsexercent sur la vitesse de combinaison une influence considérable. Si nous représentonspar 1 la constante de vitesse de la réaction précitée dans l’hexane C6H14, cette constantepour la même combinaison dans CH3aaCOaaC6H5, toutes choses égales d’ailleurs sera847.7. La di¤érence est énorme, mais, dans ce cas, elle n’atteint pas encore le maxi-mum. . . . Vous voyez que les dissolvants, soi-disant indi¤érents ne sont pas inertes; ilsmodifient profondément l’acte de la combinaison chimique. Cet énoncé est riche enconséquences pour la théorie 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 unfavourable 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]: acetyl-dibenzoylmethane 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]: phenyl-nitromethane). Thus, Claisen wrote: ‘‘Es gibt

Verbindungen, welche sowohl in der Form aaC(OH)bbC

aa

aaCOaa wie in der Form

aaCOaaC

aa

HaaCOaa zu bestehen vermögen; von der Natur der angelagerten Reste, von

* ‘‘. . . the esterification is disturbed and decelerated on addition of neutral solvents not belongingto the reaction’’ [9].** ‘‘Now, experience shows that solvents exert considerable influence on reaction rates. If we rep-resent the rate constant of the reaction to be studied in hexane C6H14 by 1, then, all else beingequal, this constant for the same reaction in CH3aaCOaaC6H5 will be 847.7. The increase is enor-mous, but in this case it has not even reached its maximum. . . . So you see that solvents, in spite ofappearing at first to be indi¤erent, are by no means inert; they can greatly influence the course ofchemical 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 givesthe same product as that obtained from the coupling reaction of 1-naphthol with benzenediazoniumsalts. This phenomenon, that the substrate can react either as phenylhydrazone or as a hydroxyazocompound, depending on the reaction circumstances, was called Ortsisomerie by Zincke [11]. Lateron, the name tautomerism, with a di¤erent meaning however from that accepted today, wasintroduced by Laar [12]. For a description of the development of the concept of tautomerism, seeIngold [13].

31 Introduction

der Temperatur, bei den gelösten Substanzen auch von der Art des Lösungsmittels hängtes ab, welche von den beiden Formen die beständigere ist’’ [14]*). The study of the keto-enol equilibrium of ethyl formylphenylacetate in eight solvents led Wislicenus to theconclusion that the keto form predominates in alcoholic solution, the enol form in tri-chloromethane or benzene. He stated that the final ratio in which the two tautomericforms coexist 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 solventsinto two groups according to their ability to isomerize tautomeric compounds. His clas-sification reflects, to some extent, the modern division into protic and aprotic solvents.The e¤ect of solvent on constitutional and tautomeric isomerization equilibria waslater studied in detail by Dimroth [19] (using triazole derivatives, e.g. 5-amino-4-methoxycarbonyl-1-phenyl-1,2,3-triazole) and Meyer [20] (using ethyl acetoacetate).

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 basisof 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 resultof electrical fields. These fields in turn originate from the dipolar properties of the mol-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 (neoteric ¼ recent, new, modern), constitut-ing a class of novel solvents with desirable, less hazardous, new properties [35, 36]. The

* ‘‘There are compounds capable of existence in the form aaC(OH)bbC

aa

aaCOaa as well as in the

form aaCOaaC

aa

HaaCOaa; it depends on the nature of the substituents, the temperature, and fordissolved compounds, also on the nature of the solvent, which of the two forms will be the morestable’’ [14].** A survey of older works of solvent e¤ects on UV/Vis absorption spectra has been given bySheppard [21].*** It should be noted that the now generally accepted meaning of the term solvatochromism dif-fers from that introduced by Hantzsch (cf. Section 6.2).

4 1 Introduction

term neoteric solvents covers supercritical fluids, ionic liquids, and also perfluorohydro-carbons (as used in fluorous biphasic systems). In addition, water, often consideredincompatible with organic synthesis, in recent decades has attracted increasing interestas an environmentally benign and cheap solvent for a multitude of organic reactions[46]. Table A-14 in Chapter A.10 (Appendix) collects some recommendations for thesubstitution of hazardous solvents, together with the relevant literature references; seealso Chapter 8.

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]. A representative recent example of a highly e‰cient solvent-free organic synthesis is the (S)-proline-catalysed stereoselective aldol reaction betweencyclohexanone and 4-nitrobenzaldehyde, applying a very simple mechano-chemicaltechnique such as ball milling [42].

However, not all organic reactions can be carried out in the absence of a solvent;some organic reactions even proceed explosively in the solid state! Therefore, solventswill still be useful in mediating and moderating chemical reactions and this book onsolvent e¤ects will certainly not become superfluous in the foreseeable future.

51 Introduction

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 (g ¼ 1) [1, 171]. Solute/solvent mixtures Aþ 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 ofAðxAÞ in the liquid mixture, i.e. xA ¼ 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 onlyby its macroscopic physical constants such as boiling point, vapour pressure, density,cohesive pressure, index of refraction, relative permittivity, thermal conductivity, surfacetension, etc. From the molecular-microscopic point of view, a solvent is a discontinuumwhich consists of individual, mutually interacting solvent molecules, characterized bytheir molecular properties such as dipole moment, electronic polarizability, hydrogen-bond donor (HBD) and hydrogen-bond acceptor (HBA) capability, electron-pair donor(EPD) and electron-pair acceptor (EPA) capability, etc. According to the extent of theseintermolecular solvent/solvent interactions, there are highly structured solvents (e.g.,water with its strong directional hydrogen bonds, forming an intermolecular networkwith cavities) and less structured solvents (e.g., hydrocarbons with their weak nondirec-tional dispersion forces, filling the available space in a more regular manner) [173].

The interactions between species in solvents (and in solutions) are at once toostrong to be treated by the laws of the kinetic theory of gases, yet too weak to be treatedby the laws of solid-state physics. Thus, the solvent is neither an indi¤erent medium inwhich the dissolved material di¤uses in order to distribute itself evenly and randomly,nor does it possess an ordered structure resembling a crystal lattice. Nevertheless, thelong-distance ordering in a crystal corresponds somewhat to the local ordering in a liq-uid. Thus, neither of the two possible models – the gas and crystal models – can be ap-plied to solutions without limitation. There is such a wide gulf between the two modelsin terms of conceivable and experimentally established variants, that it is too di‰cult todevelop a generally valid model for liquids. Due to the complexity of the interactions,the structure of liquids – in contrast to that of gases and solids – is the least-known ofall aggregation states. Therefore, the experimental and theoretical examination of the


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

structure of liquids is among the most di‰cult tasks of physical 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 thegaseous 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 replacedby 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 interpretationof its anomalies [15, 304]. Fig. 2-1 should make clear the complexity of the inner struc-ture 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, 15]. The currently accepted view of the structure of liquid water treats it asa dynamic three-dimensional hydrogen-bonded network, without a significant numberof non-bonded water molecules, that retains several of the structural characteristics ofice (i.e. tetrahedral molecular packing with each water molecule hydrogen-bondedto four nearest neighbours), although the strict tetrahedrality is lost. Its dynamic be-haviour 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 [15, 17, 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]. The inner structure of self-associated methanol isdominated by hydrogen-bonded ring clusters (preferentially hexamers; no chains) inwhich the monomers participate as both H-bond donor and acceptor [413].

* The amusing story of ‘‘polywater,’’ which excited the scientific community for a few years duringthe late 1960’s and early 1970’s, has been reviewed by Franks [175]. It turned out that polywaterwas not a new and more stable form of pure water, but merely dirty water. The strange propertiesof polywater were due to high concentrations of siliceous material dissolved from quartz capillariesin which it was produced.

8 2 Solute-Solvent Interactions

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

92.1 Solutions

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 chemicalsimilarity 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 at25 �C, 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 (s ¼ 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 inter-action between solvent and solute molecules that determines the mutual solubility. Acompound A dissolves in a solvent B only when the intermolecular forces of attractionKAA 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 canbe related to the so-called polarity*) of A and B. Denoting compounds with large inter-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).

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

The solubilities of ethane and methane are higher in nonpolar tetrachloro-methane, whereas the opposite is true for chloromethane and dimethyl ether. A surveyof the reciprocal miscibility of some representative examples of organic solvents is givenin Fig. 2-2.

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

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


Table 2-2. Solubilities of methane, ethane, chloromethane, and dimethyl ether intetrachloromethane (nonpolar solvent) and acetone (polar solvent) [22].

Solute Solute polarity Solubility/(mol � m�3) at 25 �C

in CCl4 in CH3COCH3

CH4 nonpolar 29 25CH3CH3 nonpolar 220 130CH3Cl polar 1700 2800CH3OCH3 polar 1900 2200

Fig. 2-2. Miscibility of organic solvents [23]. miscible in all proportions; – – – – limitedmiscibility; . . . . . . . little miscibility; without line: immiscible.

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

Solute A Solvent B Interaction

A � � �A B � � �B A � � �B

Solubility ofA in B

Nonpolar nonpolar weak weak weak can be higha)Nonpolar polar weak strong weak probably lowb)Polar nonpolar strong weak weak probably lowc)Polar polar strong strong strong can be higha)

a) Not much change for solute or solvent.b) Di‰cult to break up B � � �B.c) Di‰cult to break up A � � �A.

112.1 Solutions

a deeper and more detailed understanding of the diverse rules determining the solubilityof organic compounds in various solvents, see references [312–316].

The solubility parameter d of Hildebrand [4, 24], as defined in Eq. (2-1), can oftenbe used in estimating the solubility of non-electrolytes in organic solvents.

d ¼ DUvVm

� �1=2¼ DHv � R � T

Vm

� �1=2ð2-1Þ

In this equation, Vm is the molar volume of the solvent, and DUv and DHv are themolar energy and the molar enthalpy (heat) of vaporization to a gas of zero pressure,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 (d ¼ 0 for the gasphase). As a rule, it has been found that a good solvent for a certain non-electrolyte hasa d value close to that of the solute [20, 24, 25]; cf. Table 3-3 in Section 3.2 for a collec-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.

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, silicone oil, water, aniline, perfluoro(dimethylcyclohexane), whitephosphorus, gallium, and mercury. This system is stable indefinitely at 45 �C; 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 sili-cone oil, 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. Even mixtures of some hydrophilic ionic liquids andsome specific hydrophobic ionic liquids can be immiscible and give rise to two phases. Astable tetraphasic solvent mixture containing two immiscible ionic liquids is formed by(from top to bottom) n-pentane, tri(n-hexyl)-n-tetradecylphosphonium chloride, water,and 1-ethyl-3-methylimidazolium bis(trifluoromethanesulfonyl)amide [414].

The complexity of multi-phasic liquid systems can be nicely demonstrated withthe following experiment, called ‘‘James-Bond-Cocktail’’ [452]: starting from n-per-fluoroheptane, a saturated aqueous potassium carbonate solution, methanol, and tol-uene (coloured with suitable dyes), a three-phase solvent system is obtained, which aftershaking (not stirring!) develops four layers; addition of n-pentane changes the sequenceof the two upper layers.

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

Intermolecular forces are those which can occur between closed-shell molecules [26, 27].These noncovalent interactions are also called van der Waals forces, since the Dutchscientist J. D. van der Waals recognized them (1873, in his PhD thesis) as the reason for


the non-ideal behaviour of real gases. Intermolecular forces are usually classified intotwo distinct categories. The first category comprises the so-called directional, induction,and dispersion forces, which are non-specific and cannot be completely saturated ( justas Coulomb forces between ions cannot). The second group consists of hydrogen-bonding forces, and charge-transfer or electron-pair donor–acceptor forces. The lattergroup are specific, directional forces, which can be saturated and lead to stoichiometricmolecular compounds. For the sake of completeness, in the following the Coulombforces between ions and electrically neutral molecules (with permanent dipole moments)will be considered first, even though they do not belong to the intermolecular forces inthe narrower sense.

2.2.1 Ion-Dipole Forces [28, 185]

Electrically neutral molecules with an unsymmetrical charge distribution possess a per-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 m ¼ 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

U ion-dipole ¼ � 14p � e0 �

z � e � m � cos yr2

ð2-2Þ*)

where e0 is the permittivity of a vacuum, z � e the charge on the ion, r the distance fromthe ion to the center of the dipole, and y the dipole angle relative to the line r joiningthe ion and the centre of the dipole. Cos y ¼ 1 for y ¼ 0�, i.e. in this case the dipoleis 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 l ¼ 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 (z ¼ þ1, radius ¼ 95 pm) near a water molecule(radiusA140 pm; m ¼ 5:9 � 10�30 Cm), the maximum interaction energy calculated byEq. (2-2) amounts to U ¼ 39k � T or 96 kJ � mol�1 at 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 � 10�30 Cm (i.e. Coulombmeter). Among the solvents listed in theAppendix, Table A-1, hexamethylphosphoric triamide is the one with the highest dipole

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

132.2 Intermolecular Forces

moment (m ¼ 18:48 � 10�30 Cm), followed by propylene carbonate (m ¼ 16:7 � 10�30Cm), and sulfolane (m ¼ 16:05 � 10�30 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(tbp ¼ 155 �C/3 Torr) with a large relative permittivity (er ¼ 64:6 at 25 �C), has a dipolemoment of m ¼ 35:7 � 10�30 Cm (¼ 10.7 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 sol-vents, where solvated species such as Na(OH2)

lm and Cl(H2O)

mn (for solutions of NaCl

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]

3l or 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 possessinga permanent dipole moment m due to their unsymmetrical charge distribution. Whentwo dipolar molecules are optimally oriented with respect to one another at a distance ras 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.

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 favoured leads to a net attraction, which is strongly temperature dependent,according to Eq. (2-3) (kB ¼ Boltzmann constant; T ¼ absolute temperature) [29].

Udipole-dipole ¼ � 1ð4p � e0Þ2� 2m

21 � m22

3kB � T � r6 ð2-3Þ

As the temperature increases, the angle-averaged dipole/dipole interaction energybecomes less negative until at very high temperatures all dipole orientations are equally

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


populated and the potential energy is zero. This Boltzmann-averaged dipole/dipoleinteraction is usually referred to as the orientation or Keesom interaction [29]. Accordingto Eq. (2-3), for pairs of dipolar molecules with m ¼ 3:3 � 10�30 Cm (¼ 1 D), at a sepa-ration of 500 pm, the average interaction energy is about �0.07 kJ � mol�1 at 25 �C.This is clearly smaller than the average molar kinetic energy of 3/2 k � T ¼ 3:7kJ � mol�1 at 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 (n ¼ 1) is a pointcharge and a monopole moment represents an overall charge (e.g. of an ion Naþ orCl�). A dipole (n ¼ 2; e.g. H2O, H3CaaCOaaCH3) is an array of partial charges withno monopole moment (i.e. no charge). A quadrupolar molecule (n ¼ 4; e.g. CO2, C6H6)has neither a net charge nor a dipole moment, and an octupolar molecule (n ¼ 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 thechair-configured, non-dipolar solvent 1,4-dioxane, which often behaves like a polar sol-vent even though its relative permittivity is low (er ¼ 2:2), is caused by its large nonidealquadrupolar charge distribution [411].

For the calculation of quadrupole moments for water and some organic solvents,which are experimentally not available, see reference [415].

2.2.3 Dipole-Induced Dipole Forces [32]

The electric dipole of a molecule possessing a permanent dipole moment m can inducea 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 electronic polarizability a of the apolar molecule experiencing the induction ofthe permanent dipole. The net dipole/induced dipole energy of interaction for two dif-ferent molecules, each possessing a permanent dipole mo

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