front matter'. in: solvents and solvent effects in ... · organic synthesis workbook ii 2001....

Christian Reichardt

Solvents and

Solvent Eects in

Organic Chemistry

Solvents and Solvent Effects in Organic Chemistry, Third Edition. Christian ReichardtCopyright 8 2003 WILEY-VCH Verlag GmbH & Co. KGaA, WeinheimISBN: 3-527-30618-8

Related Titles from WILEY-VCH

Bittner, C. et al.

Organic Synthesis Workbook II2001. 3-527-30415-0

Jessop, P. G., Leitner, W. (Eds.)

Chemical Synthesis Using Supercritical Fluids1999. ISBN 3-527-29605-0

Wasserscheid, P., Welton, T. (Eds.)

Ionic Liquids in Synthesis2002. ISBN 3-527-30515-7

Drauz, K., Waldmann, H. (Eds)

Enzyme Catalysis in Organic Synthesis2002. ISBN 3-527-29949-1

Christian Reichardt

Solvents andSolvent Eects inOrganic ChemistryThird, Updated and Enlarged Edition

Prof. Dr. Christian ReichardtFachbereich Chemieder Philipps-Universitat MarburgHans-Meerwein-Strae35032 MarburgGermanye-mail: [email protected]

This book was carefully produced. Nevertheless, author and publisher do not warrant the infor-mation contained therein to be free of errors. Readers are advised to keep in mind that state-ments, data, illustrations, procedural details or other items may inadvertently be inaccurate.

First Reprint 2004

Library of Congress Card No.: applied for

A catalogue record for this book is available from the British Library.

Bibliographic information published by Die Deutsche BibliothekDie Deutsche Bibliothek lists this publication in the Deutsche Nationalbibliografie; detailed biblio-graphic data is available in the Internet at http://dnb.ddb.de.

ISBN 3-527-30618-8

6 2003 WILEY-VCH Verlag GmbH & Co. KGaA, WeinheimPrinted on acid-free paper.All rights reserved (including those of translation in other languages). No part of this book may bereproduced in any form by photoprinting, microfilm, or any other means nor transmitted ortranslated into machine language without written permission from the publishers. Registerednames, trademarks, etc. used in this book, even when not specifically marked as such, are not to beconsidered unprotected by law.

Composition: Asco Typesetters, Hong Kong. Printing: Strauss Osetdruck GmbH, MorlenbachBookbinding: J. Schaer GmbH & Co. KG, GrunstadtPrinted in the Federal Republic of Germany.

To Mariaand 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 in1990 no longer available, and considering the progress that has been made during thelast decade in the study of solvent eects 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 eects 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 eects 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 eects 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 solventeects.

Great progress has been made during the last decade in theoretical treatments ofsolvent eects 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 dierential 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. Mockel, Muhlhausen/Germany. For their assistance in drawing formulae, preparing the indices, and provid-ing me with dicult to obtain literature, I thank Mr. G. Schafer (technician), Mrs. S.Schellenberg (secretary), and Mrs. B. Becht-Schroder (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. Cvitas, K. Homann, N. Kallay, and K. Kuchitsu: Quantities, Units and Symbols inPhysical Chemistry, 2nd ed., Blackwell Science Publications, London, 1993.c) P. Muller: 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

Preface to the Third EditionVIII

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 eects 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 Eects 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. Forsterling,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 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 insuciently 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 aect 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 eects 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 dicult choice.

The present book is based on an earlier paperback Losungsmitteleekte 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 eects.

A knowledge of the physico-chemical principles of solvent eects 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 eects 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 eects 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 eects in organic chemistry. I am indebted to Professors W. H. Pirkle,Urbana/Illinois, D. Seebach, Zurich/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: Losungsmitteleekte in der organischen Chemie. 2nd edition. Verlag Chemie,Weinheim 1973;Eets 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), Izdatelstvo Khimiya, Leningrad 1973.

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

Preface to the First EditionXII

Contents

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

2 Solute-Solvent Interactions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 5

2.1 Solutions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 52.2 Intermolecular Forces . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 102.2.1 Ion-Dipole Forces . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 102.2.2 Dipole-Dipole Forces. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 112.2.3 Dipole-Induced Dipole Forces . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 132.2.4 Instantaneous Dipole-Induced Dipole Forces . . . . . . . . . . . . . . . . . . . . . . . . . 132.2.5 Hydrogen Bonding . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 152.2.6 Electron-Pair Donor/Electron-Pair Acceptor Interactions (EPD/EPA

Interactions) . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 192.2.7 Solvophobic Interactions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 272.3 Solvation . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 302.4 Selective Solvation . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 382.5 Micellar Solvation (Solubilization) . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 422.6 Ionization and Dissociation . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 46

3 Classification of Solvents . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 57

3.1 Classification of Solvents according to Chemical Constitution . . . . . . . . . 573.2 Classification of Solvents using Physical Constants . . . . . . . . . . . . . . . . . . . . 623.3 Classification of Solvents in Terms of Acid-Base Behaviour. . . . . . . . . . . . 733.3.1 Brnsted-Lowry Theory of Acids and Bases . . . . . . . . . . . . . . . . . . . . . . . . . . 733.3.2 Lewis Theory of Acids and Bases . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 793.4 Classification of Solvents in Terms of Specific Solute/Solvent

Interactions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 823.5 Classification of Solvents using Multivariate Statistical Methods . . . . . . . 84

4 Solvent Eects on the Position of Homogeneous Chemical Equilibria . . . . 93

4.1 General Remarks . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 934.2 Solvent Eects on Acid/Base Equilibria . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 954.2.1 Brnsted Acids and Bases in Solution . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 954.2.2 Gas-Phase Acidities and Basicities. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 994.3 Solvent Eects on Tautomeric Equilibria . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1064.3.1 Solvent Eects on Keto/Enol Equilibria . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1064.3.2 Solvent Eects on other Tautomeric Equilibria . . . . . . . . . . . . . . . . . . . . . . . 1134.4 Solvent Eects on other Equilibria . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1214.4.1 Solvent Eects on Brnsted Acid/Base Equilibria . . . . . . . . . . . . . . . . . . . . . 1214.4.2 Solvent Eects on Lewis Acid/Base Equilibria . . . . . . . . . . . . . . . . . . . . . . . . 1234.4.3 Solvent Eects on Conformational Equilibria . . . . . . . . . . . . . . . . . . . . . . . . . 1264.4.4 Solvent Eects on cis/trans or E/Z Isomerization Equilibria . . . . . . . . . . . 1324.4.5 Solvent Eects on Valence Isomerization Equilibria . . . . . . . . . . . . . . . . . . . 1354.4.6 Solvent Eects on Electron-Transfer Equilibria . . . . . . . . . . . . . . . . . . . . . . . 1374.4.7 Solvent Eects on Host/Guest Complexation Equilibria . . . . . . . . . . . . . . . 139

5 Solvent Eects on the Rates of Homogeneous Chemical Reactions. . . . . . 147

5.1 General Remarks . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1475.2 Gas-Phase Reactivities . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1555.3 Qualitative Theory of Solvent Eects on Reaction Rates. . . . . . . . . . . . . . 1625.3.1 The HughesIngold Rules . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1635.3.2 Solvent Eects on Dipolar Transition State Reactions . . . . . . . . . . . . . . . . 1735.3.3 Solvent Eects on Isopolar Transition State Reactions. . . . . . . . . . . . . . . . 1875.3.4 Solvent Eects on Free-Radical Transition State Reactions . . . . . . . . . . . 1995.3.5 Limitations of the HughesIngold Rules . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2155.4 Quantitative Theories of Solvent Eects on Reaction Rates . . . . . . . . . . . 2185.4.1 General Remarks . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2185.4.2 Reactions between Neutral, Apolar Molecules . . . . . . . . . . . . . . . . . . . . . . . 2195.4.3 Reactions between Neutral, Dipolar Molecules. . . . . . . . . . . . . . . . . . . . . . . 2255.4.4 Reactions between Neutral Molecules and Ions . . . . . . . . . . . . . . . . . . . . . . 2335.4.5 Reactions between Ions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2345.5 Specific Solvation Eects on Reaction Rates . . . . . . . . . . . . . . . . . . . . . . . . . 2375.5.1 Influence of Specific Anion Solvation on the Rates of SN and other

Reactions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2385.5.2 Protic and Dipolar Aprotic Solvent Eects on the Rates of SN

Reactions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2435.5.3 Quantitative Separation of Protic and Dipolar Aprotic Solvent Eects

for Reaction Rates by Means of Solvent-Transfer Activity Coecients 2545.5.4 Acceleration of Base-Catalysed Reactions in Dipolar Aprotic Solvents 2595.5.5 Influence of Specific Cation Solvation on Rates of SN Reactions . . . . . . 2625.5.6 Solvent Influence on the Reactivity of Ambident Anions. . . . . . . . . . . . . . 2695.5.7 Solvent Eects on Mechanisms and Stereochemistry of Organic

Reactions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2735.5.8 Influence of Micellar and Solvophobic Interactions on Reaction Rates

and Mechanisms . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2925.5.9 Liquid Crystals as Reaction Media . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2985.5.10 Solvent Cage Eects . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3035.5.11 External Pressure and Solvent Eects on Reaction Rates . . . . . . . . . . . . . 3085.5.12 Solvent Isotope Eects . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3155.5.13 Reactions in Biphasic Solvent Systems and in Neoteric Solvents. . . . . . . 317

6 Solvent Eects on the Absorption Spectra of Organic Compounds . . . . . . 329

6.1 General Remarks . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3296.2 Solvent Eects on UV/Vis Spectra . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3306.2.1 Solvatochromic Compounds . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3306.2.2 Theory of Solvent Eects on UV/Vis Absorption Spectra . . . . . . . . . . . . . 3406.2.3 Specific Solvent Eects on UV/Vis Absorption Spectra . . . . . . . . . . . . . . . 3486.2.4 Solvent Eects on Fluorescence Spectra. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3526.2.5 Solvent Eects on ORD and CD Spectra . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3596.3 Solvent Eects on Infrared Spectra . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3636.4 Solvent Eects on Electron Spin Resonance Spectra . . . . . . . . . . . . . . . . . . 369

ContentsXIV

6.5 Solvent Eects on Nuclear Magnetic Resonance Spectra . . . . . . . . . . . . . . 3756.5.1 Nonspecific Solvent Eects on NMR Chemical Shifts . . . . . . . . . . . . . . . . . 3756.5.2 Specific Solvent Eects on NMR Chemical Shifts . . . . . . . . . . . . . . . . . . . . . 3816.5.3 Solvent Eects on Spin-Spin Coupling Constants . . . . . . . . . . . . . . . . . . . . . 387

7 Empirical Parameters of Solvent Polarity . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 389

7.1 Linear Gibbs Energy Relationships. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3897.2 Empirical Parameters of Solvent Polarity from Equilibrium

Measurements. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3967.3 Empirical Parameters of Solvent Polarity from Kinetic Measurements . 4027.4 Empirical Parameters of Solvent Polarity from Spectroscopic

Measurements. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 4117.5 Empirical Parameters of Solvent Polarity from other Measurements . . . 4437.6 Interrelation and Application of Solvent Polarity Parameters . . . . . . . . . . 4457.7 Multiparameter Approaches . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 452

Appendix . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 471

A. Properties, Purification, and Use of Organic Solvents. . . . . . . . . . . . . . . . . . 471A.1 Physical Properties . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 471A.2 Purification of Organic Solvents. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 471A.3 Spectroscopic Solvents. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 479A.4 Solvents as Reaction Media. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 488A.5 Solvents for Recrystallization . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 488A.6 Solvents for Extraction and Partitioning (Distribution) . . . . . . . . . . . . . . . . 490A.7 Solvents for Adsorption Chromatography . . . . . . . . . . . . . . . . . . . . . . . . . . . . 492A.8 Solvents for Acid/Base Titrations in Non-Aqueous Media . . . . . . . . . . . . . 496A.9 Solvents for Electrochemistry . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 496A.10 Toxicity of Organic Solvents . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 500

References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 509

Figure and Table Credits . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 581

Subject Index . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 583

Author Index . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 599

Contents XV

List of Abbreviations

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

NA Avogadro constant 6:0221 1023 mol1c0 speed of light in vacuum 2:9979 108 m s1e0 absolute permittivity of vacuum

[ 1=m0 c02; m0 permeability ofvacuum]

8:8542 1012C2 J1 m1

e elementary charge 1:6022 1019 Ch Planck constant 6:6261 1034 J sR gas constant 8.3145 J K1 mol1

(or 0.08206L atm K1 mol1)

kB Boltzmann constant ( R=NA) 1:3807 1023 J K1Vm standard molar volume of an ideal

gas (at t 0 C and p 100 kPa)22.711 L mol1

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 ln10 lg x; lg decadic logarithm)

2.303

Abbreviations and Symbols for Unitsa,b)

bar bar ( 105 Pa 105 N m2) pressurecg/g centigram/gram weight percent

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

cmol/mol centimol/mol mole percent

cm centimetre (102 m) lengthcm3 cubic centimetre

(millilitre mL; 106 m3)volume

C coulomb electric charge

a) I. Mills, T. Cvitas, K. Homann, N. Kallay, and K. Kuchitsu: Quantities, Units and Symbols inPhysical Chemistry. 2nd ed., Blackwell Scientific Publications, London, 1993.b) 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.

C degrees centigrade (Celsius) temperaturedm3 cubic decimetre (litre L; 103 m3) volumeJ joule energy

kJ kilojoule (103 J) energy

K kelvin temperature

L, l litre (1 dm3; 103 m3) volumem metre length

min minute time

mol mole amount of substance

MPa megapascal (106 Pa) pressure

mT millitesla (103 T) magnetic flux density(magnetic field)

nm nanometre (109 m) lengthPa pascal (1 N m2 105 bar) pressurepercent (%) part per hundred (102) dimensionless fractionppm part per million (106) dimensionless fractions second time

Abbreviations and Symbols for Propertiesc)

ai activity of solute i

a1H ESR hyperfine coupling constant(coupling with 1H)

mT ( 103 T)

Aj the solvents anion-solvating tendencyor acity (Swain)

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

a electric polarizability of a molecule,polarizability volume

C2 m2 J1 or 4pe0 cm3

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

B empirical parameter of solvent Lewisbasicity (Palm and Koppel)

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

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

List of AbbreviationsXVIII

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

Bj the solvents 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 ( 106 Pa)

ci; ci molar concentration of solute i mol L1CA;CB Lewis acidity and Lewis basicity

parameter (Drago)

cmc critical micelle concentration mol L1DHA molar bond-dissociation energy of the

bond between H and AkJ mol1

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

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

kcal mol1

DNN normalized solvent donor number(Marcus)

d; dH Hildebrands solubility parameter MPa1=2

d chemical shift of NMR signals ppm

d solvent polarizability correction term(Taft and Kamlet)

E energy, molar energy kJ mol1E electric field strength V m1E enol constant (K. H. Meyer)

E empirical parameter of solvent Lewisacidity (Palm and Koppel)

EA;Ea Arrhenius activation energy kJ mol1EA;EB Lewis acidity and Lewis basicity

parameter (Drago)

EA electron anity kJ mol1ENB empirical solvent Lewis basicity

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

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

kcal mol1

List of Abbreviations XIX

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 mol1 or kcal mol1

ET30 empirical solvent polarity parameter,based on the intramolecular CTabsorption of a pyridinium-N-phenolate betaine dye (Dimroth andReichardt)

kcal mol1

ENT normalized ET30 solvent polarityparameter (Reichardt)

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

kcal mol1

EPA electron-pair acceptor

EPD electron-pair donor

er relative permittivity (e=e0)(dielectric constant)

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 mol1DG0 standard molar Gibbs energy of

activationkJ mol1

DGsolv standard molar Gibbs energy ofsolvation

kJ mol1

DGhydr standard molar Gibbs energy ofhydration

kJ mol1

DGt X;O!S,DGt X;W!S

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

kJ mol1

gi activity coecient of solute i

DH standard molar enthalpy change kJ mol1DH0 standard molar enthalpy of activation kJ mol1DHv molar enthalpy (heat) of

vapourizationkJ mol1

H0 acidity function (Hammett)

HBA hydrogen-bond acceptor

List of AbbreviationsXX

HBD hydrogen-bond donor

HOMO highest occupied molecular orbital

Ei; I ; IP ionization energy kJ mol1I gas-chromatographic retention index

(Kovats)

J NMR spin-spin coupling constant Hz

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

(L mol1)n1 s1

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

(L mol1)n1 s1

k0 in Hammett equations the rateconstant of unsubstituted substrates

(L mol1)n1 s1 withn 1 or 2

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

(mol L1)Sv

Ka;Kb acid and base ionization constants (mol L1)SvKauto autoionization ion product,

autoprotolysis constantmol2 L2

KAssoc;KDissoc,Kion;KT

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

(mol L1)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 ( 109 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)

m permanent electric dipole moment ofa molecule

C m (or D)

mind induced electric dipole moment of amolecule

C m (or D)

List of Abbreviations XXI

mi standard chemical potential of solute i kJ mol1myi standard chemical potential of solute i

at infinite dilutionkJ mol1

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

N empirical parameter of solventnucleophilicity (Winstein andGrunwald)

N nucleophilicity parameter for(nucleophile solvent)-systems(Ritchie)

n frequency Hz, s1

n frequency in the gas phase or in aninert reference solvent

Hz, s1

~nn wavenumber ( 1=l) cm1W empirical solvent polarity parameter,

based on a Diels-Alder reaction(Berson)

p pressure Pa ( 1N m2),bar ( 105 Pa)

P measure of solvent polarizability(Palm and Koppel)

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

PA proton anity kJ mol1Py empirical solvent polarity parameter,

based on the p ! p emission ofpyrene (Winnik)

Po=w 1-octanol/water partition coecient(Hansch and Leo)

pH lg[H3O], lg c(H3O)(abbreviation of potentia hydrogeniior puissance dhydrogene (Sorensen1909)

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

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

List of AbbreviationsXXII

pazo 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 ( 102 m)

r distance between centres of two ionsor molecules

cm ( 102 m)

r density (mass divided by volume) g cm3r; rA Hammett reaction resp. 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 K1 mol1DS0 standard molar entropy of activation J K1 mol1Sp solvophobic power of a solvent

(Abraham)

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

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

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

s Hammett substituent constant

s NMR screening constant

t Celsius temperature C

T thermodynamic temperature K

tmp melting pointC

tbp boiling pointC

U internal energy kJ

DUv molar energy of vapourization kJ mol1

List of Abbreviations XXIII

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

DV0 molar volume of activation cm3 mol1xi; xi mole fraction of i xi ni=

Pn

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 mol1

OySX;WySX solvent-transfer activity coecient of

a solute X from a reference solvent(O) or water (W) to anothersolvent (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 positive for cations,negative for anions

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

kcal mol1

List of AbbreviationsXXIV

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 eectus mirificos adscribunt. Atque illud

quidem Menstruum vocaverunt.*)

Hermannus Boerhaave (16681738)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 alchemists search for a universal solvent, the so-called Alkahest or Men-struum universale, as it was called by Paracelsus (14931541), 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 anyAlkahest, 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 (15771644) 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 sucientlyanalyzed 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 eects of their science. And this they call the solvent (menstruum).** Even if the once famous scholar J. B. Van Helmont (15771644) claimed to have prepared thisAlkahest 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.


of a substance in a solvent remained a rather mysterious process. The famous Russianpolymath Lomonosov (17111765) 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 (18301901) [28], the Dutch physical chemistvant Ho (18521911) [5], and the Swedish scientist Arrhenius (18591927) [6]. Raoultsystematically studied the eects 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 Raoults law [28].

The diculty in explaining the eects 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 vant 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 vantHo (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 (18531932), Nernst (18641941), Lewis (18751946), Debye (18841966), E. Huckel (18961980), and Bjerrum (18791958). 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 Pean de Saint-Gilles in 1862 in connection with their studies on theesterification of acetic acid with ethanol: . . . letherification est entravee et ralentie parlemploi des dissolvants neutres etrangers a la reaction [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, lexperience montre que ces dissolvantsexercent sur la vitesse de combinaison une influence considerable. Si nous representonspar 1 la constante de vitesse de la reaction precitee dans lhexane C6H14, cette constantepour la meme combinaison dans CH3aaCOaaC6H5, toutes choses egales dailleurs sera847.7. La dierence est enorme, mais, dans ce cas, elle natteint pas encore le maxi-

* . . . the esterification is disturbed and decelerated on addition of neutral solvents not belongingto the reaction [9].

1 Introduction2

mum. . . . Vous voyez que les dissolvants, soi-disant indierents ne sont pas inertes; ilsmodifient profondement lacte de la combinaison chimique. Cet enonce est riche enconsequences pour la theorie 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 gibt

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

aa

aaCOaa wie in der Form

aaCOaaC

aa

HaaCOaa zu bestehen vermogen; von der Natur der angelagerten Reste, vonder Temperatur, bei den gelosten Substanzen auch von der Art des Losungsmittels hangtes ab, welche von den beiden Formen die bestandigere 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, wherebyhe 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 eect of sol-vent on constitutional and tautomeric isomerization equilibria was later studied in detail

* 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 indierent, 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 gives thesame 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 dierent meaning however from that accepted today, wasintroduced by Laar [12]. For a description of the development of the concept of tautomerism, seeIngold [13].*** 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].

1 Introduction 3

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 eect 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 dyestus, namely chlorophyll, fuchsin, anilinegreen, cyanine, quinizarin, and egg yolk in twelve dierent solvents. The albeit limited validity of Kundts rule, e.g. found in the cases of 4-hydroxyazobenzene [24] and ace-tone [25], led to the realization that the eect 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 eects 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 dicult to contain. Therefore, the intro-duction of cleaner technologies has become a major concern throughout both academiaand industry [3134]. This includes the development of environmentally benign newsolvents, sometimes called neoteric solvents (neoteric 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.

For the development of a sustainable chemistry based on clean technologies, thebest solvent would be no solvent at all. For this reason, considerable eorts 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 [3742]. 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 eects will certainly not become superfluous in the foreseeablefuture.

* A survey of older works of solvent eects on UV/Vis absorption spectra has been given bySheppard [21].** It should be noted that the now generally accepted meaning of the term solvatochromism diersfrom that introduced by Hantzsch (cf. Section 6.2).

1 Introduction4

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 dierently 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 coecients g are close to unity (g 1) [1, 171]. Solute/solvent mixtures A Bthat obey Raoults 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 ApA to its vapour pressure as a pure liquid (pA) is equal to the mole fraction ofAxA in the liquid mixture, i.e. xA pA=pA. Many mixtures obey Raoults 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 physical constants such as density, dielectric constant, index of refraction etc., but asa 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 (andin 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 indierent medium in which the dissolved material diuses 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 dicult 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 dicult tasks ofphysical chemistry [27, 172174].

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 diraction 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 aninfinitely dilute solution.


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 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 [815, 15a]*). Numerous dierent 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 [304]. Fig. 2-1 should make clear the complexity of the inner structureof 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 itas a dynamic three-dimensional hydrogen-bonded network, without a significant num-ber of non-bonded water molecules, that retains several of the structural characteristicsof ice (i.e. tetrahedral molecular packing with each water molecule hydrogen-bondedto 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, 177179]. 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 indierent 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 suces. When a chemical

* The amusing story of polywater, which excited the scientific community for a few years duringthe late 1960s and early 1970s, 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.

2 Solute-Solvent Interactions6

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

2.1 Solutions 7

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 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 rearma-tion of the old like dissolves like rule is the dierential solubility of fullerene (C60),consisting of a three-dimensional delocalized 60p-electron system, in solvents such asmethanol (s 0:01 mg/mL) and 1-chloronaphthalene (s 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.

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) Dicult to break up B B.c) Dicult 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 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. Fora deeper and more detailed understanding of the diverse rules determining the solubilityof organic compounds in various solvents, see references [312316].

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 DUvVm

1=2 DHv R T

Vm

1=22-1

In this equation, Vm is the molar volume of the solvent, and DUv and DHv are themolar energy and the molar enthalpy (heat) of vapourization to a gas of zero pressure,

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

Solute Solute polarity Solubility/(mol m3) 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.

2.1 Solutions 9

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 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 dierent 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 paran oil, silicon 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 (paran) 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, 182184]

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 donoracceptor 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 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 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 14p e0

z e m cos yr2

2-2*)

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 center of the dipole. Cos y 1 for y 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 0) in vacuum. For typical interatomic spacings (rA300400 pm), theion-dipole interaction is much stronger than the thermal energy k T at 300 K. Forthe monovalent sodium cation (z 1, radius 95 pm) near a water molecule(radiusA140 pm; m 5:9 1030 Cm), the maximum interaction energy calculated byEq. (2-2) amounts to U 39k T or 96 kJ mol1 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 1030 Cm (i.e. Coulombmeter). Among the solvents listed in theAppendix, Table A-1, hexamethylphosphoric triamide is the one with the highest dipolemoment (m 18:48 1030 Cm), followed by propylene carbonate (m 16:7 1030Cm), and sulfolane (m 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(tbp 155 C/3 Torr) with a large relative permittivity (er 64:6 at 25 C), has a dipolemoment of m 35:7 1030 Cm (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 sucientlystable 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.

* 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.

2.2 Intermolecular Forces 11

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) (kB Boltzmann constant; T absolute temperature) [29].

Udipole-dipole 14p e02 2m

21 m22

3kB T r6 2-3

As the temperature increases, the angle-averaged dipole/dipole interaction energybecomes less negative until at very high temperatures all dipole orientations are equallypopulated 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 3:3 1030 Cm (1 D), at a sepa-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 3:7kJ mol1 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 1) is a pointcharge and a monopole moment represents an overall charge (e.g. of an ion Na orCl). A dipole (n 2; e.g. H2O, H3CaaCOaaCH3) is an array of partial charges withno monopole moment (i.e. no charge). A quadrupolar molecule (n 4; e.g. CO2, C6H6)has neither a net charge nor a dipole moment, and an octupolar molecule (n 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 oftwo dipole molecules.


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

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 polarizability a of the apolar molecule experiencing the induction of the per-manent dipole. The net dipole/induced dipole energy of interaction for two dierentmolecules, each possessing a permanent dipole moment m1 and m2 and polarizabilities a1and a2, often referred to as the induction or Debye interaction [32], is given by Eq. (2-4).

Udipole-induced dipole 14p e02 a1 m

22 a2 m21r6

2-4

For a dipolar molecule of m 3:3 1030 Cm (1 D; e.g. HaaCl) separated from amolecule of polarization volume a 10 1030 m3 (e.g. C6H6) by a distance of 300 pm,the temperature-independent interaction energy is about 0.8 kJ/mol [26d].

Similarly, a charged particle such as an ion introduced into the neighbourhood ofan 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 aorded by the ion with charge z e. Theenergy of such an interaction is given by Eq. (2-5).

U ion-induced dipole 14p e02 z

2 e2 a2 r4 2-5

The 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-

* The induced dipole moment is defined as m ind 4p e0 a E (e0 permittivity of vacuum; a elec-tric polarizability of the molecule; E electric field strength).


actions can be expressed as

Udispersion 14p e02 3a1 a2

2r6 I1 I2

I1 I2

2-6a

where a1 and a2 are the polarizabilities and I1 and I2 are the ionization potentials of thetwo dierent interacting species [33]. When applied to two molecules of the same sub-stance, Eq. (2-6a) reduces to Eq. (2-6b).

Udispersion 14p e02 3a

2 I4r6

2-6b

Dispersion forces are extremely short-range in action (depending on 1=r6!).Dispersion forces are universal for all atoms and molecules; they alone are

responsible 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 40 C consists of 8% orientational energy,14% inductional energy, and 78% dispersion energy [35]. Two molecules witha 3 1030 m3, I 20 1019 J, and r 3 1010 m have an interaction potential of11.3 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 magnitudeof 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 ofr 3 1010 m and oriented as in Fig. 2-3a, the interaction energy is only 5.3 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 a2 term 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 equationof 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 (n 1:6273), iodobenzene (n 1:6200), aniline(n 1:5863), and diphenyl ether (n 1:5763); of all organic solvents, carbon disulfide(n 1:6275) and diiodomethane (n 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 , I

m, SCNm or the picrateanion are significantly larger than for the smaller anions like Fm, HOm, or R2N

m [36].Perfluorohydrocarbons have unusually low boiling points because tightly held electronsin fluorine have only a small polarizability.


2.2.5 Hydrogen Bonding [3746, 187190, 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 :YaaR 0 according to Eq. (2-7).

2-7

RaaXaaH is the proton donor and :YaaR 0 makes available an electron pairfor the bridging bond. Thus, hydrogen bonding can be regarded as a preliminarystep in a Brnsted acid-base reaction which would lead to a dipolar reaction productRaaXm HaaYlaaR 0. 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 pairsOaaH 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 dierent 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, lone-pair directionality) has been reviewed [194].

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 enthalpiesof 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 (thehydrogen bonds are denoted by dotted lines).

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

* Bond dissociation enthalpies outside these limits are, however, known. Examples of weak, nor-mal, and strong hydrogen bonds are found in the following pairs: phenol/benzene (DH 5 kJ/mol) [47], phenol/triethylamine (DH 37 kJ/mol) [47], and trichloroacetic acid/triphenylphos-phane oxide (DH 67 kJ/mol) [48]. An extremely strong hydrogen bond is found in Me4NHF2(DH 155 kJ/mol) [38]. The strength of a hydrogen bond correlates with the basicity of theproton-acceptor and the acidity of the proton-donor molecule. Compounds with very strong hy-drogen bonds have been reviewed [320].


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 YaaR 0; (d) due to the reduced elec-tron density at H-atoms involved in hydrogen bonds, they are deshielded, resulting insubstantial downfield shifts of their 1H NMR signals; (e) in hetero-molecular hydrogenbonds, a shift of the Brnsted acid/base equilibrium RaaXaaH YaaR 0 S RaaXm HaaYlaaR 0 to 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 [4246]. The hydrogen bond can be describedas 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 asRaaXdmaaHdl YdmaaR 0. 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 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 bondsof the type Fdm H Fdm cannot be explained in terms of the electrostatic model.When the XaaY distance is suciently 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, whichdier only in the position of the proton*).

2-8

* The term protomeric structure was obviously introduced in analogy to the well-knownmesomeric structures, which are used to describe the electronic ground state of aromatic com-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 referredto as dipolar aprotic (e.g. CH3SOCH3, CH3CN, CH3NO2) are in fact not aprotic.

front matter'. in: solvents and solvent effects in ... · organic synthesis workbook ii 2001....

Documents