Contents
1 Introduction
2 Solute-Solvent Interactions
2.1 Solutions
2.2 Intermolecular Forces [26, 27, 182-184]
2.3 Solvation [49, 85–98, 98a, 233–241]
2.4 Preferential Solvation [89, 94, 96, 118–120, 241, 249, 250]
2.5 Micellar Solvation (Solubilization) [96, 128, 130-132, 220, 262-267]
2.6 Ionization and Dissociation [49, 96, 134-139, 241, 278, 279]
3 Classification of Solvents
3.1 Classification of Solvents according to Chemical Constitution
3.2 Classification of Solvents using Physical Constants
3.3 Classification of Solvents in Terms of Acid-Base Behaviour
3.4 Classification of Solvents in Terms of Specific Solute/Solvent Interactions
3.5 Classification of Solvents using Multivariate Statistical Methods
4 Solvent Effects on the Position of Homogeneous Chemical Equilibria
4.1 General Remarks
4.2 Solvent Effects on Acid/Base Equilibria [8-13]
4.3 Solvent Effects on Tautomeric Equilibria
4.4 Solvent Effects on Other Equilibria
5 Solvent Effects on the Rates of Homogeneous Chemical Reactions
5.1 General Remarks [1-29, 452-463]
5.2 Gas-Phase Reactivities
5.3 Qualitative Theory of Solvent Effects on Reaction Rates
5.4 Quantitative Theories of Solvent Effects on Reaction Rates
5.5 Specific Solvation Effects on Reaction Rates
6 Solvent Effects on the Absorption Spectra of Organic Compounds
6.1 General Remarks
6.2 Solvent Effects on UV/Vis Spectra [5-17]
6.3 Solvent Effects on Infrared Spectra
6.4 Solvent Effects on Electron Spin Resonance Spectra
6.5 Solvent Effects on Nuclear Magnetic Resonance Spectra
7 Empirical Parameters of Solvent Polarity
7.1 Linear Gibbs Energy Relationships
7.2 Empirical Parameters of Solvent Polarity from Equilibrium Measurements
7.3 Empirical Parameters of Solvent Polarity from Kinetic Measurements
7.4 Empirical Parameters of Solvent Polarity from Spectroscopic Measurements
7.5 Empirical Parameters of Solvent Polarity from Other Measurements
7.6 Interrelation and Application of Solvent Polarity Parameters
7.7 Multiparameter Approaches
8 Solvents and Green Chemistry
8.1 Green Chemistry
8.2 Reduction of Solvent Use
8.3 Green Solvent Selection
8.4 Non-Traditional Solvents
8.5 Outlook
Appendix A. Properties, Purification, and Use of Organic Solvents
A.1 Physical Properties
A.2 Purification of Organic Solvents
A.3 Spectroscopic Solvents
A.4 Solvents as Reaction Media
A.5 Solvents for Recrystallization
A.6 Solvents for Extraction and Partitioning (Distribution)
A.7 Solvents for Adsorption Chromatography
A.8 Solvents for Acid/Base Titrations in Non-Aqueous Media
A.9 Solvents for Electrochemistry
A.10 Toxicity of Organic Solvents
References
Figure and Table Credits
Subject Index
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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
Cover
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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To Maria
and in memory of my parents
C. R.
To Mike
and my parents
T. W.
Preface to the Fourth Edition
About 40 years ago, in 1969, a German paperback entitled Lösungsmitteleffekte in der organischen Chemie, written by the first author, was published by Verlag Chemie in Weinheim. Based on this paperback and its second edition in 1973, an enlarged English edition called Solvent Effects in Organic Chemistry appeared in 1979, followed by a second English edition in 1988 with the now enlarged title Solvents and Solvent Effects in Organic Chemistry. A first and second reprint in 2004 and 2005 of the third, updated and enlarged English edition of this book, published in 2003, demonstrate the continuing common interest in the study of solvent effects on chemical reactions and physical processes. This and the progress that has been made in recent years in this field of research encouraged us to present now to the interested reader a fourth, again updated and enlarged, edition of this book. This was only possible because a junior author helped the meanwhile retired senior author with the preparation of the manuscript for this new edition, particularly in writing the new Chapter 8.
This new chapter deals with the relationship between solvents and green chemistry, the classification of solvents by their environmental impact, and the replacement of traditional by non-traditional solvents for chemical reactions.
During the seven years after publication of the third edition in 2003, the number of solvent-dependent processes studied has increased to such an extent (particularly in the field of ionic liquids) that only a careful selection of instructive and representative examples could be additionally included in this fourth edition. The literature has been covered up to 2009, partly up to 2010. New references have been added at the end of the reference list of each chapter.
Consistent use of the nomenclaturea), symbolsb), termsc) and SI unitsd) recommended by the respective IUPAC Commissions has again been made in this fourth edition.
For useful comments and valuable suggestions we thank many colleagues, in particular Prof. Dr. N. O. Mchedlov-Petrossyan, Kharkov/Ukraine, Dr. T. Rager, Basel/Switzerland, and Prof. Dr. G. N. Papatheodorou, Rio/Greece. For their assistance in providing us with difficult to obtain literature and in preparing the final manuscript, 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, for her continuous support of this project. T.W. thanks the final-year Imperial College Chemistry students and Green Chemistry Master students for helpful discussions.
We both express our thanks to the staff of Wiley-VCH Verlag GmbH, Weinheim, particularly to Dr. Elke Maase and Dr. Stefanie Volk, for their help and excellent work in turning the manuscript into this final book.
| Marburg (Lahn) London Summer 2010 |
Christian Reichardt Thomas Welton |
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, and J.-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 and Symbols 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 Recommendations 1994, 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.
Preface to the Third Edition
Meeting the demand for the second edition of this book, which is – despite a reprint in 1990 – no longer available, and considering the progress that has been made during the last decade in the study of solvent effects in experimental and theoretical organic chemistry, this improved third edition is presented to the interested reader.
Following the same layout as in the second edition, all topics retained have been brought 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 effects on host/guest complexation equilibria and reactions in biphasic solvent systems and neoteric solvents, respectively. More than 900 new references have been added, giving preference to review articles, and many older ones have been deleted. New references either replace 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 effects in all areas of organic chemistry, only some illustrative examples from the didactic and systematic point of view could be selected. This book is not a monograph covering all relevant literature in this field of research. The author, responsible for this subjective selection, apologizes in advance to all chemists whose valuable work on solvent effects is not mentioned in this book. However, using the reviews cited, the reader will find easy access to the full range of papers published in a certain field of research on solvent effects.
Great progress has been made during the last decade in theoretical treatments of solvent effects by various quantum-chemical methods and computational strategies. When indicated, relevant references are given to the respective solution reactions or absorptions. However, a critical evaluation of all the theoretical models and methods used to calculate the differential solvation of educts, activated complexes, products, ground and excited states, is outside the expertise of the present author. Thus, a book on all kinds of theoretical calculations of solvent influences on chemical reactions and physical absorptions has still to be written by someone else.
Consistent use of the nomenclature,a) symbols,b) terms,c) and SI unitsd) recommended by the IUPAC commissions has also been made in this third edition.
For comments and valuable suggestions I have to thank many colleagues, in particular 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 providing me with difficult to obtain literature, I thank Mr. G. Schäfer (technician), Mrs. S. Schellenberg (secretary), and Mrs. B. Becht-Schröder (librarian), all at the Department of Chemistry, Philipps University, Marburg/Germany. Special thanks are due to the staff of Wiley-VCH Verlag GmbH, Weinheim/Germany, particularly to Dr. Elke Westermann, 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 the manuscript, but also for her constant encouragement and support during the writing of this book.
| Marburg (Lahn), Spring 2002 | Christian Reichardt |
a) G. J. Leigh, H. A. Favre, and W. V. Metanomski: Principles of Chemical Nomenclature – A Guide to IUPAC Recommendations, Blackwell Science Publications, London, 1998.
b) I. Mills, T. Cvitaš, K. Homann, N. Kallay, and K. Kuchitsu: Quantities, Units and Symbols in Physical Chemistry,2nd ed., Blackwell Science Publications, London, 1993.
c) P. Müller: Glossary of Terms used in Physical Organic Chemistry – IUPAC Recommendations 1994, 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.
Preface to the Second Edition
The response to the first English edition of this book, published in 1979, has been both gratifying and encouraging. Its mixed character, lying between that of a monograph and a textbook, has obviously made it attractive to both the industrial and academic chemist as well as the advanced student of chemistry.
During the last eight years the study of solvent effects on both chemical reactions and absorption spectra has made much progress, and numerous interesting and fascinating examples have been described in the literature. In particular, the study of ionic reactions in the gas phase – now possible due to new experimental techniques – has allowed direct comparisons between gas-phase and solution reactions. This has led to a greater understanding of solution reactions. Consequently, Chapters 4 and 5 have been enlarged to include a description of ionic gas-phase reactions compared to their solution counterparts.
The number of well-studied solvent-dependent processes, i.e. reactions and absorptions in solution, has increased greatly since 1979. Only a representative selection of the more instructive, recently studied examples could be included in this second edition.
The search for empirical parameters of solvent polarity and their applications in multiparameter equations has recently been intensified, thus making it necessary to rewrite large parts of Chapter 7.
Special attention has been given to the chemical and physical properties of organic solvents commonly used in daily laboratory work. Therefore, all Appendix Tables have been improved; some have been completely replaced by new ones. A new well-referenced table on solvent-drying has been added (Table A-3). Chapter 3 has been enlarged, in particular by the inclusion of solvent classifications using multivariate statistical methods (Section 3.5). All these amendments justify the change in the title of the book to Solvents and Solvent Effects in Organic Chemistry.
The references have been up-dated to cover literature appearing up to the first part of 1987. New references were added to the end of the respective reference list of each chapter from the first edition.
Consistent use of the nomenclature, symbols, terms, and SI units recommended by the IUPAC commissions has also been made in the second edition.*)
I am very indebted to many colleagues for corrections, comments, and valuable suggestions. Especially helpful suggestions came from Professors H.-D. Försterling, Marburg, J. Shorter, Hull/England, and R. I. Zalewski, Poznań/Poland, to whom I am very grateful. For critical reading of the whole manuscript and the improvement of my English 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 preparation 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 covalent bonds and are relatively involatile, thereby often rendering the gas-phase unsuitable as a reaction medium. Of the thousands of reactions known to occur in solution only few have been studied in the gas-phase, even though a description of reaction mechanisms is much simpler for the gas-phase. The frequent necessity of carrying out reactions in the presence of a more or less inert solvent results in two main obstacles: The reaction depends on a larger number of parameters than in the gas-phase. Consequently, the experimental results can often be only qualitatively interpreted because the state of aggregation in the liquid phase has so far been insufficiently studied. On the other hand, the fact that the interaction forces in solution are much stronger and more varied than in the gas-phase, permits to affect the properties and reactivities of the solute in manifold modes.
Thus, whenever a chemist wishes to carry out a chemical reaction he not only has to take into consideration the right reaction partners, the proper reaction vessels, and the appropriate reaction temperature. One of the most important features for the success of the planned reaction is the selection of a suitable solvent. Since solvent effects on chemical reactivity have been known for more than a century, most chemists are now familiar with the fact that solvents may have a strong influence on reaction rates and equilibria. Today, there are about three hundred common solvents available, nothing to say of the infinite number of solvent mixtures. Hence the chemist needs, in addition to his intuition, some general rules and guiding-principles for this often difficult choice.
The present book is based on an earlier paperback ‘‘Lösungsmitteleffekte in der organischen Chemie’’ [1], which, though following the same layout, has been completely rewritten, greatly expanded, and brought up to date. The book is directed both toward the 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 the other hand wishes to draw conclusions about reaction mechanisms from the observed solvent effects.
A knowledge of the physico-chemical principles of solvent effects is required for proper bench-work. Therefore, a description of the intermolecular interactions between dissolved molecules and solvent is presented first, followed by a classification of solvents derived therefrom. Then follows a detailed description of the influence of solvents on chemical equilibria, reaction rates, and spectral properties of solutes. Finally, empirical parameters of solvent polarity are given, and in an appendix guidelines to the everyday choice of solvents are given in a series of Tables and Figures.
The number of solvent systems and their associated solvent effects examined is so enormous that a complete description of all aspects would fill several volumes. For example, in Chemical Abstracts, volume 85 (1976), approximately eleven articles per week were quoted in which the words ‘‘Solvent effects on...’’ appeared in the title. In the present book only a few important and relatively well-defined areas of general importance have been selected. The book has been written from the point of view of practical use for the organic chemist rather than from a completely theoretical one.
In the selection of the literature more recent reviews were taken into account mainly. 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 the character of a monograph but also to some extent that of a textbook. In order to help the reader in his use of the literature cited, complete titles of the review articles quoted are given. The literature up until December 1977 has been considered together with a few papers from 1978. The use of symbols follows the recommendations of the Symbols Committee of the Royal Society, London, 1971 [2].
I am very grateful to Professor Karl Dimroth, Marburg, who first stimulated my interest in solvent effects in organic chemistry. I am indebted to Professors W. H. Pirkle, Urbana/Illinois, D. Seebach, Zürich/Switzerland, J. Shorter, Hull/England, and numerous other colleagues for helpful advice and information. Thanks are also due to the authors and publishers of copyrighted materials reproduced with their permission (cf. Figure and Table credits on page 495). For the careful translation and improvement of the English manuscript I thank Dr. Edeline Wentrup-Byrne, Marburg. Without the assistance and patience of my wife Maria, this book would not have been written.
| Marburg (Lahn), Summer 1978 | Christian Reichardt |
References
[1] C. Reichardt: Lösungsmitteleffekte in der organischen Chemie.2nd edition. Verlag Chemie, Weinheim 1973;
Effets de solvant en chimie organique (translation of the first-mentioned title into French, by 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, London, in 1971.
List of Abbreviations
Abbreviations and Recommended Values of Some Fundamental Constants and Numbersa,b)
| NA | Avogadro constant | 6:0221 · 1023 mol–1 |
| C0 | speed of light in vacuum | 2:9979 · 108 m · s–1 |
| ε0 | permittivity of vacuum [= 1/(μ0 · c02); μ0 = permeability of vacuum] | 8:8542 · 10–12 F · m–1 |
| e | elementary charge | 1:6022 · 10–19 C |
| h | Planck constant | 6:6261 · 10–34 J · s |
| R | molar gas constant | 8.3145 J · K–1 · mol–1 (or 0.08206 L · atm · K–1 · mol–1) |
| kB | Boltzmann constant (= R/NA) | 1:3807 · 10–23 J · K–1 |
| Vm | standard molar volume of an ideal gas (at t = 0 ºC and p = 100 kPa) | 22.711 L · mol–1 |
| T0 | zero of the Celsius scale | 273.15 K |
| π | ratio of the circumference to the diameter of a circle | 3.1416 |
| e | exponential number and base of natural logarithms (ln) | 2.7183 |
| ln 10 | natural logarithm of ten (ln x = ln 10 · lg x; lg = decadic logarithm) | 2.303 |
Abbreviations and Symbols for Unitsa,b)
| bar | bar (= 105 Pa = 105 N · m–2) | pressure |
| cg/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 |
| cm3 | cubic centimeter (millilitre mL; 10–6 m3) | volume |
| Q | coulomb | electric charge |
| ºC | degrees Celsius | temperature |
| dm3 | cubic decimetre (litre L; 10–3 m3) | volume |
| J | joule | energy |
| kJ | kilojoule (103 J) | energy |
| K | kelvin | temperature |
| L, l | litre (1 dm3; 10–3 m3) | volume |
| m | 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) | length |
| Pa | pascal (1 N · m–2 = 10–5 bar) | pressure |
| percent (%) | part per hundred (10–2) | dimensionless fraction |
| ppm | part per million (10–6) | dimensionless fraction |
| s | second | time |
Abbreviations and Symbols for Propertiesc)
| ai | activity of solute i | |
| a(1H) | ESR hyperfine coupling constant (coupling with 1H) | Hz or mT (= 10–3 T) |
| Aj | the solvent’s anion-solvating tendency or ‘acity’ (Swain) | |
| AN | solvent acceptor number, based on 31P NMR chemical shift of Et3PO (Gutmann and Meyer) | |
| α | electric polarizability of a molecule or polarizability volume | C2 · m2 · J–1 or 4πε0 · cm3 |
| α | empirical parameter of solvent hydrogen-bond donor acidity (Taft and Kamlet) | |
| B | empirical parameter of solvent Lewis basicity (Palm and Koppel) | |
| BMeOD | IR based empirical parameter of solvent Lewis basicity (Palm and Koppel) | |
| BPhOH | IR based empirical parameter of solvent Lewis basicity (Koppel and Paju; Makitra) | |
| Bj | the solvent’s cation-solvating tendency or ‘basity’ (Swain) | |
| Β | empirical parameter of solvent hydrogen-bond acceptor basicity (Taft and Kamlet) | |
| c | cohesive pressure (cohesive energy density) of a solvent | MPa (= 106 Pa) |
| ci, c(i) | molar concentration of solute i | mol · L–1 |
| CA, CB | Lewis acidity and Lewis basicity parameter (Drago) | |
| cmc | critical micellisation concentration | mol · L–1 |
| DHA | molar bond-dissociation energy of the bond between H and A | |
| Dπ | empirical parameter of solvent Lewis basicity, based on a 1,3-dipolar cycloaddition reaction (Nagai et al.) | |
| DN | solvent donor number (Gutmann) [=–ΔH(D-SbCl5)] | kcal · mol–1 |
| DNN | normalized solvent donor number (Marcus) | |
| δ, δH | Hildebrand’s solubility parameter | MPa1/2 |
| δ | chemical shift of NMR signals | ppm |
| δ | solvent polarizability correction term (Taft and Kamlet) | |
| E | energy, molar energy | kJ · mol–1 |
| E | electric field strength | V · m–1 |
| E | enol constant (K. H. Meyer) | |
| E | empirical parameter of solvent Lewis acidity (Palm and Koppel) | |
| EA, Ea | Arrhenius activation energy | kJ · mol–1 |
| EA, EB | Lewis acidity and Lewis basicity parameter (Drago) | |
| EA | electron affinity | kJ · mol–1 |
E |
empirical solvent Lewis basicity parameter, based on the n → π* absorption of an aminyloxide radical (Mukerjee; Wrona) | |
| EK | empirical solvent polarity parameter, based on the d →π* absorption of a molybdenum complex (Walther) | kcal · mol–1 |
E |
empirical solvent polarity parameter, based on the d →π* absorption of a tungsten complex (Lees) | |
| ET | molar electronic transition energy, molar electronic excitation energy | kJ · mol–1 or kcal · mol–1 |
| ET(30) | empirical solvent polarity parameter, based on the intramolecular CT absorption of a pyridinium-N-phenolate betaine dye (Dimroth and Reichardt) | kcal · mol–1 |
E |
normalized ET(30) solvent polarity parameter (Reichardt) | |
E |
empirical solvent polarity parameter, based on the n →π* absorption of an S-oxide (Walter) | kcal · mol–1 |
| EPA | electron-pair acceptor | |
| EPD | electron-pair donor | |
| εr | relative permittivity (= ε/ε0) (‘‘dielectric constant’’) | 1 |
| Φ | empirical solvent polarity parameter, based on the n →π* absorption of ketones (Dubois) | |
| G | IR based empirical solvent polarity parameter (Schleyer and Allerhand) | |
| ΔGº | standard molar Gibbs energy change | kJ · mol–1 |
| ΔG≠ | standard molar Gibbs energy of activation | kJ · mol–1 |
ΔG |
standard molar Gibbs energy of solvation | kJ · mol–1 |
ΔG |
standard molar Gibbs energy of hydration | kJ · mol–1 |
ΔG (X, O→S), |
standard molar Gibbs energy of | kJ · mol–1 |
| ΔG (X, W→S) |
transfer of solute X from a reference solvent (O) or water (W) to another solvent (S) | |
| yi | activity coefficient of solute i | |
| ΔHº | standard molar enthalpy change | kJ · mol–1 |
| ΔH≠ | standard molar enthalpy of activation | kJ · mol–1 |
| ΔHv | molar enthalpy (heat) of vaporization | kJ · mol–1 |
| H0 | acidity function (Hammett) | |
| HBA | hydrogen-bond acceptor | |
| HBD | hydrogen-bond donor | |
| HOMO | highest occupied molecular orbital | |
| Ei, I, IP | ionization energy | kJ · mol–1 |
| I, Ic | ionic strength (concentration basis) (= 1/2 ∑ci · zi2) | mol · L–1 |
| I | gas-chromatographic retention index (Kováts) | |
| J | NMR spin-spin coupling constant | Hz |
| k | rate constant for monomolecular (n = 1) and bimolecular (n = 2) reactions | (L · mol–1)n–1 · s–1 |
| k0 | rate constant in a reference solvent or in the gas phase for monomolecular (n = 1) and bimolecular reactions (n = 2) | (L · mol–1)n–1 · s–1 |
| k0 | in Hammett equations the rate constant of unsubstituted substrates | (L · mol–1)n–1 · s–1 with n = 1 or 2 |
| K, Kc | equilibrium constant (concentration basis; v = stoichiometric number) | (mol · L–1)Σv |
| Ka, Kb | acid and base ionization constants | (mol · L–1)Σv |
| Kauto | autoionization ion product, autoprotolysis constant | mol2 · L–2 |
| KAssoc, KDissoc, | equilibrium constants of association, | (mol · L–1)Σv |
| Kion, KT | dissociation, ionization, resp. tautomerization reactions | |
| 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 | |
| λ | wavelength | nm (= 10–9 m) |
| m | mass of a particle | g |
| Mr | relative molecular mass of a substance (‘‘molecular weight’’) | |
| M | miscibility number (Godfrey) | |
| MH | microscopic hydrophobicity parameter of substituents (Menger) | |
| μ | empirical solvent softness parameter (Marcus) | C · m (or D) |
| μ | permanent electric dipole moment of a molecule | C · m (or D) |
| μind | induced electric dipole moment of a molecule | kJ · mol–1 |
μ |
standard chemical potential of solute i | kJ · mol–1 |
μ |
standard chemical potential of solute i at infinite dilution | |
| n, nD | refractive index (at sodium D line) (= c0/c) | |
| N | empirical parameter of solvent nucleophilicity (Winstein and Grunwald) | |
| N+ | nucleophilicity parameter for (nucleophile + solvent)-systems (Ritchie) | |
| ν | frequency | Hz, s–1 |
| νº | frequency in the gas phase or in an inert reference solvent | Hz, s–1 |
 |
wavenumber (= 1/λ) | cm–1 |
| Ω | empirical solvent polarity parameter, based on a Diels-Alder reaction (Berson) | |
| p | pressure | Pa (= 1N · m–2), |
| P | measure of solvent polarizability (Palm and Koppel) | bar (= 105 Pa) |
| P | empirical solvent polarity parameter, based on19 F NMR measurements (Taft) | |
| PA | proton affinity | kJ · mol–1 |
| Py | empirical solvent polarity parameter, based on the π* → π emission of pyrene (Winnik) | |
| Po/w | 1-octanol/water partition coefficient (Hansch and Leo) | |
| pH | -lg c(H3O+), or more precisely -lg a(H3O+) abbreviation of potentia or pondus hydrogenii, power of hydrogen, or puissance d’hydrogène (Sørensen 1909). The pH scale ranges usually from 1 to 14, but is open-ended, allowing for pH values below 0 or above 14! |
|
| pK | -lg K | |
| π | internal pressure of a solvent | MPa (= 106 Pa) |
| π* | empirical solvent dipolarity/polarizability parameter, based on the π → π* absorption of substituted aromatics (Taft and Kamlet) | |
π |
empirical solvent dipolarity/polarizability parameter, based on the π → π* absorption of azo merocyanine dyes (Buncel) | |
| πx | hydrophobicity parameter of substituent X in H5C6-X (Hansch) | |
| r | radius of sphere representing an ion or a cavity | cm (= 10–2 m) |
| r | distance between centres of two ions or molecules | cm (= 10–2 m) |
| ρ | density (mass divided by volume) | g · cm–3 |
| ρ, ρA | Hammett reaction or absorption constants | |
| S | generalized for solvent | |
| S | empirical solvent polarity parameter, based on the Z-values (Brownstein) | |
 |
lg k2 for the Menschutkin reaction of tri-n-propylamine with iodomethane (Drougard and Decroocq) | |
| ΔSº | standard molar entropy change | J · K–1 · mol–1 |
| ΔS≠ | standard molar entropy of activation | J · K–1 · mol–1 |
| Sp | solvophobic power of a solvent (Abraham) | |
| SA | empirical parameter of solvent hydrogen-bond donor acidity (Catalán) | |
| SB | empirical parameter of solvent hydrogen-bond acceptor basicity (Catalán) | |
| SPP | empirical parameter of solvent dipolarity/polarizability, based on the π → π* absorption of substituted 7-nitrofluorenes (Catalán) | |
| σ | Hammett substituent constant | |
| σ | NMR screening constant | |
| t | Celsius temperature | ºC |
| T | thermodynamic temperature | K |
| tmp | melting point | ºC |
| tbp | boiling point | ºC |
| U | internal energy | kJ |
| ΔUv | molar energy of vaporization | kJ · mol–1 |
| Vm,Vm,i | molar volume (of i) | cm3 · mol–1 |
| ΔV≠ | molar volume of activation | cm3 · mol–1 |
| xi, x(i) | mole fraction of i (xi = ni/Σn) amount-of-substance fraction | 1 |
| X | empirical solvent polarity parameter, based on an SE2 reaction (Gielen and Nasielski) | |
| χR, χB | empirical solvent polarity parameters, based on the π → π* absorption of merocyanine dyes (Brooker) | kcal · mol–1 |
 |
solvent-transfer activity coefficient for the transfer of a solute X from a reference solvent (O) or water (W) to another solvent (S) | |
| Y | empirical parameter of solvent ionizing power, based on t-butyl chloride solvolysis (Winstein and Grunwald) | |
| YOTs | empirical parameter of solvent ionizing power, based on 2-adamantyl tosylate solvolysis (Schleyer and Bentley) | |
| Y | measure of solvent polarization (Palm and Koppel) | |
| zi | charge number of an ion i (zi = Qi/e) | positive for cations, negative for anions |
| Z | empirical solvent polarity parameter, based on the intermolecular CT absorption of a substituted pyridinium iodide (Kosower) | kcal · mol–1 |
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 and Symbols 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.
c) P. Müller: Glossary of Terms used in Physical Organic Chemistry – IUPAC Recommendations 1994. Pure Appl. Chem. 66, 1077 (1994).
“Agite, Auditores ornatissimi, transeamus alacres ad aliud negotii! quum enim sic satis 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 triumphant, cui artis suae, prae aliis omnibus effectus mirificos adscribunt. Atque illud quidem Menstruum vocaverunt.”*)
Hermannus Boerh−ve (1668–1738)
De menstruis dictis in chemia, in:
Elementa Chemiae (1733) [1, 2].
1
Introduction
The development of our knowledge of solutions reflects to some extent the development of chemistry itself [3]. Of all known substances, water was the first to be considered as a solvent. As far back as the time of the Greek philosophers there was speculation about the nature of solution and dissolution. The Greek alchemists considered all chemically active liquids under the name ‘‘Divine water’’. In this context the word ‘‘water’’ was used to designate everything liquid or dissolved. The Greek philosopher Thales of Miletus (ca. 640–546 bc) asserted that water is the origin out of which everything arose and into everything resolved itself.
From these ancient times, a familiar and today often cited quotation of the famous Greek philosopher Aristotle (384–322 bc) was handed down, which reads in Latin Corpora 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 to Hedvall [44], this seems to be a misinterpretation of the original text given in Greek as 
(Ta hygra mikta malista ton somaton), which is probably taken from Aristotle’s work De Generatione et Corruptione [45]. According to Hedvall, this statement should be better read as „ . . . it is chiefly the liquid substances which react’’ [44] or „ . . . for instance liquids are the type of bodies most liable to mixing’’ [45c]. In this somewhat softened version, Aristotle’s statement is obviously less distinct and didactic. With respect of the many solid/solid reactions known today, it is quite understandable that solid-state chemists were not very happy with the common first version of Aristotle’s statement [43, 44].
The alchemist’s search for a universal solvent, the so-called ‘‘Alkahest’’ or ‘‘Menstruum universale’’, as it was called by Paracelsus (1493–1541), indicates the importance given to solvents and the process of dissolution. Although the eager search of 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 chemical rule that ‘‘like dissolves like’’ (similia similibus solvuntur). However, at that time, the words solution and dissolution comprised all operations leading to a liquid product and it was still a long way to the conceptual distinction between the physical dissolution of a salt or of sugar in water, and the chemical change of a substrate by dissolution, for example, of a metal in an acid. Thus, in the so-called chemiatry period (iatrochemistry period), it was believed that the nature of a substance was fundamentally lost upon dissolution. Van Helmont (1577–1644) was the first to strongly oppose this contention. He claimed that the dissolved substance had not disappeared, but was present in the solution, although in aqueous form, and could be recovered [4]. Nevertheless, the dissolution of a substance in a solvent remained a rather mysterious process. The famous Russian polymath Lomonosov (1711–1765) wrote in 1747: ‘‘Talking about the process of dissolution, it is generally said that all solvents penetrate into the pores of the body to be dissolved and gradually remove its particles. However, concerning the question of what forces cause this process of removal, there does not exist any somehow reasonable explanation, 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 persons, namely the French researcher Raoult (1830–1901) [28], the Dutch physical chemist van’t Hoff (1852–1911) [5], and the Swedish scientist Arrhenius (1859–1927) [6]. Raoult systematically studied the effects of dissolved nonionic substances on the freezing and boiling point of liquids and noticed in 1886 that changing the solute/solvent ratio produces precise proportional changes in the physical properties of solutions. The observation that the vapour pressure of solvent above a solution is proportional to the mole fraction of solvent in the solution is today known as Raoult’s law [28].
The difficulty in explaining the effects of inorganic solutes on the physical properties of solutions led in 1884 to Arrhenius’ theory of incomplete and complete dissociation of ionic solutes (electrolytes, ionophores) into cations and anions in solution, which was only very reluctantly accepted by his contemporaries. Arrhenius derived his dissociation theory from comparison of the results obtained by measurements of electroconductivity and osmotic pressure of dilute electrolyte solutions [6].
The application of laws holding for gases to solutions by replacing pressure by osmotic pressure was extensively studied by van’t Hoff, who made osmotic pressure measurements another important physicochemical method in studies of solutions [5].
The integration of these three basic developments established the foundations of modern solution theory and the first Nobel prizes in chemistry were awarded to van’t Hoff (in 1901) and Arrhenius (in 1903) for their work on the osmotic pressure and the theory of electrolytic dissociation in dilute solutions, respectively.
The further development of solution chemistry is connected with the pioneering work of Ostwald (1853–1932), Nernst (1864–1941), Lewis (1875–1946), Debye (1884– 1966), E. Hϋckel (1896–1980), and Bjerrum (1879–1958). More detailed reviews on the development of modern solution chemistry can be found in references [29, 30].
The influence of solvents on the rates of chemical reactions [7, 8] was first noted by Berthelot and Péan de Saint-Gilles in 1862 in connection with their studies on the esterification of acetic acid with ethanol: ‘‘... l’éthérification est entravée et ralentie par l’emploi des dissolvants neutres étrangers a` la réaction’’ [9]*). After thorough studies on the reaction of trialkylamines with haloalkanes, Menschutkin in 1890 concluded that a reaction cannot be separated from the medium in which it is performed [10]. In a letter to Prof. Louis Henry he wrote in 1890: ‘‘Or, l’expérience montre que ces dissolvants exercent sur la vitesse de combinaison une influence considérable. Si nous représentons par 1 la constante de vitesse de la réaction précitée dans l’hexane C6H14 , cette constante pour la même combinaison dans CH3−CO−C6H5, toutes choses égales d’ailleurs sera 847.7. La différence est énorme, mais, dans ce cas, elle n’atteint pas encore le maximum. . . . Vous voyez que les dissolvants, soi-disant indifférents ne sont pas inertes; ils modifient profondément l’acte de la combinaison chimique. Cet énoncé est riche en conséquences pour la théorie chimique des dissolutions’’ [26]**). Menschutkin also discovered that, in reactions between liquids, one of the reaction partners may constitute an unfavourable solvent. Thus, in the preparation of acetanilide, it is not without importance whether aniline is added to an excess of acetic acid, or vice versa, since aniline in this case is an unfavourable reaction medium. Menschutkin related the influence of solvents 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 compounds (Claisen [14]: acetyl-dibenzoylmethane and tribenzoylmethane; Wislicenus [15]: methyl and ethyl formylphenylacetate; Knorr [16]: ethyl dibenzoylsuccinate and ethyl diacetylsuccinate) and the nitro-isonitro tautomerism of primary and secondary nitro compounds (Hantzsch [17]: phenyl-nitromethane). Thus, Claisen wrote: ‘‘Es gibt Verbindungen, welche sowohl in der Form 
wie in der Form 
zu bestehen vermögen; von der Natur der angelagerten Reste, von der Temperatur, bei den gelösten Substanzen auch von der Art des Lösungsmittels ha¨ngt es ab, welche von den beiden Formen die besta¨ndigere ist’’ [14]*). The study of the keto-enol equilibrium of ethyl formylphenylacetate in eight solvents led Wislicenus to the conclusion that the keto form predominates in alcoholic solution, the enol form in trichloromethane or benzene. He stated that the final ratio in which the two tautomeric forms 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 solvents into two groups according to their ability to isomerize tautomeric compounds. His classification reflects, to some extent, the modern division into protic and aprotic solvents. The effect of solvent on constitutional and tautomeric isomerization equilibria was later studied in detail by Dimroth [19] (using triazole derivatives, e.g. 5-amino-4methoxycarbonyl-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 by the 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 phenomenon solvatochromism***) [22]. The search for a relationship between solvent effect and solvent property led Kundt in 1878 to propose the rule, later named after him, that increasing dispersion (i.e. increasing index of refraction) is related to a shift of the absorption maximum towards longer wavelength [23]. This he established on the basis of UV/Vis absorption spectra of six dyestuffs, namely chlorophyll, fuchsin, aniline green, cyanine, quinizarin, and egg yolk in twelve different solvents. The – albeit limited – validity of Kundt’s rule, e.g. found in the cases of 4-hydroxyazobenzene [24] and acetone [25], led to the realization that the effect of solvent on dissolved molecules is a result of electrical fields. These fields in turn originate from the dipolar properties of the molecules in question [25]. The similarities in the relationships between solvent effects on reaction rate, equilibrium position, and absorption spectra has been related to the general 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 of interest because many of the solvents commonly used in laboratories and in the chemical industry are considered as unsafe for reasons of environmental protection. On the list of damaging chemicals, solvents rank highly because they are often used in huge amounts and because they are volatile liquids that are difficult to contain. Therefore, the introduction of cleaner technologies has become a major concern throughout both academia and industry [31–34]. This includes the development of environmentally benign new solvents, sometimes called neoteric solvents (neoteric = recent, new, modern), constituting a class of novel solvents with desirable, less hazardous, new properties [35, 36]. The term neoteric solvents covers supercritical fluids, ionic liquids, and also perfluorohydrocarbons (as used in fluorous biphasic systems). In addition, water, often considered incompatible with organic synthesis, in recent decades has attracted increasing interest as 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 the substitution of hazardous solvents, together with the relevant literature references; see also Chapter 8.
For the development of a sustainable chemistry based on clean technologies, the best solvent would be no solvent at all. For this reason, considerable efforts have recently been made to design reactions that proceed under solvent-free conditions, using modern 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 efficient solvent-free organic synthesis is the (S)-proline-catalysed stereoselective aldol reaction between cyclohexanone and 4-nitrobenzaldehyde, applying a very simple mechano-chemical technique such as ball milling [42].
‘‘Well then, my dear listeners, let us proceed with fervor to another problem! Having sufficiently analyzed in this manner the four resources of science and nature, which we are about to leave ( fire, water, air, and earth) we must consider a fifth element which can almost be considered the most essential part of chemistry itself, which chemists boastfully, no doubt with reason, prefer above all others, and because of which they triumphantly celebrate, and to which they attribute above all others the marvellous effects of their science. And this they call the solvent (menstruum).’’
* Even if the once famous scholar J. B. Van Helmont (1577–1644) claimed to have prepared this ‘‘Alkahest’’ in a phial, together with the adherents of the alkahest theory he was ridiculed by his contemporaries who asked in which vessel he has stored this universal solvent.
*‘‘. . . the esterification is disturbed and decelerated on addition of neutral solvents not belonging to the reaction’’ [9].
** ‘‘Now, experience shows that solvents exert considerable influence on reaction rates. If we represent the rate constant of the reaction to be studied in hexane C6H14 by 1, then, all else being equal, this constant for the same reaction in CH3−CO−C6H5 will be 847.7. The increase is enormous, but in this case it has not even reached its maximum. . . . So you see that solvents, in spite of appearing at first to be indifferent, are by no means inert; they can greatly influence the course of chemical reactions. This statement is full of consequences for the chemical theory of dissolutions’’ [26].
*** The first observation of a tautomeric equilibrium was made in 1884 by Zincke at Marburg [11]. He observed that, surprisingly, the reaction of 1,4-naphthoquinone with phenylhydrazine gives the same product as that obtained from the coupling reaction of 1-naphthol with benzenediazonium salts. This phenomenon, that the substrate can react either as phenylhydrazone or as a hydroxyazo compound, depending on the reaction circumstances, was called Ortsisomerie by Zincke [11]. Later on, the name tautomerism, with a different meaning however from that accepted today, was introduced by Laar [12]. For a description of the development of the concept of tautomerism, see Ingold [13].
* ‘‘There are compounds capable of existence in the form form 
as well as in the form 
it depends on the nature of the substituents, the temperature, and for dissolved compounds, also on the nature of the solvent, which of the two forms will be the more stable’’ [14].
** A survey of older works of solvent effects on UV/Vis absorption spectra has been given by Sheppard [21].
*** It should be noted that the now generally accepted meaning of the term solvatochromism differs from that introduced by Hantzsch (cf. Section 6.2).