Lewis Basicity and Affinity Scales: Data and Measurement

Lewis Basicity and Affinity Scales Lewis Basicity and Affinity Scales Data and Measurement CHRISTIAN LAURENCE Dépar...

Author: Christian Laurence | Jean-Francois Gal

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Lewis Basicity and Affinity Scales

Lewis Basicity and Affinity Scales Data and Measurement

CHRISTIAN LAURENCE Département de Chimie, Université de Nantes – CNRS, France

JEAN-FRANC ¸ OIS GAL Institut de Chimie, Université de Nice-Sophia Antipolis – CNRS, France

A John Wiley and Sons, Ltd., Publication

This edition first published 2010 C 2010 John Wiley & Sons Ltd Registered office John Wiley & Sons Ltd, The Atrium, Southern Gate, Chichester, West Sussex, PO19 8SQ, United Kingdom For details of our global editorial offices, for customer services and for information about how to apply for permission to reuse the copyright material in this book please see our website at www.wiley.com. The right of the author to be identified as the author of this work has been asserted in accordance with the Copyright, Designs and Patents Act 1988. All rights reserved. No part of this publication may be reproduced, stored in a retrieval system, or transmitted, in any form or by any means, electronic, mechanical, photocopying, recording or otherwise, except as permitted by the UK Copyright, Designs and Patents Act 1988, without the prior permission of the publisher. Wiley also publishes its books in a variety of electronic formats. Some content that appears in print may not be available in electronic books. Designations used by companies to distinguish their products are often claimed as trademarks. All brand names and product names used in this book are trade names, service marks, trademarks or registered trademarks of their respective owners. The publisher is not associated with any product or vendor mentioned in this book. This publication is designed to provide accurate and authoritative information in regard to the subject matter covered. It is sold on the understanding that the publisher is not engaged in rendering professional services. If professional advice or other expert assistance is required, the services of a competent professional should be sought. The publisher and the author make no representations or warranties with respect to the accuracy or completeness of the contents of this work and specifically disclaim all warranties, including without limitation any implied warranties of fitness for a particular purpose. This work is sold with the understanding that the publisher is not engaged in rendering professional services. The advice and strategies contained herein may not be suitable for every situation. In view of ongoing research, equipment modifications, changes in governmental regulations, and the constant flow of information relating to the use of experimental reagents, equipment, and devices, the reader is urged to review and evaluate the information provided in the package insert or instructions for each chemical, piece of equipment, reagent, or device for, among other things, any changes in the instructions or indication of usage and for added warnings and precautions. The fact that an organization or Website is referred to in this work as a citation and/or a potential source of further information does not mean that the author or the publisher endorses the information the organization or Website may provide or recommendations it may make. Further, readers should be aware that Internet Websites listed in this work may have changed or disappeared between when this work was written and when it is read. No warranty may be created or extended by any promotional statements for this work. Neither the publisher nor the author shall be liable for any damages arising herefrom. Library of Congress Cataloging-in-Publication Data Laurence, Christian. Lewis basicity and affinity scales : data and measurement / Christian Laurence, Jean-François Gal. p. cm. Includes bibliographical references and index. ISBN 978-0-470-74957-9 1. Acids–Basicity. I. Gal, Jean-François. II. Title. QD477.L38 2009 546 .24–dc22

2009030807

A catalogue record for this book is available from the British Library. ISBN 9780470749579 Typeset in 10/12pt Times by Aptara Inc., New Delhi, India. Printed and bound in Great Britain by CPI Antony Rowe, Chippenham, Wiltshire

To Eliane Valérie, Anne Hugo, Clément, Simon and Louis C. Laurence

In memory of my mother who enabled me to fulfill my vocation as a researcher J.-F. Gal

Contents Preface 1

2

page ix

Lewis Basicity and Affinity Measurement: Definitions and Context 1.1 The Brönsted Definition of Acids and Bases 1.2 Scales of Brönsted Basicity and Affinity in Solution 1.3 Scales of Brönsted Basicity and Affinity in the Gas Phase 1.4 The Lewis Definition of Acids and Bases 1.5 Quantum Chemical Descriptions of Lewis Acid/Base Complexes 1.5.1 Valence-Bond Model 1.5.2 Perturbation Molecular Orbital Theory 1.5.3 Variational Supermolecular Method and Energy Decomposition Schemes 1.5.4 Natural Bond Orbital Theory 1.5.5 Quantum Theory of Atoms in Molecules 1.6 Measurement of Lewis Basicity 1.6.1 Gas-phase Reactions 1.6.2 Solution Reactions 1.6.3 Standard State Transformations 1.6.4 Choice of Solvent 1.7 Measurement of Lewis Affinity 1.8 The Role of the Solvent 1.9 Spectroscopic Scales of Basicity (Affinity) 1.10 Polybasic Compounds 1.11 Attempts at a Quantitative Formulation of the Lewis Definition of Acids and Bases 1.11.1 Hard and Soft Acids and Bases 1.11.2 The ECW and ECT Models 1.11.3 The Beta and Xi Equation 1.11.4 A Chemometric Approach 1.11.5 Quantum Chemical Descriptors for Basicity Scales 1.12 Concluding Remarks and Content of Chapters 2–7 References

1 2 3 6 6 10 10 10

The Donor Number or SbCl5 Affinity Scale 2.1 Structure of SbCl5 Complexes 2.2 Definition of the Donor Number Scale 2.3 Experimental Determination of the Donor Number

71 71 73 73

12 17 18 20 21 22 23 24 24 29 34 38 42 42 47 52 53 56 58 60

viii

Contents

2.4 2.5

The Donor Number Scale: Data Critical Discussion References

74 80 81

3 The BF3 Affinity Scale 3.1 Structure of BF3 Complexes 3.2 Definition of the BF3 Affinity Scale 3.3 Experimental Determination of the BF3 Affinity Scale 3.4 The BF3 Affinity Scale: Data 3.5 Discussion 3.5.1 Medium Effects 3.5.2 Hardness of BF3 3.5.3 Comparison of the BF3 and SbCl5 Affinity Scales 3.5.4 Computation of the BF3 Affinity 3.6 Conclusion References

85 86 88 89 90 102 102 102 103 104 105 106

4 Thermodynamic and Spectroscopic Scales of Hydrogen-Bond Basicity and Affinity 4.1 Structure of Hydrogen-Bonded Complexes 4.2 Hydrogen-Bond Basicity Scales: Early Works 4.3 The 4-Fluorophenol Hydrogen-Bond Basicity Scale 4.3.1 Definition 4.3.2 Method of Determination 4.3.3 Polyfunctional Hydrogen-Bond Acceptors 4.3.4 Data 4.3.5 Range of Validity of the Scale 4.4 Hydrogen-Bond Affinity Scales: Early Studies 4.5 The 4-Fluorophenol Affinity Scale 4.6 Comparison of 4-Fluorophenol Affinity and Basicity Scales 4.7 Spectroscopic Scales 4.7.1 Infrared Shift of Methanol 4.7.2 Solvatochromic Shifts of 4-Nitrophenol and 4-Nitroaniline 4.8 Conclusion References

111 113 117 119 119 119 120 121 167 168 170 185 188 188 210 221 221

5 Thermodynamic and Spectroscopic Scales of Halogen-Bond Basicity and Affinity 5.1 Structure of Halogen-Bonded Complexes 5.2 The Diiodine Basicity Scale 5.2.1 Definition of the Scale 5.2.2 Methods for the Determination of Diiodine Complexation Constants 5.2.3 Temperature Correction 5.2.4 Solvent Effects 5.2.5 Data

229 231 237 237 238 239 239 243

Contents

6

7

ix

5.3 Is the Diiodine Basicity Scale a General Halogen-Bond Basicity Scale? 5.4 The Diiodine Affinity Scale 5.5 Spectroscopic Scales 5.5.1 Infrared Shifts of ICN, I2 and ICl 5.5.2 The Blue Shift of the Diiodine Visible Band 5.6 Conclusion References

283 285 286 286 306 309 309

Gas-Phase Cation Affinity and Basicity Scales 6.1 Cations as Lewis Acids in the Gas Phase 6.2 Structure of Cation/Molecule Adducts 6.3 Experimental Techniques for Measuring Gas-Phase Cation Affinities and Basicities 6.3.1 High-Pressure Mass Spectrometry (HPMS) 6.3.2 Collision-Induced Dissociation Threshold (CIDT) 6.3.3 Ligand-Exchange Equilibrium Measurements in Trapping Devices 6.3.4 Selected Ion Flow Tube (SIFT) 6.3.5 Kinetic Method 6.3.6 Radiative Association Kinetics (RAK) 6.3.7 Blackbody Infrared Radiative Dissociation (BIRD) 6.3.8 Vaporization and Lattice Energies 6.4 Ion Thermochemistry Conventions 6.5 Lithium, Sodium, Potassium, Aluminium, Manganese, Cyclopentadienylnickel, Copper and Methylammonium Cations Affinity and Basicity Scales 6.5.1 Lithium 6.5.2 Sodium 6.5.3 Potassium 6.5.4 Aluminium 6.5.5 Manganese 6.5.6 Cyclopentadienylnickel 6.5.7 Copper 6.5.8 Methylammonium 6.6 Significance and Comparison of Gas-Phase Cation Scales 6.6.1 Properties of Cations and Significance of MCB and MCA Scales 6.6.2 Relationship of MCA with MCB 6.6.3 The Computation of MCB and MCA Scales 6.6.4 MCA and MCB Scales and the Concept of a Cation/π Interaction 6.6.5 Conventional Versus Ionic Hydrogen-Bond Basicity and Affinity Scales 6.6.6 Comparison of Cation Basicity Scales References

323 323 326

The Measurement of Lewis Basicity and Affinity in the Laboratory 7.1 Calorimetric Determination of the BF3 Affinity of Pyridine by Gas/Liquid Reaction

334 334 335 336 337 337 338 338 339 339

340 340 346 353 354 354 360 366 371 370 370 381 382 383 386 387 389 401 401

x

Contents

7.1.1

7.2

7.3

7.4

7.5

7.6

7.7

Introduction: Principles and Difficulties in the Calorimetric Measurement of the Enthalpy of a Gas/Liquid Reaction 7.1.2 Reagents and Equipment 7.1.3 Experiment 7.1.4 Results Calorimetric Determination of the BF3 Affinity of Pyridine by Liquid/Liquid Reaction 7.2.1 Introduction: Measuring Relative BF3 Affinity by Ligand Exchange in Solution 7.2.2 Reagents and Equipment 7.2.3 Experiment 7.2.4 Results Determination by FTIR Spectrometry of the Complexation Constants of 4-Fluorophenol with Isopropyl Methyl Ketone and Progesterone 7.3.1 Introduction: Recognition of Progesterone by its Receptor 7.3.2 Reagents and Equipment 7.3.3 Experiment 7.3.4 Results and Discussion Determination by FTIR Spectrometry of the Complexation Enthalpy and Entropy of 4-Fluorophenol with Cyclopropylamine 7.4.1 Introduction 7.4.2 Reagents and Equipment 7.4.3 Experiment 7.4.4 Results 7.4.5 Comparison with Theoretical Calculations FTIR Determination of the OH Shift of Methanol Hydrogen Bonded to Pyridine, Mesitylene and N -Methylmorpholine 7.5.1 Introduction 7.5.2 Reagents and Equipment 7.5.3 Experiment 7.5.4 Results and Discussion Solvatochromic Shifts of 4-Nitrophenol upon Hydrogen Bonding to Nitriles 7.6.1 Introduction 7.6.2 Reagents and Equipment 7.6.3 Experiment 7.6.4 Results and Discussion Determination of the Complexation Constant of Diiodine with Iodocyclohexane by Visible Spectrometry 7.7.1 Introduction: Measuring the Weak Diiodine Basicity of Haloalkanes 7.7.2 The Rose–Drago Method 7.7.3 Reagents and Equipment 7.7.4 Experiment 7.7.5 Results and Discussion: Illustration of the HSAB Principle

401 403 404 405 406 406 407 407 407 408 408 409 409 410 413 413 414 414 414 417 418 418 418 419 419 420 420 421 421 422 424 424 424 426 426 427

Contents

7.8 Determination of the Complexation Enthalpy and Entropy of Diiodine with Dimethyl Sulfoxide by Visible Spectrometry 7.8.1 Introduction 7.8.2 Reagents and Equipment 7.8.3 Experiment 7.8.4 Results 7.8.5 Discussion 7.9 FTIR Determination of the Shift of the I C Stretching of Iodine Cyanide upon Halogen Bonding to Phosphine Chalcogenides 7.9.1 Introduction 7.9.2 Reagents and Equipment 7.9.3 Experiment 7.9.4 Results and Discussion: ∆ν(ICN) as a Spectroscopic Scale of Halogen-Bond Affinity 7.10 Blue Shift of the Visible Diiodine Transition Upon Halogen Bonding to Pyridines 7.10.1 Introduction 7.10.2 Reagents and Equipment 7.10.3 Experiment 7.10.4 Results and Discussion: Substituent Effects 7.11 Mass Spectrometric Determination of the Gas-Phase Lithium Cation Basicity of Dimethyl Sulfoxide and Methyl Phenyl Sulfoxide by the Kinetic Method 7.11.1 The Kinetic Method 7.11.2 Reagents and Equipment 7.11.3 Experiment 7.11.4 Data Treatment 7.11.5 Discussion: Substituent Effects References Index

xi

429 429 429 429 430 432 434 434 435 435 435 436 436 437 437 438

439 439 440 441 442 444 445 447

Preface In 1923, G.N. Lewis proposed an electronic definition of acids and bases founded on electron pair sharing. Compared with the protonic definition of Brönsted based on proton exchange, the new Lewis definition broadened considerably the field of acid/base reactions. It incorporates heterolysis, coordination, solvation, complexation, hydrogen-bond formation, halogen-bond formation and electrophilic and nucleophilic reactions into acid/base chemistry. It is therefore not surprising that discussions of Lewis acidity and Lewis basicity appear in almost every textbook on general, organic and inorganic chemistry. A major criticism, however, is often made of the Lewis definition. Contrary to the Brönsted definition, it is not possible to establish any universal order of acid or base strength. In the Brönsted definition, the proton is used as the reference and the quantitative study of proton exchange reactions between bases, by means of electrochemical or spectrometric methods, enabled the strength of several thousand Brönsted acids and bases to be measured unambiguously. Thermodynamic databases for proton exchange in the gas phase (NIST webbook, http://webbook.nist.gov.) or in water (e.g. D.D. Perrin, Dissociation Constants of Organic Bases in Aqueous Solution, Butterworths, London, 1965, and supplement, 1972) are therefore extremely useful in analytical, organic and inorganic chemistry and in biochemistry. In contrast, in the Lewis definition there is no single reference that is naturally operational. Since there is no obvious reason to choose one reference rather than another, there are potentially as many acidity or basicity scales as possible references. Concerning basicity, however, which is the subject of this book, the statistical treatment of various scales of Lewis basicity, as well as theoretical studies, show that a limited number of factors influence the strength of Lewis bases. Consequently, the judicious choice of a few reference Lewis acids should allow Lewis basicity scales to be constructed and used as a general guide to basicity. Although none of these scales can be considered as universal, each will have a domain of validity that is sufficiently wide to be useful in many branches of chemistry and biochemistry. It is with this objective that this book presents thermodynamic and spectroscopic data on the strength of Lewis bases coming from both the literature and our laboratories. We do not aim to provide exhaustive scales but rather a selective guide. From a mass of data, sometimes contradictory and often lacking consistency, we have chosen typical reference Lewis acids. These are SbCl5 , BF3 , 4-FC6 H4 OH, CH3 OH, 4-NO2 C6 H4 OH, 4-NO2 C6 H4 NH2 , I2 , ICl, ICN, Li+ , Na+ , K+ , Al+ , Mn+ , CpNi+ , Cu+ and CH3 NH+ 3 . This choice is justified in the first chapter. For each of these acids, we have selected only the data determined by the most accurate techniques and/or the most reliable methods. In cases of doubt, some measurements from the literature have been repeated in our laboratories. Additional measurements have also been carried out in order to fill significant gaps. Finally, data have

xiv

Preface

been made homogeneous either by means of the usual thermodynamic relationships, in order to refer to the same standard state and to the same temperature, or by means of extrathermodynamic relationships, specifically established, in order to refer them to the same solvent. In all, more than 2400 equilibrium constants of acid/base reactions, and thus of Gibbs energies, about 1500 complexation enthalpies and, for spectroscopic scales, nearly 2000 infrared and ultraviolet shifts of absorption bands upon complexation are gathered together in this book. We expect that this collection of data will enable experimental chemists to understand and predict better the numerous chemical, physical and biological properties that depend upon Lewis basicity. Indeed, basicity parameters can be introduced as explanatory variables in many linear free energy relationships, linear solvation energy relationships, structure– property relationships and structure–biological activity relationships. Chemometricians may be able to construct many kinds of data matrices from the Lewis basicity scales in this book. For example, matrices might refer to a given family of bases (e.g. nitriles), or a given property (e.g. complexation enthalpies), or a given type of complex (e.g. hydrogen-bonded complexes), or a combination of families, properties and types of complexes. The analysis of judiciously constructed matrices by appropriate chemometric methods should enable the intrinsic fundamental effects in the data matrix to be extracted and used to predict the Lewis basicity and many basicity-dependent properties. In addition, measured experimental basicities and affinities are essential to computational chemists for the validation, calibration and establishment of reliable computational methods for quantifying and explaining intermolecular forces and the chemical bond. In fact, the formation of Lewis acid/base adducts covers a wide variety of bond-forming processes from the weak hydrogen bond to the strong dative bond or ion/molecule bond. Scepticism about the quantitative usefulness of the Lewis concept of acids and bases is still frequently encountered in chemistry textbooks. We expect that this book will demonstrate that quantitative data exist for acid/base systems other than those involving proton donors and acceptors, and will encourage textbook authors to go beyond a mere qualitative presentation of the Lewis acid/base concepts. The hope for a quantitative Lewis acid/base chemistry has motivated our research at the Universities of Nantes and Nice-Sophia Antipolis for more than 40 years. During this long period, many people have contributed to this research. We thank our colleagues in Nantes and Nice (M. Berthelot, J. Graton, M. Helbert-Nicolet, B. Illien, J.-Y. Le Questel, P.C. Maria and P. Nicolet) for their generous help, the technical staff of our laboratories (F. Besseau, M. Decouzon, M. Luçon and A. Planchat) for their skillful assistance and our post-doctoral, PhD and Master’s Degree students, whose names appear in the references, for their efforts. We especially acknowledge the following collaborations. Prof. Michel Berthelot contributed very significantly to the construction of the following scales: 4-fluorophenol basicity and affinity, diiodine basicity, methanol, I2 and ICl shifts. Most of the 4-fluorophenol affinity values were obtained by François Besseau. Dr Michèle Decouzon dedicated much time to the maintenance of and measurements made with the FTICR mass spectrometer in Nice. Dr Mohamed Jamal El Ghomari (Marrakech) contributed to a first version of the diiodine basicity scale. Dr Maryvonne Helbert carried out the measurements of the methanol, I2 and ICl spectroscopic shifts and many diiodine and 4-fluorophenol complexation constants. Dr Josef Kaczmarek (Gdansk) carried out the measurements of the ICN

Preface

xv

spectroscopic shifts. Dr Maryvonne Luçon measured many diiodine complexation constants and IR shifts of 4-nitroaniline, constructed the methanol, pyrrole and cyanoacetylene affinity and basicity scales and recorded the FTIR and UV spectra of laboratory experiments 7.3 to 7.10. Prof. Pierre-Charles Maria participated very actively in the development of the BF3 affinity and lithium cation basicity scales. Dr Raphael Notario and Prof. Jose Luis Abboud (Madrid) helped in the construction of the diiodine basicity scale. Prof. Pierre Nicolet co-directed the work on the 4-nitrophenol and 4-nitroaniline solvatochromic scales. Prof. Ewa Raczyñska (Warsaw) studied the hydrogen-bond basicity of amidines and other bases. Prof. Manuel Yán˜ ez (Madrid) performed high-level ab initio calculations on Al+ , Mn+ and CpNi+ complexes in order to anchor the corresponding affinity and/or basicity scales. The electronic version of the manuscript was produced entirely by Dr Maryvonne Luçon. Dr Charly Mayeux prepared the figures illustrating the 3D geometry of the complexes. Mrs Carol Robins revised the English of the whole manuscript. We wish to express our gratitude to them for their assistance in the preparation of this book. We have benefited from the consummate professionalism of the editorial staff at Wiley. It was a pleasure to work with all of them. Finally, without the understanding and patience of our wives, Eliane and Juliette, we could never have accomplished this work.

1 Lewis Basicity and Affinity Measurement: Definitions and Context Two definitions of acids and bases are used nowadays, the Brönsted definition and the Lewis definition. This book deals with the quantitative behaviour of Lewis bases. However, since Lewis bases are also Brönsted bases, this chapter begins with a short presentation of the Brönsted definition and of the quantitative behaviour of Brönsted bases [1]. The Lewis definition and the many ways for its quantification will then be studied. This introductory chapter is intended to help in the understanding and use of the tables in Chapters 2–6, which contain quantitative data on Lewis basicity and affinity, and not to discuss the Lewis acid/base concept in depth. This subject has been excellently treated in a book [2] and a review [3] by Jensen, and books and chapters by Mulliken and Person [4], Gur’yanova et al. [5], Drago [6], Finston and Rychtman [7] and Weinhold and Landis [8], to quote just a few. As far as possible, we have followed the IUPAC recommendations for the names and symbols of physical and chemical quantities (http://goldbook.iupac.org/) and have used the international system of units (SI) and the recommended values of the fundamental physical constants (http://physics.nist.gov/cuu/). Units that are not part of the SI have been used ˚ = 10−10 m), in appropriate contexts. These are: litre (1 l = 10−3 m3 ), a˚ ngström (1 A −19 −30 J), Debye (1 D ≈ 3.336 × 10 C m) and bar electronvolt (1 eV ≈ 1.602 18 × 10 (1 bar = 105 Pa). In tabulating thermodynamic and spectroscopic basicity scales, 1 : 1 complexation constants, Gibbs energies, enthalpies, entropies and ultraviolet (UV) and infrared (IR) spectral shifts are therefore given in l mol−1 (identical with dm3 mol−1 ), kJ mol−1 , J K−1 mol−1 and cm−1 , respectively. Logarithms of equilibrium constants (log K) are to base 10 and without units since the calculated quantity is log (K/1 l mol−1 ). In naming compounds, we have not always followed the nomenclature rules. We have sometimes preferred the common name found in most chemical catalogues. For clarity, the Lewis Basicity and Affinity Scales: Data and Measurement C 2010 John Wiley & Sons, Ltd

Christian Laurence and Jean-François Gal

2


Table 1.1 Symbols for families of Lewis bases used in the graphs. Family Carbon π bases Aromatics, alkenes, alkynes Oxygen bases Single-bonded Carbonyls Sulfinyls Phosphoryls, arsine oxides N-Oxides Nitros, sulfonyls

Symbol

Family

Symbol

×

Nitrogen bases sp-Hybridized

sp2 -Hybridized sp3 -Hybridized Phosphorus, arsenic bases Sulfur bases Single-bonded

Family

Symbol

Selenium bases Single-bonded, seleno-carbonyls, selenophosphoryls −

Thiocarbonyls Thiophosphoryls

Halogen bases Fluoroalkanes

+

Chloroalkanes Bromoalkanes Iodoalkanes Miscellaneous bases

name is followed, in most tables in Chapters 2–6, by a formula that allows the drawing of the structure, or by the drawing itself. In the graphs, in order to facilitate the identification of family-dependent trends, bases are labelled as summarized in Table 1.1, unless otherwise stated in the legend of the graph.

1.1

The Brönsted Definition of Acids and Bases

A powerful definition of acids and bases was proposed in 1923 by J.N. Brönsted [9], namely an acid is a species capable of donating a proton, and a base is a species capable of accepting a proton. This can be expressed by the scheme A B + H+

(1.1)

where the acid A and the base B are termed a conjugate acid/base pair. Equation 1.1 represents a hypothetical scheme used for defining an acid and a base rather than a reaction. Indeed, reaction 1.1 cannot actually occur in a solvent because the bare proton H+ cannot exist in solution, and cannot be studied directly in the gas phase because of the extremely large endoergic values involved. The only reactions between Brönsted acids and Brönsted bases that can be observed in solution and studied directly in the gas phase are reactions of proton exchange between two conjugate acid/base pairs A1 /B1 and A2 /B2 A1 + B2 A2 + B1

(1.2)

For example, in aqueous solutions, the acid CH3 COOH reacts with water acting as a base: CH3 COOH + H2 O H3 O+ + CH3 COO−

(1.3)

Lewis Basicity and Affinity Measurement: Definitions and Context

3

Table 1.2 Some types of neutral Bronsted acidsa . ¨ O H acids Inorganic oxyacids Carboxylic acids Phenols, alcohols, water N H acids C H acids S H acids X H acids a

HNO3 , H2 SO4 , H3 PO4 , HClO4 RCOOH ArOH, ROH, H2 O ArNH2 , RSO2 NH2 , RCONH2 , HNCS, HNCO, HN3 HC N, RC CH, HC(NO2 )3 ArSH, H2 S HF, HCl, HBr, HI

In the formulae, R is an alkyl group and Ar an aryl group.

and the base NH3 reacts with water acting as an acid: − H2 O + NH3 NH+ 4 + OH

(1.4)

In the gas-phase reaction 1.5, the proton is exchanged between the ammonium ion/ammonia and the pyridinium ion/pyridine pairs: + NH+ 4 + C5 H5 N C5 H5 NH + NH3

(1.5)

Any compound containing hydrogen can, in principle, be regarded as a Brönsted acid, but in many of them (e.g. most hydrocarbons) the tendency to lose a proton is so small that they do not show acidic behaviour under ordinary conditions. Examples of neutral Brönsted acids are given in Table 1.2. The same kind of practical restriction should be applied to Brönsted bases. Neutral molecules or atoms can attach a proton in the gas phase because of the tremendous acidity of the bare proton: even rare gases may be protonated in the gas phase. For the liquid phase, superacid systems (such as HF/SbCl5 that are more acidic than 100% sulfuric acid) can also protonate many molecules [10]. For example, the protonated form of methane, CH5 + , which was discovered in the gas phase by mass spectrometry in the 1950s, has also been reported in superacid solutions. However, the important bases in chemistry are (i) anions and (ii) molecules containing elements of groups 15 and 16 with unshared electron pair(s).

1.2

Scales of Brönsted Basicity and Affinity in Solution

Brönsted definitions are easily translated into quantitative measurements. The equilibrium constant of reaction 1.2, K = (A2 )(B1 )/(A1 )(B2 ), where parentheses denote activities, is equal to the ratio of the hypothetical constants (B1 )(H+ )/(A1 ) and (B2 )(H+ )/(A2 ). K will therefore measure the ratio of the acid strengths of A1 and A2 , or the ratio of the base strengths of B2 and B1 . Since these two ratios are equal, it is not necessary to give separate definitions of base strength and acid strength. The base strength of any base B is usually given by the acid strength of its conjugate acid A. Thus, for the pair C5 H5 NH+ /C5 H5 N, the base strength of pyridine is described in terms of the acid strength of the pyridinium ion. It is not possible to measure the absolute strength of an acid or base in solution but strengths can be measured relative to some standard pair, A◦ /B◦ . The acid strength of the

4


studied pair, A/B, is then given by the equilibrium constant of the reaction A + B◦ A◦ + B

(1.6)

The standard pair, A◦ /B◦ , is usually chosen to be the acid/base pair of the solvent. In aqueous solutions, the pair H3 O+ /H2 O is commonly preferred to the other possible pair, H2 O/OH− . The strength of any acid A is then measured by the equilibrium constant of the reaction A + H2 O H3 O+ + B

(1.7)

When measurements are made in dilute aqueous solution, the concentration of water remains essentially constant and its activity can be taken as unity. The strength of the acid A is then measured by the acid dissociation constant: K a = (B) (H3 O+ )/(A)

(1.8)

The strength of a neutral base B is described in terms of the K a of its conjugate acid BH+ , usually denoted K BH+ : K BH+ = (B) (H3 O+ )/(BH+ )

(1.9)

Since the observed equilibrium constants vary over many powers of 10, the convention to use the operator p ≡ −log10 was adopted, leading to the quantity pK BH+ : pK BH+ = − log10 K BH+

(1.10)

Clearly, a large positive value of pK BH+ describes a strong and a small or negative value describes a weak Brönsted base. Tables of pK BH+ in aqueous solution have been compiled by Perrin [11]. They cover the literature until 1972 and contain more than 7000 organic bases, mainly sp2 and sp3 nitrogen bases. Many of the carbon, oxygen, sulfur and sp nitrogen bases are not protonated in dilute acid solutions, so that solutions with variable concentrations of a strong acid have to be used. In such media, K BH+ values cannot be calculated without formulating some extrathermodynamic assumption. The pK BH+ values of many weak bases have been carefully measured by Scorrano et al. [12–21]. Table 1.3 gives selected pK BH+ values. The pK BH+ is directly converted into the Gibbs energy change of the proton exchange as follows: ∆G ◦ = ln(10)RTpK BH+

(1.11)

The literature is poorer in enthalpies of proton exchange reactions. However, Arnett et al. have established an enthalpic scale of Brönsted basicity [22–25] (i.e. a Brönsted affinity scale) from the heats of protonation of many bases in fluorosulfuric acid. The heats of protonation (ionization), ∆H i , correspond simply to the heat of transfer of the base from infinite dilution in the inert solvent CCl4 to infinite dilution in the (often) completely protonating solvent HSO3 F. A surprisingly good correlation (r = 0.986, n = 55, s = 5.4 kJ mol−1 ) is obtained [23] between the enthalpies of protonation and the corresponding aqueous pK BH+ values. Selected values of ∆H i are given in Table 1.3.


5

Table 1.3 Thermodynamic parameters for protonation of organic bases: pK BH+ in water, ∆Hi (kJ mol−1 ) in fluorosulfuric acid and GB and PA (kJ mol−1 ) in the gas phase. Base Hexamethylbenzene Methylamine Ethylamine Dimethylamine Diethylamine Di-n-butylamine Trimethylamine Triethylamine Tri-n-butylamine Quinuclidine Triphenylamine 3,5-Dichloropyridine 2-Bromopyridine 2-Chloropyridine 3-Bromopyridine Quinoline Pyridine 4-Methylpyridine 2,6-Dimethylpyridine 2,4,6-Trimethylpyridine Aniline N,N-Dimethylaniline Methanol Ethanol Water Dimethyl ether Diethyl ether Tetrahydrofuran Benzaldehyde Acetophenone Benzophenone Acetone Diethyl ketone Dicyclopropyl ketone Methyl acetate Methyl propionate Methyl benzoate N,N-Dimethylacetamide N,N-Dimethylformamide N-Methylpyrrolidone Tetramethylurea Dimethyl sulfoxide Hexamethylphosphoric triamide Triphenylphosphine oxide Pyridine N-oxide Nitrobenzene N,N-Dimethylthioacetamide Methyl sulfide Ethyl sulfide Triphenylphosphine

pK BH+ −14.65 10.65 10.68 10.78 11.02 11.25 9.80 10.72 9.93 11.15 0.67 0.90 0.72 2.85 4.85 5.20 6.03 6.72 7.43 4.60 5.15 −2.05 −1.94 −1.74 −2.48 −2.39 −4.48 −3.87 −4.71 −3.06 −3.88 −2.40 −3.90 −4.37 −7.05 −0.21 −1.13 −0.71 −0.14 −1.54 −0.97 0.8 −2.25 −6.99 −6.68

−∆Hi 193.9 195.9 197.4 199.5 194.1 196.8 205.7 189.2 191.6 79.9 128.4 126.2 132.5 144.9 150.9 161.3 163.4 170.3 178.5 142.3 157.7 79.9 68.6 79.8 82.0 67.4 79.1 70.7 79.9 87.0

133.9 123.4 131.0 157.3 119.7 87.4 139.7 27.6 79.5 120.0

GB

PA

836.0 864.5 878.0 896.5 919.4 935.3 918.1 951.0 967.6 952.5 876.4

860.6 899.0 912.0 929.5 952.4 968.5 948.9 981.8 998.5 983.3 908.9

873.0 869.0 878.2 921.4 898.1 915.3 931.1

904.8 900.9 910.0 953.2 930.0 947.2 963.0

850.6 909.2 724.5 746.0 660.0 764.5 801.0 794.7 802.1 829.3 852.5 782.1 807.0 850.6 790.7 799.2 819.5 877.0 856.6 891.6 899.6 853.7 928.7 876.4 892.9 769.5 894.4 801.2 827.0 940.4

882.5 941.1 754.3 776.4 691.0 792.0 828.4 822.1 834.0 861.1 882.3 812.0 836.8 880.4 821.6 830.2 850.5 908.0 887.5 923.5 930.6 884.4 958.6 906.2 923.6 800.3 925.3 830.9 856.7 972.8

6

1.3


Scales of Brönsted Basicity and Affinity in the Gas Phase

Various mass spectrometric techniques permit the study of proton transfer reactions in the gas phase and the definition of Brönsted basicity scales free of solvent effects [26]. The gas-phase basicity GB and the proton affinity PA of a base B are defined as the standard Gibbs energy change and the standard enthalpy change, respectively, of the formal deprotonation reaction 1.12: BH+ → B + H+ ∆H ◦ = proton affinity = PA ∆G ◦ = gas-phase basicity = GB

(1.12)

Unfortunately, this terminology, currently in use, is not completely correct since an affinity is a chemical potential (a ∆G value) whereas the proton affinity is an enthalpy. An alternative terminology for ∆H ◦ might be ‘enthalpy of basicity’, but it seems unrealistic to propose a change of terminology now considering the accepted practice. The absolute basicity and affinity cannot be obtained directly because the gas-phase reaction 1.12 is extremely endoergic and endothermic. It is common practice to resort to thermodynamic cycles, involving enthalpies of formation and dissociation thresholds, to calculate absolute PAs. The transformation of absolute PAs into absolute GBs (Equation 1.13) requires the evaluation of the entropy of basicity (Equation 1.14) (mainly through quantum chemical calculations today): GB = PA − T ∆S ◦ ∆S ◦ = S ◦ (B) + S ◦ (H+ ) − S ◦ (BH+ )

(1.13) (1.14)

The number of absolute PA and GB values that can be accurately evaluated is very limited. In fact, most parts of the scales are obtained by measuring the relative basicity of an unknown using a reference base B◦ of known GB. Relative basicities, designated ∆GB, are obtained from equilibrium constants K of the proton exchange reaction 1.15 between bases B and B◦ : BH+ + B◦ B◦ H+ + B ∆GB = −RT lnK = GB(B) − GB(B◦ )

(1.15) (1.16)

The known basicities span a very wide range of about 1300 kJ mol−1 from He to Cs2 O. However, the basicity of the majority of organic bases falls within 700–1000 kJ mol−1 . A selection is presented in Table 1.3. Thousands of PA and GB values have been critically compiled by Hunter and Lias [27, 28].

1.4

The Lewis Definition of Acids and Bases

In the original Lewis definition (1923 [29], 1938 [30]), acids are electron-pair acceptors and bases are electron-pair donors. The fundamental reaction between a Lewis acid A and a Lewis base B is the formation of a complex (or adduct or coordination compound or addition compound) A–B (reaction 1.17): A + : B A–B

(1.17)


7

In this reaction, the unshared electron pair of the base forms a coordinate covalent bond (or dative bond or dipolar bond) with an electron-deficient atom of the acid. The archetype of a Lewis acid/base reaction is BF3 + : NH3 F3 B–NH3

(1.18)

BF3 is a Lewis acid because the boron atom has only six electrons in its valence shell and, having room for eight, can accept the lone pair of the nitrogen atom of ammonia. The proton is a Lewis acid because it can accept an electron pair into its empty 1s atomic orbital. It follows that all Brönsted bases are Lewis bases. All Brönsted acids are also Lewis acids because they are hydrogen-bond donors, that is, electron acceptors (see below). However, a much wider range of species can be classified as Lewis acids than can be classified in the Brönsted scheme. The translation of Lewis’s definition into quantummechanical terms by Mulliken (1952) [31] further widened the definition, so as to include those reagents that donate or accept a fraction, possibly very small, of an electron. With this extension, the compounds in Tables 1.4 and 1.5 are considered as Lewis acids (electron

Table 1.4 Examples of Lewis acids. Metals: M Cations Proton: H+ Metallic: Mn+ Organometallic: CH3 Hg+ Halogens: I+ Carbocations: CH3 + Covalent metal halides, hydrides or alkyls: MX n , MHn , MRn Group 4: TiCl4 Group 8: FeCl3 Group 12: ZnCl2 , CdI2 , HgCl2 Group 13: BF3 , BCl3 , BH3 , BMe3 AlCl3 , AlMe3 GaCl3 , GaH3 , GaMe3 Group 14: SnCl4 Group 15: SbCl3 , SbCl5 Halogen-bond donors Dihalogens: I2 , Br2 , Cl2 Interhalogens: ICl, IBr, ClF, BrCl Organic halogens: IC N, ICF3 , IC CR Hydrogen-bond donors (Bronsted acids) ¨ OH: RCOOH, ArOH, ROH, H2 O NH: RCONH2 , ArNH2 , HNCS CH: CHCl3 , RC CH SH: ArSH XH: HF, HCl π Acceptors SO2 , SO3 Ethylenic, acetylenic, aromatic hydrocarbons substituted with electron-withdrawing groups Quinones

8


Table 1.5 Important neutral and anionic Lewis (Bronsted) bases classified by their donor ¨ atom(s). Hydrogen Carbon

Nitrogen

Phosphorus Arsenic Oxygen

Sulfur

Selenium Tellurium Halogens

Anion H− Alkenes, alkynes, arenes, cyclopropanes Carbon monoxide CO Isonitriles R N C Carbanions, CN− , RC C− sp3 NH3 , primary amines RNH2 , secondary amines RR NH, tertiary amines RR R N sp2 Six-membered aromatic N-heterocycles (e.g. pyridines) Five-membered aromatic N-heterocycles (e.g. imidazoles) Imines R2 C NR , amidines R2 NC NR , oximes R2 C NOR sp Nitriles RC N Anions NH2 − , N3 − , SCN− , OCN− Phosphines RR R P Arsines RR R As Water H2 O, alcohols ROH, ethers ROR , peroxides ROOR Carbonyls RCOR : aldehydes, ketones, lactones, esters, carbonates, amides, lactams Sulfonyls RR SO2 Sulfinyls RR SO Seleninyls RR SeO Phosphoryls RR R PO Arsine oxides RR R AsO Amine oxides RR R NO Nitrosos R N O Nitros RNO2 Isocyanates R N C O Anions OH− , RO− , RCOO− , CO3 2− , ClO4 − , PO4 3− , HPO4 2− , H2 PO4 − , SO4 2− , HSO4 − , SO3 2− , NO3 − , NO2 − H2 S, thiols RSH, thioethers RSR , disulfides RSSR Thiocarbonyls RCSR : thioamides, thioureas Isothiocyanates RN C S Thiophosphoryls RR R PS Anions SH− , RS− , SCN− Selenoethers RSeR Selenocarbonyls RCSeR : selenoamides, selenoureas Selenophosphoryls RR R PSe Telluroethers RTeR Haloalkanes RF, RCl, RBr, RI Anions F− , Cl− , Br− , I−

acceptors) and Lewis bases (electron donors), respectively, and the following reactions (1.19)–(1.27) are considered today as Lewis acid/base reactions: Co2+ + 6 H2 O [Co(OH2 )6 ]2+ Ag+ + C6 H6 Ag+ · · · C6 H6 Ni + 4 CO Ni(CO)4 −

CO2 + OH SiF4 + 2 F−

HCO− 3 SiF2− 6

(1.19) (1.20) (1.21) (1.22) (1.23)


I2 + Et2 O I−I · · · OEt2 (NC)2 C = C(CN)2 + C6 H6 (NC)2 C = C(CN)2 , C6 H6 CH3 OH + CH3 C≡N CH3 OH · · · N≡CCH3 SO3 + C5 H5 N O3 S · · · NC5 H5

9

(1.24) (1.25) (1.26) (1.27)

In addition to the elementary reaction A + B → AB, other Lewis acid/base reactions are as follows: (i) Displacement reactions of one Lewis base by another: A−B1 + B2 A−B2 + B1

(1.28)

All Brönsted proton transfer reactions fit into this type. (ii) Displacement reactions of one Lewis acid by another: A1 −B + A2 A2 −B + A1

(1.29)

An interesting example [32] is the displacement of water hydrogen bonded to polyamines (or their N-oxides) by a halogen-bond donor:

N

N

HOH + ICF2CF2I

N

N

ICF2CF2I

+ H2O

(1.30)

This displacement reaction can be used to obtain hygroscopic bases in anhydrous form. (iii) Double displacement reactions: A1 −B1 + A2 − B2 A1 −B2 + A2 −B1

(1.31)

Many heterolytic reactions can be classified within this group, insofar as reactants and products are formally dissected into Lewis acids and bases. For example, in reaction 1.32 the reactants methanol and hydrogen iodide are both formally the products of the acids CH3 + and H+ and the bases OH− and I− : CH3 −OH + H−I CH3 −I + H−OH

(1.32)

Most reactions occurring in an amphoteric solvent αβ, with acid site α and basic site β coordinated to the two reactants, fall into this category: A · · · βα + βα · · · B A−B + βα · · · βα

(1.33)

With the extension of the original Lewis definition, and its application to many fields of chemistry, many terms specific to those fields have been substituted for the very general ‘Lewis acid’ and ‘Lewis base’ terms. These synonyms are collected in Table 1.6.

10


Table 1.6 Synonyms of ‘Lewis acid’ and ‘Lewis base’ used in various fields of chemistry. Field

Lewis acid synonym

Lewis base synonym

¨ Bronsted acid/base chemistry General chemistry Organic chemistry (kinetics) Coordination chemistry Ionic bond Cation solvation Anion solvation Hydrogen bonding Halogen bonding

Proton donor Electron acceptor Electrophile Central metallic atom (ion) Cation Cation Acidic solvent Hydrogen-bond donor Halogen-bond donor

Proton acceptor Electron donor Nucleophile Ligand Anion Basic solvent Anion Hydrogen-bond acceptor Halogen-bond acceptor

1.5 1.5.1

Quantum Chemical Descriptions of Lewis Acid/Base Complexes Valence-Bond Model

According to the Mulliken valence-bond model [4, 31], the complex AB between a Lewis acid A and a Lewis base B may be described by the wavefunction ΨAB = a Ψ0 (A, B) + b Ψ1 (A−− B+ )

(1.34)

The no-bond wavefunction Ψ 0 describes a structure in which the binding of A and B is effected by electrostatic forces such as those between permanent dipoles of A and B, the permanent dipole of A(B) and the induced dipole of B(A), and fluctuating dipoles of A and B (London dispersion forces). The dative-bond wavefunction Ψ 1 corresponds to a structure, sometimes called a charge-transfer structure, in which an electron has been transferred from the base B (the donor) to the acid A (the acceptor). Equation 1.34 shows that, by varying the ratio of weighting coefficients a and b, all degrees of electron donation are possible. The energy EAB of a simple 1 : 1 adduct associated with the wavefunction Ψ AB is given by the second-order perturbation theory E AB = E 0 −

(β 01 − E 0 S01 )2 E1 − E0

(1.35)

where E 0 = Ψ0 |H |Ψ0 , E 1 = Ψ1 |H |Ψ1 , β01 = Ψ0 |H |Ψ1 , S01 = Ψ0 |Ψ1 and H represent the energy of the no-bond structure, the energy of the dative-bond structure, the resonance integral, the overlap integral and the total exact Hamiltonian respectively. It can be seen from Equation 1.35 that the energy of the acid/base complex is the sum of an electrostatic energy term (E0 ) and a second term that is due to covalency (in the case where A and B are neutral, Ψ 1 corresponds to a covalent binding involving the odd electrons in A− and B+ ). 1.5.2

Perturbation Molecular Orbital Theory

A general description of chemical reactivity has been given by Klopman [33]. It can be applied [2] to a simple Lewis acid–base reaction A + B → AB. Klopman [33] and Jensen


11

[2] used a perturbation molecular orbital (MO) theory and wavefunction Ψ mn (AB): Ψmn (AB) = a Ψn (A) + b Ψm (B)

(1.36)

where a and b are weighting coefficients and Ψ m and Ψ n are expressed as linear combinations of atomic orbitals (AOs), φ: crm φr Ψn = csn φs (1.37) Ψm = r

s

crm (csn )

where is the AO coefficient of the mth (nth) MO at atom r (s). If the interaction occurs primarily between a single donor atom r on B and a single acceptor atom s on A, the energy change, ∆E, upon formation of the adduct AB in the gas phase can be approximated by the sum of three terms: occ unocc crm csn βr s 2 Qr Q s +2 + ∆E(repulsion) (1.38) ∆E ≈ Rr s Em − En m n base acid

The first term is the coulombic attraction between the total net opposite charges Qr and Qs of the interacting atoms r and s of B and A at a distance Rrs . The second term is a second-order orbital perturbation due to the attractive interactions between the occupied orbitals on B (A) and the unoccupied orbitals on A(B). β rs is the resonance integral between the AOs of r and s, and Em (En ) is the energy of orbital m (n) of B (A) in the field of A (B). The last term is a first-order orbital perturbation due to the repulsive interactions between the filled orbitals of A and B. The largest contribution to the double summation over orbital pairs in the second term will arise when the denominator is smallest. This occurs when the energies of the highest occupied MO (HOMO) of the base and the lowest unoccupied MO (LUMO) of the acid become closer. By considering only this contribution in the double summation (the so-called frontier orbitals approximation), Equation 1.38 simplifies to 2 2 crHOMO csLUMO βr s Qr Q s + + ∆E(repulsion) (1.39) ∆E ≈ Rr s E (HOMO)B − E (LUMO)A On the basis of Equation 1.39, the Lewis acid/base reactions can be divided into the categories of ‘charge-controlled’, those dominated by the first term, and ‘orbital-controlled’, those dominated by the second term. The factors determining the category are the following: Category E(HOMO) − E(LUMO) HOMO and LUMO overlap Polarity of A and B

Charge control Large Poor High

Orbital control Small Good Weak

Recent examples of the application of Klopman’s approach to molecular complexes are the partitioning of the binding energies of SO3 complexes with nitrogen bases [34] and of halogen-bonded complexes with diversified bases [35]. The results show that, for the Cl I· · ·B complexes, the electrostatic (charge-controlled) bonding is predominant relative

12


to covalent (orbital-controlled) bonding, whereas for B· · ·SO3 complexes the electrostatic contribution to the bonding is comparable to the covalent contribution. Despite evident shortcomings, Klopman’s theory was popular in acid/base chemistry books and reviews in the 1970s because it reproduces the qualitative features of the concept of hard and soft Lewis acids and bases (see below). More satisfactory treatments of the perturbative calculation and partition of intermolecular energies are employed nowadays. An excellent summary of these approaches is available in the textbook by Stone [36]. The contributions collectively known as symmetry-adapted perturbation theory (SAPT) have been reviewed by Szalewicz and Jeziorski [37]. An intermolecular perturbation theory (IMPT) for the region of moderate overlap [38] is of particular interest in Lewis acid/base chemistry since it emphasizes the computation of charge-transfer energies [39] and is implemented in a freely available quantum chemistry package [40]. As an example [41], the electronic interaction energy components of the hydrogen-bonded complex MeOH· · ·NEt3 , calculated at the IMPT/6–31G(d)//B3LYP/6–31+G(d,p) level, are (in kJ mol−1 ) −53.6 (electrostatic), +54.5 (exchange-repulsion), –6.4 (polarization), −7.6 (charge transfer), −19.6 (dispersion) and −32.7 (total). The perturbation calculations of intermolecular energies are tedious, even for small molecules. On the other hand, the perturbation approach is advantageous (over the variational supermolecule approach, see below) because (i) the small interaction energy is calculated directly rather than as a difference between two large, almost identical, numbers (EAB and EA + EB ), (ii) the individual contributions to the intermolecular energy have a clear physical meaning, (iii) SAPT terms are free of the basis set superposition error (BSSE, see below) and (iv) each term can be evaluated using a different basis set, the most appropriate and economic for that particular component. 1.5.3

Variational Supermolecular Method and Energy Decomposition Schemes

In this method, the system of the two interacting molecules, A and B, is treated as a supermolecule and their interaction energy, ∆E, is the energy of the supermolecule minus the energies of the isolated molecules: ∆E = E AB − (E A + E B )

(1.40)

The energy of the supermolecule AB is calculated by solving approximately the Schrödinger equation: H ΨAB = E AB ΨAB

(1.41)

by the variational method, in the same way as for the isolated molecules A and B. These calculations can be routinely performed using quantum chemistry program packages such as CAPDAC, GAMESS, GAUSSIAN, JAGUAR or SPARTAN, to cite the most popular in the chemistry community. However, they are fraught with difficulties [36, 42–45], as discussed below. Method The applicability of the self-consistent field (SCF) or Hartree–Fock (HF) method is evidently limited because electron correlation, and consequently the dispersion interaction


13

which is wholly a correlation effect, are absent from any SCF calculation. A popular method in recent years has been the use of density functional theory (DFT) based on the theorem ‘the energy is a functional of the electron density’. The literature contains many proposed functionals (for example, B3LYP). However, no functional for the correlation energy that yields values within so-called chemical accuracy (1 kcal mol−1 = 4.184 kJ mol−1 ) seems yet to exist. Hence density functional approaches must be used cautiously to study molecular interactions where the contribution of dispersion energy is significant. A satisfactory approach to the correlation problem is the Møller–Plesset perturbation theory (MPn for perturbation theory to order n). MP2 calculations are significantly more expensive than SCF calculations but the MP2 method based on the ‘resolution of identity’ (RI) is about one order of magnitude faster than the exact MP2 method (with almost identical interaction energies for both methods). Therefore, the RI-MP2 method (implemented in the TURBOMOLE package) is well suited for the study of large complexes [46]. Atomic Orbital Basis Sets These are the atomic orbitals used in the LCAO-MO process (φ in Equation 1.37). When choosing an atomic orbital basis set among the basis set libraries available in the quantum chemistry packages, one must specify (i) the size and nature of the primary core and valence basis, (ii) the so-called polarization functions added to the basis and (iii) the socalled diffuse functions that still augment the basis set. For example, 6–311++G∗∗ is a split-valence (represented by 6–311G) basis set plus polarization (represented by ∗∗ ) and diffuse functions (represented by ++). Large basis sets require a long computer time (roughly proportional to the fourth power of the number of basis functions), but the energy comes closer to the HF limit and the BSSE is less serious. BSSE arises from the fact that the supermolecule is described by a set that is formed by superposition of the basis sets of the two molecules A and B, that is, by a larger basis set than those of molecules A and B. A larger basis set of the supermolecule inevitably yields a larger EAB value and, consequently, a larger ∆E value. This artificial increase in EAB must be corrected by the counterpoise method [47, 48], in which the energies of the molecules A and B are calculated in the basis set of the supermolecule. The BSSE follows from the equation BSSE = E cβ (B) − E cα∪β (B) + E cα (A) − E cα∪β (A)

(1.42)

where α(β) means the A (B) basis set, α∪β is the basis set of the supermolecule and the subscript c denotes that the energy of the molecule A (B) has been calculated using its geometry within the complex (supermolecule). Geometry Optimization Calculating the geometry of a supermolecule with N nuclei is an extraordinarily complex problem, since it is necessary to explore a correlated free energy hypersurface (FES) with 3N – 6 degrees of freedom (paying special attention to the up to six coordinates describing the relative configuration of molecules A and B) by a counterpoise-corrected gradient optimization procedure. The aim is not only to locate the global minimum but also to identify other significant free energy minima for subsequent comparison of theoretical basicity with experimental results. When N is large and the FES contains many saddle

14


points and energy minima separated by low-energy barriers, this issue requires the use of methods of molecular dynamics and statistical mechanical simulations. In order to save computing time, several approximations are often used in the geometry optimization procedure: (i) It is assumed that entropy does not play an important role and the potential energy surface (PES) is considered instead of the FES. (ii) It is common to optimize the geometry at a DFT level and then to apply a higher (better correlated) level to compute the energy. Such calculations are indicated by a double slash (for example, MP2/6–311++G∗∗ //B3LYP/6–31+G∗∗ ). (iii) The BSSE is not a priori included in the geometry optimization cycles but a a posteriori in the correction of the interaction energy. (iv) Libraries of molecular and supermolecular structures (such as the MOGADOC database [49] for the gas phase and the CSD [50] for the solid state) are searched to find starting geometries. Then the gradient optimization procedure localizes the nearest energy minimum and stops the calculation. Since there is no guarantee that the minimum so located is the global minimum, it is necessary to restart the optimization from a different geometry. After several trials, it is hoped that the main portions of PES have been properly sampled. For example, a CSD search furnishes three conformations for the hydrogen-bonded complexes of carboxylate anions: syn 1, anti 2 and bifurcated 3. Three starting geometries for the hydrogen-bonded complex of methanol with the formate anion might thus be

C O

H

H

H

C O

O

C H

O

H

O

O CH3

O H O

O

CH3

H3C

1

2

3

A B3LYP/6–31+G∗∗ optimization indicates (in vacuo) that 1 is the deepest minimum, 2 is a local minimum and 3 is a saddle point [41]. Anharmonicity After an optimized structure has been obtained, the vibrational frequencies are calculated to characterize the structure as a local minimum (with no imaginary frequencies) and provide the data needed to evaluate (i) the zero-point energy, (ii) the thermal vibrational contribution to enthalpy and (iii) the vibrational entropy. Ab initio harmonic vibrational frequencies are typically larger than the values measured spectroscopically. A major source of this disagreement is the neglect of anharmonicity effects in the calculations (errors also arise because of the incomplete incorporation of electron correlation and the use of finite basis sets). It is customary to scale ab initio values for the imperfect calculations. A set of scale factors has been recommended [51]. A better approach is to compute the anharmonic


15

corrections, but the methods recently implemented within quantum chemistry program packages are demanding of computer resources. Relativity The properties of molecules and supermolecules containing light atoms (Z < 36) are described accurately within the non-relativistic approximation. The relativistic character of inner shell electrons in heavier elements is successfully hidden inside effective core potentials (ECPs). The ECP basis set LANL2DZ∗ is a set with double-zeta functions for each of the two outermost shells augmented by polarization functions, which was shown to yield reasonably good results for the complexes of diiodine (see Chapter 5). Comparison with Experimental Data The comparison of calculated interaction energies with the experimental enthalpies is essential since it identifies the method and basis set that yield reliable energies at minimal computational expense, and assesses the validity of approximations. The interaction (binding) enthalpy at 298 K and 105 Pa and the interaction energy at 0 K are related as follows: ◦ ∆H298 = ∆E el + ∆ZPVE + ∆E vib,therm + ∆E rot + ∆E trans + ∆n RT

(1.43)

The operator ∆ corresponds to the change upon the reaction A + B → AB. The term ∆Eel is the BSSE-corrected electronic interaction energy. The second term is the zeropoint vibrational energy contribution. The vibrational term, ∆Evib,therm , is the change in the vibrational energies in going from 0 to 298 K. The last terms are thermal terms, which account for the loss of rotational and translational degrees of freedom and the change in the number of moles of perfect gas (∆n = −1 for a 1 : 1 complexation). ∆Erot and ∆Etrans are classically equal to ± 12 RT for each degree of freedom gained or lost in the reaction. The dissociation energy of the A B bond is given by D0 = −∆E el − ∆ZPVE

(1.44)

and the Gibbs energy of complexation, ∆G ◦298 , is ◦ ◦ ∆G ◦298 = ∆H298 − T ∆S298

(1.45)

◦ where the entropy of complexation, ∆S298 , is the sum of the translational, rotational and vibrational entropy changes: ◦ = ∆Str + ∆Srot + ∆Svib ∆S298

(1.46)

Taking the hydrogen-bonded complex of methanol with methylamine, CH3 OH· · ·NH2 CH3 , as an example, the various contributions to ∆H ◦ and ∆S◦ , calculated at the MP2/aug-cc-pVTZ//B3LYP/6–31+G∗∗ level, are listed in Table 1.7. There is a very good agreement between the calculated [41] and experimental [52] enthalpies since the difference lies within the experimental error. The agreement is far less good for entropies. The largest source of error arises in the vibrational term, calculated in the harmonic approximation and sensitive to the level of theory (frequencies calculated at the B3LYP/6–31+G∗∗ level).

16

Lewis Basicity and Affinity Scales Table 1.7 Thermodynamic parameters of the reaction CH3 OH + CH3 NH2 → CH3 OH· · ·NH2 CH3 at 298 K and 105 Pa. Energies in kJ mol−1 , entropies in J K−1 mol−1 . ∆Eel BSSE ∆ZPVE Do = ∆Evib,therm ∆Etrans ∆Erot −RT ◦ ∆H298 = exp.

−36.68 6.09 6.58 24.0 10.38 −3.72 −3.72 −2.48 −23.54 −23.4 ± 1.0

∆Svib ∆Strans ∆Srot

92.57 −143.25 −52.81

◦ ∆S298 = exp.

−103.5 −73 ± 4

Energy Decomposition Schemes The supermolecule method is capable of assessing the binding strength of species A and B through the dissociation energy, D0 . It shows that there is a continuum of binding energies from very weak van der Waals bonds (for example, between two rare gas atoms in a molecular beam) to strong dative bonds. However, in contrast to perturbation methods, it cannot determine the origin of the binding forces that hold the species A and B together. Physical intuition suggests that there are five different types of electric interactions between the species A and B as they approach one another [53]. At a large separation between the two species, the most important interaction arises from the direct electrostatic (ES) (coulombic) interaction between the unperturbed charge distributions of A and B. It is customary to express the charge distribution in terms of a multipole expansion: charge, dipole, quadrupole and so on. Then the direct electrostatic interaction between A and B is the sum of the interactions of their respective multipole moments. When the separation decreases, an attractive term related to the mutual polarization (POL) of the two charge distributions will become significant. It arises from the deformation of the electron cloud of A induced by the electrical field produced by B (and vice versa). Within the same range of mutual distance, an attractive interaction originates from an intermolecular correlation between the instantaneous fluctuations in the charge distributions of A and B. London called this inherently quantum mechanical phenomenon dispersion (DISP). When the species A and B approach so closely that their electron clouds overlap, the Pauli exclusion principle keeps electrons of the same spin away from each other. Therefore, electron density is removed from the overlap region. The positively charged nuclei are thus incompletely shielded from each other and mutually repel. This interaction is described as exchange-repulsion or just exchange (EX). Another short-range interaction has been described by Mulliken in interpreting the bonding that can occur between a species that has a low ionization energy (an electron donor) and one that has a high electron affinity (an electron acceptor). When the electron clouds of the approaching species overlap, a portion of the electrons may shift from one of the two species to the other (essentially from B to A). This attractive charge transfer (CT) interaction should occur for a complex to be designated as a Lewis acid/base complex. The


17

question of how much charge is transferred in adducts is, however, very complex and will not be addressed here. According to this description of the interaction effects, the whole interaction energy ∆E could be divided into separate and additive terms: ∆E = ∆E ES + ∆E POL + ∆E DISP + ∆E EX + ∆E CT

(1.47)

There is no unequivocal operative definition of these terms allowing numerical computations. Consequently, many energy partition schemes have been proposed. Most are based on ideas presented first in 1971 by Morokuma [54, 55], who decomposed the variational Hartree–Fock interaction energy into ∆EES , ∆EPOL , ∆EEX and ∆ECT (note that this list does not include ∆EDISP , which must be calculated separately). DFT-calculated interaction energies have also been partitioned (see a review [56] on various classes of donor/acceptor complexes of transition metals and main group elements). 1.5.4

Natural Bond Orbital Theory

The natural bond orbital (NBO) method of Weinhold et al. [8, 57] provides a scheme appropriate to the analysis of Lewis acid/base interactions [8, 58] as it emphasizes the calculation of delocalization of electron density into unoccupied orbitals. The atomic orbital basis set is transformed into two sets of orbitals. The first set, consisting of core, lone pair and bond orbitals, is mathematically chosen to include the highest possible percentage of the electron density, and thus provides the most accurate possible ‘natural Lewis structure’ picture of the wavefunction. Since each valence bonding NBO must be paired with a corresponding valence antibonding NBO, the Lewis-type (donor) NBOs are complemented by a set of non-Lewis-type (acceptor) NBOs that are formally empty in an idealized Lewis structure picture. Weak occupancies of the valence antibonds correspond to irreducible departures from the idealized Lewis picture. These occupancies are a direct consequence of delocalizing interactions with the occupied donor orbitals. The energetic stabilization due to such donor/acceptor interactions can be estimated by second-order perturbation theory. An interesting example is provided by the NBO analysis of the water dimer H2 O· · ·HOH, where the left and right molecules behave as the Lewis base and the Lewis acid, respectively. The interaction energy is decomposed into charge transfer (CT) and no charge transfer (NCT) as follows: ∆E = ∆E NCT + ∆E CT

(1.48)

∆ECT is defined as the loss in interaction energy that results when the energy is calculated after deleting the Fock matrix elements between the occupied orbitals of one molecule and the unfilled orbitals of the other. A large basis set HF calculation gives values of −17.2, +10.0 and −27.2 kJ mol−1 for ∆E, ∆ENCT and ∆ECT , respectively. The main origin of the charge transfer component ∆ECT is identified by a second-order perturbative analysis of the Fock matrix as an n → σ ∗ charge transfer from one of the oxygen lone pairs, n, of the electron donor molecule (hydrogen-bond acceptor) to the proximate OH antibond, σ ∗ , of the electron acceptor molecule (hydrogen-bond donor), yielding a stabilization energy of −27.6 kJ mol−1 , in good agreement with the value of ∆ECT . Accordingly, the NBO occupancies indicate that the proximate σ ∗ orbital has increased in occupancy by 0.0083 e

18


from its monomer value of 0.0000 and that the proximate lone pair has decreased in occupancy by 0.0081 e from its monomer value of 1.9988. A correlated wavefunction confirms the n → σ ∗ picture of the water dimer and shows that dispersion is not important in this complex (∼20% of the total interaction energy). The values of the CT term calculated by the NBO scheme are very much larger than those found by most other methods. It should not be concluded that the NBO energy decomposition analysis is wrong in principle. The divergence should rather be attributed to a different operational definition of the charge transfer energy. A unified NBO donor/acceptor perspective is given to the complexes of various Lewis acids (BF3 , transition metals and ions, hydrogen-bond donors, Br2 , NO+ , tricarbonylchromium, tetracyanoethylene, etc.) in a book by Weinhold and Landis [8]. This analysis emphasizes the importance of orbital interactions in the formation and properties of complexes. A nice example is the description of rare gases as unusual Lewis bases towards BeO because suitable matching orbitals can be found on this strong Lewis acid. In the linear complex He· · ·BeO, the filled nHe orbital interacts strongly with the backside of ∗ antibond and the vacant n∗Be orbital (sp hybrid). The calculated interaction energy the σBeO amounts to −26.8 kJ mol−1 . 1.5.5

Quantum Theory of Atoms in Molecules

The theory of ‘atoms in molecules’ (AIM) of Bader [59, 60] offers criteria enabling the existence and the nature of bonds to be determined from the topological analysis of the electron density. This theory has not only provided new insights into the understanding of intramolecular bonds but has also been successful in the field of intermolecular bonds, such as hydrogen bonds [61], halogen bonds [62], van der Waals bonds [63] and more strongly bound donor/acceptor complexes [64]. The definitions sufficient to understand the topology of a molecule, or in our case a complex, are illustrated by Figure 1.1 and given in the following:

A

B

Figure 1.1 Electron density contour of a diatomic molecule AB overlaid with trajectories of ∇ρ. The bond path is defined by two trajectories starting at the bond critical point (denoted by a square) and ending at the nuclei. The atomic basins Ω(A) and Ω(B) are separated by a pair of trajectories originating at infinity and terminating at the bond critical point.


19

r Electron density function ρ(r): a three-dimensional function defined such that ρ(r)dτ is the probability of finding an electron in the elemental volume dτ at point r. This function may be determined experimentally by analysing X-ray diffraction data of solids. r Gradient of the electron density ∇ρ(r): a vector that points in the direction of maximum increase in the density. r Gradient vector field of the electron density: obtained by the trajectories traced out by the gradient vectors. r Critical point: special point where ∇ρ(r) vanishes (for instance, a nucleus). r Nucleus: attractor of the gradient vector field of ρ(r) where trajectories starting at infinity terminate. r Basin: space containing all trajectories leading to an attractor. r Atom: union of a nucleus and its basin. r Bond critical point: critical point found between nuclei that appear linked by a saddle in ρ(r). There is a unique pair of trajectories that originate at such a critical point and terminate, each one, at the neighbouring nuclei. They define the bond path along which the electron density is a maximum with respect to any lateral displacement. r Laplacian of the electron density, ∇ 2 ρ(r): scalar derivative of the gradient vector field of ρ(r). It determines where electronic charge is locally concentrated, ∇ 2 ρ < 0, and depleted, ∇ 2 ρ > 0. The Laplacian can be partitioned into energy densities. It can be demonstrated that there is a local virial theorem 1 2 (1.49) ∇ ρ(r ) = 2G(r ) + V (r ) 4 (in atomic units), where G(r) is the kinetic energy density and V(r) the potential energy density. The local energy density H(r) is the sum of the kinetic and the potential energy density. r Important properties that can be evaluated at the bond critical point (b): ρ b , ∇ 2 ρ b , Gb , V b and H b . r Atomic properties requiring an integration over the atomic basin (Ω): E(Ω) (energy), N(Ω) (electron population) and v(Ω) (volume). r Bond order (BO) can be calculated according to Angyan et al. [65] The topological AIM properties have been extensively used for rationalizing Lewis acid/base reactions. A few examples follow. The Lewis basic site of a species is identified as a critical point where −∇ 2 ρ is maximum (local charge concentration) with high potential energy, while the Lewis acidic site is identified as a critical point where −∇ 2 ρ is minimum (local charge depletion) with high kinetic energy. A Lewis acid/base reaction corresponds to the combination of a charge concentration (a ‘lump’ in the zero value surface of ∇ 2 ρ) in the valence shell of the base with a charge depletion (a ‘hole’ in the zero value surface of ∇ 2 ρ) in the valence shell of the acid. The alignment of the two critical points (i.e. the complementarity of the ‘lump’ and ‘hole’) provides a guide to the relative angle of approach of the acid and base molecules. A set of criteria have been proposed to establish weak intermolecular bonding interactions, such as CH4 · · ·OH2 (hydrogen bond?), F2 · · ·CO (halogen bond?) or Ar· · ·CO2 (van der Waals bond). For example, the criteria necessary to conclude that a hydrogen bond X H· · ·B is present are as follows [61]: (i) the existence of a bond critical point and a bond path H· · ·B, (ii) an appropriate value of ρ b (the typical range is between 0.002 and 0.034 au),

20


(iii) an appropriate value of ∇ 2 ρ b (a range of +0.05 au for the π -complex FH· · ·HC CH to +0.14 au for the O complex FH· · ·OH2 can be found in ref. [62]), (iv) mutual penetration of hydrogen and acceptor atom, (v) loss of charge of the hydrogen atom, (vi) energetic destabilization of the hydrogen atom and (vii) decrease in the hydrogen atom’s volume. Hydrogen bonds (like ionic bonds, van der Waals bonds and halogen bonds) are considered as closed-shell interactions because they show ∇ 2 ρ b > 0. Hence they differ from covalent bonds, metallic bonds and dative bonds considered as shared-shell interactions with ∇ 2 ρ b < 0. However, strong hydrogen bonds are also characterized by ∇ 2 ρ b < 0. Thus, hydrogen bonding can cover the spectrum of bonding interactions, from closed- to shared-shell. Numerous correlations have been established between ρ b , ∇ 2 ρ b , Gb , V b , BO and so on and the electronic interaction energy. However, good correlations are generally restricted to structures in which the same pair of atoms are interacting. Multivariate correlations are required to obtain good family-independent correlations [66]. The carbonyl complexes of Cr, Fe and Ni and the metallocene complexes of Fe, Al+ and Ge have been studied [64] to illustrate the complementary nature of the quantum theory of atoms in molecules and molecular orbital theory in the understanding of the metal–carbon bonds. An alternative partition of the molecular space is provided by topological analysis of the electron localization function [67], which yields basins related to the local pairing of electrons. By this method, Silvi et al., investigated the 1 : 1 complexes of the first series transition metals and various ligands [68], and hydrogen bonding [69]. They also studied the protonation site of simple bases and established the following rule [70]: protonation occurs in the most populated basin of the base that yields the least topological change of the localization gradient field.

1.6

Measurement of Lewis Basicity

According to the glossary of terms used in physical organic chemistry published by the International Union of Pure and Applied Chemistry [71], Lewis basicity is defined as follows: The thermodynamic tendency of a substance to act as a Lewis base. Comparative measures of this property are provided by the equilibrium constants for Lewis adduct formation for a series of Lewis bases with a common reference Lewis acid.

Thus, the Lewis basicity of a series of bases Bi will be measured by the thermodynamic equilibrium constants K of the equilibria Bi + A◦ Bi A◦

(1.50)

where A◦ is a reference Lewis acid. These equilibrium constants are referred to as complex formation, complexation, association, stability or binding constants. The reciprocal quantity is a dissociation or instability constant: Bi A◦ Bi + A◦

(1.51)

A Lewis basicity scale can be constructed from this set of equilibrium constants. The basicity scale may be expressed as K, or as pK or as ∆G. In order that the definition and


21

notation of Lewis scales be the same as those of the aqueous Brönsted scale pK BH+ , the Lewis basicity towards the acid A◦ is, at this point, defined as pK BA◦ = −log10 K BA◦ = −logK(reaction 1.51) = + logK(reaction 1.50)

(1.52)

In this way, a Lewis base will be strong if it forms a complex with a large association constant, that is, a low dissociation constant, and a large positive value of pK BA0 will describe a strong Lewis base towards A◦ . 1.6.1

Gas-phase Reactions

For the ideal-gas (a good approximation for real gas mixtures at relatively low pressures) reaction A(g) + B(g) AB(g), the standard pressure equilibrium constant K p◦ is K p◦ =

(PAB /P ◦ ) (PB /P ◦ ) (PA /P ◦ )

(1.53)

where PAB , PA and PB are the partial pressures at equilibrium in bar of the complex, the acid and the base, respectively, and the standard pressure P◦ is taken to be 1 bar. The total pressure of the system at equilibrium, P, is given by the sum of the partial pressures: P = PA + PB + PAB

(1.54)

From the stoichiometry of the reaction, PA and PB can be related to the initial pressures P0,A and P0,B . In the manometric method, the pressures P0,A , P0,B and P are successively measured in a vacuum line kept at a fixed temperature using a precision manometer. Thus, the equilibrium constant can be calculated using the equation P0,A + P0,B − P ◦ P◦ (1.55) Kp = P − P0,B P − P0,A Examples of determinations of pressure equilibrium constants by the manometric method can be found in the literature [52, 72] for the complexes of B(CH3 )3 and CH3 OH with amines. The equilibrium constant can also be expressed in terms of concentration C by using the ideal-gas relation: n (1.56) P = RT = CRT V where R is the ideal-gas constant (0.083 145 l bar mol−1 K−1 ), n the number of moles, T the temperature (K) and V the volume (l). We have ◦ −1 CAB RT −1 ◦ C RT = K (1.57) K p◦ = c CA CB P ◦ P◦ where K c◦ is the standard concentration equilibrium constant CAB C ◦ ◦ Kc = ◦ ◦ CA C CB C and C◦ is taken to be 1 mol l−1 .

(1.58)

22


Since the standard state of an ideal gas is defined as having 1 bar pressure, the standard complexation Gibbs energy ∆G◦ is directly related to K p◦ : ∆G ◦ = −RT lnK p◦ but indirectly related to 1.6.2

K c◦

(1.59)

through (1.57).

Solution Reactions

Most complexation reactions between neutral molecules are carried out in a solvent since it is easier to work on solutions than on gases. Moreover, the number of neutral systems that can be studied by gas-phase techniques is limited. The ion–molecule reactions in the gas phase are treated in Chapter 6. For the reaction A(solution) + B(solution) AB(solution), the activity equilibrium constant K a is defined as aAB (1.60) Ka = aA aB where aAB , aA and aB are the activities at equilibrium of the complex, the acid and the base, respectively. Practical expressions for the activities depend on the scale chosen to specify the composition of the solution. If the mole-fraction scale is used for the solutes, then the mole-fraction scale activity is ax = γx x

(1.61)

where γ x is the mole-fraction scale activity coefficient and x the mole fraction. The equilibrium constant K a becomes xAB γx,AB (1.62) Kx = xA xB γx,A γx,B and the standard Gibbs energy of complexation is ∆G ◦x = −RT lnK x

(1.63)

The subscript x on K and ∆G◦ indicates that the mole-fraction scale is used. The concentration scale is most commonly used for determining equilibrium constants of Lewis acid/base reactions in organic solvents. The concentration scale activity of a solute is ac = γc C/C ◦

(1.64)

where γ c is the concentration-scale activity coefficient and C◦ = 1 l mol−1 . The equations for K c◦ and ∆G◦ are K c◦ =

CAB /C ◦ γc,AB (CA /C ◦ ) (CB /C ◦ ) γc,A γc,B

(1.65)

∆G ◦c = −RT lnK c◦

(1.66)

In non-electrolyte solutions, if the solution is dilute, it is a good approximation to set the activity coefficients equal to one. The equilibrium constants Kx in (1.62) and K c◦ in (1.65)


23

then reduce to the expressions (1.67) and (1.68) for ideally dilute solutions: xAB xA xB

(1.67)

(CAB /C ◦ ) (CA /C ◦ )(CB /C ◦ )

(1.68)

Kx ≈ K c◦ ≈

It is often convenient to work with an equilibrium constant that omits the C◦ in (1.68). We define the concentration equilibrium constant K c as Kc =

CAB CA CB

(1.69)

K c has the dimensions of l mol−1 . The validity of the approximations (1.67) and (1.68) decreases as the complexation constant is small, owing to the high base and/or acid concentration that should be used to shift the equilibrium towards the complex formation. Unfortunately, it is very difficult, and for practical purposes impossible, to measure all the individual activity coefficients γ AB , γ A and γ B . Thus, it is hoped that the ratio γ AB /γ A γ B remains close to unity, or at least constant. For example, the complexation reaction CH3 CON(CH3 )2 + I2 (CH3 )2 NC(CH3 )=O · · · I2

(1.70)

of complexation constant K c = 6.8 l mol−1 has been studied [73] in CCl4 over the concentration range 0.002–0.40 mol l−1 of base for solutions of 10−3 mol l−1 of I2 . The activity coefficient of CH3 CON(CH3 )2 in this concentration range varies from 1.0 to 0.77. Since I2 is very dilute and non-polar, γI2 = 1. One cannot measure γ ab . Consequently, since the values of K c do not show a concentration dependence on going from 0.002 to 0.40 mol l−1 , it may be assumed that γAB /γ B is approximately one, or at least constant. 1.6.3

Standard State Transformations

The relationship between the mole fraction xi and the concentration Ci (mol l−1 ) of a solute i (A, B and AB) in a solvent S is given by n j Mj Ci xi = (1.71) 1000 d nj or, in terms of molar volumes V (l mol−1 ), xi = Ci

n j Vj nj

(1.72)

where d is the density of the solution in g cm–3 , n j is the sum of the number of moles (j = A, B, AB and S) and Mj (g mol–1 ) is the molar mass of constituent j. In very dilute solutions of solute i, Equation 1.71 becomes xi ≈

C i Ms 1000 ds

(1.73)

24


where ds and M s are the density and the molar mass of the solvent, respectively. For constituent i in a solution containing high concentrations of solute B, Equation 1.72 becomes

(1.74) xi ≈ Ci Vs + xB VB − Vs Thus, K c and K x are related by expressions

K c ≈ K x Vs + xB VB − Vs

(1.75)

when xB is appreciable, or K c ≈ K x Vs = K x

Ms 1000 ds

(1.76)

when all solutes are very dilute. Consequently, the standard Gibbs energies ∆G ◦x and ∆G ◦c are related by (1.77) ∆G ◦c ≈ ∆G ◦x − RT ln Vs /1 mol l−1 The conversion factor Vs relating Kx to K c , and the correction term −RT lnVs relating ∆G ◦x to ∆G ◦c are given in Table 1.8. 1.6.4

Choice of Solvent

The criteria for choosing a solvent fall into three classes: (i) the solubility of A, B and AB, (ii) those specific to the experimental technique, for example IR or UV transparency when using IR or UV spectrometric methods, and (iii) the avoidance of specific solvation effects. In fact, if a solvent S is specifically associated with species A, B and AB as shown by the equilibrium ASa + BSb ABS p + qS

(1.78)

where q = a + b − p, the equilibrium constant is defined by q

[ABS p ] xS Kc = [ASa ][BSb ]

(1.79)

where brackets signify equilibrium concentrations and xs is the mole fraction of solvent S. In this description of the formation of the complex in the ‘active’ solvent S, the constant q K c appears to be in error by the factor xS . It therefore appears necessary to choose solvents as ‘inert’ as permitted by criterion (i). Very roughly, the ‘inertness’ of a solvent without significant hydrogen-bond acceptor site(s) is indicated by a low relative permittivity, a low refractive index and a low value of Reichardt’s parameter ET (30) (low hydrogen-bond acidity) [75]. These solvent parameters are listed in Table 1.8 for the most often used solvents in determinations of complexation constants.

1.7

Measurement of Lewis Affinity

A number of groups (Drago [76, 77], Gutmann [78, 79], Maria and Gal [80], Arnett et al. [81]) have proposed measuring Lewis (Brönsted) basicity from the negative enthalpies of the complexation (protonation) reactions (1.50). In order to follow the IUPAC definition of

b

a

0.679 46 0.773 89 1.614 32 1.584 36 1.316 78 1.479 70 1.2458

dS

References [74, 75]. α RT 2 = RT 2 d lnK c /dT − RT 2 d lnK x /dT .

Heptane Cyclohexane Tetrachloroethylene Carbon tetrachloride Dichloromethane Chloroform 1,2-Dichloroethane

Solvent 100.203 84.161 165.834 153.823 84.933 119.378 98.960

MS 0.1475 0.1088 0.1027 0.0971 0.0645 0.0807 0.0794

V¯S 1.260 1.22 1.02 1.229 1.391 1.26 1.21

α 1.94 2.02 2.34 2.30 9.02 4.89 10.74

ε 1.388 1.426 1.505 1.460 1.424 1.446 1.445

n

30.9 30.8 31.8 32.4 40.7 39.1 41.3

ET (30)

4.745 5.500 5.641 5.781 6.795 6.240 6.279

−RTlnV¯S

0.933 0.902 0.754 0.908 1.028 0.931 0.894

αRT 2 b

Table 1.8 Properties of common solventsa (at 25◦ C unless otherwise stated): density dS (g cm–3 ), molar mass MS (g mol–1 ), molar volume V¯S (l mol–1 ), thermal expansivity α (10−3 K−1 ), relative permittivity ε (at 20◦ C), refractive index n (at 20◦ C), Reichardt’s ET (30) and correction terms RTlnV¯S and αRT 2 (kJ mol–1 ).

26


basicity [71], we recommend reserving the measurement of basicity to Gibbs energies of complexation (protonation) and using the name ‘affinity’ for the corresponding enthalpies. The names ‘enthalpy of basicity’ [26] and ‘enthalpimetric basicity’ [81] have been suggested. However, the term ‘affinity’ is in common use, mainly in the field of ion/base reactions, as exemplified by the proton affinity scale and the many scales of metal–cation affinities (see Chapter 6). Hereafter, the negative enthalpies of complexation of bases with the Lewis acids BF3 , 4-FC6 H4 OH or I2 will be named boron trifluoride, 4-fluorophenol and diiodine affinities, respectively. However, the historical name ‘donor number’ [78, 79] will also be used for the SbCl5 affinity. The two most commonly employed techniques for obtaining complexation enthalpies are based on the temperature dependence of equilibrium constants or calorimetric procedures. In the latter, the heat evolved when the acid and base are mixed in the reaction cell of a calorimeter is measured. The molar enthalpy of complexation, ∆H ◦ , is related to the measured heat output, Q, corrected for the heats of dilution, the equilibrium concentration of the complex, [AB], and the volume of the solution in litres, V, by the relation ∆H ◦ =

Q [AB] V

(1.80)

If an undetermined number of moles of AB have been formed, and if the equilibrium constant K c is known from a previous study (for example, by a spectrometric method), [AB] can be calculated by Equation 1.81 obtained by resolving Equation 1.82: 12 −1 1 −1 2 − 4[A0 ][B0 ] [A0 ] + [B0 ] + K c − [A0 ] + [B0 ] + K c (1.81) [AB] = 2 [AB] (1.82) Kc = ([A0 ] − [AB])([B0 ] − [AB]) where [A0 ] and [B0 ] represent the initial concentrations of acid and base. The value of ∆H ◦ can also be found without using a separately determined value of K c . ∆H ◦ and K c are considered as two adjustable parameters in a series of equations obtained by varying [B0 ] or [A0 ], and determined by iterative calculations. The high sensitivity of modern solution calorimeters offers great potential for the simultaneous determination of ∆H ◦ and K c , but attention must be paid to the stability of these optimized values in the iterative procedure. The alternative procedure consists in measuring the equilibrium constant at several temperatures. The temperature dependence of the pressure equilibrium constant for an ideal-gas reaction is given by the van’t Hoff Equation 1.83: (1.83) ∆H ◦ = RT 2 d lnK p◦ /dT Since d(T −1 ) = −T −2 dT, Equation 1.83 can be expressed in a form convenient for plotting: d ln K p◦ d (1/T )

=−

∆H ◦ R

(1.84)

Assuming that ∆H ◦ is constant over the temperature range involved, a plot of lnK p◦ against 1/T should be a straight line whose slope multiplied by −R gives the enthalpy ∆H ◦ . Having obtained ∆G◦ from Equation 1.59, we can calculate the value of ∆S◦ , the entropy of


27

complexation, from the equation ∆S ◦ =

(∆H ◦ − ∆G ◦ ) T

(1.85)

In the case of solution studies, the van’t Hoff Equation 1.83 can be applied to Kx values. We obtain the same ∆H ◦ value as that obtained from calorimetric measurements (Equation 1.80). However, application of Equation 1.83 to K c values leads to a different result: RT 2 (d ln K c /dT ) ≈ ∆H 0 + α RT 2

(1.86)

A demonstration of Equation 1.86 is given in ref. [82]. In the correction term αRT 2 , α is the coefficient of thermal expansion of the solvent: α = V −1 (dV /dT )

(1.87)

Table 1.8 gives values of the correction term at 298 K for common organic solvents. This term amounts to ∼1 kJ mol−1 and thus cannot be neglected for weak enthalpies. When applying the van’t Hoff equation, a temperature range as large as technically permitted must be used because the error δ in ∆H ◦ is inversely proportional to the temperature range T − T as shown [83] by Equation 1.88: δ ≈ 2R

T T ε T − T

(1.88)

where ε is the maximum relative error in the equilibrium constant (assuming that the slope of the van’t Hoff line is controlled by the first and last points). The error σ in ∆S◦ calculated by Equation 1.85 may be written as [83] T − T 1 + (1.89) σ ≈δ T 2T T The influence of the temperature range on the errors in ∆H ◦ and ∆S◦ is shown by the numerical examples in Table 1.9. Many correlations between the complexation enthalpy and the Gibbs energy of complexation have been proposed, for example the SbCl5 complexes [78], the diiodine complexes [4] or the hydrogen-bonded complexes [84]. They have been explained either by a quasiisoentropic behaviour of the complexation reaction or by an enthalpy–entropy compensation effect (Equation 1.90): ∆H ◦ = β∆S ◦ + constant

(1.90)

Table 1.9 Effect of the temperature (K) range on the errors δ in the enthalpy (kJ mol–1 ) and σ in the entropy (J K–1 mol–1 ). ε 5% 5% 5%

T

T

δ

σ

293 268 268

313 328 343

3.8 1.2 1.0

13.4 5.0 4.2

Remark Inadequate range Technically permitted range for CCl4 Technically permitted range for C2 Cl4

28


This linear relationship, in which the slope β has the dimension of temperature, gives ∆G ◦ =

β−T ∆H ◦ + constant β

(1.91)

There is a danger of misuse of these extrathermodynamic relationships for essentially two reasons. First, they are often obtained from incorrect statistical treatments. When the primary experimental quantities are the equilibrium constants, which are measured at different temperatures, any error that makes ∆H ◦ greater also makes ∆S◦ greater, as indicated by Equation 1.89. The propagation of errors will tend to distribute enthalpy and entropy estimates along a line characterized by a slope equal to the harmonic mean of the experimental temperatures [85]. This artefact has been pointed out many times [86–88]. Several correct statistical treatments have been advanced [85–89]. For example, the fair value of the correlation coefficient (0.951) of the enthalpy–entropy correlation (plotted in Figure 1.2) for the complexation of seven amines with I2 has been taken to imply a chemical causation [90]. A correct statistical treatment shows [85] that the 95% confidence interval for β is (850, 147) and includes the harmonic mean of the experimental temperatures, 298 K. Thus, the hypothesis that the observed enthalpy–entropy compensation is just a consequence of the propagation of experimental errors cannot be rejected at the 5% level of significance. The second reason is that the enthalpy–entropy compensation is generally limited to (i) acid complexes of specific base types (i.e. correlations are family dependent) and (ii) unhindered bases. An example is given for a set of complexes of boron acids with nitrogen and phosphorus bases [72, 91]. Table 1.10 shows the continuous increase in the quality of the ∆H–∆G correlation when the sample of 31 nitrogen and phosphorus bases is gradually restricted to 21 unhindered primary and secondary amines. The limited correlation is displayed in Figure 1.3. The molecular origin of the relations between ∆S and ∆H for a series of bases with a standard acid has been examined [92]. Investigation of the various terms contributing to ∆H in Equation 1.43 and to ∆S in Equation 1.46 suggests that the linear relations are

-∆H / kJ mol-1

55 45 35 25 15 30

50

70 -1

-∆S / J mol K

-1

Figure 1.2 Enthalpy–entropy plot for the complexation of diiodine with NH3 , MeNH2 , EtNH2 , BuNH2 , Me2 NH, Et2 NH and piperidine in heptane.


29

Table 1.10 Influence of the type of bases on the quality of the ∆H–∆G correlation for boron acid complexesa . Sample of bases

n

r

sb

Amines, pyridines and phosphines Amines and pyridines All amines Only unhindered primary and secondary amines

31 27 24 21

0.932 0.947 0.956 0.966

3.7 2.9 2.8 2.6

a The correct regression should be of ∆H◦ on ∆G◦ evaluated at the harmonic mean of experimental temperatures, T hm . Here, ∆G◦ is at 393 K, a temperature often close to the T hm s of the 31 experiments. b kJ mol−1 .

due to a linear relation between the vibrational entropy change, ∆Svib , and the dissociation energy, D0 . The latter relation can occur if the logarithms of A–B force constants in the complex are linearly related to D0 .

1.8

The Role of the Solvent

Ideally, the reaction A + B → AB should be carried out in the gas phase in order to measure an intrinsic Lewis basicity (affinity), free of solvent effects. The development of mass spectrometric techniques has enabled many equilibrium constants of cation exchange between bases to be measured and a number of Lewis basicity scales in the gas phase [93] to be constructed. However, there are only a few determinations of equilibrium constants for the formation of uncharged complexes in the gas phase. The construction of Lewis basicity scales towards uncharged Lewis acids has essentially been carried out in solution. It is well known that the equilibrium constant and complexation enthalpy are usually widely affected by intermolecular interactions of the species A, B and AB with the solvent

BMe3 affinity/kJ mol−1

100

75

50 −10

0

10 20 BMe3 basicity/kJ mol−1

30

Figure 1.3 Correlation of affinity and basicity for the complexes of BMe3 with primary and secondary amines in the gas phase.

30


Table 1.11 Complexation constants (l mol−1 ) and complexation enthalpies (kJ mol−1 ) of some Lewis acid/base complexes in the gas phase and in certain solvents. Gas phase

−∆H◦

Solvent

Ref.

Complexes more stable in the gas phase than in solution SO2 trans-2-Butene 3.1 0.08 TCNEb p-Xylene 280 33.9 7.2 I2 Benzene 4.5 8.4 0.24c I2 Diethyl ether 6.4 18.8 0.90c I2 Diethyl sulfide 226 35.1 195c CH3 OH Trimethylamine 20.8 28.9 4.6d CF3 CH2 OH Acetone 47.3 30.9 7.3

14.1 6.8c 17.6c 35.2c 23.7d 21.1

Hexane CH2 Cl2 Heptane Heptane Heptane C2 Cl4 CCl4

[95, 96] [97, 98] [99] [99, 100] [101] [52] [102]

Complexes more stable in solution than in the gas phase Trimethylamine 340 40.6 2550 SO2 BF3 Trimethylamine 111.3 BF3 Trimethylphosphine 79.1

46 129.5 87.4

Heptane CH2 Cl2 CH2 Cl2

e

Acid

Kc a

Base

−∆H◦

Solution Kc a

[103]

e

At 25◦ C. Tetracyanoethylene. c Chapter 5. d M. Lucon, University of Nantes, personal communication. e Chapter 3. a

b

[94]. Weak complexes generally have larger complexation constants and affinities in the gas phase than in solution, while the reverse is often observed for strong complexes [94]. Numerical examples are given in Table 1.11. The values of thermodynamic functions in the gas and solution phases can be compared in the following cycle o

o

o

o

Bg + Ag

ABg

∆Hg , ∆Gg

Bs + As

ABs

∆Hs , ∆Gs

(1.92)

The relation between the gas-phase complexation energies, ∆Hg◦ or∆G ◦g , and the solution complexation energies, ∆Hs◦ or ∆G ◦s , is (where Y = G or H) ◦ ◦ ◦ (AB) − ∆Yg→s (B) − ∆Yg→s (A) ∆Ys◦ = ∆Yg◦ + ∆Yg→s

(1.93)

◦ (i) denote the energies of transfer of individual species i (i = A, B or AB) where ∆Yg→s from the gas phase into the solvent phase. The situation |∆Hg◦ | > |∆Hs◦ | and K g > K s is found when ∆H ◦ (AB) < ∆H ◦ (A) + ∆H ◦ (B) (1.94) g→s g→s g→s

and ∆G 0

g→s (AB)

< ∆G 0

g→s (A)

+ ∆G 0g→s (B)

(1.95)


31

The reverse situation, |∆Hg◦ | < |∆Hs◦ | and K g < K s , is observed when inequalities converse to (1.94) and (1.95) apply. The latter solvent effect has been attributed [103] mainly to the existence of a strong dipole–(induced) dipole interaction between the very polar complex AB (for example µ = 4.95 D for Me3 N–SO2 and 6.02 D for Me3 N–BF3 ) [103, 104] and the polar (polarizable) solvent; solute dipole/solvent forces are not so stabilizing for the weakly polar monomers A and B (for example µ = 0.86 D for Me3 N, 1.6 D for SO2 and 0 D for BF3 ). The question now is: ‘In spite of significant solvent effects, are the solution Lewis basicity scales closely related to the intrinsic gas-phase Lewis basicity scales?’ This is an important question for computational chemists who need to identify the computational methods that yield reliable basicities. A relative comparison of gas-phase computed basicities with solution experimental basicities would avoid the difficult and approximate modelling of the solvent effect [105]. However, this comparison requires that experimental gas-phase and solution basicities (affinities) be strongly correlated. This correlation appears to exist for BF3 affinities and hydrogen-bond basicities. Equation 1.96 BF3 affinity(gas) = −15.8(±6.0) +1.02(±0.07) [BF3 affinity(CH2 Cl2 )] n = 8, r = 0.985, s = 4.8 kJ mol−1

(1.96)

and Figure 1.4 compare BF3 affinities measured in the gas phase and in the solvent CH2 Cl2 (data in Chapter 3). Equation 1.97 CF3 CH2 OH basicity(gas) = 0.82(±0.11) + 0.80(±0.07)[4-FC6 H4 OH basicity(CCl4 )] n = 11, r = 0.971, s = 0.19 kJ mol−1

(1.97)

BF3 affinity (gas) / kJ mol

-1

and Figure 1.5 compare the CF3 CH2 OH hydrogen-bond basicity measured in the gas phase [106] with the 4-fluorophenol basicity scale measured in CCl4 (data in Chapter 4). The latter comparison is not as direct as the former but is, however, correct insofar as alcohol hydrogen-bond basicities measured in CCl4 are strongly correlated with the 4-fluorophenol basicity in CCl4 (see Chapter 4). 140 120 100

1

80

2

4 3 6

60

5 7

40 20

8

0 30

50

70

90

110

130

BF3 affinity (CH2Cl2) / kJ mol-1

Figure 1.4 Comparison of BF3 affinities measured in the gas phase and in CH2 Cl2 . 1 Me3 N, 2 Me3 P, 3 tetrahydrofuran, 4 tetrahydropyran, 5 Me2 O, 6 MeCOOEt, 7 Et2 O, 8 tetrahydrothiophene.

Lewis Basicity and Affinity Scales CF3CH2OH basicity (gas)

32

4 1

3,5 3 2,5

4

2 1,5 1

7

8 11

10

0,5 -0,5

3 2

6 5

9

0,5

1,5

2,5

3,5

4-FC6H4OH basicity (CCl4)

Figure 1.5 Comparison of the trifluoroethanol basicity (as logKc ) in the gas phase with the 4-fluorophenol basicity (pKBHX ) in CCl4 . 1 Tetramethylguanidine, 2 Me3 N, 3 Et3 N, 4 pyridine, 5 NH3 , 6 tetrahydrofuran, 7 Me2 CO, 8 Et2 O, 9 MeOH, 10 C2 H5 COH, 11 CF3 CH2 OH. Because Me3 N and Et3 N behave well, the deviation of NH3 is attributed to experimental errors, rather than to solvation effects.

In the thermodynamic cycle (1.92), the Gibbs energies of transfer ∆G ◦g→s (i) can be calculated by theoretical methods. Two main classes of methods have been developed for modelling solvent effects. Molecular dynamics and Monte Carlo methods use a discrete representation of the solvent molecules whereas, in the second class of so-called SCRF (self-consistent reaction field) methods, the solvent is represented as a dielectric continuum surrounding the solute cavity. These methods are outside the scope of this book. They are described in reviews [105, 107, 108] and books [109, 110]. Examples of their application to Lewis acid/base complexes can be found in the following references: hydrogen-bonded complexes [111–113], BF3 and BH3 complexes [114] and diiodine complexes [115]. When constructing a basicity scale in solution, the measurements must be carried out in one standard solvent because complexation constants may depend strongly on the nature of the solvent. Representative variations of complexation constants are given in Table 1.12 for hydrogen-bonded complexes [116] and in Table 1.13 for diiodine complexes [117]. It is seen that K c [4-FC6 H4 OH–(Me2 N)3 PO] is 27 times larger in cyclohexane than in CH2 Cl2 , while K c (I2 –N-methylimidazole) is 14 times smaller in CHCl3 than in 1,2-dichlorobenzene. The Table 1.12 Logarithm (to base 10) of complexation constants (l mol−1 ) of 4-fluorophenol with bases at 25◦ C in the indicated solvents. Base 1,4-Dioxane Triethylamine Pyridine N,N-Dimethylformamide Dimethyl sulfoxide Hexamethylphosphoric triamide

c-C6 H12

CCl4

0.83 1.99 2.03 2.30 2.56 3.8

0.73 1.98 1.86 2.10 2.54 3.60

C6 H5 Cl 1,2-C6 H4 Cl2 0.55 1.84 1.60 1.74 2.20 3.06

0.45 1.93 1.63 1.70 2.18 3.06

ClCH2 CH2 Cl CH2 Cl2 0.09 1.70 1.29 1.27 1.65 2.92

0.14 1.67 1.26 1.18 1.44 2.37


33

Table 1.13 Logarithm (to base 10) of complexation constants (l mol−1 ) of diiodine with bases at 25◦ C in the indicated solvents. Base

n-C7 H16 CCl4 C6 H5 Cl 1,2-C6 H4 Cl2 ClCH2 CH2 Cl CH2 Cl2 CHCl3

Tetrahydrofuran N,NDimethylformamide Methyl dithiovalerate Triphenylphosphine oxide Triphenylphosphine sulfide N-Methylimidazole Triphenylphosphine selenide a

0.39 0.81

0.12 0.46

−0.07 0.44

−0.24 0.72

0.10 0.16

−0.26 −0.02

−0.44 −0.22

1.12

1.10

1.30

1.49

1.46

1.38

1.25

a

1.38

1.17

1.27

0.96

0.97

0.89

a

2.26

2.26

2.49

2.32

2.13

2.86

2.67 3.48

3.10 3.72

3.35 4.09

2.65 3.88

2.19 3.65

a

3.18 3.93

Not soluble enough.

values of complexation Gibbs energies in two solvents S1 and S2 can be compared in the thermodynamic cycle (1.98): o

BS1+ A S1

ABS1

∆GS1

BS2 + A S2

AB S2

∆GS2

o

(1.98)

Equation 1.99 shows that the difference ∆G ◦S2 − ∆G ◦S1 depends on the Gibbs energies of transfer: ∆G ◦S2 = ∆G ◦S1 + ∆G ◦S1 →S2 (AB) − ∆G ◦S1 →S2 (B) − ∆G ◦S1 →S2 (A)

(1.99)

∆G ◦S1 →S2 (i), of species i (i = A, B, AB) from the solvent S1 to the solvent S2 . These Gibbs energies of transfer can themselves be considered to have five components ∆G ◦S1 →S2 (i) = ∆G ◦S1 →S2 (i)(cavity)+∆G ◦S1 →S2 (i)(electrostatic)+∆G ◦S1 →S2 (i)(dispersion)+ ∆G ◦S1 →S2 (i)(repulsion) + ∆G ◦S1 →S2 (i)(hydrogen bonding) (1.100) The cavity term is the positive Gibbs energy required to form the solute cavity within the solvent. The other terms correspond to the intermolecular interactions already described between two molecules and generalized here to solute/solvent interactions. Unfortunately, it is not always technically possible to study a series of complexation reactions with bases spanning a wide range of properties in the same solvent. Hence the results of measurements of an extended basicity scale are often obtained in different solvents. Since there is little hope that the various terms of Equations 1.99 and 1.100 cancel so that ∆G ◦S1 ≈∆G ◦S2 , methods must be found to refer the data originally determined in a given solvent to the standard solvent. The transformation can be done by means of linear Gibbs energy relationships in the form of Equation 1.101: log K c (in a standard solvent S1 ) = a logK c (in a given solvent S2 ) + b

(1.101)

34


log K c (in CCl4)

3

Pyridines

Carbonyls

Amines

2

Sulfonyls

1

0 0

1

2

3

log K c (in CH2Cl2)

Figure 1.6 Family-dependent transformation of 4-fluorophenol basicities in CH2 Cl2 to 4fluorophenol basicities in the standard solvent CCl4 . Different lines are found for sulfonyls [118], carbonyls [119], pyridines [119] and amines [120].

These relationships have been shown to exist for hydrogen-bonded complexes [116] and for diiodine complexes (see Chapter 5), but their validity domain is limited to structurally related bases. For example, the plot of logK c (in CCl4 ) versus logK c (in CH2 Cl2 ) for the hydrogen-bonded complexes of 4-fluorophenol with bases does not give a single linear relationship (Equation 1.101) but a series of lines corresponding to various families of oxygen and nitrogen bases, as shown by Figure 1.6. The secondary values calculated by Equation 1.101 are of lower accuracy than the primary values measured in the standard solvent, since the standard deviation of the estimate (often about 0.05–0.10 logK unit) must be added to the experimental uncertainty (often about 0.05 logK unit for hydrogen bonding). On the other hand, it has been observed [80] that the relative BF3 affinities for 13 representative bases are almost insensitive to the change of measurement medium from dichloromethane to nitrobenzene (see Chapter 3). A few scales of Lewis affinity and some spectroscopic scales of Lewis basicity (see below) have been constructed by carrying out the reaction A + B AB on a dilute solution of the acid in pure, liquid base as solvent. This pure base method will be studied in Chapter 4. It gives ‘solvent basicity scales’ which are not strictly equivalent to ‘solute basicity scales’ measured on a dilute solution of the acid and the base in an ‘inert’ solvent.

1.9

Spectroscopic Scales of Basicity (Affinity)

Besides direct measurement of the thermodynamic quantities K and ∆H discussed above, spectroscopic estimates of basicity (affinity) have been proposed. Their main attraction is the ease with which they can be carried out. Moreover, many systems that do not possess the physical properties needed for a thermodynamic study can be characterized by spectroscopic parameters. Spectroscopic scales of basicity (affinity) are based on the change of a spectrochemical property (NMR, UV–Vis, IR, etc.) of the Lewis acid upon complexation


35

with a series of Lewis bases, and on the assumption that this change is dominated by the strength of the interaction. Factors other than those that control the intermolecular bond strength, such as coupling of vibrational modes or neighbour anisotropic effects, are assumed negligible or quasi-constant along the studied series of complexes. Other conditions required for such spectrochemical data to acquire the status of a spectroscopic scale of basicity (affinity) are as follows: (i) measurement in standard conditions of medium and temperature; (ii) reference to a numerous and diversified sample of bases; (iii) correlation to the Gibbs energy of complexation, for a spectroscopic scale of Lewis basicity or to the enthalpy of complexation, for a spectroscopic scale of Lewis affinity. NMR data [116, 121] for the limiting 19 F chemical shift of 4-fluorophenol upon hydrogen bonding with Lewis bases fulfil these conditions: (i) They have been measured in dilute solutions of 4-FC6 H4 OH and base in CCl4 at 25 ◦ C. (ii) They have been obtained for about 60 oxygen, nitrogen and sulfur bases. (iii) 96% of the variance of the 19 F chemical shift is explained by the thermodynamic basicity (as logK c = pK BHX ; see Chapter 4), as shown in Figure 1.7. Unfortunately, many spectroscopic scales are not as rigorously defined as the 19 F chemical shift scale and are flawed by a number of shortcomings. The first of these comes from measurements that have not been thermoregulated, notwithstanding that the system is temperature dependent. For example, in the domain of UV–Vis measurements, thermochromism goes together with solvatochromism [122]. A second weakness comes from the supposed equivalence of solvent scales with solute scales on the basis of a poor interpretation of statistics. For example, the change in the OH infrared stretching frequency

5

∆ / ppm

4

3

2

1

0 0

1

2

3

4

5

pKBHX Figure 1.7 Plot of the limiting 19 F chemical shift versus the logarithm of the complexation constant for hydrogen-bonded complexes of 4-FC6 H4 OH with 59 diversified bases (r = 0.978, s = 0.13 ppm).

36


of methanol, ∆ν(OH), upon complexation with a series of bases can be measured in the pure bases or in dilute solutions of the bases in some inert medium. The two scales are apparently satisfactorily correlated (r = 0.97, n = 72) [123]. However, their differences are not randomly distributed. They are significantly correlated (r = 0.94, n = 26) with a parameter measuring the polarity of the pure base [123]. Thus, infrared ∆ν(XH) spectroscopic scales measured in the pure bases are contaminated by the polarity of the pure base. Methods such as the solvatochromic comparison method [122], the pure base calorimetric method [124], the solvatovibrational comparison method [125] and the solvatomagnetic comparison method [126] have been developed for subtracting the non-specific effects of the base. However, the scales constructed in this way are still not strictly equivalent to solute scales. This important point can be illustrated [126] by the comparison between the 19 F enhanced chemical shift of 4-fluorophenol hydrogen-bonded to bases in CCl4 , ∆2 , and the same quantity measured in the media of pure bases by the solvatomagnetic comparison method, ∆1 . For 19 diverse bases, only 72% of the variance of ∆1 is explained by ∆2 . Moreover, the ∆2 – ∆1 differences are not randomly distributed. They are related (n = 19, r = 0.9) to the Onsager function, (ε − 1)/(2ε + 1), that is, to the reaction field of the medium. Clearly, the solvent NMR spectroscopic scale of hydrogen-bond basicity, ∆1 , is not equivalent to the solute one, ∆2 . Two other examples concerning the complexation enthalpy of 4-fluorophenol and the solvatochromism of 4-nitrophenol, given in Chapter 4, support this conclusion. The last shortcoming concerns the relationships between spectroscopic and thermodynamic scales of basicity (affinity). It appears that they often have a more limited validity domain than claimed by their advocates. As shown in Chapters 4 and 5, these relationships are often family dependent, that is, they are limited to bases that are structurally related. In the correlation of two basicity-dependent properties, it is often found that the organic bases separate into families that are distinguished by their basic sites as follows: (a) (b) (c) (d) (e)

aromatic carbon bases; nitriles and oxygen bases, which have a rather high dipole moment in common; sp2 nitrogen bases; sp3 nitrogen bases; single-bonded sulfur bases (generally distinct from double-bonded sulfur bases).

The family (b) contains the most often encountered Lewis bases in chemistry: nitriles, nitros, carbonyls, sulfonyls, sulfinyls, phosphoryls and ethers. Samples of bases heavily loaded with such highly dipolar bases may show an apparently general relationship between spectroscopic and thermodynamic scales. For example, the acetylacetonato-N,N,N ,N tetramethylethylenediaminocopper(II) perchlorate complex (Figure 1.8) has been proposed [127] as a solvatochromic indicator of Lewis basicity. Indeed, the wavenumber ν˜ of the dx z , d yz → dx 2 −y 2 transition [128] of this square-planar complex at about 550 nm is blue shifted when the basicity of the axial ligands increases, and a satisfactory relationship (n = 12, r = 0.966) exists between the donor number (see Chapter 2) of the ligands and ν˜ . However, 11 ligands out of 12 belong to the family (b). The relationship of Figure 1.8 must therefore be limited to the estimation of the donor number of ligands in this family. Any extrapolation to other families could be unreliable.


37

20 19

2

-3

10 ν / cm

-1

1

3

18

4 6

17

N L O

7

5

Cu

9 8

10

11

16

+

-

ClO4

N L O

12 15 0

50

100

DN / kJ mol

150

-1

Figure 1.8 Relationship between the wavenumber of the highest energy transition of Cu(tmen)(acac)L2 ClO4 and the donor number (DN) of ligands L. 1 CH3 NO2 , 2 C6 H5 NO2 , 3 propylene carbonate, 4 CH3 COCH3 , 5 CH3 COOC2 H5 , 6 CH3 CN, 7 c-(CH2 )4 O, 8 (MeO)3 PO, 9 HCONMe2 , 10 Me2 SO, 11 (Me2 N)3 PO, 12 C5 H5 N.

Other types of family dependence may be found between two basicity-dependent properties that are very dissimilar. For the complexes of SO3 with Lewis bases [129], the comparison of the frequency shift of the asymmetric stretching vibration of the SO bonds upon coordination, ∆ν as (SO3 ), with the proton affinity, PA, splits the sample of bases into families of first- and second-row bases (Figure 1.9). For the spectroscopic nucleophilicity, N, defined [130, 131] from the force constant of the hydrogen bond X H· · ·B, the family of sp nitrogen bases (N2 , C2 N2 , HCN, HCCCN, MeCN, t-BuCN) exhibits a behaviour different from that of oxygen bases (Figure 1.10).

140

∆ν as (SO3 ) / cm

-1

120

9

100

8

80

4

7

60 3

40 6

1

20

2

5

0 600

700

800

900

1000

-1

PA / kJ mol

Figure 1.9 Two regression lines are found in the comparison of ∆ν as (SO3 ) with the proton affinity, the line of first-row bases (1 H2 O, 2 Me2 O, 3 NH3 , 4 NMe3 ) and that of second-row bases (5 PF3 , 6 H2 S, 7 PH3 , 8 Me2 S, 9 PMe3 ).

38

Lewis Basicity and Affinity Scales 14 12 10 N

8 6 4 2 0 400

500

600

700

800

PA / kJ mol

900

1000

-1

Figure 1.10 Lack of a general relationship between the spectroscopic nucleophilicity N and the proton affinity PA (n = 18, r = 0.585). The line is a least-squares fit (r = 0.974) to the data of six sp nitrogen bases ( ).

1.10

Polybasic Compounds

Many Lewis bases bear more than one possible Lewis acid acceptor site. They are called polyfunctional bases, polybases or, in coordination chemistry, multidentate ligands. For instance (Scheme 1.1), (i) coordination of the ligand SCN− to cations may occur at the nitrogen or at the sulfur atom, (ii) nicotine may be hydrogen bonded on the amino nitrogen or on the pyridine nitrogen [132], (iii) there is competition between the two chalcogen atoms X and Y of the shown compounds to coordinate diiodine [133] and (iv) BF3 may form a dative bond with each nitrogen N1, N3 or N7 of adenine [134]. In order to measure the basicity of such polyfunctional compounds, it is necessary to determine the relative stability of the various possible isomeric complexes. Consider, as the simplest example, the interaction of a base B1 B2 which has two binding sites, B1 and B2 , with a Lewis acid A which has a single binding site. Two 1 : 1 complexes can be formed by the interaction of the Lewis acid at either of the sites B1 or B2 , giving the isomeric 1 : 1 complexes C1 (A–B1 B2 ) and C2 (B1 B2 –A) of equilibrium

Me

NH2

Y

S

C

N

N

Me

N

N

NH

HN

S

C

N

(i)

3

CH3

N

7

1

N

N

H

X

(ii)

(iii) X,Y = O,S,Se

Scheme 1.1 Examples of polyfunctional bases.

(iv)


39

constants K 1 and K 2 : B1 B2 + A C1

K1 =

[C1 ] [B1 B2 ][A]

(1.102)

B1 B2 + A C2

K2 =

[C2 ] [B1 B2 ][A]

(1.103)

A 1 : 2 (1 Base:2 Acids) complex may be formed by the addition of a second Lewis acid to either of the 1 : 1 complexes. This possibility can be ruled out if the concentration of base is chosen in large excess to the concentration of the acid. Most experimental methods are not able to determine the equilibrium concentrations [C1 ] and [C2 ] separately and furnish only the sum [C1 ] + [C2 ]. It follows that experimental methods yield only the total (or gross) equilibrium constant K t : Kt =

[C1 ] + [C2 ] [B1 B2 ][A]

(1.104)

Its relationship to the individual complexation constants K 1 and K 2 is evidently Kt = K1 + K2

(1.105)

This can be generalized to any number of isomeric 1 : 1 complexes: Kt =

n

Ki

(1.106)

i=1

Three cases are encountered for the determination of the various Ki s: all sites have the same basicity, one site is significantly more basic than the other(s) and more than one site contributes significantly to the total constant. If all sites have the same basicity, then K t = nKi and each individual constant is 1 (1.107) K t and logK i = logK t − log n n The coefficient 1/n is called a statistical factor: when the 1 : 1 complex is formed there are n possibilities of placing the Lewis acid, but in the dissociation there is only one possibility of losing it. Values for some nitrogen bases are given in Table 1.14. The statistical factor represents a correction which should be applied to the experimental constant in order to obtain a true constant for a single functional group. Since the statistical factor is independent of temperature, it is a correction to the entropy of complexation: Ki =

∆Si◦ = ∆St◦ − R ln n

(1.108)

The statistical factor can also be deduced from the symmetry numbers [135, 136] by the equation σe (C) σi (C) 1 = n σe (A) σi (A) σe (B1 B2 ) σi (B1 B2 )

(1.109)

(in the absence of a racemic mixture), where σ e is the external symmetry number (the number of different ways in which the molecule can be rotated around symmetry axes into a configuration indistinguishable from the original) and σ i the internal symmetry number

40

Lewis Basicity and Affinity Scales Table 1.14 Statistical factors for the complexation constant of N-heteroaromatics and amines. Lewis base

1/n

Lewis base

1/n

N

N

1

1 N

N

N

1/2

1/2

N

N

N N

N

1/3

N

N

1/4

N

(the number of equivalent positions that a molecule can attain by rotation of its parts around single bonds). For instance, in the hydrogen bonding of 1,3,5-triazine (Scheme 1.2), the statistical factor 1/n = (2.1)/(1.1)(6.1) = 1/3. The result would be the same if phenol was used instead of hydrogen fluoride, by considering the hydroxyl group as freely rotating and having the symmetry of a cone. In the complexation of the same base with diiodine, the statistical factor becomes 1/6 because σ e (I2 ) = 2. However, for the complexation of diiodine, the contribution 12 brought by diiodine to the statistical factor is often neglected when only diiodine complexes are being compared with each other (see the pK BI2 scale in Chapter 5). If one complex, say C1 , is significantly more stable than the others, K 1 dominates and K t ≈ K 1 and logK t ≈ logK 1 . Consider, for example, the diiodine complexation of the antithyroid drug carbimazole (Scheme 1.3). There are five possible locations that might be considered as diiodine acceptor sites: the two nitrogen atoms, the two oxygen atoms and the sulfur atom. Three of them, the two nitrogens and the ether oxygen, may be ruled out because of the strong electron-withdrawing effects of the thiocarbonyl and carbonyl groups. From the value K c = 3.3 l mol−1 of the diiodine complexation constant of Me2 NCOOEt, it is clear that the contribution of the carbonyl group to the measured value K t = 1411 l mol−1 (in CCl4 ) is negligible. IR and UV studies confirm that diiodine binds almost exclusively to

N N

N

+

H

F

F

H

N

N

N

σe

6

1

2

σi

1

1

1

Scheme 1.2 External and internal symmetry numbers of 1,3,5-triazine, HF and their complex.


41

K4 ≈ 0

K5 ≈ 0

K2 Cl− > Br− > I− , whereas towards Ag+ and Hg2+ it is I− > Br− > Cl− > F− (Table 1.15). Pearson [150–155] has explained reversals in basicity (acidity) such as those illustrated in Table 1.15 by qualifying the acids and bases by their so-called hardness (or softness, which is the inverse of hardness), in addition to their strength. Hard and soft acids and bases were originally defined as follows: r Soft bases: the donor atoms are of high polarizability, low electronegativity, easily oxidized and associated with low-lying empty orbitals suitable for π back-donation. r Hard bases: the donor atoms are of low polarizability, high electronegativity, difficult to oxidize and associated with empty orbitals of high energy. r Soft acids: the acceptor atoms are of low positive charge, large size, high polarizability and have several easily excited outer electrons. r Hard acids: the acceptor atoms are of high positive charge, small size, low polarizability and do not have easily excited outer electrons. Accordingly, acids and bases were divided into three classes, labelled hard, soft and borderline. This classification is shown in Table 1.16. Its usefulness arises from the generalization that hard acids prefer to coordinate to hard bases and soft acids to soft bases, all things being otherwise equal. This is the so-called Hard and Soft Acid and Base (HSAB) principle. A possible quantitative statement of the HSAB principle would be the four-parameter Equation 1.110: log K = SA SB + σA σB

(1.110)

where K is the equilibrium constant of the acid/base reaction, SA is a strength factor for the acid and SB for the base and σ A and σ B are parameters that measure the softness of the acid and of the base, respectively. In spite of many efforts, the lack of ad hoc experimental data for establishing separate scales of strength and of softness has prevented a numerical evaluation of Equation 1.110. The HSAB principle was, originally, entirely empirical. Its first theoretical rationale was given by Klopman’s perturbation theory [3, 33]. It was proposed [2, 3] to equate the strength and softness terms of Equation 1.110 to the charge-controlled and orbital-controlled terms

44


Table 1.16 Classification of some Lewis acids and bases into hard, borderline and soft. Hard

Borderline

Soft

Pd2+ , Pt2+

Be2+ , Mg2+ , Ca2+ , Sr2+

Acids Fe2+ , Co2+ , Ni2+ , Cu2+ , Zn2+ B(Me)3 , GaH3

Sc3+ , La3+ , Ce4+ , UO2 2+ Ti4+ , Cr3+ , Mn2+ , Fe3+ , Co3+

R3 C+ , NO+ , Bi3+ SO2

H+ , Li+ , Na+ , K+

BF3 , BCl3 , Al3+ , Al(Me)3 , Ga3+ CO2 , RCO+ , Si4+ SO3 HX (hydrogen-bond donor) NH3 , RNH2 H2 O, OH− , ROH, R2 O CH3 COO− , CO3 2− , NO3 − , SO4 2− F− , Cl−

Bases C5 H5 N, N3 − NO2 − , SO3 2− Br−

Cu+ , Ag+ , Au+ , Hg+ , Hg2+ , HgCl2 BH3 , Ga(Me)3 π Acceptors (trinitrobenzene, quinones, tetracyanoethylene) Br2 , Br+ , I2 , I+ , ICN

H− C2 H4 , C6 H6 , CN− , CO SCN− , R2 S, RSH, RS− R3 P, R3 As I−

of Equation 1.39, respectively. In short, quoting Jensen [2, 3], ‘soft acids are characterized by low-lying acceptor orbitals and soft bases by high-lying donor orbitals. Hard acids and hard bases have the opposite properties. Strong acids are characterized by large net positive charge densities at the acceptor atom and strong bases by large net negative charge densities at the donor atom. Weak acids and bases have small net charge densities’. Nevertheless, this interpretation does not give a precise definition of hardness and does not enable a numerical scale of hardness to be established. This has been done by Parr and Pearson [156] within the framework of density functional theory. Thus, the hardness of a chemical system (atom, molecule, supermolecule, ion, radical) has been identified [156] as the partial second derivative of the electronic energy E with respect to the number of electrons N, keeping constant the potential v due to the fixed nuclei: η=

1 2

∂2 E ∂N2

(1.111) v

Remember that the electronegativity χ is rigorously defined [157] as the negative of the first derivative of the E(N) curve: χ =−

∂E ∂N

(1.112) v


45

Using the energies of N, (N + 1) and (N − 1) electron systems, operational and approximate definitions of η and χ are obtained [156] as 1 (IE − EA) (1.113) 2 1 χ ≈ (IE + EA) (1.114) 2 where IE and EA are the ionization energy and electron affinity, respectively. Whereas the electronegativity defines whether a compound is a net electron acceptor or donor and must be constant everywhere in the chemical system, the hardness determines how easily the number of electrons can be changed and it need not be constant, that is it can have local values [158]. Local values determine which parts of a compound are the most reactive. Values of χ and η, calculated from experimental values of IE and EA, are given in Table 1.17 for some atoms, monoatomic cations and organic molecules. A more extensive list can be found in ref. [159] for inorganic species and ref. [160] for organic species. For anions and polyatomic cations, see ref. [161]. For molecules, the entries in Table 1.17 are arranged in order of decreasing electronegativity so that Lewis acids start the list and bases are at the bottom (see below). Cations are ordered by decreasing hardness (increasing softness) and atoms by atomic number. Because electronegativity and hardness are intimately related to the gain and loss of electrons by chemical compounds, they are fundamental properties for studying Lewis acid/base reactions. Examples are given below. The numerical values of η are generally satisfactory when compared with the qualitative HSAB classification of Table 1.16. For instance, the cations classified hard acids, such as Li+ , Mg2+ and Al3+ , have large η values, those classified borderline, such as Ni2+ and Cu2+ , have intermediate values, and those classified soft, such as I+ , Pd2+ and Ag+ , have small values. Moreover, the expected increased hardness with increased oxidation state is shown by the η values of Fe, Fe2+ and Fe3+ . The electronegativity determines which of two molecules, C and D, is the electron donor and which the acceptor when they form a complex. The direction of electron transfer is given by the difference η≈

(IEC − EAD ) − (IED − EAC ) = 2(χC − χD )

(1.115)

A positive value for the difference means that it costs less energy to transfer an electron from D to C. Thus, any molecule high in the list in Table 1.17 is a Lewis acid towards a molecule lower in the list, and vice versa. Hence the electronegativity of HF shows that it is an acid towards NH3 , by forming the hydrogen bond F H· · ·NH3 , but a base towards BF3 , by forming the complex HF–BF3 . A rigorous proof of the HSAB principle has been given [162], provided a restriction be added: hard likes hard and soft likes soft, among potential partners of a given electronegativity. A partial and approximate deduction of the HSAB principle can be found by considering the energy decrease, ∆E, resulting from the transfer of electron density from the base B to the acid A when a complex AB is formed. This transfer is driven by the electronegativity difference, χA − χB , and inhibited by the sum of hardness, ηA + ηB [156]. The principle of electronegativity equalization [163] states that electron density flows until a single average value of electronegativity exists everywhere in the complex.

46


Table 1.17 Some experimental values of electronegativity and hardness in eV. Atom H Li B C N O F Na Al Si P S Cl K V Cr Fe Co Ni Cu Se Br Rb Zr Nb Mo Rh Pd Ag Sn Sb Te I Ba Pt Au

χ

η

Cation

χ

η

Molecule

χ

η

7.17 3.00 4.29 6.27 7.27 7.53 10.41 2.85 3.21 4.76 5.62 6.22 7.31 2.41 3.64 3.76 4.03 4.26 4.44 4.48 5.89 7.60 2.34 3.63 3.88 3.92 4.30 4.44 4.44 4.30 4.84 5.49 6.76 2.6 5.6 5.8

6.42 2.38 4.01 5.00 7.27 6.08 7.01 2.30 2.77 3.38 4.86 4.12 4.70 1.92 3.11 3.05 3.87 3.60 3.24 3.25 3.86 4.24 1.85 3.21 2.99 3.17 3.16 3.88 3.14 3.05 3.79 3.52 3.70 2.6 3.5 3.5

Al3+ Li+ Mg2+ Sc3+ Na+ Ca2+ La3+ Fe3+ Ba2+ Rb+ Zn2+ Tl3+ Cd2+ Mn2+ Ni2+ Pb2+ Cu2+ Hg2+ Fe2+ Ti2+ Ag+ Pd2+ Cu+ Au+ Br+ I+

74.2 40.5 47.6 49.3 26.2 31.6 34.6 43.7 22.8 15.8 28.8 40.3 27.2 24.4 26.7 23.5 28.6 26.5 23.4 20.6 14.6 26.2 14.0 14.9 16.7 14.8

45.8 35.1 32.5 24.6 21.1 19.7 15.4 13.1 12.8 11.7 10.8 10.5 10.3 9.3 8.5 8.5 8.3 7.7 7.3 7.0 6.9 6.8 6.3 5.7 4.9 4.3

SO3 N2 Cl2 SO2 Br2 O2 BF3 CO I2 HF CH3 I CH3 Br HCl C2 H4 C5 H5 N C2 H2 C6 H6 PH3 CH3 Cl CH3 F H2 O NH3 (CH3 )3 As (CH3 )3 P (CH3 )2 S (CH3 )2 O (CH3 )3 N

7.2 7.0 7.0 6.7 6.6 6.3 6.2 6.1 6.0 5.0 4.9 4.8 4.7 4.4 4.4 4.4 4.0 4.0 3.8 3.2 3.1 2.9 2.8 2.8 2.7 2.0 1.5

5.5 8.6 4.6 5.6 4.0 5.9 9.7 7.9 3.4 11.0 4.7 5.8 8.0 6.2 5.0 7.0 5.2 6.0 7.5 9.4 9.5 7.9 5.5 5.9 6.0 8.0 6.3

An approximate value of ∆E is given [156] by ∆E ≈ −

(χA − χB )2 4 (ηA + ηB )

(1.116)

If both A and B are soft, ηA + ηB is a small number and, for a given reasonable electronegativity difference, ∆E is substantial and stabilizing. This explains the HSAB principle in part: soft acids prefer soft bases. Equation 1.116 refers only to the first step of the interaction between A and B. A complete expression for the interaction energy, within the framework of the local HSAB principle, can be found in the literature [164, 165].


47

10

E / eV

LUMO η

0 LUMO

−χ

η −χ

η

η HOMO

-10

HOMO I2

Figure 1.11

NH3

Frontier orbital energy diagram for I2 and NH3 .

The concepts of hardness and electronegativity can be considered within the framework of the frontier molecular orbital theory. According to Koopmans’ approximation, the frontier orbital energies are given by −E HOMO ≈ IE and

− E LUMO ≈ EA

(1.117)

Therefore, further estimates of η and χ may be written as η≈

1 (E LUMO − E HOMO ) 2

1 χ ≈ − (E LUMO + E HOMO ) 2

(1.118) (1.119)

Figure 1.11 shows the HOMO, the LUMO, the electronegativity and the hardness for the molecules I2 and NH3 . It is seen that ammonia will readily transfer electrons to I2 (since χI2 − χNH3 = 3.1 eV), leading to the halogen-bonded complex I–I· · ·NH3 . Equation 1.118 identifies the hardness with half the energy gap between HOMO and LUMO. This identification agrees with the earlier empirical definition of softness [150–155], which emphasizes polarizability. Indeed, the theory of polarizability shows that a small HOMOLUMO difference corresponds to a high polarizability. 1.11.2

The ECW and ECT Models

The reversals that occur in affinity scales for a series of Lewis bases with a change in the reference Lewis acid inspired Drago et al. [6, 76, 77] to formulate the double-scale enthalpy Equation 1.120: −∆H = E A E B + CA CB

(1.120)

48


where ∆H is the enthalpy of the reaction A + B → AB measured in the gas phase or in ‘inert’ solvents such as alkanes or CCl4 . The Lewis affinity of each neutral acid A is measured by two empirically determined parameters, EA and CA (E stands for ‘electrostatic’ and C for ‘covalent’), and the Lewis affinity of each neutral base B by two parameters, EB and CB . The E and C Equation 1.120 relates to the HSAB principle in that, for there to be a strong enthalpy of complexation, the two molecules must ‘match’ in the sense that both must have a large E parameter or a large C. Theoretically, the E and C model is consistent [3] with Klopman’s perturbation theory of intermolecular interactions, with the EA EB and CA CB terms corresponding to the chargecontrolled and orbital-controlled contributions to Equation 1.39, respectively. Marks and Drago [166] have shown that, by starting with the Mulliken valence-bond model (which yields the energy of the complex AB as the sum of electrostatic and covalent terms), an equation of the form of Equation 1.120 can be obtained by introducing certain approximations. The EA EB and CA CB terms also show [167] some correspondence with, respectively, the electrostatic, ∆EES , and polarization plus charge transfer, ∆EPOL + ∆ECT , terms of the Morokuma analysis. These comparisons give some support to relating the E and C parameters to electrostatic and covalent interactions, respectively. A detailed description of how the E and C parameters are determined is presented in the literature [76, 168–170]. The E and C values which best reproduce several hundreds of experimental enthalpy data are summarized in Table 1.18. They are√expressed in (kJ mol−1 )1/2 by multiplying the literature values, in (kcal mol−1 )1/2 , by 4.184. With diiodine arbitrarily chosen as the reference acid and assigned parameters EA = 1.02 and CA = 4.09 (kJ mol−1 )1/2 , the CA and CB values cannot be compared directly with the EA and EB values. However, the relative magnitudes of the CB (EB ) parameters give an indication of the relative covalent (electrostatic) affinities of the Lewis bases. In some cases [171, 172], a constant term must be added to the E and C equation, which takes the form −∆H = E A E B + CA CB + W

(1.121)

The term W, which is usually zero for bases, is a constant contribution to the enthalpies of reaction for a particular acid that is independent of the bases with which it reacts. For example, in the case of the dimeric acid Al2 Cl6 reacting to form B–AlCl3 adducts, W corresponds to half the enthalpy of cleaving the dimer; in the case of the acid (CF3 )3 COH forming hydrogen-bonded complexes (CF3 )3 COH· · ·B, W incorporates the enthalpy of breaking an OH· · ·F3 C intramolecular hydrogen bond. Thus, the E and C portion of Equation 1.121 corresponds to the exothermic interaction of the base with the monomeric acid or the free OH group of the hydrogen-bond donor. One can attempt [169, 170, 173] to interpret basicity scales or spectroscopic scales of basicity with the ECW model by rewriting Equation 1.121 as ∆G or ∆ν = eE B + cCB + w

(1.122)

where ∆G is a Gibbs energy of complexation and ∆ν a spectral shift upon complexation of a reference Lewis acid with a series of Lewis bases; e and c are the responses of the quantity measured to the electrostatic and covalent parameters of the bases; the constant w can arise from a variety of effects.


49

Table 1.18 Selected parameters (kJ mol–1 )1/2 of the ECW model. Acid

EA

CA

I2 IBr ICl C6 H5 OH 4-FC6 H4 OH CF3 CH2 OH c-(C4 H4 )NH CHCl3

1.02 2.45 5.97 4.64 4.70 4.23 2.82 3.19

4.09 6.73 3.40 2.19 2.27 2.17 1.39 0.90

B(CH3 )3 Al(CH3 )3 Ga(C2 H5 )3 In(CH3 )3 (CH3 )3 SnCl BF3 (g) SO2

5.93 17.71 14.22 13.50 5.87 12.48 1.04

7.36 7.53 3.03 4.40 1.45 5.87 3.19

Base

EB

CB

Base

EB

CB

Carbonyl, thionyl, phosphoryl compounds CH3 COCH3 3.56 c-(CH2 )4 CO 4.13 CH3 COOCH3 3.33 CH3 COOC2 H5 3.31 CH3 CON(CH3 )2 4.81 HCON(CH3 )2 4.48 (CH3 )2 SO 4.91 c-(CH2 )4 SO 4.99 (C6 H5 )3 PO 5.30 (C2 H5 O)3 PO 5.13 [(CH3 )2 N]3 PO 5.87 N-Oxides C5 H5 NO 4.68 4-CH3 C5 H4 NO 4.75 4-CH3 OC5 H4 NO 4.79 Ethers (C2 H5 )2 O 3.68 (C4 H9 )2 O 3.87 O(C2 H4 )2 O 3.80 c-(CH2 )4 O 3.35 (CH3 )2 O 3.44 Thioethers (CH3 )2 S 0.51 (C2 H5 )2 S c-(CH2 )4 S c-(CH2 )5 S

0.49 0.53 0.70

Acid

EA

CA

Nsp bases 2.58 1.80 1.94 2.00 2.68 2.68 3.01 3.35 3.42 2.25 3.11 4.77 5.26 6.18 3.33 3.42 2.64 4.46 3.07 7.67 8.02 8.33 7.79

ClCH2 C N CH3 C N C6 H5 C N (CH3 )2 NC N Nsp2 bases N-Methylimidazole 4-(CH3 )2 NC5 H4 N 4-CH3 OC5 H4 N 4-CH3 C5 H4 N 3-CH3 C5 H4 N C5 H5 N 3-ClC5 H4 N 4-N CC5 H4 N Nsp3 bases HC(C2 H4 )3 N (C2 H5 )3 N (CH3 )3 N (CH3 )2 NH c-(CH2 )5 NH CH3 NH2 C2 H5 NH2 NH3 Miscellaneous bases (CH3 )3 P (CH3 O)3 P (CH3 )2 Se (C6 H5 )3 PS CH3 Cl C6 H6

3.42 3.35 3.38 3.93

0.68 1.45 1.53 1.88

2.37 3.93 3.74 3.74 3.70 3.64 3.40 3.13

10.06 9.06 7.83 7.63 7.51 7.24 6.30 6.01

1.64 2.70 2.48 3.68 2.95 4.42 4.79 4.73

13.75 11.72 11.48 8.61 10.08 6.38 6.75 4.17

0.51 0.27 0.10 0.72 5.20 1.43

11.88 9.88 8.67 7.47 0.20 0.92

The ECW equation does not apply to gas-phase cation affinities. According to Drago et al. [174], a transfer-energy component is significant in the enthalpy of cation/molecule reactions M+ (g) + B(g) → MB+ (g)

50


and is accommodated by replacing W with a separate RA T B term: −∆H = E A E B + CA CB + RA TB

(1.123)

where RA is called the receptance of the cation and T B the transference of the Lewis base. An extensive list of RA values for cations and T B values for bases can be found in ref. [174]. Semi-empirical justifications for adding an electron-transfer term are presented in refs [174–176]. Equation 1.123 generally fits well the H+ , Li+ , K+ , Mg+ , Al+ , Mn+ , CH3 + , (CH3 )3 Sn+ , CpNi+ (Cp = cyclopentadiene), NO+ , H3 O+ and various ammonium cation affinities, but there are a number of systems that are not correlated by the ECT model. An example of a good fit is given by the proton affinity of dimethyl ether: PA(Me2 O) experimental = 792 kJ mol−1 PA(Me2 O) calculated = E A E B (316.3) + CA CB (81.8) + RA TB (397.7) = 796 kJ mol−1 whereas a bad fit is provided by the potassium cation affinity of dimethyl sulfoxide: KCA(Me2 SO) experimental = 146 kJ mol−1 KCA(Me2 SO) calculated = E A E B (37.9) + CA CB (0.6) + RA TB (56.5) = 95 kJ mol−1 The ECT model has been little used in the analysis of gas-phase ion chemistry. In contrast, the ECW model is generally found helpful in many fields of solution chemistry and biochemistry, as shown in several reviews [77, 173, 177–179] and in a book [6] and by its introduction in many textbooks of (mainly inorganic) chemistry (e.g. ref. [180]). The ECW model is particularly useful for showing that there is no inherent oneparameter order of Lewis affinity. Factoring and rearranging Equation 1.121 (with W = 0) lead to [179] (CB − E B ) CA − E A CB + E B −∆H = + (1.124) CA + E A 2 2 CA + E A Equation 1.124 shows that if −∆H/(CA + EA ) is plotted against RA = (CA − EA )/(CA + EA ), each base of Table 1.18 is represented by a straight line of slope (CB − EB )/2 and intercept (CB + EB )/2. The lines for quinuclidine, dimethyl selenide, diethyl sulfide, triphenylphosphine sulfide, pyridine, pyridine N-oxide, tetrahydrofuran, hexamethylphosphoric triamide, dimethyl sulfoxide, N,N-dimethylformamide, acetone, acetonitrile and benzene (representative of Nsp3 , Nsp2 , Nsp, O, S, Se and C bases) are drawn in Figure 1.12. The quantity RA on the horizontal axis ranges from +1 for EA = 0, that is, for a purely covalent (virtual) acid, to −1 for CA = 0, that is, for a purely electrostatic (virtual) acid. Among common neutral Lewis acids, the hydrogen-bond donor CHCl3 (RA = −0.56) is very electrostatic and the halogen-bond donor I2 (RA = +0.60) very covalent. The affinity order of bases toward any reference acid whose RA value is known can be determined from the plot of Figure 1.12. By just drawing a vertical line at the RA value, the proper sequence of affinity is observed as the intersections of the various lines. In Figure 1.12, vertical lines have been drawn for the acids CHCl3 and I2 . The affinity orders towards CHCl3 and I2 are listed in Table 1.19, together with the experimental enthalpy data for comparison. When the lines drawn for two different bases cross, the affinity order changes toward acids on opposite sides of the intersection. Due to the large number of intersections of base lines, it is immediately apparent that the ranking of Lewis bases will often change from acid to acid. However, if the base lines do not cross over the range of RA from −1 to +1, the affinity


51

HC(C2H4)3N

•

-ΔH / (CA+EA) / (kJ/mol)1/2

10

(CH3)2Se (C2H5)2S (C6H5)3PS C5H5N

• • •

5

• • • •

• •

• •

• • • •

••

- 0.56

-1

c-(CH2)4O ((CH3)2N)3PO (CH3)2SO HCON(CH3)2 CH3COCH3

• CH3CN

•

0

C5H5NO

•

0

-0.60

C6H6

1

RA = (CA-EA) / (CA+EA)

Figure 1.12

The E and C plot of Equation 1.124 for a series of different bases.

order for these bases will never change. For example, towards any neutral Lewis acid, the affinity order of the following bases is predicted to remain the same: (Me2 N)3 PO > Me2 SO > HCONMe2 > Me2 CO > MeC ≡ N > C6 H6 Although useful, the ECW model is not always successful. For example, it cannot incorporate the acid SbCl5 [181], which is unfortunate because SbCl5 is the reference Lewis acid of the donor number scale (see Chapter 2). A number of E and C values are also contaminated by model and/or experimental errors. A regrettable example concerns the EB

52


Table 1.19 Orders of affinity of Lewis bases towards the reference acids CHCl3 and I2 given by the E and C plot. Comparison with experimental affinities (kJ mol −1 ). Chloroform EC ranking [(CH3 )2 N]3 PO C5 H5 NO (CH3 )2 SO C5 H5 N HC(CH2 CH2 )3 N HCON(CH3 )2 c-(CH2 )4 O CH3 COCH3 CH3 C≡N (C6 H5 )3 PS (C2 H5 )2 S (CH3 )2 Se C6 H6

Diiodine

CHCl3 affinitya In C6 H12 20.5

EC ranking HC(C2 H4 )3 N (CH3 )2 Se (C2 H5 )2 S C5 H5 N (C6 H5 )3 PS C5 H5 NO c-(CH2 )4 O [(CH3 )2 N]3 PO (CH3 )2 SO HCON(CH3 )2 CH3 COCH3 CH3 C≡N C6 H6

17.2 15.1 9.8 7.1 8.4

I2 affinityb In heptane

In CCl4

43.1c 35.2 32.8 22.2 30.5 28.8 20.9 15.3 13.4d 6.8

36.0 34.6 32.3 30.5 23.6 19.9 16.9 15.5 8.8 6.4

a

From ref. [84]. From Chapter 5. c Value for piperidine. d Value for butyronitrile. b

and CB values of the halogen base CH3 Cl, which lead to an unacceptable 4-fluorophenol affinity: 4-FC6 H4 OH affinity (CH3 Cl) calc. = (4.70 × 5.20) + (2.27 × 0.20) = 24.9 kJ mol−1 4-FC6 H4 OH affinity (C5 H11 Cl) exp . = 6.96 and 8.95 kJ mol−1 in CCl4 and C6 H12 , respectively

1.11.3

The Beta and Xi Equation

Another dual-parameter equation has been introduced by Kamlet, Taft et al. [182] for correlating Lewis basicity: BDP = bβ + eξ + BDP0

(1.125)

where BDP is a basicity-dependent property, such as a spectral shift, a Gibbs energy or an enthalpy, β is a hydrogen-bond basicity parameter, whose definition is given in Chapter 4, ξ is an empirical parameter defined to quantify the extent of family dependence observed in BDP versus β plots, b and e are the regression coefficients and BDP0 is the intercept. Equation 1.125 formally and conceptually resembles the E and C Equations 1.120 and 1.121, with β corresponding to EB and ξ to CB . Indeed, hydrogen bonding is mainly electrostatic in origin and electrostatic bonding of the Lewis acid/base complexes must increase in strength with β. Further, the empirical parameter ξ can be interpreted as a coordinate covalency parameter. This follows from the increase in ξ with decrease in the charge on the hydrogen atom in BH+ (shown in Table 1.20), since this charge decrease


53

Table 1.20 Comparison of the coordinate covalency parameter ξ with the charge q(H+ ) on hydrogen in BH+ . ξ

Family PO bases CO and SO bases Nitriles Single-bonded O bases Pyridines Amines a

q(H+ )a

−0.20 0.00 0.10 0.20 0.60 1.00

[BH]+

0.62

[H2 C=OH]+

0.57 0.55 0.47

[H3 O]+ [C5 H5 NH]+ [NH4 ]+

Mulliken charge calculated at the HF/6–31G∗ level.

depends on the ability of the base B to transfer a charge into the empty 1s orbital of the proton, that is, to form a coordinate covalent bond with the proton. There is, however, a difference between the parameters CB and ξ : CB is a property of the individual base whereas ξ is taken to be a property of a whole family of bases. Sixteen diverse BDPs of neutral oxygen and nitrogen bases have been correlated by the β and ξ equation. These include hydrogen-bond, BF3 and I2 affinities and also proton affinity in the gas phase [183] and Gibbs energy of proton transfer to the aqueous bases from aqueous NH4 + . Because steric effects, solvent effects, entropic effects and strong conjugative π -electron donation to the base centre can lead to severe deviations from Equation 1.125, this equation is not expected to be highly precise in general. The main merits of the β and ξ equation are (i) to demonstrate the usefulness of hydrogen-bond parameters as electrostatic parameters of Lewis basicity and (ii) to show that a proper separation of Lewis bases into families is a convenient way to simplify the generalized quantitative treatment of Lewis basicity. However, the parameter β, which was determined by a somewhat vague method, has to be replaced by a more clearly defined hydrogen-bond basicity parameter, such as those listed in Chapter 4.

1.11.4

A Chemometric Approach

By the method of principal component analysis (PCA), Maria, Gal et al. [184] determined the dimensionality of basicity scales, affinity scales and spectroscopic scales of basicity. They studied a data matrix of 22 neutral oxygen and nitrogen bases and 10 scales. These include: (i) IR and UV spectroscopic shifts upon hydrogen bonding, (ii) Gibbs energies of hydrogen bonding and halogen bonding, (iii) enthalpies of complexation with BF3 , SbCl5 and I2 and (iv) enthalpies of proton transfer. PCA shows that 95% of the total variance of the data could be accounted for by only two factors, F 1 and F 2 . From the correlation of F 1 with proton affinity and F 2 with potassium cation affinity, it was proposed that F 1 represents a blend of electrostatic and covalent effects and F 2 mainly electrostatic effects. The third factor, representing in part the steric hindrance of complexation, is of marginal importance. Hence the correlation of a BDP with F 1 and F 2 by means of Equation 1.126: BDP = S1 F1 + S2 F2 + BDP0

(1.126)

54

Lewis Basicity and Affinity Scales Table 1.21 Covalent/electrostatic character of eight affinity scales given by chemometric (θ ), E and C (CA /EA ) and β and ξ (e/b) analysis. Comparison with the hardness ηA of the acid. Reference Lewis acid

θ (◦ )

CA /EA

CHCl3 4-FC6 H4 OH SbCl5 (CF3 )2 CHOH BF3 HSO3 F SO2 I2

60 42 40 29 −4 −22 −45 −51

0.28 0.47 (2.9)a 0.46 0.47 3.1 4.0

e/b

ηA (eV)

0.31 0.26b

(9.5)c

0.57 0.57b

9.7

0.84

5.6 3.4

a

Calculated from four enthalpies only. Calculated here from data in ref. [184]: SbCl5 affinity = 38.0(±3.0)β + 9.8(±1.8)ξ − 0.6(±2)(n = 18; r = 0.960; s = 2.3kcal mol−1 ) −∆Hi = 38.6(±4.1)β + 22.1(±2.3)ξ − 0.4(±3)(n = 20; r = 0.956; s = 3.2kcal mol−1 ) c Hardness of H2 O. b

yields the sensitivities S1 and S2 of the BDP to the factors F 1 and F 2 , respectively. Any BDP may thus be represented by a point of coordinates S1 and S2 in a two-dimensional space. In 1/2 and makes polar coordinates, the corresponding vector has a magnitude ρ = S12 + S22 an angle θ = arctan(S2 /S1 ) with the S1 -axis; ρ is related to the strength of the Lewis acid and θ to the electrostatic/covalent character of the BDP. If θ is close to 90◦ , the BDP presents a character like that of F 2 , that is mainly electrostatic. A decrease in θ corresponds to a larger dependence on F 1 , that is an increase in covalent character. θ values for eight affinity scales are summarized in Table 1.21 and compared with the covalent/electrostatic ratios of the EC and βξ models and with the acid hardness. There is good agreement between the four approaches, if one excludes SbCl5 and BF3 , apparently anomalous in the EC model. PCA also supports the separation of bases into the same families as those of the βξ model, as shown by the F 1 versus F 2 plot in Figure 1.13. Hence the multiple correlation of F 1 with F 2 and ξ (Equation 1.127) is expected by the physical significance attributed to these parameters: F1 = 3.11(±0.27)F2 + 1.90(±0.13) ξ − 0.37(±0.04)

(1.127)

n = 20(O and N bases), r = 0.964, s = 0.12. The application of Equation 1.126 to the thermodynamic functions of a variety of Lewis acid/base reactions shows [184] that their responses to the electrostatic/covalent character increase in the order −∆S < −∆H < −∆G. A rather covalent character is found for the spectroscopic scale of hydrogen-bond basicity ∆ν(OH), and the sequence −∆S < ∆ν(OH) < −∆H < −∆G is observed for the formation of hydrogen bonds. A thorough analysis [185] of hydrogen-bond basicity scales by means of Equation 1.126 yields θ values ranging from 53◦ (4-fluorophenol complexes in CH2 Cl2 ) to 86◦ (Ph2 NH complexes in CCl4 ), depending on the reference hydrogen-bond donor and the


55

1 0.8 0.6 0.4

F1

0.2

PhNMe2

0 -0.2 -0.4 -0.6 -0.8 PhNO2

-1 -0.5

-0.3

-0.1

0.1

0.3

F2

Figure 1.13 F1 vs F2 plot showing the separation of bases into the families of the βξ treatment: PO bases ( ), CO ( ) and SO bases ( ), nitrile ( ), ethers ( ), pyridines ( ) and amine ( ). The ξ parameters of PhNO2 ( ) and PhNMe2 ( ) are unknown, but the positions of PhNO2 near oxygen bases and of PhNMe2 near the amine appear satisfactory.

solvent. This variation in θ precludes the construction of any general scale of hydrogenbond basicity. However, a reasonably general scale could be set up either by choosing a reference hydrogen-bond donor having a θ value in the middle of the 53–86◦ range, for example 4-fluorophenol [121] in CCl4 (θ = 70◦ ), or by averaging proper hydrogenbond basicity scales [186]. Indeed, plots of logK for bases against OH (H2 O, alcohols, phenols) and strong NH (imides, amides, HNCS) hydrogen-bond donors versus logK for bases against 4-fluorophenol are family independent. Only CH (CHCl3 , alk-1-ynes) and weak NH (aromatic amines) hydrogen-bond donors show a moderate family dependence in such linear Gibbs energy relationships. Panchenko et al. found [187] that the family independence can be restored by using the two-scale Equation 1.128: logK = b pK BHX + c pK BI2 + logK 0

(1.128)

where pK BHX is the 4-fluorophenol basicity (see Chapter 4), pK BI2 is the diiodine basicity (see Chapter 5), b and c are regression coefficients and log K 0 is the intercept. For example, the application of Equation 1.128 to the diphenylamine basicity yields [187] logK = 0.73(±0.06) pK BHX − 0.25(±0.04) pK BI2 − 0.73 n = 17, r = 0.952, s = 0.15.

(1.129)

56


The correlation coefficient r = 0.952 can be compared with the value r = 0.767 found in the correlation of diphenylamine basicity with the single 4-fluorophenol basicity. In the same vein, a reasonably general hydrogen-bond acidity scale has been established [188]. Consequently, a reasonably general treatment of hydrogen-bond complexation constants (as logK values) in CCl4 has been reported [189] for more than a 1000 hydrogen-bond donor/acceptor pairs (corresponding to 89 hydrogen-bond donors and 215 hydrogen-bond acceptors) in the form of Equation 1.130: logK = 7.354(±0.019) α2H β2H − 1.094(±0.007)

(1.130)

n = 1312, r = 0.9956, s = 0.093 where α2H is a hydrogen-bond acidity parameter [188] and β2H a hydrogen-bond basicity parameter [186], both scaled to zero for non-hydrogen-bonding molecules (e.g. alkanes). The subscript 2 and the superscript H indicate that these parameters are specific to the solute (that is, they are not solvent scales) and differ from their congeners α and β. Similar treatments are also successful in the gas phase [106] (Equation 1.131) and in CCl3 CH3 [190] (Equation 1.132): logK = 9.13(±0.32) α2H β2H − 0.87(±0.11)

(1.131)

n = 23, r = 0.987, s = 0.20 logK = 6.86(±0.15) α2H β2H − 1.14(±0.07)

(1.132)

n = 84, r = 0.980, s = 0.16. A comparison of the slopes and intercepts of Equations 1.130–1.132 shows that hydrogenbond complexation increases with the medium in the order CCl3 CH3 < CCl4 < gas. 1.11.5

Quantum Chemical Descriptors for Basicity Scales

The previous Sσ , EC, βξ and F 1 F 2 two-term equations for interpreting and predicting basicities, affinities and spectroscopic scales of basicity rely on empirical parameters, either not clearly defined (S, ξ , E, C, β) or needing advanced statistical procedures for their determination (F 1 , F 2 ). Moreover, if new systems are added to the original database, the numerical values of the parameters will change, as seen by the significant evolution of E, C and β values over the years. Lastly, since many key systems have not yet been measured or are not measurable, the number of available values is disappointingly low (22 for F 1 , F 2 ) compared with the large number of bases that are of interest to chemists. These shortcomings of empirical approaches and the exponential growth of computer power in the last four decades have led to empirical parameters being replaced by calculated quantum chemical descriptors of basicity. The use of computationally derived parameters in the study of basicity, or related fields, has seen much investigation. See, for example the QSPR (quantitative structure–property relationship) [191], TLSER (theoretical linear solvation energy relationship) [192], GIPF (general interaction properties function) [193] and related methods [194–196]. The analysis and prediction of Lewis basicity by a modified GIPF method [197] is presented below.


57

Table 1.22 Standardized coefficientsa and statistics of Equation 1.133. α

β

γ

ε

n

r

s

−0.659 −0.573

−0.346 −0.341 −0.429 −0.540

−0.388 −0.591 −0.780 −0.524

3.2 194 17.1 206.4

25 37 26 36

0.979 0.966 0.950 0.975

0.4 29 1.0 3.6

Property Phenol affinity ∆ν(OH) Diiodine affinity Proton affinity a b

b

−0.266

Regression coefficients have been standardized to allow direct comparisons. Not significant.

In this method, the Lewis bases are characterized by three HF/6–31G computed properties: r The spatial minimum in the electrostatic potential, V min , associated with the electrondonor heteroatom. The electrostatic potential is the energy of interaction of a point positive charge with the nuclei and electrons of a molecule. r The polarization potential, PVmin , evaluated at the position of V min . It gives the energy of electronic reorganization of the molecule as a result of its interaction with a point positive charge. r I¯S,min , the surface minimum in the average local ionization energy (the molecular surface is defined by the 0.001 au contour of the electron density). V min , PVmin and I¯S,min are closely related to the electrostatic, polarization and chargetransfer components, respectively, of the Morokuma decomposition of proton affinity. It has been shown [197] that the spectroscopic scale ∆ν(OH) (related to methanol), the phenol affinity, the diiodine affinity and the proton affinity of 42 nitrogen, oxygen and sulfur bases can be correlated by the triple-scale Equation 1.133: −∆H (or ∆νOH) = αVmin + β PV min + γ I¯S,min + ε

(1.133)

The regression coefficients α, β and γ , the intercept ε and the statistics are given in Table 1.22. Taking into account the variety of bases and BDPs, the correlations are satisfactory. The calculation of affinity scales purely from the electronic structure of bases appears a promising method. An even better method, but of greater computational difficulty (see above), is the calculation of the properties of model complexes. For example, ∆Eel , ∆H or ∆G calculated for the complexation of the small hydrogen-bond donors H2 O [198] or HF [199] were found to be successful descriptors of the basicity towards larger ones (that is of greater computational cost) such as 4-FC6 H4 OH. Ultimately, the affinity or the basicity scales themselves might be computed. This has been done successfully for the methanol affinity in the gas phase [41]. However, the MP2/aug-cc-pVTZ//B3LYP/6–31+G(d,p) costly level required to obtain good agreement with experimental affinities has limited the scale to a few small Lewis bases. Proton and cation affinities and basicities in the gas phase are now computed on a routine basis (see Chapter 6). Nevertheless, the size of Lewis bases and cations and the number of bases studied are inversely proportional to the level of theory, that is, to the agreement with experimental data. As far as extended solution basicity scales towards usual Lewis acids

58


are concerned, it is not realistic to expect their reliable computation for a number of years or even decades, mainly because of the difficult calculation of solvent effects and the vibrational entropy of complexes. The construction of basicity scales remains a task for experimental chemists.

1.12

Concluding Remarks and Content of Chapters 2–7

There are virtually as many possible Lewis basicity scales as there are Lewis acids. This is a dramatic consequence of the Lewis definition of acids, which has considerably enlarged the number of chemical species showing an acid character and made the proton lose the status of reference acid. Mathematically, the aqueous Brönsted basicity scale corresponds to a one-column data matrix. In fact, several columns are needed to take into account the influence of the medium but two media have been studied most, water and the gas phase. So the two most filled, and consequently most used, columns are the pK BH+ and GB scales. In the field of proton affinity, no column attains the degree of completeness and reliability of the gas-phase PA column. There is no need for spectroscopic scales of Brönsted basicity since the thermodynamic scales are satisfactory. The mathematics of the Lewis acid/base concept is that of a data matrix of m rows and n columns. Data are complexation constants, as logK or ∆G. Each row corresponds to a Lewis acidity scale towards a reference base Bi◦ (i = 1 to m) and each column corresponds to a Lewis basicity scale towards a reference acid A◦j (j = 1 to n). For a rigorous treatment, the data measured in different media cannot be mixed in the same data matrix. In the matrix measuring Lewis affinity, the data are complexation enthalpies. There are extrathermodynamic relationships (isoequilibrium relationships or enthalpy–entropy compensation law) which allow transformations between blocks of the affinity and basicity matrices. In the principal component analysis of Lewis basicity, this justifies, somewhat, the mixing of affinity columns and basicity columns in a unified basicity–affinity matrix. At first glance, the size and completeness of the currently available basicity and/or affinity data matrices are rather disappointing. It appears that much information is lacking in this fundamental field of chemistry. To give an order of magnitude, the affinity matrix supporting the EC model, limited to neutral Lewis acids and bases but mixing enthalpies measured in ‘poorly solvating solvents’ and the gas phase, had a 43 (bases) × 31 (acids) order and an occupancy rate of 21% in the 1971 version [168] and a 50 × 43 order and a rate of 23% in the 1994 version [6]. One can conclude from these figures that, whereas the number of reference Lewis acids (partially) studied appears correct for constructing basicity scales, that of Lewis bases is very low compared with the ∼7000 bases of the pK BH+ database [11] and ∼3000 bases of the GB database [27, 28]. There are two main reasons for the paucity of data available for constructing basicity and affinity scales. First, the observation by Lewis himself in 1938 [30] that ‘in studying acids and bases, we find that the relative strength depends not only upon the chosen solvent but also upon the particular base or acid used for reference’ very soon discouraged many chemists from measuring Lewis basicity (acidity) quantitatively, all the more so since statements similar to that of Lewis have become commonplace in the chemistry literature.


59

Second, there are in practice not many systems in which a Lewis reference acid reacts with a series of diversified bases for which a series of reliable equilibrium constants can be measured easily by known physicochemical techniques over a large range of values (that is, from possibly very large to possibly very low equilibrium constants), and in the same conditions of temperature and medium. This is why chemists have turned to the determination of spectroscopic scales of Lewis basicity (affinity). Fortunately, there are a number of theoretical, statistical and empirical reasons to believe that Lewis basicity (affinity) depends on a limited number of factors. From the quantum chemical point of view, the acid/base interaction energy can be partitioned into five terms (electrostatic, dispersion, polarization, charge transfer and exchange–repulsion). By a principal component analysis [184], ∼99% of the variance of an affinity/basicity data matrix can be explained by three factors, the first two being by far the most important. A number of experimental affinity and basicity scales, and of spectroscopic scales of basicity, can be correlated by two parameters, using the EC or βξ equations, or three quantum chemical descriptors of basicity [197]. However, these statistical and empirical approaches are limited to systems where steric effects and π back-bonding are not important. Due to the existence of a limited number of explanatory variables for predicting and possibly interpreting Lewis basicity, a limited number of scales should be sufficient to quantify the Lewis definitions of acids and bases. The search for these few basicity scales in the literature gave the following results. In spite of a very large amount of data, listed in two books by Sillen and Marten (1964, second edition [148], and 1971, supplement [149]), the stability constants of metal-ion complexes with organic and inorganic ligands, measured mainly in aqueous solution, could not furnish any useful basicity scale. They did, however, help to establish the important HSAB classification and principle (Pearson, 1963–1969) [150–155]. Among numerous thermodynamic measurements on the complexes of covalent metal halides, antimony pentachloride was chosen by Gutmann (1966–1968) [78] as the reference Lewis acid for constructing an SbCl5 affinity scale and developing the ‘donor number’ concept. This scale is critically presented in Chapter 2. Following the use of boron acids by Brown et al. (1953–1955) [91] in his classical work on steric effects in the complexation of amines with trimethylboron, an extensive set of calorimetric measurements was provided by the studies of BF3 complexes by Gal, Maria et al. (1970–1992) [80]. Boron trifluoride is the archetype of Lewis acids in the original Lewis definition. The promising BF3 affinity scale is presented in Chapter 3. The importance of hydrogen bonding in chemistry, biochemistry and physics has led to the thermodynamic and spectroscopic measurement of thousands of hydrogen-bonded systems, reviewed in an excellent book by Joesten and Schaad (1974) [84]. The largest and most reliable set of equilibrium constants for hydrogen bonding to a series of bases refers to the acid 4-fluorophenol (Taft et al., 1969–1972) [116, 121]. This 4-fluorophenol basicity scale, and the corresponding 4-fluorophenol affinity scale pioneered by Arnett et al. (1969–1974) [23], have been considerably extended by Berthelot, Laurence et al. (1988–2008). The construction of the 4-FC6 H4 OH affinity scale enables a comparison with the SbCl5 and BF3 ones. The observation by Badger and Bauer (1937) [200] that the O H stretching frequencies are shifted upon hydrogen bonding, and that this infrared shift is correlated with the enthalpy of hydrogen bonding, has justified the construction of several IR spectroscopic scales of hydrogen-bond basicity. Methanol is one of the most

60


convenient reference acids for this purpose and infrared shifts, ∆ν(OH), of this probe have been systematically measured by Berthelot, Laurence et al. (1976–1988). By means of the solvatochromic comparison method, Kamlet and Taft (1976) [201] have set up the β scale of solvent hydrogen-bond basicity, a key parameter in the linear solvation energy relationships developed for the quantitative study of solvent effects. Improved and extended β(OH) and β(NH) scales defined from the UV shifts on hydrogen bonding of 4-nitrophenol and 4-nitroaniline, respectively, have been constructed by Laurence, Nicolet et al. (1986). Extended scales of 4-FC6 H4 OH basicity and affinity, ∆ν(OH), β(OH) and β(NH), are tabulated in Chapter 4. The discovery of charge-transfer bands in the UV spectra of diiodine complexes (Benesi and Hildebrand, 1949) [202] and the development of the underlying theory (Mulliken, 1952) [31] initiated a wealth of thermodynamic and spectroscopic measurements on diiodine complexes, mainly in the period 1949–1980. Complementary measurements by Berthelot, Guiheneuf, Laurence et al. (1970–2002) and Abboud et al. (1973–2004) enabled a homogeneous scale of diiodine basicity to be constructed. In addition, recommended values of diiodine affinity have been compiled from the literature (Laurence, 2006), for comparison with the SbCl5 , BF3 and 4-FC6 H4 OH affinity scales. UV and/or IR shifts upon complexation of the acids I2 , ICl and ICN have also been systematically measured by Berthelot, Laurence, Nicolet et al. (1981–1985). These thermodynamic and spectroscopic scales will allow the recent concept of a halogen bond to be treated quantitatively. They can be found in Chapter 5. The development of mass spectrometric techniques has led to the construction of not only the well-known GB and PA scales, but also the lithium cation basicity scale (Taft et al., 1990; Burk, Koppel, Gal et al., 2000) [93] and many metal cation basicity and affinity scales in the gas phase. A selection of the most informative scales is presented in Chapter 6. Because basicity is strongly related to the structure of complexes, Chapters 2–6 also contain a description of the main structural features of each kind of complex studied in the book. Experimental structures of complexes were taken mostly from the Cambridge Structural Database (CSD) [50] for the solid-state structures and from the MOGADOC database [49] for the gas-phase structures. The last chapter gives examples of reliable experimental determinations of most of the scales tabulated in this book. This should allow any chemist in academia or industry, and also any graduate or upper-level undergraduate student in chemistry, familiar with IR and UV spectrometry, calorimetry and mass spectrometry, to determine new values for the scales and the molecules of interest to them.

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179. Cramer, R.E. and Bopp, T.T. (1977) The great E & C plot. A graphical display of the enthalpies of adduct formation for Lewis acids and bases. J. Chem. Educ., 54, 612–613. 180. Cotton, F.A., Wilkinson, G. and Gaus, P.L. (1987) Basic Inorganic Chemistry, John Wiley & Sons, Inc., New York. 181. Drago, R.S., Kroeger, M.K. and Stahlbush, J.R. (1981) An E and C analysis of donor numbers and Soret band shifts in adducts of zinc tetraphenylporphine. Inorg. Chem., 20, 306– 308. 182. Kamlet, M.J., Gal, J.F., Maria, P.C. and Taft, R.W. (1985) Linear solvation energy relationships. Part 32. A coordinate covalency parameter, ξ , which, in combination with the hydrogen bond acceptor basicity parameter, β , permits correlation of many properties of neutral oxygen and nitrogen bases (including aqueous pK a ) . J. Chem. Soc., Perkin Trans. 2, 1583–1589. 183. Kamlet, M.J., Doherty, R.M., Abboud, J.L.M. et al. (1986) Solubility: a new look. Chemtech, 16, 566–576. 184. Maria, P.C., Gal, J.F., De Franceschi, J. and Fargin, E. (1987) Chemometrics of solvent basicity: multivariate analysis of the basicity scales relevant to nonprotogenic solvents. J. Am. Chem. Soc., 109, 483–492. 185. Abraham, M.H., Grellier, P.L., Prior, D.V. et al. (1989) Hydrogen-bonding. Part 4. An analysis of solute hydrogen-bond basicity in terms of complexation constants (logK) using F1 and F2 factors, the principal components of different kinds of basicity. J. Phys. Org. Chem., 2, 243– 254. 186. Abraham, M.H., Grellier, P.L., Prior, D.V. et al. (1990) Hydrogen bonding. Part 10. A scale of solute hydrogen-bond basicity using logK values for complexation in tetrachloromethane. J. Chem. Soc., Perkin Trans. 2, 521–529. 187. Panchenko, B.V., Oleinik, N.M., Sadovskii, Y.S. et al. (1980) On relationships between constants of hydrogen-bonded complex formation. I. Org. React. (Tartu), 17, 65–87. 188. Abraham, M.H., Grellier, P.L., Prior, D.V. et al. (1989) Hydrogen bonding. Part 7. A scale of solute hydrogen-bond acidity based on logK values for complexation in tetrachloromethane. J. Chem. Soc., Perkin Trans. 2, 699–711. 189. Abraham, M.H., Grellier, P.L., Prior, D.V. et al. (1988) A general treatment of hydrogen bond complexation constants in tetrachloromethane. J. Am. Chem. Soc., 110, 8534– 8536. 190. Abraham, M.H., Berthelot, M., Laurence, C. and Taylor, P.J. (1998) Analysis of hydrogen-bond complexation constants in 1,1,1-trichloroethane: the α2H β2H relationship. J. Chem. Soc., Perkin Trans. 2, 187–192. 191. Katritzky, A.R., Dobchev, D.A. and Karelson, M. (2006) Physical, chemical, and technological property correlation with chemical structure: the potential of QSPR. Z. Naturforsch., B: Chem. Sci., 61, 373–384. 192. Lowrey, A.H., Cramer, C.J., Urban, J.J. and Famini, G.R. (1995) Quantum chemical descriptors for linear solvation energy relationships. Comput. Chem., 19, 209–215. 193. Hagelin, H., Murray, J.S., Brinck, T. et al. (1995) Family-independent relationships between computed molecular surface quantities and solute hydrogen bond acidity/basicity and soluteinduced methanol O H infrared frequency shifts. Can. J. Chem., 73, 483–488. 194. Platts, J.A. (2000) Theoretical prediction of hydrogen bond donor capacity. Phys. Chem. Chem. Phys., 2, 973–980. 195. Lamarche, O. and Platts, J.A. (2003) Complementary nature of hydrogen bond basicity and acidity scales from electrostatic and atoms in molecules properties. Phys. Chem. Chem. Phys., 5, 677–684. 196. Catalan, J., Palomar, J., Diaz, C. and de Paz, J.L.G. (1997) On solvent basicity: analysis of the SB scale. J. Phys. Chem. A, 101, 5183–5189. 197. Brinck, T. (1997) Modified interaction properties function for the analysis and prediction of Lewis basicities. J. Phys. Chem. A, 101, 3408–3415. 198. Rablen, P.R., Lockman, J.W. and Jorgensen, W.L. (1998) Ab initio study of hydrogen-bonded complexes of small organic molecules with water. J. Phys. Chem. A, 102, 3782–3797. 199. Lamarche, O. and Platts, J.A. (2002) Theoretical prediction of the hydrogen-bond basicity pK HB . Chem. Eur. J., 8, 457–466.


69

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2 The Donor Number or SbCl5 Affinity Scale The first calorimetric measurements on the reaction of SbCl5 with Lewis bases in solution seem to have been made in 1963 by Olofsson [1]. Later, this author published a series of papers between 1963 and 1973 on the enthalpy of complexation of SbCl5 with numerous carbonyl bases [2–11], and also a few ethers [11, 12], methanol [12], water [13] and nitrobenzene [11]. Gutmann extended this series to a host of different Lewis bases [14, 15] and proposed, in 1966, the concept of donor number (DN) [16–18] to express quantitatively the Lewis basicity of solvents. Although the DN scale has been proposed and extensively used [19–21] as a solvent parameter, it relies on measurements made on dilute solutions of bases. It is, therefore, a solute scale and not a solvent scale of Lewis affinity.

2.1

Structure of SbCl5 Complexes

The SbCl5 molecule shows D3h symmetry in its trigonal bipyramidal configuration. By complexation with one molecule of Lewis base, B, it distorts to an irregular octahedron as shown by the X-ray structures of the solid complexes listed in Table 2.1. The X-ray structure of a representative complex is shown in Figure 2.1. It can be seen ˚ than that the Sb O bond of the HCONMe2 –SbCl5 complex [22] is much shorter (2.048 A) ˚ and slightly longer than the sum of the sum of the van der Waals radii (2.2 + 1.4 = 3.6 A) ˚ of the antimony and oxygen atoms. The Sb O the covalent radii (1.41 + 0.56 = 1.97 A) vector points in the direction of a putative sp2 lone pair of oxygen (SbOC = 124.5◦ ). The preferred lone pair of oxygen is the less sterically hindered one, on the hydrogen side. Interestingly, SbCl5 coordinates to the nitrogen and not to the sulfur atom of the polyfunctional bases S2 N2 (CSD reference code GIRXEK) and S4 N4 [23]. According to the HSAB rule that hard acids prefer hard bases, and since nitrogen bases are classified as hard or borderline whereas sulfur bases are soft, this site preference enables SbCl5 to be Lewis Basicity and Affinity Scales: Data and Measurement C 2010 John Wiley & Sons, Ltd


72


Table 2.1 Solid-state structures of SbCl5 complexes: references and intermolecular distances ˚ d (A). Lewis base

Ref.a

d (Sb B)

Bond

N C C N ClCOCH2 CH2 COCl C3 Cl4 O3 (tetrachloroethylene carbonate) C6 H5 COCl 3-MeC6 H4 COCl S2 N2 ClC N 4-MeC6 H4 COCl C10 H15 BrO (endo-3-bromocamphor) MeC N S4 N4 C6 H5 COCH2 Br C16 H18 O (5-phenyl-2-adamantanone) i-PrN C Ni-Pr ClSOMe ClSOMe i-Pr2 NC N Me2 N N O Cl3 PO Ph2 SO Cl2 SeO C6 H7 NO (3-methylpyridine N-oxide) Me2 SO HCONMe2 Ph3 PO H2 O (H-bonded to dioxane) Me3 PO C12 H8 AsClO (10-chlorophenoxarsine oxide) C17 H21 NO2 [(4-tert-butyl-1-pyridinio)(4methoxyphenyl)methoxy]

RAGRAT SUCCSB ETCBSB BZCSBC20 PBATUC GIRXEK RAGREX TOLCSB VIVFOV SBCACN [23] VIVFIP FORSAG DUYYAX ZORFAN VURKAU CIDCUN NEYHII [24] JOHHIX [25] ULOYAV CURDUO SBCLMF PINSIO DOXSBC ACMEPO POXCSB FOCFIM

2.663 2.424 2.402 2.317 2.295 2.29 2.286 2.254 2.236 2.232 2.17 2.167 2.163 2.153 2.151 2.149 2.145 2.113 2.11 2.091 2.08 2.067 2.052 2.048 2.039, 2.035 2.025 1.996 1.991 1.945

(Sb (Sb (Sb (Sb (Sb (Sb (Sb (Sb (Sb (Sb (Sb (Sb (Sb (Sb (Sb (Sb (Sb (Sb (Sb (Sb (Sb (Sb (Sb (Sb (Sb (Sb (Sb (Sb (Sb

a

Reference code of the CSD database [26] or bibliographic reference.

Figure 2.1

Structure of the complex of SbCl5 with N,N-dimethylformamide.

Nsp) O) O) O) O) N) Nsp) O) O) Nsp) N) O) O) Nsp2 ) O) O) Nsp) O) O) O) O) O) O) O) O) O) O) O) O)

The Donor Number or SbCl5 Affinity Scale

73

classified a hard acid. Other structural and thermodynamic studies also confirm the hard character of SbCl5 [21].

2.2

Definition of the Donor Number Scale

The formation of the 1 : 1 complex B–SbCl5 at 298 K in a dilute solution of 1,2dichloroethane can be written Cl

Cl Cl

Cl

Sb

Cl +

Cl

B

Cl Sb

Cl

Cl

Cl B

(2.1)

The donor number (DN), or donicity, of the Lewis base was defined [16–18] as the negative enthalpy (in kcal mol−1 ) of this reaction: DN = −∆H (2.1)

(2.2)

In the terminology of this book, DN is the SbCl5 affinity, denoted SbCl5 A. For solubility reasons, the SbCl5 complexation could not be accomplished in a less polar solvent. For example, the complexes of the following bases were not soluble enough in CCl4 : dimethyl sulfoxide, dimethylacetamide, tetramethylurea, trihexylamine and octylamine [11].

2.3

Experimental Determination of the Donor Number

Heats were measured by thermometry in an adiabatic calorimeter [1] by breaking a sealed glass ampoule containing the pure base, (i) first in the pure solvent and (ii) second in a solution of excess SbCl5 . The first measurement gives the enthalpy of solution of the base in the solvent, ∆H(2.3), which is generally small and may be positive or negative. It corresponds to the reaction B(pure) → B(soln)

(2.3)

The second measurement corresponds to the reaction of enthalpy ∆H(2.4): B(pure) + SbCl5 (soln) → B−SbCl5 (soln)

(2.4)

∆H(2.1) is obtained from the difference ∆H(2.4) – ∆H(2.3). Typically, the reaction mixtures in the calorimeter were 1,2-dichloroethane solutions containing 0.01–0.03 mol l−1 of B and usually about 0.5 mol l−1 of SbCl5 . In cases where complexation was not complete in the concentration range used, it was necessary to determine the equilibrium constant, K. Both the enthalpy of complexation and the equilibrium constant could be calculated simultaneously from the calorimetric data [6]. K values could also be determined independently by a spectrometric method [14].

74

2.4


The Donor Number Scale: Data

Over 100 SbCl5 affinity values are collected in Table 2.2. The sample of Lewis bases contains mainly oxygen bases and, among these, mainly carbonyl compounds. Only nine nitrogen bases (seven nitriles, one pyridine, one amine) have been studied. The scale lacks second-row bases, carbon bases and halogen bases. We have not reported the values estimated indirectly from correlations between DN and solvent basicity-dependent properties, such as nuclear magnetic resonance chemical shifts of the 23 Na nucleus [27] or of the chloroform proton [28]. These values would be valid only if there was no separation into families of bases in the correlation chart, which is rarely found. As shown in this book, family-dependent correlations between basicity-dependent properties are generally the norm. When different values are given by different authors for the same base, we have selected the value that we consider the most reliable on the basis of critical discussions in the literature. For example, Gutmann’s DN values for weak bases seem underestimated because, apparently, no correction was made to take into account the dissociation of the complex. Such is the case of nitrobenzene: the DN value (converted into kJ mol−1 ) of 18.4 is too low compared with the value of 34 kJ mol−1 given by Olofsson [11]. The latter value includes the equilibrium constant K c = 20 l mol−1 in determining the equilibrium concentration of the complex C6 H5 NO2 –SbCl5 [11], while Gutmann’s value is calculated under the incorrect assumption that SbCl5 is totally complexed. In the case of triethylamine, the Menschutkin reaction (2.5): Et3 N + ClCH2 CH2 Cl → [Et3 NCH2 CH2 Cl]+ Cl−

(2.5)

with the chlorinated solvent is catalysed by SbCl5 , as shown by conductivity measurements. Accordingly, the measurement was carried out [29] in a less reactive solvent, CH2 Cl2 , and at a lower temperature, −55◦ C. These reliable conditions give a value for NEt3 that is half the indirect value ordinarily quoted [19–21]. This discrepancy shows that serious problems arise from inferring donor numbers indirectly. Major anomalies have also been reported for phosphoramides. Bollinger et al.’s values [31] seem overestimated compared with the values given by Gutmann [21] and by Buchikhin et al. [32] for hexamethylphosphoric triamide and by Ozari and JagurGrodzinski for three N-substituted phosphoramides [33]. Differences range from +16.7 to +54.8 kJ mol−1 . The high solvent permittivity and the strong basicity of phosphoramides might favour the reaction − 2 B−SbCl5 → SbCl4 B+ 2 + SbCl6

(2.6)

The formation of these ions certainly contributes to the total heat quantity measured. Consequently, we have selected the lower values. In any case, the DN values of phosphoramides must be used with great caution. Equilibrium constants and Gibbs energies of complexation are given in Table 2.3. These data can be used to study the existence of a linear relation between ∆H ◦ and ∆G◦ for the complexation of SbCl5 . A plot of ∆H ◦ versus ∆G◦ values for 28 Lewis bases (mostly carbonyl compounds) is shown in Figure 2.2. Only 84% of the variance of ∆H ◦ is explained by ∆G◦ . The lack of a linear ∆H−∆G relationship has been attributed to pronounced


75

Table 2.2 SbCl5 affinities, SbCl5 A (kJ mol−1 ), of Lewis bases in 1,2-dichloroethane at 298 K. No. Lewis base

Formula

SbCl5 A

Ref.

Amine 1 Triethylaminea

NEt3

132.6 ± 2.5

[29]

Pyridine 2 Pyridine

C5 H5 N

142.3 ± 1.7

[30]

Nitriles 3 Chloroacetonitrileb 4 Benzonitrile 5 Acetonitrile 6 Benzyl cyanide 7 Isobutyronitrile 8 Propionitrile 9 Butyronitrile

ClCH2 C N C6 H5 C N CH3 C N C6 H5 CH2 C N (CH3 )2 CHC N CH3 CH2 C N CH3 CH2 CH2 C N

38.8 54.4 ± 0.8 61.1 ± 0.4 63.2 64.4 67.4 69.5

[30] [30] [30] [15] [15] [15] [15]

Water, alcohol 10 Water 11 Methanol (0.01–0.04 mol l−1 )

H2 O CH3 OH

101.7 ± 1.3 79.8 ± 0.1

[13] [12]

Ethers 12 1,4-Dioxane 13 Di-n-propyl ether 14 Methyl propyl ether 15 Diethyl ether 16 Tetrahydrofuran

C4 H8 O2 (CH3 CH2 CH2 )2 O CH3 OCH2 CH2 CH3 (CH3 CH2 )2 O c-(CH2 )4 O

61.9 74.6 ± 0.1 75.7 ± 0.1 80.3 88.0 ± 0.2

[19] [12] [12] [14] [11]

Carbonyl compounds 17 Acetyl chloride 18 Tetrachloroethylene carbonate 19 Benzoyl fluoride 20 Benzoyl chloride 21 Methyl trifluoroacetate 22 Ethyl trichloroacetate 23 1,2-Dichloroethylene carbonate 24 Propionyl chloride 25 Ethyl dichloroacetate 26 2,2,4,4-Tetramethylpentan-3-one 27 Acetic anhydride 28 2,2,4-Trimethylpentan-3-one 29 2,2,5,5-Tetramethylhexan-3-one 30 Ethyl chloroacetate 31 Methyl 2,2-dimethylpropanoate 32 Ethyl 2,2-dimethylpropanoate 33 Isopropyl 2,2-dimethylpropanoate 34 2,2-Dimethylpentan-3-one 35 Propylene carbonate 36 2,2,6,6-Tetramethylheptan-4-one 37 Dimethyl carbonate 38 Diethyl carbonate 39 2-Methylpentan-3-one 40 2,4-Dimethylpentan-3-one

CH3 COCl C3 Cl4 O3 C6 H5 COF C6 H5 COCl CF3 COOCH3 CCl3 COOCH2 CH3 C3 H2 Cl2 O3 CH3 CH2 COCl Cl2 CHCOOCH2 CH3 (CH3 )3 CCOC(CH3 )3 CH3 C(O)OC(O)CH3 (CH3 )3 CCOCH(CH3 )2 (CH3 )3 CCOCH2 C(CH3 )3 ClCH2 COOCH2 CH3 (CH3 )3 CCOOCH3 (CH3 )3 CCOOCH2 CH3 (CH3 )3 CCOOCH(CH3 )2 CH3 CH2 COC(CH3 )3 C4 H6 O3 (CH3 )3 CCH2 COCH2 C(CH3 )3 CH3 OCOOCH3 CH3 CH2 OCOOCH2 CH3 CH3 CH2 COCH(CH3 )2 (CH3 )2 CHCOCH(CH3 )2

2.9 3.3 9.6 9.6 11.3 ± 1.7 12.76 ± 0.29 13.4 13.8 ± 0.8 39.12 ± 0.21 40.2 ± 1.3 43.9 48.53 ± 0.13 49.29 ± 0.42 53.56 ± 0.04 53.68 ± 0.25 54.52 ± 0.17 56.5 ± 1.3 61.00 ± 0.04 62.4 ± 0.3 62.76 ± 0.42 63.47 ± 0.13 66.86 ± 0.21 67.61 ± 0.13 67.78 ± 0.13

[15] [19] [19] [14] [8] [6] [19] [7] [6] [9] [15] [9] [9] [6] [8] [6] [8] [3] [11] [9] [1] [1] [3] [3]

(Continued)

76


Table 2.2 (Continued) No.

Lewis base

Formula

SbCl5 A

Ref.

41 42 43 44 45 46 47

Methyl propionate Ethyl isobutyrate 2,6-Dimethylheptan-4-one Methyl acetate Ethylene carbonate Pentan-3-one N,N-Dimethyl-2,2,2trifluoroacetamide Ethyl butyrate Ethyl propionate 3,3-Dimethylbutan-2-one Ethyl formate Acetone 3-Methylbutan-2-one Ethyl acetate Dimethylcarbamoyl chloride Butan-2-one γ -Butyrolactone Pentan-2-one Isopropyl acetate Cyclohexanone Ethyl dimethylcarbamate N,N-Dimethylformamide 1-Methyl-2-pyrrolidinone N,N -Dimethylethyleneurea (DMEU) N,N-Dimethylacetamide N-Methylcaprolactam 1,1,3-Trimethylurea 1,1,3,3-Tetramethylurea 1,3-Dimethylurea N,N-Diethylformamide N,N-Diethylacetamide

CH3 CH2 COOCH3 (CH3 )2 CHCOOCH2 CH3 (CH3 )2 CHCH2 COCH2 CH(CH3 )2 CH3 COOCH3 C3 H4 O3 CH3 CH2 COCH2 CH3 CF3 CON(CH3 )2

67.91 ± 0.04 68.49 ± 0.08 68.53 ± 0.08 68.53 ± 0.13 68.6 69.16 ± 0.10 69.37 ± 0.08

[8] [6] [9] [1] [15] [3] [5]

CH3 CH2 CH2 COOCH2 CH3 CH3 CH2 COOCH2 CH3 CH3 COC(CH3 )3 HCOOCH2 CH3 CH3 COCH3 CH3 COCH(CH3 )2 CH3 COOCH2 CH3 ClCON(CH3 )2 CH3 COCH2 CH3 C4 H6 O2 CH3 COCH2 CH2 CH3 CH3 COOCH(CH3 )2 C6 H10 O CH3 CH2 OCON(CH3 )2 HCON(CH3 )2 C5 H9 NO C5 H10 N2 O

70.12 ± 0.21 70.37 ± 0.04 70.92 ± 0.13 71.13 ± 0.29 71.25 ± 0.17 71.46 ± 0.13 71.46 ± 0.21 72.05 ± 0.04 72.93 ± 0.13 73.18 ± 0.04 73.22 ± 0.08 73.35 ± 0.21 74.43 ± 0.04 93.6 ± 0.13 111.3 118.0 115.8 ± 0.3

[1] [6] [2] [8] [1] [2] [1] [5] [2] [8] [9] [1] [9] [5] [14] [31] [10]

CH3 CON(CH3 )2 C7 H13 ON (CH3 )2 NCONHCH3 (CH3 )2 NCON(CH3 )2 CH3 NHCONHCH3 HCON(CH2 CH3 )2 CH3 CON(CH2 CH3 )2

116.3 ± 0.3 113.4 122.6 ± 0.1 124.0 ± 0.2 124.1 ± 0.2 129.3 134.7

[4] [19] [10] [4] [10] [15] [15]

CH3 NO2 C6 H5 NO2

11.3 34 ± 3

[15] [11]

ClSO2 Cl c-(CH2 )4 SO2

0.4 61.9

[14] [15]

ClSOCl ClSeOCl C2 H4 O3 S CH3 SOCH3 CH3 (CH2 )7 SO(CH2 )7 CH3 [CH3 (CH2 )7 OC6 H5 ]2 SeO

1.7 51.0 64.0 124.7 126.9 ± 2.4 187.7 ± 0.9

[14] [14] [15] [14] [32] [32]

48 49 50 51 52 53 54 55 56 57 58 59 60 61 62 63 64 65 66 67 68 69 70 71

Nitro compounds 72 Nitromethane 73 Nitrobenzene Sulfonyl compounds 72 Sulfuryl chloride 75 Tetramethylene sulfone, sulfolane Sulfinyl, seleninyl compounds 76 Thionyl chloride 77 Selenium oxychloride 78 Ethylene sulfite 79 Dimethyl sulfoxide 80 Di-n-octyl sulfoxide 81 Dioctyloxyphenyl selenoxide

(Continued)


77

Table 2.2 (Continued) No. Lewis base

Formula

SbCl5 A

Ref.

Amine oxides 82 Pyridine N-oxide 83 4-Methylpyridine N-oxide 84 Tri(n-octyl)amine N-oxide

C5 H5 NO 4-CH3 C5 H4 NO [CH3 (CH2 )7 ]3 NO

142.4 ± 0.7 151.9 ± 1.7 218.9 ± 5.9

[32] [30] [32]

Cl3 PO C6 H5 (F)2 PO C6 H5 (Cl)2 PO (C6 H5 )2 ClPO (CH3 O)3 PO (CH3 CH2 CH2 CH2 O)3 PO (CH3 CH2 O)3 PO (CH3 CH2 O)2 PON(CH3 )2

49.0 68.6 77.4 93.7 96.2 99.2 100.4 123.6

[14] [19] [19] [14] [14] [15] [31] [31]

CH3 (OctO)2 PO Oct3 PO [(CH3 )2 N]3 PO [(CH3 )2 N]2 POOCH2 CH3

130.6 ± 0.3 141.5 ± 0.3 162.3 197.3

[32] [32] [15] [31]

[(CH3 CH2 )2 N]3 PO [(CH3 )2 N]2 PON(CH2 CH3 )2

198.5 204.8

[31] [31]

[(CH2 )4 N]2 PON(CH3 )2

206.7

[31]

[(CH3 )2 N]2 PON(CH2 )4

213.8

[31]

[(CH3 CH2 )2 N]2 PON(CH2 )4

218.1

[31]

(CH3 CH2 NH)2 PON(CH3 )2

220.1

[31]

[(CH2 )5 N]3 PO [(CH3 )2 N]2 PON(CH2 )3

222 223.0

[31] [31]

[(CH3 )2 N]2 PONHCH2 CH3

223.3

[31]

[(CH2 )4 N]3 PO [(CH2 )3 N]2 PON(CH3 )2

229.2 229.2

[31] [31]

[(CH2 )4 N]2 PON(CH2 CH3 )2

230.9

[31]

[(CH2 )3 N]3 PO (CH3 CH2 NH)3 PO [(CH3 )2 N]2 PON(CH2 )2

233.2 256.4 263.8

[31] [31] [31]

[(CH2 )2 N]2 PON(CH3 )2

266.3

[31]

[(CH2 )2 N]3 PO

383

[31]

Phosphorus compounds 85 Phosphoryl chloride 86 Difluorophenylphosphine oxide 87 Dichlorophenylphosphine oxide 88 Chlorodiphenylphosphine oxide 89 Trimethyl phosphate 90 Tri(n-butyl) phosphate 91 Triethyl phosphate 92 Diethyl N,N-dimethylphosphoramidate 93 Di-n-octyl methylphosphonate 94 Tri-n-octylphosphine oxide 95 Hexamethylphosphoric triamide 96 Ethyl tetramethylphosphorodiamidate 97 Hexaethylphosphoric triamide 98 Diethyltetramethylphosphoric triamide 99 Bis(pyrrolidino)dimethylaminophosphine oxide 100 Bis(dimethylamino)pyrrolidinophosphine oxide 101 Bis(diethylamino)pyrrolidinophosphine oxide 102 Bis(ethylamino)dimethylaminophosphine oxide 103 Tripiperidinophosphine oxide 104 Bis(dimethylamino)azetidinophosphine oxide 105 Bis(dimethylamino)ethylaminophosphine oxide 106 Tripyrrolidinophosphine oxide 107 Bis(azetidino)dimethylaminophosphine oxide 108 Bis(pyrrolidino)diethylaminophosphine oxide 109 Triazetidinophosphine oxide 110 Triethylaminophosphine oxide 111 Bis(dimethylamino)aziridinophosphine oxide 112 Bis(aziridino)dimethylaminophosphine oxide 113 Triaziridinophosphine oxide a b

In CH2 Cl2 at −55◦ C. Calculated from the relationship between ∆H values in CCl4 and in 1,2-dichloroethane. See text.

78


Table 2.3 Thermodynamic data for SbCl5 /Lewis bases equilibria in 1,2-dichloroethane at 298 K. No.

Lewis base

Kc a

21 24 22 76 85 73 25 4 29 26 5 31 75 35 32 28 30 54 52 10 15 42 49 16 89 90 65 79

CF3 COOCH3 CH3 CH2 COCl CCl3 COOCH2 CH3 ClSOCl Cl3 PO C6 H5 NO2 Cl2 CHCOOCH2 CH3 C6 H5 CN (CH3 )3 COCH2 C(CH3 )3 (CH3 )3 CCOC(CH3 )3 CH3 CN (CH3 )3 CCOOCH3 c-(CH2 )4 SO2 C4 H6 O3 (CH3 )3 CCOOCH2 CH3 (CH3 )3 CCOCH(CH3 )2 ClCH2 COOCH2 CH3 CH3 COOCH2 CH3 CH3 COCH3 H2 O (CH3 CH2 )2 O (CH3 )2 CHCOOCH2 CH3 CH3 CH2 COO CH2 CH3 c-(CH2 )4 O (CH3 O)3 PO (CH3 CH2 CH2 CH2 O)3 PO CH3 CON(CH3 )2 CH3 SOCH3

0.61 ± 0.1 14 ± 1 1.7 ± 0.2 2 5 20 ± 3 49 ± 2 139 265 ± 35 330 ± 80 714 780 ± 80 833 909 2300 ± 100 3700 ± 300 3800 ± 500 6.7 × 104 7.1 × 104 1.8 × 105 3.2 × 105 (8.8 ± 2.6).105 (3.0 ± 0.7).106 8.3 × 106 1.9 × 109 2.9 × 1010 3.2 × 1010 1.3 × 1011

a b

−∆Gob

Ref.

−1.23 6.54 1.32 1.71 4.00 7.43 9.65 12.23 13.83 14.38 16.29 16.51 16.67 16.89 19.19 20.37 20.43 27.54 27.71 30.02 31.39 33.93 36.97 39.50 52.99 59.68 59.93 63.36

[8] [7] [6] [18] [18] [11] [6] [36] [9] [9] [36] [8] [36] [36] [6] [9] [6] [36] [36] [36] [18] [6] [6] [36] [36] [36] [18] [18]

l mol−1 . ∆G◦ /kJ mol−1 = –RTlnKc .

differences in steric interaction in these complexes [34]. A more optimistic view was, however, presented by Marcus [35] (‘the linear dependence of the enthalpy and the Gibbs free energy for the reaction of donor solvents with antimony pentachloride is affirmed’), although the percentage of explained variance (87%) is about the same and the number of bases (14) is more limited in his correlation. Table 2.4 compares the influence of CCl4 , a solvent of low relative permittivity (2.30), and of 1,2-dichloroethane, of much higher relative permittivity (10.74), on the enthalpies of complexation. The SbCl5 affinities are generally higher in the more polar solvent. Complex formation is accompanied by a large increase in dipole moment. For example, PhCHO–SbCl5 has a dipole moment of 8.65 D [37] while the dipoles of SbCl5 and PhCHO are about 0 and 2.80 D, respectively. Therefore, greater stabilization of the complex is expected in the more polar ClCH2 CH2 Cl. The enthalpy difference between CCl4 and ClCH2 CH2 Cl decreases when the SbCl5 affinity decreases and falls to zero for the weakest bases. However, the ranking of Lewis base strengths remains the same in the two solvents,


79

140 120

-∆H °/ kJ mol

-1

100 80 60 40 20 0 -10

0

10

20

30

40

50

60

70

-1

-∆G ° / kJ mol

Figure 2.2 Plot of ∆H◦ versus ∆G◦ for the reaction of SbCl5 with different Lewis bases. Bases are labelled as follows: nitriles ( ), water and ethers ( ), carbonyls ( ), nitro and sulfonyl ( ), sulfinyls ( ), and phosphoryls ( ).

as shown by the following linear complexation enthalpy relationship: SbCl5 affinity(ClCH2 CH2 Cl) = 1.24 (±0.03) SbCl5 affinity(CCl4 ) −11.1 (±1.7) n = 9, r = 0.997, s = 1.1 kJ mol−1

(2.7)

Table 2.4 Influence of solvent on the SbCl5 affinities (kJ mol−1 ) of Lewis bases at 298 K.a SbCl5 affinity Lewis base (CH3 )2 NCOOCH2 CH3 c-(CH2 )4 O (CH3 CH2 CH2 )2 O CH3 COCH2 CH2 CH3 CH3 COOCH2 CH3 C4 H6 O3 c ClCH2 COOCH2 CH3 C6 H5 CNe C6 H5 NO2 a

CCl4 83.8 ± 0.2 80.8 ± 0.5 68.8 ± 0.1 67.4 ± 0.1 66.5 ± 0.1 58 ± 1d 53.6 ± 0.3 53.1 ± 0.8 36 ± 2

Ref. [11]. Difference is affinity in ClCH2 CH2 Cl minus affinity in CCl4 . Propylene carbonate. d At 40◦ C. e Ref. [30]. f A difference of 4.6 ± 2.1 is found in ref. [30]. b c

ClCH2 CH2 Cl

Differenceb

93.6 ± 0.1 88.0 ± 0.2 74.6 ± 0.2 73.2 ± 0.2 71.5 ± 0.2 62.4 ± 0.3 53.6 ± 0.3 54.4 ± 0.8 34 ± 3

9.8 ± 0.3 7.2 ± 0.7 5.8 ± 0.3 5.8 ± 0.3 5.0 ± 0.3f 4.4 ± 1.3 0.0 ± 0.6 1.3 ± 1.6 −2 ± 5

80


The enthalpies of complexation of SbCl5 with oxo bases in the solvents C6 H6 , C6 H5 NO2 , MeCOOBu and MeCOMe also correlate [32] well with the enthalpy measured in ClCH2 CH2 Cl, as shown by Equations 2.8–2.11: SbCl5 affinity in C6 H6 = (0.85 ± 0.03)DN + (1.41 ± 0.74)

(2.8)

n = 4, r = 0.998 SbCl5 affinity in C6 H5 NO2 = (0.86 ± 0.06)DN − (3.67 ± 1.91)

(2.9)

n = 5, r = 0.992 SbCl5 affinity in MeCOOBu = (0.87 ± 0.06)DN − (14.80 ± 2.08)

(2.10)

n = 8, r = 0.985 SbCl5 affinity in MeCOMe = (0.73 ± 0.05)DN − (20.65 ± 1.97)

(2.11)

n = 9, r = 0.986

2.5

Critical Discussion

The DN scale has historical merits. It was among the first attempts to establish Lewis base strength quantitatively. It was widely used in the field of solution, coordination and organic chemistry to describe the Lewis basicity of solvents. However, after 40 years of use, convergent critical analyses have pointed out the serious limitations of DN. They concern (i) the choice of SbCl5 as a reference Lewis acid, (ii) the sample of Lewis bases studied, (iii) the quality of calorimetric measurements and (iv) the domain of validity of the scale. The choice of SbCl5 appears unfortunate for technical and theoretical reasons. From the experimental point of view, SbCl5 catalyses the reaction of chlorinated solvents with amines, which prevents the easy and accurate study of this very important class of nitrogen bases. In the same way, the frequently used solvent hexamethylphosphoric triamide does not give reliable results with SbCl5 . From the perspective of the theoretical prediction of DN, the 136 electrons of SbCl5 and the position of Sb in the fourth row of the periodic table make quantum chemical calculations on a series of SbCl5 complexes difficult. Semiempirical PM3 calculations of complexation enthalpies have been performed [38] on SbCl5 complexes with 12 bases. At this low calculation level, only 43% of the variance of gasphase calculated enthalpies is explained by experimental enthalpies in CH2 ClCH2 Cl. At the ab initio or DFT level, we are aware of only two studies: the first [39] on five nitriles and pyridine and the second [40] on H2 S complexes with SbCl5 . For the pyridine–SbCl5 complex, the gas-phase enthalpy, calculated without BSSE, is −115.5 kJ mol−1 , while the experimental value in ClCH2 CH2 Cl amounts to −142.3 kJ mol−1 . It must be noted that antimony is described with the use of an effective core potential that incorporates major relativistic effects. In the end, the nature of the Sb B dative bond does not seem to have been studied in depth and no theoretical prediction of the SbCl5 affinity can easily be made for Lewis bases.


81

Concerning the sample of bases, the chemical variety is very low since 90% of them are oxygen bases. There is a serious lack of sp3 - and sp2 -hybridized nitrogen bases, and the carbon π bases, the halogen bases and the sulfur bases are totally missing. In addition, a number of the calorimetric measurements seem rather unreliable, probably because of the experimental problems caused by the high reactivity and catalytic effects of SbCl5 , or because the complexation was incorrectly assumed to be complete for weak bases. Lastly, the domain of validity of the scale seems to be more limited than was claimed by its originator. Many correlations with the DN scale, established with a restricted sample of bases, have been unduly generalized. As an example, a fair correlation (r = 0.984) was found between DN and the frequency shift, ∆ν(OD), of the OD stretching vibration of deuteromethanol, upon hydrogen bonding with 15 Lewis bases. It was concluded [41] that high donor numbers could be predicted with reasonable confidence from this correlation. We have re-examined the relationship by using a larger and more diverse sample of 61 Lewis bases with the equivalent scale ∆ν(OH) based on the shift of the OH stretching vibration of methanol (see Chapter 4). The correlation coefficient falls to 0.904 and the standard deviation of the estimate of DN becomes 16 kJ mol−1 . Clearly, the DN/∆ν(OH) correlation has no practical usefulness. For the above reasons, we discourage the use of the DN scale when precise and robust scales are available. In particular, the affinity scale constructed with the reference Lewis acid BF3 is an alternative for measuring the strength of dative bonds. This scale is presented in the following chapter.

References 1. Olofsson, G., Lindqvist, I. and Sunner, S. (1963) Adduct between antimony pentachloride and carbonyl compounds. A calorimetric study. Acta Chem. Scand., 17, 259–265. 2. Olofsson, G. (1964) Heats of formation of adducts between antimony pentachloride and some ketones and carboxylic esters. Acta Chem. Scand., 18, 11–17. 3. Olofsson, G. (1965) Heats of formation of adducts of antimony pentachloride and some ketones. Acta Chem. Scand., 19, 2155–2159. 4. Olofsson, G. (1964) Heats of formation of adducts between antimony pentachloride and N,Ndimethylacetamide and tetramethylurea. Acta Chem. Scand., 18, 1022–1023. 5. Olofsson, G. (1967) Heats of formation of adducts between antimony pentachloride and substituted N,N-dimethylamides. Acta Chem. Scand., 21, 93–97. 6. Olofsson, G. (1967) Enthalpies of formation and stability constants for adducts between antimony pentachloride and substituted ethyl acetates. Acta Chem. Scand., 21, 1892–1902. 7. Olofsson, G. (1967) Heats of reaction of antimony pentachloride with propionyl chloride and ethyl chlorocarbonate. Acta Chem. Scand., 21, 1114–1116. 8. Olofsson, G. (1967) Enthalpies of formation of adducts between antimony pentachloride and various esters. Acta Chem. Scand., 21, 2143–2150. 9. Olofsson, G. (1967) Adducts between antimony pentachloride and various ketones. Calorimetric study. Acta Chem. Scand., 21, 2415–2422. 10. Olofsson, G. (1971) Calorimetric study of the donor properties of N,N -methylureas. Acta Chem. Scand., 25, 691–694. 11. Olofsson, G. and Olofsson, I. (1973) Comparison between enthalpies of formation of antimony pentachloride adducts in carbon tetrachloride and in 1,2-dichloroethane solution. J. Am. Chem. Soc., 95, 7231–7233.

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12. Olofsson, G. (1968) Calorimetric studies of adduct formation between antimony pentachloride and methanol, methyl propyl ether, and dipropyl ether. Acta Chem. Scand., 22, 1352– 1353. 13. Olofsson, G. (1967) Enthalpy of adduct formation between antimony pentachloride and water. Acta Chem. Scand., 21, 1887–1891. 14. Gutmann, V., Steininger, A. and Wychera, E. (1966) Donor strengths in 1,2-dichloroethane. Monatsh. Chem., 97, 460–467. 15. Gutmann, V. and Scherhaufer, A. (1968) Donor strengths in 1,2-dichloroethane. IV. Monatsh. Chem., 99, 335–339. 16. Gutmann, V. and Wychera, E. (1966) Effect of the solvent on coordination reactions in nonaqueous systems. Rev. Chim. Minér., 3, 941–951. 17. Gutmann, V. and Wychera, E. (1966) Coordination reactions in nonaqueous solutions. The role of the donor strength. Inorg. Nucl. Chem. Lett., 2, 257–260. 18. Gutmann, V. (1968) Coordination Chemistry in Non-Aqueous Solutions, Springer-Verlag, New York. 19. Gutmann, V. (1976) Empirical parameters for donor and acceptor properties of solvents. Electrochim. Acta, 21, 661–670. 20. Gutmann, V. (1977) Solvent concepts. Chemtech, 7, 255–263. 21. Gutmann, V. (1978) The Donor–Acceptor Approach to Molecular Interactions, Plenum Press, New York. 22. Brun, L. and Branden, C.I. (1966) The crystal structure of SbCl5 ·HCON(CH3 )2 . Acta Crystallogr., 20, 749–758. 23. Neubauer, D. and Weiss, J. (1960) The crystal and molecular structure of SbCl5 ·S4 N4 . Z. Anorg. Allg. Chem., 303, 28–38. 24. Lindqvist, I. and Branden, C.I. (1958) The crystal structure of POCl3 ·SbCl5 . Acta Chem. Scand., 12, 134. 25. Hermodsson, Y. (1967) The crystal structure of antimony pentachloride selenium oxychloride. Acta Chem. Scand., 21, 1313–1327. 26. Allen, F.H. (2002) The Cambridge Structural Database: a quarter of a million crystal structures and rising. Acta Crystallogr., B58, 380–388. 27. Popov, A.I. (1975) Alkali metal NMR and vibrational spectroscopic studies on solvates in nonaqueous solvents. Pure Appl. Chem., 41, 275–289. 28. Hahn, S., Miller, W.M., Lichtenthaler, R.N. and Prausnitz, J.M. (1985) Donor number estimation for oxygen- and nitrogen-containing solvents via proton NMR shift of chloroform. J. Solution Chem., 14, 129–137. 29. Taft, R.W., Pienta, N.J., Kamlet, M.J. and Arnett, E.M. (1981) Linear solvation energy relationships. 7. Correlations between the solvent-donicity and acceptor-number scales and the solvatochromic parameters π ∗ , α, and β. J. Org. Chem., 46, 661–667. 30. Lim, Y.Y. and Drago, R.S. (1972) Lewis acidity of antimony pentachloride. Inorg. Chem., 11, 202–204. 31. Bollinger, J.C., Yvernault, G. and Yvernault, T. (1983) Complexation enthalpies of some phosphoramides with antimony(V) chloride. Effect of the structure on the Gutmann donor number. Thermochim. Acta, 60, 137–147. 32. Buchikhin, E.P., Chekmarev, A.M., Agafonov, Y.A. et al. (2002) Solvent effect on the efficiency of donor-acceptor interaction. Russ. J. Gen. Chem., 72, 1049–1052. 33. Ozari, Y. and Jagur-Grodzinski, J. (1974) Donor strength of N-substituted phosphoramides. J. Chem. Soc., Chem. Commun., 295–296. 34. Olofsson, G. (1968) Thermodynamic properties and frequency shifts of some molecular adducts of antimony pentachloride and molecules containing the carbonyl group. Acta Chem. Scand., 22, 377–388. 35. Marcus, Y. (1984) The effectivity of solvents as electron pair donors. J. Solution Chem., 13, 599–624. 36. Gutmann, V. and Mayer, U. (1967) Donor strengths in 1,2-dichloroethane. III. Monatsh. Chem., 98, 294–297. 37. Webster, M. (1966) Addition compounds of Group V pentahalides. Chem. Rev., 66, 87–118.


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38. Poleshchuk, O.K., Koput, J., Latosinska, J.N. and Nogaj, B. (1996) A study of electronic structures of SbCl5 ·L and SnCl4 ·L2 complexes by the PM3 method. J. Mol. Struct., 380, 267–275. 39. Klapotke, T.M., Noth, H., Schutt, T. et al. (2001) Synthesis and characterization of the Lewis acid-base complexes SbCl5 ·LB (LB = ICN, BrCN, ClCN, 1/2(CN)2 , NH2 CN, pyridine) – a combined theoretical and experimental investigation. The crystal structures of SbCl5 ·NCCl and SbCl5 ·NCCN·SbCl5 . Z. Anorg. Allg. Chem., 627, 1582–1588. 40. Poleshchuk, O.K., Dolenko, G.N., Koput, Y. and Latoshinska, I. (1997) Ab initio calculations of the coordination-induced electronic effects in donor-acceptor complexes. Russ. J. Coord. Chem., 23, 601–608. 41. Griffiths, T.R. and Pugh, D.C. (1979) Correlations among solvent polarity scales, dielectric constant and dipole moment, and a means to reliable predictions of polarity scale values from current data. Coord. Chem. Rev., 29, 129–211.

3 The BF3 Affinity Scale Since the discovery of the complex H3 N–BF3 by Gay-Lussac in 1809 [1], a great number of complexes of boron trifluoride with Lewis bases have been studied. Indeed, BF3 is an archetypical Lewis acid. It is clear that the central boron is electron deficient since it has only six electrons in its outer shell and the highly electronegative fluorine atoms further decrease its σ electron density. The thermodynamics of formation of BF3 complexes has been extensively investigated. However, the thermodynamic data collected in reviews prior to 1970 [2–4] do not allow the construction of homogeneous Lewis basicity or affinity scales because they do not refer to the same reaction and/or to the same medium. In fact, the types of reactions investigated fall into four categories: LB−BF3 (g) LB(solution) + BF3 (solution) LB1 (solution) + LB2 −BF3 (solution) LB(solution) + BF3 (g)

LB(g) + BF3 (g)

(3.1)

LB−BF3 (solution) LB1 −BF3 (solution) + LB2 (solution)

(3.2) (3.3)

LB−BF3 (solution)

(3.4)

and the medium varies from the gas phase to solutions in different solvents (benzene, dichloromethane, diethyl ether, tetrahydrofuran or nitrobenzene). In Equations 3.1–3.4, LB stands for Lewis base to avoid confusion with the symbol B of boron. At the beginning of the 1970s, the calorimetry group of the University of Nice-Sophia Antipolis determined systematically the enthalpy of complexation of gaseous BF3 with a large number of carbonyl bases in standardized conditions (reaction 3.4 in CH2 Cl2 ). They aimed to establish structure–basicity relationships in the family of carbonyl bases. In the 1980s, they further extended their work to new families of oxygen and nitrogen liquid bases in order to build a ‘solvent Lewis basicity scale for nonprotogenic solvents’ [5]. Since measurements were not carried out on the pure base but on dilute solutions of the base, and enthalpies and not Gibbs energies were determined, it appears more appropriate today to call this scale a solute scale of Lewis affinity rather than a solvent scale of Lewis basicity. Lewis Basicity and Affinity Scales: Data and Measurement C 2010 John Wiley & Sons, Ltd


86


In 1993, at the end of these calorimetric studies, the enthalpies of complexation of BF3 had been obtained for 348 Lewis bases. The construction of the BF3 affinity scale from this set of homogeneous data is presented below. For a better understanding of structure–affinity relationships, a brief summary of the structure of BF3 complexes is given first.

3.1

Structure of BF3 Complexes

The molecular structure of BF3 complexes has been determined in both the gaseous and solid phases. Solid-state structures, retrieved mainly from the CSD [6], are listed in Table 3.1, while Table 3.2 corresponds to gas-phase structures. In both tables, the complexes are arranged according to the intermolecular distance d between the boron atom and the electron donor atomic site of the base. The geometries of the BF3 complexes can be described by three structural characteristics: (i) the distortion of the BF3 fragment upon complexation, (ii) the length of the dative bond and (iii) the directionality (spatial orientation) of the dative bond. Upon complexation, the geometry of the Lewis base is not much altered. In contrast, the geometry of the BF3 moiety is significantly distorted through lengthening of the B F bonds and a reorganization from a trigonal planar to a pyramidal geometry. The extent of the latter change can be measured by the decrease in the FBF angle from the trigonal value of 120◦ to a smaller value. The energy necessary to distort the BF3 fragment has been calculated by various theoretical methods [22–25] and constitutes a significant part of the binding electronic energy. The length of the dative bond is intermediate between the sum of the van der Waals and covalent radii of boron and the atomic donor site, and approaches the covalent bond distance for stronger complexes. For example, the shortest dative-bond length in Tables 3.1 and 3.2, found in the ˚ is very close to the sum of covalent radii of boron and Ph3 AsO–BF3 complex (1.486 A), ˚ oxygen (1.477 A). The spatial orientation of the BF3 fragment with respect to the Lewis base moiety appears to be governed mainly by the direction of a putative lone pair of the Lewis atomic site. Thus, in the complexes of BF3 with nitriles, pyridines and amines, the B N vector points in the direction of the sp, sp2 and sp3 nitrogen lone pair, respectively. In addition, in the complexes of BF3 with carbonyl bases the B O vector points towards one of the two sp2 -like orbitals of oxygen [26, 27]. The X-ray structure of a representative complex is shown in Figure 3.1. In this ˚ than the pyridine–BF3 complex [28], the B N bond has a much shorter length (1.602 A)

Figure 3.1

˚ Structure of the complex of BF3 with pyridine (distances in A).

The BF3 Affinity Scale

87

Table 3.1 Solid-state structures of BF3 complexes: references and intermolecular distances ˚ d (A). Lewis base

Ref.a

Triethylphosphine 8-(2,6-Dimethylpiperidino)-2-methyl-1,9diazaborabicyclo(4.3.0)nonane 1,3-Dimethyl-1,3-pyrazolidine Iodoacetonitrile Chloroacetonitrile Fluoroacetonitrile Acetonitrile Acetonitrile Benzonitrile 5-Benzyl-2-methyl-1,2,4,3-triazaphosphole Pyridine Quinuclidine Ammoniab 1-Methyl-5-benzyl-1,2,4,3-triazaphosphole Benzylamine 4-Dimethylaminopyridine 2-Methylacrolein Trimethylamine 4,5-Diisopropyl-1,2,4,3-triazaphosphole 1-[bis(Diisopropylamino)boryl]imidazole Triethylphosphino(trimethylsilyl)imine Methylamine 2,3-Methylenedioxybenzaldehyde Benzylideneacetone Methyl cinnamate Ammonia-d 3 Imidazole Waterc Dibenzylideneacetone Diphenyl-(2-hydroxy-3,5-di-tert-butylphenyl) phosphinate Methanol 5,5-Pentamethylene-2-phenyl-4,5dihydrooxazole N-oxide Triphenylphosphine oxide Waterd Waterd Watere Triphenylphosphoranylimine 3-Diethylamino-3-phenylacrolein 2,5-Bis(trimethylsilyl)thiophene oxide Triphenylarsine oxide

HUVROF BANSIS

2.028 (B P) 1.655 (B Nsp3 )

OCIBAD [7] [7] [7] MCYNBF MCYNBF01 [8] KAXKUP CEZJIA LATZIP CEDZAM VUTTEJ DEMCON JUDMIE VUSRAC TMAMBF 10 SUBWUH IKAPIT TOPWAW [9, 10] RUGYEX MOZCEJ MOZCAF [11] TARPOR SIXFOU MOZCIN ZINPIV

1.651, 1.650 (B 1.643 (B Nsp) 1.649, 1.640 (B 1.635 (B Nsp) 1.633 (B Nsp) 1.630 (B Nsp) 1.630 (B Nsp) 1.605, 1.608 (B 1.604, 1.602 (B 1.601 (B Nsp3 ) 1.598 (B Nsp3 ) 1.597 (B Nsp2 ) 1.595 (B Nsp3 ) 1.589(B Nsp2 ) 1.587 (B O) 1.586 (B Nsp3 ) 1.585 (B Nsp2 ) 1.580, 1.575 (B 1.578 (B Nsp2 ) 1.58 (B Nsp3 ) 1.567 (B O) 1.558 (B O) 1.556 (B O) 1.554 (B Nsp3 ) 1.545 (B Nsp2 ) 1.534 (B O) 1.533 (B O) 1.528 (B O)

BEDVUB VOLNEP

1.524 (B O) 1.517 (B O)

JAWFUI LEYKYIJ LEYKYIJ 01 NIYGAD YODGIH JADBEW GIYBEV JAWGAP

1.516 (B O) 1.509 (B O) 1.509 (B O) 1.508 (B O) 1.505 (B Nsp2 ) 1.495, 1.496 (B O) 1.494 (B O) 1.486 (B O)

a

Reference code in the CSD database [6] or bibliographic reference. Amminetrifluoroboron 18-crown-6 dichloromethane solvate. c 18-Crown-6 aqua-trifluoroborate clathrate toluene solvate. d 18-Crown-6 bis(aqua-trifluoroboron) dihydrate. e Dicyclohexano-18-crown-6 bis(monoaquatrifluoroboron). b

d (boron–Lewis base)

Nsp3 ) Nsp)

Nsp2 ) Nsp2 )

Nsp2 )

88

Lewis Basicity and Affinity Scales Table 3.2 Gas-phase structures of BF3 complexes: references ˚ and intermolecular distances d (A). Lewis base

Ref.

d (boron–Lewis base)

Carbon monoxide Dinitrogen Cyanogen Hydrogen fluoride Hydrogen cyanide Acetonitrile Phosphine Dimethyl ether Ammonia Pyridine Trimethylamine

[12] [12] [13] [14] [15] [16] [17] [18] [19] [20] [21]

2.886 (B C) 2.875 (B Nsp) 2.647 (B Nsp) 2.544 (B F) 2.473 (B Nsp) 2.011 (B Nsp) 1.921 (B P) 1.75 (B O) 1.673 (B Nsp3 ) 1.669 (B Nsp2 ) 1.636 (B Nsp3 )

˚ but slightly longer than the sum of the covalent radii sum of the van der Waals radii (3.0 A) ˚ of the boron and nitrogen atoms. The mean of the BNC angles (120.7◦ ) indicates (1.517 A) that the B N vector points in the direction of the putative sp2 lone pair of nitrogen. The FBF angles (110.5◦ , 111.8◦ , 111.4◦ ) become close to tetrahedral. The B F distances (1.330, ˚ in the solid state [29]). ˚ are longer than in the free BF3 molecule (1.287 A 1.347, 1.357 A) A comparison of the solid- and gas-phase structures of the same complex shows that ˚ is the dative bond is usually shorter in the solid phase. The greatest shortening (0.84 A) observed in the HCN–BF3 complex [30]. This shortening has been studied theoretically [8, 31–33]. One of the suggested mechanisms of shortening is the interaction of the dipole moment induced upon complexation [34] with the crystal dipolar field. This large effect of the medium on the dative-bond length must have consequences on the strength of the dative bond and, specifically, on the BF3 affinity values. It is expected that higher affinities will be found on going from the gas phase to more polar media. In the context of the BF3 affinity scale, a further interesting structural feature is the importance of steric effects in determining the structure of BF3 complexes. For example, the complexes of benzaldehydes with BF3 [27, 35] have the phenyl group anti to the BF3 . In other words, the B O bond points towards the less hindered sp2 lone pair of the carbonyl oxygen. In three BF3 –ketimine complexes [36], the BF3 group orients towards the less sterically demanding side of the imine function. The steric strain between BF3 and Lewis bases has been explained by orbital interaction theory [37].

3.2

Definition of the BF3 Affinity Scale

The formation of the 1 : 1 complex between BF3 and a Lewis base, LB, can be written as F

F +

B F

F

LB

LB

B

F F

(3.5)


89

By far the most important body of data on the thermodynamics of this reaction has been provided by the group of Gal and Maria at the University of Nice-Sophia Antipolis. For technical reasons, these authors chose to study reaction 3.5 between gaseous BF3 and the Lewis base in a dilute solution of dichloromethane at 298 K. The affinity scale can be defined as the negative enthalpy of reaction 3.6: BF3 (g) + LB(CH2 Cl2 solution) LB−BF3 (CH2 Cl2 solution) BF3 A = −∆H ◦ (3.6)

(3.6) (3.7)

Since the enthalpy of solution of gaseous BF3 in CH2 Cl2 is −10.0 ± 3.0 kJ mol−1 [38], the enthalpy of reaction 3.8: BF3 (CH2 Cl2 solution) + LB(CH2 Cl2 solution) LB−BF3 (CH2 Cl2 solution)

(3.8)

is easily deduced by adding 10 kJ mol−1 to ∆H ◦ (3.6).

3.3

Experimental Determination of the BF3 Affinity Scale

A complete account of the calorimetric method used at the University of Nice-Sophia Antipolis is given in ref. [39] and a detailed example is provided in Section 7.1. Therefore, the description of the method given here is limited to the essential features. Heat evolved in the complexation reaction 3.6 is measured in a heat-flux microcalorimeter thermoregulated at 298 K. The measuring cell contains ∼3 cm3 of a dilute solution of base in CH2 Cl2 . The base concentration depends on its strength: it usually ranges from 0.2 mol l−1 for strong bases to 1 mol l−1 for weak bases. Aliquots in the range (1–3) × 10-4 mol of gaseous BF3 are added to the solution of base by means of a vacuum line. Each addition of a quantity of BF3 (n mol) generates a quantity of heat, Q. When the reaction is complete, the enthalpy of complexation for each addition, ∆H ◦ , is defined as the Q/n ratio. The method is equivalent to a discontinuous titration of the base by the acid BF3 . One titration provides 6–8 ∆H ◦ values. Their mean and the corresponding confidence limits, usually at the 95% level, are calculated. The precision is fairly good (0.2–0.5% within a set, 0.5–1% between sets) and the accuracy is estimated as 1–2%. The stoichiometry of the BF3 complex is obtained by observing the saturation of the solution contained in the calorimeter. The number of moles of BF3 coordinated to 1 mol of base can also be measured by observation of a break in the slope of the curve Q (total heat produced) versus n (number of moles of added BF3 ). The titration curve in Figure 3.2a corresponds to nearly quantitative complex formation, resulting from a large equilibrium constant K for reaction 3.6. When K becomes smaller, the complex may dissociate significantly at the working concentration and a continuous curve, rather than straight-line parts, is obtained (Figure 3.2b). From the part of the curve where the concentration of boron trifluoride in solution remains low (at higher concentrations, gaseous BF3 is released) the curve can be modelled so as to determine K and ∆H ◦ [38] simultaneously.

Q


Q

90

(a)

nc

n

(b)

n

Figure 3.2 Q = f(n) curves: (a) high complexation constant; (b) low complexation constant. nc is the quantity of added BF3 for complete complexation.

3.4

The BF3 Affinity Scale: Data

The values of 348 BF3 affinities are collected in Table 3.3. Such a database corresponds to the most extended set of Lewis affinities ever measured against a covalent metal halide as reference Lewis acid. The sample of Lewis bases contains mainly oxygen bases, of which 215 are carbonyl compounds, and nitrogen bases. There is a very limited number of secondrow bases. The scale lacks π carbon bases and halogen bases and, more generally, any weak Lewis base. Evidently, this is due to the low formation constants of the BF3 complexes of such weak bases [40], which prevent a straightforward utilization of the calorimetric method described in Section 3.3. An IR spectroscopic van’t Hoff method yields BF3 affinities of 9.3, 10.0, 11.8, 13.2 and 16.8 kJ mol−1 for ethyne [41], ethene [42], propene [42], propyne [41] and fluoromethane [43], respectively. These values, determined in liquid argon, cannot be mixed with those measured in CH2 Cl2 , but are indicative of the weak BF3 affinities of alkenes, alkynes and fluoroalkanes. The weak enthalpies of solution of BF3 in benzene [44], toluene [44] and dichloromethane [38] (−2.9, −3.3 and −10.0 kJ mol−1 , respectively) also testify the weak BF3 affinities of aromatics and chloroalkanes. Most BF3 affinities in Table 3.3 are primary values measured in dichloromethane. This solvent was chosen because it can dissolve most BF3 complexes, while being rather ‘inert’ (low hydrogen-bond acidity, Lewis basicity and relative permittivity). Nitrobenzene also exhibits good solvent properties towards BF3 complexes, and some measurements were run on representative bases, in both C6 H5 NO2 and CH2 Cl2 . A very good linear relationship was observed [5] between the enthalpy changes of reaction 3.4 in dichloromethane and nitrobenzene: BF3 A = −∆H ◦ (3.4, CH2 Cl2 solution) = 0.958 [−∆H ◦ (3.4, C6 H5 NO2 solution)] −0.31 (3.9) n = 12, r = 0.9970, s = 2.26 kJ mol–1 This relation, established for bases 41, 49, 50, 75, 79, 82, 84, 116, 287, 305, 311 and 339 in Table 3.3, covers a wide range of reactivity, about 100 kJ mol−1 , from dimethyl sulfite to 4-dimethylaminopyridine: hence it can be safely used for calculating secondary BF3 affinities, as described below. In order to minimize the dissociation of weakly stable complexes, some measurements were made on pure bases (bulk liquids). They correspond to reaction: LB(pure) + BF3 (g) LB−BF3 (pure base)

(3.10)


91

Table 3.3 BF3 affinities (kJ mol−1 ) of Lewis bases in dichloromethane at 298 K and 1 atm. No.

Lewis base

Formula

BF3 affinity

Ref.

4-NO2 C6 H4 NMe2 4-N CC6 H4 NMe2 4-COOEtC6 H4 NMe2 C6 H5 NMe2 4-MeC6 H4 NMe2 NEt3 c-HexNMe2 c-(CH2 )4 NMe NMe3

78.13 ± 0.53 87.12 ± 0.26 92.32 ± 0.32 109.16 ± 0.76 111.17 ± 0.52 135.87 ± 1.09b 137.85 ± 0.59 139.51 ± 0.77 139.53 ± 1.79 142.76 ± 1.08b

[45] [45] [46] [47] [45] [47] [46] [5] [46] [46]

a

150.01 ± 3.48 153.44 ± 0.91b

[46] [46]

80.10 ± 0.36 82.46 ± 0.39 95.37 ± 0.28 96.20 ± 0.29 96.28 ± 0.48 96.76 ± 0.29 97.73 ± 0.58 97.93 ± 0.52 101.03 ± 0.29 103.35 ± 0.41 109.73 ± 0.53 110.36 ± 0.26 113.02 ± 0.21 113.27 ± 0.31 115.55 ± 0.31 115.75 ± 0.36 116.44 ± 0.31 118.08 ± 0.59 118.67 ± 0.42 118.79 ± 0.39 119.62 ± 0.79 119.63 ± 0.20 119.71 ± 0.76 120.10 ± 0.57 120.98 ± 0.69 121.29 ± 0.52 123.44 ± 0.47 126.60 ± 0.61 127.67 ± 0.48 128.08 ± 0.50 129.50 ± 0.49 130.93 ± 0.39 133.93 ± 0.49

[48] [48] [48] [48] [48] [48] [47] [48] [47] [48] [48] [48] [48] [48] [48] [48] [48] [48] [48] [48] [48] [48] [48] [48] [48] [48] [5] [48] [48] [47] [48] [5] [48]

11

Amines N,N-Dimethyl-4-nitroaniline N,N-Dimethyl-4-cyanoaniline N,N-Dimethyl-4-carbethoxyaniline N,N-Dimethylaniline N,N-Dimethyl-4-toluidine Triethylamine N,N-Dimethylcyclohexylamine 1-Methylpyrrolidine Trimethylamine 1,4-Diazabicyclo[2.2.2]octane (DABCO) Quinuclidine

12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44

Six-membered aromatic N-heterocycles 2-tert-Butylpyridine 2-t-BuC5 H4 N 2-Trifluoromethylpyridine 2-CF3 C5 H4 N 2-Bromopyridine 2-BrC5 H4 N 2-Cyanopyridine 2-N CC5 H4 N 2-Chloropyridine 2-ClC5 H4 N 2-Fluoropyridine 2-FC5 H4 N 2,6-Dimethylpyridine 2,6-Me2 C5 H3 N 2-Dimethylaminopyridine 2-Me2 NC5 H4 N 2,4,6-Trimethylpyridine 2,4,6-Me3 C5 H2 N 2-Phenylpyridine 2-PhC5 H4 N 2-Methoxypyridine 2-MeOC5 H4 N 3-Cyanopyridine 3-N CC5 H4 N Pyrimidine 1,3-N2 C4 H4 4-Cyanopyridine 4-N CC5 H4 N 2-Vinylpyridine 2-CH2 CHC5 H4 N 4-Trifluoromethylpyridine 4-CF3 C5 H4 N 3-Trifluoromethylpyridine 3-CF3 C5 H4 N 3-Bromopyridine 3-BrC5 H4 N 2-Isopropylpyridine 2-i-PrC5 H4 N 3-Chloropyridine 3-ClC5 H4 N 2-n-Propylpyridine 2-n-PrC5 H4 N 3-Acetylpyridine 3-CH3 COC5 H4 N 3-Carbomethoxypyridine 3-COOMeC5 H4 N 2-Ethylpyridine 2-EtC5 H4 N 4-Carbomethoxypyridine 4-COOMeC5 H4 N 4-Acetylpyridine 4-CH3 COC5 H4 N 2-Methylpyridine 2-MeC5 H4 N 3-Phenylpyridine 3-PhC5 H4 N 3-Methoxypyridine 3-MeOC5 H4 N Pyridine C5 H5 N 4-Phenylpyridine 4-PhC5 H4 N 3-Methylpyridine 3-MeC5 H4 N 4-tert-Butylpyridine 4-t-BuC5 H4 N

1 2 3 4 5 6 7 8 9 10

a

(Continued)

92



Lewis base

Formula

BF3 affinity

Ref.

45 46 47 48 49

4-Methylpyridine 4-Methoxypyridine 3-Dimethylaminopyridine 4-Morpholinopyridine 4-Dimethylaminopyridine

4-MeC5 H4 N 4-MeOC5 H4 N 3-Me2 NC5 H4 N 4-(OC4 H8 N)C5 H4 N 4-Me2 NC5 H4 N

134.17 ± 0.59 135.27 ± 0.45 136.68 ± 0.15 144.47 ± 0.52 151.55 ± 0.76

[47] [48] [48] [48] [48]

50

Five-membered aromatic N-heterocycle a Thiazole

118.35 ± 0.49

[5]

Amidine 1,8-Diazabicyclo[5.4.0]undec7-ene (DBU)

a

159.36 ± 0.86

[46]

ClCH2 C N MeOCH2 C N MeSCH2 C N ClCH2 CH2 C N 2-MeC6 H4 C N MeSC N C6 H5 C N PhCH2 C N 4-MeC6 H4 C N 4-MeOC6 H4 C N i-PrC N MeC N 1-AdamC N n-BuC N t-BuC N EtC N c-HexC N n-PrC N c-PrC N Me2 NC N c-(CH2 )5 NC N

44 ± 2 52.07 ± 0.16 52.68 ± 0.21 53.09 ± 0.59 54.38 ± 0.42 55.30 ± 0.42 55.44 ± 0.28 56.61 ± 0.23 58.05 ± 0.74 59.11 ± 0.44 60.07 ± 0.31 60.39 ± 0.46 60.50 ± 0.43 60.75 ± 0.24 60.92 ± 0.13 60.95 ± 0.21 61.16 ± 0.33 61.18 ± 0.28 63.17 ± 0.38 77.23 ± 0.61 78.15 ± 0.19

[49] [49] [49] [49] [46] [49] [49] [49] [46] [46] [49] [49] [49] [49] [49] [49] [49] [49] [49] [49] [46]

EtOCH2 CH2 C N

64.79 ± 0.23 68.63 ± 0.43 74.09 ± 0.27 74.67 ± 0.19 76.61 ± 0.39 78.57 ± 0.39 78.77 ± 0.38 79.42 ± 0.27 83.55 ± 0.20 85.36 ± 0.46 87.78 ± 0.38 90.40 ± 0.28 92.25 ± 0.23 92.83 ± 0.30

[49] [46] [47] [46] [5] [5] [47] [5] [46] [5] [5] [47] [46] [46]

51

52 53 54 55 56 57 58 59 60 61 62 63 64 65 66 67 68 69 70 71 72

Nitriles Chloroacetonitrile Methoxyacetonitrile (Methylthio)acetonitrile 3-Chloropropionitrile 2-Tolunitrile Methyl thiocyanate Benzonitrile Benzyl cyanide 4-Tolunitrile 4-Methoxybenzonitrile Isobutyronitrile Acetonitrile 1-Adamantanecarbonitrile Valeronitrile tert-Butyl cyanide Propionitrile Cyclohexyl cyanide n-Butyronitrile Cyclopropyl cyanide Dimethylcyanamide 1-Piperidinecarbonitrile

73 74 75 76 77 78 79 80 81 82 83 84 85 86

Ethers 3-Ethoxypropionitrile 1,3-Dioxolane 1,4-Dioxane 1,2-Dimethoxyethane Diisopropyl ether Di-n-butyl ether Diethyl ether Di-n-propyl ether Dimethyl ether Tetrahydropyran Oxepane Tetrahydrofuran 3-Methyltetrahydrofuran 2-Methyltetrahydrofuran

a a

MeOCH2 CH2 OMe i-Pr2 O n-Bu2 O Et2 O n-Pr2 O Me2 O c-(CH2 ) 5 O a

c-(CH2 )4 O a a

(Continued)


93

Formula

BF3 affinity

Ref.


Lewis base Carbonyl compounds

87 88 89 90 91 92 93 94 95

Aldehydes 4-Nitrobenzaldehyde 4-Cyanobenzaldehyde Acetaldehyde 4-Chlorobenzaldehyde Benzaldehyde 4-Fluorobenzaldehyde 4-Methybenzaldehyde 4-Methoxybenzaldehyde 4-Dimethylaminobenzaldehyde

4-NO2 C6 H4 COH 4-N CC6 H4 COH MeCOH 4-ClC6 H4 COH PhCOH 4-FC6 H4 COH 4-MeC6 H4 COH 4-MeOC6 H4 COH 4-Me2 NC6 H4 COH

62.32 ± 0.18 65.06 ± 0.44 69.57 ± 1.23 73.51 ± 0.54 74.88 ± 1.00 76.14 ± 0.36 79.62 ± 0.56 84.81 ± 0.42 103.21 ± 0.38

[50] [50] [46] [50] [5] [50] [50] [50] [50]

96 97 98 99 100 101 102 103 104 105 106 107 108 109 110 111 112 113 114 115 116 117 118 119 120

Aliphatic ketones Di-tert-butyl ketone Chloroacetone Phenoxyacetone Methoxyacetone Diisopropyl ketone Benzyl methyl ketone Di-n-butyl ketone Diethyl ketone tert-Butyl methyl ketone Di-n-propyl ketone Neopentyl methyl ketone sec-Butyl methyl ketone n-Nonyl methyl ketone n-Hexyl methyl ketone sec-Hexyl methyl ketone n-Butyl methyl ketone Isopentyl methyl ketone Isopropyl methyl ketone Isobutyl methyl ketone 1-Adamantyl methyl ketone Acetone Ethyl methyl ketone n-Propyl methyl ketone Cyclopropyl methyl ketone Dicyclopropyl ketone

t-BuCO-t-Bu MeCOCH2 Cl PhOCH2 COMe MeOCH2 COMe i-PrCO-i-Pr PhCH2 COMe n-BuCO-n-Bu EtCOEt t-BuCOMe n-PrCO-n-Pr neo-PenCOMe sec-BuCOMe n-NonCOMe n-HexCOMe sec-HexCOMe n-BuCOMe i-PenCOMe i-PrCOMe i-BuCOMe 1-AdamCOMe MeCOMe MeCOEt n-PrCOMe c-PrCOMe c-PrCOc-Pr

31.32 ± 0.41c 57.49 ± 0.51d 60.24 ± 0.80d 67.46 ± 0.84d 68.07 ± 0.74 68.65 ± 0.45d 70.70 ± 0.57 72.28 ± 0.23 72.83 ± 0.37 73.28 ± 0.49 73.32 ± 0.42d 73.73 ± 0.50d 74.15 ± 0.54 74.30 ± 0.45 74.47 ± 0.42 74.60 ± 0.46d 74.84 ± 0.63d 74.84 ± 0.25 75.09 ± 0.34d 75.59 ± 0.21 76.03 ± 0.21 76.07 ± 0.33 76.19 ± 0.37 83.20 ± 0.53d 87.48 ± 0.38

[5] [51] [51] [51] [5] [51] [5] [5] [5] [5] [51] [51] [46] [46] [46] [51] [51] [5] [51] [5] [47] [5] [5] [51] [5]

121 122 123 124 125 126 127 128

Cycloalkanones 2,2,6-Trimethylcyclohexanone Cyclobutanone 2,5-Dimethylcyclopentanone 2,2-Dimethylcyclohexanone 2,2,4-Trimethylcyclopentanone 2,2-Dimethylcyclopentanone 2-Methylcyclohexanone Cyclohexanone

C9 H16 O c-(CH2 )3 CO C7 H12 O C8 H14 O C8 H14 O C7 H12 O C7 H12 O c-(CH2 )5 CO

55.95 ± 2.03 70.30 ± 0.16 70.64 ± 1.21 73.93 ± 0.70 75.09 ± 0.98 75.25 ± 1.24 76.03 ± 0.78 76.36 ± 0.82

[46] [5] [46] [52] [52] [46] [52] [47] (Continued)

94



Lewis base

Formula

BF3 affinity

Ref.

129 130 131 132 133 134 135 136 137 138

3,3,5-Trimethylcyclohexanone Norcamphor 2-Methylcyclopentanone 3-Methylcyclohexanone Cyclopentanone 2,4,4-Trimethylcyclopentanone 4-Methylcyclohexanone Cycloheptanone 2,3-Dimethylcyclopentanone Bicyclo[3.2.1]octan-2-one

C9 H16 O

76.97 ± 0.37 77.22 ± 0.33 77.22 ± 0.86 77.30 ± 0.45 77.44 ± 0.45 77.50 ± 0.29 77.51 ± 1.27 77.57 ± 0.33 77.60 ± 0.95 77.71 ± 0.44

[52] [46] [52] [52] [5] [52] [52] [5] [46] [46]

139 140 141 142 143 144 145 146 147 148 149 150 151 152

4-Substituted camphors 4-Nitrocamphor 4-Cyanocamphor 4-Bromocamphor 4-Chlorocamphor 4-Iodocamphor 4-Acetoxycamphor 4-Carbomethoxycamphor 4-Ethynylcamphor 4-Methoxycamphor 4-Phenylcamphor 4-Vinylcamphor Camphor 4-Methylcamphor 4-Ethylcamphor

63.65 ± 0.22 64.99 ± 0.21 67.73 ± 0.13 68.20 ± 0.29 68.72 ± 0.25 71.17 ± 0.28 71.71 ± 0.25 71.96 ± 0.32 72.55 ± 0.25 74.59 ± 0.20 75.37 ± 0.34 77.23 ± 0.20 77.48 ± 0.19 78.19 ± 0.20

[53] [46] [53] [53] [46] [53] [53] [53] [46] [46] [46] [53] [53] [53]

153 154 155 156 157 158 159 160 161 162 163 164 165 166 167 168 169 170 171 172 173 174

Substituted acetophenones, propiophenones, butyrophenones 3-Nitrobutyrophenone 3-NO2 C6 H4 COPr 3-Nitroacetophenone 3-NO2 C6 H4 COMe 4-Nitroacetophenone 4-NO2 C6 H4 COMe 4-Cyanoacetophenone 4-C NC6 H4 COMe 2 -Phenylacetophenone C6 H5 COCH2 C6 H5 4-Bromopropiophenone 4-BrC6 H4 COEt 4-Chloropropiophenone 4-ClC6 H4 COEt 3-Carbomethoxyacetophenone 3-COOMeC6 H4 COMe Propiophenone C6 H5 COEt 4-Chlorobutyrophenone 4-ClC6 H4 COPr 4-Acetylacetophenone 4-MeCOC6 H4 COMe Butyrophenone C6 H5 COPr 3-Methoxyacetophenone 3-MeOC6 H4 COMe 4-Chloroacetophenone 4-ClC6 H4 COMe 4-Bromoacetophenone 4-BrC6 H4 COMe Acetophenone C6 H5 COMe 4-Methylbutyrophenone 4-MeC6 H4 COPr 3-Methylacetophenone 3-MeC6 H4 COMe 4-Phenylacetophenone 4-PhC6 H4 COMe 4-Methylacetophenone 4-MeC6 H4 COMe 4-Methoxypropiophenone 4-MeOC6 H4 COEt 4-Methoxybutyrophenone 4-MeOC6 H4 COPr

62.78 ± 0.36 65.10 ± 0.90 67.07 ± 0.50 67.08 ± 0.25 67.1 ± 1.2 68.29 ± 0.54 68.74 ± 0.52 68.80 ± 0.48 70.31 ± 0.62 70.53 ± 0.31 70.58 ± 1.18 72.56 ± 0.35 72.68 ± 0.45 73.03 ± 0.31 73.09 ± 0.40 74.52 ± 0.15 75.02 ± 0.33 76.35 ± 0.47 77.03 ± 0.18 77.82 ± 0.33 78.17 ± 0.42 80.16 ± 0.44

[46] [46] [50] [50] [54] [46] [46] [46] [46] [46] [50] [46] [46] [50] [50] [5] [46] [46] [55] [50] [46] [46]

a

C6 H10 O C7 H12 O c-(CH2 )4 CO C8 H14 O C7 H12 O c-(CH2 )6 CO C7 H12 O a

a a a a a a a a a a a a a a

(Continued)


95


Lewis base

Formula

BF3 affinity

Ref.

175 176

4-Methoxyacetophenone 4-Dimethylaminoacetophenone

4-MeOC6 H4 COMe 4-Me2 NC6 H4 COMe

83.01 ± 0.43 98.84 ± 0.23

[50] [50]

C10 H16 O

81.04 ± 0.92

[56]

C9 H14 O

C7 H10 O C8 H12 O C10 H16 O

81.46 ± 0.79 81.71 ± 0.84 82.42 ± 0.38 84.10 ± 1.09 84.94 ± 1.38

[56] [52] [52] [52] [56]

C7 H10 O C8 H12 O C9 H14 O C7 H10 O C8 H12 O

84.94 ± 0.96 85.35 ± 0.33 87.45 ± 0.42 89.54 ± 0.17 90.37 ± 0.92

[56] [56] [56] [52] [52]

64.46 ± 0.53

[55]

72.55 ± 0.57

[55]

77.28 ± 0.18

[55]

77.87 ± 0.33

[55]

79.71 ± 0.26

[55]

80.29 ± 0.48

[55]

82.20 ± 0.55

[55]

83.50 ± 0.49

[55]

83.58 ± 0.34 83.67 ± 0.24

[55] [55]

84.87 ± 0.34

[55]

85.11 ± 0.35

[55]

85.41 ± 0.43

[55]

85.44 ± 0.69

[55]

86.06 ± 0.26

[55]

177 178 179 180 181 182 183 184 185 186 187 188 189 190 191 192 193 194 195 196 197 198 199 200 201 202

Cyclohexenones 4,4,6,6-Tetramethyl-2-cyclohexen1-one 4,4,6-Trimethyl-2-cyclohexen-1-one 2-Cyclohexen-1-one 4-Methyl-2-cyclohexen-1-one 4,4-Dimethyl-2-cyclohexen-1-one 2-Methyl-5-isopropyl-2-cyclohexen1-one 2-Methyl-2-cyclohexen-1-one 2,5-Dimethyl-2-cyclohexen-1-one 2,4,4-Trimethyl-2-cyclohexen-1-one 3-Methyl-2-cyclohexen-1-one 3,5-Dimethyl-2-cyclohexen-1-one

a

3-Substituted 5,5-dimethyl-2-cyclohexen-1-ones a 3-Cyano-5,5-dimethyl-2cyclohexen-1-one a 3-Carbomethoxy-5,5-dimethyl-2cyclohexen-1-one a 3-Bromo-5,5-dimethyl-2cyclohexen-1-one a 3-Chloro-5,5-dimethyl-2cyclohexen-1-one a 3-Cyanomethyl-5,5-dimethyl-2cyclohexen-1-one a 3-(4-Nitrophenyl)-5,5-dimethyl-2cyclohexen-1-one a 3-(4-Cyanophenyl)-5,5-dimethyl-2cyclohexen-1-one a 3-Carbethoxymethyl-5,5-dimethyl2-cyclohexen-1-one a 5,5-Dimethyl-2-cyclohexen-1-one a 3-Benzoyloxy-5,5-dimethyl-2cyclohexen-1-one a 3-(4-Bromophenyl)-5,5-dimethyl-2cyclohexen-1-one a 3-Acetoxy-5,5-dimethyl-2cyclohexen-1-one a 3-(4-Chlorophenyl)-5,5-dimethyl-2cyclohexen-1-one a 3-(2-Oxobutyl)-5,5-dimethyl-2cyclohexen-1-one a 3-(4-Fluorophenyl)-5,5-dimethyl-2cyclohexen-1-one

(Continued)

96



Lewis base

Formula

BF3 affinity

Ref.

203

3-Phenyl-5,5-dimethyl-2cyclohexen-1-one 3-(4-Phenylphenyl)-5,5-dimethyl-2cyclohexen-1-one 3-(4-Methylphenyl)-5,5-dimethyl-2cyclohexen-1-one 3-(4-Methoxyphenyl)-5,5-dimethyl2-cyclohexen-1-one 3,5,5-Trimethyl-2-cyclohexen-1-one (isophorone) 3-Ethylthio-5,5-dimethyl-2cyclohexen-1-one 3-Methoxy-5,5-dimethyl-2cyclohexen-1-one 3-Ethoxy-5,5-dimethyl-2cyclohexen-1-one 3-Diphenylamino-5,5-dimethyl-2cyclohexen-1-one 3-Dibenzylamino-5,5-dimethyl-2cyclohexen-1-one 3-Methylphenylamino-5,5dimethyl-2-cyclohexen-1-one 3-Morpholino-5,5-dimethyl-2cyclohexen-1-one 3-(3,5-Dimethylmorpholino)-5,5dimethyl-2-cyclohexen-1-one 3-Methybenzylamino-5,5-dimethyl2-cyclohexen-1-one 3-Thiomorpholino-5,5-dimethyl-2cyclohexen-1-one 3-(N-Methylpiperazino)-5,5dimethyl-2-cyclohexen-1-one 3-Ethybenzylamino-5,5-dimethyl-2cyclohexen-1-one 3-Dimethylamino-5,5-dimethyl-2cyclohexen-1-one 3-(3,5-Dimethylpiperidino)-5,5dimethyl-2-cyclohexen-1-one 3-Heptamethyleneimino-5,5dimethyl-2-cyclohexen-1-one 3-Piperidino-5,5-dimethyl-2cyclohexen-1-one 3-Pyrrolidino-5,5-dimethyl-2cyclohexen-1-one 3-Diethylamino-5,5-dimethyl-2cyclohexen-1-one 3-Azetidino-5,5-dimethyl-2cyclohexen-1-one

a

86.48 ± 0.29

[55]

a

87.16 ± 0.33

[55]

a

88.79 ± 0.28

[55]

a

90.44 ± 0.33

[55]

a

90.56 ± 0.41

[5]

a

96.64 ± 0.80

[55]

a

99.89 ± 0.20

[55]

a

101.25 ± 0.35

[55]

a

120.75 ± 0.34

[57]

a

122.85 ± 0.47

[57]

a

127.32 ± 0.74

[57]

a

128.06 ± 0.36

[57]

a

128.90 ± 0.46

[57]

a

129.11 ± 0.71

[57]

a

129.25 ± 0.44

[57]

a

129.60 ± 0.39

[57]

a

130.45 ± 0.62

[57]

a

132.43 ± 0.47

[57]

a

132.51 ± 0.81

[57]

a

132.76 ± 0.54

[57]

a

134.15 ± 0.60

[57]

a

134.74 ± 0.44

[57]

a

135.05 ± 0.39

[57]

a

136.00 ± 0.49

[57]

204 205 206 207 208 209 210 211 212 213 214 215 216 217 218 219 220 221 222 223 224 225 226

(Continued)


97

Formula

BF3 affinity

Ref.

C6 H5 COC6 H5

62.86 ± 0.51 67.11 ± 1.58 69.7 ± 2.2 70.44 ± 0.49 70.7 ± 8.6 72.84 ± 0.48

[46] [54] [38] [46] [38] [46]

a

77.77 ± 1.01

[46]

a

79.73 ± 0.62

[46]

a

79.98 ± 0.55 80.01 ± 0.81 81.30 ± 0.44

[46] [46] [46]

83.01 ± 1.03 88.20 ± 0.85 115.56 ± 0.65 115.89 ± 0.64

[54] [46] [5] [46]

a

128.61 ± 0.51

[46]

3-BrC6 H4 COOEt 4-BrC6 H4 COOEt HCOOPh C6 H5 COOMe FCH2 COOEt 3-MeC6 H4 COOEt C6 H5 COOEt 4-FC6 H4 COOEt 4-MeC6 H4 COOEt

MeCOOEt

56.0 ± 0.8 57.8 ± 0.9 57.84 ± 0.44 59.4 ± 1.1 59.53 ± 0.27 61.0 ± 0.7 61.2 ± 0.8 61.7 ± 2.0 63.0 ± 1.5 64.19 ± 0.39 64.7 ± 2.0 66.4 ± 1.1b 67.63 ± 0.38 69.76 ± 0.11 71.03 ± 0.35 71.17 ± 0.29 72.79 ± 0.33 74.96 ± 0.70 75.13 ± 1.23 75.55 ± 0.31

[38] [38] [58] [5] [46] [38] [5] [38] [38] [5] [38] [46] [5] [5] [47] [5] [5] [58] [46] [47]

CCl3 CONMe2 CF3 CONMe2 ClCONMe2 N CCONMe2

55.74 ± 0.25 61.59 ± 0.33 72.20 ± 0.58 74 ± 1

[58] [58] [46] [46]


Lewis base

242

Miscellaneous conjugated ketones Benzophenone 9-Fluorenone 2,2-Dimethyl-1-tetralone 7-Nitro-1-tetralone Pivalophenone 1-(4-Nitrophenyl)-3-(2-thienyl)-2propen-1-one-E 1-(4-Chlorophenyl)-3-(2-thienyl)-2propen-1-one-E 1-Phenyl-3-(2-thienyl)-2-propen-1one-E 1-Tetralone 7-Methoxy-1-tetralone 1-(4-Methoxyphenyl)-3-(2-thienyl)2-propen-1-one-E 1-Indanone 6-Methoxy-1-tetralone 2,6-Dimethyl-γ -pyrone 4-Dimethylaminoantipyrine (aminopyrine) Antipyrine (phenazone)

243 244 245 246 247 248 249 250 251 252 253 254 255 256 257 258 259 260 261 262

Esters, carbonates Ethyl 3-bromobenzoate Ethyl 4-bromobenzoate Phenyl formate Methyl benzoate Ethyl fluoroacetate Ethyl 3-methylbenzoate Ethyl benzoate Ethyl 4-fluorobenzoate Ethyl 4-methylbenzoate Propylene carbonate Ethyl 4-methoxybenzoate Ethylene carbonate Dimethyl carbonate Methyl formate Diethyl carbonate Ethyl formate Methyl acetate Methyl cyclopropanecarboxylate γ -Butyrolactone Ethyl acetate

263 264 265 266

Amides, ureas N,N-Dimethyltrichloroacetamide N,N-Dimethyltrifluoroacetamide N,N-Dimethylcarbamoyl chloride N,N-Dimethylcarbamoyl cyanide

227 228 229 230 231 232 233 234 235 236 237 238 239 240 241

a

C12 H14 O C10 H9 O3 N C6 H5 CO-t-Bu a

C11 H12 O2

a a

C11 H12 O2

a a

a

4-MeOC6 H4 COOEt a

MeOCOOMe HCOOMe EtOCOOEt HCOOEt MeCOOMe c-PrCOOMe a

(Continued)

98



Formula

BF3 affinity

Ref.

267 268 269 270 271 272

MeOCONMe2 EtOCONMe2 HCONPh2 t-BuCONMe2 4-NO2 C6 H4 CONMe2 4-CF3 C6 H4 CONMe2

81.36 ± 0.46 84.16 ± 0.22 92.45 ± 0.49 92.73 ± 0.54 94.20 ± 0.57 98.38 ± 0.40

[58] [58] [58] [58] [50] [50]

a

4-BrC6 H4 CONMe2

98.78 ± 0.59 98.93 ± 0.38

[50] [5]

4-FC6 H4 CONMe2 HCON(Me)Ph C6 H5 CONMe2 4-MeC6 H4 CONMe2 4-MeOC6 H4 CONMe2

100.08 ± 0.35 100.83 ± 0.42 101.75 ± 0.24 102.70 ± 0.60 103.19 ± 0.57

[50] [58] [50] [50] [50]

Me2 NCONMe2 4-Me2 NC6 H4 CONMe2

108.62 ± 0.22 [47] 108.76 ± 0.34 [50]

HCONMe2 EtCONMe2

110.49 ± 0.18 110.56 ± 1.21 111.42 ±0.48 112.08 ± 0.99 112.13 ± 0.29

[47] [58] [46] [5] [5]

112.14 ± 0.41 112.56 ± 0.36 113.20 ± 0.35 113.61± 0.25 114.16 ± 0.57 116.94 ± 0.33

[47] [47] [5] [5] [5] [5]

273 274 275 276 277 278 279 280 281 282 283 284 285 286 287 288 289 290 291 292

N,N-Dimethyl methyl carbamate N,N-Dimethyl ethyl carbamate N,N-Diphenylformamide N,N,2,2,2-Pentamethylacetamide 4-Nitro-N,N-dimethylbenzamide 4-Trifluoromethyl-N,Ndimethylbenzamide 4-Bromo-N,N-dimethylbenzamide 1,3-Dimethyl-2-imidazolidinone (DMEU) 4-Fluoro-N,N-dimethylbenzamide N-Methyl-N-phenylformamide N,N-Dimethylbenzamide 4-Methyl-N,N-dimethylbenzamide 4-Methoxy-N,Ndimethylbenzamide Tetramethylurea 4-Dimethylamino-N,Ndimethylbenzamide N,N-Dimethylformamide N,N-Dimethypropionamide N-Methyl-ε-caprolactam 1-Isopropyl-2-pyrrolidinone 1,3-Dimethyl-3,4,5,6-tetrahydro2(1H)-pyrimidinone (DMPU) N,N-Dimethylacetamide 1-Methyl-2-pyrrolidinone N,N-Diethylformamide N,N-Diethylacetamide 1-Formylpiperidine 1-Methyl-2-pyridone

Carboxamidates 293 4-Nitro-N,N,Ntrimethylammoniobenzamidate 294 4-Cyano-N,N,Ntrimethylammoniobenzamidate 295 4-Trifluoromethyl-N,N,Ntrimethylammoniobenzamidate 296 4-Bromo-N,N,Ntrimethylammoniobenzamidate 297 4-Fluoro-N,N,Ntrimethylammoniobenzamidate 298 N,N,NTrimethylammoniobenzamidate 299 4-Methyl-N,N,Ntrimethylammoniobenzamidate

a a a

MeCONMe2

a

HCONEt2 MeCONEt2 HCON(CH2 )5 a

4-NO2 C6 H4 CON− N+ Me3

108.34 ± 0.58 [59]

4-N ≡ CC6 H4 CON− N+ Me3 110.83 ± 0.41 [59] 4-CF3 C6 H4 CON− N+ Me3

112.86 ± 0.37 [59]

4-BrC6 H4 CON− N+ Me3

113.14 ± 0.42 [59]

4-FC6 H4 CON− N+ Me3

113.36 ± 0.44 [59]

C6 H5 CON− N+ Me3

115.96 ± 0.31 [59]

4-MeC6 H4 CON− N+ Me3

116.90 ± 0.57 [59] (Continued)


99


Formula

300 4-Methoxy-N,N,Ntrimethylammoniobenzamidate 301 N,N,NTrimethylammonioacetamidate

4-MeOC6 H4 CON− N+ Me3 117.57 ± 0.40 [59]

Nitro compounds 302 Nitromethane 303 Nitrobenzene 304 305 306 307 308 309 310 311 312 313

Sulfinyl and sulfonyl compounds Ethylene sulfite Dimethyl sulfite Tetramethylene sulfone N,N,N ,N -Tetraethylsulfamide Diethyl sulfite Diphenyl sulfoxide Methyl phenyl sulfoxide Dimethyl sulfoxide Di-n-butyl sulfoxide Tetramethylene sulfoxide

BF3 affinity

Ref.

MeCON− N+ Me3

129.33 ± 0.95 [46]

MeNO2 PhNO2

37.63 ± 0.56c 35.79 ± 1.40c

[5] [47]

a

50.0 ± 0.5 51.27 ± 0.46 51.32 ± 0.29 55.08 ± 0.28 55.13 ± 0.77 90.34 ± 0.45 97.37 ± 0.26 105.34 ± 0.36 107.60 ± 0.46 108.10 ± 0.47

[46] [5] [5] [46] [5] [60] [46] [47] [5] [46]

65.3 ± 4.7b

[61]

66.41 ± 0.34 67 ± 3b

[46] [61]

(MeO)2 SO c-(CH2 )4 SO2 Et2 NSO2 NEt2 (EtO)2 SO Ph2 SO PhSOMe Me2 SO (n-Bu)2 SO c-(CH2 )4 SO

Phosphoryl compounds a 314 2-Phenoxy-5,5-diphenyl-2-oxo1,3,2-dioxaphosphorinane 315 Triphenyl phosphate (PhO)3 PO 316 2-Chloro-5,5-diphenyl-2-oxo-1,3,2- a dioxaphosphorinane a 317 2-Phenoxy-5,5-dimethyl-2-oxo1,3,2-dioxaphosphorinane a 318 2-Phenoxy-4-methyl-2-oxo-1,3,2dioxaphospholane a 319 2-Phenoxy-2-oxo-1,3,2dioxaphospholane a 320 2-Phenoxy-4-methyl-2-oxo-1,3,2dioxaphosphorinane 321 2-Chloro-5,5-dimethyl-2-oxo-1,3,2- a dioxaphosphorinane a 322 2-Phenoxy-4,4,6-trimethyl-2-oxo1,3,2-dioxaphosphorinane a 323 2-Phenoxy-2-oxo-1,3,2dioxaphosphorinane 324 2-Ethoxy-5,5-dimethyl-2-oxo-1,3,2- a dioxaphosphorinane a 325 2-Methoxy-5,5-dimethyl-2-oxo1,3,2-dioxaphosphorinane a 326 2-Ethoxy-2-oxo-1,3,2dioxaphospholane 327 2-Phenyl-5,5-diphenyl-2-oxo-1,3,2- a dioxaphosphorinane

73.12 ± 0.74b [61] 74.77 ± 0.79b [61] 76.57 ± 0.96b [62] 76.7 ± 2.2b

[46]

78.1± 4.3b

[61]

79.6± 2.2b

[63]

82± 5b

[61]

82.09 ± 0.91b [61] 82.25 ± 0.75b [61] 83.62 ± 0.71b [62] 84.4 ± 4.1b

[61] (Continued)

100



Lewis base

328 329

338 339 340 341 342 343

Trimethyl phosphate 2-Phenyl-5,5-dimethyl-2-oxo-1,3,2dioxaphosphorinane 2-Methoxy-2-oxo-1,3,2dioxaphospholane 3-(Diethoxyphosphoryl)-5,5dimethyl-2-cyclohexen-1-one 2-(Dimethylamino)-5,5-diphenyl-2oxo-1,3,2-dioxaphosphorinane 2-Phenyl-2-oxo-1,3,2dioxaphosphorinane 2-Phenyl-2-oxo-1,3,2dioxaphospholane Diisopropylphosphonate 2-Dimethylamino-2-oxo-1,3,2dioxaphospholane 2-Dimethylamino-5,5-dimethyl-2oxo-1,3,2-dioxaphosphorinane Triphenylphosphine oxide Hexamethylphosphoric triamide Tripiperidinophosphine oxide Triethylphosphine oxide Trimethylphosphine oxide Tripyrrolidinophosphine oxide

344 345

330 331 332 333 334 335 336 337

BF3 affinity

Ref.

(MeO)3 PO

84.75 ± 0.22 86 ± 13b

[47] [61]

a

86.07 ± 0.33b

[62]

a

87.36 ± 0.23

[46]

a

87.8 ± 8.3b

[46]

a

89.0 ± 1.3b

[61]

a

89.2 ± 2.0b

[62]

(i-PrO)2 POH

89.56 ± 0.72 91.12 ± 0.54b

[46] [62]

a

92.5 ± 1.6b

[61]

Ph3 PO (Me2 N)3 PO [c-(CH2 )5 N]3 PO Et3 PO Me3 PO [c-(CH2 )4 N]3 PO

103.3 ± 1.50 117.53 ± 0.45 118.14 ± 0.92 119.28 ± 0.74 119.68 ± 0.45 122.52 ± 0.14

[46] [47] [5] [46] [46] [5]

Thioether Tetrahydrothiophene

c-(CH2 )4 S

51.62 ± 0.20

[46]

Thiocarbonyl compound 1,1,3,3-Tetramethyl-2-thiourea

Me2 NCSNMe2

60.96 ± 0.40

[46]

a

66.8 ± 1.4b

[63]

a

67.4 ± 1.3b

[63]

Me3 P

97.43 ± 0.30

[46]

a

a

Thiophosphoryl compounds 2-Phenoxy-5,5-dimethyl-2-thio1,3,2-dioxaphosphorinane 2-Phenoxy-4,4,6-trimethyl-2-thio1,3,2-dioxaphosphorinane

346 347

Phosphine Trimethylphosphine

348 a

Formula

Formulae: N N

N

O

N

N

O

O

N

10

11

50

O O

R

O O

O

S

51

74

75

R= NO2, C≡N, Br, Cl, I, 139 140 141 142 143 R= COOMe, C≡CH, OMe, C6H5, 145 146 147 148

O

O

83

85

86

OCOMe 144 CH=CH2, H, Me, Et 149 150 151 152

130 O 6

1

5

2 3

4

138

179 (Continued)


101

Table 3.3 (Continued) O

R= C N, 188 R

COOMe, 189

CH2COOEt, 195

H, 196

CH2COEt, 201 Me, 207

Br, 190

OCOPh, 197

4-FC6H4, 202

SEt, 208

Cl, 191

OMe, 209

CH2C N, 192

4-NO2C6H4, 193

4-BrC6H4, 198

OCOMe, 199

Ph, 203

4-C6H5C6H4, 204

OEt, 210

NPh2, 211

4-C NC6H4 194

4-ClC6H4, 200

4-MeC6H4, 205

4-MeOC6H4, 206

N(CH2Ph)2, 212

N(Me)Ph, 213

N(Me)CH2Ph O

N(Et)CH2Ph S

N O

N

MeN

N

N

214

215

216

217

218

NMe2

NEt2

N

N

N

219 N

N

220

221

222

O

223

O H C

H C

C

224

225

226 O

R= NO2, Cl, H, MeO 232 233 234 237

R

8

S

2

1

7

3

6 5

228

235 O

O

NMe O

O

N

O

N

238

O

Ph

240

241

N N

6 R

P 2 O

5

1 O

5

P2

4

O 3

252

O

254

261

N

274

O

O

O

N

S

O

O

X= R=

PhO, PhO, Me, H, 318 319

288

PhO, Cl, Me, Me, 317 321

292

PhO, EtO, H, Me, 323 324

304

MeO, Ph, Me, Ph, 325 327

EtO, MeO H, H, 326 330

Ph, H, 334

Ph, Me, 329

Me2N, Ph, Ph, H, 332 333

Me2N Me 337

Me2N H 336

O

O

O

320

O

OPh

P

S

O O

322

PO(OEt)2

331

O

OPh

P

P

O

d

PhO, Cl, Ph, Ph, 314 316

OPh

P

c

O

O

286

O

OPh

b

R

N

O

4

X O

N

N

O O

O

O

O

242

N

X= R=

R O 3

N Ph

N

285

1 O

X

O

iPr

O

284

4

S

O

346

Secondary value calculated from measurements in nitrobenzene and relationship (3.9). See text. Secondary value calculated from measurements in the bulk liquid and relationship (3.9). See text. Published values have been corrected (×0.988 011) after improvement of the calibration technique.

O

347

102


Thus −∆H ◦ (3.10) = 37.68 kJ mol−1 when gaseous BF3 is complexed with bulk nitrobenzene. Since bulk nitrobenzene plays the role of solute and solvent, Equation 3.9 can be used to calculate a BF3 affinity of nitrobenzene in CH2 Cl2 of 35.79 kJ mol−1 . A series of experiments carried out on bulk nitromethane gave −∆H ◦ (3.10) = 39.60 kJ mol−1 . Assuming that bulk nitromethane may be considered (in this context) as nitromethane dissolved in nitrobenzene, a BF3 affinity of nitromethane in CH2 Cl2 of 37.63 kJ mol−1 can be calculated from Equation 3.9. Di-tert-butyl ketone was also investigated as a bulk liquid and corrected in the same way. For these weak Lewis bases, the corrections are very small. Equation 3.9 also allows the calculation of secondary BF3 affinities for triethylamine 6, DABCO 10 and quinuclidine 11. These tertiary amines react slowly with dichloromethane. Few measurements are possible immediately after preparation of the solution, but the reduced calorimetric stabilization time and side-reactions induce a decrease in accuracy and a possible overestimation of their BF3 affinities. Consequently, the complexation of these amines with gaseous BF3 was performed in nitrobenzene, and their BF3 affinities in CH2 Cl2 were obtained from Equation 3.9. In the case of 11, the uncertainty about the value directly measured in CH2 Cl2 is largely reduced by this means. To understand how the BF3 affinity depends on self-association, the hydrogen-bonded solvents formamide, methanol and water were also studied as bulk liquids. The results [46] are –∆H ◦ (3.10) = 115.32 ± 0.90, 109.20 ± 0.44 and 104.40 ± 0.36 kJ mol–1 , respectively. There are two obvious terms in these enthalpies: (i) the enthalpy of complexation of BF3 and (ii) the enthalpy due to other interactions and to the reorganization of the liquid structure. It would be necessary to correct for the latter terms to isolate the BF3 affinity of these self-associated compounds.

3.5 3.5.1

Discussion Medium Effects

The use of CH2 Cl2 as a solvent for measuring Lewis affinity scales has been criticized [64] because CH2 Cl2 , a weak C H hydrogen-bond donor, is itself a Lewis acid. However, even if hydrogen bonds between CH2 Cl2 and solutes and/or their BF3 adducts exist, it is sufficient that the energy involved is weak and roughly proportional to the solute’s Lewis affinity to establish a relative affinity scale. This is probably the case for the BF3 /solute/CH2 Cl2 system, as shown by the linear relationship 3.9 between BF3 complexation enthalpies measured in CH2 Cl2 and in the non-hydrogen-bond donor solvent C6 H5 NO2 . Moreover, BF3 affinities measured in the gas phase, collected in Table 3.4, show a very similar trend to those measured in CH2 Cl2 [see Equation 1.96]. These observations on medium effects indicate that the BF3 affinity scale is free from large and unpredictable solvation effects. 3.5.2

Hardness of BF3

In his 1963 empirical classification of Lewis acids, Pearson put BF3 in the category of hard acids (see Section 1.11.1 in Chapter 1). In fact, semi-quantitative studies in the gas phase [3] show that the BF3 basicity of methyl derivatives of the elements of groups 15 and 16 decreases in the order N > P > As > Sb

(group 15)


103

Table 3.4 Comparison of gas-phase and CH2 Cl2 solution data for the enthalpy of complexation (kJ mol−1 ) of BF3 with Lewis bases. Lewis base Trimethylamine Trimethylphosphine Tetrahydrofuran Tetrahydropyran Dimethyl ether Ethyl acetate Diethyl ether Tetrahydrothiophene a

∆H◦ (3.1) gas

Ref.

−∆H◦ (3.8) CH2 Cl2 solutiona

111.3 79.1 70.3 64.5 57.1 53.6 49.9 21.8

[65] [2] [66] [66] [67] [68] [67] [69]

129.5 87.4 80.4 75.4 73.6 65.6 68.8 41.6

Corrected for the BF3 enthalpy of solution in CH2 Cl2 = −10 kJ mol-1 .

and O > S > Se

(group 16)

This is confirmed by the quantitative BF3 affinity scale, which shows that trimethylamine and tetrahydrofuran have a higher BF3 affinity than trimethylphosphine and tetrahydrothiophene, respectively. Moreover, the thiocarbonyl and thiophosphoryl sulfur bases show lower BF3 affinities than the corresponding carbonyl and phosphoryl oxygen bases. Indeed, the decrease in basicity (affinity) in descending groups 15 and 16 of the periodic table is characteristic of hard basicity scales. The hard character of the BF3 affinity scale has been confirmed theoretically. The theoretical absolute hardness of BF3 , η = 9.7 eV, is fairly high compared with other molecular Lewis acids (see Table 1.17). The hardness of the interaction of BF3 with NH3 , NMe3 and CO has been studied [70] through the local Hard–Soft Acid–Base principle. More generally, the nature of the dative bond in BF3 complexes has been characterized by various means: valence-bond method [25], calculation of the amount of intermolecular electron transfer [71] and partitioning of the interaction energy [24, 72]. 3.5.3

Comparison of the BF3 and SbCl5 Affinity Scales

Comparison of the BF3 and SbCl5 affinities collected in Tables 3.3 and 2.2, respectively, shows that thermal effects are roughly equal for the same Lewis base. The relationship 3.11 confirms a linear trend with a slope of about 1 and a low value of the intercept: SbCl5 A = −4.83 + 1.09 BF3 A

(3.11)

n = 36, r = 0.963, s = 7.4 kJ mol−1 A better fit should be expected since BF3 and SbCl5 (i) are both hard Lewis acids, (ii) have similar Lewis acid strengths and (iii) have similar steric requirements. The variance of SbCl5 A not explained by BF3 A is partly attributable to the experimental difficulties encountered in the measurement of SbCl5 affinities (see Section 2.4 in Chapter 2).

104

Lewis Basicity and Affinity Scales NH2

NH2 N

S

N

Me

Me

2

O O

A

N

N 3

N H B

N

N

H C

Scheme 3.1 Sulfinimine, cytosine and adenine.

3.5.4

Computation of the BF3 Affinity

The existence of numerous theoretical studies on the Lewis affinity of BF3 and on the nature of the dative bond in BF3 complexes, generally performed at high theoretical levels, illustrates that the choice of BF3 as a reference Lewis acid for constructing an affinity scale of Lewis bases is well founded, not only from the experimental but also from the computational point of view. Indeed, the electronic structure of BF3 is very simple, since this trigonal planar molecule contains only four first-row atoms. Calculations on the thermodynamics of BF3 complexes can significantly improve our knowledge of Lewis affinity, as illustrated below. It is possible to compute the BF3 Lewis affinity, at 0 and/or 298 K in the gas phase, of bases that are difficult to study experimentally, such as weak carbon π bases [41, 42, 73], water [74] and diaminocarbenes [75]. For polyfunctional bases, the main site of BF3 fixation can be determined theoretically. For example, in sulfinimine (A) [76], cytosine (B) [77] and adenine (C) [78] (Scheme 3.1), the main Lewis basic sites towards BF3 are the nitrogen atom, the carbonyl oxygen O2 and the pyrimidine nitrogen N3, respectively. Further, the small size of BF3 enables the BF3 affinity in the gas phase for a large series of Lewis bases to be calculated at a high level of theory. For example, as early as 1994, ∆H ◦ (3.1) was computed at the MP2/TZ2P//MP2/TZ2P level for eight diverse carbon, nitrogen and oxygen bases [79], and at the MP3/6–31G∗∗ //HF/6–31G∗∗ level for 17 oxygen bases [37]. There is a very good relative agreement, illustrated in Figure 3.3, between the calculated and experimental BF3 affinities for nine diverse oxygen and nitrogen bases, common to the experimental and theoretical sets, since 99% of the variance of the experimental BF3 affinity scale in CH2 Cl2 solution is explained by calculations corresponding to the gas phase. More recent calculations [80] of BF3 affinities have been performed at several computational levels for a series of eight Lewis bases for which BF3 affinities had been measured in the gas phase (see Table 3.4). This choice of Lewis bases allowed an absolute comparison between computed and experimental BF3 affinities for both first- and second-row bases. It was concluded that B3LYP/6–311+G∗∗ calculations seem to be the best compromise between speed and accuracy for predicting trends in BF3 affinities, whereas the Complete-Basis-Set QB3 method [81] can be recommended for predicting absolute BF3 affinities, the mean absolute deviation being 10 kJ mol−1 . Difficulties were encountered, however, in applying the BSSE correction by the counterpoise method and in reproducing the BF3 affinity of Me3 P, a second-row base.


105

150 Exp. BF3 affinities / kJ mol

-1

99

100 8 4 5 2

7 6

3 1

50

0 0

50

100

150 -1

Calc. BF3 affinities / kJ mol

Figure 3.3 Comparison of BF3 affinities measured in CH2 Cl2 solution and calculated ab initio in vacuo. 1 MeCN, 2 HCOOMe, 3 MeCHO, 4 PhCHO, 5 Me2 CO, 6 MeCOOMe, 7 Me2 O, 8 c-(CH2 )4 O, 9 Me3 N.

Consequently, the Complete-Basis-Set QB3 results explain only 84% of the variance of the experimental results. The relative agreement rises to 93% when Me3 P is excluded.

3.6

Conclusion

The BF3 affinity scale represents an improved version of the SbCl5 affinity scale (DN scale). The concept is similar but the methodology produces cleaner complexation reactions and consequently more accurate complexation enthalpies. Additionally, the BF3 affinity database is more comprehensive and more varied than the SbCl5 affinity database. Finally, compared with SbCl5 the electronic structure of BF3 is much simpler. These advantages should lead to remarkable advances in our knowledge of dative-bond affinity. The statistical comparisons with hydrogen-bond, halogen-bond and cation affinity scales (see Chapters 4–6) should become more reliable (see refs. [5, 47] for preliminary statistical studies). Structure–dative-bond affinity relationships should be better established and, from easier quantum chemistry calculations, less empirically explained in terms of electronic and steric effects than in the past. Finally, dative-bond BF3 affinities should be computed more extensively and these computed affinities might gradually substitute for experimental ones. However, in spite of impressive hardware and software progress in computational chemistry and of the development of improved implicit and explicit models for condensed phases [82], difficulties remain in computing accurate absolute BF3 affinities, both in solution (P. Burk, University of Tartu, personal communication) and in the gas phase [80]. Undoubtedly, the comparison of the experimental accurate BF3 affinities presented here with computed values should be of help in choosing the most appropriate computational strategy and in validating new solvation models in the field of dative bond modelling.

106


References 1. Gay-Lussac, L.J. (1809) Mémoire sur la combinaison des substances gazeuses, les unes avec les autres, Mémoires de Physique et de Chimie de la Société d’Arcueil, 2, 207–234. 2. Brown, H.C., McDaniel, D.H. and Hafliger, O. (1955) Dissociation constants, in Determination of Organic Structures by Physical Methods (eds E.A. Braude and F.C. Nachod), Academic Press, New York, pp. 634–643. 3. Stone, F.G.A. (1958) Stability relations among analogous molecular addition compounds of Group III elements. Chem. Rev., 58, 101–129. 4. Satchell, D.P.N. and Satchell, R.S. (1969) Quantitative aspects of the Lewis acidity of covalent metal halides and their organo derivatives. Chem. Rev., 69, 251–278. 5. Maria, P.C. and Gal, J.F. (1985) A Lewis basicity scale for nonprotogenic solvents: enthalpies of complex formation with boron trifluoride in dichloromethane. J. Phys. Chem., 89, 1296–1304. 6. Allen, F.H. and Taylor, R. (2004) Research applications of the Cambridge structural database (CSD). Chem. Soc. Rev., 33, 463–475. 7. Phillips, J.A., Halfen, J.A., Wrass, J.P. et al. (2006) Large gas–solid structural differences in complexes of haloacetonitriles with boron trifluoride. Inorg. Chem., 45, 722–731. 8. Phillips, J.A., Giesen, D.J., Wells, N.P. et al. (2005) Condensed-phase effects on the structural properties of C6 H5 CN–BF3 and (CH3 )3 CCN–BF3 : IR spectra, crystallography, and computations. J. Phys. Chem. A, 109, 8199–8208. 9. Geller, S. and Hoard, J.L. (1950) Structures of molecular addition compounds. I. Monomethyl amine–boron trifluoride, H3 CH2 N·BF3 . Acta Crystallogr., 3, 121–129. 10. Hoard, J.L., Geller, S. and Owen, T.B. (1951) Structures of molecular addition compounds. V. Comparison of four related structures. Acta Crystallogr., 4, 405–407. 11. Penner, G.H., Ruscitti, B., Reynolds, J. and Swainson, I. (2002) Structure and dynamics of ND3 BF3 in the solid and gas phases: a combined NMR, neutron diffraction, and ab initio study. Inorg. Chem., 41, 7064–7071. 12. Janda, K.C., Bernstein, L.S., Steed, J.M. et al. (1978) Synthesis, microwave spectrum, and structure of ArBF3 , BF3 CO, and N2 BF3 . J. Am. Chem. Soc., 100, 8074–8079. 13. Leopold, K.R., Fraser, G.T. and Klemperer, W. (1984) Rotational spectroscopy of molecular complexes of boron fluoride with ethanedinitrile, carbon dioxide, and nitrous oxide. J. Am. Chem. Soc., 106, 897–899. 14. Phillips, J.A., Canagaratna, M., Goodfriend, H. et al. (1995) Microwave and ab initio investigation of HF–BF3 . J. Am. Chem. Soc., 117, 12549–12556. 15. Reeve, S.W., Burns, W.A., Lovas, F.J. et al. (1993) Microwave spectra and structure of hydrogen cyanide–boron trifluoride: an almost weakly bound complex. J. Phys. Chem., 97, 10630–10637. 16. Dvorak, M.A., Ford, R.S., Suenram, R.D. et al. (1992) van der Waals vs. covalent bonding: microwave characterization of a structurally intermediate case. J. Am. Chem. Soc., 114, 108–115. 17. Odom, J.D., Kalasinsky, V.F. and Durig, J.R. (1975) Spectra and structure of phosphorus–boron compounds. XI. Microwave spectrum, structure, dipole moment, and barrier to internal rotation in phosphine–trifluoroborane. Inorg. Chem., 14, 2837–2839. 18. Iijima, K., Yamada, T. and Shibata, S. (1981) Temperature effects of the molecular structure of dimethyl ether–boron trifluoride as studied by gas-phase electron diffraction. J. Mol. Struct., 77, 271–276. 19. Fujiang, D., Fowler, P.W. and Legon, A.C. (1995) Geometric and electric properties of the donor–acceptor complex H3 N–BF3 . J. Chem. Soc., Chem. Commun., 113–114. 20. Iijima, K., Noda, T., Maki, M. et al. (1986) The molecular structures of complexes of pyridine with boron trifluoride, boron trichloride and boron tribromide as studied by gas phase electron diffraction. J. Mol. Struct., 144, 169–179. 21. Cassoux, P., Kuczkowski, R.L. and Serafini, A. (1977) Microwave spectrum and structure of trimethylamine–boron trifluoride. Me3 N·BX3 adduct stability and the reorganization energies of trifluoroborane and borane. Inorg. Chem., 16, 3005–3008. 22. Branchadell, V., Sbai, A. and Oliva, A. (1995) Density functional study of complexes between Lewis acids and bases. J. Phys. Chem., 99, 6472–6476.


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47. Maria, P.C., Gal, J.F., De Franceschi, J. and Fargin, E. (1987) Chemometrics of solvent basicity: multivariate analysis of the basicity scales relevant to nonprotogenic solvents. J. Am. Chem. Soc., 109, 483–492. 48. Berthelot, M., Gal, J.F., Laurence, C. and Maria, P.C. (1984) Chemometrics of basicity. 1. Comparison of the basicity of o-, m-, and p-substituted pyridines toward boron trifluoride and methanol. J. Chim. Phys. Phys.-Chim. Biol., 81, 327–331. 49. Maria, P.C., Gal, J.F. and Taft, R.W. (1987) Proton affinities versus enthalpies of complexation with boron trifluoride in solution, in the nitrile series: role of the polarizability. New J. Chem., 11, 617–621. 50. Gal, J.F., Geribaldi, S., Pfister-Guillouzo, G. and Morris, D.G. (1985) Basicity of the carbonyl group. Part 12. Correlations between ionization potentials and Lewis basicities in aromatic carbonyl compounds. J. Chem. Soc., Perkin Trans. 2, 103–106. 51. Derrieu, G., Gal, J.F., Elegant, L. and Azzaro, M. (1974) Carbonyl group basicity. IV. Application of the Taft–Pavelich equation on the enthalpies of complexation of ketones (CH3 COR) with boron trifluoride. C. R. Acad. Sci., Ser. IIC, 279, 705–707. 52. Azzaro, M., Derrieu, G., Elegant, L. and Gal, J.F. (1975) Basicity of the carbonyl group. V. Applicability of the Taft–Pavelich equation to cyclic systems with reference to the complexation enthalpy of cyclic ketones using boron trifluoride. J. Org. Chem., 40, 3155–3157. 53. Gal, J.F., Morris, D.G. and Rouillard, M. (1992) Basicity of the carbonyl group. Part 13. 4Substituted camphors as models for transmission of polar effects. A calorimetric, infrared and 1 H NMR study. J. Chem. Soc., Perkin Trans. 2, 1287–1293. 54. Gal, J.F., Elegant, L. and Azzaro, M. (1974) Basicity of the carbonyl group. II. Effect of substituents on the enthalpy of complex formation between substituted aromatic carbonyl compounds and boron trifluoride. Bull. Soc. Chim. Fr., 411–414. 55. Azzaro, M., Gal, J.F., Geribaldi, S. and Loupy, A. (1982) Transmission of substituent effects through unsaturated systems. 5. Brønsted and Lewis basicities of β-substituted α,β-unsaturated ketones. J. Org. Chem., 47, 4981–4984. 56. Elegant, L., Paris, C., Gal, J.F. and Azzaro, M. (1974) Basicity of the carbonyl group. Enthalpies of complexation between a series of cyclohexenones and boron trifluoride. Thermochim. Acta, 9, 210–212. 57. Azzaro, M., Gal, J.F., Geribaldi, S. et al. (1983) Basicity of the carbonyl group. Part 9. Lewis and Brønsted basicities of enamino ketones. J. Chem. Soc., Perkin Trans. 2, 57–65. 58. Berthelot, M., Gal, J.F., Helbert, M. et al. (1985) Chemometrics of basicity. Part 2. Comparison of the basicity of carbonyl compounds with respect to boron trifluoride and methanol. J. Chim. Phys. Phys.-Chim. Biol., 82, 427–432. 59. Gal, J.F. and Morris, D.G. (1978) Basicity of the carbonyl group. Part 6. Calorimetric and spectrometric study of complexation of para-substituted N-ammoniobenzamidates by boron trifluoride. J. Chem. Soc., Perkin Trans. 2, 431–435. 60. Kamlet, M.J., Gal, J.F., Maria, P.C. and Taft, R.W. (1985) Linear solvation energy relationships. Part 32. A coordinate covalency parameter, ξ , which, in combination with the hydrogen bond acceptor basicity parameter, β, permits correlation of many properties of neutral oxygen and nitrogen bases (including aqueous pK a ) . J. Chem. Soc., Perkin Trans. 2, 1583–1589. 61. Maria, P.C., Elegant, L., Azzaro, M. et al. (1971) Phosphorus-containing heterocycles. X. Properties of a series of 2-oxo-1,3,2-dioxaphosphorinanes toward boron trifluoride. Measurement of the addition enthalpies. Bull. Soc. Chim. Fr., 3750–3753. 62. Maria, P.C., Elegant, L., Azzaro, M. et al. (1972) Phosphorus-containing heterocycles. XV. Behavior of a series of 2-oxo-1,3,2-dioxaphospholanes toward boron trifluoride determined from enthalpy of addition. Thermochim. Acta, 4, 505–511. 63. Maria, P.C., Elegant, L., Azzaro, M. et al. (1972) Calorimetric study of the comparative reactivity of 2-oxo- and 2-thiodioxaphosphorinanes with boron trifluoride. Tetrahedron Lett., 1485–1486. 64. Drago, R.S., Nusz, J.A. and Courtright, R.C. (1974) Solvation contributions to enthalpies measured in methylene chloride. J. Am. Chem. Soc., 96, 2082–2086. 65. Bauer, S.H. and McCoy, R.E. (1956) Energetics of the boranes. III. The enthalpy and heat capacity of trimethylamine-trifluoroborane as determined by the drop method. J. Phys. Chem., 60, 1529–1532.


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4 Thermodynamic and Spectroscopic Scales of Hydrogen-Bond Basicity and Affinity Water, alcohols, phenols, hydrogen halides, carboxylic acids, primary and secondary amines and amides, thiols, 1-alkynes and so on constitute a special class of Lewis acids. They all contain an XH group in which the hydrogen atom is bonded to an electronegative atom X and bears a partial positive charge. Therefore, the hydrogen is attracted, by means of electrostatic forces, toward high electron density regions, mainly the lone pairs of Lewis bases B, and a hydrogen bond (HB) X H · · · B is formed. In addition to this electrostatic interaction, there is a charge transfer from the base to the antibonding orbital σ ∗ of the X H sigma bond. The respective contributions of electrostatic forces and charge transfer are a matter of debate. A modern electrostatic model of the hydrogen bond provides a near-quantitative description of structure and properties for a wide range of typical hydrogen bonds [1]. In contrast, NBO analysis suggests that charge transfer of n→σ ∗ type is a characteristic feature of hydrogen bonding [2, 3]. In any case, the existence of even a small charge transfer allows the hydrogen-bonding interaction to be considered as Lewis acid/base in nature [4]. The hydrogen-bond donor (HBD) XH is the electron acceptor and the hydrogen-bond acceptor (HBA) B is the electron donor. In the HSAB concept, HBDs are classified as hard Lewis acids [4] from the simple observation that the hydrogen bond is stronger when B is fluorine, oxygen or nitrogen than when it is iodine, sulfur or phosphorus, respectively. The reaction 4.1 of formation of a hydrogen bond between the molecules XH and B: X−H + B X−H · · · B

(4.1)

can be used for the construction of a Lewis basicity scale, if a reference HBD is chosen and if the equilibrium constant K of the reaction is measured for a series of bases B under the same conditions: physical state (gas or solution in a given solvent), temperature and pressure. Such a scale, as logK or ∆G, must be named a hydrogen-bond basicity scale. Lewis Basicity and Affinity Scales: Data and Measurement C 2010 John Wiley & Sons, Ltd


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Similarly, the measurement of the enthalpy of reaction 4.1 leads to a hydrogen-bond affinity scale. Due to the importance of the hydrogen bond in physical, chemical and biological sciences, a great number of hydrogen-bond equilibrium constants and enthalpies have been measured. The first text devoted entirely to hydrogen bonding, by Pimentel and McClellan (1960) [5], contains a table of nearly 300 entries of thermodynamic data for HB formation. A book by Joesten and Schaad (1974) [6] gives an appendix of thermodynamic data from 1960 to 1973 containing nearly 2000 entries. Among this accumulation of measurements, only those of Taft [7–9] and Arnett [10, 11] aimed at building hydrogen-bond basicity and affinity scales. Later, the groups of Abraham [12, 13] and Laurence [14] worked towards the same objective. The formation of a hydrogen-bonded complex X H · · · B induces significant changes in the IR, UV and NMR spectra of the subunits XH and B. The shift to lower wavenumbers, upon hydrogen bonding, of the stretching vibration of the X H bond is the most spectacular of these spectral changes and has been the most studied. A relationship between these IR shifts, ∆ν(XH), for a given HBD in a series of hydrogen-bonded complexes with various bases, and the corresponding enthalpy of hydrogen bonding has been advocated by several workers, but challenged by others. It appears today that the ∆H–∆ν relationship is not general but family dependent [15, 16]. Nevertheless, spectroscopic scales, such as ∆ν(OH) of phenol, have been used as substitutes for hydrogen-bond basicity scales and have proved useful in many treatments of the effect of basic solvents on thermodynamic, kinetic or spectral properties. Similarly, UV spectroscopic scales have been defined from solvatochromic shifts upon hydrogen bonding of the electronic transitions of probes such as 4-nitroaniline and 4nitrophenol. Among these scales, it is essential to make a distinction between solute and solvent scales. A solute scale is obtained when the HBD probe is studied in dilute solutions of a series of bases in the same given solvent, whereas a solvent scale comes from measurements on binary systems of the HBD probe dissolved in the pure bases. The same distinction must be made for hydrogen-bond affinity scales. When the heats of reaction of, for example, 4-fluorophenol with a series of bases are measured in a ternary system (4-FC6 H4 OH/B/solvent), a solute scale is obtained, whereas measurements on the binary system 4-FC6 H4 OH/B, corrected by the heats of solution of a similar but non-HBD probe such as 4-FC6 H4 OMe (the so-called pure base calorimetric method) [10], furnish a solvent scale. The degree of equivalence of solute and solvent scales will be considered later. A number of scales proposed in the literature for measuring the hydrogen-bond basicity (affinity) of Lewis bases are summarized in Table 4.1. In this book, we will select four of them. The first is a thermodynamic scale of hydrogen-bond basicity built from the equilibrium constant of the reaction 4.2 of 4-fluorophenol with a series of Lewis bases in CCl4 at 25 ◦ C: 4-FC6 H4 OH + B 4-FC6 H4 OH · · · B

(4.2)

The second is a hydrogen-bond affinity scale constructed from the enthalpy of the above reaction. The third is a spectroscopic scale constructed from the IR wavenumber shift of the OH stretching vibration of methanol, upon hydrogen bonding with a series of Lewis bases diluted in CCl4 . The fourth is a solvatochromic scale based on the hydrogen-bond shifts of the π → π ∗ transition of 4-nitrophenol at 286 nm. Before presenting these scales in depth, we give some information about the structure of hydrogen-bonded complexes.

Thermodynamic and Spectroscopic Scales of Hydrogen-Bond Basicity and Affinity

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Table 4.1 Hydrogen-bond basicity, affinity and spectroscopic scales. Basicity-dependent property

Number of bases Ref.

Symbol

Solvent

log Kf , pKHB pKHB

CCl4 CCl4

75 117

[7, 9] [8]

log Kβ log K BH , β H 2

CCl3 CH3 CCl4

90 215

[12] [13]

Enthalpies 4-FC6 H4 OH + B Complex 4-FC6 H4 OH + B → Complex Pyrrole + B → Complex

SB

CCl4 Pure base Pure base

40 57 35

[10, 11] [10, 11] [17]

IR shifts ν(OD), methanol-OD ν(OH), phenol ν(OH), methanol

∆ν D B ∆ν(OH)

Pure base CCl4 CCl4

89 198 573

[18] [19] [20, 21]

Equilibrium constants 4-FC6 H4 OH + B Complex OH donors + B Complex (statistical average) 4-NO2 C6 H4 OH + B Complex OH, NH, CH donors + B Complex (statistical average)

Bathochromic shifts of the longest wavelength π → π ∗ transition Shift of 4-nitrophenol upon hydrogen β sm CCl3 CH3 bonding to B in a solvent Enhanced shift of 5-nitroindoline relative SB Pure base to 1-methyl-5-nitroindoline Enhanced shift of 4-nitroaniline relative ∆˜ν (NH2 – NMe2 ) Pure base to N,N-dimethyl-4-nitroaniline Enhanced shift of 4-nitrophenol relative ∆˜ν (OH – OMe) Pure base to 4-nitroanisole Miscellaneous basicity-dependent β Undefined properties

4.1

85

[12]

202

[22]

189

[23, 24]

191

[23, 24]

164

[25–27]

Structure of Hydrogen-Bonded Complexes

The hydrogen bond may be described in terms of r, d, θ and Φ as shown in Figure 4.1. The distances d and r measure the length of the hydrogen bond and the XH bond, respectively. The latter varies significantly on going from the free to the bonded XH group. This modification induces a change in the XH force constant and, consequently, in the XH stretching frequency. The angles θ and Φ are called the linearity and directionality of the hydrogen bond, respectively.

Φ R

Figure 4.1

B

d

r H

X

q

Radial and angular descriptors of the hydrogen-bond geometry.

114


In the gas phase, the structures of a large number of hydrogen halide complexes X H · · · B have been investigated by rotational spectroscopy [28]. A number of rules have been derived from these structures: (i) Generally, the hydrogen bond approaches linearity (θ = 180◦ ). (ii) The axis of the HX molecule coincides with the supposed axis of a nonbonding electron pair as conventionally envisaged. For example, the H F molecule lies along the C2 axis of the C5 H5 N subunit in the pyridine–hydrogen fluoride complex. (iii) If B has π-bonding pairs, the axis of the HX molecule intersects the internuclear axis of the atoms forming the π-bond and is perpendicular to the plane of symmetry of the π -bond. Thus, a T-shaped geometry is observed for the acetylene–hydrogen chloride complex. (iv) In furan and thiophene, B has both nonbonding and π -bonding pairs. The hydrogen halides give either n complexes of C2v symmetry (as in the case of HF and HCl with furan) or π complexes of Cs symmetry, with the HX subunit sampling the π electron density near to the heteroatom (as in the case of HF, HCl and HBr with thiophene and of HBr with furan). In addition to hydrogen halides, water is the most studied HBD in the gas phase. The structures of the complexes with water, dinitrogen, carbon monoxide, ozone, benzene, ethane, formaldehyde, formamide, 1,4-dioxane, ethylene oxide, tetrahydropyran, difluoromethane, pyrazine, pyrimidine, pyridazine, benzonitrile, quinuclidine, ammonia, methylamine, trimethylamine and so on can be found in the Mogadoc database [29]. In the water–morpholine complex [30], water is hydrogen bonded to the nitrogen and not to the oxygen, as predicted by the higher HB basicity of amines than that of ethers. Two possible hydrogen-bonded structures can be observed for complexes between amphoteric molecules. For example, phenol and water can form either complex A, where phenol acts as the HBD, or B, where water serves as the HBD (Scheme 4.1). The hydrogenbond complexation constants corresponding to each conformer can be empirically predicted [31, 32] (see above, Equation 1.131) from a knowledge of the hydrogen-bond basicities and acidities of phenol and water. In terms of Gibbs energies of complexation, it is found that A is 7.8 kJ mol−1 more stable than B. Indeed, it is complex A that is observed in the microwave spectrum of the phenol–water system [33]. Observed spectra and empirically calculated Gibbs energies also agree for the systems phenol/methanol [34], methanol/aniline [35] and pyrrole/water [36]. H O

H O

H

O H

H

H O

A

B

Scheme 4.1 Phenol is the hydrogen-bond donor in A and the hydrogen-bond acceptor in B.


115

Figure 4.2 Structure of the solid 1 : 1 complex 2-methylpyridine-pentachlorophenol (r and d ˚ θ in degrees). in A,

Another empirical equation has been established to predict the strength of hydrogen bonds between small amphoteric molecules such as H2 O, HF, HCl, HBr, HC N, HC CH and NH3 [37], from the intermolecular stretching force constants. This equation can be used to quantify the spectroscopic observation that water is the HBA in its complexes with HCl, HBr, HC N and HC CH and is the HBD with NH3 . The structures of many hydrogen-bonded complexes between larger molecules have been studied in the solid state by X-ray and neutron diffraction methods. The latter is the most useful since it allows the precise location of the hydrogen-bonded hydrogen in the X H · · · B moiety to be determined. For example, in the complex of pentachlorophenol with 2-methylpyridine [38] (Figure 4.2), the position of hydrogen shows that the molecular O H · · · N (and not the ionic O− · · · H N+ ) complex is formed. However, the very large elongation of the O H bond means that the incipient proton transfer has already reached an advanced stage. The hydrogen bond is not perfectly linear (OH · · · N = 167.5◦ ) and its directionality is governed by the axis of the Nsp2 lone pair (HNC4 = 172.7◦ ). When several HBA sites are present in the molecule, the structure shows which has the highest hydrogen-bond basicity. In the complex of 4,4 -bis(dimethylamino)benzophenone with pentafluorophenol [39] (FUYNEP), the phenol is hydrogen bonded to the carbonyl group and not to the nitrogen atom. In contrast, in the 1 : 1 complex with CF3 SO3 H, one of the two nitrogen atoms has been protonated by the acid. This illustrates that the hydrogenbonding site is not always the protonation site or, in other words, that there is no general connection between Brönsted and hydrogen-bond basicity scales. Other information about the structure of hydrogen-bonded complexes relies on crystal structures that, in many cases, involve molecules with both HBA and HBD functional groups. Thus, from the large number of structures in the CSD database (250 000 in October 2001) [40], it is possible to perform statistical surveys of hydrogen-bonded geometries directed to specific classes, for example the complexes with nitriles [41] or those with sulfoxides [42]. Statistical methods lead to averaged radial and angular parameters of the hydrogen bond. The statistical treatment is of vital importance because the hydrogen-bond geometry is easily deformed by other interactions in the crystal. If a sufficient number of structures are examined, chemically significant trends may be observed in the averaged data. For example, selenoxides are shown to be stronger HBAs than sulfoxides [42]. Also, the

116

Lewis Basicity and Affinity Scales Csp3 O

H

ECsp

3

O

H

Csp3 O

E

H

E

Csp3

Figure 4.3

Csp3 Csp3

From left to right, the elements E correspond to groups 17, 16 and 15, respectively.

conjugation of the nitrogen lone pair of amino groups with electron-withdrawing groups, such as cyano or carbonyl in the push–pull compounds N T C N [43] and N T C O [44] (where T is a group transmitting the conjugation), lowers the basicity of the amino nitrogen so dramatically that the carbonyl and cyano groups are the main hydrogen-bond acceptor sites and not the amino nitrogen. A search of the CSD has been performed [45] for the hydrogen-bonded complexes between OH donors and the alkyl derivatives of elements E of groups 15, 16 and 17. The various fragments searched for are shown in Figure 4.3. The author wished to study how basicity varies with position in the periodic table. The search for E = I, Te, As and Sb was done in vain, given the very low basicity of iodoalkanes, tellurides, arsines and stibines. One structure was found for F, Se and P. The number of OH · · · ClCsp3 , OH · · · BrCsp3 and OH · · · S(Csp3 )2 contacts is also limited. Only the OH · · · N(Csp3 )3 and, mostly, OH · · · O(Csp3 )2 contacts are very common. Averaged hydrogen-bond lengths and linearities are given in Table 4.2 and compared with data calculated in vacuo at the B3LYP/LANL2DZ level for the complexes of water and the methyl derivatives of elements E. A satisfactory agreement is found between averaged solid-state structural data and in vacuo calculated geometries, the correlation coefficient being 0.954 for the distances and 0.796 for the linearities. Correlations between HB length, inversely related to HB strength, and chemical descriptors of elements E, such as electronegativity, are clear-cut only within a period because, Table 4.2 Comparison between B3LYP/LANL2DZ-calculated geometries of HOH · · · EMen complexes in vacuo and corresponding solid-state geometric CSD data. θ/◦

d/A˚ HBA moiety C C C C C C C C a

F Cl Br O C S C Se C N(C) C P(C) C

Number of contacts. KOVCAZ. DIWNIG. d TOCHIC. b c

na

CSD

Calc.

CSD

Calc.

1b 10 5 880 23 1c 64 1d

2.02 2.61 2.70 1.91 2.53 2.65 1.83 2.30

1.99 2.54 2.79 1.87 2.44 2.57 1.88 2.52

152 137 148 161 134 147 164 151

148 136 131 175 156 152 172 171


117

3 Br

Se

Cl

2.5 d/Å

S

P 2

F O N

1.5 5

10

χ/eV Figure 4.4 Plot of the crystallographic HB length, d, versus the Mulliken electronegativity of the HB acceptor atom E in O H · · · E hydrogen bonds.

within a group, the radii of elements are too different. Figure 4.4 shows the increase in HB length, that is, the decrease in HB strength, with increasing electronegativity in a given period, in contradiction with Pauling’s statement [46] that ‘the strength of the [hydrogen] bond should increase with an increase in the electronegativity of the two bonded atoms’ (X and E in X H · · · E). The crystallographic orders of strength are amines > ethers > fluoroalkanes, phosphines > sulfides > chloroalkanes and selenides > bromoalkanes. Many other relationships between geometric parameters of the HB and basicity can be found in a book [47] and a review [48] on the hydrogen bond in the solid state.

4.2

Hydrogen-Bond Basicity Scales: Early Works

The first set of equilibrium constants for the hydrogen bonding of a large number of diverse Lewis bases to a given hydrogen-bond donor, in the same conditions of temperature and solvent, which should permit the construction of a Lewis basicity scale, refers to phenol at 20 ◦ C in CCl4 . The association constants between phenol and about 200 Lewis bases (reaction 4.3): C6 H5 OH + B C6 H5 OH · · · B

(4.3)

were determined mostly by Gramstad and co-workers [49–59] in Norway, mainly between 1960 and 1969, using a near-infrared (NIR) spectroscopic method. This method employs the first overtone of the stretching O H vibration of phenol at 7052 cm−1 (1418 nm).

118


We believe that these data are less reliable than a subsequent set referring to 4fluorophenol at 25 ◦ C in CCl4 (reaction 4.2), which were obtained jointly by the groups of Taft (University of California at Irvine) by a 19 F NMR method [7] and Schleyer (University of Princeton) by an IR method [10] between 1969 and 1972. The IR method gives more reliable results than the NIR method for the obvious reason that it employs the fundamental O H stretching vibration of 4-FC6 H4 OH at 3614 cm−1 , which has a molar absorption coefficient 84 times higher than that of the phenol overtone. Taking into account the different cell lengths in the NIR and IR methods (10 and 1 cm, respectively), the IR absorbances remain 8.4 times more sensitive to concentration changes upon complexation than the NIR absorbances. Confirmation is given [7] by the excellent agreement obtained between the values determined for 13 4-FC6 H4 OH/base complexes by the IR and 19 F NMR methods, and by a third, calorimetric, method. The equilibrium constants of reaction 4.2 are found in three papers. The first [10] collects 25 values, called K f , determined by the IR method, ranging from K f = 1.9 l mol−1 for diphenyl ether to 3600 l mol−1 for hexamethylphosphoric triamide. The second paper [7] reports 35 new K f values determined by the 19 F NMR method, ranging from 261 l mol−1 for tetramethylurea to 1.3 l mol−1 for diethyl sulfide. In the third paper [9], revised values are given for 19 bases and new values for 20 bases. The scale is extended up to K f = 4570 l mol−1 for triphenylarsine oxide. The construction, use and validity of a basicity scale is explicitly carried out in a fourth paper [8]. Taft and co-workers examined the possibility that logK f for hydrogen bonding by other OH hydrogen-bond donors to a variety of bases is linearly related to (logK f )0 for hydrogen bonding of the same bases to 4-fluorophenol. In other words, they studied the existence of linear free energy relationships in the field of hydrogen bonding. Application of Equation 4.4: logK f = m(logK f )0 + c

(4.4)

to ethanol, methanol, 1-naphthol, phenol, hexafluoro-2-propanol and 2,2,2-trifluoroethanol gave excellent results. The success of Equation 4.4 makes it useful to define a scale, called pK HB , which measures the relative base strength in hydrogen-bonded complex formation with OH HBDs. Because of the availability of extensive and reliable data, the standard reaction selected to define the pK HB value is the formation of the hydrogen-bonded complex of 4-FC6 H4 OH, CCl4 , 25 ◦ C, that is pK HB ≡ log(K f )0

(4.5)

In the absence of primary data, a secondary value may be obtained from Equation 4.6: pK HB =

logK f − c m

(4.6)

using literature data for one of the OH donors obeying Equation 4.4. In this manner, Taft and co-workers defined an extensive set of 140 pK HB values (80 primary and 60 secondary), ranging from −0.30 for benzoyl fluoride to 3.66 for triphenylarsine oxide. Few further determinations of equilibrium constants of the complexation of 4fluorophenol were reported between 1972 and 1988. Then, Laurence, Berthelot and coworkers (University of Nantes) began to extend the pK HB scale systematically to almost


119

all families of organic bases. The results published in a series of 33 papers (until 2005) [60–92], and also unpublished results from Laurence’s group, are presented below.

4.3 4.3.1

The 4-Fluorophenol Hydrogen-Bond Basicity Scale Definition

Let K c (Equation 4.7) be the equilibrium constant for the hydrogen bonding of 4fluorophenol to a variety of Lewis bases in CCl4 at 25 ◦ C, where the subscript c indicates that the concentration scale is used: K c /l mol−1 = C(complex)/C(base) × C(4-FC6 H4 OH)

(4.7)

Currently, the lowest and highest primary K c values are 0.14 l mol−1 for the very weak π base 2,3-dimethylbut-2-ene [75] and 120 000 l mol−1 for the tetrabutylammonium cyanate ion pair [93], respectively. A secondary K c value of 288 400 l mol−1 has been reached for trimethylamine N-oxide [94]. Since K c values vary over many orders of magnitude, it is more convenient to use a logarithmic scale. In the original symbol, pK HB , the subscript HB stands for ‘Hydrogen Bond’. It does not specify if pK HB is a hydrogen-bond basicity or acidity scale. This ambiguity can be cleared up by returning to a symbol similar to pK BH+ , which measures the basicity towards H+ and is defined as the antilog of the dissociation constant of the species BH+ . We propose to call pK BHX the scale measuring the basicity towards HX (where X stands for 4-FC6 H4 O) and to define pK BHX by Equation 4.8, as the antilog of the dissociation constant, 1/K c , of the complex BHX (B · · · HX): (4.8) pK BHX = − log10 1/ K c /1 l mol−1 = log10 K c /1 l mol−1 With these definitions, the strongest Brönsted base and the best hydrogen-bond base have the largest pK BH+ and pK BHX values, respectively. For example, pK BHX = −0.85 for 2,3dimethylbut-2-ene and 5.46 for trimethylamine N-oxide. Thus, at present, the Lewis basicity towards 4-fluorophenol extends over a range of 6.3 pK BHX units. This range corresponds to a Gibbs energy variation of 36 kJ mol−1 (Equations 4.9 and 4.10). δ∆G ◦ = −RT δ ln K c δ∆G ◦ /kJ mol−1 = −5.708 δpK BHX (at 298 K) 4.3.2

(4.9) (4.10)

Method of Determination

The K c values have been obtained by measuring the equilibrium concentration of 4fluorophenol from the absorbance of the ν(OH) infrared band at 3614 cm−1 by means of a Fourier transform infrared (FTIR) spectrometer. The well-known advantages of FTIR spectrometry over dispersive IR spectrometry, used in the early studies [9, 10], should provide more accurate absorbance measurements in a larger domain. Consequently, it is easier to adjust the initial concentrations of 4-fluorophenol and base in order to form the complex of 1 : 1 stoichiometry almost exclusively and avoid the self-association of 4-fluorophenol and the formation of complexes of higher stoichiometry. Examples of such side-effect

120

Lewis Basicity and Affinity Scales O

Ar

O

Ar

H N

H

O

H

Ar

Ar

O

O

O

H

N

H

N

H

Figure 4.5 FC6 H4 .

O Ar

Ar

Examples of 1(base) : 2(ArOH) hydrogen-bonded complexes. Ar stands for 4-

complexes are shown in Figure 4.5. Typically, the initial concentration of 4-fluorophenol is kept under the limit of 4 × 10−3 mol l−1 and the initial concentration of base is maintained in excess, so that 20–80% of 4-fluorophenol is hydrogen bonded. Moreover, using curvefitting software, absorbance corrections may be made when the OH band of the complex overlaps the key band of free 4-fluorophenol at 3614 cm−1 . Hence there is still only one unknown, K c , when FTIR deconvoluted absorbances are treated. This is not the case with either the NMR method, the observed time-average 19 F NMR shift depending on two unknowns, K c and the limiting 19 F NMR shift, or the calorimetric method, where the evolved heats depend on both K c and the enthalpy of complexation. From separate determinations by various operators on different FTIR spectrometers under varied experimental conditions, pK BHX values are believed to be accurate within ±0.02–0.05. For the 56 compounds common to the FTIR pK BHX and the IR, NMR or calorimetric pK HB sets, 29 values do not differ by more than ±0.05 (the maximum FTIR experimental error) and 11 values by no more than ±0.10 (the added assumed maximum experimental errors). Significant pK HB − pK BHX differences are found for Ph3 AsO (−0.49), Ph2 O (+0.33), (PhCH2 )3 N (−0.27), PhOMe (+0.25), N-methylimidazole (−0.22), Nmethylpyridone (−0.19), (HC CCH2 )3 N (−0.17) and Et2 S (−0.14). For the reasons given above, we recommend choosing the pK BHX value. 4.3.3

Polyfunctional Hydrogen-Bond Acceptors

If a molecule (polybase) contains n HBA sites, n 1 : 1 complexes are formed through equilibria of individual constants K c (i) (Equation 4.11): K c (i) = C(complex i) /C(4-FC6 H4 OH) × C(polybase)

(4.11)

In the usual IR method, the equilibrium molar concentration of hydrogen-bonded 4fluorophenol is measured, that is, the sum C(t) of the molar concentrations C(complex i) of each 1 : 1 complex (Equation 4.12): C(t) =

i

C(complex i)

(4.12)


121

N O

N

N

N

Cl

Me N

Scheme 4.2 Cotinine, pyrazine and 3-chloroquinuclidine, from left to right.

Therefore, the observed equilibrium constant is the ratio given in Equation 4.13: C(complex i)/C(4-FC6 H4 OH) × C(polybase)

(4.13)

i

This ratio is evidently related to the individual constants through Equation 4.14: K c (i) K c (t) =

(4.14)

i

which shows that the observed constant K c (t) is the sum of the individual constants (t stands for total). Since the logarithm of a sum is not equal to the sum of logarithms, logK c (t) is devoid of thermodynamic interest. It is therefore highly desirable to achieve the determination of the various K c (i)s. A number of methods have been proposed for separating K c (t) into individual K c (i)s. They have been applied to the following polybases: progesterone [84], nicotine [95], lindane [88], cotinine [89] and arylamines [90]. For example, the total equilibrium constant for the hydrogen bonding of 4-fluorophenol to cotinine (Scheme 4.2), K c (t) = 197 ± 14 l mol−1 , has been separated into the constants for the carbonyl oxygen, K c (O) = 145 ± 10 l mol−1 , and for the pyridine nitrogen, K c (N) = 42 ± 12 l mol−1 . Thus, on the pK BHX scale, the hydrogen-bond basicity of cotinine must be described by two values, pK BHX = 2.16 and pK BHX = 1.62, for the oxygen and nitrogen sites, respectively. In the case of polybases with n equivalent basic sites, it is easy to put logK c (t) on the pK BHX scale by a −logn statistical correction. For example, the total complexation constant of pyrazine (Scheme 4.2) is 16.6 l mol−1 . The statistically corrected quantity logK c (t) −log2 = 0.92 refers to the hydrogen-bond basicity per nitrogen atom and can be used for comparison with monobasic pyridines [80]. Many polybases have sites of such different HBA strengths that logK c (t) can often be considered as the pK BHX of the most basic site. For example, 3-chloroquinuclidine (Scheme 4.2) has two potential HBA sites, the nitrogen and the chlorine atom, but, considering the value K c = 0.5 l mol−1 of chlorocyclohexane, the contribution of the chlorine atom to K c (t) = 99.3 l mol−1 can be neglected. One can safely attribute the hydrogen-bond basicity to the nitrogen atom only and assign a pK BHX value of log99.3 = 1.97 [86] to the molecule. 4.3.4

Data

A collection of about 1000 pK BHX values, representing the most comprehensive Lewis basicity scale so far known, is given in Tables 4.3–4.26. To make the data presentation

Formula C16 H10 a C14 H10 a Ph3 CH C6 H6 C10 H8 a C6 H5 C6 H5 Ph2 CH2 PhSiMe3 PhMe PhEt Ph-i-Pr Ph-c-Hex Ph-t-Bu C6 H4 Me2 C16 H20 a C6 H4 Me2 C6 H4 Me2 C6 H3 t-Bu3 C6 H3 Me3 C6 H3 -i-Pr3 C6 H2 Me4 C6 H2 Me4 C6 HMe5 C6 Me6 C13 H21 Na C14 H23 Na C4 H4 Sa

Lewis base

Arenes Pyrene Phenanthrene Triphenylmethane Benzene Naphthalene Biphenyl Diphenylmethane Trimethylphenylsilane Methylbenzene, toluene Ethylbenzene Isopropylbenzene, cumene Cyclohexylbenzene tert-Butybenzene 1,4-Dimethylbenzene, p-xylene 1-Phenyladamantane 1,2-Dimethylbenzene, o-xylene 1,3-Dimethylbenzene, m-xylene 1,3,5-Tri-tert-butylbenzene 1,3,5-Trimethylbenzene, mesitylene 1,3,5-Triisopropylbenzene 1,2,4,5-Tetramethylbenzene, durene 1,2,3,4-Tetramethylbenzene, prehnitene Pentamethylbenzene Hexamethylbenzene

Hindered pyridines 2,6-Di-tert-butylpyridine 2,6-Di-tert-butyl-4-methylpyridine Thiophene

No.

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24

25 26 27

0.55 0.68 0.78

0.74 0.69 0.95

Kc (t)

−0.54 −0.45 −0.50

−0.56b −0.53b −0.50b −0.49 −0.48b −0.47b −0.41b −0.36 −0.36 −0.36 −0.34 −0.32 −0.32 −0.29 −0.28 −0.27 −0.28 −0.23 −0.18 −0.18 −0.15 −0.14 −0.07 0.02

pKBHX

3.08 2.57 2.85

3.20 3.03 2.85 2.80 2.74 2.68 2.34 2.05 2.05 2.05 1.94 1.83 1.83 1.66 1.60 1.54 1.60 1.31 1.03 1.03 0.86 0.80 0.40 −0.11

∆G◦

Table 4.3 Equilibrium constants Kc (t) (l mol−1 ) and hydrogen-bond basicity scales, pKBHX and ∆G◦ (kJ mol−1 ), for the complexes of 4-fluorophenol with carbon π bases, in CCl4 at 25 ◦ C.

[77] [77] [97]

[75] [75] [75] [75] [75] [75] [75] [75] [75] [75] [75] [75] [75] [75] [75] [75] [75] [75] [75] [75] [75] [75] [75] [75]

Ref.

2

5

15

25

tBu But

26

N

Me

tBu

PhCH CH2 PhC(Me) CH2 PhCH CHMe

PhC CMe

27

S

b

Statistically corrected, value per π sextet. c Statistically corrected, value per double bond. d Statistically corrected, value per triple bond. e Polybases. f The total equilibrium constant has been decomposed into the contributions of the triple bond and the aromatic ring.

1

N

Vinylbenzene, styrene α-Methylstyrene trans-β-Methylstyrene

39 40 41

But

1-Phenyl-1-propyne

38

Formulae:

PhC CH

Arylacetylenes, styrenese Ethynylbenzene, phenylacetylene

37

a

HC BuC EtC PrC

Alkynes 1,5-Hexadiyne 1-Hexyne 3-Hexyne 2-Hexyne

33 34 35 36 CCH2 CH2 C CH CH CEt CMe

C6 H8 Me2 C CMe2 C6 H10 C7 H12 Pent CH CH2

Alkenes 1,4-Cyclohexadiene 2,3-Dimethyl-2-butene Cyclohexene 1-Methyl-1-cyclohexene 1-Heptene

28 29 30 31 32

0.44 0.48 0.55

0.91

0.60

0.74

0.26

−0.45 (C C)f −0.60 (Ph)f −0.19 (C C)f −0.57 (Ph)f

−0.43d −0.22 −0.10 −0.09

−0.88c −0.85 −0.82 −0.74 −0.67

2.57 3.42 1.08 3.25

2.45 1.26 0.57 0.51

5.02 4.85 4.68 4.22 3.82

[75] [75] [75] [75] [75] [75] [75]

[75] [75] [75] [96]

[75] [75] [75] [75] [75]

124


Table 4.4 Equilibrium constants Kc (t) (l mol−1 ) and hydrogen-bond basicity scales, pKBHX and ∆G◦ (kJ mol−1 ), for the complexes of 4-fluorophenol with ammonia and primary amines, in C2 Cl4 a at 25 ◦ C. No. Lewis base

Formula

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32

CF3 CH2 NH2 3-NO2 C6 H4 CH2 NH2 3,5-Cl2 C6 H3 CH2 NH2 N CCH2 CH2 NH2 3,5-F2 C6 H3 CH2 NH2 3-CF3 C6 H4 CH2 NH2 3-ClC6 H4 CH2 NH2 HC CCH2 NH2 3-FC6 H4 CH2 NH2 NH3 c-PrNH2 C6 H5 CH2 NH2 H2 C CHCH2 NH2 3-MeOC6 H4 CH2 NH2 3-MeC6 H4 CH2 NH2 PhCH2 CH2 NH2 EtNH2 n-PrNH2 MeNH2 n-BuNH2 H2 N(CH2 )4 NH2 H2 N(CH2 )6 NH2 i-PrNH2 H2 NCH2 CH2 NH2 MeO(CH2 )3 NH2 CH3 (CH2 )15 NH2 t-BuNH2 n-OctNH2 MeOCH2 CH2 NH2 c-HexNH2 1-Adam-NH2 H2 N(CH2 )3 NH2

2,2,2-Trifluoroethylamine 3-Nitrobenzylamine 3,5-Dichlorobenzylamine 3-Aminopropionitrile 3,5-Difluorobenzylamine 3-(Trifluoromethyl)benzylamine 3-Chlorobenzylamine Propargylamine 3-Fluorobenzylamine Ammonia (in CCl4 ) Cyclopropylamine Benzylamine Allylamine 3-Methoxybenzylamine 3-Methylbenzylamine Phenethylamine Ethylamine n-Propylamine Methylamine n-Butylamine 1,4-Diaminobutane 1,6-Diaminohexane Isopropylamine Ethylenediamine 3-Methoxypropylamine n-Hexadecylamine tert-Butylamine n-Octylamine 2-Methoxyethylamine Cyclohexylamine Adamantane-1-amine 1,3-Diaminopropane

Kc (t) pKBHX

26

325 323 354 182

196 413

0.71 1.26 1.27 1.28b 1.28 1.43 1.55 1.57 1.58 1.74 1.74 1.88 1.94 1.94 1.97 2.16 2.17 2.19 2.20 2.21 2.21c 2.21c 2.22 2.25c 2.25b 2.26 2.26 2.27 2.28b 2.29 2.31 2.31c

∆G◦

Ref.

−4.05 −7.19 −7.25 −7.31 −7.31 −8.16 −8.85 −8.96 −9.02 −9.93 −9.93 −10.73 −11.07 −11.07 −11.24 −12.33 −12.39 −12.50 −12.56 −12.61 −12.61 −12.61 −12.67 −12.84 −12.84 −12.90 −12.90 −12.96 −13.01 −13.07 −13.19 −13.19

[98] [99] [99] [82] [99] [99] [99] [82] [99] [98] [82] [82] [82] [99] [99] [82] [82] [82] [98] [82] [82] [82] [82] [82] [82] [82] [82] [82] [82] [82] [82] [82]

a C2 Cl4 was preferred to the definition solvent CCl4 because a few amines precipitate in CCl4 solutions. The equation logKc (in C2 Cl4 ) = 1.005pKBHX (in CCl4 ) + 0.002 (n = 12, r = 0.995, s = 0.05) correlates 12 values obtained both in CCl4 and C2 Cl4 . It shows that this solvent does not cause any variation in pKBHX greater than the experimental error. Consequently, for the sake of simplicity, the same symbol, pKBHX , has been retained for logKc (in C2 Cl4 ), whether the measurements were made in CCl4 or C2 Cl4 . b Value corrected for the presence of a second hydrogen-bond acceptor group. c Statistically corrected to put the pKBHX value on a per nitrogen basis.

Diallylamine Morpholine

N-Methylbenzylamine (in C2 Cl4 ) 2,2,6,6-Tetramethylpiperidine 2-Phenylpyrrolidine Diisopropylamine N-Methylallylamine 1,2,3,4-Tetrahydroisoquinoline (in C2 Cl4 ) Di-n-butylamine Piperazine N-Methylphenethylamine 1,2,3,6-Tetrahydropyridine N-Methylisopropylamine

11 12

13 14 15 16 17 18 19 20 21 22 23

Bis(2-chloroethyl)amine Dibenzylamine 3-(Methylamino)propionitrile

4 5 6

2-(3-Trifluoromethylphenyl)pyrrolidine N-Methylpropargylamine 2-(3-Fluorophenyl)pyrrolidine Thiomorpholine

Thiazolidine

3

7 8 9 10

1,1,1,3,3,3-Hexamethyldisilazane (Methylamino)acetonitrile

Lewis base

1 2

No.

i-PrNHMe

b

PhCH2 CH2 NHMe

b

n-Bu2 NH

b

i-Pr2 NH H2 C CHCH2 NHMe

b

PhCH2 NHMe

b

(H2 C CHCH2 )2 NH

b

b

b

HC CCH2 NHMe

b

(ClCH2 CH2 )2 NH (PhCH2 )2 NH N CCH2 CH2 NHMe

b

(Me3 Si)2 NH N CCH2 NHMe

Formula

261

73

48

32

15

11

Kc (t)

∆G◦ 2.57 −3.77 −4.51 −6.28 −1.83 −6.79 −7.65 −7.82 −5.31 −7.88 −9.36 −9.42 −9.53 −1.94 −9.70 −10.16 −6.28 −10.39 −10.73 −11.02 −11.42 −11.42 −11.64 −12.04 −12.04 −12.22 −12.33 −12.56

pKBHX −0.45 0.66 (N)c 0.79 (C N) 1.10 (N)c 0.32 (S) 1.19 1.34 1.37 (N)c 0.93 (C N) 1.38 1.64 1.65 1.67 (N)c 0.34 (S) 1.70 1.78 (N)c 1.10 (O) 1.82 1.88 1.93 2.00 2.00 2.04 2.11 2.11d 2.14 2.16 2.20

[85] [85] [85] [85] [85] [85] [85] [85] [85] [95] [85] [85] [85] [85] [85] [85] [85] [85] [85] [95] [85] [85] [85] [85] [85] [85] [99] [85]

Ref.

(Continued)

Table 4.5 Equilibrium constants Kc (t) (l mol−1 ) and hydrogen-bond basicity scales, pKBHX and ∆G◦ (kJ mol−1 ), for the complexes of 4-fluorophenol with secondary amines, in CCl4 or C2 Cl4 a at 25 ◦ C.

H

27

31

N

22

NH

7 X: 3-CF3 9 X: 3-F 15 X: H

H

7, 9, 15

H

N

CHX

N

3

H

N

S S

10

34

NH

NH

O

12

b

b

b

35

H

N

NH

14

H

N

36

NH

MeNHCH2 CH2 NHMe (MeOCH2 CH2 )2 NH

b

EtNHMe Et2 NH Me2 NH

b

t-BuNHMe c-HexNHMe n-BuNHMe

Formula

Value corrected for the presence of a second hydrogen-bond acceptor site. d Statistically corrected to put the pKBHX value on a per nitrogen basis.

c

b

See footnote a in Table 4.4. Formulae:

N-Methyl-tert-butylamine N-Methylcyclohexylamine N-Methylbutylamine Azepane, hexamethylenimine N-Methylethylamine Diethylamine Dimethylamine (in C2 Cl4 ) 2-Methylaziridine N,N -Dimethylethylenediamine Bis(2-methoxyethyl)amine Piperidine Pyrrolidine (in C2 Cl4 ) Azetidine

24 25 26 27 28 29 30 31 32 33 34 35 36

a

Lewis base

No.

Table 4.5 (Continued)

18

NH

394 225

Kc (t) 2.21 2.24 2.24 2.24 2.25 2.25 2.26 2.28 2.29d 2.31c 2.38 2.59 2.59

pKBHX

Ref. [85] [85] [85] [85] [85] [85] [85] [85] [85] [85] [85] [85] [85]

∆G◦ −12.61 −12.79 −12.79 −12.79 −12.84 −12.84 −12.90 −13.01 −13.07 −13.19 −13.59 −14.78 −14.78

Tripropargylamine N-Methyl-2-(3-trifluoromethyl)pyrrolidine N-Ethyl-N-methyl-3-(trifluoromethyl)benzylamine N,N-Diisopropylethylamine N-Methyl-2-(3-fluorophenyl)pyrrolidine N,N-Dicyclohexylethylamine N-Ethyl-N-methyl-3-fluorobenzylamine 3-(Dimethylamino)propionitrile

N,N-Dimethyl-3-(trifluoromethyl)benzylamine 1,2,2,6,6-Pentamethylpiperidine N,N-Dimethyl-3-fluorobenzylamine Hexamethylenetetramine Triallylamine N-Methyl-2-phenylpyrrolidine N-(2-Chloroethyl)pyrrolidine Tri-n-propylamine N-Ethyl-N-methylbenzylamine N,N-Dimethyl-3-chloropropylamine Tri-n-butylamine N-Methylmorpholine

Tri-n-octylamine 1,3,5-Trimethylhexahydrotriazine N,N-Dimethylbenzylamine N,N-Dimethylpropargylamine 4-Chloro-N-methylpiperidine 3-Chloromethyl-N-methylpiperidine

13 14 15 16 17 18 19 20 21 22 23 24

25 26 27 28 29 30

Tribenzylamine N,N-Diisopropyl-3-pentylamine N,N-Diisopropylisobutylamine N,N-Dimethylaminoacetonitrile

Lewis base

5 6 7 8 9 10 11 12

1 2 3 4

No.

a

a

PhCH2 NMe2 HC CCH2 NMe2

a

n-Oct3 N

a

n-Pr3 N PhCH2 N(Me)Et ClCH2 CH2 CH2 NMe2 n-Bu3 N

a a

(H2 C CHCH2 )3 N

a

3-FC6 H4 CH2 NMe2

a

3-CF3 C6 H4 CH2 NMe2

c-Hex2 NEt 3-FC6 H4 CH2 N(Me)Et N CCH2 CH2 NMe2

a

3-CF3 C6 H4 CH2 N(Me)Et i-Pr2 NEt

(HC CCH2 )3 N

a

(PhCH2 )3 N i-Pr2 NCHEt2 i-Pr2 N-i-Bu N CCH2 NMe2

Formula

115

46

85

23

7.5

10.4

1.3

Kc (t)

∆G◦ 8.56 1.94 −1.71 −3.77 −4.34 −4.74 −5.25 −5.88 −5.99 −6.22 −6.51 −6.51 −6.56 −5.25 −6.62 −7.02 −7.25 −7.59 −7.65 −7.88 −8.28 −8.39 −8.39 −8.79 −8.85 −8.90 −5.48 −8.96 −9.02 −9.08 −9.13 −9.70 −9.93

pKBHX ∼−1.5(N)b −0.34 0.30 0.66 (N)b 0.76 (C N) 0.83 (N)b 0.92 1.03 1.05 1.09 1.14 1.14 1.15 (N)b 0.92 (C N) 1.16 1.23 1.27 1.33c 1.34 1.38 1.45 1.47 1.47 1.54 1.55 1.56 (N)b 0.96 (O) 1.57 1.58c 1.59 1.60 1.70 1.74

[86] [86] [86] [86] [86] [86] [95] [99] [86] [95] [86] [99] [86] [86] [99] [86] [99] [86] [86] [95] [86] [86] [99] [86] [86] [86] [86] [86] [86] [86] [86] [86] [86]

Ref

(Continued)

Table 4.6 Equilibrium constants Kc (t) (l mol−1 ) and hydrogen-bond basicity scales, pKBHX and ∆G◦ (kJ mol−1 ), for the complexes of 4-fluorophenol with tertiary amines, in CCl4 at 25 ◦ C.

48

N

31

NMe

NMe

49

N

32

NMe

6 X: 3-CF3 9 X: 3-F 18 X: H

Cl

Me

34

N

14

N

N

N

16

36

Me

N N

b Value corrected for the presence of a second hydrogen-bond acceptor c Statistically corrected to put the pK BHX value on a per nitrogen basis.

Ph

Me

N

CHX

site.

38

Bu

N

19

CH CH Cl

N

N-Methyl-1,2,3,4-tetrahydroisoquinoline N,N’-Dimethylpiperazine N,N-Dimethylallylamine 3-Chloroquinuclidine Triethylamine N-Methyl-1,2,3,6-tetrahydropyridine N,N,N ,N -Tetramethylethylenediamine N-Butylpyrrolidine N,N,N ,N -Tetramethylhexane-1,6-diamine N,N-Dimethylisopropylamine N-Methylpiperidine Trimethylamine N,N-Dimethylcyclohexylamine N,N-Dimethylethylamine N-Methylpyrrolidine Diazabicyclooctane (DABCO) Tropane 4-Phenylquinuclidine Quinuclidine

31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49

a Formulae:

Lewis base

(Continued)

No.

Table 4.6

41

O

NMe

24

a a

a a a

NMe

N

45

Me

MeN

NMe3 c-HexNMe2 EtNMe2

a

N

26

Me

N

46

NMe

Cl

29 N

47

Me N

NMe

429

225

Me2 N(CH2 )6 NMe2 i-PrNMe2

a

211

150

Kc (t)

Me2 NCH2 CH2 NMe2

a

NEt3

a

H2 C CHCH2 NMe2

a a

Formula 1.80 1.88c 1.92 1.97 1.98 2.02 2.02c 2.04 2.05c 2.11 2.11 2.13 2.15 2.17 2.19 2.33c 2.39 2.46 2.71

pKBHX

Ref [86] [86] [86] [86] [86] [86] [86] [86] [86] [86] [86] [86] [86] [86] [86] [86] [86] [86] [86]

∆G◦ −10.27 −10.73 −10.96 −11.24 −11.30 −11.53 −11.53 −11.64 −11.70 −12.04 −12.04 −12.16 −12.27 −12.39 −12.50 −13.30 −13.64 −14.04 −15.47


129

Table 4.7 Total equilibrium constants, Kc (t) (l mol−1 ), and hydrogen-bond basicity scales, pKBHX and ∆G◦ (kJ mol−1 ), for the π and N complexes of 4-fluorophenol with arylamines, in CCl4 , at 25 ◦ C [90]. No.

Lewis base

Formula

Kc (t)

1

Primary arylamines 3-Chloroaniline

3-ClC6 H4 NH2

1.7

2

3-Fluoroaniline

3-FC6 H4 NH2

2.0

3

Aniline

C6 H5 NH2

3.5

4

4-Toluidine

4-MeC6 H4 NH2

4.3

5

2,6-Dimethylaniline

2,6-Me2 C6 H3 NH2

3.8

6

2,6-Diethylaniline

2,6-Et2 C6 H3 NH2

3.4

7

2,6-Diisopropylaniline

2,6-i-Pr2 C6 H3 NH2

3.3

8

2,4,6-Tri-tert-butylaniline

2,4,6-t-Bu3 C6 H2 NH2

1.8

9

Secondary arylamines 4-Chloro-N-methylaniline

4-ClC6 H4 NHMe

1.5

10

Diphenylamine

Ph2 NH

1.1

11

N-Methylaniline

C6 H5 NHMe

2.5

12

2,3-Dihydroindole

a

6.9

13

N-Methyl-4-toluidine

4-MeC6 H4 NHMe

3.5

14

1,2,3,4-Tetrahydroquinoline

a

5.8

15

Tertiary arylamines 4-Bromo-N,N-dimethylaniline

4-BrC6 H4 NMe2

1.9

16

N-Phenylpiperidine

c-(CH2 )5 NC6 H5

5.3

17

¨ Troger’s base

a

29.1

18

Proton sponge

a

1.4

19

N,N-Dimethylaniline

C6 H5 NMe2

3.2

20

N-Phenylpyrrolidine

c-(CH2 )4 NC6 H5

2.3

21

N,N-Dimethyl-4-toluidine

4-MeC6 H4 NMe2

5.7

22

N,N-Dimethyl-3-toluidine

3-MeC6 H4 NMe2

3.5

pKBHX (π ) pKBHX (N)

∆G◦

−0.43 0.13 −0.38 0.20 −0.26 0.46 −0.18 0.56 −0.08 0.47 −0.03 0.39 −0.02 0.37 —

2.45 −0.74 2.17 −1.14 1.48 −2.63 1.03 −3.20 0.46 −2.68 0.17 −2.23 0.11 −2.11

−0.41 0.05 0.30b −1.05 −0.16 0.26 −0.13 0.79 −0.10 0.43 −0.09 0.70

2.34 −0.29 1.71 5.99 0.91 −1.48 0.74 −4.51 0.57 −2.45 0.51 −4.00

−0.42 0.17 −0.27 0.68 −0.26b 1.15c −0.24b −0.55c −0.15 0.39 −0.08 0.16 −0.07 0.69 −0.05 0.41

2.40 −0.97 1.54 −3.88 1.48 −6.56 1.37 3.14 0.86 −2.23 0.46 −0.91 0.40 −3.94 0.29 −2.34

130



Lewis base

Formula

Kc (t)

pKBHX (π ) pKBHX (N)

23

N,N,N ,N -Tetramethylbenzene-1, 4-diamine

4-Me2 NC6 H4 NMe2

27.7

−0.02

0.11

24

Julolidine

a

3.5

25

N,N-Diethylaniline

C6 H5 NEt2

1.9

−6.45 −0.11 −2.23 0.57 −0.29

26

Triphenylamine

N(C6 H5 )3

1.0

1.13c 0.02 0.39 −0.10 0.05 —

a

Formulae: N

Me

Me Me

N

N

N

14

17

18

Statistically corrected to put the pKBHX (π ) value on a per π sextet basis. c Statistically corrected to put the pKBHX (N) value on a per nitrogen basis. b

Me

N

H

H

12

Me

Me N

N

24

∆G◦

3,5-Dichloropyridine Pyrazine 4-Cyanopyridine

7 8 9

2-Fluoropyridine 2-Bromopyridine 2-Chloropyridine Pyrimidine 2-N,N-Dimethylaminopyrimidine 2,2 -Bipyridine 7,8-Benzoquinoline Phenazine 2,5-Dimethylpyrazine 3-Bromopyridine 3-Chloropyridine 3-Fluoropyridine 3-Iodopyridine

5-Bromopyrimidine 3-Cyanopyridine

5 6

10 11 12 13 14 15 16 17 18 19 20 21 22

Pentafluoropyridine 2,6-Difluoropyridine 1,3,5-Triazine 2-Cyanopyridine

Lewis base

1 2 3 4

No.

Me2 C4 H2 N2 3-BrC5 H4 N 3-ClC5 H4 N 3-FC5 H4 N 3-IC5 H4 N

a

a

2-Me2 NC4 H3 N2 (C5 H4 N)2

a

2-FC5 H4 N 2-BrC5 H4 N 2-ClC5 H4 N

4-N CC5 H4 N

3,5-Cl2 C5 H3 N

a

BrC4 H3 N2 3-N CC5 H4 N

2-N CC5 H4 N

a

C5 F5 N 2,6-F2 C5 H3 N

Formula

33 39

23 25 28

17 11

8 10

6 7

Kc (t)

∆G◦ 2.80 −0.80 −1.83 −2.74 −3.48 −3.37 −4.68 −3.03 −4.85 −5.25 −5.25 −2.68 −5.42 −5.88 −5.99 −6.11 −6.28 −6.56 −6.62 −6.96 −7.36 −7.48 −7.48 −7.71 −7.82

pKBHX ∼−0.49 0.14 0.32c 0.48 (N)d 0.61 (C N) 0.59c 0.82 (N)d 0.53 (C N) 0.85 0.92c 0.92 (N)d 0.47 (C N) 0.95 1.03 1.05 1.07c 1.10b 1.15c 1.16 1.22c 1.29c 1.31 1.31 1.35 1.37

[80] [80] [80] [80] [80] [80] [80] [80] [80] [80] [80] [80] [80] [80] [80] [80] [80] [80] [80] [80] [99] [80] [80] [80] [80]

Ref.

(Continued)

Table 4.8 Equilibrium constants, Kc (t) (l mol−1 ), and hydrogen-bond basicity scales, pKBHX and ∆G◦ (kJ mol−1 ), for the complexes of 4-fluorophenol with six-membered N-heteroaromatics, in CCl4 at 25 ◦ C.

Lewis base

4-Acetylpyridine

3-Benzoylpyridine

2-tert-Butylpyridine 2-Phenylpyridine Methyl nicotinate

4,6-Dimethylpyrimidine 4-Chloropyridine 2-Aminopyrimidine 5-N,N-Dimethylaminopyrimidine 2-N,N-Dimethylaminopyridine 2-Vinylpyridine Pyridazine 2-Isopropylpyridine Pyridine Phenanthridine 2-n-Butylpyridine Quinoline Isoquinoline 2-Ethylpyridine Acridine 4-Vinylpyridine 4-Phenylpyridine Phthalazine 3-Methylpyridine, 3-picoline 3-Ethylpyridine 2-Methylpyridine, 2-picoline 4-Methylpyridine, 4-picoline

No.

23

24

25 26 27

28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49


3-MeC5 H4 N 3-EtC5 H4 N 2-MeC5 H4 N 4-MeC5 H4 N

a

4-CH2 CHC5 H4 N 4-PhC5 H4 N

a

2-EtC5 H4 N

a

a

2-n-BuC5 H4 N

a

2-i-PrC5 H4 N C5 H5 N

a

Me2 C4 H2 N2 4-ClC5 H4 N 2-H2 NC4 H3 N2 5-Me2 NC4 H3 N2 2-Me2 NC5 H4 N 2-CH2 CHC5 H4 N

2-t-BuC5 H4 N 2-PhC5 H4 N 3-COOMeC5 H4 N

3-COPhC5 H4 N

4-CH3 COC5 H4 N

Formula

186

89

71 76

59

31

31

32

Kc (t) 1.41(N)d 0.78 (O) 1.41(N)d 0.72 (O) 1.42 1.43 1.45 (N)d 0.50 (O) 1.47c 1.54 1.55c 1.58b 1.61 1.65 1.65c 1.76 1.86 1.87 1.88 1.89 1.94 1.94 1.95 1.95 1.96 1.97c 2.00 2.01 2.03 2.07

pKBHX

Ref. [80] [80] [80] [80] [80] [80] [80] [80] [99] [80] [80] [80] [80] [80] [80] [80] [80] [80] [7] [80] [80] [80] [80] [80] [80] [80] [80] [80] [80] [80]

∆G◦ −8.05 −4.45 −8.05 −4.11 −8.11 −8.16 −8.28 −2.85 −8.39 −8.79 −8.85 −9.02 −9.19 −9.42 −9.42 −10.05 −10.62 −10.67 −10.73 −10.79 −11.07 −11.07 −11.13 −11.13 −11.19 −11.24 −11.42 −11.47 −11.59 −11.82

4-Ethylpyridine 4-tert-Butylpyridine 2-Methylaminopyridine 2-Aminopyridine 4-Methoxypyridine 2,6-Dimethylpyridine, 2,6-lutidine 3-Aminopyridine 3,5-Dimethylpyridine, 3,5-lutidine 2,4-Dimethylpyridine, 2,4-lutidine 3,4-Dimethylpyridine, 3,4-lutidine 2,4,6-Trimethylpyridine, 2,4,6-collidine 9-N,N-Dimethylaminoacridine 3-N,N-Dimethylaminopyridine 4-N,N-Dimethylaminoquinoline 4-Aminopyridine N-Methyl-N-pyridin-4-ylhydrazine 4-(4-Methylpiperidino)pyridine 4-Piperidinopyridine 4-Methylaminopyridine 4-N,N-Dimethylaminopyridine 4-N,N-Diethylaminopyridine 4-Pyrrolidinopyridine 4-N,N-Dimethylaminopyrimidine 1,7-Phenanthroline Quinazoline 2-Methoxypyridine

Substituted 1,10-phenanthrolines 1,10-Phenanthroline

50 51 52 53 54 55 56 57 58 59 60 61 62 63 64 65 66 67 68 69 70 71 72 73 74 75

76

a

2-MeOC5 H4 N

a

a

4-(CH2 )5 NC5 H4 N 4-MeNHC6 H4 N 4-Me2 NC5 H4 N 4-Et2 NC5 H4 N 4-(CH2 )4 NC5 H4 N 4-Me2 NC4 H3 N2

a

a

4-H2 NC5 H4 N

a

3-Me2 NC5 H4 N

a

4-EtC5 H4 N 4-t-BuC5 H4 N 2-H(Me)NC5 H4 N 2-H2 NC5 H4 N 4-MeOC5 H4 N Me2 C5 H3 N 3-H2 NC5 H4 N Me2 C5 H3 N Me2 C5 H3 N Me2 C5 H3 N Me3 C5 H2 N

1148e

129b 74 35 10

3.06

2.07 2.11 2.11 2.12 2.13 2.14 2.20 2.21 2.21 2.24 2.29 2.31b 2.43b 2.43b 2.56 2.58 2.68 2.68 2.69 2.80 2.89 2.93 — — — — −17.47

−11.82 −12.04 −12.04 −12.10 −12.16 −12.22 −12.56 −12.61 −12.61 −12.79 −13.07 −13.19 −13.87 −13.87 −14.61 −14.73 −15.30 −15.30 −15.35 −15.98 −16.50 −16.72

(Continued)

[94]

[80] [80] [80] [80] [9] [80] [80] [80] [80] [80] [80] [80] [80] [80] [80] [80] [80] [80] [96] [80] [80] [80] [80] [80] [80] [80]

40

66

N

8

3

N

N

N

N

N

Me

N

N

73

42

N

13

N

N

16

74

45

N

N

N

N

N 3 2

4

1

N

N

5

76

61

NMe

17

N

N

6

10

N 9

7 8

34

N

N

NMe

63

N

C24 H16 N2 C14 H12 N2 C24 H16 N2 C16 H16 N2 C26 H20 N2 C20 H22 N4

Formula

N

37

1820e 2188e 2570e 2884e 4074e 7943e

Kc (t) 3.26 3.34 3.41b 3.46 3.61b 3.90b

pKBHX

Ref. [94] [94] [94] [94] [94] [94]

∆G◦ −18.61 −19.06 −19.46 −19.75 −20.61 −22.26

d

c

Secondary value. Statistically corrected to put the pKBHX value on a per nitrogen basis. Value corrected for the presence of a second hydrogen-bond acceptor group. e In spite of the presence of two nitrogen atoms, no statistical correction is made because there is only one 1 : 1 three-centre hydrogen-bonded complex [94].

b

N

N

Formulae:

4,7-Diphenyl-1,10-phenanthroline 4,7-Dimethyl-1,10-phenanthroline 2,9-Diphenyl-1,10-phenanthroline 3,4,7,8-Tetramethyl-1,10-phenanthroline 2,9-Dimethyl-4,7-diphenyl-1,10-phenanthroline 4,7-Dipyrrolidino-1,10-phenanthroline

77 78 79 80 81 82

a

Lewis base

No.



135

Table 4.9 Equilibrium constants, Kc (t) (l mol−1 ), and hydrogen-bond basicity scales, pKBHX and ∆G◦ (kJ mol−1 ), for the complexes of 4-fluorophenol with five-membered N-heteroaromatics, in CCl4 at 25 ◦ C. Formulaa

No. Lewis base 1 2 3 4

Pyrrole 1-Methylpyrrole 2,1,3-Benzothiadiazole 1-(Dimethylamino)pyrrole

C4 H4 NH C4 H4 NMe C6 H4 N2 S C4 H4 NNMe2

5 6 7 8 9 10 11 12

1,2-Benzoisoxazole Isoxazole 2-Phenylbenzoxazole Oxazole Benzothiazole Thiazole 2-Methylbenzoxazole 2,5-Dimethyl-1,3,4thiadiazole 1-Methylpyrazole 1-Acetylimidazole 2,4,5-Trimethyloxazole 5-Bromo-1methylimidazole 4-Butyl-1,2,4-triazole 1-Vinylimidazole Imidazole 1-Methylimidazole 1,5-Dicyclohexylimidazole

C7 H5 NO C3 H3 NO C13 H9 NO C3 H3 NO C7 H5 NS C3 H3 NS C8 H7 NO C4 H6 N2 S

13 14 15 16 17 18 19 20 21 a

Ref.

−0.86 −1.31 −1.43 −1.48 −3.88 −4.62 −6.74 −7.42 −7.71 −7.82 −8.90 −9.19

[14] [96] [96] [96] [96] [96] [21] [76] [21] [96] [21] [96] [96]

1.84 1.86 2.02 2.22

−10.50 −10.62 −11.53 −12.67

[96] [100] [96] [96]

2.29b 2.35 2.42d 2.72 3.12

−13.07 −13.41 −13.81 −15.53 −17.81

[96] [21] [96] [96] [96]

pKBHX

0.15 0.23 3.5 0.25b 3.5 0.26 (pyrrole)c 0.23 (NMe2 ) 0.68 0.81 1.18 1.30 1.35 1.37 1.56 81 1.61b

C4 H6 N2 C3 H3 N2 COMe C3 (Me)3 NO C3 H2 BrN2 Me C6 H11 N3 C3 H3 N2 CH CH2 C3 H3 N2 H C3 H3 N2 Me C15 H24 N2

389

Formulae: N

N H

1

N

N

N

O

3

5 N

N N

N

Me

R

13

N

N S

b

∆G◦

Kc (t)

14 18 19 20

O

6 R: COMe R: CH=CH2 R: H R: Me

N

N

Ph O

O

7

S

8 N Br

N Me

16

Statistically corrected to put the pKBHX value on a per nitrogen basis. c Value corrected for the presence of a second hydrogen-bond acceptor group. d Secondary value.

9 N

N

N Bu

17

S

10

N

N Me

S

12

Me

136


Table 4.10 Hydrogen-bond basicity scales, pKBHX and ∆G◦ (kJ mol−1 ), for the complexes of 4-fluorophenol with amidines, in CCl4 at 25 ◦ C. Lewis base: R1 R2 N C(R3 ) N R4 No.

R1

Formamidines (R3 = H) 1 Me 2 Me 3 Me 4 Me 5 Me 6 Me 7 Me 8 Me 9 Me 10 Me 11 Me 12 Me 13 Me 14 Me 15 Me 16 Me 17 Me 18 Me 19 Me 20 Me 21 Me 22 Me 23 Me 24 Me 25 Me 26 Me 27 Me 28 Me 29 Me 30 Me 31 Me 32 Me 33 Me 34 Me 35 Me

R2

R3

R4

Me Me Me Me Me Me Me Me Me Me Me Me Me Me Me Me Me Me Me Me Me Me Me Me Me Me Me Me Me Me Me Me Me Me Me

H H H H H H H H H H H H H H H H H H H H H H H H H H H H H H H H H H H

3,5-(NO2 )2 C6 H3 4-NO2 C6 H4 4-N CC6 H4 2-BrC6 H4 4-CF3 C6 H4 4-MeCOC6 H4 CH2 CF3 2-MeC6 H4 4-BrC6 H4 4-FC6 H4 C6 H5 3-CF3 C6 H4 CH2 3,5-Cl2 C6 H3 CH2 4-MeC6 H4 4-MeOC6 H4 3-ClC6 H4 CH2 4-ClC6 H4 CH2 CH2 CH2 C N t-Am CH2 C CH 4-Me2 NC6 H4 C6 H5 CH2 4-MeC6 H4 CH2 c-Pr 3-MeOC6 H4 CH2 t-Bu NMe2 CH2 CH CH2 i-Bu 1-Adam n-Pr c-Hex i-Pr Me CH2 CH2 OMe

Ph Ph Ph Ph Ph Ph Ph Ph

4-BrC6 H4 4-BrC6 H4 4-BrC6 H4 3-MeOC6 H4 4-MeC6 H4 4-MeC6 H4 4-MeOC6 H4 4-MeC6 H4

Benzamidines (R3 = phenyl) 36 4-BrC6 H4 Me 37 Ph Me 38 4-MeOC6 H4 Me 39 Ph Me 40 4-BrC6 H4 Me 41 Ph Me 42 Ph Me 43 4-MeOC6 H4 Me

pKBHX

∆G◦

Ref.

∼0.60 1.20a 1.33a 1.37 1.43 1.54a 1.60 1.63 1.65 1.80 1.90 1.99 2.00 2.07 2.08 2.10 2.12 2.12a 2.26 ∼2.30 2.32 2.35 2.36 2.36 2.40 2.41 2.43 2.48 2.52 2.52 2.59 2.59 2.60 2.70 2.74a

−3.42 −6.85 −7.59 −7.82 −8.16 −8.79 −9.13 −9.30 −9.42 −10.27 −10.85 −11.36 −11.42 −11.82 −11.87 −11.99 −12.10 −12.10 −12.90 −13.13 −13.24 −13.41 −13.47 −13.47 −13.70 −13.76 −13.87 −14.16 −14.38 −14.38 −14.78 −14.78 −14.84 −15.41 −15.64

[64] [60] [60] [60] [64] [60] [64] [60] [60] [64] [60] [64] [60] [60] [64] [60] [60] [64] [64] [64] [64] [60] [60] [64] [64] [94] [64] [64] [60] [64] [60] [64] [64] [64] [64]

0.99 1.11 1.23 1.41 1.42 1.52 1.53 1.63

−5.65 −6.34 −7.02 −8.05 −8.11 −8.68 −8.73 −9.30

[62] [62] [62] [62] [62] [62] [62] [62]


137

Table 4.10 (Continued) Lewis base: R1 R2 N C(R3 ) N R4 ∆G◦

Ref.

1.99 2.19 2.62 2.72

−11.36 −12.50 −14.95 −15.53

[62] [62] [62] [62]

3-ClC6 H4 4-BrC6 H4 3-ClC6 H4 4-BrC6 H4 Ph 4-BrC6 H4 Ph Ph 4-MeC6 H4 Ph 4-MeC6 H4

1.05 1.19 1.24 1.36 1.38 1.43 1.44 1.65 1.75 1.76 1.82

−5.99 −6.79 −7.08 −7.76 −7.88 −8.16 −8.22 −9.42 −9.99 −10.05 −10.39

[62] [62] [62] [62] [62] [62] [62] [62] [62] [62] [62]

i-Pr Me H

1.06 3.16 3.21 3.48

−6.05 −18.04 −18.32 −19.86

[94] [94] [94] [94]

2.50c −14.27

[96]

No. R1

R2

R3

R4

pKBHX

44 45 46 47

Me Me Me Me

4-NO2 C6 H4 4-BrC6 H4 Ph 4-MeC6 H4

Me Me Me Me

Me Me Me Me Me Me Me Me Me Me Me

Me Me Me Me Me Me Me Me Me Me Me

Me Me Me Me

Acetamidines (R3 = methyl) 48 3-ClC6 H4 49 4-BrC6 H4 50 Ph 51 Ph 52 4-BrC6 H4 53 4-MeC6 H4 54 3-ClC6 H4 55 Ph 56 Ph 57 4-MeC6 H4 58 4-MeC6 H4 Guanidines 59 i-Pr 60 Me 61 Me 62 7-Methyl-1,5, 7-triazabicyclo[4.4.0] dec-5-ene (MTBD)b

i-Pr N(i-Pr)2 Me NMe2 Me NMe2

Miscellaneous 63 2-Dimethylamino-3, 3-dimethylazirineb 64 1,8-Diazabicyclo[5.4.0] undec-7-ene (DBU)b 65 1,5-Diazabicyclo[4.3.0] non-5-ene (DBN)b Tautomerizing amidines R1 HNC(R3 ) NR4 R1 N C(R3 )NHR4d 66 Ph H Me 4-NO2 C6 H4 67 Ph H Ph 3-NO2 C6 H4 68 4-BrC6 H4 H Ph 4-BrC6 H4 69 Ph H Ph 3-ClC6 H4 70 4-ClC6 H4 H Me 4-ClC6 H4 71 Ph H Ph 4-ClC6 H4 72 Ph H Me 3-BrC6 H4 73 Ph H Ph 4-FC6 H4 74 Ph H Ph Ph 75 Ph H Me 4-BrC6 H4 76 3-MeC6 H4 H Ph Ph 77 4-MeC6 H4 H Ph Ph 78 4-MeC6 H4 H Me 4-BrC6 H4 79 4-ClC6 H4 H H 4-ClC6 H4

3.85

−21.98 [100]

3.89

−22.20 [100]

0.83 0.86 0.87 1.05 1.07 1.13 1.24 1.25 1.29 1.36 1.40 1.47 1.58 1.63

−4.74 −4.91 −4.97 −5.99 −6.11 −6.45 −7.08 −7.14 −7.36 −7.76 −7.99 −8.39 −9.02 −9.30

[68] [68] [68] [68] [68] [68] [68] [68] [68] [68] [68] [68] [68] [68]

(Continued)

138


Table 4.10 (Continued) Lewis base: R1 R2 N C(R3 ) N R4 No.

R1

R2

R3

R4

pKBHX

∆G◦

Ref.

80 81 82 83 84 85

Ph Me 4-MeC6 H4 4-MeC6 H4 Ph 4-MeC6 H4

H H H H H H

Me Ph Me Me H H

Ph Ph Ph 4-MeC6 H4 Ph 4-MeC6 H4

1.66 1.83 1.85 2.01 2.13 2.22

−9.48 −10.45 −10.56 −11.47 −12.16 −12.67

[68] [68] [68] [68] [68] [68]

a b

Value corrected for the presence of a second hydrogen-bond acceptor group. Formulae: N

N

N

N

Me N

N

NMe2

Me

N

N

Me

62 c

63

64

65

Secondary value. d The percentage of the right side tautomer varies from 50 to 81%.

b

R

N

8 9 12 16

Secondary value.

C

Formulae:

R: 3-CF3C6H4 R: 3-FC6H4 R: Ph R: Me

3

N OC H

2-Propanone O-phenyloxime N-Benzylideneaniline 3-Cyclohexen-1-one O-decyloxime N-Benzylidenebenzylamine N-Benzylidene-tert-butylamine Neopentylideneisobutylamine N-Benzylidenemethylamine 2-(3-Trifluoromethylphenyl)-1-pyrroline 2-(3-Fluorophenyl)-1-pyrroline Isobutylideneisopropylamine Benzophenone imine, diphenylketimine 2-Phenyl-1-pyrroline 3-Phenylprop-2-enylideneisopropylamine trans-Hexa-2,4-dienylideneisopropylamine trans-Butenylideneisopropylamine 2-Methyl-1-pyrroline

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16

a

Lewis base

No.

a

PhCH CH CH N-i-Pr MeCH CH CH CH CH N-i-Pr MeCH CH CH N-i-Pr

a

i-PrCH N-i-Pr Ph2 C NH

a

a

PhCH NCH2 Ph PhCH N-t-Bu t-BuCH N-i-Bu PhCH NMe

a

Me2 C NOPh PhCH NPh

Formula

Ref. [96] [94] [96] [94] [96] [96] [94] [99] [99] [96] [96] [99] [96] [96] [96] [14]

∆G◦ −4.57 −4.97 −6.39 −6.74 −7.36 −7.71 −8.50 −8.68 −9.48 −10.16 −10.27 −11.30 −11.42 −11.93 −12.22 −14.78

pKBHX 0.80b 0.87 1.12b 1.18 1.29 1.35b 1.49 1.52 1.66 1.78b 1.80 1.98 2.00b 2.09b 2.14b 2.59

Table 4.11 Hydrogen-bond basicity scales, pKBHX and ∆G◦ (kJ mol−1 ), for the complexes of 4-fluorophenol with imines and oximes, in CCl4 at 25 ◦ C.

140


Table 4.12 Hydrogen-bond basicity scales, pKBHX and ∆G◦ (kJ mol−1 ), for the complexes of 4-fluorophenol with nitriles, in CCl4 at 25 ◦ C. No.

Lewis base

Formula

pKBHX

∆G◦

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41

Trichloroacetonitrile Pentafluorobenzonitrile Dibromoacetonitrile Cyanogen bromide Chloroacetonitrile α,α,α-Trifluoro-m-tolunitrile α,α,α-Trifluoro-p-tolunitrile 2-Fluorobenzonitrile 4-Chlorobenzonitrile 2-Chlorobenzonitrile α-Bromo-o-tolunitrile Acrylonitrile 2-Bromobenzonitrile 4-Fluorobenzonitrile Methyl thiocyanate Phenyl cyanate Benzonitrile Benzyl cyanide o-Tolunitrile 2,6-Dimethylbenzonitrile Acetonitrile-d 3 Allyl cyanide Butyronitrile Hexanenitrile Acetonitrile Propionitrile Trimethylsilyl cyanide 4-Methoxybenzonitrile Isobutyronitrile Trimethylacetonitrile 1-Adamantanecarbonitrile Cyclopropyl cyanide 2-Methoxybenzonitrile 4-Dimethylaminobenzonitrile 4-Morpholinecarbonitrile Cyanamide Dimethylcyanamide 1-Piperidinecarbonitrile Diethylcyanamide 1-Pyrrolidinecarbonitrile trans-3Dimethylaminoacrylonitrile Diisopropylcyanamide N,N-Dimethyl-N’ cyanoformamidine N-Methyl-N’ propylcyanoguanidine

CCl3 C N C6 F5 C N Br2 CHC N BrC N ClCH2 C N 3-CF3 C6 H4 C N 4-CF3 C6 H4 C N 2-FC6 H4 C N 4-ClC6 H4 C N 2-ClC6 H4 C N 2-BrCH2 C6 H4 C N H2 C CHC N 2-BrC6 H4 C N 4-FC6 H4 C N MeSC N PhOC N PhC N PhCH2 C N 2-MeC6 H4 C N 2,6-Me2 C6 H3 C N CD3 C N CH2 CHCH2 C N n-PrC N n-PentC N MeC N EtC N Me3 SiC N 4-MeOC6 H4 C N i-PrC N t-BuC N 1-AdamC N c-PrC N 2-MeOC6 H4 C N 4-Me2 NC6 H4 C N

−0.26 0.01 0.19 0.19 0.39 0.53 0.54 0.64 0.66 0.67 0.69 0.70 0.71 0.72 0.73 0.77 0.80 0.81 0.83 0.86 0.87 0.87 0.89 0.89 0.91 0.93 0.93 0.97 1.00 1.00 1.00 1.03 1.06 1.23 1.34b 1.40 1.56 1.58 1.63 1.66 1.70

1.48 −0.06 −1.08 −1.08 −2.23 −3.03 −3.08 −3.65 −3.77 −3.82 −3.94 −4.00 −4.05 −4.11 −4.17 −4.40 −4.57 −4.62 −4.74 −4.91 −4.97 −4.97 −5.08 −5.08 −5.19 −5.31 −5.31 −5.54 −5.71 −5.71 −5.71 −5.88 −6.05 −7.02 −7.65 −7.99 −8.90 −9.02 −9.30 −9.48 −9.70

[66] [101] [66] [66] [66] [66] [66] [66] [96] [66] [66] [66] [66] [66] [66] [66] [66] [66] [66] [96] [66] [96] [66] [66] [66] [96] [101] [9] [66] [96] [66] [66] [66] [9] [96] [94] [67] [67] [67] [94] [67]

i-Pr2 NC N Me2 NCH NC N

1.74 2.09

−9.93 −11.93

[94] [67]

a

2.22c

−12.67

[67]

42 43 44

a

H2 NC N Me2 NC N c-(CH2 )5 NC N Et2 NC N c-(CH2 )4 NC N Me2 NCH CHC N

Ref.


141


Lewis base

Formula

pKBHX

∆G◦

Ref.

45

N,N-Dimethyl-N’ cyanoacetamidine N,N,N-Tri-nbutylammoniocyanamidate

Me2 NC(Me) NC N

2.24

−12.79

[67]

n-Bu3 N+ N− C N

3.24

−18.49

[72]

46 a

Formulae: Me

O

NC

N

H H

N C N

35 b c

N C N

nPr

44

The contribution of the oxygen acceptor site is assumed to be small. Secondary value.

142


Table 4.13 Hydrogen-bond basicity scales, pKBHX and ∆G◦ (kJ mol−1 ), for the complexes of 4-fluorophenol with water, alcohols and phenols, in CCl4 at 25 ◦ C [61]. No.

Lewis base

Formula

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26

1,1,1,3,3,3-Hexafluoro-2-propanol 3-Trifluoromethylphenol 2,2,2-Trifluoroethanol 4-Fluorophenol 2,2,2-Trichloroethanol Phenol 3-Methylphenol 4-Methylphenol Propargyl alcohol 2-Chloroethanol 2-Bromoethanol 2-Fluoroethanol Water Allyl alcohol Methanol Benzyl alcohol Phenethyl alcohol Ethylene glycol n-Propanol n-Butanol Ethanol n-Octanol 2-Propanol tert-Butanol Cyclohexanol 1-Adamantanol

(CF3 )2 CHOH 3-CF3 C6 H4 OH CF3 CH2 OH 4-FC6 H4 OH CCl3 CH2 OH C6 H5 OH 3-MeC6 H4 OH 4-MeC6 H4 OH HC CCH2 OH ClCH2 CH2 OH BrCH2 CH2 OH FCH2 CH2 OH H2 O CH2 CHCH2 OH MeOH PhCH2 OH PhCH2 CH2 OH HOCH2 CH2 OH n-PrOH n-BuOH EtOH n-OctOH i-PrOH t-BuOH c-C6 H11 OH

a

a

Formula: OH

b c

Secondary value. Statistically corrected to put the pKBHX value on a per oxygen basis.

pKBHX

∆G◦

∼−0.96b ∼−0.36b −0.28b −0.13 −0.13b −0.07b 0.01b 0.03b 0.38b 0.50 0.54b 0.55b 0.65 0.79b 0.82 0.86b 0.97b ∼0.97b,c 1.00b 1.02b 1.02 1.04b 1.06 1.14 1.14b 1.27

5.48 2.05 1.60 0.74 0.74 0.40 −0.06 −0.17 −2.17 −2.85 −3.08 −3.14 −3.71 −4.51 −4.68 −4.91 −5.54 −5.54 −5.71 −5.82 −5.82 −5.94 −6.05 −6.51 −6.51 −7.25


143

Table 4.14 Hydrogen-bond basicity of ethers and peroxides: complexation constants with 4-fluorophenol, Kc (t) (l mol−1 ), in CCl4 at 25 ◦ C, pKBHX and ∆G◦ (kJ mol−1 ) [81]. No.

Lewis base

3 4 5

Ethers Hexamethyldisiloxane 1,1,1,3,3,3-Hexafluoro-2propyl methyl ether Furan Dichloromethyl methyl ether Bis(2-chloroethyl) ether

6 7 8 9 10

1,3,5-Trioxane Ethyl vinyl ether Tetramethyl orthocarbonate 3,4-Dihydro-2H-pyran 2-Chloroethyl ethyl ether

11 12 13 14 15 16

1,3-Dioxolane 2,3-Dihydrofuran Trimethyl orthoformate Dimethoxymethane 1,3-Dioxane Dibenzyl ether

17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38

Diallyl ether 1,4-Dioxane Di-tert-butyl ether Di-n-butyl ether Propylene oxide Diethyl ether 1,2-Dimethoxyethane tert-Butyl ethyl ether 1,2-Diethoxyethane Diisopropyl ether 15-Crown-5 12-Crown-4 Cyclohexene oxide tert-Butyl methyl ether 18-Crown-6 3-Methyltetrahydropyran Tetrahydropyran Tetrahydrofuran 2-Methyltetrahydrofuran Oxetane, trimethylene oxide Cineole, eucalyptol 2,2,5,5Tetramethyltetrahydrofuran 2,3-Diadamant-2-yl oxirane Epichlorhydrin

1 2

39 40

Formula

Kc (t)

∼−0.53 ∼−0.41

(Me3 Si)2 O MeOCH(CF3 )2 a

MeOCHCl2 (ClCH2 CH2 )2 O

0.8 1.9

a

3.2

EtOCH CH2 (MeO)4 C

7.2

EtOCH2 CH2 Cl

3.6

a

5.6

a

a

(MeO)3 CH (MeO)2 CH2

a

(PhCH2 )2 O

10.5 7.6 8.5 5.2

(H2 C CHCH2 )2 O

a

11

t-Bu2 O n-Bu2 O a

Et2 O MeOCH2 CH2 OMe t-BuOEt EtOCH2 CH2 OEt i-Pr2 O a a a

21 25 66 54

t-BuOMe a a

95

a a a a a a a a

pKBHX

2.8

∆G◦ 3.03 2.34

−0.40 ∼−0.26 (O)b −0.03 (O)b −0.34 (Cl) 0.02c 0.10 0.26c 0.41 0.44 (O)b −0.11 (Cl) 0.45c 0.53 0.55c 0.58c 0.63c 0.65 (O)b −0.41 (π ) 0.70 0.73c 0.75 0.88 0.97 1.01 1.02c 1.08 1.09c 1.11 1.12c 1.13c 1.13 1.19 1.20c 1.22 1.23 1.28 1.34 1.36 1.38 1.43

2.28 1.48 0.17 1.94 −0.11 −0.57 −1.48 −2.34 −2.51 0.63 −2.57 −3.03 −3.14 −3.31 −3.60 −3.71 2.34 −4.00 −4.17 −4.28 −5.02 −5.54 −5.77 −5.82 −6.16 −6.22 −6.34 −6.39 −6.45 −6.45 −6.79 −6.85 −6.96 −7.02 −7.31 −7.65 −7.76 −7.88 −8.16

1.44 —

−8.22 (Continued)

144



Lewis base

41 42 43

Peroxides tert-Butyl peroxide 3,4-Diadamant-2-yl dioxetan Ascaridole

a

Formula

Kc (t)

t-BuOO-t-Bu

2.7 8.5 17

a a

∆G◦

pKBHX

−0.74 −3.60 −5.25

0.13c 0.63c 0.92c

Formulae: O

O

O

O

O

O

3

O

11

12

O

O

O O

O

O

O

O

9

15

18

21

O

O O

27

O

28

O

O O

O

O

O

O

6

O

O

O

O

O

O

O

O

O

O

29

31

32 CH2Cl

O

33

34 O

35

36

37

38

O

O O

O

39 b

40

42

43

Corrected for the presence of a second hydrogen-bond acceptor site. c Statistically corrected to put the pKBHX value on a per oxygen basis. However peroxides 42 and 43 may form three-centre hydrogen bonds.


145

Table 4.15 Total equilibrium constants, Kc (t) (l mol−1 ), for the π and oxygen complexes of 4-fluorophenol with aromatic ethers, in CCl4 at 25 ◦ C. Hydrogen-bond basicity scales, pKBHX and ∆G◦ (kJ mol−1 ) [96]. No. Lewis base

Formula

1

4-Chloroanisole

2 3

1,2-(Methylenedioxy) benzene, 1,3-benzodioxole Diphenyl ether

4

Xanthene

5

Anisole, methoxybenzene

6

Phenetole, ethoxybenzene

7

1,4-Dimethoxybenzene

8

2,3-Dihydro-1,4-benzodioxin

9

2,3-Dihydrobenzofuran 4-Methylanisole

11

2-Methylanisole

12

3-Methylanisole

13


14

2,6-Dimethylanisole

15

2-tert-Butyl-5-methylanisole

16


a

∆G◦ 3.42 1.43 3.03 2.57 2.45 4.57 2.45 9.76 2.11 0.40 1.94 0.06 1.83 −0.68 1.83 1.31 1.71 −1.20 1.66 −0.46 1.60 2.11 1.54 −0.29 1.48 0.46 1.20 −1.83 0.91 8.22 −6.62

Formulae: O

O

O

2 b

pKBHX (π ) pKBHX (O)

0.8 −0.60 −0.25 a 1.0 −0.53 −0.45b Ph2 O 0.9 −0.43c −0.80 a 0.8 −0.43c −1.71 C6 H5 OMe 1.3 −0.37 −0.07 C6 H5 OEt 1.4 −0.34 −0.01 4-MeOC6 H4 OMe 3.1 −0.32 0.12b a 1.7 −0.32 −0.23b a 2.1 −0.30 0.21 4-MeC6 H4 OMe 1.7 −0.29 0.08 2-MeC6 H4 OMe 1.0 −0.28 −0.37 3-MeC6 H4 OMe 1.7 −0.27 0.05 3-MeOC6 H4 OMe 2.2 −0.26 −0.08b 2,6-Me2 C6 H3 OMe 2.7 −0.21 0.32 2-t-Bu-5-MeC6 H3 OMe 0.7 −0.16 −1.44 2-MeOC6 H4 OMe 14.5 1.16 (O)d 4-ClC6 H4 OMe

10

Kc (t)

O

4

O

O

8

9

Statistically corrected to put the pKBHX (O) value on a per oxygen basis. Statistically corrected to put the pKBHX (π ) value on a per π sextet basis. d Without statistical correction, because of a three-centre hydrogen bond. The π contribution becomes negligible. c

Formula 2-HOC6 H4 COH 4-ClC6 H4 COH MeCOH PhCOH 4-MeOC6 H4 COH 4-EtOC6 H4 COH 2-MeOC6 H4 COH PhCH CHCOH 4-MeNC6 H4 COH MeCOCF3 MeCOCl3 MeCOCOMe MeCOCHCl2 ClCH2 COCH2 Cl t-BuCO-t-Bu PhCH2 COCH2 Ph i-BuCO-i-Bu s-BuCO-s-Bu i-PrCO-i-Pr 1-AdCO-t-Bu EtCO-n-Pr n-PrCO-n-Pr EtCOEt MeCO-t-Bu MeCO-n-Pr MeCO-i-Bu 1-Adam-CO-1-Adam

Lewis base

Aldehydes Salicylaldehyde 4-Chlorobenzaldehyde Acetaldehyde Benzaldehyde 4-Methoxybenzaldehyde 4-Ethoxybenzaldehyde 2-Methoxybenzaldehyde trans-Cinnamaldehyde 4-Dimethylaminobenzaldehyde

Aliphatic ketones 1,1,1-Trifluoropropan-2-one 1,1,1-Trichloropropan-2-one Biacetyl 1,1-Dichloropropan-2-one 1,3-Dichloropropan-2-one 2,2,4,4-Tetramethylpentan-3-one Dibenzyl ketone 2,6-Dimethylheptan-4-one 3,5-Dimethylheptan-4-one 2,4-Dimethylpentan-3-one 1-Adamantyl tert-butyl ketone Hexan-3-one Heptan-4-one Pentan-3-one 3,3-Dimethylbutan-2-one Pentan-2-one 4-Methylpentan-2-one Di-(1-Adamantyl)ketone

1 2 3 4 5 6 7 8 9

10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27

No.

−0.06 0.00 0.23b 0.25 0.32 0.96 1.00c 1.07c 1.07c 1.08c 1.08 1.13c 1.14c 1.14c 1.17c 1.17c 1.17c 1.17

0.36 0.63 0.65 0.78 1.10 1.10 1.11 1.13 1.53

pKBHX

Table 4.16 Hydrogen-bond basicity scales, pKBHX and ∆G◦ (kJ mol−1 ), of aldehydes and ketones, in CCl4 at 25 ◦ C.

0.34 0.00 −1.31 −1.43 −1.83 −5.48 −5.71 −6.11 −6.11 −6.16 −6.16 −6.45 −6.51 −6.51 −6.68 −6.68 −6.68 −6.68

−2.05 −3.60 −3.71 −4.45 −6.28 −6.28 −6.34 −6.45 −8.73

∆G◦

[79] [79] [79] [79] [79] [79] [79] [79] [79] [79] [79] [79] [79] [79] [79] [79] [79] [79]

[76] [79] [79] [79] [7] [100] [76] [79] [9]

Ref.

c-(CH2 )5 CO c-(CH2 )6 CO c-(CH2 )7 CO 3-NO2 C6 H4 COMe 2-OHC6 H4 COMe 4-NO2 C6 H4 COMe 4-N CC6 H4 COMe 3-CF3 C6 H4 COMe 4-CF3 C6 H4 COMe

Ring-substituted acetophenones 3-Nitroacetophenone 2-Hydroxyacetophenone 4-Nitroacetophenone 4-Cyanoacetophenone 3-Trifluoromethylacetophenone 4-Trifluoromethylacetophenone

49 50 51 52 53 54

a

a

c-(CH2 )3 CO c-(CH2 )10 CO c-(CH2 )14 CO c-(CH2 )11 CO c-(CH2 )4 CO MeCH(CH2 )4 CO

Cycloalkanones Cyclobutanone Cycloundecanone Cyclopentadecanone Cyclododecanone Cyclopentanone 2-Methylcyclohexanone Camphor 5-α-Androstane-17-one Cyclohexanone Cycloheptanone Cyclooctanone

38 39 40 41 42 43 44 45 46 47 48

MeCO-n-Bu MeCOMe MeCO-i-Pr n-BuCO-n-Bu MeCOEt MeCO-s-Bu MeCO-c-Hex MeCO-1-Adam MeCO-c-Pr c-PrCO-c-Pr

Hexan-2-one Propan-2-one, acetone 3-Methylbutan-2-one Nonan-5-one Butan-2-one 3-Methylpentan-2-one Cyclohexyl methyl ketone 1-Adamantyl methyl ketone Methyl cyclopropyl ketone Dicyclopropyl ketone

28 29 30 31 32 33 34 35 36 37

0.53d 0.56 0.57d 0.60d 0.72 0.78

1.00 1.20 1.22 1.23 1.27 1.27c 1.31 1.36 1.39 1.41 1.45

1.18c 1.18 1.20c 1.21 1.22 1.22c 1.24c 1.30 1.32c 1.36

−3.03 −3.20 −3.25 −3.42 −4.11 −4.45

−5.71 −6.85 −6.96 −7.02 −7.25 −7.25 −7.48 −7.76 −7.93 −8.05 −8.28

−6.74 −6.74 −6.85 −6.91 −6.96 −6.96 −7.08 −7.42 −7.53 −7.76

(Continued)

[79] [76] [79] [79] [79] [79]

[79] [79] [79] [79] [79] [79] [79] [96] [79] [79] [79]

[79] [79] [79] [96] [79] [79] [79] [79] [79] [79]

3-ClC6 H4 COMe 3-FC6 H4 COMe 3-COMeC6 H4 COMe 2-ClC6 H4 COMe 4-COMeC6 H4 COMe 4-ClC6 H4 COMe 4-FC6 H4 COMe 3-MeC6 H4 COMe C6 H5 COMe 3-MeOC6 H4 COMe 4-MeSC6 H4 COMe 4-i-PrC6 H4 COMe 4-MeC6 H4 COMe 4-EtC6 H4 COMe 4-t-BuC6 H4 COMe 4-(1-Adam)C6 H4 COMe 4-MeOC6 H4 COMe 2-MeOC6 H4 COMe 4-NH2 C6 H4 COMe O(CH2 )4 NC6 H4 COMe c-(CH2 )5 NC6 H4 COMe 4-NMe2 C6 H4 COMe 4-NEt2 C6 H4 COMe 2-HOC6 H4 COC6 H5 2,4,6-Me3 C6 H2 COC6 H2 -2 ,4 ,6 -Me3 C6 H5 COC6 H5 4-MeOC6 H4 COC6 H5 4-MeOC6 H4 COC6 H4 -4 -OMe

Ring-substituted benzophenones 2-Hydroxybenzophenone Bis(mesityl) ketone Benzophenone 4-Methoxybenzophenone 4,4 -Dimethoxybenzophenone

78 79 80 81 82

Formula

3-Chloroacetophenone 3-Fluoroacetophenone 1,3-Diacetylbenzene 2-Chloroacetophenone 1.4-Diacetylbenzene 4-Chloroacetophenone 4-Fluoroacetophenone 3-Methylacetophenone Acetophenone 3-Methoxyacetophenone 4-(Methylthio)acetophenone 4-Isopropylacetophenone 4-Methylacetophenone 4-Ethylacetophenone 4-tert-Butylacetophenone 4-(1-Adamantyl)acetophenone 4-Methoxyacetophenone 2-Methoxyacetophenone 4-Aminoacetophenone 4-Morpholinoacetophenone 4-Piperidinoacetophenone 4-Dimethylaminoacetophenone 4-Diethylaminoacetophenone

Lewis base

55 56 57 58 59 60 61 62 63 64 65 66 67 68 69 70 71 72 73 74 75 76 77

No.


0.49 1.01c 1.07 1.27 1.49

0.82 0.83 0.86b 0.90 0.92b 0.93 1.00 1.10 1.11 1.16 1.21 1.21 1.24 1.25 1.25 1.30 1.33 1.34 1.50c 1.61 1.71 1.76 1.82

pKBHX

[76] [79] [79] [79] [79]

[79] [79] [79] [79] [79] [79] [79] [79] [79] [79] [79] [79] [79] [79] [79] [79] [79] [76] [79] [79] [79] [79] [79]

−4.68 −4.74 −4.91 −5.14 −5.25 −5.31 −5.71 −6.28 −6.34 −6.62 −6.91 −6.91 −7.08 −7.14 −7.14 −7.42 −7.59 −7.65 −8.56 −9.19 −9.76 −10.05 −10.39 −2.80 −5.77 −6.11 −7.25 −8.50

Ref.

∆G◦

Miscellaneous conjugated ketones Benzil 1,4-Benzoquinone But-3-yn-2-one 9,10-Phenanthrenequinone Acetylacetone (keto–enol tautomer) Propiophenone Butyrophenone 9-Fluorenone 2-Acetylnaphthalene Anthrone Thioxanthen-9-one 3-Chloro-5,5-dimethyl-2-cyclohexen-1-one α,α -Diphenylbenzotropone Tropolone Xanthone 1,4-Androstadiene-3,17-dione

trans-4-Phenylbut-3-en-2-one α,α -Dimethylbenzotropone Acetylferrocene 3,5,5-Trimethyl-2-cyclohexen-1-one, isophorone Benzotropone 10-Methyl-9(10H)-acridone Tropone Flavone γ -Pyrone N,N-Dimethyl-N -benzoylformamidine

102 103 104 105 106 107 108 109 110 111

4-Dimethylaminobenzophenone 4,4 -Bis(dimethylamino)benzophenone 4,4 -Bis(diethylamino)benzophenone

86 87 88 89 90 91 92 93 94 95 96 97 98 99 100 101

83 84 85

Me2 NCH NCOPh

a

a

a

a

a

a

C5 H5 FeC5 H4 COMe

PhCH CHCOMe

a

a

a

a

a

a

a

a

a

a

MeCOCH C(Me)OH PhCOEt PhCO-n-Pr

a

HC CCOMe

a

PhCOCOPh

4-Me2 NC6 H4 COC6 H5 4-Me2 NC6 H4 COC6 H4 -4 -NMe2 4-Et2 NC6 H4 COC6 H4 -4 -NEt2 0.44b 0.51b 0.68 0.70b,c 0.90 1.04 1.04 1.09 1.13 1.16 1.18c 1.21 1.30c 1.31 1.36 1.36 (O17) 1.84 (O3) 1.38 1.48c 1.65 1.74 1.88c 1.92c 1.97 1.99 2.03 2.10

1.67 1.93c 2.33 −2.51 −2.91 −3.88 −4.00 −5.14 −5.94 −5.94 −6.22 −6.45 −6.62 −6.74 −6.91 −7.42 −7.48 −7.76 −7.76 −10.50 −7.88 −8.45 −9.42 −9.93 −10.73 −10.96 −11.24 −11.36 −11.59 −11.99

−9.53 −11.02 −13.30

(Continued)

[79] [79] [79] [79] [79] [96] [96] [79] [79] [79] [79] [79] [79] [76] [9] [96] [96] [79] [79] [9] [79] [79] [79] [79] [7] [79] [63]

[79] [79] [79]

99

H

O

X

O

O

3

101

95 96 100 107

45

17

O

X CH2 S O NMe

O

87

O

O

108

O

O

X

89

O

O

O

109

97 105 114

O

Ph

X 3-Cl 3-Me 3-NMe2

Statistically corrected to put the pKBHX value on a per carbonyl basis. c Secondary value. d Corrected for the presence of a second hydrogen-bond acceptor site.

b

O

44

O

Formulae:

X

93

O

O

O

X

Diphenylcyclopropenone 2,6-Dimethyl-γ -pyrone 3-Dimethylamino-5,5-dimethyl-2-cyclohexen-1-one

112 113 114

a

Lewis base

No.


110 113

O

94

X H Me

X

X

Ph

98 103 106

a

a

a

112

O

X Ph Me H

Ph

COMe

Formula 2.30 2.50 2.92

pKBHX

Ref. [79] [7] [79]

∆G◦ −13.13 −14.27 −16.67

Ethyl propiolate Diethyl oxalate Methyl formate Ethyl formate Ethyl chloroacetate Vinylene carbonate Ethyl fluoroacetate Diethyl terephthalate Ethyl 4-bromobenzoate 2-Chloroethyl acetate Dimethyl carbonate Ethyl methyl carbonate Triacetin β-Propiolactone Diethyl carbonate Methyl benzoate Ethyl benzoate β-Butyrolactone Methyl acetate Methyl 3,3-dimethylacrylate Ethyl 2,2-dimethylpropanoate Ethyl 4-methylbenzoate

Ethyl 4-cyanobenzoate

8

9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30

2,2,2-Trifluoroethyl acetate Ethyl trifluoroacetate Methyl trichloroacetate Ethyl trichloroacetate Methyl salicylate 2,2,2-Trichloroethyl acetate Ethyl 4-nitrobenzoate

Lewis base

MeCOOMe Me2 CH CHCOOMe t-BuCOOEt 4-MeC6 H4 COOEt

a

EtOCOOEt PhCOOMe PhCOOEt

a

FCH2 COOEt 4-EtOOCC6 H4 COOEt 4-BrC6 H4 COOEt MeCOOCH2 CH2 Cl MeOCOOMe MeOCOOEt (CH3 COOCH2 )2 CHOCOCH3

a

HC CCOOEt EtOOCCOOEt HCOOMe HCOOEt ClCH2 COOEt

4-N CC6 H4 COOEt

MeCOOCH2 CF3 CF3 COOEt CCl3 COOMe CCl3 COOEt 2-HOC6 H4 COOMe MeCOOCH2 CCl3 4-NO2 C6 H4 COOEt

Formula

∆G◦ 0.40 −0.46 −0.63 −0.86 −1.83 −2.68 −2.80 −0.46 −3.03 −3.77 −3.60 −3.71 −3.71 −3.77 −3.82 −3.94 −4.22 −4.34 −4.45 −4.57 −4.68 −4.79 −4.79 −4.91 −5.02 −5.08 −5.37 −5.54 −5.71 −5.82 −5.94 −5.99

pKBHX −0.07 0.08 0.11 0.15 0.32 0.47 0.49 (CO)b 0.08 (NO2 ) 0.53 (CO)b 0.66 (C≡N) 0.63 0.65c 0.65 0.66 0.67 0.69 0.74 0.76c 0.78 0.80 0.82 0.84 0.84c 0.86 0.88 0.89 0.94 0.97 1.00 1.02 1.04 1.05

Hydrogen-bond basicity scales, pKBHX and ∆G◦ (kJ mol−1 ), of esters, lactones and carbonates, in CCl4 at 25 ◦ C.

1 2 3 4 5 6 7

No.

Table 4.17

(Continued)

[99] [71] [71] [71] [76] [99] [71] [71] [71] [71] [71] [71] [71] [71] [71] [71] [71] [71] [71] [99] [71] [71] [71] [71] [71] [71] [71] [71] [71] [99] [71] [71]

Ref.

14

22

O

O

HC

26

O

O

O

47

O

O

Me

48

O

O

49

O

O

c

Value corrected for the presence of a second hydrogen-bond acceptor group. Statistically corrected to put the pKBHX value on a per carbonyl basis. d Secondary value.

b

O

O

O

Formulae:

Ethyl phenylacetate Ethyl 1-adamantanecarboxylate Ethyl acetate Ethyl propionate Ethyl isobutyrate tert-Butyl acetate Ethyl isovalerate Ethyl butyrate Ethyl cyclopropanecarboxylate Butyl acetate Ethyl 4-methoxybenzoate (E)-Ethyl cinnamate Isobutyl acetate Isopropyl acetate Propyl acetate sec-Butyl acetate Propylene carbonate Coumarin γ -Butyrolactone γ -Valerolactone Ethyl 4-dimethylaminobenzoate Methyl 2-methoxybenzoate δ-Valerolactone ε-Caprolactone Ethyl 3-dimethylaminoacrylate

31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55

a

Lewis base

(Continued)

No.

Table 4.17

Me

50

O

O

53

O

O

Me2 NCH CHCOOEt

a

a

4-Me2 NC6 H4 COOEt 2-MeOC6 H4 COOMe

a

a

a

a

PhCH2 COOEt 1-AdamCOOEt MeCOOEt EtCOOEt i-PrCOOEt MeCOO-t-Bu i-BuCOOEt n-PrCOOEt c-PrCOOEt MeCOO-n-Bu 4-MeOC6 H4 COOEt PhCH CHCOOEt MeCOO-i-Bu MeCOO-i-Pr MeCOO-n-Pr MeCOO-s-Bu

Formula

54

O O

1.05 1.06 1.07 1.08 1.09 1.10 1.11d 1.11d 1.12 1.13d 1.13 1.14 1.14d 1.15 1.15d 1.16d 1.22 1.30 1.32 1.43 1.45 1.49 1.57 1.63 2.09

pKBHX

Ref. [71] [71] [71] [71] [71] [71] [71] [71] [71] [71] [71] [71] [71] [71] [71] [71] [71] [71] [71] [71] [71] [76] [71] [71] [71]

∆G◦ −5.99 −6.05 −6.11 −6.16 −6.22 −6.28 −6.34 −6.34 −6.39 −6.45 −6.45 −6.51 −6.51 −6.56 −6.56 −6.62 −6.96 −7.42 −7.53 −8.16 −8.28 −8.50 −8.96 −9.30 −11.93

Lewis base

Dimethylcarbamoyl chloride N,N-Dimethyl-2,2,2-trifluoroacetamide N,N-Diethyl-2,2,2-trifluoroacetamide N,N-Dimethylsalicylamide S-Methyl diethylthiocarbamate Acetanilide Phenyl dimethylcarbamate N,N-Dimethyl-2-chloroacetamide N-Methylformanilide Formamide Benzamide Ethyl dimethylcarbamate N-(4-Ethoxyphenyl)acetamide, phenacetin N-Formylmorpholine N,N -Diphenylacetamide N-Methylformamide N,N-Diethylnicotinamide

N-Methylbenzamide N,N -Diisopropyl-2,2-dimethylpropionamide Acetamide N,N-Dicyclohexyl-2,2-dimethylpropionamide N,N,2,2-Tetramethylpropionamide N,N-Dimethylformamide N,N-Diethylformamide N-Acetylmorpholine N-Formylpiperidine N,N-Di-n-butylformamide N-Methylacetanilide

No.

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17

18 19 20 21 22 23 24 25 26 27 28 HCON(CH2 )5 HCON-n-Bu2 MeCONMePh

a

PhCONHMe t-BuCON-i-Pr2 MeCONH2 t-BuCON-c-Hex2 t-BuCONMe2 HCONMe2 HCONEt2

MeCONPh2 HCONHMe 3-NEt2 COC5 H4 N

a

ClCONMe2 CF3 CONMe2 CF3 CONEt2 2-HOC6 H4 CONMe2 MeSCONEt2 MeCONHPh PhOCONMe2 ClCH2 CONMe2 HCONMePh HCONH2 PhCONH2 EtOCONMe2 MeCONHC6 H4 OEt

Formula 1.00 1.04 1.06 1.27 1.56 1.69 1.70 1.74 1.74 1.75b 1.81 1.83 1.83b 1.93 1.94 1.96 1.98 (O) 1.63 (N) 2.03 2.03 2.06b 2.06 2.10 2.10 2.15 2.16 2.17 2.17 2.19

pKBHX

Ref. [65] [65] [9] [76] [9] [94] [65] [65] [65] [94] [96] [65] [96] [94] [65] [65] [100] [100] [65] [65] [94] [65] [65] [65] [94] [94] [94] [94] [65]

∆G◦ −5.71 −5.94 −6.05 −7.25 −8.90 −9.65 −9.70 −9.93 −9.93 −9.99 −10.33 −10.45 −10.45 −11.02 −11.07 −11.19 −11.30 −9.30 −11.59 −11.59 −11.76 −11.76 −11.99 −11.99 −12.27 −12.33 −12.39 −12.39 −12.50

Table 4.18 Hydrogen-bond basicity scales, pKBHX and ∆G◦ (kJ mol−1 ), of amides, lactams, carbamates, ureas and carboxamidates, in CCl4 at 25 ◦ C.

29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57

N,N-Dicyclohexylpropionamide N,N-Dimethylbenzamide N,N-Diisopropylformamide N-Methylpropionamide N,N-Dicyclohexylisobutyramide N,N-Dimethylisobutyramide N-Ethylacetamide N-Methylacetamide N,N-Dimethyladamantane-1-carboxamide N-Formylpyrrolidine N,N-Dimethylpropionamide 1-Methyl-2-pyrrolidone N,N-Dicyclohexylacetamide 1,1,3,3-Tetraethylurea N,N-Dimethylacetamide 1,1,3,3-Tetramethylurea 1-Acetylpiperidine 4-(Dimethylamino)antipyrine N,N -Dimethyl-N,N -ethyleneurea (DMEU) N,N-Diethylacetamide N,N-Diisopropylacetamide N,N-Dimethyl-2-methoxybenzamide 1-Methyl-2-pyridone N-Methylcaprolactam 1-Methyl-2-piperidone 1-Acetylpyrrolidine Antipyrine N,N -Dimethyl-N,N -propyleneurea (DMPU) 3-[(Dimethylamino)methylene]-1,1-dimethylurea Me2 NCH NCONMe2

a

a

MeCON(CH2 )4

a

a

a

MeCONEt2 MeCON-i-Pr2 2-MeOC6 H4 CONMe2

a

a

MeCON-c-Hex2 Et2 NCONEt2 MeCONMe2 Me2 NCONMe2 MeCON(CH2 )5

a

EtCON-c-Hex2 PhCONMe2 HCON-i-Pr2 EtCONHMe i-PrCON-c-Hex2 i-PrCONMe2 MeCONHEt MeCONHMe 1-Adam-CONMe2 HCON(CH2 )4 EtCONMe2

2.22 2.23 2.24 2.24 2.24 2.26 2.29 2.30 2.30 2.31 2.36 2.38 2.41 2.43 2.44 2.44 2.45 2.45 2.46 2.47 2.47 2.48 2.50 2.53 2.60 2.61 2.79 2.79 2.92

−12.67 −12.73 −12.79 −12.79 −12.79 −12.90 −13.07 −13.13 −13.13 −13.19 −13.47 −13.59 −13.76 −13.87 −13.93 −13.93 −13.98 −13.98 −14.04 −14.10 −14.10 −14.16 −14.27 −14.44 −14.84 −14.90 −15.93 −15.93 −16.67

(Continued)

[65] [65] [94] [65] [65] [65] [65] [65] [96] [94] [65] [65] [65] [65] [65] [65] [94] [96] [65] [65] [94] [76] [99] [65] [65] [94] [100] [65] [100]

b

Me

O

55

N

Me

Ph

NCOMe

N

25

O

Secondary values.

53

Me

N

O

NCOH

O

24

Formulae:

Me

N

O

56

40

Me

N

N

O

Me

NMe

Me

46

O

N N

Me

Ph

Me

N

47

O

N Me

N

51

Me

O

52

Me

N O

4-CF3 C6 H4 CON− N+ Me3 C6 H5 CON− N+ Me3 4-MeOC6 H4 CON− N+ Me3 Me3 C6 H2 CON− N+ Me3 C9 H19 CON− N+ Me3 t-BuCON− N+ Me3 1-Adam CON− N+ Me3 MeCON− N+ Me3 i-PrCON− N+ Me3 EtCON− N+ Me3

Carboxamidates N,N,N-Trimethylammonio 4-trifluoromethylbenzamidate N,N,N-Trimethylammonio benzamidate N,N,N-Trimethylammonio 4-methoxybenzamidate N,N,N-Trimethylammonio mesitylamidate N,N,N-Trimethylammonio nonamidate N,N,N-Trimethylammonio trimethylacetamidate N,N,N-Trimethylammonio adamantyl-1-carboxamidate N,N,N-Trimethylammonio acetamidate N,N,N-Trimethylammonio isobutanamidate N,N,N-Trimethylammonio propanamidate

58 59 60 61 62 63 64 65 66 67

a

Formula

Lewis base

No.


2.66b 2.93b 3.02b 3.31b 3.36b 3.36b 3.40b 3.42b 3.49b 3.58b

pKBHX −15.18 −16.72 −17.24 −18.89 −19.18 −19.18 −19.41 −19.52 −19.92 −20.43

∆G◦

[69] [94] [69] [94] [69] [69] [69] [69] [69] [69]

Ref.

156


Table 4.19 Hydrogen-bond basicity scales, pKBHX and ∆G◦ (kJ mol−1 ), of nitro and nitroso compounds, in CCl4 at 25 ◦ C. No.

Lewis base

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23

Nitro compounds 3-Chloronitrobenzene 4-Chloronitrobenzene Nitromethane 1,3-Dimethyl-2-nitrobenzene Nitrobenzene Nitroethane N-nitrocamphorimine 4-Nitrotoluene 2-Nitropropane 4-(1-Adamantyl)nitrobenzene 2-Methyl-2-nitropropane 1,2-Dimethyl-4-nitrobenzene 4-Nitroanisole Dimethylnitramine N,N-Dimethyl-4-nitroaniline N,N-Diethyl-4-nitroaniline 1-Morpholino-2-nitroethylene 1-Dimethylamino-2-nitroethylene 1-Nitro-2-piperidinoethylene 1-Diethylamino-2-nitroethylene 1-Nitro-2-pyrrolidinoethylene Pyridinium nitramidate 1-Methylpiperidinium nitramidate

24 25 26 27 28

Nitroso compounds Nitrosobenzene N-Nitrosodiphenylamine N-Nitrosodimethylamine 4-Nitroso-N,N-diethylaniline 1-Nitrosopyrrolidine

a b

Formula

pKBHX a

∆G◦

Ref.

3-ClC6 H4 NO2 4-ClC6 H4 NO2 MeNO2 Me2 C6 H3 NO2 PhNO2 EtNO2

−0.10c −0.01c −0.03 −0.01 0.00 0.05c 0.04 0.12c 0.11 0.12 0.13c 0.16 0.20 0.52 0.53c 0.60 0.94c 1.17 1.25 1.28 1.32 1.39c 1.55c

0.57 0.06 0.17 0.06 0.00 −0.29 −0.23 −0.68 −0.63 −0.68 −0.74 −0.91 −1.14 −2.97 −3.03 −3.42 −5.37 −6.68 −7.14 −7.31 −7.53 −7.93 −8.85

[70] [70] [70] [96] [70] [70] [70] [70] [70] [96] [70] [70] [70] [74] [74] [70] [74] [70] [70] [74] [74] [74] [74]

0.15 0.64c 1.31c 1.33 1.49c

−0.86 −3.65 −7.48 −7.59 −8.50

[96] [96] [96] [96] [96]

b

4-MeC6 H4 NO2 i-PrNO2 b

t-BuNO2 Me2 C6 H3 NO2 4-MeOC6 H4 NO2 Me2 NNO2 4-NMe2 C6 H4 NO2 4-NEt2 C6 H4 NO2 OC4 H8 NCH CHNO2 Me2 NCH CHNO2 c-(CH2 )5 NCH CHNO2 Et2 NCH CHNO2 c-(CH2 )4 NCH CHNO2 C5 H5 N+ N− NO2 b

C6 H5 NO Ph2 NNO Me2 NNO 4-NEt2 C6 H4 NO c-(CH2 )4 NNO

Statistically corrected to put the pKBHX values of nitro compounds on a per oxygen basis. Formulae: Me +N

NO

NNO

7 c

10

Secondary values.

23

NNO


157

Table 4.20 Hydrogen-bond basicity scales, pKBHX and ∆G◦ (kJ mol−1 ), of sulfonyl compounds, in CCl4 at 25 ◦ C [78]. No. Lewis base

Formula

pKBHX a

∆G◦

1 2 3 4 5 6 7 8 9 10

EtOSO2 OEt MeSO2 OEt PhSO2 NMe2 PhSO2 Ph MeSO2 NMe2 MeSO2 Me c-(CH2 )4 SO2 Et2 NSO2 NEt2 n-BuSO2 -n-Bu PhSO2 N CHNMe2

0.50 0.72 0.89 0.91 1.00 1.10b 1.17 1.17 1.22 1.51

−2.85 −4.11 −5.08 −5.19 −5.71 −6.28 −6.68 −6.68 −6.96 −8.62

4-MeC6 H4 SO2 N SMe2

1.73b

−9.87

4-MeC6 H4 SO2 N− N+ Me3

2.14b

−12.22

n-OctSO2 N− N+ Me3

2.42b

−13.81

11 12 13 a b

Diethyl sulfate Ethyl methanesulfonate N,N-Dimethylbenzenesulfonamide Diphenyl sulfone N,N-Dimethylmethanesulfonamide Dimethyl sulfone Tetramethylene sulfone, sulfolane N,N,N ,N -Tetraethylsulfamide Di-n-butyl sulfone N,N-Dimethyl-N phenylsulfonylformamidine N-[4-Tolylsulfonylimino(dimethyl)-λ4 sulfane] N,N,N-Trimethylammoniotoluene-4sulfonamidate N,N,NTrimethylammoniooctanesulfonamidate

Statistically corrected to put the pKBHX values on a per oxygen basis. Secondary values.

158


Table 4.21 Hydrogen-bond basicity scales, pKBHX and ∆G◦ (kJ mol−1 ), of sulfinyl and seleninyl compounds, in CCl4 at 25 ◦ C [96]. No.

Lewis base

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25

Thionyl chloride Ethylene sulfite Dimethyl sulfite Diethyl sulfite Di-n-butyl sulfite Di-n-propyl sulfite Methyl 4-nitrophenyl sulfoxide Bis(4-chlorophenyl) sulfoxide Diphenyl sulfoxide N,N-Dimethybenzenesulfinamide Bis(4-methylphenyl) sulfoxide Methyl phenyl sulfoxide N,N,4-Trimethylbenzenesulfinamide Dibenzyl sulfoxide N,N-Dimethylmethanesulfinamide Tetramethylene sulfoxide Diisopropyl sulfoxide Dimethyl sulfoxide tert-Butyl isobutyl sulfoxide n-Butyl tert-butyl sulfoxide Di-n-butyl sulfoxide sec-Butyl tert-butyl sulfoxide n-Butyl sec-butyl sulfoxide Dibenzyl selenoxide Dimethyl selenoxide

a

Formula: O

O S O

b c

Secondary value. Ref. [7].

Formula

pKBHX

∆G◦

Cl2 SO

−0.38 0.87 0.94 1.07 1.08b 1.16b 1.58c 1.68 2.04 2.15b 2.21 2.24 2.24b 2.43 2.43b 2.47 2.51b 2.54 2.63b 2.64b 2.65 2.65b 2.66b 3.30 3.43

2.17 −4.97 −5.37 −6.11 −6.16 −6.62 −9.02 −9.59 −11.64 −12.27 −12.61 −12.79 −12.79 −13.87 −13.87 −14.10 −14.33 −14.50 −15.01 −15.07 −15.13 −15.13 −15.18 −18.84 −19.58

a

(MeO)2 SO (EtO)2 SO (n-BuO)2 SO (n-PrO)2 SO 4-NO2 C6 H4 SOMe (4-ClC6 H4 )2 SO Ph2 SO PhSONMe2 (4-MeC6 H4 )2 SO PhSOMe 4-MeC6 H4 SONMe2 (PhCH2 )2 SO MeSONMe2 c-(CH2 )4 SO i-Pr2 SO Me2 SO t-BuSO-i-Bu n-BuSO-t-Bu (n-Bu)2 SO s-BuSO-t-Bu n-BuSO-s-Bu (PhCH2 )2 SeO Me2 SeO


159

Table 4.22 Hydrogen-bond basicity scales, pKBHX and ∆G◦ (kJ mol−1 ), of oxides of organoderivatives of group 15 elements, in CCl4 at 25 ◦ C [94, 96]. No. Lewis base

Formula

pKBHX

∆G◦

1 2 3 4 5 6 7 8 9 10 11 12

Amine oxides 4-Nitropyridine N-oxide 3,5-Dichloropyridine N-oxide 4-Chloropyridine N-oxide Pyridine N-oxide 4-Phenylpyridine N-oxide 3-Methylpyridine N-oxide 4-Methylpyridine N-oxide 4-Methoxypyridine N-oxide Triethylamine N-oxide N-methylpiperidine N-oxide Tri(n-octyl)amine N-oxide Trimethylamine N-oxide

4-NO2 C5 H4 NO 3,5-Cl2 C5 H3 NO 4-ClC5 H4 NO C5 H5 NO 4-PhC5 H4 NO 3-MeC5 H4 NO 4-MeC5 H4 NO 4-MeOC5 H4 NO Et3 NO c-(CH2 )5 N(Me)O n-Oct3 NO Me3 NO

1.05a 1.56 2.44 2.72 2.85a 2.92a 3.12 3.70a 4.79a 5.05a 5.06a 5.46a

−5.99 −8.90 −13.93 −15.53 −16.27 −16.67 −17.81 −21.12 −27.34 −28.83 −28.88 −31.17

13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28

Phosphoroso compounds Trimethylammonio-P,P-diphenylphosphinamidate Tripiperidinophosphine oxide Triethylphosphine oxide Tri(n-butyl)phosphine oxide Hexamethylphosphoric triamide (HMPA) Tri(n-octyl)phosphine oxide Trimethylphosphine oxide Triphenylphosphine oxide Diethyl methylphosphonate Triethyl phosphate Tri(n-butyl) phosphate Trimethyl phosphate Chlorodiphenylphosphine oxide Triphenyl phosphate Dichlorophenylphosphine oxide Phosphoryl chloride

Me3 N+ N− POPh2 [c-(CH2 )5 N]3 PO Et3 PO n-Bu3 PO (Me2 N)3 PO n-Oct3 PO Me3 PO Ph3 PO Me(EtO)2 PO (EtO)3 PO (n-BuO)3 PO (MeO)3 PO Ph2 ClPO (PhO)3 PO PhCl2 PO Cl3 PO

4.33a 3.74 3.66 3.63 3.60 3.59 3.53 3.16 2.81 2.68 2.66 2.50 2.17 1.89 1.26 0.56

−24.72 −21.35 −20.89 −20.72 −20.55 −20.49 −20.15 −18.04 −16.04 −15.30 −15.18 −14.27 −12.39 −10.79 −7.19 −3.20

29 30

Arsine oxides Triphenylarsine oxide Triethylarsine oxide

Ph3 AsO Et3 AsO

4.15 4.89a

−23.69 −27.91

a

Secondary values.

2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32

1

No.

Table 4.23

Methyl isothiocyanate Ethylenetrithiocarbonate Thiocamphor N,N-Dimethylthiocarbamoyl chloride Methyl N,N-dimethyldithiocarbamate O-Methyldimethylthiocarbamate N,N-Dimethylthiobenzamide N,N-Dimethylthioformamide Benzyl[(dimethylamino)methylene]dithiocarbamate Methyl[(dimethylamino)methylene]dithiocarbamate N-methylthioacetamide 4,N,N-Trimethylthiobenzamide N,N-Dimethyl-4-methoxythiobenzamide N,N-Dimethylthioacetamide N,N-Dimethylamino-N’-thiobenzoylformamidine N,N-Dimethyl-N’,N’-diethylthiourea N,N-Dimethyl-4-aminothiobenzamide N,N,N’ N -Tetramethylthiourea Benzyl[1-(dimethylamino)ethylidene]dithiocarbamate ε-Thiocaprolactam 3-[(Dimethylamino)methylene]-1,1-dimethylthiourea N-Methyl-N,N -propylenethiourea 3-[1-(Dimethylamino)ethylidene]-1,1-dimethylthiourea 1-Methyl-1,3-thiazolidine-2-thione 1,3-Thiazolidine-2-thione N,N -Dimethyl-N,N -ethylenethiourea 1,3-Oxazolidine-2-thione 1-Methyl-1,3-imidazolidine-2-thione 1,4,5-Trimethylthiazoline-2-thione 1,4,5-Trimethyloxazoline-2-thione 1,3,4,5-Tetramethylimidazoline-2-thione

Phenyl isothiocyanate

Lewis base

a

a

a

a

a

a a

a

Me2 NC(Me) NCSNMe2

a

Me2 NCH NCSNMe2

a

ClCSNMe2 MeSCSNMe2 MeOCSNMe2 PhCSNMe2 HCSNMe2 Me2 NCH NCSSCH2 Ph Me2 NCH NCSSMe MeCSNHMe 4-MeC6 H4 CSNMe2 4-MeOC6 H4 CSNMe2 MeCSNMe2 Me2 NCH NCSPh Et2 NCSNMe2 4-NH2 C6 H4 CSNMe2 Me2 NCSNMe2 Me2 NC(Me) NCSSCH2 Ph

a a

MeN C S

PhN C S

Formula

∆G◦ 3.14 3.60 0.29 −1.71 −1.94 −2.85 −4.00 −4.79 −5.82 −5.99 −6.28 −6.39 −6.51 −6.56 −6.62 −6.96 −7.02 −7.36 −7.59 −7.71 −8.50 −9.13 −10.22 −11.42 −11.76 −5.19 −7.53 −7.53 −7.88 −10.73 −6.45 −7.02 −10.22

pKBHX −0.55 (S) −0.63 (π ) −0.05 0.30 0.34 0.50 0.70b 0.84b 1.02 1.05 1.10 1.12 1.14 1.15 1.16 1.22 1.23 1.29 1.33 1.35 1.49 1.60 1.79 2.00 2.06b 0.91b 1.32b 1.32 1.38b 1.88b 1.13b 1.23b 1.79b

Hydrogen-bond basicity scales, pKBHX and ∆G◦ (kJ mol−1 ), of thiocarbonyl compounds [73] and isothiocyanates, in CCl4 at 25 ◦ C [87].

3

S

X

S

S

b Secondary

S

a Formulae:

33 34 35 36 37 38 39 40 41 42 43 44

R

R

values.

X

N

N

S

4

36 37 38 39 40 41

25 26 27 28 29

S

N

R Me Me H Me H H

R Me H Me H H

21

S

H

N

S

23

X S O S NMe O NMe

X S S NMe O NMe

H

Me

N

X

X

Me

S

S

N

N R

Me

R

42 43 44

30 31 32 33 34 35 R Me Me Me

R Me Me Me H H H

4,5-Dimethylthiazoline-2-thione 4,5-Dimethyloxazoline-2-thione 1,4,5-Trimethylimidazoline-2-thione 3-Methylbenzothiazole-2-thione 3-Methylbenzoxazole-2-thione Benzothiazole-2-thione 1,3-Dimethylbenzimidazole-2-thione Benzoxazole-2-thione 1-Methylbenzimidazole-2-thione 1-Methylthiazoline-2-thione 1,3-Dimethylimidazoline-2-thione 1-Methylimidazoline-2-thione, methimazole

X S NMe NH

X S O NMe S O NMe

a

a

a

a

a

a

a

a

a

a

a

a

1.81b 1.81b 2.17b 0.66b 0.78b 1.15b 1.19b 1.24b 1.53b 0.83b 1.55b 2.11b

−10.33 −10.33 −12.39 −3.77 −4.45 −6.56 −6.79 −7.08 −8.73 −4.74 −8.85 −12.04

162


Table 4.24 Equilibrium constants, Kc (t) (l mol−1 ), and hydrogen-bond basicity scales, pKBHX and ∆G◦ (kJ mol−1 ), for the complexes of 4-fluorophenol with sulfides, thiols and disulfides, in CCl4 at 25 ◦ C [92]. Lewis base

Formula

Kc (t)

pKBHX

∆G◦

1

Diphenyl sulfide

Ph2 S

0.74

2 3 4 5

Methyl disulfide Ethyl disulfide 1,3-Dithiolane Chloromethyl methyl sulfide

Me2 S2 Et2 S2 MeSCH2 Cl

0.65 0.79 0.85 0.69

6

Thioanisole

PhSMe

0.81

7 8 9 10

Bis(methylthio)methane Ethanethiol 1,4-Dithiane Ethyl vinyl sulfide

MeSCH2 SMe EtSH

1.23

EtSCH CH2

1.45 0.88

11

2-Chloroethyl methyl sulfide

MeSCH2 CH2 Cl

1.07

12 13 14

2-Propanethiol 1,3-Dithiane Benzyl methyl sulfide

a

15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31

Ethylene sulfide Allyl methyl sulfide Propylene sulfide Methyl sulfide Trimethylene sulfide Methyl n-propyl sulfide n-Butyl methyl sulfide Methyl n-octyl sulfide Ethyl methyl sulfide Pentamethylene sulfide n-Butyl sulfide Cyclohexyl methyl sulfide Ethyl sulfide tert-Butyl methyl sulfide Isopropyl sulfide Tetrahydrothiophene tert-Butyl sulfide

−0.93 (S)b −0.51 (π ) −0.49c −0.40c −0.38c −0.37 (S)b −0.59 (Cl) −0.32 (S)b −0.48 (π ) −0.22c −0.16 −0.14c −0.13 (S)b −0.82 (π ) −0.12 (S)b −0.52 (Cl) −0.10 −0.06c −0.02 (S)b −0.42 (π ) 0.03 0.08 0.10 0.12 0.13 0.16 0.17 0.17 0.18 0.23 0.23 0.24 0.25 0.25 0.30 0.32 0.40

5.31 2.91 2.80 2.28 2.17 2.11 3.37 1.83 2.74 1.26 0.91 0.80 0.74 4.68 0.68 2.97 0.57 0.34 0.11 2.40 −0.17 −0.46 −0.57 −0.68 −0.74 −0.91 −0.97 −0.97 −1.03 −1.31 −1.31 −1.37 −1.43 −1.43 −1.71 −1.83 −2.28

No.

a

a

a

i-PrSH PhCH2 SMe c-(CH2 )2 S CH2 CHCH2 SMe

a

Me2 S c-(CH2 )3 S n-PrSMe n-BuSMe n-OctSMe EtSMe c-(CH2 )5 S n-Bu2 S c-HexSMe Et2 S t-BuSMe i-Pr2 S c-(CH2 )4 S t-Bu2 S

Formulae: S

S S

S

S

S S

4 b c

9

13

17

Corrected for a second hydrogen-bond acceptor site. Statistically corrected to put the pKBHX value on a per sulfur basis.

1.74 1.33


163

Table 4.25 Hydrogen-bond basicity scales, pKBHX and ∆G◦ (kJ mol−1 ), of haloalkanes, in CCl4 at 25 ◦ C [83, 88, 91]. No. 1 2 3 4 5 6 7 8

Lewis base Fluoroalkanes 1,3-Difluoropropane 1-Fluoropentane 2-Fluoroadamantane 1,3-Difluoroadamantane 1-Fluorooctane Fluorocyclohexane 1-Fluoroadamantane syn-2,4-Difluoroadamantane

9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24

Chloroalkanes 1,1,1-Trichloroethane Dichloromethane 1,2,3-Trichloropropane 1,1-Dichloroethane 1,2-Dichloroethane 1,2-Dichloropropane 1,3-Dichloropropane trans-1,2-Dichlorocyclohexane 1,4-Dichlorobutane 1-Chlorobutane 1,5-Dichloropentane 1-Chloropentane 2-Chloropropane 2-Chloro-2-methylpropane Chlorocyclohexane 1-Chloroadamantane

25 26 27 28 29 30 31 32 33 34 35 36 37

Bromoalkanes Dibromomethane 1,2-Dibromoethane 1,3-Dibromopropane Bromocyclopropane 1,4-Dibromobutane Bromoethane 1-Bromopropane 1-Bromopentane 1-Bromobutane 2-Bromopropane Bromocyclohexane 2-Bromo-2-methylpropane 1-Bromoadamantane

38 39 40 41 42

Iodoalkanes Diiodomethane 1,2-Diiodoethane 1,3-Diiodopropane Iodomethane Iodoethane

Formula

pKBHX

∆G◦

FCH2 CH2 CH2 F CH3 (CH2 )4 F

−0.27b −0.06 −0.02 0.00b 0.02 0.09 0.31 0.48c

1.54 0.34 0.11 0.00 −0.11 −0.51 −1.77 −2.74

−1.15b −0.80b −0.76b −0.72b −0.61b −0.61b −0.51b −0.50b −0.47b −0.41 −0.39b −0.38 −0.30 −0.28 −0.27 −0.18

6.56 4.57 4.34 4.11 3.48 3.48 2.91 2.85 2.68 2.34 2.23 2.17 1.71 1.60 1.54 1.03

CH2 Br2 BrCH2 CH2 Br BrCH2 CH2 CH2 Br c-C3 H5 Br Br(CH2 )4 Br CH3 CH2 Br Br(CH2 )2 CH3 Br(CH2 )4 CH3 Br(CH2 )3 CH3 CH3 CHBrCH3 c-C6 H11 Br CH3 C(CH3 )BrCH3

−0.70b −0.63b −0.53b −0.47 −0.47b −0.40 −0.38 −0.35 −0.34 −0.30 −0.25 −0.22 −0.17

4.00 3.60 3.03 2.68 2.68 2.28 2.17 2.00 1.94 1.71 1.43 1.26 0.97

CH2 I2 ICH2 CH2 I ICH2 CH2 CH2 I CH3 I CH3 CH2 I

−0.68b −0.65b −0.51b −0.47 −0.47

3.88 3.71 2.91 2.68 2.68

a a

CH3 (CH2 )7 F c-C6 H11 F a a

CCl3 CH3 CH2 Cl2 ClCH2 CHClCH2 Cl Cl2 CHCH3 ClCH2 CH2 Cl ClCH2 CHClCH3 ClCH2 CH2 CH2 Cl a

Cl(CH2 )4 Cl Cl(CH2 )3 CH3 Cl(CH2 )5 Cl Cl(CH2 )4 CH3 CH3 CHClCH3 CH3 C(CH3 )ClCH3 c-C6 H11 Cl a

a

(Continued)

164



Lewis base

Formula

pKBHX

∆G◦

43 44 45 46 47 48

1,4-Diiodobutane 1-Iodopentane 2-Iodopropane 2-Iodo-2-methylpropane Iodocyclohexane 1-Iodoadamantane

I(CH2 )4 I I(CH2 )4 CH3 CH3 CHICH3 CH3 C(CH3 )ICH3 c-C6 H11 I

−0.46b −0.37 −0.37 −0.33 −0.32 −0.19

2.63 2.11 2.11 1.88 1.83 1.08

a

Formulae: F

X

F

F

3 b

a

4

7 24 37 48

Cl

X F Cl Br I

F

F

8

Statistically corrected to put the pKBHX value on a per halogen basis. c Not statistically corrected because of a three-centre hydrogen-bond.

Cl

16


165

Table 4.26 Hydrogen-bond basicity scales, pKBHX and ∆G◦ (kJ mol−1 ), of miscellaneous bases, in CCl4 at 25 ◦ C. No. Lewis base

Formula

pKBHX

∆G◦

Ref.

Thiophosphoroso and selenophosphoroso compounds Trimethyl thiophosphate (MeO)3 PS 0.58 Triethyl thiophosphate (EtO)3 PS 0.76 Triphenylphosphine selenide Ph3 PSe 0.94 Triphenylphosphine sulfide Ph3 PS 1.00 Hexamethylphosphorothioic (Me2 N)3 PS 1.33b triamide Triethylphosphine sulfide Et3 PS 1.46 Tri(n-octyl)phosphine sulfide n-Oct3 PS 1.54

−3.31 −4.34 −5.37 −5.71 −7.59

[21] [102] [96] [96] [96]

−8.33 −8.79

[21] [97]

8 9 10

Alkyl derivatives of group 15 elements Triethylstibine Et3 Sb Triethylarsine Et3 As Triethylphosphine Et3 P

−0.47b −0.13b 0.68b

2.68 0.74 −3.88

[96] [96] [96]

11 12

Alkyl derivatives of group 16 elements Methyl selenide Me2 Se n-Butyl telluride n-Bu2 Te

−0.01 0.03b

0.06c −0.17

[96] [96]

13 14 15

Imides N-Ethylmaleimide N-Methylsuccinimide Azobenzene

0.68c 1.06c ∼−0.22 (N) ∼−0.53 (π )

−3.88 −6.05 1.26 3.03

[96] [14] [96] [96]

16

Halobenzenes Bromobenzene

PhBr

17

Chlorobenzene

PhCl

18

Iodobenzene

PhI

∼−0.66 (π ) ∼−0.92 (Br) ∼−0.65 (π ) ∼−1.02 (Cl) ∼−0.64 (π ) ∼−0.92 (I)

3.77 5.25 3.71 5.82 3.65 5.25

[96] [96] [96] [96] [96] [96]

19

Polybases Progesterone

a

20

Nicotine

a

21

Nornicotine

a

22

Cotinine

a

1.75 (O3) 1.20 (O20) 2.03 (Nsp2 ) 1.11 (Nsp3 ) 2.09 (Nsp2 ) 1.52 (Nsp3 ) 2.16 (O) 1.62 (N)

−9.99 −6.85 −11.59 −6.34 −11.93 −8.68 −12.33 −9.25

[84] [84] [95] [95] [95] [95] [89] [89]

NBu4 + I− NBu4 + MeSO3 −

2.80 3.90

−15.98 −22.26

[14] [14]

NBu4 + Cl− NBu4 + N3 −

4.30 4.51

−24.54 −25.74

[14] [93]

1 2 3 4 5 6 7

23 24 25 26

Ion pairs Tetrabutylammonium iodide Tetrabutylammonium methylsulfonate Tetrabutylammonium chloride Tetrabutylammonium azide

a a

PhN NPh

(Continued)

166



Lewis base

Formula

27 28

Tetrabutylammonium cyanate Tetrabutylammonium acetate

NBu4 + OCN− NBu4 + MeCOO−

a

Ref.

5.08 5.60

−29.00 −31.96

[93] [14]

Formulae: 20

O

∆G◦

pKBHX

N

O

O

N

Et

Me

13

14

O

O

R 20 Me N

21 H

3

N

N

R

Me

O

N

22

O

19 b

Secondary values. c Statistically corrected to put the pKBHX value on a per carbonyl basis.

logical and easy to scan, the bases are divided into families according to the kind of HBA atom: for example, the family of oxygen bases. These are further divided into subfamilies according to the functionality of this atom, for example, the sub-family of ketones. Within each sub-family, the values depend on the effects of substituents on the function. To illuminate the connection between structure and hydrogen-bond basicity, the values are arranged either in an increasing (decreasing) order or grouped according to a given effect, for example, ring-substituted acetophenones are grouped together. The adopted ordering system gives the following contents for Tables 4.3–4.26: Table 4.3 – Carbon π bases: aromatic, acetylenic and ethylenic hydrocarbons [75]. Hindered pyridines [77]. For the π basicity of aromatic amines, aromatic ethers, aromatic thioethers and halobenzenes, see Tables 4.7, 4.15, 4.24 and 4.26 respectively. Table 4.4 – Nitrogen bases: primary amines [82]. Table 4.5 – Nitrogen bases: secondary amines [85]. Table 4.6 – Nitrogen bases: tertiary amines [86]. Table 4.7 – Nitrogen and π bases: aromatic amines [90]. Table 4.8 – Nitrogen bases: six-membered aromatic N-heterocycles [80], with the peculiar set of 1,10-phenanthrolines, which form three-centre hydrogen bonds. Table 4.9 – Nitrogen bases: five-membered aromatic N-heterocycles [96]. Table 4.10 – Nitrogen bases: amidines [60, 62, 64, 68]. Table 4.11 – Nitrogen bases: imines, oximes [94, 96]. Table 4.12 – Nitrogen bases: nitriles [66, 67, 72]. Table 4.13 – Oxygen bases: water, alcohols, phenols [61]. Table 4.14 – Oxygen bases: ethers, peroxides [81]. Table 4.15 – Oxygen and π bases: aromatic ethers [96]. Table 4.16 – Oxygen bases: aldehydes and ketones [79]. Table 4.17 – Oxygen bases: esters, lactones and carbonates [71]. Table 4.18 – Oxygen bases: amides, lactams, carbamates, ureas [65] and carboxamidates [69].


167

Table 4.19 – Oxygen bases: nitroalkanes, nitroaromatics, nitramines, nitramidates [70, 74] and nitroso compounds. Table 4.20 – Oxygen bases: sulfones, sulfate, sulfonate, sulfonamides, sulfonamidates [78]. Table 4.21 – Oxygen bases: sulfoxides, selenoxides, sulfites, sulfinamides [96]. Table 4.22 – Oxygen bases: oxides of organoderivatives of elements of group 15: phosphoryl bases, amine N-oxides, arsine oxides [94, 96]. Table 4.23 – Sulfur bases: thioketone, thioamides, thioureas [73] and isothiocyanates [87]. Table 4.24 – Sulfur bases: thiols, thioethers, disulfides [92], aromatic thioethers (sulfur and π bases) [96]. Table 4.25 – Halogen bases: fluoro-, chloro-, bromo- and iodoalkanes [83, 88, 91]. Table 4.26 – Miscellaneous bases: thiophosphoroso and selenophosphoroso compounds, alkyl derivatives of group 15 and 16 elements, imides, azobenzene, halobenzenes, polybases and ion pairs. It is highly desirable to determine the pK BHX values experimentally and obtain the socalled primary values. However, the pK BHX values of many bases are not accessible by direct measurement, because compounds are either commercially unavailable or not yet synthesized, or sparingly soluble in CCl4 or too strongly basic towards 4-fluorophenol. The last situation arises because the FTIR method does not give accurate values when more than 80% of 4-fluorophenol is hydrogen bonded, that is, when K c is greater than ∼5000 l mol−1 (pK BHX > 3.7). In order to overcome the solubility problem and to lower the equilibrium constant, the better solvent CH2 Cl2 and the weaker HBD methanol have been used instead of CCl4 and 4-fluorophenol, respectively. The excellent linear correlations, sometimes family dependent, between logK c values in the two solvents or with the two HBDs enable secondary pK BHX values to be calculated. In the literature, there are also homogeneous sets of complexation constants of families of bases with OH HBDs similar to 4-fluorophenol, for example phenol [49–58]. If precise linear correlations can be established between logK c values of bases common to the 4-fluorophenol and phenol sets, these literature data can be used to estimate secondary pK BHX values for bases not measured with 4-fluorophenol. Other methods for obtaining secondary pK BHX values are correlations with OH infrared shifts upon hydrogen bonding, or with σ substituent constants, or with quantum chemical descriptors of hydrogen-bond basicity such as electrostatic potentials or electronic energy variations upon complexation. In all cases, the accuracy is lower for secondary than for primary pK BHX values, and can be estimated from the statistics of the linear regression, for example the standard error of the fit. 4.3.5

Range of Validity of the Scale

Delimitation of the range of validity of the pK BHX scale is a crucial issue for its use in hydrogen bonding studies. Both the solvent and the reference HBD validity ranges need to be specified. Carbon tetrachloride has been selected for pK BHX measurements by means of the FTIR method because it is devoid of significant HBA ability and is transparent to IR light in the OH stretching zone. This solvent has a relative permittivity ε of 2.23 and its Onsager function, (ε − 1)/(2ε + 1), which ranks the strength of the solvent reaction field, is 0.23. However, physical, chemical and biological surroundings of relevance for HB interactions can range all the way from the gas phase (ε = 1) to the aqueous phase (ε = 78), corresponding to

168


values for the Onsager function from 0 to 0.49. It is necessary to address the question of how the ranking order of HBA strength may be affected by such variations. The study [9] of the hydrogen bonding of 4-fluorophenol (reaction 4.2) in several solvents of relative permittivity ranging from 2.02 (cyclohexane) to 10.36 (1,2-dichloroethane) (Onsager function from 0.2 to 0.43) shows that linear free energy relationships (logK c in a given solvent versus pK BHX in CCl4 ) are always obeyed by oxygen and sp-hybridized nitrogen bases. However, sp2 and sp3 nitrogen bases gain basicity relative to oxygen and sp nitrogen bases as the reaction field rises. This has been explained [9] by an increased extent of proton sharing in hydrogen-bonded complexes, permitted by the action of polar solvents. However, the resulting increase in basicity does not exceed 4 kJ mol−1 . The pK BHX database has been scaled to 4-fluorophenol but other HBDs are chemically and biochemically relevant. It is important to know if CH, NH (and even NH+ ) and other OH donors rank the bases in the same order as, or different from, 4-fluorophenol. The study [8, 12, 13] of the existence of linear free energy relationships between pK BHX , or a very similar quantity, and logK c of equilibrium 4.1 for various hydrogen-bond donors shows that (i) OH donors (phenols, alcohols and water), (ii) NH+ donors such as n-Bu3 NH+ and (iii) strong NH donors such as amides and imides rank all the bases in the same way. However, CH donors, such as chloroform and alk-1-ynes, and weak NH donors, such as aliphatic amines, anilines and pyrroles, rank only the oxygen and the sp nitrogen bases in the same order. The sp2 and sp3 nitrogen bases lose basicity relative to oxygen bases as the hydrogen-bond acidity of the hydrogen-bond donor falls. In conclusion, the pK BHX scale is the most reliable and comprehensive hydrogen-bond basicity scale so far measured. Although not universal towards all HBDs in all solvents, it is a reasonably general scale. In particular, the pK BHX scale may be applied to the hydrogenbonded complexes of water, the most ubiquitous hydrogen-bond donor encountered in physics, chemistry and biology. Moreover, by extension of the comparison of 4-fluorophenol with strong NH donors such as amides, this scale can be applied to the hydrogen-bonded complexes of the peptide backbone of proteins, that is, to many recognitions of drugs by proteinic receptors [84].

4.4

Hydrogen-Bond Affinity Scales: Early Studies

In spite of a large number of determinations of hydrogen-bond enthalpies, particularly in the period 1960–1980, there were only two attempts to build a scale of hydrogen-bond affinity. The earlier one was made by Arnett and co-workers [10, 11], using two different methods. The first is referred to as the ‘high-dilution calorimetric method’. In this case, a small quantity of base is injected into a dilute CCl4 solution of 4-fluorophenol. The observed heat of reaction is corrected for the heat due to the solution of the base in pure CCl4 . It is related to the enthalpy of hydrogen-bond formation (∆H ◦ ) by the expression Q = ∆H ◦ [C] V

(4.15)

where V is the volume in litres of the solution in the calorimeter and [C] is the equilibrium molar concentration of the hydrogen-bonded complex, which is calculated from the equilibrium constant K c . Values of K c are derived from separate determinations by 19 F


169

NMR and/or IR spectrometry. Forty values were obtained in this way, spanning from −8.0 ± 1.7 kJ mol−1 for diphenyl ether to −39.8 ± 0.8 kJ mol−1 for quinuclidine. Values of ∆H ◦ determined by this dilution calorimetric method are very sensitive to errors in K c , especially if K c is low. In order to sidestep the determination of K c , a different approach was also used. In this alternative method, referred to as the ‘pure-base calorimetric method’, the base is used as the solvent, so complete complexing of the 4-fluorophenol and base can usually be assured. The two heat terms involved are (i) the heat due to the hydrogen-bond formation and (ii) the heat produced by other interactions. The latter contribution is obtained by using 4-fluoroanisole, which resembles 4-fluorophenol in all respects except in its ability to hydrogen bond. The enthalpy is determined through Equation 4.16 by measuring heats of solution, ∆H s , for 4-fluorophenol and 4-fluoroanisole in the pure base and the reference inert solvent CCl4 . ∆H ◦ = [∆Hs (4-FC6 H4 OH) − ∆Hs (4-FC6 H4 OMe)]base − [∆Hs (4-FC6 H4 OH) − ∆Hs (4-FC6 H4 OMe)]CCl4

(4.16)

Fifty-seven values were obtained by the pure-base method, covering a range from −4.9 ± 1.0 kJ mol−1 for thionyl chloride Cl2 SO, a weak sulfuryl base, to −37.3 ± 0.4 kJ mol−1 for triethylamine NEt3 , a strong nitrogen base. The degree of agreement between the pure-base method and the high-dilution method can be tested on a sample of 19 bases that have been studied by both methods. It is found that eight values of ∆H ◦ are almost identical within experimental errors. The 11 values that differ by more than the experimental errors are for (i) N-methylformamide, a selfassociated liquid, (ii) four bases that have strong relative permittivity (29 to 46.7), (iv) five pyridines and (v) diphenyl ether, a polyfunctional base (π and O sites). Therefore, the discrepancy between the hydrogen-bond enthalpies obtained by the two methods is not statistically distributed and is probably real. Consequently, the solute scale obtained by the high-dilution method and the solvent scale obtained by the pure-base method should not be mixed. The solvent scale must be reserved for the hydrogen-bond affinity term occurring in multiparameter equations developed for explaining and predicting solvent effects in chemistry [103]. Later, a new hydrogen-bond affinity scale was constructed by Catalan et al. [17] for 35 solvents by means of a ‘pure-solvent calorimetric method’. The reference hydrogenbond donor is pyrrole and again the base is used as the solvent. The heat produced by interactions other than hydrogen bonding is obtained by using not only N-methylpyrrole but also benzene and its C-methylated analogue, toluene, in order to take into account the increase in volume, basicity and polarity caused by N-methylation. The enthalpy is determined through Equation 4.17 by measuring heats of solution, ∆H s , for pyrrole, Nmethylpyrrole, benzene and toluene in the pure base and heats of vaporization, ∆H v , of the same compounds. ∆H ◦ = [∆Hs (pyrrole) + ∆Hs (toluene)]base − [∆Hs (N -methylpyrrole) + ∆Hs (benzene)]base − [∆Hv (pyrrole) + ∆Hv (toluene) − ∆Hv (N -methylpyrrole) − ∆Hv (benzene)]

(4.17)

The last term is constant and has the value 8.78 kJ mol−1 . In this method, no measurement in a reference inert solvent is required and the enthalpies of hydrogen bonding are anchored

170


to the gas phase. A good linear relationship is found between ∆H ◦ (4.16) for hydrogen bonding of 4-fluorophenol and ∆H ◦ (4.17) for hydrogen bonding of pyrrole: ∆H ◦ (4.17) = 0.524 (±0.035) ∆H ◦ (4.16) −2.41 (±0.86)

(4.18)

n = 13, r = 0.976, s = 1.45 kJ mol−1 As expected, the correlation coefficient increases (to 0.987) and the constant term decreases (to 1.49) if the data of Catalan et al. are treated by the ‘pure-base calorimetric method’ of Arnett (Equation 4.19): ∆H ◦ = [∆Hs (pyrrole) − ∆Hs (N -methylpyrrole)]base − [∆Hs (pyrrole) − ∆Hs (N -methylpyrrole)]CCl4

(4.19)

The above scales refer to a too narrow sampling of Lewis bases to be of wide use in chemistry. Moreover, the scales measured by the pure-base or pure-solvent methods are evidently limited to liquid bases. Therefore, Laurence, Berthelot and co-workers (University of Nantes) decided to extend the solute ∆H ◦ scale of Arnett. Their method and results [15, 16, 92, 95, 96] are presented below.

4.5

The 4-Fluorophenol Affinity Scale

In Section 4.3, we reported the equilibrium constants K c for the hydrogen bonding of 4-fluorophenol to Lewis bases in CCl4 at 25 ◦ C. These values allow the calculation of the Gibbs energy of this reaction, ∆G◦ , by means of Equation 4.20: ∆G ◦ = −RT lnK c

(4.20)

(see Chapter 1 for standard state). As shown by Equation 4.21: ∆G ◦ = ∆H ◦ − T∆S ◦ ◦

◦

(4.21)

◦

∆G is determined by ∆H , ∆S and T, which are the enthalpy, entropy and absolute temperature of the hydrogen-bond formation, respectively. To understand the thermodynamics of hydrogen-bond basicity, one must know whether this basicity is controlled by enthalpy, by entropy or by both. If it is the last case, the respective contributions of the enthalpic and entropic terms need to be determined. Unlike Arnett and Catalan, Laurence et al. chose to measure the enthalpy of hydrogen bonding from the temperature dependence of the equilibrium constant. The variation of an equilibrium constant with temperature is given by Equation 4.22: ∆H ◦ d lnK x (4.22) = dT RT 2 (called the van’t Hoff or Gibbs–Helmholtz equation when related to K p , the pressure equilibrium constant of a chemical equilibrium in a mixture of ideal gases) (see Chapter 1 for the choice of K x instead of K c ). Since d(T −1 ) = −T −2 dT, Equation 4.22 can be written as ∆H ◦ dlnK x =− (4.23) d(1/T ) R


171

ln Kx

10

6 3

1000 K / T

4

Figure 4.6 Plot of lnKx versus 1/T for the hydrogen bonding of 4-fluorophenol to trimethyl phosphate in CCl4 over the temperature range −6 to +55 ◦ C. The circles on the least-squares line represent experimental data (Table 4.27).

At a given temperature, the slope of the graph of lnKx versus 1/T multiplied by −R equals ∆H ◦ at that temperature. If ∆H ◦ is essentially constant over the temperature range plotted, the graph is a straight line. Laurence et al. carried out measurements on the temperature dependence of the equilibrium constants of reaction 4.2 over 60 ◦ C (from −5 to +55 ◦ C) in CCl4 , over 75 ◦ C (from −5 to +70 ◦ C) in C2 Cl4 , and over 55 ◦ C (from +10 to +65 ◦ C) in cyclohexane. The temperature range is governed by the boiling and melting points of solvents, solubilities at low temperatures and other technical problems. For five temperatures (for example −5, +10, +25, +40 and +55 ◦ C in CCl4 ), the squared correlation coefficient of the linear regression of lnKx on 1/T is generally greater than 0.9990 and no curvature has ever been detected. Figure 4.6 shows such a plot for the reaction 4-FC6 H4 OH + (MeO)3 PO 4-FC6 H4 OH · · · OP(OMe)3 over the temperature range −6 to +55 ◦ C. The data and the results of the regression analysis are given in Table 4.27. They show that ∆H ◦ is not sensitive to temperature in the studied temperature range. Other measurements for more negative temperatures are collected in Table 4.28. They show very good agreement between the values obtained over the ranges −20 to +25◦ C [7] and −20 to + 58 ◦ C [104] and those obtained over the range −5 to +55 ◦ C. We conclude that the hydrogen-bond enthalpies of 4-fluorophenol show no measurable temperature dependence over the ranges −20 to +58 ◦ C in CCl4 and −5 to +70 ◦ C in C2 Cl4 . There is a widespread opinion that calorimetrically determined enthalpy values are, generally, more reliable than those obtained from the temperature coefficient of equilibrium constants. For example, Ahrland [105] wrote the following about the enthalpy of complexation of metallic cations by anionic ligands in water: ‘An inherently much better way of determining ∆H ◦ is by direct calorimetric measurement. This method has been increasingly used in later years, as the futility, in many cases, of the temperature coefficient method has been realized’. In order to refute this assertion, Table 4.29 presents a comparison of van’t Hoff and calorimetric hydrogen-bond enthalpies of 4-fluorophenol in CCl4 for 27 diverse Lewis bases. It is seen that the differences between the enthalpies determined by the two methods are statistically distributed since the mean algebraic difference of −0.36 kJ mol−1 is close to zero, while the mean absolute difference of only 1.18 kJ mol−1 is mostly

172


Table 4.27 Determination of the enthalpy of hydrogen bonding of trimethyl phosphate (initial mole fraction: 6.5 × 10−4 ) to 4-fluorophenol (initial mole fraction: 4 × 10−4 ) in CCl4 . T(◦ C) Kx Regression analysis Slope Squared correlation coefficient Fisher test

−6.1 11113

8.9 5792

24.5 3145

39.4 1864

54.6 1128

3293 ± 11 0.999 97

Intercept Standard error

−3.012 ± 0.037 0.006

90 584

Number of data

5

Enthalpy ∆H◦ = (3293 ± 11) × (−8.3145) × 10−3 = −27.38 ± 0.09 kJ mol−1a A second determination with x◦ (trimethyl phosphate) = 8.8 × 10−4 and x◦ (4-fluorophenol) = 4 × 10−4 , gives ∆H◦ = –26.73 ± 0.22 kJ mol−1 .

a

within experimental errors. Interestingly, the van’t Hoff determined enthalpy for 1-methyl2-pyrrolidone (−27.36 kJ mol−1 ) is situated between two calorimetric values (−29.29 and −26.61 kJ mol−1 ). We conclude that good determinations of hydrogen-bond enthalpies by the van’t Hoff equation are as reliable as calorimetric measurements. A necessary condition for making such good determinations is to use a sufficiently wide interval of temperature. Very often this condition has not been fulfilled in the past. A safe and quick procedure for determining the temperature coefficient of hydrogenbonding equilibrium constants is to follow the absorbance of the ν(OH) band of 4fluorophenol in the FTIR spectrum of a single solution of 4-fluorophenol and base in CCl4 as a function of temperature. This FTIR method has been fully described in the literature [15] and [16]. It is adapted from the UV method first proposed by Joesten and Drago [106]. An important advantage of IR over UV measurements is that the free 4-fluorophenol band (at 3614 cm−1 ) is not overlapped by the hydrogen-bonded 4-fluorophenol band (shifted to lower frequencies). Thus, the equilibrium concentration of free 4-fluorophenol is directly obtained by reading the free OH band absorbance and applying the Beer–Lambert law.

Table 4.28 Comparison between values of hydrogen-bond enthalpies (kJ mol−1 ) of 4-fluorophenol in CCl4 obtained by the van’t Hoff equation over different temperature intervals. −∆H◦

Lewis base 3-Bromopyridine Tetrahydrofuran Ethyl acetate 1,4-Dioxane N,N-Dimethylacetamide

−20 to +25 ◦ C 25.7 24.0 20.5 20.0 −20 to +58 ◦ C 28.9

Difference −5 to +55 ◦ C 24.8 23.7 20.8 20.1

+0.9 +0.3 −0.3 −0.1

29.2

−0.3

Acetonitrile 3,5-Dichloropyridine Diethyl ether Tetrahydrofuran Cyclohexanone 3-Bromopyridine Diphenyl sulfoxide Methyl phenyl sulfoxide Trimethyl phosphate Triethyl phosphate N,N-Dimethylformamide Dimethyl sulfoxide 2,6-Dimethyl-γ -pyrone N,N-Dimethylbenzamide

b

19.40 23.86 24.20 23.74 21.36 24.81 24.86 25.82 27.01 28.42 24.85 27.82 28.89 26.82

−∆H◦ (vH) +1.83 +1.27 +0.77 +0.31 −2.91 −1.13 −1.08 −0.54 −0.19 +1.22 −2.76 +0.21 +0.02 −2.05

D

a

D = |∆H◦ (vH)| – |∆H◦ (cal)|. 1,2-Dichlorobenzene was used as solvent. Not included in the calculation of means.

17.57 22.59 23.43 23.43 24.27 25.94 25.94 26.36 27.20 27.20 27.61 27.61 28.87 28.87

Lewis base

a

−∆H◦ (cal) Dibutyl sulfoxide 1-Methyl-2-pyrrolidone Pyridine 4-Methylpyridine Quinoline Cyclopropylamine Pyridine-N-oxide Trimethylphosphine oxide Tetramethylurea 4-Dimethylaminopyridine 2,6-Dimethylpyridine 2,4,6-Trimethylpyridine Hexamethylphosphoramide Quinuclidine

Lewis base 28.87 29.29 29.71 30.54 30.75 31.38 31.38 32.22 32.64 32.64 32.64 33.05 33.47 39.75b

−∆H◦ (cal)

28.81 27.36 29.62 30.83 30.08 31.88 28.27 31.98 28.61 34.11 33.22 35.35 33.74 37.68

−∆H◦ (vH)

Table 4.29 Hydrogen-bond enthalpies (kJ mol−1 ) of 4-fluorophenol in CCl4 . Agreement between values determined by the van’t Hoff equation (vH) and calorimetric values (cal).

b

−0.06 −1.93 −0.09 +0.29 −0.67 +0.50 −3.11 −0.24 −4.03 +1.47 +0.68 +2.30 +0.27

Da

174


The FTIR spectra of such a single solution are recorded at five temperatures between −5 and +55 ◦ C. The method requires the preliminary determination of the temperature dependence of the absorption coefficient, ε, of the ν(OH) IR band. This quantity falls off linearly with increasing temperature according to Equation 4.24, in which the temperature t is in ◦ C: ε(t) = ε(25) − 0.624 (t − 25)

(4.24)

The experimental result is controlled in three ways. First, the accuracy of the initial 4fluorophenol solution concentration is tested by a calculation of ε, which must fall within the limits 238 ± 5 l mol−1 cm−1 at 25 ◦ C. Second, the equilibrium constant calculated at 25 ◦ C must match the value found during the construction of the pK BHX scale. In the latter work, K c was the average of five or six determinations in which the base concentration was varied in order to complex various quantities of 4-fluorophenol. A new solution of base is made if the two K c values obtained from the concentration and temperature variation methods differ by more than 10%. Third, the least-squares treatment of the van’t Hoff plot must give a squared correlation coefficient greater than 0.9990. About 300 values of hydrogen-bond enthalpies have been obtained by this procedure [15, 16, 88, 91, 92, 95, 96]. They are gathered in Table 4.30, which also contains values of entropies. These are calculated from the intercept of the van’t Hoff plot as shown by Equation 4.25: lnK x = −

∆H ◦ 1 ∆Sx◦ + R T R

(4.25)

The applicability of the 4-fluorophenol affinity scale to other hydrogen-bond donors is now investigated in the light of the existing literature data for the following donors, mostly in CCl4 : r OH donors: methanol [96], cyclohexanol [107], phenol [6], trifluoroethanol [108], hexafluoro-2-propanol [109] and perfluoro-tert-butanol [110]. There are also two data sets [10] for 1-butanol and phenol in the pure base medium. They will be compared with the 4-fluorophenol data obtained by the same pure-base method (see Section 4.4). r NH donors: pyrrole [96], isocyanic acid (HNCO) [111] and isothiocyanic acid (HNCS) [112, 113]. r NH+ donor: tri-n-butylammonium picrate (in 1,2-dichlorobenzene and chlorobenzene) [114, 115]. r CH donors: chloroform [116–119] (in cyclohexane), and cyanoacetylene [96]. The existence of linear enthalpy relationships is examined through Equation 4.26: −∆H ◦ (XH donor) = m [−∆H ◦ (4-FC6 H4 OH)] + c

(4.26)

where m and c are constants characteristic of the reference XH hydrogen-bond donor and the solvent. Table 4.31 summarizes the linear enthalpy Equation 4.26 results, by listing values of m and c, of the correlation coefficient r, of the standard deviation s (in kJ mol−1 ), of the number of Lewis bases n and by indicating the families of bases upon which the relationship is based. It is found that the hydrogen-bond enthalpies of (i) all OH donors studied, (ii) strong NH donors (HNCO and HNCS) and (iii) the NH+ donor are well correlated by the 4-fluorophenol affinity scale. The correlation coefficients range from 0.935

16.59 30.23 30.47 31.01

Secondary amines 1,1,1,3,3,3-Hexamethyldisilazane Bis(2-chloroethyl)amine 2-(3-Trifluoromethylphenyl)pyrrolidine N-Methylpropargylamine

6.18 9.52

Alkenes 1,4-Cyclohexadiene Cyclohexene

26.10 27.80 28.41 30.63 29.61 31.49 31.88 31.06

6.04

Thiophene Thiophene

Ammonia, primary amines (in C2 Cl4 ) 2,2,2-Trifluoroethylamine 3,5-Dichlorobenzylamine 3-(Trifluoromethyl)benzylamine Propargylamine 3-Fluorobenzylamine Ammonia (in CCl4 ) Cyclopropylamine Benzylamine

7.88 8.11 8.79

Methylbenzene, toluene 1,4-Dimethylbenzene, p-xylene 1,3,5-Trimethylbenzene, mesitylene

11.74 11.12

7.32

Arenes Benzene

Alkynes 1-Hexyne 3-Hexyne

−∆H◦

Lewis base

45.6 59.0 56.2 53.0

53.9 50.0 49.0 53.1 50.2 53.1 54.9 49.2

22.7 18.6

15.2b 27.0

10.5

13.8 13.6 14.5

13.4

−∆Sx◦ a

[16] [16] [16] [16]

[16] [99] [99] [16] [99] [16] [16] [16]

[96] [96]

[96] [96]

[92]

[96] [96] [96]

[96]

Ref.

N-Methylphenethylamine 1,2,3,6-Tetrahydropyridine N-Methylisopropylamine N-Methyl-tert-butylamine

3-Methoxybenzylamine 3-Methylbenzylamine Ethylamine Methylamine n-Butylamine (in CCl4 ) tert-Butylamine (in CCl4 ) Cyclohexylamine (in CCl4 )

2-Hexyne

1-Heptene

1,2,3,4-Tetramethylbenzene, prehnitene Pentamethylbenzene Hexamethylbenzene

Lewis base

34.92 32.08 35.01 35.09

32.34 31.69 33.84 35.13 33.70 34.20 34.26

11.72

6.77

10.32 10.77

9.10

∆H◦

57.4 46.8 55.6 56.1

52.7 49.8 51.0 56.8 52.2 52.7 51.1

21.7

18.2

16.7 16.7

13.5

[16] [16] [16] [16]

[16] [16] [16] [16] [16] [16] [16]

[96]

[96]

[96] [96]

[96]

Ref.

(Continued)

−∆Sx◦ a

Table 4.30 Enthalpies (kJ mol−1 ) and entropies (J K−1 mol−1 ) of hydrogen bonding between 4-fluorophenol and Lewis bases in CCl4 (C2 Cl4 , c-C6 H12 ).

23.83 27.15 (23.23)c

Tertiary amines N,N-Diisopropyl-3-pentylamine N,N-Diisopropylisobutylamine Tripropargylamine

N-Methyl-2-(3trifluoromethyl)pyrrolidine N-Ethyl-N-methyl-3(trifluoromethyl)benzylamine (in C2 Cl4 ) N,N-Diisopropylethylamine N-Methyl-2-(3fluorophenyl)pyrrolidine N,N-Dicyclohexylethylamine N-Ethyl-N-methyl-3fluorobenzylamine (in C2 Cl4 ) N,N-Dimethyl-3(trifluoromethyl)benzylamine (in C2 Cl4 ) 1,2,2,6,6-Pentamethylpiperidine

59.3 52.0b

35.56 33.32

60.6 56.2 74.0 63.8 71.4 56.0 54.8 70.8

29.26 28.24 33.89 31.01 33.60 28.82 28.63 34.03

67.1 65.8 42.0

53.1 58.3 58.4 70.3 55.2 62.9 55.9 52.2

31.09 32.80 33.54 37.23 33.27 35.97 34.07 33.10

2-(3-Fluorophenyl)pyrrolidine Diallylamine N-Methylbenzylamine (in C2 Cl4 ) 2,2,6,6-Tetramethylpiperidine 2-Phenylpyrrolidine Diisopropylamine N-Methylallylamine 1,2,3,4-Tetrahydroisoquinoline (in C2 Cl4 ) Di-n-butylamine Piperazine

−∆Sx◦ a

−∆H◦

Lewis base


[16]

[99]

[16] [99]

[16] [16]

[99]

[16]

[16] [16] [16]

[16] [16]

[16] [16] [16] [16] [16] [16] [16] [16]

Ref.

N,N,N ,N -Tetramethylhexane-1,6diamine

N-Methyl-1,2,3,6-tetrahydropyridine N,N,N ,N Tetramethylethylenediamine (in C2 Cl4 ) N-Butylpyrrolidine

3-Chloroquinuclidine Triethylamine

N,N-Dimethylallylamine

4-Chloro-N-methylpiperidine 3-Chloromethyl-N-methylpiperidine N-Methyl-1,2,3,4tetrahydroisoquinoline N,N -Dimethylpiperazine

Pyrrolidine (in C2 Cl4 ) Azetidine

N-Methylcyclohexylamine N-Methylbutylamine Azepane, hexamethyleneimine N-Methylethylamine Diethylamine Dimethylamine (in C2 Cl4 ) N,N’-Dimethylethylenediamine Piperidine

Lewis base

35.15

34.53

32.45 35.49

34.11 39.35

59.0b

57.3

51.5 60.6b

57.6 76.4

53.8

[16]

[16]

[16] [16]

[16] [16]

[16]

[16]

55.1b 32.93 32.84

[16] [16] [16]

[16] [16]

[16] [16] [16] [16] [16] [16] [16] [16]

Ref.

56.1 57.2 49.7

53.1 51.0

55.7 56.2 55.0 53.7 57.7 56.2 55.8b 56.4

−∆Sx◦ a

32.18 32.87 30.83

36.09 35.67

35.21 35.38 35.08 34.69 35.74 35.13 35.63 35.98

∆H◦

59.8 80.5 44.5 61.4 48.7

32.45 38.40 (28.58)c 33.13 29.27 (17.15)c 19.76 23.86 22.56 24.50 23.67 24.09 23.93 26.63 24.81 25.44 27.18 25.89 31.72 27.56 27.84

Six-membered N-heteroaromatics 2,6-Difluoropyridine 1,3,5-Triazine 3,5-Dichloropyridine Pyrazine 2-Fluoropyridine 2-Bromopyridine 2-Chloropyridine Pyrimidine 2,5-Dimethylpyrazine 3-Bromopyridine 3-Fluoropyridine

3-Chloropyridine 3-Iodopyridine 2-tert-Butylpyridine 2-Phenylpyridine 4,6-Dimethylpyrimidine

47.0 41.8 60.1 46.0 45.9b

34.5 40.9b 45.1 38.4b 44.9 40.1 40.9 40.6b 45.4b 38.0 40.1

[16] [16] [16] [99]

43.2b 65.0 64.0 64.1

26.15 32.58 32.79 33.07

[95] [96] [96] [96] [96]

[96] [96] [95] [96] [96] [96] [96] [96] [96] [96] [95]

[16] [16]

[16] [16] [96]

[99]

56.6

29.73

N,N-Dimethyl-3-fluorobenzylamine (in C2 Cl4 ) Hexamethylenetetramine Triallylamine N-Methyl-2-phenylpyrrolidine N-Ethyl-N-methylbenzylamine (in C2 Cl4 ) N,N-Dimethyl-3-chloropropylamine Tri-n-butylamine N-Methylmorpholine (O and N complex) N,N-Dimethylbenzylamine N,N-Dimethylpropargylamine Pyridine Quinoline Isoquinoline 2-Ethylpyridine 3-Methylpyridine, 3-picoline 2-Methylpyridine, 2-picoline 4-Methylpyridine, 4-picoline 2,6-Dimethylpyridine, 2,6-lutidine 3,5-Dimethylpyridine, 3,5-lutidine 2,4-Dimethylpyridine, 2,4-lutidine 2,4,6-Trimethylpyridine, 2,4,6-collidine 4-Aminopyridine 4-Methylaminopyridine 4-N,N-Dimethylaminopyridine 4-Pyrrolidinopyridine (in C2 Cl4 ) 2-Methoxypyridine

Quinuclidine

N-Methylpyrrolidine (in C2 Cl4 ) Diazabicyclooctane (DABCO) Tropane

N-Methylpiperidine Trimethylamine N,N-Dimethylcyclohexylamine N,N-Dimethylethylamine

N,N-Dimethylisopropylamine

32.66 33.51 34.10 36.30 25.77

29.62 30.08 29.66 34.01 30.01 30.51 30.83 33.32 31.89 31.82 35.35

37.68

34.79 33.20 35.66

34.04 33.48 34.75 34.47

34.67

41.9 42.0 42.1 42.5 48.7

44.2 45.2 43.4 58.3 42.5 44.6 43.9 52.3 45.8 45.0 55.5

55.8

[96] [96] [96] [96] [96]

[95] [96] [96] [95] [95] [96] [96] [96] [95] [96] [96]

[16]

[16] [16] [16]

[16] [16] [16] [16]

[16]

(Continued)

54.8 47.9b 55.3

54.4 52.4 56.1 54.6

55.9

17.1 40.6

12.20 30.50 37.52

31.08

Amidines Tetramethylguanidine

Imines N-Benzylidene-tert-butylamine

46.3

29.03 16.25 15.83 17.53 16.63 17.50 18.10 17.82 19.29 18.20 15.49 17.60 19.45 20.10 23.76

Nitriles Chloroacetonitrile 4-Chlorobenzonitrile Acrylonitrile Methyl thiocyanate Benzonitrile 2,6-Dimethylbenzonitrile Allyl cyanide Acetonitrile Propionitrile

Ethers Hexamethyldisiloxane 1,3-Dioxolane 1,3-Dioxane 1,4-Dioxane Di-n-butyl ether 43.0 31.4 33.7 33.8 43.4

27.1 20.7 25.7 22.4 24.0 24.8 23.6 28.2 23.8

45.6

28.00

2-(3-Trifluoromethylphenyl)-1pyrroline 2-(3-Fluorophenyl)-1-pyrroline

60.1

45.3

[96]

42.5b

27.86

Pyridazine Five-membered N-heteroaromatics 1-Methylpyrrole 5-Bromo-1-methylimidazole

[96] [96] [96] [96] [96]

[41] [41] [41] [96] [41] [96] [96] [41] [96]

[99]

[99]

[96]

[96]

[96] [96]

Ref.

−∆Sx◦ a

−∆H◦

Lewis base


Cyclohexene oxide tert-Butyl methyl ether 3-Methyltetrahydropyran Tetrahydropyran Tetrahydrofuran

4-Methoxybenzonitrile Isobutyronitrile Trimethylacetonitrile 4-Dimethylaminobenzonitrile Dimethylcyanamide 1-Piperidinecarbonitrile Diethylcyanamide trans-3-Dimethylaminoacrylonitrile N,N-Dimethyl-N -cyanoformamidine

2-Methyl-1-pyrroline

Benzophenone imine, diphenylketimine 2-Phenyl-1-pyrroline

1,5-Diazabicyclo[4.3.0]non-5-ene (DBN)

1-Methylimidazole

1,10-Phenanthroline

Lewis base

21.35 25.86 23.25 23.29 23.74

18.10 18.14 18.47 20.46 22.36 21.71 21.88 23.49 24.52

34.18

30.96

30.08

40.78

34.03

36.07

∆H◦

31.2 45.2 35.2 34.8 35.5

22.5 22.7 23.7 25.1 26.7 23.2 23.3 26.9 24.0

46.3

46.7

47.0

42.3

43.1

42.2

−∆Sx◦ a

[96] [96] [96] [96] [96]

[41] [96] [41] [41] [41] [96] [96] [41] [41]

[96]

[99]

[87]

[96]

[96]

[96]

Ref.

18.63 12.15 12.41 14.52 20.30 20.55 21.36 20.76

Aldehyde Benzaldehyde

Aliphatic ketones 1,1,1-Trifluoropropan-2-one 1,1,1-Trichloropropan-2-one 1,1-Dichloropropan-2-one 2,2,4,4-Tetramethylpentan-3-one 2,4-Dimethylheptan-3-one

Cycloalkanone Cyclohexanone

Acetophenone Acetophenone

Esters, carbonate Methyl formate Ethyl formate Dimethyl carbonate Methyl benzoate 17.96 17.95 17.77 19.23

23.43 25.26

23.0

Aromatic ether 1,2-Dimethoxybenzene

Miscellaneous conjugated ketones Anthrone 3-Methyl-5,5-dimethylcyclohexenone, isophorone

24.09 26.63 26.53

Diethyl ether tert-Butyl ethyl ether Diisopropyl ether

28.7 26.4 24.4 27.6

38.3 31.8

28.7

25.5

25.1 21.4 25.6 29.6 27.7

27.3

36.2

41.5 49.5 48.7

[96] [96] [96] [96]

[10] [96]

[96]

[96]

[96] [96] [96] [96] [96]

[96]

[96]

[96] [96] [96]

Methyl acetate Methyl 3,3-dimethylacrylate Ethyl acetate Ethyl 3-dimethylaminoacrylate

2,6-Dimethyl-γ -pyrone 3-Dimethylamino-5,5dimethylcyclohexenone

Pentan-3-one Propan-2-one, acetone 3-Methylbutan-2-one Nonan-5-one Butan-2-one

Oxetane, trimethylene oxide Cineole, eucalyptol 2,2,5,5-Tetramethyltetrahydrofuran

20.79 19.99 20.76 27.00

28.89 32.49

21.02 22.40 21.66 20.98 21.05

23.33 28.55 28.42

30.6 28.3 28.5 30.8

29.1 33.7

28.8 32.0 30.2 28.4 28.4

32.6 49.8 48.1

(Continued)

[96] [96] [96] [96]

[96] [96]

[96] [96] [96] [96] [7]

[96] [96] [96]

26.82 27.36

11.24 10.06 11.49 18.11 18.23 17.32 18.16 19.02 22.91 24.86 26.19 25.82 28.27

N,N-Dimethylbenzamide 1-Methyl-2-pyrrolidone

Nitro compounds Nitromethane 1,3-Dimethyl-2-nitrobenzene Nitrobenzene

Sulfonyl compounds N,N-Dimethylbenzenesulfonamide N,N-Dimethylmethanesulfonamide

Sulfites, sulfoxides, selenoxides Ethylene sulfite Dimethyl sulfite Diethyl sulfite Bis(4-chlorophenyl) sulfoxide Diphenyl sulfoxide Bis(4-methylphenyl) sulfoxide Methyl phenyl sulfoxide

Amine oxide Pyridine N-oxide 34.59 33.74

19.61 23.01 24.85

Amides, ureas Dimethylcarbamoyl chloride N-Methylformamide N,N-Dimethylformamide

Phosphoroso compounds Tri(n-butyl)phosphine oxide Hexamethylphosphoric triamide, HMPA

−∆H◦

Lewis base


27.1 24.6

23.5

22.1 24.1 24.0 24.8 25.1 25.8 24.3

24.4b 22.6b

18.2b 14.4b 19.8b

27.4 25.2

27.6 23.4 24.1

−∆Sx◦ a

[96] [96]

[96]

[96] [96] [96] [96] [96] [96] [96]

[96] [96]

[96] [96] [96]

[96] [96]

[96] [96] [96]

Ref.

Trimethyl phosphate Chlorodiphenylphosphine oxide

Dibenzyl sulfoxide Tetramethylene sulfoxide Dimethyl sulfoxide Di-n-butyl sulfoxide Dibenzyl selenoxide Dimethyl selenoxide

Tetramethylene sulfone, sulfolane

4-Nitroanisole 1-Diethylamino-2-nitroethylene

N,N-Dimethylacetamide 1,1,3,3-Tetramethylurea N,N -Dimethyl-N,N -ethyleneurea (DMEU) 1-Methyl-2-pyridone N,N -Dimethyl-N,N -propyleneurea (DMPU)

Lewis base

27.07 26.29

28.06 27.21 27.82 28.81 31.93 30.60

18.65

13.24 18.97

28.53 30.64

29.24 28.61 27.29

∆H◦

23.9 26.7

29.2 24.8 25.5 25.4 23.4 17.8

21.6b

21.9b 20.5b

28.4 30.0

32.2 30.7 24.8

−∆Sx◦ a

[96] [96]

[96] [96] [96] [96] [96] [96]

[96]

[96] [96]

[96] [96]

[96] [96] [96]

Ref.

12.0 24.4 18.2 19.2b 29.5b 18.3b 18.2 20.1b 16.1 19.6 20.3b 16.7 18.9 19.3 21.9 21.5

Thiocarbonyl compounds and isothiocyanate Methyl isothiocyanate 9.04 Thiocamphor 14.72 N,N-Dimethylthiocarbamoyl chloride 14.25 8.75 12.43 10.01 10.36 10.97 (9.98)c 11.02 11.48 (11.51)c 11.61 (12.03)c 13.00 12.95 9.71 10.47 12.8 12.3

Sulfides, thiols, disulfide Methyl disulfide 1,3-Dithiolane Bis(methylthio)methane Ethanethiol 1,4-Dithiane Ethyl vinyl sulfide 2-Propanethiol 1,3-Dithiane Benzyl methyl sulfide Ethylene sulfide Allyl methyl sulfide Propylene sulfide Methyl sulfide

Fluoroalkanes 1,3-Difluoropropane 1-Fluoropentane 2-Fluoroadamantane 1,3-Difluoroadamantane 17.9b 15.8 24.0 16.1b

21.2 26.9 25.3

31.98 30.96 28.42

Trimethylphosphine oxide Triphenylphosphine oxide Triethyl phosphate

[15] [15] [91] [91]

[92] [92] [92] [92] [92] [92] [92] [92] [92] [92] [92] [92] [92]

[87] [87] [96]

[96] [96] [96]

Fluorocyclohexane 1-Fluoroadamantane syn-2,4-Difluoroadamantane

Trimethylene sulfide Methyl n-propyl sulfide n-Butyl methyl sulfide Methyl n-octyl sulfide Ethyl methyl sulfide Pentamethylene sulfide n-Butyl sulfide Cyclohexyl methyl sulfide Ethyl sulfide tert-Butyl methyl sulfide Isopropyl sulfide Tetrahydrothiophene tert-Butyl sulfide

N,N-Dimethylthioformamide N,N-Dimethylthioacetamide

Dichlorophenylphosphine oxide Phosphoryl chloride

12.65 14.43 15.4

12.69 13.45 13.54 13.58 13.36 14.21 15.47 14.62 14.71 14.84 15.96 14.27 18.14

19.09 19.96

20.86 16.53

[15] [15] [91]

[92] [92] [92] [92] [92] [92] [92] [92] [92] [92] [92] [92] [92]

[92] [96]

[96] [96]

(Continued)

21.1 23.1 23.0

20.5 22.6 22.7 22.6 21.8 23.8 28.8 24.7 25.7 25.4 28.4 22.9 33.9

22.9 24.4

26.2 24.9

[45] [45] [45]

Alkyl derivatives of group 15 and 16 elements (in c-C6 H12 ) Triethylstibine ∼12.8 Triethylarsine ∼15.0 Triethylphosphine 21.56 36.8

Methyl selenide n-Butyl telluride Benzyl selenide

Tri(n-octyl)phosphine sulfide

Iodocyclohexane 1-Iodoadamantane

Bromocyclohexane 1-Bromoadamantane

1-Chloropentane 2-Chloropropane Chlorocyclohexane 1-Chloroadamantane

Lewis base

b

Relative to mole fraction. Statistically corrected by −Rln(n), where n is the number of (quasi-)equivalent sites, in order to put these values on a per acceptor atom basis. c To be considered cautiously because of a second (weaker) hydrogen-bonding site.

a

[96] [96]

[15] [96] [15]

Thiophosphoroso and selenophosphoroso compounds Triphenylphosphine selenide 15.90 16.0 Triphenylphosphine sulfide 16.52 16.9

4.96 4.32 6.74

Iodoalkanes 1,3-Diiodopropane 1-Iodopentane 2-Iodo-2-methylpropane

8.0b 10.1 11.2

[15] [88] [88] [15]

Ref.

[15] [15] [15]

5.29 6.54 7.15

Bromoalkanes 1,3-Dibromopropane Bromoethane 1-Bromopentane

10.4b 9.4b 10.2b 9.6b

−∆Sx◦ a

7.0b 7.7 9.6

5.82 5.75 6.25 6.41

−∆H◦

Chloroalkanes 1,3-Dichloropropane trans-1,2-Dichlorocyclohexane 1,4-Dichlorobutane 1,5-Dichloropentane

Lewis base


16.05 12.66 (12.58)c

20.71

6.23 8.52

7.45 9.68

6.96 7.57 8.81 10.17

∆H◦

30.1 19.7

20.5

7.6 13.4

10.2 16.5

10.4 12.1 15.1 18.0

−∆Sx◦ a

[45] [45] [96]

[92]

[15] [15]

[15] [15]

[15] [88] [15] [15]

Ref.


183

Table 4.31 Summary of linear enthalpy Equation 4.26 results. Reference hydrogen-bond donor

Solvent

Methanol

m

c

r

s

n

Lewis bases

CCl4

0.68

−0.2

0.978

1.3

41

Cyclohexanol

CCl4

0.66

−1.7

0.941

1.8

15

n-Butanola

Pure base

0.66

−0.4

0.994

0.2

19

Phenol

CCl4

0.87

+1.9

0.976

1.4

55

Phenola

Pure base

1.01

−0.1

1.000

0.1

26

Trifluoroethanol Hexafluoro-2propanol Perfluoro-tertbutanol Perfluoro-tertbutanolb Isocyanic acid

CCl4 CCl4

0.94 0.93

−0.1 +7.8

0.986 0.979

1.1 1.5

10 11

π , F, Cl, S, O and N bases Br, O and N bases π , O and N bases F, Cl, Br, I, S, O and N bases π , Cl, Br, I, O and N bases O and N bases S, O and N bases

CCl4

1.25

+2.7

0.935

2.3

9

O and N bases

Hexane

1.54

−2.7

0.954

4.6

5

S, O and N bases

CCl4

0.68

2.1

0.955

1.5

6

Isothiocyanic acid

CCl4

1.18

−1.4

0.989

1.0

23

Tri-nbutylammonium

1,2-Cl2 C6 H4

0.99

−2.5

0.947

1.6

7

π , S, O and N bases π , S, O and Nsp bases O and N bases

C6 H5 Cl CCl4

1.41 0.56

−10.2 −0.4

0.975 0.980

1.6 1.2

6 23

CCl4

0.70

−3.5

0.993

0.5

9

c-C6 H12 c-C6 H12 CCl4

0.57

−0.8

0.901 0.953 0.857

1.2

26 16 18

CCl4

0.62

−4.7

0.996

0.4

9

Pyrrole

Chloroform Cyanoacetylene

a b

O and N bases π , S, F, Cl, O and N bases S, O and Nsp bases S, O and N bases S and O bases π , S, O and N bases S, O and Nsp bases

Versus 4-fluorophenol in pure base medium. Versus 4-fluorophenol in cyclohexane.

to 1.000. In view of the well-known difficulty in obtaining precise values of hydrogen-bond enthalpies, most of the deviations from the linear enthalpy relationship appear to be random. Figure 4.7 illustrates typical linear enthalpy relationships. However, in the case of CH donors and pyrrole, a weak NH donor, the correlation is limited by base type. In Figure 4.8, the hydrogen-bond enthalpies of HC CC N are plotted versus the 4-fluorophenol affinity scale in CCl4 . The results may be described by three roughly parallel lines: the upper line for π bases, the middle line for sulfur, oxygen and Nsp bases and the lower line for Nsp2 and Nsp3 bases. This line separation amounts to about 5 kJ mol−1 between the middle and lower lines. Thus, whereas the cyanoacetylene affinities

184

Lewis Basicity and Affinity Scales 70

OH or NH affinity / kJ mol-1

60 50 40 30 20 10 0 0

10

20

30

40

50

-1

4-FC6 H4 OH affinity / kJ mol

Figure 4.7 Family-independent linear enthalpy relationships in the formation of hydrogenbonded complexes of OH and NH reference donors: (CF3 )3 COH, HNCS, CF3 CH2 OH and CH3 OH from left to right.

of dimethyl sulfoxide and N-methylpiperidine are essentially the same (12.4 and 12.2 kJ mol−1 , respectively), the 4-fluorophenol affinity of N-methylpiperidine is 6.2 kJ mol−1 higher than that of dimethyl sulfoxide (34.0 and 27 kJ mol−1 , respectively). A qualitatively similar family behaviour is found when comparing the chloroform and pyrrole affinities to the 4-fluorophenol affinity. It appears that easily protonatable Lewis bases (pyridines and

HCCCN affinity / kJ mol

-1

20

10

0 0

10

20

30

40

50

-1

4-FC6 H4 OH affinity / kJ mol

Figure 4.8 Limited linear enthalpy relationship for formation of hydrogen-bonded complexes of cyanoacetylene and 4-fluorophenol. The upper line represents π bases (benzene, mesitylene, hexamethylbenzene), the middle line sulfur (tetrahydrothiophene), Nsp (acetonitrile) and O (ethers, CO, SO, PO) bases, and the lower line pyridine and tertiary amine bases.


185

amines) gain hydrogen-bond affinity, compared with bases that are difficult to protonate (nitriles, sulfur and oxygen bases), on going from CH and weak NH donors to OH donors. The applicability of the 4-fluorophenol affinity scale to solvents other than the definition solvent CCl4 is difficult to test rigorously because of a lack of data. However, the data for 3-CF3 C6 H4 OH and 3-FC6 H4 OH [120–122] can be used to test the application of Equation 4.26 to the effects of solvents, because these phenols have structures and hydrogen-bond donor properties very close to those of 4-FC6 H4 OH. Table 4.32 lists the values of hydrogenbond affinities of various bases in six solvents: cyclohexane, carbon tetrachloride, benzene, 1,2-dichlorobenzene, dichloromethane and 1,2-dichloroethane. It also gives the values of the solvent polarity parameter ET [103], which increases in the above order from cyclohexane to 1,2-dichloroethane. The linear enthalpy relationship Equation 4.26 is observed (r ≥ 0.996) for all bases in c-C6 H12 , C6 H6 and 1,2-Cl2 C6 H4 . In the most polar solvents, ClCH2 CH2 Cl and CH2 Cl2 , Nsp and oxygen bases still obey Equation 4.26, whereas Nsp2 and Nsp3 nitrogen bases deviate. All deviations from Equation 4.26 are positive, that is, the hydrogen-bond affinities for these bases in polar solvents are greater than expected from Equation 4.26. The gain in affinity amounts to 2–4 kJ mol−1 for pyridine, 1-methylimidazole and triethylamine in ClCH2 CH2 Cl and CH2 Cl2 . Figure 4.9 shows the deviation of pyridine (+3.7 kJ mol−1 ) in a plot of the 3-CF3 C6 H4 OH affinity in CH2 Cl2 versus the 4-FC6 H4 OH affinity in CCl4 . In conclusion, the 4-FC6 H4 OH affinity scale defined in CCl4 ranks the Lewis bases in a reasonably general order. The ranking applies to Nsp and oxygen bases for most hydrogenbond donors in most solvents and to all bases for most OH and NH donors in media of low to medium polarity.

4.6

Comparison of 4-Fluorophenol Affinity and Basicity Scales

The question of whether ∆H or ∆G is the best thermodynamic parameter for measuring the strength of Lewis bases quantitatively has been considered in Chapter 1. An answer was proposed by Pimentel and McClellan in their 1960 review of hydrogen bonding [5]. They presented evidence supporting a monotonic relationship between hydrogen-bond enthalpies and entropies. Consequently, a linear correlation should hold between ∆H and ∆G and the hydrogen-bond affinity and basicity scales should be equivalent. This proposed relation was rationalized on the basis that ‘a higher value of (−∆H) implies stronger bonding, with a more restricted configuration in the polymer, hence greater order, leading to a larger value of (−∆S)’. Arnett’s data for 44 bases with the single hydrogen-bond donor 4-fluorophenol in CCl4 did not confirm the existence of a general ∆H−∆S correlation [11]. However, in a plot of ∆H versus ∆G, some of the data could be resolved into separate lines corresponding to the families of pyridines, amides, sulfoxides and phosphoryl bases. Unfortunately, there were not enough data for other families of bases to enable this observation to be extended. The ∆H−∆G relationship can now be re-examined from the 300 bases listed in Tables 4.3–4.26 and 4.30, for which both ∆H and ∆G have been measured. The ∆G values have been calculated from equilibrium constants obtained by a concentration variation method, while the ∆H values have been determined, in a separate experiment, by a temperature

CCl4 20.8 22.4 23.8 27.8 29.2 29.6 33.7 34.0 39.4

Lewis base

Ethyl acetate Dimethylcyanamide Dibutyl ether Dimethyl sulfoxide Dimethylacetamide Pyridine Hexamethylphosphoric triamide 1-Methylimidazole Triethylamine ET (solvent polarity parameter) 26.4

31.4

32.4

36.4 35.1 33.7 37.5 40.2 30.9 36.0 34.3

25.5

25.1 30.1

C6 H6 16.7

CCl4 21.8

26.5 28.6

C6 H12

4-FC6 H4 OH

38.9 38.0

28.9

23.8 28.0

19.7

o-Cl2 C6 H4

3-FC6 H4 OH

40.7

28.9

20.9 21.3 26.4

CH2 Cl2

36.8 41.3

26.8

18.8 22.6

15.5

(ClCH2 )2

20.9 22.6 26.4 25.9

17.6

CH2 Cl2

3-CF3 C6 H4 OH

Table 4.32 4-Fluorophenol, 3-fluorophenol and 3-trifluoromethylphenol affinities (kJ mol−1 ) of various Lewis bases in the indicated solvents.


187

3-CF3C6H4OH affinity (CH 2Cl2) / kJ mol

-1

30

pyridine

(Me2N)3PO

25

MeCONMe2 Me2SO

20 NMe2CN 15 20

25

30

35

4-FC6H4OH affinity (CCl4) / kJ mol

40 -1

Figure 4.9 Linear enthalpy relationship Equation 4.26 showing the deviation of an Nsp2 base from the line drawn by Nsp and oxygen bases.

variation method. Consequently, the ∆G and ∆H values listed in Tables 4.3–4.26 and 4.30 are based on essentially independent measurements and their direct correlation can be safely attempted. A more sophisticated statistical treatment of the data for haloalkanes [15], sulfur bases [92] and amines [16] confirms the following conclusions: (a) ∆G and ∆H are correlated by a series of crudely parallel lines, each with a slope of about unity. (b) Each line characterizes a given family of bases. In a plot of ∆H versus ∆G, the upper line is drawn by the Nsp3 bases (ammonia, primary, secondary and tertiary amines). Then follow the lines of Nsp2 bases (pyridines, azoles, imines and amidines) and of ethers. The lower lines are those of polar bases: carbonyl compounds (aldehydes, ketones, esters and amides), sulfoxides, phosphoryl compounds and nitriles. These four lines are close together and can be gathered into a single line of polar oxygen and Nsp bases, if a greater scatter is accepted. With the same prerequisite, the lines of fluoroalkanes, nitro and sulfonyl compounds and thiocarbonyl bases may also be merged into the single line of polar bases. The lines of thio- and selenophosphoryl compounds, thioethers, iodo-, bromo- and chloroalkanes are more difficult to define because of a restricted number of bases and/or a narrow range of basicity. The π bases form the lowest line. This family-dependent behaviour is shown in Figure 4.10. Table 4.33 lists the equations and the statistics for each family of bases. They are calculated by a linear regression of −∆H against −∆G assuming that the largest relative error pertains to the enthalpy. A more rigorous statistical treatment of the ∆H−∆G relationship can be found in the papers of Krug et al. [123, 124]. (c) The sterically hindered bases do not obey the ∆H−∆G relationship because of an enhanced entropic term and, consequently, a reduced ∆G value. Steric effects are mainly encountered in the families of secondary and tertiary amines, ethers and thioethers. In these compounds, there are two or three substituents directly attached to the atomic site of hydrogen bonding. If bulky, these substituents restrict the access of the OH group

188

Lewis Basicity and Affinity Scales 45 2 40 1

4-10

35

-∆H / kJ mol-1

30

25

20 3

15

14

10

5

0 -5

0

5

10

15

20

25

-1

-∆G / kJ mol

Figure 4.10 Plot of ∆H versus ∆G for hydrogen bonding of 4-fluorophenol to Lewis bases in CCl4 . The numbering of lines corresponds to numbers listed in Table 4.33. For the sake of clarity, a number of points have been omitted.

of 4-fluorophenol to the lone pair(s) of the atomic site and lead to a reduced freedom within the complex and, consequently, to a larger (negative) value of the hydrogen-bond entropy. Values of hydrogen-bond entropies of hindered bases are compared with those of similar unhindered bases in Table 4.34. In summary, the correlation of hydrogen-bond affinity and basicity scales exhibits a scatter caused by family-dependent relationships and steric effects. For unhindered bases, the scatter can be analysed into a series of nearly parallel lines for (i) amines, (ii) pyridines, (iii) polar bases and (iv) π bases.

4.7 4.7.1

Spectroscopic Scales Infrared Shift of Methanol

It is well known that the X–H stretching infrared frequency of hydrogen-bond donors is significantly lowered upon hydrogen bonding [5, 6] to Lewis bases, and that the shift increases with the hydrogen-bond strength. Therefore, many workers have suggested that


189

Table 4.33 Family-dependent relationships between enthalpies and Gibbs energies for the hydrogen bonding of 4-fluorophenol to Lewis bases, in CCl4 at 25 ◦ C. No

Family of bases

n

r

s

Slope

Intercept

1 2 3 4 5 6

Nitrogen sp3 bases Nitrogen sp2 bases Cyclic ethers Carbonyl bases Sulfinyl bases Phosphoryl bases, pyridine N-oxide Nitriles Fluoroalkanes Nitro, sulfonyl bases Thiocarbonyl bases Thiols, thioethers Chloro-, bromo-, iodoalkanes Thio-, selenophosphoryl bases π bases ‘Polar’ bases

69 37 9 32 11 10

0.891 0.988 0.970 0.989 0.997 0.995

1.2 0.7 0.6 0.8 0.3 0.6

0.99 ± 0.06 1.03 ± 0.03 1.13 ± 0.11 1.08 ± 0.03 1.10 ± 0.03 0.96 ± 0.04

22.0 ± 0.7 18.0 ± 0.3 15.1 ± 0.7 13.6 ± 0.3 12.2 ± 0.3 13.7 ± 0.6

18 7 8 4 20 18 3 7 90

0.974 0.929 0.986 0.975 0.972 0.845 0.999 0.967 0.986

0.6 0.8 0.7 0.8 0.4 0.9 ∼0 0.4 1.0

0.98 ± 0.06 0.99 ± 0.18 1.19 ± 0.08 1.19 ± 0.19 1.64 ± 0.09 2.00 ± 0.32 1.39 ± 0.04 1.23 ± 0.15 1.07 ± 0.02

13.1 ± 0.4 12.4 ± 0.3 11.3 ± 0.4 11.8 ± 0.9 11.9 ± 0.1 11.0 ± 0.7 8.5 ± 0.3 10.4 ± 0.2 12.8 ± 0.2

7 8 9 10 11 12 13 14 4–10

the ν(XH) shifts can be used as a spectroscopic measurement of the hydrogen-bond acceptor strength of Lewis bases, and a number of infrared scales have been proposed [18–21, 125]. For reasons given below, we advocate the scale based on the ν(OH) frequency of methanol measured in the ‘inert’ solvent CCl4 at 298 K. This scale, derived according to Equation 4.27, is presented in Table 4.35. ∆ν(OH) = ν(OH, ‘free’ methanol in CCl4 ) −ν(OH, hydrogen-bonded methanol in CCl4 ) ∆ν(OH)/cm−1 = 3644 − ν(MeOH · · · B in CCl4 )

(4.27)

The methanol scale presents the following advantages: r The OH band of the complex does not overlap bands of the base. r The OH band maximum is easy to measure because the band shape is quasi-GaussoLorentzian. Table 4.34 Comparison of hydrogen-bond entropies (J K−1 mol−1 ) for hindered and similar but unhindered Lewis bases. Hindered base 2,2,5,5-Tetramethyltetrahydrofuran tert-Butyl sulfide N-Benzylidene-tert-butylamine 2,2,6,6-Tetramethylpiperidine 1,2,2,6,6-Pentamethylpiperidine 2,6-Dimethylpyridine 2-tert-Butylpyridine

−∆Sx◦ 48.1 33.9 60.1 70.3 70.8 52.3 60.1

Unhindered base Tetrahydrofuran Methyl sulfide Benzophenone imine Piperidine 1-Methylpiperidine Pyridine 2-Methylpyridine

−∆Sx◦

Difference

35.5 21.5 47.0 56.4 54.4 44.2 44.6

12.6 12.4 13.1 13.9 16.4 8.1 15.5

33 42 74 50 68 33 36 38 46 74

Thiophene Thiophene

Alkenes 1,4-Cyclohexadiene 2,3-Dimethyl-2-butene Cyclohexene 1-Methyl-1-cyclohexene 1-Heptene 3,3-Dimethyl-1-butene

Alkynes 1,5-Hexadiyne 1-Hexyne 3-Hexyne 244 284

24

Hindered pyridines 2,6-Di-tert-butylpyridine

Ammonia and primary amines 2,2,2-Trifluoroethylamine Propargylamine

28 33 35 37 36 37 43 43

∆ν(OH)

Arenes Benzene Diphenylmethane Methylbenzene, toluene Ethylbenzene Isopropylbenzene, cumene tert-Butybenzene 1,4-Dimethylbenzene, p-xylene 1,3-Dimethylbenzene, m-xylene

Lewis base

[21] [21]

[21] [21] [21]

[21] [21] [21] [21] [21] [21]

[97]

[77]

[21] [21] [21] [21] [21] [21] [21] [21]

Ref.

Isopropylamine tert-Butylamine

2-Hexyne 3,3-Dimethyl-1-butyne

2-Methyl-1-propene cis-2-Pentene 2-Methyl-1-butene 2,3-Dimethyl-1-butene cis-2-Octene 2,3,3-Trimethyl-1-butene

2,6-Di-tert-butyl-4-methylpyridine

1,3,5-Trimethylbenzene, mesitylene 1,3,5-Triisopropylbenzene 1,2,4,5-Tetramethylbenzene, durene 1,2,3,4-Tetramethylbenzene, prehnitene Pentamethylbenzene Hexamethylbenzene Cyclopropylbenzene

Lewis base

351 355

72 50

47 47 51 54 55 63

32

50 51 57 56 63 71 35

∆ν(OH)

[21] [21]

[21] [21]

[21] [21] [21] [21] [21] [21]

[77]

[21] [21] [21] [21] [21] [21] [21]

Ref.

Table 4.35 Shift (cm−1 ) of the infrared stretching wavenumber of the OH group of methanol upon hydrogen bonding to Lewis bases in CCl4 .

310 324 327 351 345 346 175 270 274 334 344 337 348 336 361 359 356 360 370 412 381 396 374 376 401 385 394 383

Cyclopropylamine Benzylamine Allylamine Ethylamine Methylamine n-Butylamine

Secondary amines 1,1,1,3,3,3-Hexamethyldisilazane (Methylamino)acetonitrile Thiazolidine Bis(2-chloroethyl)amine Dibenzylamine 3-(Methylamino)propionitrile 2-(3-Trifluoromethylphenyl)pyrrolidine N-Methylpropargylamine 2-(3-Fluorophenyl)pyrrolidine Thiomorpholine Diallylamine Morpholine N-Methylbenzylamine 2,2,6,6-Tetramethylpiperidine 2-Phenylpyrrolidine Diisopropylamine N-Methylallylamine 1,2,3,4-Tetrahydroisoquinoline Di-n-butylamine Piperazine N-Methylphenethylamine 1,2,3,6-Tetrahydropyridine [85] [85] [21] [21] [21] [85] [16] [16] [16] [21] [21] [21] [21] [21] [16] [21] [16] [21] [16] [21] [16] [99]

[21] [21] [21] [21] [21] [21] N-Methylisopropylamine N-Methyl-tert-butylamine N-Methylcyclohexylamine N-Methylbutylamine Azepane, hexamethylenimine N-Methylethylamine Diethylamine Dimethylamine 2-Methylaziridine N,N -Dimethylethylenediamine Bis(2-methoxyethyl)amine Piperidine Pyrrolidine Azetidine N-Ethylbenzylamine N-Isopropylbenzylamine N-tert-Butylbenzylamine N-Ethylcyclohexylamine Dicyclohexylamine N-Isopropylcyclohexylamine N-tert-Butylcyclohexylamine N-Isopropyl-tert-butylamine

Cyclohexylamine Adamantane-1-amine Hydrazine 2-Chloroethylamine 2-Aminobutane

395 406 402 395 403 394 389 389 327 407 396 405 411 406 374 375 388 402 405 399 396 390

357 357 294 298 348

(Continued)

[16] [16] [21] [16] [21] [16] [21] [21] [85] [16] [85] [21] [21] [21] [21] [21] [21] [21] [21] [21] [21] [21]

[21] [21] [21] [21] [21]

[99] [21] [16] [86] [99] [21] [99] [21]

366 417 372 419 377 368 355 416 368 335 375 386 386 427 384 397 430 383 356 397 367

N,N-Dimethyl-3-fluorobenzylamine Hexamethylenetetramine Triallylamine N-Methyl-2-phenylpyrrolidine N-(2-Chloroethyl)pyrrolidine Tri-n-propylamine N-Ethyl-N-methylbenzylamine N,N-Dimethyl-3-chloropropylamine Tri-n-butylamine N-Methylmorpholine 1,3,5-Trimethylhexahydrotriazine N,N-Dimethylbenzylamine N,N-Dimethylpropargylamine

[99] [86] [86] [16] [86] [86] [99] [86] [86] [21] [86] [21] [86]

[86] [86] [16]

Ref.

392 282 368

∆ν(OH)

Tertiary amines N,N-Diisopropylisobutylamine N,N-Dimethylaminoacetonitrile N-Methyl-2(3-trifluoromethyl)pyrrolidine N-Ethyl-N-methyl-3(trifluoromethyl)benzylamine N,N-Diisopropylethylamine N-Methyl-2(3-fluorophenyl)pyrrolidine N,N-Dicyclohexylethylamine N-Ethyl-N-methyl-3fluorobenzylamine 3-(Dimethylamino)propionitrile N,N-Dimethyl3(trifluoromethyl)benzylamine 1,2,2,6,6-Pentamethylpiperidine

Lewis base


402 426 415

N,N,N ,N -Tetramethylethylenediamine N-Butylpyrrolidine N,N,N ,N -Tetramethylhexane-1,6diamine N,N-Dimethylisopropylamine N-Methylpiperidine Trimethylamine N,N-Dimethylcyclohexylamine N,N-Dimethylethylamine N-Methylpyrrolidine Diazabicyclooctane (DABCO) Tropane 4-Phenylquinuclidine Quinuclidine 3-Quinuclidinone 2-Chloro-N,N-diethylethylamine

425 421 408 434 418 422 417 446 433 450 371 373

429 399

399 394

402

400 401 383

∆ν(OH)

Triethylamine N-Methyl-1,2,3,6-tetrahydropyridine

N,N-Dimethylallylamine 3-Chloroquinuclidine

4-Chloro-N-methylpiperidine 3-Chloromethyl-N-methylpiperidine N-Methyl-1,2,3,4tetrahydroisoquinoline N,N-Dimethylpiperazine

Lewis base

[86] [86] [21] [21] [86] [21] [86] [86] [86] [21] [21] [21]

[86]

[86] [86]

[86] [99]

[86] [86]

[86]

[86] [86] [86]

Ref.

Six-membered N-heteroaromatics Pentafluoropyridine 2,6-Difluoropyridine 1,3,5-Triazine 2-Cyanopyridine 5-Bromopyrimidine 3-Cyanopyridine 3,5-Dichloropyridine Pyrazine 4-Cyanopyridine 2-Fluoropyridine 2-Bromopyridine 2-Chloropyridine Pyrimidine 2,2 -Bipyridine 7,8-Benzoquinoline Phenazine 2,5-Dimethylpyrazine 3-Bromopyridine 3-Fluoropyridine 3-Chloropyridine 3-Iodopyridine 4-Acetylpyridine 3-Benzoylpyridine 2-tert-Butylpyridine 2-Phenylpyridine Methyl nicotinate 4-Chloropyridine 2-N,N-Dimethylaminopyridine 2-Vinylpyridine Pyridazine 2-Isopropylpyridine 40 87 147 158 171 203 200 205 211 167 192 192 213 267 270 248 248 241 240 239 244 255 248 292 270 254 248 294 287 218 307

[80] [80] [80] [80] [80] [80] [80] [80] [80] [80] [80] [80] [80] [80] [80] [80] [99] [80] [80] [80] [80] [80] [80] [80] [80] [80] [80] [80] [80] [80] [80]

Phenanthridine Quinoline Isoquinoline 2-Ethylpyridine Acridine 4-Vinylpyridine 4-Phenylpyridine Phthalazine 3-Methylpyridine, 3-picoline 3-Ethylpyridine 2-Methylpyridine, 2-picoline 4-Methylpyridine, 4-picoline 4-Ethylpyridine 4-tert-Butylpyridine 2-Aminopyridine 4-Methoxypyridine 2,6-Dimethylpyridine, 2,6-lutidine 3,5-Dimethylpyridine, 3,5-lutidine 2,4-Dimethylpyridine, 2,4-lutidine 3,4-Dimethylpyridine, 3,4-lutidine 2,4,6-Trimethylpyridine, 2,4,6-collidine 3-N,N-Dimethylaminopyridine 4-Aminopyridine N-Methyl-N-pyridine-4-ylhydrazine 4-(4-Methylpiperidino)pyridine 4-Piperidinopyridine 4-N,N-Dimethylaminopyridine 4-N,N-Diethylaminopyridine 4-Pyrrolidinopyridine 1,7-Phenanthroline Quinazoline

292 296 291 316 310 293 293 239 300 305 313 304 306 303 320 312 331 314 316 314 349 330 345 354 359 359 366 370 373 289 235

(Continued)

[80] [80] [80] [80] [80] [80] [80] [80] [80] [80] [80] [80] [80] [80] [80] [80] [80] [80] [21] [80] [80] [21] [80] [80] [80] [80] [80] [80] [80] [80] [80]

[64] [60] [60] [60]

Me Me Me Me

230 281 284 268

Me Me Me Me

R2

H H H H

R3

CH2 CH2 C N tert-Am 4-Me2 NC6 H4 C6 H5 CH2

R4

1-Methylpyrazole 1-Acetylimidazole 5-Bromo-1-methylimidazole 1-Vinylimidazole 1-Methylimidazole 1,5-Dicyclohexylimidazole 1-Benzylimidazole

3,4,7,8-Tetramethyl-1,10phenanthroline 4,7-Dipyrrolidino-1,10-phenanthroline

Formamidines (R3 = H) Me Me H 3,5-(NO2 )2 C6 H3 Me Me H 4-NO2 C6 H4 Me Me H 4-N≡CC6 H4 Me Me H 2-BrC6 H4

[21] [21] [21] [21] [96] [21] [96]

[94] [94] [94]

2-Methoxypyridine 3-Acetylpyridine 3-Phenylpyridine 2-Chloroquinoline 6-Methylquinoline 6-Methoxyquinoline 2-Methylquinoline

Lewis base

R1

69 75 123 189 216 217 217

Five-membered N-heteroaromatics Pyrrole 1-Methylpyrrole Isoxazole Oxazole Benzothiazole Thiazole 2-Methylbenzoxazole

NR4

235 243 254

Substituted 1,10-phenanthrolines 1,10-Phenanthroline 4,7-Diphenyl-1,10-phenanthroline 4,7-Dimethyl-1,10-phenanthroline

[80] [21] [21] [21] [21] [21] [21]

Ref.

Amidines R1 R2 NC(R3 ) R1 R2 R3 R4

286 228 259 298 343 227 248

∆ν(OH)

Pyridine 4-Trifluoromethylpyridine Methyl isonicotinate 4-Benzylpyridine 4-Morpholinopyridine 3-Trifluoromethylpyridine N,N-Diethylnicotinamide

Lewis base


364 392 334 371

[64] [64] [64] [60]

[96] [96] [96] [21] [21] [21] [21]

[94]

302 240 257 272 283 313 354 315

[94]

[80] [21] [21] [21] [21] [21] [21]

Ref.

261

232 252 287 202 303 304 319

∆ν(OH)

287 295 302 309 304 309 316 355 350 325 327 360 365

472

239 271 313 290 287

Miscellaneous amidines 1,8-Diazabicyclo[5.4.0]undec-7-ene (DBU)

Imines and oximes N-Benzylideneaniline N-Benzylidenebenzylamine N-Benzylidene-tert-butylamine N-Benzylidenemethylamine 2-(3-Trifluoromethylphenyl)-1pyrroline

390

4-CF3 C6 H4 4-MeCOC6 H4 CH2 CF3 2-MeC6 H4 4-BrC6 H4 4-FC6 H4 C6 H5 3-CF3 C6 H4 CH2 3,5-Cl2 C6 H3 CH2 4-MeC6 H4 4-MeOC6 H4 3-ClC6 H4 CH2 4-ClC6 H4 CH2 390

H H H H H H H H H H H H H

Guanidines Tetramethylguanidine

Me Me Me Me Me Me Me Me Me Me Me Me Me

Benzamidine (R3 = phenyl) Me Me Ph Me

Me Me Me Me Me Me Me Me Me Me Me Me Me

[94] [94] [94] [94] [99]

[96]

[60]

[60]

[64] [60] [64] [60] [60] [64] [60] [64] [60] [60] [64] [60] [60]

2-(3-Fluorophenyl)-1-pyrroline Isobutylideneisopropylamine Benzophenone imine, diphenylketimine 2-Phenyl-1-pyrroline 2-Methyl-1-pyrroline

1,5-Diazabicyclo[4.3.0]non-5-ene (DBN)

7-Methyl-1,5,7-triazabicyclo[4.4.0] dec-5-ene (MTBD)

Me Me H 4-MeC6 H4 CH2 Me Me H c-Pr Me Me H 3-MeOC6 H4 CH2 Me Me H t-Bu Me Me H NMe2 Me Me H CH2 CH CH2 Me Me H i-Bu Me Me H 1-Adam Me Me H n-Pr Me Me H c-Hex Me Me H i-Pr Me Me H Me Me Me H CH2 CH2 OMe

297 319 280 314 344

464

457

372 368 376 387 381 376 390 399 395 395 385 386 396

(Continued)

[99] [94] [94] [99] [94]

[96]

[96]

[60] [64] [64] [60] [64] [64] [60] [64] [60] [64] [60] [64] [64]

Nitriles Trichloroacetonitrile Dibromoacetonitrile Cyanogen bromide Chloroacetonitrile α,α,α-Trifluoro-m-tolunitrile α,α,α-Trifluoro-p-tolunitrile 2-Fluorobenzonitrile 4-Chlorobenzonitrile 2-Chlorobenzonitrile α-Bromo-o-tolunitrile Acrylonitrile 2-Bromobenzonitrile 4-Fluorobenzonitrile Methyl thiocyanate Phenyl cyanate Benzonitrile Benzyl cyanide, phenylacetonitrile o-Tolunitrile 2,6-Dimethylbenzonitrile Acetonitrile-d 3 Allyl cyanide Butyronitrile Hexanenitrile Acetonitrile Propionitrile Trimethylsilyl cyanide 4-Methoxybenzonitrile Isobutyronitrile Trimethylacetonitrile

Lewis base


23 39 44 48 60 59 62 68 66 64 67 65 70 69 70 72 73 76 77 75 78 79 83 75 79 80 83 81 83

∆ν(OH) [66] [66] [66] [66] [66] [66] [66] [66] [66] [66] [66] [66] [66] [66] [66] [66] [66] [66] [96] [66] [96] [66] [66] [66] [66] [66] [66] [66] [66]

Ref. N1 ,N1 -Dimethyl-N2 -cyanoformamidine N1 ,N1 -Dimethyl-N2 -cyanoacetamidine N,N,N-Tri-n-butylammoniocyanamidate Cyanoacetylene 2-Chloroacrylonitrile Bromoacetonitrile Iodoacetonitrile 3-Chloropropionitrile (Methylthio)acetonitrile 4-Chlorobutanenitrile 5-Chloropentanenitrile Pentanenitrile, valeronitrile Cyclohexyl cyanide 4-Nitrobenzonitrile α-Bromo-p-tolunitrile 4-Bromobenzonitrile 4-Phenylbenzonitrile 3-Nitrobenzonitrile 3-Cyanobenzonitrile 3,5-Dichlorobenzonitrile 3-Bromobenzonitrile 3-Chlorobenzonitrile α-Bromo-m-tolunitrile 3,5-Dimethoxybenzonitrile 3-Methoxybenzonitrile 3-Methylbenzonitrile 2-Cyanobenzonitrile α,α,α-Trifluoro-o-tolunitrile 2-Nitrobenzonitrile

Lewis base 149 161 244 43 47 59 65 69 71 75 77 80 83 46 70 70 75 51 54 56 66 64 67 72 73 76 48 59 62

∆ν(OH)

[66] [66] [72] [66] [66] [66] [66] [66] [66] [66] [66] [66] [66] [66] [66] [66] [66] [66] [66] [66] [66] [66] [66] [66] [66] [66] [66] [66] [66]

Ref.

87 86 82 100 107 117 122 124 129 96 76 116 126 129 120 121 79 16 78 60 70 82 83 117 94 86 89 117 107

1-Adamantanecarbonitrile Cyclopropyl cyanide 2-Methoxybenzonitrile 4-Dimethylaminobenzonitrile 4-Morpholinecarbonitrile Dimethylcyanamide 1-Piperidinecarbonitrile Diethylcyanamide trans-3-Dimethylaminoacrylonitrile

Water, alcohols 2-Chloroethanol Water Methanol Phenethyl alcohol Ethylene glycol n-Propanol n-Butanol

Ethers Hexamethyldisiloxane Furan Bis(2-chloroethyl) ether 1,3,5-Trioxane Ethyl vinyl ether Tetramethyl orthocarbonate 3,4-Dihydro-2H-pyran 2-Chloroethyl ethyl ether 1,3-Dioxolane 2,3-Dihydrofuran Trimethyl orthoformate Dimethoxymethane 1,3-Dioxane [81] [81] [81] [81] [81] [81] [81] [81] [81] [81] [81] [81] [81]

[21] [21] [61] [61] [21] [61] [61]

[66] [66] [66] [66] [21] [66] [66] [66] [66]

15-Crown-5 12-Crown-4 Cyclohexene oxide tert-Butyl methyl ether 18-Crown-6 3-Methyltetrahydropyran Tetrahydropyran Tetrahydrofuran 2-Methyltetrahydrofuran Oxetane, trimethylene oxide Cineole, eucalyptol 2,2,5,5-Tetramethylhydrofuran 2,3-Diadamant-2-yl oxirane

Ethanol n-Octanol 2-Propanol tert-Butanol Cyclohexanol 1-Adamantanol

4-Fluorophenylacetonitrile 4-Chlorophenylacetonitrile 4-Bromophenylacetonitrile 4-Methylphenylacetonitrile 4-Methoxyphenylacetonitrile 4-Aminophenylacetonitrile 3-(Trifluoromethyl)phenylacetonitrile 3-Chlorophenylacetonitrile 3-Methylphenylacetonitrile

141 143 118 166 134 143 157 158 160 155 185 185 180

120 122 129 135 133 141

71 74 68 75 77 84 67 69 77

(Continued)

[81] [81] [81] [81] [81] [96] [81] [81] [81] [81] [81] [81] [81]

[61] [61] [61] [61] [61] [61]

[66] [66] [66] [66] [66] [66] [66] [66] [66]

65 93 123 77 83 79 84 84 89

Aldehydes Benzaldehyde 4-Methoxybenzaldehyde 4-Dimethylaminobenzaldehyde

Aliphatic ketones 2,2,4,4-Tetramethylpentan-3-one 2,4-Dimethylpentan-3-one 1-Adamantyl tert-butyl ketone Pentan-3-one 3,3-Dimethylbutan-2-one Di(1-adamantyl) ketone

61 59 77

101 68

Peroxides tert-Butyl peroxide 3,4-Diadamant-2-yl dioxetan

Aromatic ethers 4-Chloroanisole Diphenyl ether Anisole, methoxybenzene

129 127 126 168 154 104 150 125 165 133 164

∆ν(OH)

Dibenzyl ether Diallyl ether 1,4-Dioxane Di-tert-butyl ether Di-n-butyl ether Propylene oxide Diethyl ether 1,2-Dimethoxyethane tert-Butyl ethyl ether 1,2-Diethyoxyethane Diisopropyl ether

Lewis base


[79] [21] [79] [21] [21] [79]

[79] [79] [79]

[21] [96] [21]

[81] [81]

[81] [81] [81] [81] [81] [81] [81] [81] [81] [81] [81]

Ref.

Propan-2-one, acetone 3-Methylbutan-2-one Butan-2-one 1-Adamantyl methyl ketone Methyl cyclopropyl ketone Dicyclopropyl ketone

4-Methylbenzaldehyde 2-Naphthaldehyde

1,4-Dimethoxybenzene 2,3-Dihydro-1,4-benzodioxin 3-Chloroanisole

Ascaridole

Epichlorhydrin n-Butyl vinyl ether Dimethyl ether Morpholine Di-n-propyl ether 2,5-Dimethyltetrahydrofuran 1,2-Epoxybutane 7-Oxabicyclo[2.2.1]heptane Bis(2-Methoxyethyl) ether, diglyme

Lewis base

115 86 91 87 107 110

83 63

83 62 61

80

72 71 134 149 153 171 109 151 123

∆ν(OH)

[79] [21] [79] [79] [21] [79]

[21] [21]

[96] [21] [21]

[81]

[81] [21] [21] [21] [21] [21] [21] [21] [21]

Ref.

46 87 58 80 105 89

Miscellaneous conjugated ketones Benzil 1,4-Benzoquinone But-3-yn-2-one Propiophenone 9-Fluorenone 2-Acetylnaphthalene 119

87 97 95

Ring-substituted benzophenones Benzophenone 4-Methoxybenzophenone 4,4 -Bis(methoxy)benzophenone

3-Chloro-5,5dimethylcyclohexenone

65 68 67 77 64 71 84 80 88 92

89 83 85 121 103 126

Ring-substituted acetophenones 4-Cyanoacetophenone 3-Trifluoromethylacetophenone 4-Trifluoromethylacetophenone 3-Chloroacetophenone 2-Chloroacetophenone 1,4-Diacetylbenzene 4-Chloroacetophenone 4-Fluoroacetophenone 3-Methylacetophenone Acetophenone

Cycloalkanones Cycloundecanone Cyclopentadecanone Cyclododecanone Cyclopentanone Camphor Cyclohexanone

[79]

[96] [21] [79] [96] [79] [79]

[21] [79] [79]

[79] [79] [79] [79] [79] [79] [79] [79] [79] [79]

[79] [79] [79] [79] [79] [79]

2,2-Dimethylpropiophenone 2-Methylpropiophenone 3-Bromo-5,5-dimethylcyclohexenone 5,5-Dimethylcyclohexenone 3-Phenyl-5,5-dimethylcyclohexenone 3-Methylthio-5,5dimethylcyclohexenone 3-Methoxy-5,5-dimethylcyclohexenone

4-Dimethylaminobenzophenone 4,4 -Bis(diethylamino)benzophenone

3-Methoxyacetophenone 4-(Methylthio)acetophenone 4-Isopropylacetophenone 4-Methylacetophenone 4-Ethylacetophenone 4-tert-butylacetophenone 4-(1-Adamantyl)acetophenone 4-Methoxyacetophenone 4-Morpholinoacetophenone 4-Dimethylaminoacetophenone

3-Nitrocamphor 3-Bromocamphor 3-Chlorocamphor 3-Ethynylcamphor 3-Phenylcamphor 3-Ethylcamphor

179

75 82 116 135 144 169

123 160

88 87 98 102 86 92 106 111 117 134

48 75 79 85 103 106

(Continued)

[21]

[21] [21] [21] [21] [21] [21]

[79] [79]

[79] [79] [79] [79] [79] [79] [79] [79] [79] [79]

[21] [21] [21] [21] [21] [21]

Esters, lactones, carbonates Ethyl trifluoroacetate Methyl trichloroacetate Ethyl trichloroacetate Ethyl 4-nitrobenzoate Ethyl 4-cyanobenzoate Ethyl propiolate Diethyl oxalate Methyl formate Ethyl formate Ethyl chloroacetate Vinylene carbonate Ethyl fluoroacetate Diethyl terephthalate Ethyl 4-bromobenzoate Dimethyl carbonate Ethyl methyl carbonate 13 25 29 57 60 50 43 52 58 52 42 54 69 72 68 71

185 214 220 248

[71] [71] [71] [71] [71] [71] [71] [96] [71] [71] [71] [71] [71] [71] [71] [71]

[79] [79] [79] [79]

[79]

167

γ -Pyrone Diphenylcyclopropenone 2,6-Dimethyl-γ -pyrone 3-Dimethylamino-5,5dimethylcyclohexenone

[79]

154

3-Methyl-5,5dimethylcyclohexenone, isophorone Flavone

[79] [79]

Ref.

119 124

∆ν(OH)

Xanthone trans-4-Phenylbut-3-en-2-one

Lewis base


Ethyl phenylacetate Ethyl 1-adamantanecarboxylate Ethyl acetate Ethyl propionate Ethyl isobutyrate Ethyl cyclopropanecarboxylate Ethyl 4-methoxybenzoate (E)-Ethyl cinnamate Isopropyl acetate Propyl acetate Propylene carbonate Coumarin γ -Butyrolactone γ -Valerolactone Ethyl 4-dimethylaminobenzoate δ-Valerolactone

3-Diethylamino-5,5dimethylcyclohexenone Cyclohexyl phenyl ketone α-Tetralone 1-Indanone Chromone

3-Ethoxy-5,5-dimethylcyclohexenone 3-Diphenylamino-5,5dimethylcyclohexenone 3-Morpholino-5,5dimethylcyclohexenone

Lewis base

71 85 83 81 83 90 88 83 85 78 73 88 93 92 106 107

85 109 129 151

255

226

191 207

∆ν(OH)

[71] [71] [71] [71] [71] [71] [71] [71] [71] [99] [71] [71] [71] [71] [71] [71]

[21] [21] [21] [21]

[21]

[21]

[21] [21]

Ref.

[65] [65]

[65] [96] [65] [65] [65] [65] [65] [65] [65] [65] [96]

112 133 123 137 130 130 145 152 161 160 161 150 153 158

N-Methylformanilide

Ethyl dimethylcarbamate N-Formylmorpholine N,N-Diphenylacetamide N-Methylformamide N-Methylbenzamide N,N-Diisopropyl-2,2dimethylpropionamide N,N-Dicyclohexyl-2,2dimethylpropionamide N,N,2,2-Tetramethylpropionamide N,N-Dimethylformamide N,N-Diethylformamide N-Acetylmorpholine

[65]

[65] [65]

[71] [71] [71] [71] [71] [71] [71] [99] [71] [71]

70 81

65 49 76 76 81 54 77 78 83 85

Amides, lactams, carbamates and ureas Dimethylcarbamoyl chloride N,N-Dimethyl-2,2,2trifluoroacetamide Phenyl dimethylcarbamate N,N-Dimethyl-2-chloroacetamide

Triacetin β-Propiolactone Diethyl carbonate Methyl benzoate Ethyl benzoate β-Butyrolactone Methyl acetate Methyl 3,3-dimethylacrylate Ethyl 2,2-dimethylpropanoate Ethyl 4-methylbenzoate

N,N-Diphenyl-2-methylpropionamide N,N-Diethylbenzamide N,N-Dimethyl-4-nitrobenzamide N,N-Dimethyl-4trifluoromethylbenzamide

Antipyrine N,N -Dimethyl-N,N -propyleneurea (DMPU) 2,4-bis(dimethylamino)-1-oxo-3-azabutadiene Diphenylcarbamoyl chloride Diethylcarbamoyl chloride N,N-Dimethyl-2,2,2-trichloroacetamide N,N-Diphenylformamide N,N-Diphenylbenzamide N,N-Diphenyl-2,2dimethylpropionamide N,N-Diphenyl-4-methoxybenzamide

N-Methylcaprolactam 1-Methyl-2-piperidone

ε-Caprolactone Ethyl 3-dimethylaminoacrylate 4-Nitrophenyl acetate Phenyl formate tert-Butyl benzoate Methyl cyclohexanecarboxylate Methyl cyclopropanecarboxylate Ethyl 4-fluorobenzoate Ethyl 3-methylbenzoate (S)-Ethyl methylcarbonate

131 167 139 144

121

48 74 81 100 116 118

230

216 210

183 194

98 148 32 34 75 81 85 76 83 59

(Continued)

[21] [65] [65] [65]

[65]

[65] [65] [65] [65] [65] [65]

[100]

[100] [65]

[65] [65]

[71] [71] [71] [71] [71] [71] [71] [71] [71] [71]

[69] [94] [69] [69] [69] [69]

239 263 278 285 289 335

[65] [65]

184 192

Carboxamidates N,N,N-Trimethylammonio 4-trifluoromethylbenzamidate N,N,N-Trimethylammonio benzamidate N,N,N-Trimethylammonio 4-methoxybenzamidate N,N,N-Trimethylammonio trimethylacetamidate N,N,N-Trimethylammonio adamantyl-1-carboxamidate N,N,N-Trimethylammonio acetamidate

[65] [65] [65] [65] [65] [65] [65] [65]

171 167 166 185 175 179 177 183

N,N-Dimethylisobutyramide N-Methylacetamide N,N-Dimethylpropionamide 1-Methyl-2-pyrrolidone 1,1,3,3-Tetraethylurea N,N-Dimethylacetamide 1,1,3,3-Tetramethylurea N,N -Dimethyl-N,N -ethyleneurea (DMEU) N,N-Diethylacetamide 1-Methyl-2-pyridone

[65] [65] [65] [65] [65]

Ref.

152 150 170 159 174

∆ν(OH)

N-Formylpiperidine N-Methylacetanilide N,N-Dicyclohexylpropionamide N,N-Dimethylbenzamide N,N-Dicyclohexylisobutyramide

Lewis base


N,N,N-Trimethylammonio isobutanamidate N,N,N-Trimethylammonio propanamidate N,N,N-Trimethylammonio 4-cyanobenzamidate N,N,N-Trimethylammonio 4-fluorobenzamidate N,N,N-Trimethylammonio 4-methylbenzamidate

Methyl dimethylcarbamate Ethyl diethylcarbamate

N,N-Dimethyl-4-bromobenzamide N,N-Dimethyl-4-fluorobenzamide N,N-Dimethyl-4-methylbenzamide N,N-Dimethyl-4-methoxybenzamide N,N-Dimethyl-4dimethylaminobenzamide N,N-Diethyl-4-methoxybenzamide 1,1,3,3-Tetraphenylurea N,N-Diphenyl-N ,N -diethylurea N,N-Diphenyl-N ,N -dimethylurea N,N -Diethyl-N,N -diphenylurea Phenyl diphenylcarbamate Methyl diphenylcarbamate Ethyl diphenylcarbamate

Lewis base

274

252

223

311

291

131 143

174 126 152 153 159 82 100 103

152 158 168 171 185

∆ν(OH)

[69]

[94]

[94]

[69]

[69]

[65] [65]

[65] [65] [65] [65] [65] [65] [65] [65]

[65] [65] [65] [65] [65]

Ref.

178 211 264 278 287 291

68 74

N,N-Dimethylmethanesulfonamide Dimethyl sulfone

Amine oxides 4-Nitropyridine N-oxide 3,5-Dichloropyridine N-oxide 4-Chloropyridine N-oxide Pyridine N-oxide 4-Phenylpyridine N-oxide 3-Methylpyridine N-oxide

35 52 61 63

Sulfonyl compounds Diethyl sulfate Ethyl methanesulfonate N,N-Dimethylbenzenesulfonamide Diphenyl sulfone

62 65 75 140 158 169 175

74

Nitroso compounds Nitrosobenzene

Sulfinyl, seleninyl compounds Ethylene sulfite Dimethyl sulfite Diethyl sulfite Bis(4-chlorophenyl) sulfoxide Diphenyl sulfoxide Bis(4-methylphenyl) sulfoxide Methyl phenyl sulfoxide

22 26 29 28 30 31 35

Nitro compounds Nitromethane Nitrobenzene Nitroethane N-Nitrocamphorimine 2-Nitropropane 2-Methyl-2-nitropropane 1,2-Dimethyl-4-nitrobenzene

[21] [94] [21] [21] [21] [21]

[96] [96] [96] [96] [96] [96] [21]

[78] [78]

[78] [78] [78] [78]

[96]

[70] [70] [70] [70] [70] [70] [70]

4-Methylpyridine N-oxide 4-Methoxypyridine N-oxide Triethylamine N-oxide N-Methylpiperidine N-oxide Tri(n-octyl)amine N-oxide Trimethylamine N-oxide

Dibenzyl sulfoxide Tetramethylene sulfoxide Dimethyl sulfoxide Di-n-butyl sulfoxide Dibenzyl selenoxide Dimethyl selenoxide

Tetramethylene sulfone, sulfolane N,N,N ,N -Tetraethylsulfamide Di-n-butyl sulfone N,N-Dimethyl-N phenylsulfonylformamidine N-Trimethylammoniooctanesulfonamidate Dimethyl sulfate

1-Nitrosopyrrolidine

4-Nitroanisole Dimethylnitramine N,N-Diethyl-4-nitroaniline 1-Dimethylamino-2-nitroethylene 1-Nitro-2-piperidinoethylene 1-Diethylamino-2-nitroethylene

306 338 473 472 488 474

189 220 205 223 315 331

163 28

77 77 86 92

111

38 55 60 93 98 104

(Continued)

[21] [21] [94] [94] [94] [94]

[96] [96] [21] [96] [96] [96]

[78] [21]

[78] [78] [78] [78]

[96]

[70] [74] [70] [70] [70] [74]

[21] [94] [94] [94]

287 263 239 206 365

Arsine oxide Triphenylarsine oxide

Disulfides, thiols, thioethers Diphenyl sulfide Methyl disulfide Ethyl disulfide 1,3-Dithiolane

[97] [97] [97] [97]

[73] [73] [73] [73]

157 162 167 172

89 63 75 100

[87] [73] [73] [73]

53 105 130 111

[94]

[94] [94] [21]

288 282 274

Thiocarbonyl compounds, isothiocyanates Methyl isothiocyanate Ethylene trithiocarbonate Thiocamphor N,N-Dimethylthiocarbamoyl chloride N,N-Dimethylthiobenzamide N,N-Dimethylthioformamide 4-N,N-Trimethylthiobenzamide N,N-Dimethyl-4methoxythiobenzamide

[94]

Ref.

352

∆ν(OH)

Phosphoroso compounds Trimethylammonio-P,Pdiphenylphosphinamidate Tripiperidinophosphine oxide Triethylphosphine oxide Hexamethylphosphoric triamide (HMPA) Tri(n-octyl)phosphine oxide Trimethylphosphine oxide Triphenylphosphine oxide Diethyl methylphosphonate

Lewis base


n-Butyl methyl sulfide Methyl n-octyl sulfide Ethyl methyl sulfide Pentamethylene sulfide

Ethyl isothiocyanate Butyl isothiocyanate Tetramethylthiuram disulfide Tetraethylthiuram disulfide

N,N-Dimethylthioacetamide N,N-Dimethyl-N ,N -diethylthiourea N,N,N ,N -Tetramethylthiourea N,N -Dimethyl-N,N -ethylenethiourea

Triphenyl phosphate Dichlorophenylphosphine oxide Phosphoryl chloride Diethyl chlorophosphate

Tri(n-butyl) phosphate Trimethyl phosphate Chlorodiphenylphosphine oxide

Triethyl phosphate

Lewis base

140 139 140 142

56 56 129 130

184 190 194 180

126 91 49 127

189 173 159

189

∆ν(OH)

[97] [97] [97] [97]

[21] [21] [21] [21]

[73] [73] [73] [73]

[94] [94] [94] [21]

[94] [94] [94]

[94]

Ref.

10 18 19 24 25 29 34 24 24 32 33 38

Chloroalkanes 1,3-Dichloropropane 1-Chlorobutane 1-Chloropentane 2-Chloropropane 2-Chloro-2-methylpropane

Bromoalkanes 1,3-Dibromopropane Bromocyclopropane Bromoethane 1-Bromobutane 2-Bromopropane

100 108 90 113 106 94 107 123 119 123 125 135 137 136

Fluoroalkanes 1-Fluorooctane Fluorocyclohexane

Thioanisole Bis(methylthio)methane Ethanethiol 1,4-Dithiane 2-Chloroethyl methyl sulfide 2-Propanethiol 1,3-Dithiane Benzyl methyl sulfide Ethylene sulfide Allyl methyl sulfide Propylene sulfide Methyl sulfide Trimethylene sulfide Methyl n-propyl sulfide

[21] [21] [21] [21] [21]

[21] [21] [21] [21] [21]

[21] [21]

[97] [97] [97] [97] [97] [97] [97] [97] [97] [97] [97] [97] [97] [97]

Bromocyclohexane 2-Bromo-2-methylpropane 1-Bromoadamantane Bromomethane

Chlorocyclohexane Chloroform Dichloromethane Chloroethane

2-Fluoro-2-methylpropane

n-Butyl sulfide Cyclohexyl methyl sulfide Ethyl sulfide tert-Butyl methyl sulfide Isopropyl sulfide Tetrahydrothiophene tert-Butyl sulfide Hydrogen sulfide Methanethiol Benzyl mercaptan 2-Methyl-2-propanethiol Cyclohexanethiol 4-Bromothioanisole

41 43 49 25

33 3 7 24

21

148 145 146 152 159 146 176 44 84 88 102 103 89

(Continued)

[21] [21] [21] [21]

[21] [21] [21] [21]

[21]

[21] [97] [97] [97] [97] [21] [97] [21] [21] [21] [21] [21] [21]

Ion pairs Tetrabutylammonium iodide Tetrabutylammonium methylsulfonate Tetrabutylammonium chloride

Alkyl derivatives of group 16 elements Methyl selenide [96] [96] [96]

257 266 361

132

[21]

[21] [21]

81 96

Organo derivatives of group 15 elements Trimethyl phosphite Triethyl phosphite

[21] [21] [21] [21]

Ref.

[21] [21] [21] [21] [21]

25 28 31 37

∆ν(OH)

Thiophosphoroso and selenophosphoroso compounds Trimethyl thiophosphate 111 Triethyl thiophosphate 120 Triphenylphosphine selenide 158 Triphenylphosphine sulfide 145 Triethylphosphine sulfide 195

Iodoalkanes 1,2-Diiodoethane 1,3-Diiodopropane Iodomethane Iodoethane

Lewis base


Tetrabutylammonium acetate

Tetrabutylammonium azide Tetrabutylammonium cyanate

Ethyl selenide

Triphenylarsine

Tri(n-octyl)phosphine sulfide Diethyl chlorothiophosphate Trimethylphosphine sulfide Tri(n-butyl)phosphine sulfide Triethylselenophosphate

2-Iodopropane 2-Iodo-2-methylpropane Iodocyclohexane 1-Iodobutane

Lewis base

450

360 446

145

104

199 77 191 200 141

43 49 48 40

∆ν(OH)

[96]

[96] [96]

[21]

[21]

[21] [21] [21] [21] [21]

[21] [21] [21] [21]

Ref.


207

Methanol basicity /kJ mol-1

14

9

4

-1

-6 0

100

200

300

400

500

∆ν (OH) / cm-1

Figure 4.11

Plot of methanol basicity versus ∆ν(OH).

r Careful measurements have been made in the same conditions in the same laboratory. In particular, care has been taken to (i) check the wavenumber calibration, (ii) regulate the temperature at 298 K, (iii) purify and desiccate the solvent, (iv) use a long-pathlength cell (1–4 cm) in order to work with very dilute solutions of methanol and base and (v) deconvolute overlapping free and bonded OH bands, in the case of weak bases, and overlapping multiple bonded OH bands, in the case of polyfunctional bases. r About 800 ∆ν(OH) values are available spanning from ∼3 cm−1 for the very weak chloro base CHCl3 to 488 cm−1 for the strong oxygen base Oct3 NO. Almost all kinds of organic bases are represented in Table 4.35 and, in each family, large structural variations have been studied. For example, in the pyridine family, ∆ν(OH) varies from 40 cm−1 (pentafluoropyridine) to 373 cm−1 (4-pyrrolidinopyridine). It was claimed [126] that infrared scales could provide the basicity descriptor for multiparameter linear solvation energy relationships, which relate the variation of a solventdependent property to the polarizability, polarity, acidity and basicity of the solvent. This implies that infrared scales are related to thermodynamic scales. The Gibbs energy and enthalpy of the reaction 4.28 have been measured for 41 (42 for ∆G) diverse Lewis bases B [96] in CCl4 at 298 K. CH3 OH + B CH3 OH · · · B

(4.28)

As shown in Figure 4.11, the methanol spectroscopic scale is very weakly correlated to the methanol basicity (n = 42, r = 0.737). Only 54% of the variance of ∆G is explained by ∆ν. Some resolution of the scatter diagram is, however, achieved by noting good correlations of ∆G with ∆ν(OH) when separating the families of pyridines, ethers and polar bases (nitrile, sulfoxide, sulfone, carbonyls and phosphoryls) which obey Equations 4.29, 4.30 and 4.31,

208


respectively. ∆G = 0.040∆ν(OH) − 8.75

(4.29)

n = 5, r = 0.997, s = 0.2 kJ mol

−1

∆G = 0.048∆ν(OH) − 6.68

(4.30)

n = 4, r = 0.993, s = 0.2 kJ mol ∆G = 0.046∆ν(OH) − 3.32

−1

(4.31) −1

n = 12, r = 0.988, s = 0.6 kJ mol

The last relation explains why ∆ν(OH) may behave as a successful descriptor of basicity if the sample of bases is restricted to such nitrogen sp and oxygen bases. Such was the case with the correlation analysis of the solvent effect on the rate of reaction between benzoic acid and diazodiphenylmethane in 44 solvents [127, 128]. A ∆ν(OH) spectroscopic scale described satisfactorily the stabilization of benzoic acid by hydrogen bonding to basic solvents (PhCOOH · · · B), that is, rate deceleration with increasing solvent basicity. The correlation between ∆ν(OH) and the methanol affinity scale is more successful since 84% of the variance of the hydrogen-bond enthalpy can be explained by ∆ν(OH). However, Equation 4.32: −∆H (kJ mol−1 ) = 0.0438 ∆ν (OH) (cm−1 ) +7.73

(4.32)

−1

n = 41, r = 0.915, s = 2.52 kJ mol

does not seem to be a satisfactory method for providing reliable methanol affinities. Obviously, it cannot be used for hydrogen bonds whose negative enthalpies are less than 7.7 kJ mol−1 , for instance the case of methanol hydrogen bonded to π , halogen and sulfur bases. For stronger nitrogen and oxygen bases, which have methanol affinities falling between 10 and 26 kJ mol−1 , a standard deviation amounting to 2.5 kJ mol−1 reduces the utility of Equation 4.32. However, Figure 4.12 suggests that several lines are generated for several series of bases (e.g. π bases, ethers, pyridines and polar bases). Such family-dependent ∆H−∆ν(OH) and ∆G−∆ν(OH) relationships have proved useful for detecting steric effects [80] and three-centre hydrogen bonds [91], and for calculating the individual hydrogen-bond affinities (basicities) of molecules with more than one hydrogen-bonding site [16]. Examples are given below. The departure of 2-substituted pyridines from the behaviour of 3- and 4-substituted pyridines in a ∆G−∆ν(OH) plot illustrates the sensitivity of hydrogen-bond basicity to steric effects [80]. For 24 3- and 4-substituted pyridines, the 4-FC6 H4 OH basicity is well correlated with the ∆ν(OH) shifts of methanol (Equation 4.33): − ∆G (kJ mol−1 ) = 0.067 ∆ν (OH) (cm−1 ) − 8.43

(4.33)

−1

n = 24, r = 0.998, s = 0.2 kJ mol

Pyridines substituted at the ortho position by bulky substituents fall below the regression line. These deviations ∆∆G have been attributed mainly to steric effects. Indeed, they are well correlated by the steric upsilon (υ) parameter of the substituent (Equation 4.34): ∆∆G (kJ mol−1 ) = 2.04 υ + 0.1

(4.34) −1

n = 7(H, Me, Et, i-Pr, t-Bu, vinyl, Ph), r = 0.998, s = 0.2 kJ mol


209

Methanol affinity / kJ mol-1

30

20

10

0 0

100

200

300

400

500

∆ν (OH) /cm-1

Figure 4.12

Plot of methanol affinity versus ∆ν(OH).

In a series of complexes of 4-fluorophenol with fluoroalkanes, the ∆ν(OH) value of syn-2,4-difluoroadamantane appears abnormally weak (42 cm−1 ) compared with its 4FC6 H4 OH affinity (15.4 kJ mol−1 ) and does not obey the ∆H − ∆ν(OH) relationship Equation 4.35: −∆H (kJ mol−1 ) = 0.125 ∆ν(OH) (cm−1 ) + 5.5

(4.35)

n = 6, r = 0.984, s = 0.3 kJ mol−1 This has been attributed [91] to the formation of a three-centre hydrogen-bonded complex (Figure 4.13). Indeed, the topological analysis of the electron density shows two separated bond critical points between the two fluorine atoms and the hydrogen atom. In morpholine, there are two sites for accepting a hydrogen bond. Therefore, this molecule forms two 1 : 1 complexes simultaneously with OH donors: the N complex and the O complex [16] (Figure 4.14).

Figure 4.13 Three-centre hydrogen-bonded complex of 4-fluorophenol with syn-2,4difluoroadamantane.

210


(a)

Figure 4.14

(b)

N complex (b) and O complex (a) of morpholine with an OH donor.

There is no simple experimental method for measuring the hydrogen-bond affinity of each site. However, two ∆ν(OH) values are obtained from the IR spectrum of the complex: ∆ν(OH · · · O) = 149 cm−1 and ∆ν(OH · · · N) 360 cm−1 . From the ∆H−∆ν(OH) relationship Equation 4.36: −∆H (kJ mol−1 ) = 0.076 ∆ν(OH) (cm−1 ) + 5

(4.36)

n = 29, r = 0.944, s = 0.65 kJ mol−1 found for the family of secondary amines, we obtain ∆H = −32.0 kJ mol−1 for the N complex. This calculated value agrees well with the experimental values of piperazine and piperidine since the hydrogen-bond affinity decreases when the electronegativity of the γ ring atom increases from carbon to nitrogen to oxygen (Scheme 4.3). 4.7.2

Solvatochromic Shifts of 4-Nitrophenol and 4-Nitroaniline

With hydrogen-bond formation, the S0 → S1 transition of phenols, corresponding to the A1g → B2u transition of benzene, undergoes a bathochromic shift. For example, the longest wavelength π → π ∗ transition of 4-fluorophenol at 281.1 nm in CCl4 (absorption coefficient ∼3000 l mol−1 cm−1 ) is red shifted to 286.5 nm on hydrogen bonding with Oct3 PO. This electronic shift of 671 cm−1 (8 kJ mol−1 ) is equal to W 0 − W 1 + w1 , where W 0 and W 1 are the energies of hydrogen-bond formation in the ground and excited states, respectively, and w1 is the excitation implied by the Franck–Condon principle [129]. Since a red shift is observed, then W 0 < W 1 (i.e. the hydrogen bond is strengthened in the excited state) and the red shift is less by w1 than the difference W 1 − W 0 .

O

N

32.0 (calc.)

H < H N

N H

33.3 (exp.)

< H2C

N H

34.8 (exp.)

Scheme 4.3 Comparison of the 4-fluorophenol affinities (kJ mol−1 ) of, from left to right, morpholine, piperazine and piperidine.

Thermodynamic and Spectroscopic Scales of Hydrogen-Bond Basicity and Affinity O

211

H

gas NO2

Heptane CCl4

-1 ν∼ /cm

o-Cl2C6H4

DMSO

MeO -1 ν∼ /cm

NO2

∆ ν∼ (vdW)

Figure 4.15 The solvatochromic comparison method. In a plot of the corresponding ν˜ values of a hydrogen-bond donor probe, 4-NO2 C6 H4 OH, vs a very similar but non-hydrogen-bond donor probe, 4-NO2 C6 H4 OMe, non- or very weak hydrogen-bond acceptor solvents draw a so-called comparison line with a very high correlation coefficient from the gas phase to polyhalogenated benzenes, because the van der Waals (vdW) effects of these solvents are similar for the two probes. Hydrogen-bond acceptor solvents (e.g. dimethyl sulfoxide) are displaced below the comparison line because of an enhanced solvatochromic shift caused by hydrogen bonding. The contribution of the hydrogen bond, ∆˜ν (HB), to the total solvatochromic shift, ν˜ (gas) − ν˜ (DMSO), of 4-NO2 C6 H4 OH is calculated as shown in the figure.

Greater electronic frequency shifts are observed when intramolecular charge transfer occurs upon excitation in push–pull compounds such as 4-nitroaniline and 4-nitrophenol. For example, in cyclohexane, free 4-nitrophenol absorbs at 285 nm and 4-nitrophenol hydrogen bonded to triethylamine at 307 nm. The shift amounts to 2515 cm−1 (30 kJ mol−1 ). Kamlet and Taft [25] proposed using the two probes 4-nitrophenol and 4-nitroaniline to construct a solvatochromic scale of solvent hydrogen-bond acceptor strength and devised a method for measuring the hydrogen-bond shifts of their longest wavelength π → π ∗ electronic transition (at 283.6 and 318.8 nm in heptane, respectively), which they refer to as the solvatochromic comparison method. In this method, the solutes 4-nitrophenol and 4-nitroaniline act as reference hydrogen-bond donors and the base is used as the solvent, so the solutes can be assumed to be entirely hydrogen bonded (however, see below). The solvatochromic shifts of 4-nitroaniline and 4-nitrophenol are compared with those of N,N-diethyl-4-nitroaniline and 4-nitroanisole, respectively, in order to isolate the

212


hydrogen-bond contribution from the total shift. The method is outlined in Figure 4.15. The hydrogen-bond shifts of 4-nitrophenol and 4-nitroaniline were claimed to be proportional to one another, proportional to limiting 19 F NMR shifts of 4-fluorophenol in CCl4 , and linear with the logarithm of the hydrogen-bonding constants of 4-fluorophenol and phenol in CCl4 . This gave support to defining a β scale of hydrogen-bond acceptor strength from the five sets of properties: hydrogen-bond shifts of 4-nitrophenol and 4-nitroaniline, log K of 4-fluorophenol and phenol and ∆19 F of 4-fluorophenol. The scale was fixed by setting β = 0 for cyclohexane and β = 1 for hexamethylphosphoric triamide. Thirty-nine β values were originally calculated in this way [25]. New aniline probes [130], and correlations with the β scale of the literature data concerning numerous basicity-dependent properties [26] have allowed the number of available β values to be extended to 162 [27]. Thus, the β scale has been defined by mixing (i) numerous OH, NH2 and NH reference hydrogen-bond donors (and even other Lewis acids [26]), (ii) numerous physicochemical properties (solvatochromic shifts, infrared shifts, complexation constants and complexation enthalpies) and (iii) two different standard states: the pure base (hydrogen-bond donor in the pure base as solvent) and the dilute base (hydrogen-bond donor in a diluted solution of the base in an ‘inert’ solvent). Thus, it is a ‘wide-range’ but ‘low-resolution’ scale. It was used as the basicity descriptor of linear solvation energy relationships, which were claimed to describe hundreds of physicochemical properties and reactivity parameters of many diverse types [131–133]. The concept of linear solvation energy relationships was criticized by statisticians [134]. Moreover, there is a discussion in the literature [23] concerning the use of averaged and statistically optimized parameters instead of experimentally derived parameters, which are based on a single and well-understood basicity-dependent reference process. For these reasons, the definition of solvatochromic scales is revisited in the following. Laurence et al. [23, 24, 135] measured two solvatochromic scales for about 200 bases, each based on one reference hydrogen-bond donor, 4-nitrophenol or 4-nitroaniline, one standard state, the pure base, and one physical property, the enhanced bathochromic shift of the longest wavelength π → π ∗ transition (with intramolecular charge-transfer character) upon hydrogen bonding. The solvatochromic comparison method was improved and ‘inert’ solvents were carefully chosen [135] to fix the comparison lines of Equations 4.37 and 4.38: ν˜ (OH) = 1.0381 ν(OMe) ˜ − 384

(4.37)

n = 32 (‘inert’ solvents and gas phase), r = 0.9959, s = 77 cm−1

ν(NH ˜ ˜ 2 ) = 0.9752 ν(NMe 2 ) + 3278

(4.38)

n = 32 (‘inert’ solvents and gas phase), r = 0.9957, s = 96 cm−1 where ν˜ (OH), ν˜ (OMe), ν(NH ˜ ˜ 2 ) and ν(NMe 2 ) stand for the maximum wavenumber of the longest wavelength π → π ∗ transition of 4-nitrophenol, 4-nitroanisole, 4-nitroaniline and N,N-dimethyl-4-nitroaniline, respectively, in the various solvents at 25 ◦ C. The ∆ν˜ shift attributable to hydrogen bonding with the basic solvent is extracted from the total


213

4-Fluorophenol affinity/kJmol-1

50 40 30 20 10 0 0

500

1000

1500

2000

2500

∆ν (OH-OMe) /cm-1

Figure 4.16 Plot of 4-fluorophenol affinity (in pure base) versus hydrogen-bond shifts of the longest wavelength π → π ∗ transition of 4-nitrophenol.

solvatochromic shift by means of Equations 4.39 and 4.40: ∆ν˜ (OH − OMe) = [1.0381 ν(OMe) ˜ − 384] − ν˜ (OH)

˜ ˜ (NH2 ) ∆˜ν (NH2 − NMe2 ) = [0.9752 ν(NMe 2 ) + 3728] − ν

(4.39)

(4.40)

They are given in Table 4.36. Contrary to previous claims [25], ∆ν˜ (OH – OMe) and ∆˜ν (NH2 – NMe2 ) are not strongly correlated in general. For 127 solvents common to the two sets, the correlation coefficient is only r = 0.786. Only π , halogen, Nsp and O bases give a good correlation (n = 91, r = 0.978), which is not obeyed by ethers, pyridines and amines. As shown in Figure 4.16, the ∆˜ν (OH – OMe) scale is well correlated (n = 41, r = 0.977) with the 4-FC6 H4 OH affinity scale measured by Arnett [11] in pure base media. The temperature independence of ∆ν˜ (OH – OMe) from 0 to 50–100 ◦ C confirms that ∆˜ν (OH – OMe) is an enthalpy-like property [23]. The correlation of the ∆˜ν (NH2 – NMe2 ) scale with the pyrrole affinity scale measured by Catalan [17] in pure base media is less clear-cut. As shown in Figure 4.17, tetrahydrofuran and dioxane (but not anisole) and triethylamine (but not aniline and pyridine) deviate significantly above the regression line drawn with 20 π , Nsp, Nsp2 and O bases (r = 0.987). Thermosolvatochromism studies show [23] that ∆ν˜ (NH2 – NMe2 ) generally decreases when the temperature increases from 0 to 50–100 ◦ C, indicating instability of the complexes at high temperatures and, for a number of bases, incomplete complexation at 25 ◦ C. Infrared studies of the symmetric and asymmetric NH2 stretching vibrations of 4-nitroaniline show (ref. [136] and M. Lucon, University of Nantes, personal communication) that a class of solvents (e.g. pyridine, hexamethylphosphoric triamide) gives mainly 1 : 2 complexes, whereas another class gives either a mixture of 1 : 1 and 1 : 2 complexes (e.g. tetrahydrofuran) or mainly 1 : 1 complexes (e.g. triethylamine) (Figure 4.18). The

100 290 170 470 100

Thiophene Thiophene

Alkenes 1,4-Cyclohexadiene Dicyclopentadiene 1,5-Cyclooctadiene

Arylacetylene Ethynylbenzene, phenylacetylene

Secondary amines Diallylamine Morpholine N-Methylbenzylamine Diisopropylamine 2270 2470 2450 2230

2430 3230

500

Hindered pyridine 2,6-Di-tert-butylpyridine

Primary amines Propargylamine tert-Butylamine n-Butylamine

220 180 230 270

1860 1740 1730 1880

1710 2600 2560

90

300 280 330

110

80 200 340 480

N-Methylcyclohexylamine Diethylamine Piperidine N-(tert-Butyl)benzylamine

Allylamine Cyclohexylamine

trans,trans,cis-1,5,9-Cyclododecatriene Cycloheptatriene

1,3,5-Trimethylbenzene, mesitylene 1,2,3,4-Tetramethylbenzene, prehnitene 1-Methylnaphthalene

∆˜ν ∆˜ν (OH − OMe) (NH2 − NMe2 ) Lewis base

Arenes Benzene Diphenylmethane Methylbenzene, toluene 1,4-Dimethylbenzene, p-xylene

Lewis base

2800 2670 3610

500

390 480

2590 2500 2530 1220

2180 2490

290 620

570 650 280

∆˜ν ∆˜ν (OH − OMe) (NH2 − NMe2 )

Table 4.36 Bathochromic shift (cm−1 ) of the longest wavelength π → π ∗ transition of 4-nitrophenol and 4-nitroaniline upon hydrogen bonding to Lewis bases in pure base media, obtained by the solvatochromic comparison method [24].

370 480 1030 620 1170 1260

Nitriles Trichloroacetonitrile Chloroacetonitrile Benzonitrile Benzyl cyanide, phenylacetonitrile Butyronitrile Hexanenitrile 480 400 780 750 930 960

2880

370 890

Guanidine Tetramethylguanidine

200

330 850 1190 1270 1490 1800

Six-membered N-heteroaromatics Pentafluoropyridine 2,6-Difluoropyridine 2-Fluoropyridine 2-Bromopyridine 3-Bromopyridine Pyridine

Five-membered N-heteroaromatics Pyrrole N-Methylpyrrole

1340 1510 1840

2440 2030 2050

Tertiary amines Tri-n-butylamine N,N-Dimethylbenzylamine N,N-Dimethylpiperazine 90 870 1230 1240 1520 1830

750

Primary arylamine Aniline

Acetonitrile Propionitrile Trimethylacetonitrile Dimethylcyanamide Valeronitrile, pentanenitrile

N-Methylimidazole

4-Methylpyridine, 4-picoline 2,6-Dimethylpyridine, 2,6-lutidine Quinoline 3,4-Dimethylpyridine, 3,4-lutidine 2,4,6-Trimethylpyridine, 2,4,6-collidine Pyrimidine

Triethylamine N-Methylpiperidine N,N-Dimethylcyclohexylamine

800 850 950 1120 940

1960 2000

1960 1910

2320 2220 2330

(Continued)

1030 1110 1120 1370 1200

2080

1590 2050 1950 1420

1940

1770 1980 2090

450

180 650 1420 1450 810 960 940

Ethers Furan Bis(2-chloroethyl) ether Cineole, eucalyptol 2,2,5,5-Tetramethyltetrahydrofuran 1,3-Dioxolane Dibenzyl ether 1,4-Dioxane

Aromatic ethers Anisole, methoxybenzene

450 600 320 1110 1100 1170 910 1500 1440 1330 1390 1330 1740 1810 1640

360

790 1950 1860 980 840 920

140 1030 630 960 1120 1070 440 1710 1690 1170 1580 1550 2050 2280 1990

Phenetole, ethoxybenzene

Di-tert-butyl ether Di-n-butyl ether Diethyl ether Diisopropyl ether Tetrahydropyran Tetrahydrofuran 2,5-Dimethyltetrahydrofuran (cis + trans)

n-Octanol 2-Propanol tert-Butanol Cyclohexanol 2-Methoxyethanol 2-Ethoxyethanol n-Pentanol 2-Butoxyethanol 3-Hydroxypropionitrile Propargyl alcohol 2-Phenoxyethanol n-Decanol n-Hexanol tert-Amyl alcohol

∆˜ν ∆˜ν (OH − OMe) (NH2 − NMe2 ) Lewis base

Water, alcohols, phenol 2,2,2-Trifluoroethanol 2,2,2-Trichloroethanol 3-Methylphenol 2-Chloroethanol 2-Bromoethanol 2-Fluoroethanol Water Allyl alcohol Methanol Benzyl alcohol Phenethyl alcohol Ethylene glycol n-Propanol n-Butanol Ethanol

Lewis base


790

1550 1340 1220 1370 1180 1210 1340

1930 1870 2120 1430 1390 1460 1810 1550 1010 1100 1210 1960 1800 1960

370

1110 1360 1450 1560 1510 1340 1840

2220 2360 2520 2200 1740 1770 2250 1780 910 830 1130 2160 2290 2350

∆˜ν ∆˜ν (OH − OMe) (NH2 − NMe2 )

Amides, lactams and ureas N-Methylformanilide Formamide N-Methylformamide

Esters, lactones, carbonates Methyl trichloroacetate Ethyl trichloroacetate Ethyl chloroacetate Dimethyl carbonate Triacetin Diethyl carbonate Ethyl benzoate Methyl acetate Ethyl acetate Propylene carbonate γ -Butyrolactone δ-Valerolactone n-Propyl formate 1290 1300 1550

660 610 860 820 840 930 970 970 1010 820 940 1100

1100

Cycloalkanones Cyclopentanone

Acetophenone Acetophenone

570 1090 1150

Aliphatic ketones 2,4-Dimethylpentan-3-one 1,1,1-Trichloropropan-2-one 2,2,4,4-Tetramethylpentan-3-one Pentan-3-one

1610 1300

1030 1500 1120 1180 1070 1230 1460 1080

460 490 1070 810

1240

1420

1560 440 1480 1540

N-Methylacetamide N,N-Dimethylpropionamide 1-Methyl-2-pyrrolidone

Methyl formate Methyl trifluoroacetate Methyl propionate Methyl butyrate Methyl valerate Methyl caproate Methyl caprylate Methyl caprate Methyl oleate Methyl linoleate 4,5-Dichloro-1,3-dioxolan-2-one Triolein

Cyclohexanone

Propan-2-one, acetone Butan-2-one Hexachloroacetone

1780 1610 1560

830 630 980 1040 1040 1040 1070 1280 1300 1160 410 1030

1080

1010 1130 320

(Continued)

2200 2330 2170

430

1470

1400 1440 330

700 930

Sulfinyl compounds Dimethyl sulfite Diethyl sulfite

Disulfides, thioethers Methyl sulfide Trimethylene sulfide Pentamethylene sulfide Ethyl disulfide Methyl disulfide 690 680 780

2030 1490 1580

690

Sulfonyl compounds Tetramethylene sulfone, sulfolane

Phosphoroso compounds Hexamethylphosphoric triamide (HMPA) Triethyl phosphate Tri(n-butyl) phosphate

740 420

1450 1580 1500 1850 1810

∆˜ν (OH − OMe)

Nitro compounds Nitromethane Nitroethane

N,N-Dimethylformamide N,N-Diethylformamide N-Formylpiperidine N-Methylpropionamide N-Ethylacetamide

Lewis base


590 690 620 350 410

2760 2090 2120

1170 1260

1020

640 810

2050 2140 2150 2460 2480

∆˜ν (NH2 − NMe2 )

Thioanisole n-Butyl sulfide Ethyl sulfide Isopropyl sulfide Tetrahydrothiophene

Trimethyl phosphate Phosphoryl chloride Diethyl chlorophosphate

Tetramethylene sulfoxide Dimethyl sulfoxide

Dimethyl sulfate

2-Nitropropane 2-Methyl-2-nitropropane

1,1,3,3-Tetraethylurea N,N-Dimethylacetamide 1,1,3,3-Tetramethylurea N,N-Diethylacetamide 2-Pyrrolidone

Lewis base

890 770 860 760

1340 670 1160

1560 1500

450

550 650

1620 1610 1590 1700 910

∆˜ν (OH − OMe)

370 930 850 920 760

1590

1820

2080 2020

710

780 680

2310 2210 2320 2340

∆˜ν (NH2 − NMe2 )

220 280 330 260 210 190 280 220 730 220 590

Chloroalkanes 1,4-Dichlorobutane 1-Chlorobutane 1-Chloropropane 1,10-Dichlorodecane

Bromoalkanes Bromoethane 1-Bromopropane

Iodoalkanes Iodomethane Iodoethane

Miscellaneous bases Ethyl propiolate Propargyl chloride Formic acid 900 70

130 180

160 190

190 360 240 220

Acetic acid Propionic acid Valeric acid

1-Iodobutane 1-Iodopropane

1-Bromobutane

1000 870 1070

430

240

890 710 910

190 140

270

220


Pyrrole affinity/kJ mol-1

25

NEt3

20 THF

15

Dioxane

Pyridine

10 Anisole Aniline

5 0 0

500

1000

1500

2000

2500

3000

∆ν (NH2-NMe2)/cm

-1

Figure 4.17 Plot of pyrrole affinity (in pure base) versus hydrogen-bond shifts of the longest wavelength π → π ∗ transition of 4-nitroaniline.

latter class of solvents will therefore produce lower hydrogen-bond shifts than the first one and corresponding data points will be upshifted from the ∆H−∆˜ν (NH2 – NMe2 ) regression line. The ∆˜ν (OH – OMe) scale is devoid of this stoichiometric problem and is the solvatochromic scale of choice for the correlation of solvent basicity-dependent properties. It can be scaled in a range from 0 (solvents obeying the comparison Equation 4.37) to 1 by dividing by the 2030 cm−1 shift of HMPA: β(OH) = ∆ν˜ (OH − OMe)/2030

(4.41)

It must be recalled that β(OH) is a solvent scale and that solvent and solute scales of hydrogen-bond acceptor strength are not entirely equivalent. For example, for 23 common bases, β(OH) explains only 91% of the variance of a solute scale constructed [12] with the same property, the hydrogen-bond shift of the same π → π ∗ transition of 4-nitrophenol, but in the solvent CCl3 CH3 instead of the pure base.

B

B

H O2N

N

O2N

H

Figure 4.18

H

B

N H

1 : 2 (left) and 1 : 1 (right) complexes of 4-nitroaniline with basic solvents.


4.8

221

Conclusion

Three extended, diverse and reliable scales are currently available for hydrogen-bonding studies: r the solute 4-FC6 H4 OH basicity scale, a reasonably general scale for Gibbs energydependent properties; r the solute 4-FC6 H4 OH affinity scale, correlated with basicity by family-dependent relationships; r the solvent spectroscopic scale ∆˜ν (OH – OMe) of 4-nitrophenol, correlated with the pure-base hydrogen-bond affinity. An infrared scale based on the OH shift of methanol can be applied to the study of steric effects, three-centre hydrogen bonds and bases with multiple sites.

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5 Thermodynamic and Spectroscopic Scales of Halogen-Bond Basicity and Affinity It has been known for a long time that the heavy halogen atoms, iodine, bromine and chlorine, can function as electron-acceptor sites, that is, act as Lewis acids, and engage in molecular interactions with electron-donor sites of Lewis bases. Indeed, in 1863, the formation of a substance with the composition I2 NH3 was described [1] and in 1896 it was reported [2] that the addition of trimethylamine to dibromine produced crystals with the composition (CH3 )3 NBr2 . In 1949, Benesi and Hildebrand deduced the existence of the 1 : 1 diiodine–benzene and diiodine–mesitylene complexes from the electronic absorption spectrum of diiodine dissolved in these aromatic solvents [3]. The experimental observations were given a theoretical basis in 1952 by Mulliken [4]. In 1954, Hassel and Hvoself determined the structure of the solid 1 : 1 complex formed by dibromine and 1,4-dioxane [5]. The first gas-phase microwave study of a complex involving a halogen atom as an acceptor was the determination by Legon et al., in 1975, of the structure of the 1 : 1 complex formed by iodotrifluoromethane and trimethylamine [6]. Dumas et al. in 1983 introduced [7] the term ‘halogen bonding’ to designate the interaction between the positive end of heavier halogens in carbon–halogen bonds and electron-rich portions of other molecules. This term was used again by Blackstock et al. in 1987 [8] and Murray et al. in 1994 [9] for the complexes of CCl4 and CBr4 with Lewis bases. Particularly important reviews on the concept of halogen bonding were published by Legon in 1999 (gaseous complexes of dihalogens with small Lewis bases) [10] and Metrangolo and Resnati in 2001 and 2005 (solid complexes of organic halides) [11, 12]. The halogen bond is usually represented as Y X· · ·B, where X is the halogen, the notation · · · indicates a halogen bond, the Lewis base B is the halogen-bond acceptor (electron donor), the Lewis acid X Y is the halogen-bond donor (electron acceptor) and

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Y is any (preferably electron-withdrawing) organic (e.g. CF3 or C N) or inorganic (e.g. halogen or N3 ) group. Many structural [10, 13, 14], energetic [14, 15] and spectroscopic [10, 15] properties of the halogen bond run parallel to those of the hydrogen bond. Thus, by analogy with the hydrogen bond, it seems worthwhile to introduce the concept of the halogen bond to the chemical world. Indeed, halogen bonding is now recognized as playing a key role in the field of synthetic chemistry [12, 16–18], supramolecular chemistry [11, 12], material science [19–24] and bioorganic chemistry [25, 26]. Theoretical [27–31] and experimental [10] investigations of the halogen bond show that, like the hydrogen bond, it is predominantly electrostatic in origin. However, because of the greater number of electrons around the halogens than around the hydrogen, polarization [32] and dispersion [33] forces contribute more to the halogen bond than to the hydrogen bond. Moreover, charge transfer seems greater in the halogen interactions [34]. It is therefore not surprising to observe that Lewis bases cannot always be ordered in the same sequence when the equilibrium constants of halogen-bond and hydrogen-bond formation are compared. For example, towards diiodine, the sulfur bases are stronger than the analogous oxygen bases, whereas towards 4-fluorophenol, the oxygen bases are stronger than the analogous sulfur bases. These differences between halogen-bond and hydrogen-bond basicities, together with the growing importance of the halogen bond in chemistry, advocate the construction of a halogen-bond basicity scale. Because there are hundreds of equilibrium constants for the formation of diiodine complexes with Lewis bases (Equation 5.1): I2 + B B · · · I − I

(5.1)

in the literature, whereas far fewer data are available for other halogen-bond donors, diiodine is the obvious choice as the reference Lewis acid for such a construction. At this point, the question of the generality of the diiodine scale should be raised. The test for this generality consists in searching for the existence of linear Gibbs energy relationships between complexation constants referring to diverse halogen-bond donors. Values of the enthalpy of reaction 5.1 can also be gathered to build a halogen-bond affinity scale. Studies of the X Y stretching vibration in complexes of XY with different Lewis bases reveal a characteristic decrease in frequency as the strength of the base increases [35, 36]. Hence spectroscopic scales of halogen-bond basicity can be built [37] in the manner described in Chapter 4 for the O H stretching vibration in hydrogen-bonded complexes. Spectroscopic scales based on the shifts of the ν(I I) band of diiodine at 211 cm−1 , the ν(I Cl) band of iodine monochloride at 376 cm−1 and the ν(I CN) band of iodine cyanide at 485 cm−1 will be presented and compared with thermodynamic basicity and/or affinity scales. Even more characteristic is the blue shift observed for the π g →σ u transition of diiodine (at 521 nm) upon complexation. This blue shift explains the various colours of diiodine in solvents of different basicities. There appears to be a correlation between the strength of the complex and the blue shift of the diiodine visible spectrum [38, 39]. The generality of the diiodine blue shift as a spectroscopic indicator of halogen-bond basicity (affinity) will be analysed. Before defining and tabulating extensively these thermodynamic and spectroscopic scales, we begin with an introduction to the geometry of halogen-bonded complexes.

Thermodynamic and Spectroscopic Scales of Halogen-Bond Basicity and Affinity

5.1

231

Structure of Halogen-Bonded Complexes

The distances d and r and angles α, θ and ϕ that describe the halogen-bond geometry are defined in Figure 5.1. The distances d and r measure the length of the halogen bond and the X Y bond, respectively. The angle α gives a measure of the linearity of the halogen bond. The pseudo-polar angles θ and ϕ describe the position of the halogen X with respect to the Lewis base B, the so-called directionality of the halogen bond. For example, in a complex of a carbonyl base with a dihalogen (Figure 5.1b), if the halogen bond points in the direction of a putative sp2 lone pair on oxygen and is perfectly linear, then θ = 0◦ , ϕ = 120◦ and α = 180◦ . The geometry of halogen bonding has been studied using theoretical calculations, rotational spectroscopy and X-ray crystal diffraction. Representative results are presented below. In spite of the large number of core electrons of heavy halogens, with the associated relativistic effect, ab initio calculations are becoming increasingly used for optimizing the geometry of halogen-bonded complexes. Optimizations carried out at the MP2/LANL2DZ∗ level were shown to reproduce experimental results well for the complexes of dihalogens with pyridine [40], thiocarbonyl bases [41] and carbonyl compounds [42–44]. Theoretical calculations have revealed many geometric features that can influence Lewis basicity, such as: r The existence of two stereoisomeric complexes with thiocarbonyl [41] and carbonyl bases [42], one planar on the non-bonding sulfur (oxygen) lone pair and the other perpendicular on the sulfur (oxygen) π electrons. The perpendicular complex is favoured by bulky substituents on the (thio)carbonyl group and by the π character of the HOMO of the base. r The formation of an N H· · ·I intramolecular hydrogen bond in the dihalogen complexes of thioamides [41] and lactams [44] which assists the sulfur (oxygen)–iodine coordination and enhances the Lewis basicity [44]. Such a chelation of the internal iodine atom by a head-on approach of sulfur and a side-on approach of the NH group leads to a very high complexation constant [45] for the complex of diiodine with methimazole(1methylimidazole-2-thione). Thus methimazole functions as a ‘diiodine sponge’ and acts as an efficient anti-thyroid drug [45]. z

B α

θ d

d

R

X x

α

ϕ r

Y

a

C

y

O

Y X

r

θ

ϕ

R

b

Figure 5.1 (a) Radial and angular descriptors of the halogen-bond geometry; (b) example of a carbonyl base.

232


a

b

c

d

Figure 5.2 MP2/LANL2DZ∗ optimized geometries of halogen-bonded complexes: (a) thiourea–diiodine (planar complex with NH· · ·I hydrogen bond); (b) tetramethylthiourea– diiodine (perpendicular complex); (c) oxiran-2-one–iodine monochloride (anti conformer); (d) γ -butyrolactone–iodine monochloride (syn conformer). Distances are in angstr oms and ˚ ¨ angles in degrees.

r The existence of a syn–anti isomerism in the dihalogen complexes of lactones [43] and lactams [44]. The structures of representative complexes are displayed in Figure 5.2. The gas-phase structures of about 60 complexes of F2 , Cl2 , Br2 , BrCl, ClF and ICl with small Lewis bases such as CO, C2 H2 , C2 H4 , C6 H6 , C5 H5 N, HCN, NH3 , PH3 , H2 S and H2 O have been determined by rotational spectroscopy by Legon’s group at the University of Exeter. Important generalizations about the angular and radial geometries in these complexes have been made in a review [10]. In particular, Legon’s work demonstrates that: r The angular geometries of Y X· · ·B are isostructural with those of Y H· · ·B for a given B. r These angular geometries can be understood in terms of the halogen X aligning with the direction of the axis of the non-bonding electron pair or π -electron cloud on B. r The significant non-linearity of hydrogen bonds (α < 180◦ ) is absent in halogen bonds (α ≈ 180◦ ). The complex furan · · · ClF is the only one to contravene these rules. While HCl and HF form a hydrogen bond to the n-pair of oxygen, ClF interacts with the π -electron density. In other molecules with two potential acceptor sites, the observed geometries give the


a

b

c

d

233

Figure 5.3 Observed gas-phase geometries of halogen-bonded complexes: (a) oms acetylene· · ·Br2 ; (b) H2 O· · ·BrCl; (c) H3 N· · ·ICl; (d) H2 CO· · ·ClF. Distances are in angstr ˚ ¨ and angles in degrees.

following halogen-bond basicity orders: F n-pair(s) > π bond in vinyl fluoride [46], π bonds > S n-pair in thiophene [10], C n-pair > O n-pair in carbon monoxide [10] and exocyclic π bond > pseudo-π bonds in methylenecyclopropane [10]. The geometries of selected halogen-bonded complexes are drawn in Figure 5.3. In addition to leading to geometries, rotational spectroscopy has also provided halogenbond stretching force constants (kσ ) for many dimers B· · ·X Y. Values of kσ can be calculated from the centrifugal distortion constants [10]. Strictly, they measure the restoring force per unit infinitesimal displacement of the halogen bond. Nevertheless, they have been used as a measure of intermolecular binding strength [10, 47]. A number of kσ values are collected in Table 5.1. They enable some Lewis bases to be ranked which could not otherwise be studied because of their reactivity with dihalogens and interhalogens. For example, from the values of kσ for B· · ·ClF [48], the ranking of π and pseudo-π bases is benzene < 1,3-butadiene < allene < cyclopropane ≈ acetylene ≈ methylenecyclopropane < ethylene. The data in Table 5.1 can be compared with the kσ values of the corresponding hydrogenbonded dimers B· · ·H Y [47] by means of Equation 5.2: kσ (B· · ·HY) = a kσ (B· · ·XY) + b

(5.2)

234


Table 5.1 Intermolecular stretching force constants (kσ , N m−1 ) for dimers B· · ·X Y [10]. XY Base Ar N2 CO C2 H2 C2 H4 HCN MeCN H2 O H2 S PH3 NH3

F2

2.61 2.5 3.63 2.36 4.7

Cl2

Br2

3.68 5.6 5.9 6.55

5.0 7.8 8.8

8.0 6.3 5.5 12.71

9.7 10.0 9.8 17.9

BrCl

ClF

ICl

2.8 4.4 6.3 9.4 10.5 11.09

2.8 5.0 7.02 10.0 11.0 12.33 13.9 14.2 13.34

3.2 5.35 8.0 12.2 14.0 14.5

12.5 12.2 11.5 26.7

34.3

15.9 16.6 20.8 30.4

The results are presented in Table 5.2. At first glance, the correlation coefficients r might appear satisfactory. However, they partly result from the weight of van der Waals and weak complexes (with Ar, N2 , CO, C2 H2 and C2 H4 ) in the sample. The correlation falls if medium and strong complexes (with HCN, H2 O, H2 S, PH3 and NH3 ) are considered. In Figure 5.4, it appears that the order of strength of the B· · ·H Cl hydrogen bond is NH3 > H2 O > HCN > H2 S > PH3 whereas for the B· · ·I Cl halogen bond the order becomes NH3 > PH3 > H2 S > H2 O > HCN These observations indicate that the halogen bond will provide information complementary to that of the hydrogen bond for measuring Lewis basicity quantitatively. Many X-ray crystal structures have been reported for halogen-bonded complexes since the first determination in 1954 [5]. In the case of inorganic halogens, the structures of 141 complexes could be found [14] in the October 2002 version of the CSD. They correspond to 219 halogen bonds between the dihalogens I2 (mainly), IBr, ICl, Br2 and Cl2 , and the B atomic centres C (π ), O, S ( S , C S, PS), Se ( Se , C Se, PSe), N (sp3 , sp2 , sp),

Table 5.2 Results of the comparison of halogen-bond force constants with hydrogen-bond force constants (Equation 5.2). H Y

X Y

n

r

a

b

H H H H H H

Cl F Br Cl Br2 I Cl Cl2 Br2

9 10 5 10 9 7

0.905 0.935 0.842 0.867 0.971 0.890

0.98 ± 0.18 0.70 ± 0.09 1.82 ± 0.68 0.54 ± 0.11 1.60 ± 0.15 1.04 ± 0.24

2.7 ± 2.7 −0.3 ± 1.2 −1.5 ± 7.7 −0.4 ± 1.7 −2.2 ± 1.0 −1.9 ± 2.6

F Cl F Cl Cl Cl


235

20

k σ (B...HCl) / N m-1

NH3 H2 O HCN H2 S

C2 H2

Ar

CO N2

C2 H4

PH3

0 0

35

k σ (B...ICl) / N m-1

Figure 5.4 Comparison of intermolecular stretching force constants: hydrogen bond versus halogen bond.

P (sp3 ) and As (sp3 ). A statistical investigation [14] of this crystallographic database shows the following. Halogen-bond length. The d(X· · ·B) distances are intermediate in length between van der Waals contacts and covalent X B bonds. In the complexes of I2 with sulfur bases, ˚ lies closer to the sum of cothe mode of the distribution of S· · ·I distances (≈2.65 A) ˚ ˚ of iodine and valent radii (2.37 A) than to the sum of van der Waals radii (3.79 A) sulfur. Linearity of the halogen bond. The mean of the distribution of Y X· · ·B angles is 176◦ . Halogen bonds I I· · ·B are closer to linearity by 11–38◦ (depending on the family of B) than the corresponding O H· · ·B hydrogen bonds. Lone-pair directionality of the halogen bond. For amines, phosphines, arsines and pyridines, the direction of the halogen bond (that is, of the vector X· · ·B) is very close to the axis of the sp3 or sp2 non-bonding electron pair of nitrogen, phosphorus or arsenic. For thio(seleno)ethers R1 S(Se)R2 and thio(seleno)carbonyl bases R1 C(=S,Se)R2 , the halogen bond is quasi-perpendicular to the R1 S(Se)R2 plane (Figure 5.5a) or to the plane bisector of the angle R1 CR2 (Figure 5.5b). However, for thio(seleno)amides (ureas), the torsion angle NCS(Se)I changes from about 0 or 180◦ (the planar complex found in theoretical calculations, see above) to about 90◦ (the perpendicular complex), when bulky groups R1 and R2 hinder the S(Se) lone pairs (Figure 5.5c). For thio(seleno)phosphoryl bases R3 PS(Se), the average torsion angle RPS(Se)I around the PS(Se) bond is 57◦ , showing that the halogen bond is staggered relative to the three P R bonds (Figure 5.5d), and indicating that the three non-bonding electron pairs on the chalcogen atom are staggered relative to the three substituents on the phosphorus (in agreement with theoretical calculations [49]). Lengthening of the X Y bond. In dimers Y X· · ·B, the X Y bond is elongated compared with the free X Y molecule. The X Y bond is all the more lengthened as the halogen bond X· · ·B becomes shorter. This reciprocal relationship can be explained by a valence model of the halogen bond, in which the valence 1 of halogen is assumed to be shared between

236


a

b

c

d

Figure 5.5 X-ray structures of the diiodine complexes (CSD code in parentheses) of (a) methyl selenide (RIZMES), (b) N-methylthiazolidine-2(3H)-selenone (planar) (YEYFEN), (c) N,N dimethylimidazolidine-2-selenone (perpendicular) (REBNER) and (d) triphenylphosphine sulfide (NOFKOI). Distances are in angstr oms and angles in degrees. ˚ ¨

the bonds it forms, so that

n(X − Y) + n(X · · · B) = 1

(5.3)

where n is the bond order. From Equation 5.3 and well-known empirical expressions between bond order and bond length, analytical expressions can be found for the curves ∆r(X Y) versus d(X· · ·B), which satisfactorily fit the experimental data. The lengthening ∆r(I I) is significantly related to the diiodine basicity of B measured by pK BI2 (see Section 5.2). As displayed in Figure 5.6, the relationship is family dependent. It appears that, for complexes of the same stability, the elongation of the I I bond increases with the degree to which the base can transfer electron density into the σ ∗ (I I) molecular orbital (i.e. in the order N < S < Se < As of the donor atom of the base).


237

6 Diiodine basicity

5 4 3

Ph3As

2 1 0 0

0,1

0,2

0,3

0,4

∆r (I-I)

˚ , the lengthening of Figure 5.6 Family-dependent plot of diiodine basicity against ∆r(I I) (A) the I I bond. Nitrogen bases (squares), sulfur bases (diamonds), selenium bases (stars) and Ph3 As. For the sake of clarity, a number of points have been omitted in each family.

5.2

The Diiodine Basicity Scale

The first equilibrium constants for the formation of diiodine complexes (Equation 5.1) were measured by Benesi and Hildebrand in 1949 [3] for the π bases benzene and mesitylene. Since then, it has been found that the interaction between diiodine and Lewis bases provides a good system for investigating Lewis basicity, and organic bases of all kinds have been studied. In their 1971 review of the spectroscopy of electron donor/acceptor systems, Rao et al. [50] were able to collect more than 300 values of diiodine complexation constants, ranging from 0.026 l mol−1 for 1,2-dichlorobenzene in cyclohexane to 46 600 l mol−1 for tricyclohexylphosphine selenide in chloroform. A 1999 review [51] reports 163 values for only the sulfur and selenium bases. This large number of values, estimated today at more than 1000, makes it possible to construct an extensive diiodine basicity scale. However, the literature data sometimes show serious discrepancies and are often inhomogeneous, referring to different concentration scales, solvents and temperatures. Consequently, it is necessary to select the most reliable determinations and to correct the equilibrium constants to the same concentration scale and the same conditions of solvent and temperature. The following sections explain the construction of the diiodine basicity scale and present the results obtained for about 800 Lewis bases. 5.2.1

Definition of the Scale

Diiodine is chosen as the reference Lewis acid. The standard conditions are T = 25 ◦ C and an alkane (e.g. n-heptane or cyclohexane) is the solvent. Diiodine provides a remarkable opportunity to study halogen-bonded complexes from three regions of the electronic spectrum in which the absorption is directly related either to the concentration of the complex (the charge-transfer band, often around 240–350 nm, and the blue-shifted band, around 400–510 nm) or to the concentration of free diiodine (the visible band at 520 nm in n-heptane). In alkanes, diiodine interacts with Lewis bases B to form stable molecular

238


complexes BI2 , of 1 : 1 stoichiometry if the concentrations of B and I2 are judiciously chosen. In more polar solvents, the complex might dissociate to form ionic species. Different ionization mechanisms have been proposed. In the case of the complex with pyridine (Pyr), reaction 5.4 giving the centrosymmetric iodopyridinium cation and the linear triiodide anion is favoured [52] by an increase in the polarity of the medium: 2 PyrI2 Pyr2 I+ I− 3

(5.4)

The equilibrium expression for the formation of 1 : 1 diiodine complexes is Kc =

(CB0

Cc − Cc )(CI02 − Cc )

(5.5)

where Cc = concentration of the complex BI2 at equilibrium, CB0 = initial concentration of the base and CI02 = initial concentration of diiodine. Since K c values vary from ∼10−1 to ∼106 l mol−1 , it is convenient to use a logarithmic scale. We propose calling this scale pK BI2 , in accordance with the symbol pK BH+ for the protonation of a base B. In defining pK BI2 from Equation 5.6: pK BI2 ≡ log10 K c /1 l mol−1 (5.6) the strongest protonic bases and the strongest diiodine bases have the largest pK BH+ and pK BI2 values, respectively. The Gibbs energy of diiodine complexation at 298 K is given by Equation 5.7: ∆G ◦298 /kJ mol−1 = −5.708 pK BI2 5.2.2

(5.7)

Methods for the Determination of Diiodine Complexation Constants

These constants are usually obtained from the UV–visible spectra of the solutions. When the base does not absorb in the region studied, Equation 5.8: A = εc Cc + εI2 CI2

(5.8)

shows the relationship of the total absorbance (A) to the diiodine concentration and to the complex concentration, at any given wavelength for a cell of 1 cm optical path; εc and εI2 are the molar absorption coefficients of the complex and diiodine, respectively, and CI2 is the concentration of free diiodine at equilibrium. When only one 1 : 1 complex is formed, Equation (5.9): CI02 = Cc + CI2

(5.9)

is also characteristic of the system. The basic Equations 5.5, 5.8 and 5.9 can be combined in various ways, leading to equations in which the two unknowns, K c and εc , are always present. From 1949 to the early 1980s, the Benesi–Hildebrand, Ketelaar, Scott and Rose–Drago equations (reviewed in Chapter 6 of Foster’s book [53]) were solved graphically for a set of several solutions with different base and diiodine concentrations. Typically, CI02 = 3 × 10−3 mol l–1 and CB0 is varied so that the Cc /CB0 ratio is contained between 20 and 80%, according to the criteria of Person [54] and Deranleau [55]. Thanks to rapid developments in computing, data on absorbances and concentrations are now analysed by nonlinear least-squares curve-fitting procedures [51, 56–58].


239

IR spectrometry, NMR spectrometry and calorimetry have also been used [59–61] since IR intensities, NMR chemical shifts and complexation heats are proportional to the concentration of B or BI2 , or to a linear combination of the concentrations of these species. In these methods, the second unknown, in addition to the unknown K c , is either the IR absorption coefficient or the NMR chemical shift of the complex or the calorimetric enthalpy of complexation. The fact that the determination of diiodine complexation constants is dependent on a second unknown makes values more uncertain than the hydrogen-bond formation constants. A revision of the statistical evaluation of diiodine complexation constants obtained by the popular Benesi–Hildebrand method [53] shows [57] that the confidence interval is always much larger than previously reported. Examples of revised 95% confidence limits are (in l mol−1 ): 0.32–0.40, 0.53–1.86 and 1.10–1.32 for the complexation constants of diiodine with 1-bromobutane, benzene and dioxane, respectively. However, better 95% confidence intervals can be obtained. A careful application of the Rose–Drago method to the complexation of diiodine with carbonyl bases gives [62], for example 0.53 ± 0.04 (benzaldehyde), 1.12 ± 0.06 (acetone), 8.1 ± 0.7 (N,N-dimethylbenzamide) and 15 ± 0.4 (N,N-dimethylacetamide) (in l mol−1 , in heptane at 25 ◦ C). 5.2.3

Temperature Correction

In those cases where the literature K c values were at a temperature T other than 298 K, the lnK c versus 1/T relationship was assumed to be linear and K c at 25 ◦ C was calculated from ∆H ◦ T − 298.15 ln K c (298) = ln K c (T ) − (5.10) R 298.15T Examples of the effect of a temperature change from 20 to 25 ◦ C on pK BI2 values are given in Table 5.3. The pK BI2 of Lewis bases is seen to vary slightly with temperature. Hence it is not necessary to know the exact complexation enthalpy to apply a temperature correction. It is sufficient to estimate the complexation enthalpy from data pertaining to closely related Lewis bases. 5.2.4

Solvent Effects

The measurement of diiodine complexation constants has generally been performed in alkanes or chlorinated solvents. It has been observed [63] that the value of the equilibrium Table 5.3 Effect of a temperature change on pK B I 2 values for some Lewis bases. Lewis base Benzene Diethyl ether N,N-Dimethylacetamide Pyridine N,N-Dimethylthioformamide Triethylamine a

∆pK BI2 = pK BI2 (25 ◦ C) − pK BI2 (20 ◦ C).

−∆H◦ (kJ mol−1 )

∆pK BI2 a

pK BI2 (25 ◦ C)

6.8 17.6 22.4 31.3 43.1 50.2

−0.02 −0.05 −0.06 −0.09 −0.13 −0.15

−0.62 −0.05 1.18 2.21 3.14 3.67

240


constant is markedly affected by the solvent. Thus, the construction of a homogeneous diiodine basicity scale requires one standard solvent to be chosen and the values obtained in other solvents to be referred to that reference solvent. The inert, UV-transparent, low-toxic and common solvent heptane is the solvent of choice for evident technical and theoretical reasons. Its major limitation is the difficulty of solubilizing Lewis bases and diiodine complexes that are strongly dipolar. It may then be necessary to use media of a greater solubilization power than heptane, for example cyclohexane, CCl4 , CH2 Cl2 and CHCl3 . The comparison, made in Table 5.4, of K c values measured in heptane and in other alkanes (cyclohexane, hexane, methylcyclohexane) for a set of 30 diverse bases, shows that the differences are generally within the sum of experimental uncertainties. In the following, the K c values measured in cyclohexane, hexane and methylcyclohexane will therefore be considered equivalent to, and will be mixed with, those measured in heptane. The K c values measured in CCl4 are significantly lower than those in alkanes, except in the case of selenoethers. It was not possible to find a general relationship between the K c values determined in these two solvents. However, linear Gibbs energy relationships of good precision were found [84] for families of structurally related bases. The LFER equation is pK BI2 ≡ log K c (in alkanes) = a logK c (in CCl4 ) + b

(5.11)

Figure 5.7 illustrates typical relationships and Table 5.5 lists values of a and b (with their standard deviations sa and sb ), the squared correlation coefficient r2 , the standard deviation s and the number of bases n. The data on which the relationships are based are given in Tables 5.6–5.14 and 5.17–5.20. In the absence of data in heptane or cyclohexane for a given base to obtain the pK BI2 value from Equation 5.6, a secondary value may be obtained from Equation 5.11 using data in CCl4 . When the diiodine complexation constants cannot be measured either in alkanes or in CCl4 , solvents such as CH2 Cl2 or CHCl3 are generally used. To establish relationships between the K c values determined in alkanes and those determined in CH2 Cl2 or CHCl3 , there are only enough literature data for pyridines and thiocarbonyl compounds. The equations found are as follows: (i) for pyridines: pK BI2 ≡ logK c (in alkanes) = 0.883 (±0.019) logK c (in CH2 Cl2 , 30 ◦ C) +0.46 (±0.04) n = 8, r 2 = 0.997, s = 0.04

(5.12)

pK BI2 ≡ logK c (in alkanes) = 0.785 (±0.042) logK c (in CH2 Cl2 , 25 ◦ C) +0.58 (±0.06) n = 4, r 2 = 0.994, s = 0.04

(5.13)

pK BI2 ≡ logK c (in alkanes) = 1.032 (±0.051) logK c (in CHCl3 ) + 0.42 (±0.10) (5.14) n = 12, r 2 = 0.976, s = 0.05


241

Table 5.4 Diiodine complexation constants (l mol−1 ) measured in heptane and in various alkanes. Lewis base

LogKc

Ref.

LogKc

Heptane Benzene Dibutyl ether 4-Chloroacetophenone Diethyl ether Ethyl acetate Butyronitrile Acetophenone Hexamethylbenzene 4-Methylacetophenone 4-Methoxyacetophenone 7-Oxabicyclo[2.2.1]heptane Diphenyl sulfide Thioanisole N,N-Dimethylacetamide Pyrazine Triethoxyphosphine sulfide 3-Chloropyridine tert-Butyl phenyl sulfide Selenoanisole Quinoline n-Propyl sulfide Pyridine tert-Butyl phenyl selenide 2-Methylpyridine Isoquinoline 3-Methylpyridine 4-Methylpyridine 4-Ethylpyridine Trioctylphosphine oxide Trioctylphosphine sulfide Butyronitrile 3-Chloropyridine Quinoline Pyridine 2-Methylpyridine Isoquinoline Benzene Hexamethylbenzene a

D = logKc (in heptane) – logKc (in alkanes).

Da

Cyclohexane

−0.61 −0.16 −0.09 −0.06 −0.03 0.00 0.06 0.14 0.16 0.29 0.52 0.59 0.99 1.18 1.18 1.27 1.36 1.38 1.79 2.10 2.17 2.18 2.29 2.35 2.35 2.47 2.57 2.61 2.89 3.79 Heptane

[64] [65] [67] [69] [70] [71] [67] [72] [67] [67] [73] [74] [75] [70] [65] [78] [79] [74] [75] [79] [74] [79] [74] [80] [79] [80] [80] [79] [65] [83]

−0.59 −0.14 −0.10 −0.05 −0.07 −0.03 0.08 0.16 0.21 0.31 0.53 0.53 1.04 1.12 1.08 1.26 1.26 1.41 1.85 2.05 2.21 2.11 2.39 2.27 2.27 2.39 2.53 2.56 2.77 3.67

0.00 1.36 2.10 2.18 2.35 2.35 Heptane

[71] [79] [79] [79] [80] [79]

−0.03 1.34 2.09 2.12 2.26 2.30

[64] [72]

−0.60 0.15

−0.61 0.14

Ref.

[64] [66] [68] [65] [65] [71] [68] [72] [68] [68] [65] [74] [75] [76] [77] [65] [79] [74] [75] [79] [74] [79] [74] [81] [79] [82] [82] [79] [65] [65] Hexane

−0.02 −0.02 0.01 −0.01 0.04 0.03 −0.02 −0.02 −0.05 −0.02 −0.01 0.06 −0.05 0.06 0.10 0.01 0.10 −0.03 −0.06 0.05 −0.04 0.07 −0.10 0.08 0.08 0.08 0.04 0.05 0.12 0.12

[71] 0.03 [79] 0.02 [79] 0.01 [79] 0.06 [81] 0.09 [79] 0.05 Methylcyclohexane [64] [72]

−0.01 −0.01

242


pK BI2

2

1

0

-1 -1

0

1

2

3

log K c (in CCl4)

Figure 5.7 Family-dependent relationships between diiodine complexation constants measured in alkanes, pK BI2 , and in CCl4 , logKc (in CCl4 ): nitriles ( ), phosphoryls (◦), π bases (×), thioethers (♦), pyridines () and selenoethers (∗). The dashed line of slope unity is drawn for comparison.

(ii) for thiocarbonyls: pK BI2 ≡ logK c (in alkanes) = 0.902 (±0.012) logK c (in CH2 Cl2 ) − 0.15 (±0.04) (5.15) n = 5, r 2 = 0.999, s = 0.03 pK BI2 ≡ logK c (in alkanes) = 1.002 (±0.033) logK c (in CHCl3 ) − 0.16 (±0.09) (5.16) n = 3, r 2 = 0.999, s = 0.07 Table 5.5 Family-dependent relationships (Equation 5.11) between diiodine complexation constants measured in alkanes and in CCl4 . Family Arenes Pyridines Nitriles Ethers Carbonyls Thionyls Phosphoryls Thiocarbonyls Thioethers Thiophosphoryls Selenoethers

a ± sa

b ± sb

r2

s

n

0.875 ± 0.073 1.024 ± 0.002 1.143 ± 0.035 1.106 ± 0.042 1.125 ± 0.03 1.13 1.163 ± 0.003 1.048 ± 0.029 0.990 ± 0.010 1.137 0.985 ± 0.075

0.15 ± 0.04 0.144 ± 0.004 0.32 ± 0.01 0.22 ± 0.01 0.21 ± 0.02 0.37 0.40 ± 0.01 −0.05 ± 0.09 0.07 ± 0.02 0.09 0.00 ± 0.18

0.966 0.998 0.995 0.996 0.986

0.05 0.03 0.02 0.02 0.08

1.000 0.997 0.999

0.004 0.07 0.02

0.989

0.05

7 10 7 5 21 2 4 6 13 2 4


5.2.5

243

Data

In this way, an extensive set of pK BI2 values has been derived and is given in Tables 5.6–5.15 and 5.17–5.22. The data are divided and presented in a manner similar to that used in Chapter 4 for the pK BHX scale. When the mentioned solvent is an alkane, pK BI2 and ∆G◦ are primary values calculated through Equations 5.6 and 5.7, respectively. When the mentioned solvent is CCl4 , CH2 Cl2 or CHCl3 , Tables 5.6–5.22 give the K c values measured in the given solvent, while pK BI2 and ∆G◦ are secondary values referring to alkanes and are calculated through Equations 5.11–5.16. If several references are quoted for the same compound, the value given to K c is the average of the various determinations. The contents of Tables 5.6–5.22 are the following: Table 5.6 – Carbon π bases: aromatic hydrocarbons, alkenes, thiophenes, furans, halobenzenes (π and halogen bases). Table 5.7 – Nitrogen bases: primary, secondary and tertiary amines and aminoboranes. Table 5.8 – Nitrogen bases: six-membered N-heteroarenes, with the peculiar set of 1,10phenanthrolines, which form three-centre halogen bonds. Table 5.9 – Nitrogen bases: five-membered N-heterocycles. Table 5.10 – Nitrogen bases: nitriles. Table 5.11 – Oxygen bases: ethers. Table 5.12 – Oxygen bases: carbonyl compounds. Table 5.13 – Oxygen bases: thionyl compounds. Table 5.14 – Oxygen bases: phosphoryl compounds. Table 5.15 – Oxygen bases: amine N-oxides and pyridine N-oxides. The insolubility of most N-oxides in alkanes prevents referring the K c values measured in CCl4 , C6 H6 and/or CH2 Cl2 to the standard solvent. Table 5.16 – Miscellaneous oxygen bases. The insolubility of these bases in alkanes prevents referring the K c values measured in CCl4 , C6 H6 , CH2 Cl2 and/or CHCl3 to the standard solvent. Table 5.17 – Sulfur bases: thiocarbonyl compounds. Table 5.18 – Sulfur bases: thiols, thioethers and disulfides. Table 5.19 – Sulfur bases: thiophosphoryl compounds. Table 5.20 – Selenium bases: selenides (selenoethers), selenocarbonyl and selenophosphoryl compounds. Table 5.21 – Halogen bases: chloro-, bromo- and iodoalkanes. The complexes of diiodine with fluoroalkanes are too weak to be studied. Table 5.22 – Miscellaneous bases: tetramethylguanidine, triphenylarsine, isothiocyanates and drugs. Many Lewis bases in Tables 5.6–5.22 contain several potential sites for halogen bonding. In the usual UV–visible method, the concentration of base is maintained in excess so that it may be safely assumed that almost only 1 : 1 complexes are formed. The observed (total) constant, K c (t), is the sum of the individual constants for each 1 : 1 complex [85]. Three cases are encountered:

244

1,2,5-Trimethylbenzene 1,3,5-Triethylbenzene 1,2,3-Trimethylbenzene 1,2,4-Trimethylbenzene 1,3-Dimethylbenzene, m-xylene

1,2-Dimethylbenzene, o-xylene

Phenyltrimethylsilane 1,4-Dimethylbenzene, p-xylene

Hexaethylbenzene Cyclohexylbenzene tert-Butylbenzene Isopropylbenzene n-Propylbenzene Isobutylbenzene n-Butylbenzene Ethylbenzene Methylbenzene, toluene

12

13 14

15 16 17 18 19 20 21 22 23

Pentamethylbenzene 1,2,3,4-Tetramethylbenzene, prehnitene 1,2,3,5-Tetramethylbenzene 1,2,4,5-Tetramethylbenzene, durene 1,3,5-Trimethylbenzene, mesitylene

2 3 4 5 6

7 8 9 10 11

C6 (CH3 )6

Alkylbenzenes and phenyltrimethylsilane Hexamethylbenzene

1

C6 (C2 H5 )6 C6 H5 C6 H11 C6 H5 C(CH3 )3 C6 H5 CH(CH3 )2 C6 H5 CH2 CH2 CH3 C6 H5 CH2 CH(CH3 )2 C6 H5 (CH2 )3 CH3 C6 H5 CH2 CH3 C6 H5 CH3

C6 H5 Si(CH3 )3 C6 H4 (CH3 )2

C6 H4 (CH3 )2

C6 H3 (CH3 )3 C6 H3 (C2 H5 )3 C6 H3 (CH3 )3 C6 H3 (CH3 )3 C6 H4 (CH3 )2

C6 H(CH3 )5 C6 H2 (CH3 )4 C6 H2 (CH3 )4 C6 H2 (CH3 )4 C6 H3 (CH3 )3

Formula

Lewis base

No. Hept CCl4 CCl4 CCl4 CCl4 CCl4 Hept CCl4 CCl4 CCl4 CCl4 CCl4 Hept CCl4 Hept CCl4 cHex Hept CCl4 CCl4 CCl4 CCl4 CCl4 CCl4 CCl4 CCl4 CCl4 Hept CCl4

Solvent 1.37 1.02 0.99 0.82 0.78 0.62 0.87 0.61 0.55 0.51 0.43 0.41 0.57 0.31 0.52 0.29 0.52 0.46 0.30 0.25 0.24 0.23 0.23 0.23 0.22 0.22 0.21 0.36 0.19

Kc

−0.81 −0.40 −0.29 0.21 0.35 0.47 0.63 1.00 1.11 1.39 1.62 1.62 1.92 2.18 2.27 2.36 2.36 2.36 2.46 2.46 2.56 2.53

−0.08 −0.11 −0.18 −0.19 −0.24 −0.28 −0.28 −0.34 −0.39 −0.40 −0.41 −0.41 −0.42 −0.43 −0.43 −0.44 −0.44

−0.78

∆G◦

0.14 0.07 0.05 −0.04 −0.06

0.14

pK BI2

[72] [86] [87–89] [87, 90] [88] [87–89, 91] [90] [88, 89, 91] [87] [91] [88] [88] [90] [87–89] [90] [87–89] [92] [90] [87–89, 91] [89, 91] [87] [86, 87] [86, 87] [86, 87] [87] [86, 87] [90] [87] [86–89]

Ref.

Table 5.6 Equilibrium constants Kc (l mol−1 ) and diiodine basicity scales pK B I 2 and ∆G◦ (kJ mol−1 ) for the complexes of diiodine with arenes, heteroarenes and alkenes.

245

46 47

34 35 36 37 38 39 40 41 42 43 44 45

28 29 30 31 32 33

24 25 26 27

C6 H5 (CH2 )4 CH3 C6 H5 (CH2 )5 CH3 C6 H5 (CH2 )6 CH3 C6 H6

Halobenzenes and chlorophenylsilanes Iodobenzene C6 H5 I Chlorodimethylphenylsilane C6 H5 (CH3 )2 SiCl Dichloromethylphenylsilane C6 H5 (CH3 )SiCl2 Bromobenzene C6 H5 Br Chlorobenzene C6 H5 Cl Fluorobenzene C6 H5 F Polyfunctional and polycyclic arenes 2,2-Diphenylpropane CH3 C(C6 H5 )2 CH3 Triphenylmethane (C6 H5 )3 CH Di-p-tolylmethane (CH3 C6 H4 )2 CH2 1,2-Di-p-tolylethane CH3 C6 H4 CH2 CH2 C6 H4 CH3 1,2-Diphenylethane C6 H5 CH2 CH2 C6 H5 b Acenaphthene Stilbene C6 H5 CH CHC6 H5 Styrene C6 H5 CH CH2 Diphenylmethane C6 H5 CH2 C6 H5 Biphenyl C6 H5 C6 H5 b Naphthalene b Phenanthrene Thiophenes and furans b Benzo[b]thiophene b Thiophene ◦ b Thiophene (20 C)

n-Pentylbenzene n-Hexylbenzene n-Heptylbenzene Benzene

cHex Hept Hex

CCl4 CCl4 CCl4 CCl4 CCl4 CCl4 CCl4 CCl4 CCl4 CCl4 CCl4 CCl4

CCl4 cHex cHex CCl4 CCl4 CCl4

CCl4 CCl4 CCl4 Hept CCl4

1.15 0.55 0.61

1.64 1.32 0.83 0.68 0.50 0.42 0.31 0.31 0.28 0.26 0.24 0.20

0.34 0.33 0.24 0.11 0.07 0.06

0.20 0.17 0.17 0.24 0.16

1.48

(Continued)

[94] [95] [96] −0.26

a

a

a

a

a

a

a

a

a

a

a

a

[87] [92] [92] [87] [87] [87]

[87] [87] [87] [90] [86–89]

[93] [93] [93] [93] [89, 93] [93] [89] [89] [93] [87] [87] [87, 89]

5.27

2.66 3.02 3.02 3.54

a

−0.92

a

a

a

a

a

−0.47 −0.53 −0.53 −0.62

246 46

S

c

b

a

44

2,2,4-Trimethylpentane.

39

Polyfunctional base. Formulae:

45

47

S

52

O

C7 H12 C5 H8 c-(CH2 )3 CH2 c-(CH2 )7 CH2 c-(CH2 )5 CH2 C8 H14

b

b

53

54

CH

55

0.69 0.59 0.53 0.49 0.45 0.45 0.43 0.42 0.17

C8 H18 c C8 H18 c C8 H18 c C8 H18 c C8 H18 c C8 H18 c C8 H18 c C8 H18 c C8 H18 c

b

53 54 55 56 57 58 59 60 61

b

0.79 0.81 1.19 0.36 0.19

Hex Hex Hex Hept Hept

2-CH3 C4 H3 S 3-CH3 C4 H3 S 2,5-(CH3 )2 C4 H2 S 2-CH3 C4 H3 O

2-Methylthiophene (20 ◦ C) 3-Methylthiophene (20 ◦ C) 2,5-Dimethylthiophene (20 ◦ C) 2-Methylfuran Furan Alkenes Norbornene Methylenecyclohexane Cyclohexene Cycloheptene Cyclopentene Methylenecyclobutane Methylenecycloheptane Methylenecyclopentane Cyclooctene

Kc

48 49 50 51 52

Solvent

Formula

Lewis base

No.


2.53 4.12 0.92 1.31 1.57 1.77 1.98 1.98 2.09 2.15 4.39

−0.16 −0.23 −0.28 −0.31 −0.35 −0.35 −0.37 −0.38 −0.77

∆G◦

−0.45 −0.73

pK BI2

[98] [98] [98] [98] [98] [98] [98] [98] [98]

[96] [96] [96] [95] [95, 97]

Ref.

247

25 26 27 28 29

2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24

1

No.

Ammonia Primary amines n-Butylamine Isopropylamine tert-Butylamine sec-Butylamine Dodecylamine Isobutylamine Ethylamine Methylamine Phenethylamine DL-Norephedrine 4-Anisidine 4-Toluidine 3,4-Dimethylaniline 2-Toluidine 3-Toluidine 2,5-Dimethylaniline Aniline 2,3-Dimethylaniline 3-Anisidine 4-Bromoaniline 2,6-Dimethylaniline 3-Bromoaniline 3-Chloroaniline Secondary amines 1,3-Di-4-piperidylpropane Pyrrolidine Piperidine Diethylamine Piperazine

Lewis base

MecHex Hept Hept Hept Hept

b

b

c-(CH2 )4 NH c-(CH2 )5 NH (C2 H5 )2 NH

Hept Hept Hept Hept Hept Hept Hept Hept cHex cHex CCl4 CCl4 Hept CCl4 CCl4 CCl4 CCl4 Hept CCl4 CCl4 Hept CCl4 CCl4

Hept

Solvent

CH3 (CH2 )3 NH2 (CH3 )2 CHNH2 (CH3 )3 CNH2 C2 H5 CH(CH3 )NH2 CH3 (CH2 )11 NH2 (CH3 )2 CHCH2 NH2 CH3 CH2 NH2 CH3 NH2 C6 H5 CH2 CH2 NH2 C6 H5 CHOHCH(CH3 )NH2 4-MeOC6 H4 NH2 4-MeC6 H4 NH2 3,4-Me2 C6 H3 NH2 2-MeC6 H4 NH2 3-MeC6 H4 NH2 2,5-Me2 C6 H3 NH2 C6 H5 NH2 2,3-Me2 C6 H3 NH2 3-MeOC6 H4 NH2 4-BrC6 H4 NH2 2,6-Me2 C6 H3 NH2 3-BrC6 H4 NH2 3-ClC6 H4 NH2

NH3

Formula

46 000 7032 7001 5395 11 000

993 898 725 608 630 552 560 433 410 372 18.0 13.8 16 12.4 9.2 8.9 8.8 6.7 5.7 3.7 3.7 2.0 1.8

58

Kc a [99]

[100] [101] [100] [100] [102] [100] [99] [99] [103] [103] [104] [104] [105] [104] [104] [104] [104] [105] [104] [104] [105] [104] [104] [106] [107] [107] [99] [65]

−10.05 −17.12 −16.84 −16.32 −15.87 −15.98 −15.64 −15.70 −15.07 −14.90 −14.67 −6.85

−4.74 −3.25

−24.89 −21.98 −21.98 −21.29 −21.35

3.00 2.95 2.86 2.78 2.80 2.74 2.75 2.64 2.61 2.57 1.20c

0.83c 0.57c

4.36d 3.85 3.85 3.73 3.74d

1.76

(Continued)

Ref.

pK BI2

∆G◦

Table 5.7 Equilibrium constants Kc (l mol−1 ) and diiodine basicity scales pK B I 2 and ∆G◦ (kJ mol−1 ) for the complexes of diiodine with amines and aminoboranes.

248

Lewis base

Dimethylamine Di-n-propylamine Di-n-dodecylamine Morpholine

D-Pseudoephedrine Diisopropylamine N-Methylaniline Diphenylamine Tertiary amines Quinuclidine Diazabicyclooctane (DABCO) N-Isopropylpyrrolidine N-n-Butylpyrrolidine Tetramethyl-1,4-butanediamine Tetramethyl-1,3-propanediamine N,N-Dimethylcyclohexylamine Tetramethyl-1,6-hexanediamine Trimethylamine N,N-Dimethylbutylamine N-Ethylpiperidine Tetramethyl-1,2-ethanediamine N,N-Dimethyl-tert-butylamine Triethylamine Promazine (26◦ C) N,N-Dimethyl-sec-butylamine Tri-n-octylamine Chlorpromazine

Hexamethylenetetramine N-Isopropylpiperidine N,N-Dicyclohexylmethylamine N-Isobutylpyrrolidine

No.

30 31 32 33

34 35 36 37

38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55

56 57 58 59


c-(CH2 )5 NCH(CH3 )2 (C6 H11 )2 NCH3 c-(CH2 )4 NCH2 CH(CH3 )2

b

b

b

C2 H5 CH(CH3 )N(CH3 )2 (n-C8 H17 )3 N

b

c-(CH2 )4 NCH(CH3 )2 c-(CH2 )4 N (CH2 )3 CH3 (CH3 )2 N(CH2 )4 N(CH3 )2 (CH3 )2 N(CH2 )3 N(CH3 )2 c-C6 H11 N(CH3 )2 (CH3 )2 N(CH2 )6 N(CH3 )2 (CH3 )3 N CH3 (CH2 )3 N(CH3 )2 c-(CH2 )5 NCH2 CH3 (CH3 )2 N(CH2 )2 N(CH3 )2 (CH3 )3 CN(CH3 )2 (C2 H5 )3 N

b

b

C6 H5 CHOHCHCH3 NHCH3 ((CH3 )2 CH)2 NH C6 H5 NHCH3 (C6 H5 )2 NH

b

b

(CH3 )2 NH (C3 H7 )2 NH (C12 H25 )2 NH

Formula

Hept Hept Hept Hept MecHex MecHex Hept MecHex Hept Hept Hept MecHex Hept Hept Hept Hept Hept Hept CCl4 MecHex Hept Hept Hept

Hept Hept Hept Hept CCl4 cHex Hept CCl4 CCl4

Solvent

164 840 92 462 16 971 15 838 20 000 19 462 9765 19 000 7540 6913 6602 12 594 5244 4690 4500 4071 3388 2982 2221 10 276 2088 1658 1174

5137 3776 1460 1066 762 1082 703 20.4 1.5

Kc a

3.41d 3.32 3.22 3.07

5.22 4.69d 4.23 4.20 4.00d 3.99d 3.99 3.98d 3.88 3.84 3.82 3.80d 3.72 3.67 3.65 3.61 3.53 3.47

3.03 2.85

3.71 3.58 3.16 3.03

pK BI2

[99] [107] [102] [108] [108] [103] [107] [104] [104] [107] [106] [107] [107] [106] [106] [107] [106] [99] [107] [107] [106] [107] [99] [109] [107] [102] [109] [110] [106] [107] [107] [107]

−21.18 −20.43 −18.04 −17.28

−29.80 −26.77 −24.14 −23.97 −22.83 −22.77 −22.77 −22.72 −22.15 −21.92 −21.80 −21.69 −21.23 −20.95 −20.83 −20.61 −20.15 −19.81 −19.46 −18.95 −18.38 −17.52

−17.30 −16.27

Ref.

∆G◦

249

Tri-n-butylamine Tri-n-propylamine Tri-n-dodecylamine N,N-Diisopentyloctylamine N,N-Di-n-octylisopentylamine N,N-Diisopentylhexadecylamine Triisopentylamine Tetramethylmethanediamine N-sec-Butylpiperidine N-Isobutylpiperidine N,N-Di-n-propylisobutylamine N,N-Di-n-propyl-sec-butylamine N,N-Dimethyl-4-toluidine N,N-Diisobutylmethylamine N,N-Dimethylaniline N,N-Diisobutyl-n-propylamine N,N-Dimethyl-2-toluidine N,N-Dimethyl-2,6-xylidine Aminoboranes Piperidinodimethylborane (Dimethylamino)dimethylborane (Dimethylamino)di-n-propylborane c-(CH2 )5 NB(CH3 )2 (CH3 )2 NB(CH3 )2 (CH3 )2 NB(C3 H7 )2

(CH3 (CH2 )3 )3 N (CH3 (CH2 )2 )3 N (C12 H25 )3 N [(CH3 )2 CHCH2 CH2 ]2 NC8 H17 (C8 H17 )2 NCH2 CH2 CH(CH3 )2 [(CH3 )2 CHCH2 CH2 ]2 NC16 H33 [(CH3 )2 CHCH2 CH2 ]3 N (CH3 )2 NCH2 N(CH3 )2 c-(CH2 )5 NCH(CH3 )C2 H5 c-(CH2 )5 NCH2 CH(CH3 )2 (C3 H7 )2 NCH2 CH(CH3 )2 (C3 H7 )2 NCH(CH3 )C2 H5 4-CH3 C6 H4 N(CH3 )2 [(CH3 )2 CHCH2 ]2 NCH3 C6 H5 N(CH3 )2 [(CH3 )2 CHCH2 ]2 NC3 H7 2-CH3 C6 H4 N(CH3 )2 2,6-(CH3 )2 C6 H3 N(CH3 )2 CCl4 CCl4 CCl4

Hept Hept Hept Hept Hept Hept Hept MecHex Hept Hept Hept Hept Hept Hept Hept Hept Hept Hept 1310 287 157

1125 983 987 987 839 813 688 1374 549 219 115 100 36 28 21 13 2.2 1.8

3.05 2.99 2.99 2.99 2.92 2.91 2.84 2.84d 2.74 2.34 2.06 2.00 1.56c 1.45 1.32c 1.11 0.34c 0.26c

−17.41 −17.07 −17.07 −17.07 −16.67 −16.61 −16.21 −16.21 −15.64 −13.36 −11.76 −11.42 −8.90 −8.28 −7.53 −6.34 −1.94 −1.48 [112] [112] [112]

[99] [99] [102] [102] [102] [102] [102] [106] [107] [107] [107] [107] [111] [107] [111] [107] [111] [111]

S

N

(CH2)3NMe2

25

(CH2)3

NH

R

52 55

HN

R=H R=Cl

29

NH

O

N

56

N

N

33

May contain a contribution of the π site. d Statistically corrected to put the pK BI2 value on a per nitrogen basis.

c

HN

N

NH

N

38

N

39

N

When necessary, Kc at 298 K is calculated from a value at a different temperature by means of the van’t Hoff equation. If the enthalpy of complexation is unknown, we assume values of −32, −41 and −51 kJ mol−1 for primary, secondary and tertiary amines, respectively. b Formulae:

a

78 79 80

60 61 62 63 64 65 66 67 68 69 70 71 72 73 74 75 76 77

250

4-Vinylpyridine Isoquinoline

3-Methylpyridine

trans-4-Styrylpyridine 2-Aminopyridine cis-4-Styrylpyridine

16 17

18

19 20 21

2-Amino-4-methylpyridine 4-Isopropylpyridine 4-tert-Butylpyridine 4-Methoxypyridine 4-Ethylpyridine

5 6 7 8 9

2,5-Dimethylpyridine 2,4-Dimethylpyridine 3-Ethylpyridine 3-Aminopyridine 4-Hydroxypyridine 4-Methylpyridine

3,5-Dimethylpyridine

4

10 11 12 13 14 15

4-Dimethylaminopyridine 4-Aminopyridine 3,4-Dimethylpyridine

Lewis base

1 2 3

No.

4-PhCH CHC5 H4 N 2-NH2 C5 H4 N 4-PhCH CHC5 H4 N

3-MeC5 H4 N

a

4-CH2 CHC5 H4 N

2,5-Me2 C5 H3 N 2,4-Me2 C5 H3 N 3-EtC5 H4 N 3-H2 NC5 H4 N 4-HOC5 H4 N 4-MeC5 H4 N

2-NH2 -4-MeC5 H3 N 4-i-PrC5 H4 N 4-t-BuC5 H4 N 4-MeOC5 H4 N 4-EtC5 H4 N

3,5-Me2 C5 H3 N

4-Me2 NC5 H4 N 4-H2 NC5 H4 N 3,4-Me2 C5 H3 N

Formula Hept cHex Hept CCl4 CH2 Cl2 cHex CH2 Cl2 cHex CHCl3 CCl4 CH2 Cl2 Hept CCl4 cHex Hex CCl4 cHex CH2 Cl2 Hept CCl4 CH2 Cl2 CH2 Cl2 Hept CCl4 Hept CCl4 CH2 Cl2 Hex cHex Hex (30)

(30) (30)

(30)

(30)

(30)

(30)

Solvent (t/◦ C) 6000 1580 697 421 507 599 349 580 166 300 290 408 248 398 390 244 380 246 368 216 249 218 294 149 296 207 208 287 280 246

Kc

2.46 2.45 2.39

2.47

2.52 2.47

2.60 2.59 2.59 2.58 2.57 2.57

2.76 2.71 2.68 2.63 2.61

2.78

3.78 3.20 2.84

pK BI2

[113] [114] [84] [115] [116] [117] [116] [114] [118] [115] [116] [79] [115] [117] [79] [115] [114] [116] [80] [115] [116] [116] [80] [79] [80] [115] [116] [119] [114] [119]

−21.58 −18.27 −16.21

−14.04 −13.98 −13.64

−14.10

−14.38 −14.10

−14.84 −14.78 −14.78 −14.73 −14.67 −14.67

−15.75 −15.47 −15.30 −15.01 −14.90

−15.87

Ref.

∆G◦

Table 5.8 Equilibrium constants Kc (l mol−1 ) and diiodine basicity scales pK B I 2 and ∆G◦ (kJ mol−1 ) for the complexes of diiodine with six-membered N-heteroarenes.

251

6-Methoxyquinoline 2-Methylpyridine 4-Methyl-cis-2-styrylpyridine 2,4,6-Trimethylpyridine 6-Methylquinoline Pyridine

cis-3-Styrylpyridine trans-3-Styrylpyridine 3-Hydroxypyridine 5-Methyl-cis-2-styrylpyridine 2-Amino-6-methylpyridine 7-Azaindole 2-Ethylpyridine 2-Anilinopyridine Phenanthridine Quinoline

Cinnoline 2-Amino-4,6-dichloropyrimidine 2,6-Dimethylquinoline 3-Methyl-cis-2-styrylpyridine cis-2-Styrylpyridine 2,6-Dimethylpyridine 6-Chloroquinoline

4,4 -Bipyridine 4-Acetylpyridine 2-Methylquinoline, quinaldine 4-Chloropyridine 4-Methyl-trans-2-styrylpyridine 5-Methyl-trans-2-styrylpyridine

22 23 24 25 26 27

28 29 30 31 32 33 34 35 36 37

38 39 40 41 42 43 44

45 46 47 48 49 50

a

4-ClC5 H4 N 4-Me-2-(PhCH CH)C5 H3 N 5-Me-2-(PhCH CH)C5 H3 N

a

(C5 H4 N) 2 4-MeCOC5 H4 N

a

3-Me-2-(PhCH CH)C5 H3 N 2-PhCH CHC5 H4 N 2,6-Me2 C5 H3 N

a

a

a

a

a

2-EtC5 H4 N 2-C6 H5 NHC5 H4 N

a

3-PhCH CHC5 H4 N 3-PhCH CHC5 H4 N 3-HOC5 H4 N 5-Me-2-(PhCH CH)C5 H3 N 2-NH2 -6-MeC5 H3 N

C5 H5 N

a

2-MeC5 H4 N 4-Me-2-(PhCH CH)C5 H3 N 2,4,6-Me3 C5 H2 N

CH2 Cl2 Hept Hex Hept Hept Hept CCl4 CH2 Cl2 (30) Hex Hex CH2 Cl2 (30) Hex cHex CCl4 Hept CCl4 CHCl3 Hept CCl4 CHCl3 CCl4 CH2 Cl2 Hex Hex cHex Hept CCl4 CHCl3 Hept CCl4 CH2 Cl2 (30) Hex Hex

185 225 200 192 175 166 108 104 162 157 85 145 130 81 122 77 41 115 86 37b 135 66 95 80 71 67 31 44 63 41 27 50 50 1.81c 1.80 1.79 1.71 1.70 1.70

2.03c 2.01 1.98 1.90 1.85 1.83

2.21 2.20 2.16 2.16 2.11 2.10 2.09 2.08 2.08 2.06

2.36 2.35 2.30 2.28 2.24 2.22

−10.33 −10.27 −10.22 −9.76 −9.70 −9.70

−11.59 −11.47 −11.30 −10.85 −10.56 −10.45

−12.61 −12.56 −12.33 −12.33 −12.04 −11.99 −11.93 −11.87 −11.87 −11.76

−13.47 −13.41 −13.13 −13.01 −12.79 −12.67

(Continued)

[120] [80] [121] [65] [80] [65] [122] [116] [119] [119] [116] [121] [114] [123] [65] [123] [124] [80] [79] [125] [126] [120] [121] [119] [117] [80] [79] [127] [65] [122] [116] [121] [121]

252

Lewis base

3-Bromoquinoline 2-Isopropylpyridine (Z)-2-(α-Methylstyryl)pyridine Methyl nicotinate Phthalazine 3-Acetylpyridine 3-Fluoropyridine 5-Methylpyrimidine Quinazoline 1,5-Naphthyridine 6-Methyl-cis-2-styrylpyridine trans-2-Styrylpyridine 3-Bromopyridine

3-Chloropyridine

Pyridazine 3-Methyl-trans-2-styrylpyridine Pyrimidine 2-Dimethylaminopyridine 4-Cyanopyridine 2,2 -Bipyridine Quinoxaline (E)-2-(α-Methylstyryl)pyridine Pyrazine Tetramethylpyrazine 3-Cyanopyridine 3,5-Dichloropyridine

4-Nitropyridine 2-Bromopyridine

No.

51 52 53 54 55 56 57 58 59 60 61 62 63

64

65 66 67 68 69 70 71 72 73 74 75 76

77 78


4-O2 NC5 H4 N 2-BrC5 H4 N

Me4 C4 N2 3-N CC5 H4 N 3,5-Cl2 C5 H3 N

a

a

a

2-Me2 NC5 H4 N 4-N CC5 H4 N (C5 H4 N) 2

a

3-Me-2-(PhCH CH)C5 H3 N

a

3-ClC5 H4 N

6-Me-2-(PhCH CH)C5 H3 N 2-PhCH CHC5 H4 N 3-BrC5 H4 N

a

a

3-MeCOC5 H4 N 3-FC5 H4 N 5-MeC4 H3 N2

a

3-MeOOCC5 H4 N

a

2-i-PrC5 H4 N

a

Formula CH2 Cl2 Hept Hex Hept CHCl3 Hept Hept CCl4 CHCl3 CHCl3 Hex Hex CCl4 CH2 Cl2 Hept CCl4 CH2 Cl2 CH2 Cl2 Hex Hept cHex Hept cHex CHCl3 Hex Hept CHCl3 CH2 Cl2 Hept CH2 Cl2 CH2 Cl2 CCl4 (30) (30)

(30)

(30) (30)

(30)

Solvent (t/◦ C) 25 45 44 40 25 33 32 42 10b 21 29 27 17 13.3 24 16.6 10.9 21 12 19 9 9 8 5.9 8 15 5.4 2.8 6.5 2.3 2.5 4.4

Kc

0.81 0.80

1.35c 1.08 0.98c 0.95 0.95 0.62c 0.91c 0.89 0.88c 0.87c 0.85 0.81

1.38

1.47c 1.46 1.43 1.40

1.68 1.65 1.64 1.60 1.56c 1.52 1.51 1.51c

pK BI2

[120] [65] [121] [65] [125] [65] [65] [126] [125] [125] [121] [119] [115] [116] [80] [115] [116] [128] [121] [129] [114] [113] [82, 130] [125] [121] [65] [131] [116] [84] [116] [116] [115]

−9.59 −9.42 −9.36 −9.13 −8.90 −8.68 −8.62 −8.62

−4.62 −4.57

−7.72 −6.16 −5.59 −5.42 −5.42 −3.51 −5.19 −5.08 −5.02 −4.97 −4.85 −4.62

−7.88

−8.39 −8.33 −8.16 −7.99

Ref.

∆G◦

253

c

b

N

N

N

NH

X

6

5

N

N

Me

Me

C

C

C

53

C

2

X'

72

N

3

H

H

22 26 37 40 44 47 51

X 6-MeO 6-Me H 6-Me 6-Cl H H

73

N

N

55

X’ H H H 2-Me H 2-Me 3-Br N

N

N

N

N

83

33

N

N

59

H

N

C14 H12 N2 C12 H7 ClN2

a

C13 H10 N2 C13 H10 N2 C24 H16 N2

a

N

3 2

4

36

N 1

N

87

60 5

N

N

N 10

6

6-Me-2-(PhCH CH)C5 H3 N 3-O2 NC5 H4 N 2-ClC5 H4 N 2-FC5 H4 N

Two non-equivalent nitrogens. Statistically corrected in order to put the pK BI2 values on a per nitrogen basis.

71

39

N

Cl

17

N

6-Methyl-trans-2-styrylpyridine 3-Nitropyridine 2-Chloropyridine 2-Fluoropyridine 1,3,5-Triazine Substituted 1,10-phenanthrolines 4-Methyl-1,10-phenanthroline 5-Methyl-1,10-phenanthroline 4,7-Diphenyl-1,10-phenanthroline 1,10-Phenanthroline 2,9-Dimethyl-1,10-phenanthroline 5-Chloro-1,10-phenanthroline

Formulae:

Cl

a

84 85 86 87 88 89

79 80 81 82 83

9

7 8

37

N

65

N

N

CCl4 CCl4 CCl4 CCl4 CCl4 CCl4

67

N

38

N

N

Hex CH2 Cl2 (30) Hept Hept cHex

N

155 131 111 107 22 15

6 1.9 5 2.7 2 2.39 2.31 2.24 2.22 1.52 1.35

0.78 0.70 0.70 0.43 −0.18c

[121] [116] [80] [65] [77] [123] [123] [123] [123] [123] [123]

−4.45 −4.00 −4.00 −2.46 1.03 −13.64 −13.19 −12.79 −12.67 −8.68 −7.71

254


Table 5.9 Equilibrium constants Kc (l mol−1 ) and diiodine basicity scales pK B I 2 and ∆G◦ (kJ mol−1 ) for the complexes of diiodine with five-membered N-heterocycles (mainly heteroarenes). No.

Lewis basea

Solvent

Kc

pK BI2

∆G◦

Ref.

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28

1-(Adamant-1-yl)imidazole 2-Methyl-2-thiazoline 1-tert-Butylimidazole 1-n-Butylimidazole 1-Methylimidazole 1-n-Butylbenzimidazole 1-Methylbenzimidazole Benzimidazole 1-Methylpyrazole 1,3,5-Trimethylpyrazole 2,4-Dimethylthiazole Pyrazole 1-(Adamant-1-yl)-1,2,4-triazole 1-Methyl-1,2,4-triazole Indazole 5-(3-Chloropropyl)-4-methylthiazole 1-(Adamant-1-yl)pyrazole 5-(2-Chloroethyl)-4-methylthiazole 4-Methylthiazole 1-Benzyl-1,2,4-triazole 2-Methylbenzoxazole Benzothiazole Benzotriazole Thiazole 1,5-Cyclopentamethylenetetrazole 1-Phenylpyrazole N-Methylpyrrole 3,5-Dimethylisoxazole

Hept CCl4 Hept Hept Hept Hept CCl4 CH2 Cl2 Hept CCl4 CCl4 Hept Hept C6 H5 Cl CH2 Cl2 CCl4 Hept CCl4 CCl4 Hept CCl4 CCl4 CH2 Cl2 CCl4 CCl4 Hept Hept CCl4

2561 926 1194 1094 725 612 208 164 138 72 65 78 79 99 16 25 37 20 18 23 14 11 4 9 7.5 10 5.9 1.7

3.41 3.18 3.08 3.04 2.86 2.79 2.52 2.49 2.14 2.05 2.00 1.89

−19.46 −18.15 −17.58 −17.35 −16.32 −15.93 −14.38 −14.21 −12.22 −11.70 −11.42 −10.79

[84] [133] [84] [132] [84] [132] [134, 135] [132] [132] [134, 135] [136] [84] [132] [132] [132] [133] [132] [133] [133] [132] [134, 135] [134, 135] [132] [133] [137] [132] [95] [134, 135]

a

b b

−9.42 −9.02 −8.96 −8.50 −8.11

1.65 1.58 1.57 1.49 1.42 b

1.32 1.22

−7.53 −6.96

1.11

−6.34

1.00 0.77 0.38

−5.71 −4.40 −2.17

c

Formulae: R

R

1 3 4 5

N

N

S

R = Adam R = t-Bu R = n-Bu R = Me

N

Me N

N

R

R = n-Bu R = Me R=H

6 7 8

N N

2

9 12 17 26

R = Me R=H R = Adam R = Ph

Me Me

Me

N

N N

N Me

N

Me

10

H

R

S

N

4

13 R= Adam 14 R = Me 20 R = CH2Ph

S R

N

N

N

Me

11

16 R = ClCH2CH2CH2 18 R= ClCH2CH2 19 R = H

15 H

O

S

N

Me N

21 b c

S

N

N

22

N

N

23

N4 is the main halogen-bond acceptor site. N2 is the main halogen-bond acceptor site.

Me

N

N

2 N

N

24

25

O N

N

Me

Me

27

28

255

Isobutyronitrile

Valeronitrile, pentanenitrile

Trimethylacetonitrile

N,N-Dimethylcyanamide N,N-Diethylcyanamide

N,N,N-Tri-n-butylammoniocyanamidate

13

14

15

16 17

18

CCl4

n-Bu3 N N C ≡ N

+ −

Me2 NC N Et2 NC N

(CH3 )3 CC N

CH3 (CH2 )3 C N

(CH3 )2 CHC N

CH3 CH2 CH2 C N

CH3 C N C2 H5 C N

CCl4 CCl4 CCl4 CCl4 CCl4 CCl4 CCl4 CCl4 CCl4 Hept CCl4 CCl4 Hept CCl4 Hept CCl4 Hept CCl4 Hept CCl4 Hept Hept CCl4 Hept

Solvent

ClCH2 C N CH3 SC N CH3 (CH2 )3 SC N CH3 (CH2 )7 SC N C6 H5 SC N C6 H5 CH2 SC N 3-MeC6 H4 CH2 SC N H2 C CHC N C6 H5 C N

Formula

57

0.19 0.34a 0.38a 0.44a 0.28a 0.37a 0.32a 0.38 0.39 0.73 0.46 0.51 0.95 0.51 0.99 0.56 0.98 0.53 1.05 0.59 1.12 3.29 2.25 5.31

Kc

0.78 0.38 0.13

−0.14 −0.07 −0.02

−13.26

−4.14 0.73 2.32

−0.28 −2.95

−0.12

0.05

0.05 0.52

0.02

−0.01

0.02

2.89 1.39 1.08 0.66 1.94 1.16 1.57 0.92

−0.51 −0.24 −0.19 −0.12 −0.34 −0.20 −0.27 −0.16

0.00

∆G◦

pK BI2

At 20 ◦ C. For the calculation of pK BI2 , these values are corrected to 25 ◦ C (and to heptane), assuming a complexation enthalpy of –8 kJ mol−1 .

Butyronitrile

12

a

Acetonitrile Propionitrile

Chloroacetonitrile Methyl thiocyanate Butyl thiocyanate Octyl thiocyanate Phenyl thiocyanate Benzyl thiocyanate 3-Methylbenzyl thiocyanate Acrylonitrile Benzonitrile

Lewis base

10 11

1 2 3 4 5 6 7 8 9

No.

[65]

[138] [139] [139] [139] [139] [139] [139] [65] [65] [65] [71] [71] [71] [71] [71] [65] [65] [71] [71] [71] [71] [65] [65] [65]

Ref.

Table 5.10 Equilibrium constants Kc (l mol−1 ) and diiodine basicity scales pK B I 2 and ∆G◦ (kJ mol−1 ) for the complexes of diiodine with nitriles.

256 CH3 OCH2 CH2 CH3 CH3 O(CH2 )3 CH3 C2 H5 CH(CH3 )OCH3 (CH3 )2 CHCH2 OCH3 CH3 OC(CH3 )3 (C2 H5 )2 O

Aliphatic ethers Methyl n-propyl ether n-Butyl methyl ether sec-Butyl methyl ether Isobutyl methyl ether tert-Butyl methyl ether Diethyl ether

Ethyl n-propyl ether n-Butyl ethyl ether sec-Butyl ethyl ether tert-Butyl ethyl ether Di-n-propyl ether n-Butyl n-propyl ether tert-Butyl n-propyl ether n-Butyl isopropyl ether Di-n-butyl ether

Diisopropyl ether Isopropyl n-propyl ether Dibenzyl ether Benzyl methyl ether Dimethoxymethane 1,2-Dimethoxyethane 1,1-Dimethoxyethane 1,1,1-Trimethoxymethane Diethylene glycol dimethyl ether

Cyclic and crown ethers Propylene oxide Oxetane Tetrahydrofuran

Tetrahydropyran 1,3-Dioxane

1 2 3 4 5 6

7 8 9 10 11 12 13 14 15

16 17 18 19 20 21 22 23 24

25 26 27

28 29

c

c

c

c

c

[(CH3 )2 CH]2 O CH3 CH2 CH2 OCH(CH3 )2 (C6 H5 CH2 )2 O C6 H5 CH2 OCH3 CH3 OCH2 OCH3 CH3 OCH2 CH2 OCH3 (CH3 O)2 CHCH3 (CH3 O)3 CH (CH3 OCH2 CH2 )2 O

C2 H5 OCH2 CH2 CH3 C2 H5 O(CH2 )3 CH3 C2 H5 CH(CH3 )OC2 H5 C2 H5 OC(CH3 )3 (CH3 CH2 CH2 )2 O CH3 CH2 CH2 O(CH2 )3 CH3 CH3 CH2 CH2 OC(CH3 )3 CH3 CH2 CH2 CH2 OCH(CH3 )2 (CH3 CH2 CH2 CH2 )2 O

Formula

Lewis base

No.

Hept Hept Hept CCl4 Hept CCl4

cHex cHex cHex cHex cHex cHex CCl4 cHex cHex cHex cHex cHex cHex cHex cHex Hept CCl4 cHex cHex CCl4 CCl4 CCl4 CCl4 CCl4 CCl4 CCl4

Solvent

0.94 3.85 2.54 1.4 2.51 0.71

0.78a 0.93a 1.19a 0.85a 1.34a 0.90 0.58 0.87a 0.75a 0.75a 0.61a 0.78a 0.77a 0.29a 0.82a 0.69 0.44 0.70a 0.82a 0.48b 0.51b 0.63 0.95 0.63 0.50 1.18

Kc

1.18 0.79

1.91 0.78 1.91 3.65

−0.21 −0.14 −0.33d −0.14d −0.33d −0.64d

0.40 −0.28d

−2.28 1.58

0.15 −3.34 −2.31

0.64 1.01 1.01 1.53 0.92 0.95 3.37 0.79 0.92

−0.11 −0.18 −0.18 −0.27 −0.16 −0.17 −0.59 −0.14 −0.16

−0.03 0.59 0.40

0.92 0.48 −0.13 0.70 −0.43 0.26

∆G◦

−0.16 −0.08 0.02 −0.12 0.07 −0.05

pK BI2

[69] [69] [69] [141] [69] [142]

[140] [140] [140] [140] [140] [65] [65] [140] [140] [140] [140] [140] [140] [140] [140] [65] [141] [140] [140] [141] [141] [142] [141] [141] [141] [141]

Ref.

Table 5.11 Equilibrium constants Kc (l mol−1 ) and diiodine basicity scales pK B I 2 and ∆G◦ (kJ mol−1 ) for the complexes of diiodine with ethers.

257

1,3,5-Trioxane 2-Methyltetrahydrofuran 7-Oxabicyclo[2.2.1]heptane

12-Crown-4 15-Crown-5 18-Crown-6

Aromatic ethers 1,2-Dimethoxybenzene 1,4-Dimethoxybenzene Anisole 1,4-Benzodioxane Diphenyl ether Ethoxybenzene 1,3-Dimethoxybenzene Diphenylene dioxide (20 ◦ C)

31 32 33

34 35 36

37 38 39 40 41 42 43 44 c

C6 H5 OC6 H5 C6 H5 OC2 H5 3-CH3 OC6 H4 OCH3

c

2-CH3 OC6 H4 OCH3 4-CH3 OC6 H4 OCH3 C6 H5 OCH3

c

c

c

c

c

c

c

CCl4 CCl4 CCl4 CCl4 CCl4 CCl4 CCl4 CCl4

cHex CCl4 CCl4 Hept cHex CCl4 CCl4 CCl4 CCl4 0.61b 0.54b 0.34b 0.33b 0.30b 0.29b 0.285b 0.27b

1.56 1.00 0.35 2.98 3.36 1.89 2.3 2.18 2.8

d

b

a

O

26

O

O

O

27

35

O

O

O

28

O

O

O

O

O

36

29

O

O

O

O

O

30

O

Statistically corrected in order to put the pK BI2 value on a per oxygen basis.

O

O

34

O

O

25

O

31

40

O

O

O

O

O

32

O

44

O

O

33

O

Values at 20 ◦ C. For the calculation of pK BI2 at 25 ◦ C, a value of −17.6 kJ mol−1 is assumed for the enthalpy of complexation. These values include the contribution of the complexation constant of diiodine with the π system(s). c Formulae:

1,4-Dioxane

30

0.62 4.63 −2.71 −3.00 0.26 1.01 0.83

−0.11d −0. 81d 0.47 0.53 −0.04d −0.18d −0.15d

[141] [141] [141] [141] [141] [141] [141] [143]

[65] [65] [141] [69] [65] [65] [141] [141] [141]

258

Ethyl butanoate δ-Valerolactone ε-Caprolactone Thioester Ethyl thioacetate

19 20 21

22

γ -Butyrolactone Methyl butanoate Ethyl acetate

Acid halides Benzoyl bromide Benzoyl chloride Acetyl chloride Acetyl bromide Aldehyde Benzaldehyde Esters, lactones β-Propiolactone Ethyl trifluoroacetate Ethyl trichloroacetate Ethyl formate Dimethyl carbonate Methyl benzoate Diethyl carbonate Ethyl methyl carbonate Ethyl benzoate Methyl acetate

Lewis base

16 17 18

6 7 8 9 10 11 12 13 14 15

5

1 2 3 4

No.

MeCOSEt

a

n-PrCOOEt

a

n-PrCOOMe MeCOOEt

a

Hept

CCl4 Hept Hept Hept Hept Hept Hept Hept Hept Hept CCl4 CCl4 CCl4 Hept CCl4 CCl4 CCl4 CCl4

a

CF3 COOEt CCl3 COOEt HCOOEt MeOCOOMe PhCOOMe EtOCOOEt MeOCOOEt PhCOOEt MeCOOMe

Hept

Hept Hept Hept Hept

Solvent

PhCOH

PhCOBr PhCOCl MeCOCl MeCOBr

Formula

0.86b

0.08 0.1 0.24 0.49 0.56 0.61 0.62 0.62 0.71 0.76 0.52 0.52 0.60 0.94 0.53 0.70 0.95 1.01

0.53

0.1 0.1 0.12 0.13

Kc

5.83 5.71 3.54 1.77 1.44 1.23 1.19 1.19 0.85 0.68

−1.02 −1.00 −0.62 −0.31 −0.25 −0.21 −0.21 −0.21 −0.15 −0.12

0.04 0.19 0.22

−0.22 −1.07 −1.24

0.61 0.21 0.15

1.57

−0.28

−0.11 −0.04 −0.03

5.71 5.71 5.26 5.06

∆G◦

−1.00 −1.00 −0.92 −0.89

pK BI2

[70]

[43] [70] [70] [70] [70] [70] [70] [65] [70] [70] [144] [43] [43] [70] [65] [43] [43] [43]

[70]

[70] [70] [70] [70]

Ref.

Table 5.12 Equilibrium constants Kc (l mol−1 ) and diiodine basicity scales pK B I 2 and ∆G◦ (kJ mol−1 ) for the complexes of diiodine with carbonyl bases.

259

Ketones α,α,α-Trifluoroacetophenone 4-Trifluoromethylacetophenone 3-Trifluoromethylacetophenone 1-Acetylnaphthalene Isobutyrophenone 3-Chloroacetophenone 3-Bromoacetophenone 4-Bromoacetophenone Propiophenone 2-Acetylnaphthalene Benzophenone 4-Chloroacetophenone

4-Fluoroacetophenone Cyclobutanone

Ethyl methyl ketone Isopropyl methyl ketone Acetone

tert-Butyl methyl ketone Acetophenone

3-Methoxyacetophenone 3-Methylacetophenone 4-Methylacetophenone

23 24 25 26 27 28 29 30 31 32 33 34

35 36

37 38 39

40 41

42 43 44

3-MeOC6 H4 COMe 3-MeC6 H4 COMe 4-MeC6 H4 COMe

t-BuCOMe PhCOMe

EtCOMe i-PrCOMe MeCOMe

4-FC6 H4 COMe c-(CH2 )3 CO

PhCOCF3 4-CF3 C6 H4 COMe 3-CF3 C6 H4 COMe 1-Naphth COMe PhCO-i-Pr 3-ClC6 H4 COMe 3-BrC6 H4 COMe 4-BrC6 H4 COMe PhCOEt 2-Naphth-COMe PhCOPh 4-ClC6 H4 COMe Hept Hept Hept Hept Hept Hept Hept cHex Hept Hept Hept Hept CCl4 Hept Hept CCl4 Hept Hept Hept CCl4 Hept Hept CCl4 Hept Hept Hept CCl4

0.1 0.57 0.58 0.65 0.73 0.74 0.76 0.78 0.79 0.85 0.86 0.86 0.53 0.93 0.95 0.44 1.03 1.08 1.12 0.85 1.14 1.15 0.92 1.19 1.29 1.44 1.14 0.08 0.11 0.16

0.06 0.06

−0.43 −0.63 −0.90

−0.32 −0.35

−0.07 −0.19 −0.28

0.18 0.13

−0.03 −0.02 0.01 0.03 0.05

5.71 1.39 1.35 1.07 0.78 0.75 0.68 0.62 0.58 0.40 0.37 0.37

−1.00 −0.24 −0.24 −0.19 −0.14 −0.13 −0.12 −0.11 −0.10 −0.07 −0.07 −0.07

(Continued)

[70] [67] [70] [70] [70] [67] [67] [68] [70] [70] [70] [70] [145] [67] [65] [43] [70] [70] [70] [144] [70] [70] [145] [67] [70] [70] [145]

260

Cyclopentanone

4-Methoxyacetophenone Cycloheptanone Cyclohexanone 4H-Pyran-4-one 4H-Thiopyran-4-one Amides, lactams N,N-Dimethylcyanoformamide N,N-Dimethyltrichloroacetamide

N,N-Dimethyltrifluoroacetamide N,N-Dimethylchloroacetamide N,N-Dimethylphenylcarbamate 1-(3 -Chlorophenyl)-2-pyrrolidone 3,4-Dichloro-N,N-dimethylbenzamide 1-(4 -Chlorophenyl)-2-pyrrolidone N,N-Diisopropylpivalamide N,N-Dicyclohexylpivalamide N,N-Dimethylmethylcarbamate 2,4-Dichloro-N,N-dimethylbenzamide N,N-Dicyclohexylisobutyramide 1-(3 -Methylphenyl)-2-pyrrolidone 1-Phenyl-2-pyrrolidone 1-(3 -Methoxylphenyl)-2-pyrrolidone Azetidin-2-one 1-(4 -Ethylphenyl)-2-pyrrolidone 4-Chloro-N,N-dimethylbenzamide 1-(4 -Methylphenyl)-2-pyrrolidone 1-(2 -Chlorophenyl)-2-pyrrolidone

45

46 47 48 49 50

53 54 55 56 57 58 59 60 61 62 63 64 65 66 67 68 69 70 71

51 52

Lewis base

No.


a

a

4-ClC6 H4 CONMe2

a

a

a

a

a

t-BuCON(i-Pr)2 t-BuCON(c-Hex)2 MeOCONMe2 2,4-Cl2 C6 H3 CONMe2 i-PrCON(c-Hex)2

a

3,4-Cl2 C6 H3 CONMe2

a

CF3 CONMe2 ClCH2 CONMe2 PhOCONMe2

N CCONMe2 CCl3 CONMe2

a

a

4-MeOC6 H4 COMe c-(CH2 )6 CO c-(CH2 )5 CO

c-(CH2 )4 CO

Formula

CCl4 Hept CCl4 Hept CCl4 Hept CCl4 CCl4 CCl4 Hept Hept Hept CCl4 Hept CCl4 CCl4 CCl4 CCl4 CCl4 CCl4 CCl4 CCl4

Hept CCl4 Hept CCl4 CCl4 CCl4 CCl4

Solvent

0.26b 0.45 0.3 0.61 1.3 2.2 1.5 1.7 1.7 3 3.5 3.8 2.2 4 2.3 2.4 2.4 2.50 2.7 2.8 2.9 2.9

1.68 0.94 1.94 1.43 1.47 3.5 4.6

Kc

1.98 1.23 −1.95 −1.95 −2.34 −2.69 −2.69 −2.72 −3.11 −3.31 −3.41 −3.44 −3.54 −3.66 −3.66 −3.77 −3.98 −4.09 −4.18 −4.18

−0.21 0.34 0.34 0.41 0.47 0.47 0.48 0.54 0.58 0.60 0.60 0.62 0.64 0.64 0.66 0.70 0.72 0.73 0.73

[147] [65] [148] [65] [148] [147] [149] [150] [149] [151] [151] [147] [150] [151] [149] [149] [149] [44] [149] [150] [149] [149]

[65] [43] [67] [43] [43] [146] [146]

−1.29 −1.64 −2.21 −2.29 −4.71 −5.47

Ref.

∆G◦

−0.35

0.29 0.39 0.40 0.82 0.96

0.23

pK BI2

261

1-(4 -Methoxyphenyl)-2-pyrrolidone 2-Chloro-N,N-dimethylbenzamide N,N-Dimethylformamide

N,N-Diethylpivalamide N,N-Dimethylpivalamide N,N-Dimethylbenzamide

2-Methyl-N,N-dimethylbenzamide 4-Methyl-N,N-dimethylbenzamide N,N-Dicyclohexylpropionamide 4-Methoxy-N,N-dimethylbenzamide 1-(2 -Methylphenyl)-2-pyrrolidone 2-Methoxy-N,N-dimethylbenzamide N,N-Dimethylacetamide

Pyrrolidin-2-one N,N-Diethylacetamide

1-Methylpyrrolidin-2-one

N,N-Dicyclohexylacetamide 1-Methyl-ε-caprolactam

δ-Valerolactam, 2-piperidone ε-Caprolactam 1-Methyl-2-piperidone

72 73 74

75 76 77

78 79 80 81 82 83 84

85 86

87

88 89

90 91 92

a

a

a

a

a

MeCON(c-Hex)2

a

MeCONEt2

a

2-MeOC6 H4 CONMe2 MeCONMe2

2-MeC6 H4 CONMe2 4-MeC6 H4 CONMe2 EtCON(c-Hex)2 4-MeOC6 H4 CONMe2

t-BuCONEt2 t-BuCONMe2 PhCONMe2

2-ClC6 H4 CONMe2 HCONMe2

a

CCl4 CCl4 Hept CCl4 Hept Hept Hept CCl4 CCl4 CCl4 Hept CCl4 CCl4 CCl4 Hept CCl4 CCl4 Hept CCl4 Hept CCl4 Hept Hept CCl4 CCl4 CCl4 Hept

3.2 3.3 6.45 2.9 6.65 6.7 8.1 3.8 4.8 4.8 10 5.3 5.5 7.1 15 6.8 7.5 18.8 7.6 19.9 10.2 20 22.4 12.8 10.4 10.77 29.7 1.36 1.37 1.47

1.30 1.35

1.30

1.20 1.27

0.98 0.98 1.00 1.03 1.05 1.17 1.18

0.82 0.83 0.91

0.78 0.80 0.81

−7.74 −7.84 −8.41

−7.43 −7.71

−7.41

−6.83 −7.27

−5.59 −5.59 −5.71 −5.86 −5.97 −6.68 −6.71

−4.70 −4.72 −5.19

−4.46 −4.54 −4.62

(Continued)

[149] [150] [70] [148] [147] [147] [70] [152] [150] [150] [151] [150] [149] [150] [70] [153] [44] [70] [154] [155, 156] [44] [151] [155] [44] [44] [44] [155]

262

b

a

R

N

O

R: 85 87

Me

Polyfunctional base.

O

N

O

6

O

Formulae:

R

H Me

R:

16

O

O

56 58 64 65

O

N

R

20

3’-Cl 4’-Cl 3’-Me H

O

Me H

72 82

21

O

3’-MeO 4’-Et 4’-Me 2’-Cl

R: 89 91

66 68 70 71

O

O

N,N-Dimethyl-N ,N -diethylurea N,N -Dimethyl-N,N -propyleneurea (DMPU)

95 96

O

N

4’-MeO 2’-Me

R

49

O

O

R: 90 92

O

H Me

67

NH

50

S

O

Me2 NCONEt2

Et2 NCONEt2 Me2 NCONMe2

Ureas Tetraethylurea Tetramethylurea

93 94 a

Formula

Lewis base

No.


Me

N

96

O

N

Hept Hept CCl4 Hept cHex CCl4

Me

Solvent 11 14.3 6.5 14.3 39.6 16.7

Kc

1.16 1.60

1.04 1.16

pK BI2

−6.59 −9.12

−5.94 −6.59

∆G◦

[70] [70] [144] [147] [65] [65]

Ref.

263

a

MeSONMe2 PhSOMe Ph2 SO PhSOOEt c-(CH2 )4 SO2 (EtO)2 SO (MeO)2 SO MeOSOCl Cl2 SO

Et2 SO c-(CH2 )4 SO Me2 SO

n-Bu2 SO

Formula

Statistically corrected in order to put the pK BI2 value on a per oxygen basis.

N,N-Dimethylmethanesulfinamide Methyl phenyl sulfoxide Diphenyl sulfoxide Ethyl benzenesulfinate Sulfolane Diethyl sulfite Dimethyl sulfite Methyl chlorosulfinate Thionyl chloride

Diethyl sulfoxide Tetramethylene sulfoxide Dimethyl sulfoxide

2 3 4

5 6 7 8 9 10 11 12 13

Di-n-butyl sulfoxide

Lewis base

1

No. cHex CCl4 CCl4 CCl4 cHex CCl4 CCl4 CCl4 CCl4 CCl4 CCl4 CCl4 CCl4 CCl4 CCl4

Solvent 84.8 23.7 17.3 15.5 35.9 11.1 10.6 5.7 4.4 1.20 0.73 0.34 0.31 0.17 0.10

Kc

1.53 1.23 1.10 0.46 −0.12a −0.16 −0.20 −0.50 −0.76

1.77 1.72 1.56

1.93

pK BI2

[65] [65] [157] [157, 158] [65] [65] [159] [160] [160, 161] [161] [158] [158] [161] [161] [161]

−11.01

−8.75 −7.01 −6.29 −2.65 0.69 0.89 1.15 2.83 4.32

−10.12 −9.81 −8.88

Ref.

∆G◦

Table 5.13 Equilibrium constants Kc (l mol−1 ) and diiodine basicity scales pK B I 2 and ∆G◦ (kJ mol−1 ) for the complexes of diiodine with thionyl compounds.

264

Triethyl phosphate Trimethyl phosphate Diethyl chlorophosphonate Diethyl chloromethylphosphonate Methylphosphonic dichloride Diethyl dichloromethylphosphonate Dichlorophosphoric dimethylamide Diethylphosphoramidic dichloride Ethyl dichlorophosphinate Diphenyl chlorophosphonate Phosphoryl chloride

13 14 15 16 17 18 19 20 21 22 23

Hexamethylphosphoric triamide Pentaethylphosphonic diamide Triphenylphosphine oxide Ethyl tetraethylphosphorodiamidate Diethyl methylphosphonate

5 6 7 8 9

Bis(dimethylamino)phosphorochloridate Diethyl N,N-diethylphosphoramidate Tri-n-butyl phosphate

Tri-n-octylphosphine oxide

4

10 11 12

Trimethylphosphine oxide Triethylphosphine oxide Tri-n-butylphosphine oxide

Lewis base

1 2 3

No.

(EtO)3 PO (MeO)3 PO (EtO)2 POCl (EtO)2 POCH2 Cl MePOCl2 (EtO)2 POCHCl2 Me2 NPOCl2 Et2 NPOCl2 Cl2 POOEt (PhO)2 POCl POCl3

(Me2 N)2 POCl (EtO)2 PONEt2 (n-BuO)3 PO

(Me2 N)3 PO (Et2 N)2 POEt Ph3 PO (Et2 N)2 POOEt (EtO)2 POMe

n-Oct3 PO

Me3 PO Et3 PO n-Bu3 PO

Formula CCl4 CCl4 cHex CCl4 cHex CCl4 CCl4 CCl4 CCl4 CCl4 cHex CCl4 CCl4 CCl4 cHex CCl4 CCl4 CCl4 CCl4 CCl4 CCl4 CCl4 CCl4 CCl4 CCl4 CCl4 CCl4

Solvent 101 108 561 105 592 111 71 65.6 28.1 20.8 36.8 10.2 6.1 6.7 21.6 6.4 4.5 3.0 1.96 2.55 1.68 1.25 0.80 0.52 0.31 0.35 0.10

Kc

1.16 0.95 0.73 0.87 0.66 0.51 0.28 0.06 −0.20 −0.13 −0.77

1.31 1.35 1.33

2.55 2.51 2.08 1.93 1.57

2.77

2.72 2.76 2.75

pK BI2

[161] [162] [65] [65] [65] [65] [161] [163] [59, 108, 161, 163] [161] [65] [65] [161] [163] [65] [65] [161, 163] [162] [164] [163] [164] [163] [164] [161] [161] [161] [161]

−15.55 −15.74 −15.69

−6.60 −5.42 −4.19 −4.95 −3.75 −2.90 −1.61 −0.37 1.12 0.77 4.38

−7.46 −7.73 −7.62

−14.53 −14.31 −11.87 −11.00 −8.94

−15.82

Ref.

∆G◦

Table 5.14 Equilibrium constants Kc (l mol−1 ) and diiodine basicity scales pK B I 2 and ∆G◦ (kJ mol−1 ) for the complexes of diiodine with phosphoryl compounds.

265

N

17

NO

N

19

NO

20

N NO

Ethyl isonicotinate N-oxide Pyrimidine mono-N-oxide 4-Acetylpyridine N-oxide Pyrazine mono-N-oxide Pyridazine mono-N-oxide 4-Nitropyridine N-oxide

16 17 18 19 20 21

Formulae:

3-Methylpyridine N-oxide 2-Methylpyridine N-oxide 2,6-Dimethylpyridine N-oxide 2,4,6-Trimethylpyridine N-oxide 4-Chloropyridine N-oxide

11 12 13 14 15

a

4-tert-Butylpyridine N-oxide 4-Ethylpyridine N-oxide 4-Methylpyridine N-oxide

Tri-n-butylamine N-oxide Trimethylamine N-oxide Tribenzylamine N-oxide 4-Dimethylaminopyridine N-oxide 4-Ethoxypyridine N-oxide 4-Methoxypyridine N-oxide

2 3 4 5 6 7

8 9 10

Pyridine N-oxide

Lewis base

1

No.

4-O2 NC5 H4 NO

a

a

4-MeCOC5 H4 NO

4-EtOOCC5 H4 NO

a

3-MeC5 H4 NO 2-MeC5 H4 NO 2,6-Me2 C5 H3 NO 2,4,6-Me3 C5 H2 NO 4-ClC5 H4 NO

4-t-BuC5 H4 NO 4-EtC5 H4 NO 4-MeC5 H4 NO

n-Bu3 NO Me3 NO (PhCH2 )3 NO 4-Me2 NC5 H4 NO 4-EtOC5 H4 NO 4-MeOC5 H4 NO

C5 H5 NO

Formula cHex CCl4 C6 H6 CH2 Cl2 C6 H6 CH2 Cl2 CH2 Cl2 CH2 Cl2 CH2 Cl2 CCl4 CH2 Cl2 CH2 Cl2 CH2 Cl2 CCl4 C6 H6 CH2 Cl2 CCl4 CCl4 CCl4 CCl4 CCl4 CH2 Cl2 CH2 Cl2 CH2 Cl2 CH2 Cl2 CH2 Cl2 CH2 Cl2 CCl4 CH2 Cl2

Solvent 250 73.5 83 38.8 20 000 4770 2570 1500 260 1403 180 44 99 206 316 103 135 114 156 560 31 13.0 10.2 5.95 5.18 3.55 1.53 0.7 0.9

Kc 2.40

pK BI2

Ref. [65] [65, 165, 166] [162] [167, 168] [162] [169] [169] [167] [167] [170] [167] [167] [167] [165, 166, 170] [162] [167] [165, 166] [165, 166] [165] [171] [165, 170] [167] [167] [168] [167] [168] [168] [165, 170] [167, 170]

∆G◦ −13.69

Table 5.15 Equilibrium constants Kc (l mol−1 ) and diiodine basicity scales pK B I 2 and ∆G◦ (kJ mol−1 ) for the complexes of diiodine with amine N-oxides and pyridine N-oxides in various solvents.

266

2,4,6-Trimethylbenzonitrile N-oxide N-Benzylidenemethylamine N-oxide Dimethyl selenoxide Diphenyl selenoxide

Ethyl benzeneseleninate Diethyl selenite Dimethyl selenite Triphenylarsine oxide

Tricyclohexylarsine oxide Tri-n-octylarsine oxide Triphenylstibine oxide

5 6 7 8

9 10 11

Lewis base

1 2 3 4

No.

(c-Hex)3 AsO (n-Oct)3 AsO Ph3 SbO

PhSeOOEt (EtO)2 SeO (MeO)2 SeO Ph3 AsO

Me3 C6 H2 C NO C6 H5 CH N(Me) O Me2 SeO Ph2 SeO

Formula CCl4 CCl4 CCl4 CCl4 C6 H6 CH2 Cl2 CCl4 CCl4 CCl4 CCl4 C6 H6 CH2 Cl2 CHCl3 CCl4 CCl4 CH2 Cl2

Solvent

1.78 25.5 281 92 63 27 14.5 0.90 0.77 1171 980 218 41 15 700 12 000 3960

Kc

Table 5.16 Equilibrium constants Kc (l mol−1 ) for the complexes of diiodine with miscellaneous oxygen bases in various solvents.

[172] [169] [161] [161] [162] [173] [161] [161] [161] [59] [162] [174] [175] [176] [176] [177]

Ref.

267

Thiobenzamide Thioacetamide N,N-Dimethylthioacetamide

1-Thiobenzoylpiperidine N,N-Dimethylthioformamide

13 14 15

16 17

8 9 10 11 12

6 7

2,2,4,4-Tetramethylpentan-3-thione S-Methyl O-methyldithiocarbonate Acyclic thioamides Thiocinnamamide Cinnamothiopiperidide Thiophenylacetamide Phenylthioacetopiperidide N-Methyl-ε-thiocaprolactam

4,5-Ethylenedithio-1,3-dithiole-2thione Ethylenetrithiocarbonate 1,3-Dithianecyclohexane-2-thione Methyl dithiovalerate

2

3 4 5

Thiocamphor

Thioketones, thioesters

Lewis base

1

No.

C6 H5 CSN(CH2 )5 HCSN(CH3 )2

C6 H5 CSNH2 CH3 CSNH2 CH3 CSN(CH3 )2

a

C6 H5 CH CHCSNH2 C6 H5 CH CHCSN(CH2 )5 C6 H5 CH2 CSNH2 C6 H5 CH2 CSN(CH2 )5

(CH3 )3 CCSC(CH3 )3 CH3 OCSSCH3

C4 H9 CSSCH3

a

CCl4 CCl4 CCl4 CCl4 Hept CCl4 CH2 Cl2 CCl4 CH2 Cl2 Hept CCl4 CCl4 Hept

cHex CHCl3 Hept CCl4 CHCl3 CH2 Cl2 Hept Hept

a

a

Hept CCl4 CHCl3 CH2 Cl2 CHCl3

Solvent

a

Formula

7378 6324 5853 4959 3844 3020 12 880 2622 8193 1790 1190 1163 1380

45 33 13.2 12.6 17.7 24 4.8 3

141 152 231 370 87

Kc

3.16 3.14

3.53 3.38 3.25

4.00 3.93 3.90 3.82 3.58

0.68 0.48

1.65 1.36 1.12

1.78

2.15

pK BI2

−18.05 −17.92

−19.32 −18.57

−22.85 −22.45 −22.25 −21.81 −20.46

−3.89 −2.72

−9.44 −7.76 −6.40

−10.17

−12.27

∆G◦

Ref.

(Continued)

[181] [181] [181] [181] [155] [155] [155] [181] [182] [41] [183] [181] [184]

[180] [179] [178] [178] [178] [178] [41] [41]

[178] [178] [178] [178] [179]

Table 5.17 Equilibrium constants Kc (l mol−1 ) and diiodine basicity scales pK B I 2 and ∆G◦ (kJ mol−1 ) for the complexes of diiodine with thiocarbonyl bases.

268

N,N-Dimethylthioisobutyramide N,N,2,2Tetramethylthiopropionamide N,N-Diethyl-2,2dimethylthiopropionamide O-Methyl dimethylthiocarbamate Tolnaftate N,N-Dimethylthiocarbamoyl chloride Cyclic thioamides Pyrrolidine-2-thione 1,3-Thiazolidine-2-thione 1-Methyl-1,3-thiazolidine-2-thione 1,3-Oxazolidine-2-thione Benzoxazole-2-thione 3-Methylbenzoxazole-2-thione Benzothiazole-2-thione 3-Methylbenzothiazole-2-thione 5,5-Dimethyl-4-thiohydanthoin 3,5,5-Trimethyl-4-thiohydanthoin Thiuram disulfides Tetraethylthiuram disulfide Bis(morpholinothiocarbonyl) disulfide Tetrabenzylthiuram disulfide Dithioxamides Tetraethyl dithioxamide Tetramethyl dithioxamide Bis(morpholinothiocarbonyl)

18 19

37 38 39

36

34 35

24 25 26 27 28 29 30 31 32 33

21 22 23

20

Lewis base

No.


CH2 Cl2 CCl4 CCl4 CH2 Cl2 CH2 Cl2 CH2 Cl2 CH2 Cl2 CH2 Cl2 CH2 Cl2 CH2 Cl2 CCl4 CH2 Cl2 CH2 Cl2 CHCl3 CHCl3 CHCl3

a

a

a

a

a

a

a

a

a

a

a

a

a

a

a

a

ClCSN(CH3 )2

Hept CCl4 Hept

560 501 115

61.3

237 91.3

40 050 1821 181 705 59.1 57 567 626 69 12

155 47 14.5

372

1020 4790 427 415

CCl4 CH2 Cl2 Hept Hept Hept

Kc

Solvent

CH3 OCSN(CH3 )2

a

(CH3 )3 CCSN(C2 H5 )2

(CH3 )2 CHCSN(CH3 )2 (CH3 )3 CCSN(CH3 )2

Formula

2.29b 2.24b 1.60b

1.20b

2.12b 1.35b

4.01 3.37 2.32 2.42 1.45 1.44 2.34 2.38 1.51 0.83

2.19 1.70 1.16

2.57

2.63 2.62

pK BI2

[41] [41] [185] [41]

[186] [133] [186] [186] [187] [187] [187] [187] [188] [188] [185] [189] [189] [190] [190] [190]

−15.01 −14.94 −14.67 −12.50 −9.71 −6.63 −22.87 −19.21 −13.21 −13.84 −8.29 −8.21 −13.35 −13.57 −8.64 −4.73 −12.11 −7.72 −6.83 −13.06 −12.79 −9.14

Ref. [184] [184] [41] [41]

∆G◦

269

59 60 61 62 63 64 65 66 67 68 69

55 56 57 58

40 41 42 43 44 45 46 47 48 49 50 51 52 53 54

4-4 -Dimethoxythiocarbanilide 1-Methyl-2-thiourea 4,4 -Dimethylthiocarbanilide 1-Phenyl-1,3,3-triisopropyl-2thiourea 1-Benzyl-3-methyl-2-thiourea 1-Benzyl-1,3,3-tri-n-butyl-2-thiourea 1-Phenyl-1,3,3-triethyl-2-thiourea 1-Phenyl-1,3,3-trimethyl-2-thiourea 1,1-Diethyl-3-phenyl-2-thiourea 1-Methyl-3-phenyl-2-thiourea Thiocarbanilide 1-Phenyl-2-thiourea Thiourea 4,4 -Dibromothiocarbanilide 1-(4-Tolyl)-2-thiourea

Acyclic thioureas 1,3-Di-tert-butyl-2-thiourea 1,3-Dicyclohexyl-2-thiourea 1,3-Di-sec-butyl-2-thiourea 1,3-Di-n-hexyl-2-thiourea 1,3-Diisobutyl-2-thiourea 1,3-Di-n-butyl-2-thiourea 1,3-Di-n-propyl-2-thiourea 1-tert-Butyl-3-methyl-2-thiourea 1,1,3,3-Tetraethyl-3-thiourea 1,1-Diethyl-3,3-dimethyl-2-thiourea 1-Isopropyl-3-methyl-2-thiourea 1,3-Diisopropyl-2-thiourea 1,3-Diethyl-2-thiourea 1,3-Dimethyl-2-thiourea 1,1,3,3-Tetramethyl-2-thiourea

C6 H5 CH2 NHCSNHMe C6 H5 CH2 (n-Bu)NCSN(n-Bu)2 C6 H5 (Et)NCSNEt2 C6 H5 (Me)NCSNMe2 Et2 NCSNHC6 H5 MeNHCSNHC6 H5 C6 H5 NHCSNHC6 H5 C6 H5 NHCSNH2 H2 NCSNH2 BrC6 H4 NHCSNHC6 H4 Br 4-MeC6 H4 NHCSNH2

MeOC6 H4 NHCSNHC6 H4 OMe MeNHCSNH2 MeC6 H4 NHCSNHC6 H4 Me C6 H5 (i-Pr)NCSN-i-Pr2

t-BuNHCSNHt-Bu c-HexNHCSNHc-Hex s-BuNHCSNH-s-Bu n-HexNHCSNH-n-Hex i-BuNHCSNH-i-Bu n-BuNHCSNH-n-Bu n-PrNHCSNH-n-Pr t-BuNHCSNHMe Et2 NCSNEt2 Me2 NCSNEt2 i-PrNHCSNHMe i-PrNHCSNH-i-Pr EtNHCSNHEt MeNHCSNHMe Me2 NCSNMe2

CHCl3 CH2 Cl2 CH2 Cl2 CH2 Cl2 CH2 Cl2 CH2 Cl2 CHCl3 CH2 Cl2 CH2 Cl2 CHCl3 CCl4

CHCl3 CH2 Cl2 CH2 Cl2 CH2 Cl2 CH2 Cl2 CH2 Cl2 CH2 Cl2 CHCl3 Hept Hept CHCl3 CHCl3 CHCl3 CHCl3 Hept CCl4 CHCl3 CH2 Cl2 CHCl3 CHCl3 CHCl3 CHCl3 10 350 26 062 24 245 23 495 20 535 19 059 6570 7649 6700 2320 1187

40 860 118 417 91 541 85 661 78 138 74 682 71 589 20 820 13 570 12 940 17 130 16 430 15 520 14 960 9946 8000 13 560 37 816 14 010 13 100 11 620 27 700 3.86 3.84 3.81 3.80 3.74 3.72 3.66 3.36 3.31 3.21 3.17

3.99 3.96 3.91 3.86

4.46 4.43 4.33 4.30 4.27 4.25 4.23 4.16 4.13 4.11 4.08 4.06 4.04 4.02 4.00

(Continued)

[191] [194] [194] [194] [194] [194] [191] [194] [182] [191] [195]

−22.03 −21.91 −21.75 −21.68 −21.38 −21.21 −20.90 −19.17 −18.87 −18.32 −18.10

−22.78 −22.62 −22.32 −22.04

[191] [192] [192] [192] [192] [192] [192] [191] [41] [41] [191] [191] [191] [191] [41, 193] [183] [191] [193] [191] [191] [191] [194]

−25.44 −25.29 −24.72 −24.57 −24.36 −24.36 −24.17 −23.77 −23.59 −23.47 −23.28 −23.18 −23.04 −22.95 −22.82

270

91

90

86 87 88 89

83 84 85

82

78 79 80 81

77

70 71 72 73 74 75 76

No.

Cyclic thioureas Imidazoline-2-thione 1-Methylimidazoline-2-thione 1,3-Dimethylimidazoline-2-thione 4,5-Dimethylimidazoline-2-thione 1,4,5-Trimethylimidazoline-2-thione 4,5-Diphenylimidazoline-2-thione 1,1 -Methylenebis(3-methyl-4imidazoline-2-thione) 1,1 -Ethylenebis(3-methyl-4imidazoline-2-thione) 1,3-Imidazolidine-2-thione 1-Methyl-1,3-imidazolidine-2-thione 1-Ethyl-1,3-imidazolidine-2-thione 1,3-Dimethyl-1,3-imidazolidine-2thione 1,3-Diethyl-1,3-imidazolidine-2thione Benzimidazole-2-thione 1-Methylbenzimidazole-2-thione 1,3-Dimethylbenzimidazole-2thione 4-Methyl-1,2,4-triazole-5-thione 1-Methyl-1,2,4-triazole-5-thione 1,4-Dimethyl-1,2,4-triazole-5-thione ∆3 -4-S-Methyl-5,5dimethylimidazolidine-2-thione 2-S-Methyl-5,5dimethylimidazolidine-4-thione 4,5,6,7-Tetrathiocino[1,2-b:3,4b ]diimidazolyl-1,3,8,10tetrabutyl-2,9-dithione

Lewis base


CHCl3 CH2 Cl2 CH2 Cl2 CH2 Cl2 CH2 Cl2 CH2 Cl2 CH2 Cl2 CH2 Cl2 CH2 Cl2 CH2 Cl2 CH2 Cl2 CH2 Cl2 CH2 Cl2 CH2 Cl2 CHCl3

a a

a a

a

a a

a

a

a

a

a

a

a

a

a

a

a

a

a

a

CH2 Cl2 CH2 Cl2 CH2 Cl2 CH2 Cl2 CH2 Cl2 CH2 Cl2 CHCl3

Solvent

a

Formula

2240

640c

1587c 1861c 2017c 11 000c

4570 10 800 4740

8520

50 600 49 500 82 750 8100

22 387

49 480 84 730 106 905 73 525 112 453 12 635 6309

Kc

2.89b

3.16 3.49 3.17

3.40

4.10 4.09 4.29 3.38

3.89b

4.09 4.30 4.39 4.24 4.41 3.55 3.34b

pK BI2

[198] [198] [198] [198] [198] [187] [187] [187]

−23.39 −23.34 −24.49 −19.30 −19.41 −18.02 −19.94 −18.10

−16.51

[197]

−22.22

[201]

[200]

[199] [199] [199] [200]

[196] [196] [196] [196] [196] [196] [197]

Ref.

−23.36 −24.54 −25.06 −24.23 −25.18 −20.29 −19.08

∆G◦

271

X

R2

N

S

1

S

C

N

S

Formulae:

R1

a

102 103

94 95 96 97 98 99 100 101

93

92

S

R

S

24 25 26 27

S

S

S

C

N

X CH2 S S O

2

S

S

R2

R1

R H H Me H

S

3

S

S

R1 34 Et 36 PhCH2

S

X

N S

4

S

S

R2 Et PhCH2

R

S

28 29 30 31

X O O S S

4,5,6,7-Tetrathiocino[1,2-b:3,4b ]diimidazolyl-1,3,8,10tetraethyl-2,9-dithione 4,5,6,7-Tetrathiocino[1,2-b:3,4b ]diimidazolyl-1,3,8,10tetraphenyl-2,9-dithione 5,5-Dimethyl-2-thiohydantoin 3,5,5-Trimethyl-2-thiohydantoin 1,5,5-Trimethyl-2-thiohydantoin 1,3,5,5-Tetramethyl-2-thiohydantoin 5,5-Dimethyl-2,4-dithiohydantoin 3,5,5-Trimethyl-2,4-dithiohydantoin 1,5,5-Trimethyl-2,4-dithiohydantoin 1,3,5,5-Tetramethyl-2,4dithiohydantoin Carbimazole 6-Propyl-2-thiouracil

N

12

Me

O

R H Me H Me

S

a

N

H

Me

5

S

C

N1

Me

Me

O

2

S

S

35

3N

4

S

S

R

C

S

C

22

O

CH2 Cl2 CH2 Cl2

a

a

a

a

a

a

a

N

CH2 Cl2 CH2 Cl2 CH2 Cl2 CH2 Cl2 CH2 Cl2 CH2 Cl2 CH2 Cl2 CH2 Cl2

a

Me

CHCl3

a

a

CHCl3

a

N

R 32 H 33 Me

O

2686 52

201 172 149 30.9 143d 126d 131d 19.1d

371

598

2.95 1.40

1.93 1.87 1.82 1.20

2.11b

2.32b [203] [85] [188] [188] [200] [85] [188] [200] [200] [204] [108]

−12.04 −11.03 −10.68 −10.36 −6.84

−16.83 −8.01

(Continued)

[202]

−13.23

272

N

N

S

C

N

MeS

N

N

N

H

R1

N

102

S

R2

N

90

S

S

N

S

C

R

R

COOEt

Me

Me

S

(CH2)n

N

N

Me

S

R

H

Pr

N

N

O

S

S

N

N

R

S

R2 H Me Me

H

O

N

R

S

N

76 n = 1 77 n = 2

103

S

R

N

S

R1 83 H 84 H 85 Me

S

N

R 37 Et 38 Me

R

91 92 93

N 2 N1

S

S

39

C

C

3 4N

N

R Bu Et Ph

S

5

R

N

R

S

O

N

86 87 88

R

R

5 N1

Me Me

S

2

4

R1 H H H Me Et

N

R1 H Me Me

78 79 80 81 82

R

R

c

Statistically corrected in order to put the pK BI2 value on a per thiocarbonyl basis. Contain a contribution of a second 1 : 1 complex on N2 (86–88) or N3 (89, 90). d Contain a contribution of a second 1 : 1 complex on the second thiocarbonyl group.

b

Me

Me

R

R


R2 H Me Et Me Et

3N

Y

R2 Me H Me

S

N R

R

R

94 95 96 97 98 99 100 101

70 71 72 73 74 75

Y O O O O S S S S

H

Me

R1 H Me Me H Me H

R1 H H Me Me H H Me Me

N

Me

R2 H H Me H H H

89

S

SMe

R2 H Me H Me H Me H Me

N

R3 H H H Me Me Ph

273

Lewis base

2,2,2-Trifluoroethanethiol Hydrogen sulfide Ethanethiol 4-Methoxythiophenol 4-Methylthiophenol 3-Methylthiophenol 4-Fluorothiophenol 4-Chlorothiophenol 3-Chlorothiophenol Thiophenol 2-Chloroethyl methyl sulfide

Methyl sulfide

Ethyl methyl sulfide

Methyl n-propyl sulfide n-Butyl methyl sulfide n-Decyl methyl sulfide Methyl isopropyl sulfide Ethyl sulfide

Isopropyl sulfide n-Propyl sulfide

tert-Butyl sulfide 4-Methylthioanisole Thioanisole

No.

1 2 3 4 5 6 7 8 9 10 11

12

13

14 15 16 17 18

19 20

21 22 23

t-Bu2 S 4-MeC6 H4 SMe PhSMe

i-Pr2 S n-Pr2 S

MeS-n-Pr MeS-n-Bu MeS-n-Dec MeS-i-Pr Et2 S

MeSEt

Me2 S

CF3 CH2 SH H2 S EtSH 4-MeOC6 H4 SH 4-MeC6 H4 SH 3-MeC6 H4 SH 4-FC6 H4 SH 4-ClC6 H4 SH 3-ClC6 H4 SH PhSH ClCH2 CH2 SMe

Formula CCl4 CH2 Cl2 cHex cHex cHex cHex cHex cHex cHex cHex cHex CCl4 Hept CCl4 Hept CCl4 CCl4 CCl4 CCl4 CCl4 Hept CCl4 CCl4 cHex CCl4 CCl4 cHex cHex CCl4

Solvent 1.20 1.3 23.2 24.4b 18.2b 14.1b 9.1b 6.4b 4.7b 11.6b 12.23 10.64 78.1 72.6 119 107 112 115 120 155.4 195 166.6 184.7 161 152 159.3 19.6b 10.27b 8.57

Kc

2.26 1.29 1.01

2.32 2.21

2.10 2.11 2.13 2.24 2.29

2.08

1.89

0.15 0.11 1.37 1.39 1.26 1.15 0.96 0.81 0.67 1.06 1.09

pK BI2

−12.87 −7.38 −5.77

−13.24 −12.60

−12.01 −12.07 −12.18 −12.81 −13.07

−11.85

(Continued)

[205] [206] [207] [207] [207] [207] [207] [207] [207] [207] [65] [65] [208] [60, 209] [208] [209] [210] [209] [209] [60] [210–214] [60, 210] [60] [74] [74] [60] [215] [65] [65]

−0.87 −0.65 −7.79 −7.92 −7.19 −6.56 −5.47 −4.60 −3.84 −6.08 −6.21 −10.80

Ref.

∆G◦

Table 5.18 Equilibrium constants Kc (l mol−1 ) and diiodine basicity scales pK B I 2 and ∆G◦ (kJ mol−1 ) for the complexes of diiodine with thiols, sulfides and disulfides.

274

Lewis base

4-Fluorothioanisole 4-Chlorothioanisole 3-Chlorothioanisole 3-Fluorothioanisole 4-Dimethylaminothioanisole 2-Methoxythioanisole 4-Methoxythioanisole 3-Methylthioanisole 4-(Methylthio)thioanisole 2-Methylthioanisole 2-(Methylthio)thioanisole 4-Bromothioanisole 2-Bromothioanisole 2-Chlorothioanisole tert-Butyl phenyl sulfide

Diphenyl sulfide

4-Methylphenyl phenyl sulfide 4-Methoxyphenyl phenyl sulfide 4-Chlorophenyl phenyl sulfide 3-Methylphenyl phenyl sulfide 3-Methoxyphenyl phenyl sulfide 3-Chlorophenyl phenyl sulfide 2-Methylphenyl phenyl sulfide Bis(4-methylphenyl) sulfide Bis(4-methoxyphenyl) sulfide Bis(4-chlorophenyl) sulfide Bis(2-methylphenyl) sulfide 4-Chlorophenyl 4 -methoxyphenyl sulfide

No.

24 25 26 27 28 29 30 31 32 33 34 35 36 37 38

39

40 41 42 43 44 45 46 47 48 49 50 51


4-MeC6 H4 SPh 4-MeOC6 H4 SPh 4-ClC6 H4 SPh 3-MeC6 H4 SPh 3-MeOC6 H4 SPh 3-ClC6 H4 SPh 2-MeC6 H4 SPh 4-MeC6 H4 SC6 H4 -4 -Me 4-MeOC6 H4 SC6 H4 -4 -OMe 4-ClC6 H4 SC6 H4 -4 -Cl 2-MeC6 H4 SC6 H4 -2 -Me 4-ClC6 H4 SC6 H4 -4 -OMe

Ph2 S

4-FC6 H4 SMe 4-ClC6 H4 SMe 3-ClC6 H4 SMe 3-FC6 H4 SMe 4-Me2 NC6 H4 SMe 2-MeOC6 H4 SMe 4-MeOC6 H4 SMe 3-MeC6 H4 SMe 4-MeSC6 H4 SMe 2-MeC6 H4 SMe 2-MeSC6 H4 SMe 4-BrC6 H4 SMe 2-BrC6 H4 SMe 2-ClC6 H4 SMe PhS-t-Bu

Formula cHex cHex cHex cHex Hept Hept Hept Hept Hept Hept Hept Hept Hept Hept cHex CCl4 cHex CCl4 cHex cHex cHex cHex cHex cHex cHex cHex cHex cHex cHex cHex

Solvent

−7.66 −7.55 −5.94 −5.94 −5.28 −4.64 −3.50 −1.84 −1.84 −8.03 1.34 1.32 1.04 1.04d 0.92 0.81d 0.61 0.32 0.32 1.41

0.74 0.80 0.49 0.67 0.56 0.38 0.73 0.79 0.86 0.45 0.76 0.60

0.53

−4.23 −4.56 −2.80 −3.84 −3.18 −2.17 −4.18 −4.48 −4.89 −2.55 −4.36 −3.44

−3.03

[215] [215] [215] [215] [216] [216] [216] [216] [216] [216] [216] [216] [216] [216] [74] [74] [74] [74] [217] [217] [217] [217] [217] [217] [217] [217] [217] [217] [217] [217]

−5.42 −4.27 −3.78 −3.62

0.95 0.75 0.66 0.63

8.9b 5.6b 4.6b 4.3b 67b,c 22b 21b 11b 22b 8.4b 13b 4.1b 2.1b 2.1b 25.5b 24.1 3.4b 2.9 5.5b 6.3b 3.1b 4.7b 3.6b 2.4b 5.4b 6.1b 7.2b 2.8b 5.8b 4.0b

Ref.

∆G◦

pK BI2

Kc

275

55

56

57

58

S

c

b

59

S

n-PrSS-n-Pr i-PrSS-i-Pr t-BuSS-t-Bu t-BuSSMe t-BuSSEt

MeSSMe EtSSEt

a

a

a

a

a

a

a

a

a

a

60

S

S

61

S S

2,6-Me2 C6 H3 SC6 H3 -2 ,6 -Me2

a

Contains a contribution of the π site(s). Contains a contribution of the nitrogen site. d Statistically corrected in order to put the pK BI2 value on a per sulfur basis.

O

54

S

53

S

S

n-Propyl disulfide Isopropyl disulfide tert-Butyl disulfide tert-Butyl methyl disulfide tert-Butyl ethyl disulfide

66 67 68 69 70

S

1,3,5-Trithiane Methyl disulfide Ethyl disulfide

63 64 65

S

1,4-Dithiane 1,2-Dithiane 1,4-Thioxane

60 61 62

S

2,2,5,5-Tetramethyltetrahydrothiophene Pentamethylene sulfide

58 59

Formulae:

Tetrahydrothiophene

57

a

Bis(2,6-dimethylphenyl) sulfide Thianthrene Phenoxathiin trans-2,3-Butylene sulfide Trimethylene sulfide

52 53 54 55 56

62

O

S

S

63

S

cHex CCl4 CCl4 Hept Hept CCl4 Hept CCl4 CCl4 Hept CCl4 CCl4 CCl4 Hept CCl4 CCl4 CCl4 Hept CCl4 CCl4 CCl4 CCl4 CCl4 CCl4

S

3.7b 1.2b 0.84b 35.1 100.9 79.1 230.4 199 78.4 154.6 135.6 76.9 9.93 26.2 23.0 6.8 2.65 4.65 3.87 3.76 4.36 5.93 3.62 4.12 −2.43 −1.11 −2.09 −1.97 −2.34 −3.09 −1.88 −2.20 0.35d 0.41d 0.54d 0.33d 0.39d

−9.39 −4.36 −8.10

−11.13 −12.50

−13.49

−8.82 −11.44

−3.24

0.43d 0.20d 0.37d

1.64d 0.76d 1.42

1.95 2.19

2.36

1.55 2.00

0.57

[217] [143] [143] [211] [211] [211] [211] [60, 218] [60] [211] [60, 218] [219] [220] [108] [108] [218] [220] [210] [210] [220] [220] [220] [220] [220]

276

Tri-n-octylphosphine sulfide

Triethylphosphine sulfide Trimethylphosphine sulfide Hexamethylthiophosphoric triamide Triphenylphosphine sulfide Isopropyl diethylphosphinothioate Bis(dimethylamino)chlorophosphine sulfide Bromodiethylphosphine sulfide Triethoxyphosphine sulfide

Diethylchlorophosphine sulfide Chloro(diethoxy)phosphine sulfide Dichloro(dimethylamino)phosphine sulfide Dichloro(ethoxy)phosphine sulfide Dichlorophenylphosphine sulfide Trichlorophosphine sulfide

2 3 4 5 6 7 8 9

10 11 12 13 14 15

Lewis base

1

No.

Et2 PSCl (EtO)2 PSCl Me2 NPSCl2 EtOPSCl2 PhPSCl2 Cl3 PS

Et3 PS Me3 PS (Me2 N)3 PS Ph3 PS Et2 PSO-i-Pr (Me2 N)2 PSCl Et2 PSBr (EtO)3 PS

Oct3 PS

Formula cHex CCl4 CCl4 CCl4 CCl4 CCl4 CCl4 CCl4 CCl4 cHex CCl4 Hept CCl4 CCl4 CCl4 CCl4 CCl4

Solvent 4715 1426 1450 1230 620 152 127 13.7 13.2 17.99 10.64 3.0 1.58 0.52 0.33 0.27 0.05

Kc

0.48 0.31 −0.24 −0.46 −0.56 −1.39

3.68 3.60 3.26 2.57 2.48 1.38 1.36 1.26

3.67

pK BI2

[65] [65] [162] [221] [221] [65] [221] [221] [221] [65] [65] [83] [221] [221] [221] [221] [221]

−20.97

−2.72 −1.79 1.34 2.63 3.19 7.94

−21.02 −20.55 −18.62 −14.66 −14.15 −7.88 −7.77 −7.16

Ref.

∆G◦

Table 5.19 Equilibrium constants Kc (l mol−1 ) and diiodine basicity scales pK B I 2 and ∆G◦ (kJ mol−1 ) for the complexes of diiodine with thiophosphoryl compounds.

277

1,4-Selenoxane 1,4-Selenothiane 1,4-Diselenane 4-Dimethylaminoselenoanisole 4-Aminoselenoanisole 4-Methoxyselenoanisole 4-(Methylseleno)selenoanisole 2-Methoxyselenoanisole 4-Methylselenoanisole 4-(Methylthio)selenoanisole 3-Methylselenoanisole 3-Methoxyselenoanisole Selenoanisole

2-Methylselenoanisole 3-(Methylthio)selenoanisole 4-Fluoroselenoanisole 4-Chloroselenoanisole 4-Bromoselenoanisole 3-Chloroselenoanisole 4-Acetylselenoanisole 4-(Ethoxycarbonyl)selenoanisole

19 20 21 22 23 24 25 26

tert-Butyl phenyl selenide

5

6 7 8 9 10 11 12 13 14 15 16 17 18

c-(CH2 )4 Se c-(CH2 )5 Se (C2 H5 )2 Se (CH3 )2 Se

Selenides Tetrahydroselenophene Pentamethylene selenide Ethyl selenide Methyl selenide

1 2 3 4

2-MeC6 H4 SeMe 3-MeSC6 H4 SeMe 4-FC6 H4 SeMe 4-ClC6 H4 SeMe 4-BrC6 H4 SeMe 3-ClC6 H4 SeMe 4-MeCOC6 H4 SeMe 4-EtOCOC6 H4 SeMe

4-Me2 NC6 H4 SeMe 4-H2 NC6 H4 SeMe 4-MeOC6 H4 SeMe 4-MeSeC6 H4 SeMe 2-MeOC6 H4 SeMe 4-MeC6 H4 SeMe 4-MeSC6 H4 SeMe 3-MeC6 H4 SeMe 3-MeOC6 H4 SeMe C6 H5 SeMe

a

a

a

C6 H5 SeC(CH3 )3

Formula CCl4 CCl4 Hept cHex CCl4 Hept cHex CCl4 CCl4 CCl4 CCl4 Hept Hept Hept Hept Hept Hept Hept Hept Hept Hept CCl4 Hept Hept Hept Hept Hept Hept Hept Hept

Solvent 2203 1397 1534 494 530 195 243 240 138 198b 282 468 251 146 125 118 90 82b 77 64 62 65 50 44b 38 29 28 22 21 20

Kc

−9.70 −9.02 −8.35 −8.26 −7.66 −7.55 −7.43 1.58 1.46 1.45 1.34 1.32 1.3

−10.77 −10.31 −10.22 1.89 1.81 1.79 1.70

−12.08 −15.24 −13.70 −12.35 −10.25 −11.83 −11.15

−12.03

−13.07

−18.80 −17.68 −18.18 −15.38

∆G◦

2.12c 2.67 2.40 2.16 1.80c 2.07 1.95

2.11

2.29

3.29 3.10 3.19 2.69

pK BI2

(Continued)

[222] [222] [214] [65] [65] [74] [74] [74] [219] [219] [219] [216] [216] [216] [216] [216] [216] [216] [216] [216] [75] [75] [216] [216] [216] [216] [216] [216] [216] [216]

Ref.

Equilibrium constants Kc (l mol−1 ) and diiodine basicity scales pK B I 2 and ∆G◦ (kJ mol−1 ) for the complexes of diiodine with selenium

Lewis base

No.

Table 5.20 bases.

278

3-Bromoselenoanisole 4-(Methoxycarbonyl)selenoanisole 2-Chloroselenoanisole 2-Bromoselenoanisole Bis(4-ethoxyphenyl) selenide Bis(4-methoxyphenyl) selenide Bis(4-methylphenyl) selenide 2-Biphenyl 4 -methylphenyl selenide Bis(2-biphenyl) selenide Diphenyl selenide Bis-(4-chlorophenyl) selenide Bis(2-methylphenyl) selenide Bis(2-chlorophenyl) selenide Dibenzoselenophene Selenoamides, selenoureas 1,3-Imidazolidine-2-selenone 1-Methyl-1,3-imidazolidine-2-selenone

1-Ethyl-1,3-imidazolidine-2-selenone

1,3-Dimethyl-1,3-imidazolidine-2selenone

1,3-Diethyl-1,3-imidazolidine-2selenone

1-Methyl-1,3-oxazolidine-2-selenone 1,3-Thiazolidine-2-selenone 1-Methyl-1,3-thiazolidine-2-selenone 1-Methylbenzimidazole-2-selenone 1,3-Dimethylbenzimidazole-2selenone Benzoxazole-2-selenone 3-Methylbenzoxazole-2-selenone

27 28 29 30 31 32 33 34 35 36 37 38 39 40

41 42

43

44

45

46 47 48 49 50

51 52

Lewis base

(Continued)

No.

Table 5.20

1 340 000 2 590 000 358 000 1 370 000 315 000 344 000 397 000 620 000 368 000 21 700 42 000 23 700 861 000 774 000 8000 3210

CH2 Cl2 CH2 Cl2 CCl4 CH2 Cl2 CCl4 CH2 Cl2 CCl4 CH2 Cl2 CCl4 CH2 Cl2 CH2 Cl2 (17 ◦ C) CH2 Cl2 CH2 Cl2 CH2 Cl2 CH2 Cl2 (17 ◦ C) CH2 Cl2

a

a

a

a

a

a

a

a

a

a

a

a

a

19 17 14 13 76.9 52.6 40 37 28.6 28 8.6 5.6 4.5 1.9

Kc

Hept Hept Hept Hept CCl4 CCl4 CCl4 CCl4 CCl4 CCl4 CCl4 CCl4 CCl4 CCl4

Solvent

3-BrC6 H4 SeMe 4-MeOCOC6 H4 SeMe 2-ClC6 H4 SeMe 2-BrC6 H4 SeMe (4-EtOC6 H4 )2 Se (4-MeOC6 H4 )2 Se (4-MeC6 H4 )2 Se 2-C6 H5 C6 H5 SeC6 H4 -4 -Me (2-PhC6 H4 )2 Se (C6 H5 )2 Se (4-ClC6 H4 )2 Se (2-MeC6 H4 )2 Se (2-ClC6 H4 )2 Se

Formula

[225] [187] [225] [225] [187] [187]

d

d

d

d

d

d

d

d

[187] [187]

[225] [225] d

d

d

d

d

d

d

[216] [216] [216] [216] [223] [224] [224] [223] [223] [224] [224] [223] [223] [223]

−7.30 −7.02 −6.54 −6.36 −10.76 −9.68 −9.00 −8.82 −8.19 −8.12 −5.26 −4.19 −3.70 −1.60

[225] [225] [225] [225] [225] [225]

Ref.

∆G◦

d

1.28 1.23 1.15 1.11 1.89 1.70 1.58 1.55 1.43 1.42 0.92 0.73 0.65 0.28

pK BI2

279

Se

X

Formulae:

R

N

X

Se

49 50 51 52 53

R

Me Me H Me Me

X

NH NMe O O S

40

Se

H

Me

X

5 4 N1 2 3 N

Me

R

Y

N

CH2 Cl2 CH2 Cl2

a a

CCl4 CHCl3 CCl4 CHCl3 CH2 Cl2 CCl4

(C2 H5 )3 PSe (c-C6 H11 )3 PSe (C6 H5 )3 PSe

H

Se

54 55 56 57 58

X

41 42 43 44 45 46 47 48

Se O Se S Se

X

O Se S Se Se

Y

X NH NH NH NMe NEt O S S

(C2 H5 O)3 PSe

Hept

(Me2 N)3 PSe

a

a

R H Me Et Me Et Me H Me

CH2 Cl2 CH2 Cl2 CH2 Cl2 CH2 Cl2

a

a

18 000 46 600 2999 4476 7670 105

23 800

5850

475

65 400 20 600 1000 13 400

d

d

d

d

4.38

d

d

d

d

d

d

−24.98

[162] [227] [108] [108] [108] [162]

[226]

[85]

[85]

[187] [85] [85] [85]

Contain the contribution of a sulfur complex. c Statistically corrected to put the pK BI2 value on a per selenium basis. d The Kc values measured in CCl4 and CH2 Cl2 could not be referred to alkanes because of the too low solubility of these selenocarbonyl compounds and phosphine selenides in alkanes.

b

a

X O S Se

63

6 7 8

Triethylselenophosphate

60 61 62

59

58

57

3-Methylbenzothiazole-2-selenone 5,5-Dimethyl-2-selenohydantoin 5,5-Dimethyl-4-selenohydantoin 5,5-Dimethyl-2-seleno-4thiohydantoin 5,5-Dimethyl-2-thio-4selenohydantoin 5,5-Dimethyl-2,4-diselenohydantoin Phosphine selenides Hexamethylselenophosphoric triamide Triethylphosphine selenide Tricyclohexylphosphine selenide Triphenylphosphine selenide

53 54 55 56

280

Lewis base

2-Chloropropane 1-Bromobutane Bromocyclohexane

Iodomethane Iodoethane

2-Iodopropane 2-Iodo-2-methylpropane Iodocyclohexane

No.

1 2 3

4 5

6 7 8

(CH3 )2 CHI (CH3 )3 CI c-C6 H11 I

CH3 I CH3 CH2 I

(CH3 )2 CHCl CH3 (CH2 )3 Br c-C6 H11 Br

Formula cHex Hept CCl4 cHex Hept Hept CCl4 Hept Hept cHex CCl4

Solvent 0.11 0.21 0.26 0.38 0.23 0.36 0.29 0.43 1.33 1.33 0.97

Kc

∆G◦ 5.48 3.88 2.40 3.65 2.51 2.11 −0.68 −0.68

pK BI2 −0.96 −0.68 −0.42 −0.64 −0.44 −0.37 0.12 0.12

[65] [228] [65] [65] [228] [228] [228] [228] [228] [65] [228]

Ref.

Table 5.21 Equilibrium constants Kc (l mol−1 ) and diiodine basicity scales pK B I 2 and ∆G◦ (kJ mol−1 ) for the complexes of diiodine with haloalkanes.

281

14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29

tert-Butyl isothiocyanate n-Butyl isothiocyanate Dodecyl isothiocyanate Cyclohexyl isothiocyanate Trimethyltin isothiocyanate Triethyltin isothiocyanate Tri(n-butyl)tin isothiocyanate Tri(isopropyl)tin isothiocyanate Drugs Promethazine Chlorpromazine Chlorproethazine Vanitiolide Clotrimazole Tetrahydrozoline Triflupromazine Alimemazine Levomepromazine Ethymemazine Perphenazine Fluphenazine Isothipendyl Chlorprothixene Imipramine Clomipramine a

a

a

a

a

a

a

a

a

a

a

a

a

a

a

a

t-BuNCS n-BuNCS n-DoNCS c-C6 H11 NCS Me3 SnNCS Et3 SnNCS n-Bu3 NCS i-Pr3 NCS

CH2 CHCH2 NCS MeNCS EtNCS

Isothiocyanates Allyl isothiocyanate Methyl isothiocyanate Ethyl isothiocyanate

3 4 5

6 7 8 9 10 11 12 13

Me2 NC(Me2 N) NH Ph3 As

Formula

Tetramethylguanidine Triphenylarsine

Lewis base

1 2

No.

CCl4 CCl4 CCl4 CCl4 CCl4 CCl4 CCl4 CCl4 CCl4 CCl4 CCl4 CCl4 CCl4 CCl4 CCl4 CCl4 (20 ◦ C) (20 ◦ C) (20 ◦ C) (20 ◦ C) (20 ◦ C) (20 ◦ C) (20 ◦ C) (22 ◦ C) (20 ◦ C)

(20 ◦ C)

(20 ◦ C) (20 ◦ C)

(17 ◦ C)

CCl4 CCl4 (20 ◦ C) Hept CCl4 (20 ◦ C) CH2 Cl2 (20 ◦ C) CCl4 (20 ◦ C) CCl4 (20 ◦ C) CCl4 (20 ◦ C) CCl4 (20 ◦ C) CH2 Cl2 CH2 Cl2 CH2 Cl2 CH2 Cl2

Hept CCl4 CH2 Cl2

Solvent

4259 2221 2507 653 364 308 2095 788 738 731 1074 1115 2683 3401 4907 4545

0.86 0.87 0.98 1.1 0.65 1.18 1.26 1.41 1.76 5.34 6.15 7.10 8.04

23 415 1067 16 244

Kc

−0.01

4.37

pK BI2

0.04

[229] [230] [230]

−24.94

(Continued)

[234] [110] [234] [185] [185] [185] [185] [185] [185] [185] [185] [185] [185] [185] [185] [185]

[231] [232] [231] [231] [231] [232] [232] [232] [232] [233] [233] [233] [233]

Ref.

∆G◦

Table 5.22 Equilibrium constants Kc (l mol−1 ) and diiodine basicity scales pK B I 2 and ∆G◦ (kJ mol−1 ) for the complexes of diiodine with miscellaneous bases.

282

R1 H Cl H H

CH3CH2

28 29 30 31

N

R

N

32

CH2CH2CH2

CH2CH2CH2

O

Me

R2 CH2CH2CH2N(CH3)2 CH2CH2CH2N(CH3)2 CH2CH(CH3)CH2N(CH3)2 CH2CH2CH2NHCH3

R

N

N

17

S

C

N

a

a

a

a

a

26

33

N

N CH Ph

CH CH N CH CH OH

CH CH

O

Formula

CH CH(CH )N(CH )

N

Me

N

O

NCH2CH2OH

CH2CH2CH2N

25 CF3 S

NCH2CH2OH

CH2CH2CH2N

MeO

24 Cl

CH2CH(CH3)N(CH3)2 CH2CH2CH2N(CH3)2 CH2CH2CH2N(C2H5)2 CH2CH2CH2N(CH3)2 CH2CH(CH3)CH2N(CH3)2 CH2CH(CH3)CH2N(CH3)2 CH2CH(CH3)CH2N(CH3)2

H Cl Cl CF3 H OCH3 C2H5

HO

R2

R

R1

R

N

S

Formulae:

14 15 16 20 21 22 23

a

Trimipramine Desipramine Dipropyline Bamifylline Acetiamine

Lewis base

No.

30 31 32 33 34

(Continued)

Table 5.22

Me

18

C

N

Cl

N

N

27

CH

NH

N

34

19

C

H

Cl

Me

CHCH CH N(CH )

S

N

C

(20 ◦ C)

(20 ◦ C) (20 ◦ C) (20 ◦ C)

CH CH OCOCH

SCOCH

H

N

N

O

CCl4 CCl4 CCl4 CCl4 CCl4

Solvent 1003 2987 1715 57 216

Kc

pK BI2

∆G◦ [185] [185] [185] [185] [185]

Ref.


283

(i) The sites are equivalent or identical by symmetry. The value given in the K c column is the observed total constant. In order to obtain the pK BI2 value per basic site, a statistical correction is applied to K c (t). Examples: 1,4-dioxane 1,4-dithiane

Kc (t) (in cyclohexane) = 1.56 l mol−1 pK BI2 = log(1.56/2) = −0.11 Kc (t) (in CCl4 ) = 76.9 l mol−1 pK BI2 = 0.990log(76.9/2) + 0.07 = 1.64

(ii) One site is significantly more basic than the other(s). One can safely attribute the diiodine basicity to only the most basic site and give to pK BI2 the value logK c (t). For example, 1,4-thioxane has two potential halogen-bond acceptor sites, the sulfur and the oxygen atoms. Considering the value K c (O) = (1.56/2) = 0.78 l mol−1 for one oxygen of 1,4-dioxane, the contribution of the oxygen atom to K c (t) = 26.2 l mol−1 can be neglected. The diiodine basicity of 1,4-thioxane is controlled by the sulfur atom, which can be given the pK BI2 value of log26.2 = 1.42. (iii) Several individual constants contribute significantly to K c (t). It is then necessary to determine the various individual constants in order to assign a pK BI2 value to each significantly basic site. There seem to be no examples in the literature where the separation of K c (t) into individual constants has been achieved. For example, both the selenium and sulfur atoms of 1,4-selenothiane are significantly basic towards diiodine. This is shown by comparing the diiodine basicities of one sulfur atom of 1,4-dithiane (K c = 76.9/2 = 38.5 l mol−1 ) and one selenium atom of 1,4-diselenane (K c = 282/2 = 141 l mol−1 ). So, in the observed constant K c (t) = 198 l mol−1 , which is equal to the sum K c (Se) + K c (S) for each site of 1,4-selenothiane, the sulfur basicity cannot be neglected. While K c (Se) and K c (S) have not been individually determined, no pK BI2 value can be attributed to 1,4-selenothiane. The value logK c (t) = log198 = 2.30 is devoid of thermodynamic significance since the logarithm of a sum is not equal to the sum of logarithms.

5.3

Is the Diiodine Basicity Scale a General Halogen-Bond Basicity Scale?

The choice of diiodine as a reference halogen-bond donor to construct the diiodine basicity scale has been dictated mainly by the availability of literature data. Sets of equilibrium constants for halogen bonding of a number of different Lewis bases can also be assembled for the halogen-bond donors ICl, IBr, Br2 and ICN. The literature references and experimental conditions for the determination of most of these data are summarized in Table 5.23. The applicability of the diiodine basicity scale to reference acids of the XY type (in which X is the halogen and Y is any group) therefore needs to be examined. For a given series of bases, the possibility that logK for halogen bonding to XY acids is linearly related to logK for halogen bonding to diiodine can be investigated by means of the linear Gibbs energy Equation 5.17: pK BXY = m pK BI2 + c

(5.17)

284


Table 5.23 Literature references to determinations of complexation constants for some halogen-bond donor/Lewis base systems. Lewis base Carbon bases Aromatic compounds Alkylbenzenes Aromatic compounds Aromatic compounds Nitrogen bases Nitriles Nitriles Methylpyridines Substituted pyridines Trimethylamine Oxygen bases Ethers Aldehydes Carbonyl compounds Lactones, cyclic ketones Ethyl acetate Lactams N,N-Dimethylacetamide N,N-Diethylacetamide Pyridine N-oxides Miscellaneous oxygen bases Sulfuryl compounds Phosphoryl compounds Sulfur bases Ethyl sulfide Triphenylphosphine sulfide Ethyl isothiocyanate

Halogen-bond donor

Solvent (t/◦ C)

Br2 ICl IBr Br2 , IBr, ICl

CCl4 CCl4 CCl4 CCl4

ICl Br2 , IBr, ICl ICl, IBr Br2 , ICl, IBr ICN ICN

CCl4 (28 ◦ C) CCl4 (25 ◦ C) CCl4 (25 ◦ C) CCl4 (30 ◦ C) CCl4 (27 ◦ C) C6 H6 (25 ◦ C)

[239] [138, 139, 240-242] [243] [244] [245] [246]

ICl ICl ICN ICl, IBr ICl ICl ICl ICl Br2 , ICl ICl, IBr, ICN, IC CR Br2 , ICl, IBr IBr IBr, ICl, ICN, ICF3 ICN ICN

CCl4 (20 ◦ C) CCl4 (25 ◦ C) Hept (27 ◦ C) CS2 (30 ◦ C) CCl4 (25 ◦ C) CCl4 (25 ◦ C) CCl4 (25 ◦ C) CCl4 (25 ◦ C) CCl4 (25 ◦ C) CCl4 (25 ◦ C) CCl4 (30 ◦ C) CCl4 (25 ◦ C) CCl4 (25 ◦ C) CS2 (27 ◦ C) C6 H6

[247] [141] [248] [249, 250] [42] [43] [251] [44] [252] [154] [165] [253] [59] [254] [255]

ICN IBr, ICl, ICN, ICF3 IBr, ICl

CCl4 (25 ◦ C) CCl4 (25 ◦ C) CCl4 (20 ◦ C)

[256] [59] [257]

(25 ◦ C) (25 ◦ C) (24 ◦ C) (25 ◦ C)

Ref. [235] [89, 236] [237] [238]

where m and c are constants characteristic of the studied XY halogen-bond donor. Application of Equation 5.17 to data on ICl, IBr, Br2 and ICN gives the results summarized in Table 5.24 and illustrated in Figures 5.8 and 5.9. The correlations obtained with ICl, IBr and Br2 are good enough to conclude that Equation 5.17 is rather general for heavy dihalogens and interhalogens. However, in the case of ICN, the correlation is good only within families of related bases. One good linear plot is obtained from data on oxygen bases. A second line corresponds to the points for sp2 nitrogen bases. The location of the points for Et2 S and Ph3 PS shows that the family(ies) of sulfur bases should be distinguished from those of oxygen and nitrogen bases (see Figure 5.9). Such family-dependent behaviour is also observed for the correlations between the diiodine basicity scale and (i) IR and UV spectroscopic shifts upon halogen bonding (see below) and (ii) halogen-bond lengths [14]. However, a set of 19 F NMR chemical shifts


285

Table 5.24 Summary of linear Gibbs energy Equation 5.17 results. Halogen-bond donor XY Iodine monochloride Iodine monobromide Dibromine Iodine cyanide

m

c

r

s

n

1.79 (±0.04) 1.53 (±0.04) 0.70 (±0.03) — 0.88 (±0.07) 0.70 (±0.02)

1.42 (±0.04) 0.87 (±0.05) 0.48 (±0.05) — 0.66 (±0.10) 0.13 (±0.05)

0.975 0.992 0.986 0.638 0.971 0.998

0.33 0.21 0.16 — 0.17 0.03

93 32 19 20 (all bases) 11 (oxygen bases) 7 (nitrogen bases)

upon halogen bonding [11] appears to be rather generally correlated by pK BI2 . Indeed, the relationship ∆δ/ppm = 2.12 (±0.22) pK BI2 + 2.7 (±0.4)

(5.18)

n = 14, r = 0.941, s = 1.1 2

is observed between ∆δ = δ(in n-pentane) – δ(in a basic solvent) for the 19 F NMR chemical shifts of ICF2 CF2 I halogen bonded to a variety of 14 oxygen, nitrogen and sulfur bases and the diiodine basicity scale.

5.4

The Diiodine Affinity Scale

There have been many determinations of the enthalpy of complexation of diiodine with Lewis bases. The data measured in alkanes and/or CCl4 are collected in Table 5.25. It is difficult, however, to assemble a reliable diiodine affinity scale from these values. The reasons are that many results appear inaccurate and determinations have been made in 6

ICl

5

IBr

XY basicity

4 3 2

Br2

1 0 -1 -2 -2

0

2

4

6

Diiodine basicity Figure 5.8 Linear Gibbs energy relationship for the formation of halogen-bonded complexes of ICl, IBr, Br2 and I2 . The dashed line is of slope unity.

286


Iodine cyanide basicity

2.5 2 Ph3PS

1.5

Et2S

1 0.5 0 -0.5

0

0.5

1

1.5

2

2.5

3

Diiodine basicity Figure 5.9 Limited linear Gibbs energy relationship for formation of halogen-bonded complexes of ICN and I2 . The upper line represents oxygen bases, the lower line sp2 nitrogen bases and ∗ sulfur bases.

different solvents. The lack of accuracy is shown by the serious discrepancies observed between the results from various authors for the same system. These discrepancies may arise from (i) a too restricted range of temperature variation when enthalpies are obtained from van’t Hoff plots, (ii) an incorrect standard state when the molar concentration scale is used for equilibrium constants and (iii) ionization of the complex. The data in CCl4 are difficult to refer to alkanes (or vice versa) because there are not enough accurate enthalpies in both solvents for the same family of bases to establish useful family-dependent relationships between enthalpies in CCl4 and enthalpies in alkanes. Nevertheless, the values given in Table 5.25 may be useful for comparison with other Lewis affinity scales and with the spectroscopic scales presented later. Their comparison with theoretical diiodine affinities may also help in choosing the best calculation method(s) in halogen-bonding studies.

5.5 5.5.1

Spectroscopic Scales Infrared Shifts of ICN, I2 and ICl

The changes in the vibrational spectrum of iodine cyanide, iodine monochloride and diiodine as these molecules form halogen-bonded complexes B· · ·ICN, B· · ·ICl and B· · ·I2 with series of Lewis bases B have been extensively studied [35, 36, 273–277]. The main perturbation is a lowering of the frequency of the I CN, I Cl and I I stretching motions. These frequency shifts are greater for strong bases, such as pyridines, than for weak bases, such as benzene or dioxane. Hence correlations were attempted between frequency shifts


287

Table 5.25 Enthalpies (kJ mol−1 ) for the complexation of diiodine with Lewis bases. Lewis base Alkylbenzenes Hexamethylbenzene Pentamethylbenzene 1,2,3,4-Tetramethylbenzene 1,2,3,5-Tetramethylbenzene 1,2,4,5-Tetramethylbenzene 1,3,5-Trimethylbenzene 1,3,5-Triethylbenzene 1,2,3-Trimethylbenzene 1,2,4-Trimethylbenzene 1,3-Dimethylbenzene 1,2-Dimethylbenzene 1,4-Dimethylbenzene Hexaethylbenzene tert-Butylbenzene Isopropylbenzene n-Propylbenzene Isobutylbenzene n-Butylbenzene Ethylbenzene Toluene Benzene Thiophene and furan Thiophene 2-Methylfuran Amines Ammonia n-Butylamine tert-Butylamine sec-Butylamine Dodecylamine Isobutylamine Ethylamine Methylamine 3,4-Dimethylaniline 2,3-Dimethylaniline 2,6-Dimethylaniline Piperidine Diethylamine Dimethylamine Di-n-dodecylamine

−∆H◦

Ref.

Hept CCl4 CCl4 CCl4 CCl4 CCl4 Hept CCl4 CCl4 CCl4 CCl4 Hept CCl4 Hept CCl4 Hept CCl4 CCl4 CCl4 CCl4 CCl4 CCl4 CCl4 CCl4 Hept CCl4 Hept CCl4

15.9 15.6 13.4 11.2 12.9 12.0 12.0 12.2 11.0 11.0 10.9 9.7 10.3 9.2 9.8 8.9 9.4 7.5 8.5 8.3 8.6 8.4 8.9 7.8 7.9 8.1 6.8 6.4

[90] [91] [88] [88] [88] [88, 91] [90] [88, 91] [91] [88] [88] [90] [88] [90] [88] [90] [88, 91] [91] [86] [86] [86] [86] [86] [86] [90] [86, 258] [90] [64, 86, 91, 258]

Hept Hept

10.0 10.0

[95] [95]

Hept Hept Hept Hept Hept Hept Hept Hept Hept Hept Hept Hept Hept Hept Hept

20.1 35.1 29.5 37.6 24.3 38.8 31.0 29.7 26.5 9.9 7.9 43.1 40.6 41.0 36.4

[99] [99] [100] [100] [102] [100] [99] [99] [105] [105] [105] [99] [99] [99] [102]

Solvent

(Continued)

288


Table 5.25 (Continued) −∆H◦

Lewis base

Solvent

Trimethylamine Triethylamine Tri-n-octylamine Chlorpromazine Tri-n-butylamine Tri-n-propylamine Tri-n-dodecylamine N,N-Diisopentyloctylamine N,N-Di-n-octylisopentylamine N,N-Diisopentylhexadecylamine Triisopentylamine N,N-Dimethyl-4-toluidine N,N-Dimethylaniline N,N-Dimethyl-2-toluidine N,N-Dimethyl-2,6-xylidine Pyridines 4-Dimethylaminopyridine 4-Aminopyridine 2-Amino-4-methylpyridine 2,4-Dimethylpyridine

Hept Hept Hept CCl4 Hept Hept Hept Hept Hept Hept Hept Hept Hept Hept Hept

50.6 50.2 48.1 46.9 51.5 50.6 45.6 47.3 45.6 43.1 47.7 34.7 34.3 9.6 7.1

[99] [99] [102] [110] [99] [99] [102] [102] [102] [102] [102] [111] [111] [111] [111]

cHex cHex cHex cHex CCl4 cHex Hept cHex CCl4 Hept CCl4 Hept cHex CCl4 Hex cHex CCl4 Hex Hept cHex CCl4 CCl4 Hept CCl4 Hept

40.6 38.5 36.0 37.2 35.5 37.6 37.4 39.0 37.3 34.9 34.0 34.9 35.7 35.8 39.7 36.8 33.7 36.0 33.3 35.2 35.7 35.4 32.5 30.2 32.8

cHex

33.1

CCl4

32.3

Hex Hex

30.1 44.1

[114] [114] [114] [259] [259] [114] [80] [82, 117] [122] [80] [122] [80] [82, 117] [122] [119] [114, 259] [259] [119] [80] [81, 82, 117] [81, 122] [122] [80] [260] [79, 80, 128, 129, 261, 262] [81, 82, 117, 263] [79, 81, 122, 262, 263] [119] [119]

3-Aminopyridine 4-Methylpyridine Isoquinoline 3-Methylpyridine trans-4-Styrylpyridine 2-Aminopyridine cis-4-Styrylpyridine 2-Methylpyridine 2,4,6-Trimethylpyridine 6-Methylquinoline Pyridine

cis-3-Styrylpyridine trans-3-Styrylpyridine

Ref.

(Continued)


289


Lewis base

Solvent

2-Amino-6-methylpyridine 7-Azaindole 2-Anilinopyridine Quinoline

cHex CCl4 CCl4 Hept CCl4 Hex cHex CCl4 Hept CCl4 Hex Hept Hept CCl4 cHex Hept cHex Hept CCl4 Hept

35.1 21.3 23.4 30.2 33.3 34.3 35.4 33.0 26.8 28.5 31.5 24.7 23.4 23.0 34.3 18.9 17.8 22.2 22.2 12.9

[114] [123] [123] [80] [122, 260] [119] [263] [122, 263] [80] [122] [119] [80, 113] [128] [123] [114] [113] [82, 130] [128] [123] [80]

CCl4 CCl4 CCl4 CCl4 CCl4 Hept

36.1 27.9 29.6 28.5 22.6 18.8

[264] [133] [133] [133] [133] [95]

CCl4 CCl4 CCl4 Hept CCl4 Hept CCl4 CCl4 CCl4

6.3 8.8 9.7 13.4 11.8 12.6 11.3 8.8 11.7

[138] [138, 158] [71] [71] [71] [265] [71] [71] [242]

Hept CCl4 CCl4 Hept Hept Hept Hept CCl4 CCl4 Hept Hept

17.6 13.4 15.5 15.9 26.8 22.2 20.5 15.5 15.9 25.9 20.3

[69] [142] [142] [69] [69] [69] [69] [142] [142] [69] [73]

cis-2-Styrylpyridine 2,6-Dimethylpyridine 6-Chloroquinoline 2-Methylquinoline trans-2-Styrylpyridine 3-Chloropyridine Pyrimidine 2-Dimethylaminopyridine 4-Cyanopyridine 2,2 -Bipyridine Pyrazine 2-Chloropyridine Five-membered N-heterocycles 1-Methylimidazole 5-(3-Chloropropyl)-4-methylthiazole 5-(2-Chloroethyl)-4-methylthiazole 4-Methylthiazole Thiazole N-Methylpyrrole Nitriles Chloroacetonitrile Acetonitrile Propionitrile Butyronitrile Valeronitrile Trimethylacetonitrile N,N-Dimethylcyanamide Ethers Diethyl ether Dimethoxymethane 1,2-Dimethoxyethane Propylene oxide Oxetane Tetrahydrofuran Tetrahydropyran 1,3-Dioxane 1,4-Dioxane 2-Methyltetrahydrofuran 7-Oxabicyclo[2.2.1]heptane

Ref.

(Continued)

290


Table 5.25 (Continued) Lewis base Carbonyl compounds Benzoyl bromide Benzoyl chloride Acetyl chloride Acetyl bromide Benzaldehyde Ethyl trifluoroacetate Ethyl trichloroacetate Ethyl formate Dimethyl carbonate Methyl benzoate Diethyl carbonate Ethyl benzoate Methyl acetate Ethyl acetate Methyl thioacetate Ethyl thioacetate α,α,α-Trifluoroacetophenone 3-Trifluoromethylacetophenone 1-Acetylnaphthalene Isobutyrophenone Propiophenone 2-Acetylnaphthalene Benzophenone 4-Chloroacetophenone Ethyl methyl ketone Isopropyl methyl ketone Acetone tert-Butyl methyl ketone Acetophenone 3-Methylacetophenone 4-Methylacetophenone 4H-Pyran-4-one 4H-Thiopyran-4-one N,N-Dimethyltrichloroacetamide N,N-Dimethylchloroacetamide 3,4-Dichloro-N,N-dimethylbenzamide 4-Chloro-N,N-dimethylbenzamide 2-Chloro-N,N-dimethylbenzamide N,N-Dimethylformamide N,N-Dimethylbenzamide 4-Methyl-N,N-dimethylbenzamide 4-Methoxy-N,N-dimethylbenzamide N,N-Dimethylacetamide

Solvent Hept Hept Hept Hept Hept Hept Hept Hept Hept Hept Hept Hept Hept CCl4 Hept CCl4 Hept Hept Hept Hept Hept Hept Hept Hept CCl4 Hept Hept Hept Hept Hept Hept Hept Hept CCl4 CCl4 CCl4 CCl4 CCl4 CCl4 CCl4 Hept CCl4 Hept CCl4 CCl4 CCl4 Hept CCl4

−∆H◦ 2.9 4.6 3.8 2.9 11.7 3.3 5.2 9.0 9.2 9.6 10.7 10.0 12.1 10.5 12.6 13.2 15.3 3.8 10.9 14.2 10.9 11.3 15.1 10.9 10.1 12.3 14.0 14.4 15.3 14.2 13.4 13.6 14.2 13.4 15.5 10.5 13.8 14.6 15.9 16.7 20.9 15.5 20.3 16.7 18.8 19.2 21.5 16.3

Ref. [70] [70] [70] [70] [70] [70] [70] [70] [70] [70] [70] [70] [70] [144] [70] [144] [70] [70] [70] [70] [70] [70] [70] [70] [59] [70] [70] [70] [70] [70] [70] [70] [70] [146] [146] [148] [148] [150] [150] [150] [70] [148] [70] [150] [150] [150] [70] [148] (Continued)


291


Lewis base

Solvent

N,N-Diethylacetamide

Hept CCl4 Hept Hept Hept Hept Hept CCl4

22.0 20.1 23.1 23.1 24.1 18.8 21.8 18.2

[70] [154] [155] [155] [155] [70] [70] [144]

CCl4 CCl4 cHex CCl4 CCl4 CCl4 CCl4 CCl4 CCl4

18.4 18.0 25.3 16.9 14.2 13.4 11.7 9.2 9.2

[157] [157, 158] [65] [158, 266] [159] [160] [160] [158] [158]

CCl4 CCl4 cHex CCl4 CCl4 CCl4 CCl4 Hept CCl4 CCl4 CCl4 CCl4 CCl4 CCl4 CCl4 CCl4

23.0 21.5 30.5 19.9 19.7 18.8 10.3 12.3 15.5 16.4 13.2 15.0 8.6 10.5 11.3 6.3

[164] [163, 164] [267] [164] [163] [59, 163] [163] [268] [162] [163, 164] [162–164] [163] [163] [164] [164] [164]

CCl4 CCl4 CCl4 CCl4 CCl4 CCl4

23.6 30.1 28.1 27.7 26.8 19.5

[166, 169] [170] [166, 170] [166] [166] [170]

CCl4 CCl4 CCl4 CCl4 CCl4

16.2 22.0 27.0 27.9 35.8

[172] [169] [59] [59] [176]

1-Methylpyrrolidin-2-one 1-Methyl-ε-caprolactam 1-Methyl-2-piperidone Tetraethylurea Tetramethylurea Thionyl compounds Diethyl sulfoxide Tetramethylene sulfoxide Dimethyl sulfoxide N,N-Dimethylsulfinamide Methyl phenyl sulfoxide Diphenyl sulfoxide Sulfolane Diethyl sulfite Phosphoryl compounds Trimethylphosphine oxide Triethylphosphine oxide Hexamethylphosphoric triamide Pentaethylphosphonic diamide Triphenylphosphine oxide Diethyl N,N-diethylphosphoramidate Tri-n-butyl phosphate Triethyl phosphate Trimethyl phosphate Diethyl chloromethylphosphonate Diethyl dichloromethylphosphonate Dichlorophosphoric dimethylamide Ethyl dichlorophosphonate Phosphoryl chloride Pyridine N-oxides Pyridine N-oxide 4-Methoxypyridine N-oxide 4-Methylpyridine N-oxide 3-Methylpyridine N-oxide 2-Methylpyridine N-oxide 4-Chloropyridine N-oxide Miscellaneous oxygen bases 2,4,6-Trimethylbenzonitrile N-oxide N-Benzylidenemethylamine N-oxide Diphenyl selenoxide Triphenylarsine oxide Tricyclohexylarsine oxide

Ref.

(Continued)

292



Lewis base

Solvent

Tri-n-octylarsine oxide Thiocarbonyl bases Thiocamphor Ethylenetrithiocarbonate N-Methyl-ε-thiocaprolactam

CCl4

29.8

[176]

cHex cHex Hept CCl4 Hept CCl4 CCl4 CCl4 CCl4 CCl4 Hept CCl4 CCl4 CCl4 CCl4 CCl4 CCl4

46.0 35.6 43.9 36.8 43.1 32.6 24.3 45.5 41.8 29.1 41.4 43.9 35.6 45.2 46.0 33.9 33.1

[180] [180] [155] [155] [184] [184] [185] [186, 269] [186] [185] [193] [193] [270] [198] [198] [198] [198]

cHex cHex Hept CCl4 Hept CCl4 CCl4 CCl4 CCl4 Hept CCl4 Hept Hept Hept Hept Hept Hept Hept Hept Hept Hept Hept Hept Hept Hept Hept Hept CCl4

19.2 18.0 34.7 34.7 32.6 30.5 34.3 32.2 31.8 35.2 34.6 32.6 28.0 24.3 23.8 23.4 23.0 28.9 30.1 25.9 28.0 27.6 26.4 23.4 21.8 21.3 30.1 30.5

[207] [207] [208] [209] [208] [209] [209] [209] [209] [210–212] [210] [74] [216] [216] [216] [216] [216] [216] [216] [216] [216] [216] [216] [216] [216] [216] [74] [74]

N,N-Dimethylthioformamide Tolnaftate 1,3-Thiazolidine-2-thione 1-Methyl-1,3-thiazolidine-2-thione Tetraethylthiuram disulfide 1,1,3,3-Tetramethyl-2-thiourea 1-Methylimidazoline-2-thione 1-Methyl-1,3-imidazolidine-2-thione 1-Ethyl-1,3-imidazolidine-2-thione 1,3-Dimethyl-1,3-imidazolidine-2-thione 1,3-Diethyl-1,3-imidazolidine-2-thione Thiols, sulfides, disulfides Ethanethiol Thiophenol Methyl sulfide Ethyl methyl sulfide Methyl n-propyl sulfide n-Butyl methyl sulfide n-Decyl methyl sulfide Ethyl sulfide n-Propyl sulfide 4-(Methyl)thioanisole Thioanisole 4-Fluorothioanisole 4-Chlorothioanisole 3-Chlorothioanisole 2-Methoxythioanisole 4-Methoxythioanisole 3-Methylthioanisole 4-(Methylthio) thioanisole 2-Methylthioanisole 2-(Methylthio)thioanisole 4-Bromothioanisole 2-Bromothioanisole 2-Chlorothioanisole tert-Butyl phenyl sulfide

Ref.

(Continued)


293


Lewis base

Solvent

Diphenyl sulfide

Hept CCl4 cHex cHex cHex cHex cHex Hept Hept CCl4 Hept CCl4 Hept CCl4 CCl4 CCl4 CCl4 CCl4 Hept CCl4 CCl4 CCl4 CCl4

22.2 21.8 28.5 33.9 36.4 36.8 42.7 26.5 30.4 28.5 33.6 36.4 31.0 29.7 25.9 24.0 19.2 21.2 19.3 18.7 19.8 26.4 20.9

[74] [74] [217] [217] [217] [217] [217] [211] [211] [211] [211] [218] [211] [218] [218] [220] [218] [220] [212] [220] [220] [220] [220]

Hept CCl4 CCl4 CCl4 CCl4 CCl4 CCl4 CCl4 CCl4 CCl4

47.9 38.5 35.1 30.5 23.0 23.8 20.9 12.1 16.3 10.5

[83] [162] [164] [59] [164] [164] [164] [164] [164] [164]

CCl4 CCl4 CCl4 Hept CCl4 CCl4 CCl4 Hept Hept Hept Hept Hept Hept Hept Hept

46.4 41.0 36.0 36.4 36.4 31.8 29.3 38.5 35.1 32.6 34.3 33.5 30.5 30.1 30.1

[222] [222] [271] [74] [74] [219] [219] [216] [216] [216] [216] [216] [216] [216] [216]

4-Methylphenyl phenyl sulfide 4-Methoxyphenyl phenyl sulfide Bis(4-methylphenyl) sulfide Bis(2-methylphenyl) sulfide Bis(2,6-dimethylphenyl) sulfide trans-2,3-Butene sulfide Trimethylene sulfide Tetrahydrothiophene Pentamethylene sulfide 1,4-Dithiane 1,2-Dithiane 1,3,5-Trithiane Methyl disulfide Ethyl disulfide Isopropyl disulfide tert-Butyl disulfide tert-Butyl methyl disulfide Thiophosphoryl compounds Tri-n-octylphosphine sulfide Triethylphosphine sulfide Hexamethylthiophosphoric triamide Triphenylphosphine sulfide Bis(dimethylamino)chlorophosphine sulfide Triethoxyphosphine sulfide (Chlorodiethoxy)phosphine sulfide Dichloro(dimethylamino)phosphine sulfide (Dichloroethoxy)phosphine sulfide Trichlorophosphine sulfide Selenium bases Tetrahydroselenophene Pentamethylene selenide Methyl selenide tert-Butyl phenyl selenide 1,4-Selenoxane 1,4-Diselenane 4-Dimethylaminoselenoanisole 4-Aminoselenoanisole 4-Methoxyselenoanisole 4-(Methylseleno)selenoanisole 2-Methoxyselenoanisole 4-Methylselenoanisole 3-Methylselenoanisole 3-Methoxyselenoanisole

Ref.

(Continued)

294



Lewis base

Solvent

Ref.

Selenoanisole 2-Methylselenoanisole 4-Fluoroselenoanisole 4-Chloroselenoanisole 4-Bromoselenoanisole 3-Chloroselenoanisole 4-Acetylselenoanisole 4-(Ethoxycarbonyl)selenoanisole 3-Bromoselenoanisole 4-(Methoxycarbonyl)selenoanisole 2-Chloroselenoanisole 2-Bromoselenoanisole Bis(4-methoxyphenyl) selenide Bis(4-methylphenyl) selenide Bis(4-chlorophenyl) selenide 1-Methyl-1,3-imidazolidine-2-selenone 1-Ethyl-1,3-imidazolidine-2-selenone 1,3-Dimethyl-1,3-imidazolidine-2-selenone 1,3-Diethyl-1,3-imidazolidine-2-selenone Hexamethylselenophosphoric triamide Triethylphosphine selenide Triphenylphosphine selenide Triethylselenophosphate Arsine Triphenylarsine Isothiocyanates Allyl isothiocyanate Ethyl isothiocyanate Cyclohexyl isothiocyanate

Hept Hept Hept Hept Hept Hept Hept Hept Hept Hept Hept Hept CCl4 CCl4 CCl4 CCl4 CCl4 CCl4 CCl4 Hept CCl4 CCl4 CCl4

26.8 30.5 27.2 27.2 27.2 25.5 24.7 25.1 25.5 25.5 24.7 23.8 31.8 28.9 17.6 33 43 44 46 53.6 49.0 38.5 30.1

[216] [216] [216] [216] [216] [216] [216] [216] [216] [216] [216] [216] [271] [271] [271] [225] [225] [225] [225] [226] [162] [272] [162]

CCl4

39.3

[230]

CCl4 Hept CCl4

15.1 16.2 14.6

[231] [231, 232] [232]

upon complexation and basicity scales [245, 254, 255] or affinity scales. However, because of an insufficient number of reliable and diverse thermodynamic and spectroscopic measurements at that time, these correlations did not lead to firm conclusions on the utility of infrared spectroscopic scales for ordering halogen-bond acceptors quantitatively. One exception was the significant correlation (r = 0.95) between the ∆ν(I CN) shift and the diiodine affinity of 41 diverse Lewis bases [37]: a spectroscopic scale of soft basicity was proposed from this correlation. Since that time, additional spectroscopic measurements have been performed under standard conditions on an extended sample of bases [65, 274]. They are collected in Table 5.26. These are: r The ν(I CN) frequency shifts upon complexation of iodine cyanide with about 300 bases. The complex B· · ·ICN is formed in dichloromethane (because of the low solubility of ICN in alkanes and CCl4 ) at 15 ◦ C. Under these conditions, the carbon–iodine stretch of free iodine cyanide is at 485 cm−1 and the frequency shifts span over 100 cm−1 , from benzene (5 cm−1 ) to quinuclidine (107 cm−1 ). The ν(I· · ·B) vibration is found at


295

Table 5.26 Spectroscopic scales (cm−1 ) of halogen-bond basicity. Lewis base Alkylbenzenes Hexamethylbenzene Pentamethylbenzene 1,2,4,5-Tetramethylbenzene 1,3,5-Trimethylbenzene Aniline (π complex) 1,3-Dimethylbenzene 1,4-Dimethylbenzene tert-Butylbenzene Toluene Benzene Thiophene and furans Thiophene 2-Methylfuran Furan Alkynes 1-Hexyne 2-Hexyne 3-Hexyne Alkenes 3,3-Dimethyl-1-butene cis-2-pentene Cyclohexene 2,3-Dimethyl-1-butene 1-Methyl-1-cyclohexene 2,3-Dimethyl-2-butene Amines Ammonia n-Butylamine Isopropylamine Hexylamine tert-Butylamine sec-Butylamine Ethylamine 2,2,2-Trifluoroethylamine 2-Chloroethylamine Benzylamine Ethylenediamine Aniline (N complex) Piperidine Diethylamine Piperazine Dimethylamine Morpholine (N complex) Thiomorpholine (N complex) Bis(2-chloroethyl)amine Quinuclidine N-Methylpiperidine

∆ν(I CN)

∆ν(I I)

∆ν(I Cl)

∆ν(π g → σ u )

14.5 13.6 11.9 11.6 11.3 9.1

10.4

1070

8.0 8.3

36.4 30.5 25.7 23.2

5.3

19.9 19.0

660

3.8 4.3

16.4 13.4

580 450

840 810

6.9 4.7 7.3 5.5 4.4

9.8

12.5 17.5 18.0 13.3 18.1 20.2 21.2 21.3 22.7 54.4 69.5 72.7 73 75.5 80.2 35.6 59.5 62.5 72.7 43.8 97.9 90.1 82.8 80.8 86.8 58.6 107.2 96.3

34 32.2 34

39.5 39 37

(Continued)

296


Table 5.26 (Continued) Lewis base Diazabicyclooctane N,N-Dimethylcyclohexylamine Trimethylamine Triethylamine Quinuclidinone (N complex) Hexamethylenetetramine 3-Chloroquinuclidine Tri-n-butylamine Tri-n-propylamine N-Methylmorpholine (N complex) N,N-Diethyl-2-chloroethylamine N,N-Dimethylaniline (N complex) Pyridines 4-Dimethylaminopyridine 4-Aminopyridine 3-Dimethylaminopyridine 3,4-Dimethylpyridine 3,5-Dimethylpyridine 4-tert-Butylpyridine 4-Methoxypyridine 4-Ethylpyridine 2,4-Dimethylpyridine 3-Ethylpyridine 4-Methylpyridine Isoquinoline 2-Vinylpyridine 3-Methylpyridine 2-Aminopyridine 2-Methylpyridine 2,4,6-Trimethylpyridine 6-Methylquinoline Pyridine 2-Ethylpyridine Quinoline 2,6-Dimethylpyridine 4-Acetylpyridine (N complex) 3-Iodopyridine 2-Methylquinoline 4-Chloropyridine 3-Bromopyridine 3-Chloropyridine 3-Trifluoromethylpyridine Pyrimidine 4-Trifluoromethylpyridine 4-Cyanopyridine 2-Phenylpyridine Pyrazine 3-Cyanopyridine 3,5-Dichloropyridine

∆ν(I CN)

∆ν(I I)

∆ν(I Cl)

∆ν(π g → σ u )

94.8 95.2 99.3 95.7 81.9 81.4 89.8 99.2 97.3 76.8 80.3 54.5 91.1 78.2 74.2 66.6 69.1 65.2 68.5 64.6 59.1 65.5 66.3 63.2 66.4 65.1 80.0 61.2

5680 5280 5170 4970 4920 4870 4920

33.7

4700 29.6 29.1

4830 4700 4730

29.1

4540 4230 57.1

61.4 60.4 56.8 56.9 52.4 52.2

27.4 27.1

62.6

4560 4450 4380 4110 4000

61 50.1 50.3 43.3 43.2 44.0 44.5 44.4 38.9 39.8

21.5 19.3

16.5

4110 3930 3920 3670 3490 3780 3620 3440 3440 3390 3380 (Continued)


297

Table 5.26 (Continued) Lewis base 2-Bromopyridine 2-Chloropyridine 2-Fluoropyridine 1,3,5-Triazine 2-Cyanopyridine 2,6-Difluoropyridine Five-membered N-heterocycle 1-Methylimidazole Nitriles Trichloroacetonitrile Chloroacetonitrile 3-Chlorobenzonitrile 4-Chlorobenzonitrile Methylthiocyanate 3-Trifluoromethylphenylacetonitrile Acrylonitrile 3-Chlorophenylacetonitrile Benzonitrile Phenylacetonitrile Acetonitrile 4-Methylphenylacetonitrile 2-Methylbenzonitrile 4-Methylbenzonitrile Propionitrile Butyronitrile Isobutyronitrile Valeronitrile Trimethylacetonitrile 4-Methoxybenzonitrile 4-Dimethylaminobenzonitrile N,N-Dimethylcyanamide N,N-Diethylcyanamide Ethers tert-Butyl methyl ether Diethyl ether tert-Butyl ethyl ether Di-n-propyl ether Di-n-butyl ether Diisopropyl ether Dibenzyl ether Dimethoxymethane 1,2-Dimethoxyethane Propylene oxide Oxetane Tetrahydrofuran Tetrahydropyran 1,3-Dioxane 1,4-Dioxane 1,3,5-Trioxane

∆ν(I CN) 36.1 35.3 30.8 26.8 32.1 18.1 66.8

∆ν(I I)

13.2 13.6 13.7 13.9 14.7 13.9 14.9 15.1 17.8 20.5 22.0 18.9 19.4 18.3 18.9 24.0 18.1 16.4 15.1 15.5 16.5 23.5 22.4 21.3 14.6 16.3 8.9

∆ν(π g → σ u ) 3080 3240 2920

15.8

2600

31

8.6 10.3 12.6 12.7 11.7 12.1 10.6 11.9 13.0

∆ν(I Cl)

4.1

4.3

7.5

5.5 6.0 5.4 3.5 4.9 6.4 6.5 4.5 4.9 3.0

6.9 11.9 15.0 14.1 15.7 14.3 15.1 15.3 17.1 16.1 16.1 16.8

1260 1410 1470 1360 1450 1410 1590 1480 1610

18.3 16.8

1610 1570

17.3

1570

19.2

1700

25.4 26.8

2220 2370

21.1 21.2 22.5 21.7 22.7 19.7 20.0 16.4 20.4 19.8 24.7 23.5 24.1 16.9 19.2 11.1

1950

1720 1730 1920 2390 2280 2300 1820 2020 1340 (Continued)

298


Table 5.26 (Continued) Lewis base 2-Methyltetrahydrofuran 7-Oxabicyclo[2.2.1]heptane 1,4-Dimethoxybenzene Anisole 1,4-Benzodioxane 3-Chloroanisole Ethoxybenzene Xanthene Bis(2-chloroethyl) ether 1,3-Dioxolane 2-Chloroethyl ethyl ether Dimethyl ether 2,5-Dimethyltetrahydrofuran 2,2,5,5-Tetramethyltetrahydrofuran Carbonyl compounds Benzoyl chloride Acetyl chloride Acetyl bromide Formaldehyde Acetaldehyde Cyclohexanecarboxaldehyde Isobutyraldehyde 4-Cyanobenzaldehyde Benzaldehyde 4-Methylbenzaldehyde 4-Methoxybenzaldehyde Ethyl trifluoroacetate Ethyl chloroformate Ethyl trichloroacetate Methyl formate Ethyl formate Dimethyl carbonate Methyl benzoate Diethyl carbonate Ethyl methyl carbonate Ethyl pivalate Diethyl oxalate Ethyl benzoate Methyl acetate γ -Butyrolactone Ethyl isobutyrate Methyl cyclohexanecarboxylate Methyl cyclopropanecarboxylate Ethyl propionate Ethyl acetate Ethyl cyclopropanecarboxylate Phenyl formate Ethylene carbonate Propylene carbonate

∆ν(I CN)

∆ν(I I)

∆ν(I Cl)

∆ν(π g → σ u )

23.6 23.4 9.1 7.4 7.3 4.4 7.8 6.8 5.4 12.9 14.0 17.5 23.9 24.4

7.0 7.4

24.4 25.9 13.3 11.0

2310 2410

12.7

8.2 11.1 14.6 16.3 19.0 25.6

3.5 4.0

12.9 11.4 14.7 16.1 18.0

8.8 9.2 6.1 9.2 7.1

2.2 1.6

9.9 9.3 12.5 15.7 10.4 10.2 12.0 11.3 13.7 10.9 8.4 10.2 10.7

2.2 3.6

1030 1700 1350 2310

7.0 5.4

700 730 730

15.2 18.8 17.1 14.4 19.3

1650 1510

23.8 3.5 9.6 4.6 11.4 13.0 8.1, 15.7 10.6 9.5, 16.7 9.1, 16.7 11.2, 17.8 11.3, 17.7 16.5 18.7 11.6, 17.3 11.6, 17.9 11.4, 19.3 16.9 12.0, 19.2

1540 1690 900 650 1260 1360 1300 1260 1330 1300 1400 1220 1330 1550 1490 1470 1470 1490 1570 1590

12.8 13.8 (Continued)


299

Table 5.26 (Continued) Lewis base S-Ethyl methyl thiocarbonate Ethyl thioacetate α,α,α-Trifluoroacetophenone 1,1,1-Trichloroacetone 1,1,1-Trifluoroacetone 4-Trifluoromethylacetophenone 3-Trifluoromethylacetophenone 1-Acetylnaphthalene Isobutyrophenone 3-Chloroacetophenone 3-Bromoacetophenone Propiophenone Benzophenone 4-Chloroacetophenone 4-Fluoroacetophenone Cyclobutanone Diethyl ketone Ethyl methyl ketone Isopropyl methyl ketone Acetone tert-Butyl methyl ketone Acetophenone 3-Methoxyacetophenone 3-Methylacetophenone 1-Adamantyl methyl ketone 4-Methylacetophenone Cyclopentanone 4-Methoxyacetophenone Cyclohexanone 4H-Pyran-4-one (γ -pyrone) Methyl cyclopropyl ketone Dicyclopropyl ketone Xanthone α-Tetralone 1-Indanone Flavone 2,6-Dimethyl-γ -pyrone 3-Chloro-5,5-dimethylcyclohexenone 3-Methyl-5,5-dimethylcyclohexenone 3-Methoxy-5,5-dimethylcyclohexenone 3-Ethoxy-5,5-dimethylcyclohexenone 5,5-Dimethylcyclohexenone 2,2,4,4-Tetramethylpentan-3-one 2-Acetylnaphthalene tert-Butyl phenyl ketone Dimethylcarbamoyl chloride N,N-Dimethyltrichloroacetamide N,N-Dimethyltrifluoroacetamide N,N-Dimethylethylcarbamate

∆ν(I CN) ∆ν(I I) 14.4

4.7

∆ν(π g → σ u )

42.0 49.4 4.8 6.8 4.7

1550 2130 730 820 830 1410 1440 1640 1330 1530 1480 1480 1440 1590 1720 1810 1720 1860 1850 1850 1810 1650 1600 1700 1900 1840 2110 1980 1930

22.2 13.1, 19.7

10.9 11.7 13.9

∆ν(I Cl)

3.5 4.4

11.3, 21.8 19.1 19.8

13.3 16.9 16.5 17.5

5.2 4.9 5.0 5.1 4.6

21.5 21.8 20.3 23.4 21.6 23.3, 33.5

5.9 18.6 14.9 14.7 18.3 19.1 26.0 32.9 17.0

15.0

27.8 28.3 21.0

12.4 11.1 10.5 18.2

3.6 5.3

22.8 25.6 23.2 32.6 23.0 15.5, 24.2 17.9, 25.3 15, 25.4 26.8 36.4 42.6 23.3 29.7 25.6 10.9, 19.1 23.4 12.0, 20.3 15.0 12.4, 18.7 12.1, 19.7 21.4

1940 1750

3180

1400 1760 1380 1420 1410 2110 (Continued)

300


Table 5.26 (Continued) Lewis base N,N-Dimethylmethylcarbamate N,N-Dimethylformamide N,N-Dimethylpivalamide N,N-Dimethylbenzamide N,N-Dimethylacetamide N,N-Diethylacetamide Tetraethylurea Tetramethylurea N,N-Dimethyl-N,N -propyleneurea Diphenylcarbamoyl chloride Phenyl diphenylcarbamate Methyl diphenylcarbamate Ethyl diphenylcarbamate N,N-Diphenylformamide N,N-Diphenylisobutyramide N,N-Diphenylbenzamide N,N-Dimethylisobutyramide N-Methylformanilide N,N-Dimethylpropionamide N,N-Diethylformamide N-Formylpiperidine Sulfinyl compounds Di-n-butyl sulfoxide Tetramethylene sulfoxide Dimethyl sulfoxide Methyl phenyl sulfoxide Diphenyl sulfoxide Diethyl sulfite Dimethyl sulfite Dibenzyl sulfoxide Bis(4-methylphenyl) sulfoxide Bis(4-chlorophenyl) sulfoxide Ethylene sulfite Sulfonyl compounds Diphenyl sulfone Dimethyl sulfone Tetramethylene sulfone Di-n-butyl sulfone N,N,N ,N -Tetraethylsulfamide Phosphoryl compounds Trimethylphosphine oxide Triethylphosphine oxide Tri-n-octylphosphine oxide Hexamethylphosphoric triamide Triphenylphosphine oxide Tri-n-butyl phosphate Triethyl phosphate Trimethyl phosphate

∆ν(I CN)

∆ν(I I)

∆ν(I Cl)

∆ν(π g → σ u )

17.5 26.8 21.8 26.6 28.1 33.6 28.3 27.5

4.9 9.6

2270 2600

9.2 10.2 11.3 10.5 10.3

20.2 36.1 25.3 32.4 34.5 36.0 36.4 34.4

6.5

13.2 15.9 16.7 18.2 27.5

1910

8.1 8.7 10.6 10.0

26.7 25.3 29.2 25.0 37.1 38.3

2290 2480 2720 2830

9.9 12.0 13.2 13.6 18.3 19.1 23.1 23.9 24.1 27.8 28.2 33.7 30.8 29.2 27.2 25.3 11.1 9.5 28.1 26.8 24.0 9.6 8.3 8.5 9.6 11.1 11.5 34.7 37.6 37.7 36.2 29.9 23.7 22.8 21.9

2480 2730 2830 2600 2620 2950

13.3 13.4 9.4 3.8

31.8

3.5

16.2 17.6 15.3 12.1 7.7 7.0 5.6

42.8 42.7 35.4 26.9 26.2 23.5 (Continued)


301

Table 5.26 (Continued) Lewis base Tripiperidinophosphine oxide Chlorodiphenylphosphine oxide Triphenyl phosphate Dichlorophenylphosphine oxide Miscellaneous oxygen bases Nitromethane Pyridine N-oxide 4-Methylpyridine N-oxide Triphenylarsine oxide Thiocarbonyl bases Ethylenetrithiocarbonate N,N-Dimethylthioacetamide N,N-Dimethylthioformamide N,N-Dimethylthiocarbamoyl chloride Tetraethylthiuram disulfide Tetramethylthiuram disulfide 1,1,3,3-Tetramethyl-2-thiourea Sulfides, disulfides Methyl sulfide Ethyl sulfide Isopropyl sulfide n-Butyl sulfide Thioanisole 4-Bromothioanisole Diphenyl sulfide 1,4-Thioxane Trimethylene sulfide Tetrahydrothiophene Pentamethylene sulfide 1,4-Dithiane 1,3,5-Trithiane 1,3-Dithiane Methyl disulfide Ethyl disulfide Thiophosphoryl compounds Tri-n-octylphosphine sulfide Triethylphosphine sulfide Trimethylphosphine sulfide Triphenylphosphine sulfide Triethoxyphosphine sulfide Trimethoxyphosphine sulfide (Chlorodiethoxy)phosphine sulfide (Dichloroethoxy)phosphine sulfide Trichlorophosphine sulfide Selenium bases Ethyl selenide Methyl selenide Triphenylphosphine selenide

∆ν(I CN)

∆ν(I I)

∆ν(I Cl)

∆ν(π g → σ u )

39.5 20.6 16.1 12.5

37.0

20.4 23.9 27.0

7.2 45.0

46.7 73.1 70.0 44.9 51.5 47.0 73.8 55.2 59.0 59.6 59.7 39.9 28.0 27.0 55.3 61.1 58.8 46.6 33.0 46.5 29.8 31.2 64.4 63.2 61.4 51.6 36.1 33.2 20.0 12.0 10.9 75.4 70.3 70.2

3110 3890 2440 2510 3720 37.5 40.0 31.0 27.0 25.9 33.0 38.0 44.5 40.5

70.3 78.7

3570 3800 3033 2240

74.7 77.0

22.2 23.7

3460 3640 3790 3370

2530 3950 3210 2680

27.1

5.5

93 (Continued)

302


Table 5.26 (Continued) Lewis base Triethyl selenophosphate Haloalkanes 2-Chloropropane 2-Chloro-2-methylpropane Bromoethane 2-Bromopropane 2-Bromo-2-methylpropane Bromocyclohexane Iodomethane Iodoethane 2-Iodopropane 2-Iodo-2-methylpropane Iodocyclohexane Isothiocyanates Ethyl isothiocyanate n-Butyl isothiocyanate Arsines, phosphines Triphenylarsine Triphenylphosphine Diphenylmethylphosphine Triethylphosphine

∆ν(I CN)

∆ν(I I)

∆ν(I Cl)

∆ν(π g → σ u )

50.8 8.8 9.2 10.7 11.4 12.6 11.9 14.7 16.6 18.0 22.3 20.6 22.1 22.6

14.0

1670 1620

52.0 67.4 74.1 79.6

very low frequencies (between 80 and 110 cm−1 for the strong complexes of ICN with pyridines [277]) so that the coupling between ν(I CN) and ν(I· · ·B) can be neglected in the interpretation of the ν(I CN) shifts. r The ν (I I) frequency shifts upon complexation of diiodine with about 100 Lewis bases. The complex is formed in cyclohexane at 20 ◦ C. Although the diiodine stretch cannot be observed by IR spectroscopy, since it does not produce a dipole moment variation, the polarization of the I I bond by complex formation renders the vibration IR active. The frequency shifts are calculated from the free diiodine value of 210 cm−1 measured by Raman spectroscopy. The highest frequency shift is found for piperidine (39.5 cm−1 ). In the interpretation of these results, it must be noted that the normal coordinate describing ν(I I) is expected to mix with ν(B· · ·I) because the two bands are generally close. Hence, the frequency shift ν(I I) is not simply related to the change in force constant of the I I bond upon complexation. r The ν(I Cl) frequency shifts of iodine monochloride with about 200 Lewis bases. The complex B· · ·ICl is formed in heptane at 20 ◦ C. Many ICl complexes cannot be studied because of chemical reactions and/or ionization of the complex and/or lack of solubility in heptane. Hence the sample of bases is mainly limited to nitriles, aromatic π bases and oxygen bases. With carbonyl bases, two bands of complexes are encountered when both the perpendicular and planar complexes [42] exist in significant proportions. The frequency shifts are calculated from the free I35 Cl value at 376 cm−1 in heptane. They range from 3.5 cm−1 for CF3 COOEt to 93 cm−1 for Me2 Se. In strong complexes, the ν(B· · ·I) band becomes close to, and may be significantly coupled with, the ν(I Cl)


303

4

ICN basicity

oxygen bases 3 Pyridines 2

Ph3PS Et2S

1 0 0

20

40

60

80

∆ν (I-CN) / cm

-1

Figure 5.10

ICN basicity plotted against values of ∆ν(I CN).

band. Hence the results must be interpreted with caution because the approximation of the complexed ICl molecule to a slightly perturbed diatomic molecule may not be valid. The relationships between the IR spectroscopic shifts ∆ν(I CN), ∆ν(I Cl) and ∆ν(I I) and ICN basicity, ICl basicity and I2 basicity, respectively, are shown in Figures 5.10–5.12. It appears that logK values (found in the references in Table 5.23 for ICN and ICl complexes and in Tables 5.6–5.22 for I2 complexes) exhibit fairly good relationships with the corresponding ∆ν values (Table 5.26) if the comparisons are restricted to families of bases with similar halogen-bond acceptor sites. The ICN, ICl and I2 complexes show similar separations into families; the following families can be distinguished from left to right in plots of logK versus ∆ν: r Oxygen bases (PO, CO, NO, SO, AsO and single-bonded oxygen bases) gather with nitriles in one family. However, in the plot of diiodine basicity versus ∆ν(I I), phosphoryl bases form a family different from that of other oxygen bases.

5 oxygen bases nitriles

ICl basicity

4 3 2

π bases

1 0 -1 0

20

40

60

∆ν(I-Cl) / cm-1

Figure 5.11

ICl basicity plotted against values of ∆ν(I Cl).

304

Lewis Basicity and Affinity Scales 5 Nsp2, Nsp3 bases

4 PO bases Diiodine basicity

3

O bases nitriles

Sulfides

2 PhNMe2 (EtO)3PS

1 π bases

0

EtNCS

-1 Cl3PS

-2 0

10

20

30

40

50

∆ν (I-I) / cm

-1

Figure 5.12

Diiodine basicity plotted against values of ∆ν(I I).

r Aromatic π bases. r sp2 and sp3 nitrogen bases. r Single-bonded sulfur bases. The individual regression equations logK = a∆ν + b are assembled in Table 5.27 together with the r and s measurements of the goodness of the statistical fits. Table 5.27 Correlation equations between ∆ν (cm−1 ) and logK (K in l mol−1 ). Family of bases

a

b

r

s

n

ICN basicity versus ∆ν(I CN) Oxygen bases Pyridines

0.119 ± 0.015 0.055 ± 0.003

−1.5 ± 0.4 0.927a −1.7 ± 0.2 0.992 ICl basicity versus ∆ν(I Cl)

0.3 0.05

12 8

Oxygen bases, nitriles Aromatic π bases

0.122 ± 0.005 0.069 ± 0.006

−1.2 ± 0.1 0.968 −1.2 ± 0.1 0.981 I2 basicity versus ∆ν(I I)

0.3 0.1

38 8

Phosphoryl bases Oxygen bases, nitriles Aromatic π bases Pyridines, primary and secondary amines, N-methylimidazole Sulfides, disulfides

0.164 ± 0.008 0.171 ± 0.008 0.104 ± 0.016 0.135 ± 0.004

0.05 ± 0.10 −0.8 ± 0.1 −0.9± 0.1 −1.5 ± 0.1

0.993 0.956 0.935 0.992

0.1 0.2 0.1 0.1

8 43 8 16

0.107 ± 0.005

−2.2 ± 0.2

0.990

0.1

11

a

The measurements of K in different solvents (heptane, CCl4 , CS2 ) may explain the low correlation coefficient.


305

New complexation constants can be calculated from these correlation equations and the frequency shifts of Table 5.26. Examples: 3-dimethylaminopyridine N-formylpiperidine N,N-diethylformamide 2,6-dimethyl-γ -pyrone

∆ν(I ∆ν(I ∆ν(I ∆ν(I

I) = 33.7 cm−1 I) = 10.0 cm−1 I) = 10.6 cm−1 I) = 15.0 cm−1

pK BI2 (calc.) pK BI2 (calc.) pK BI2 (calc.) pK BI2 (calc.)

= 3.07 = 0.94 = 1.04 = 1.79

For bases with more than one site of halogen bonding, the site of major basicity can be identified by the presence of the base in a given family. For example, in Figure 5.12, thiophene and 2-methylfuran are members of the π base family whereas MeCOSEt falls on the carbonyl family line. Since Figure 5.12 discriminates between oxygen, π and sulfur bases, the presence of thiophene and 2-methylfuran in the π base family and that of MeCOSEt in the carbonyl family offers evidence (albeit indirect) of the fixation of diiodine on the π clouds of thiophene and 2-methylfuran and on the oxygen of MeCOSEt. Moreover, since MeSCN falls on the nitrile family line in Figure 5.11, this indicates that ICl complexing occurs on nitrogen and not on sulfur. The relationships between the IR spectroscopic shifts and the affinity scales are more difficult to study. There is a lack of reliable ICN and ICl affinity values. The diiodine affinity values are numerous and diverse (Table 5.25), but their reliability is difficult to assess. The largest set of data is found for the comparison of diiodine affinity and ∆ν(I CN) scales. A correlation has previously been claimed for 41 bases [37]. A careful distinction between affinity values measured in heptane and in CCl4 and the extension of the ∆ν(I CN) scale to many new bases [65] enable the enthalpy–frequency shift relationships 5.19 and 5.20 to be proposed: −∆H ◦ (BI2 ) in heptane= 0.413 (±0.018) ∆ν(I−CN) + 8.0 (±0.9) n = 77, r = 0.936, s = 4.4kJ mol−1

(5.19)

−∆H ◦ (BI2 ) in CCl4 = 0.475 (±0.018) ∆ν(I−CN) + 5.2 (±0.7) n = 71, r = 0.954, s = 3.2kJ mol−1 .

(5.20)

The relationship 5.20 is shown in Figure 5.13. It is difficult to assess the relative importance of model and experimental errors in these correlations. However, it is clear that the ∆H(BI2 )–∆ν(I CN) correlation is less family dependent than the pK BI2 –∆ν(I I) and pK BICN –∆ν(I CN) relationships shown in Figures 5.12 and 5.10. Hence the relationships 5.19 and 5.20 may support the use of ∆ν(I CN) as a spectroscopic scale of soft affinity. Indeed, diiodine is the archetype of soft Lewis acids in the Pearson classification since it has a very low absolute hardness (η = 3.4 eV). Moreover, ∆ν(I CN) values obey the HSAB principle (soft acids prefer soft bases) since they decrease with the absolute hardness of the donor atom (in a given column of the periodic table), as shown in Table 5.28. The comparison with other scales claimed to represent donor properties of soft solvents is summarized in Table 5.29. The µ scale [278] (µ for malakos = ‘soft’ in Greek) is defined as the difference between the Gibbs energies of transfer of hard sodium and potassium ions from water to a given basic solvent and the corresponding quantity for soft silver ions.

306


Diiodine affinity / kJ mol

-1

60 50 40 30 20 10 0 0

10

20

30

40

50

60

70

80

90

∆ν (I-CN) / cm

-1

Figure 5.13 Correlation between the diiodine affinity and the frequency shift of the I CN stretching upon halogen bonding to 71 Lewis bases.

The SP (softness parameter) scale [279, 280] is derived from the Gibbs energy of transfer of Ag+ from benzonitrile to a given basic solvent. The Ds (donor strength) scale [281] is based on the change in the stretching vibration frequency of HgBr2 in going from the gas phase to the solute state in a given basic solvent. There are acceptable correlations between ∆ν(I CN) and these three scales. The correlation with the Ds scale is shown in Figure 5.14. 5.5.2

The Blue Shift of the Diiodine Visible Band

The electronic configuration and the state of the outermost electrons in the ground state of the diiodine molecule are given by 2 4 . . . σg 5p (πu 5p)4 πg 5p , 1 Σg+ The lowest energy excited configuration and states are − + 2 3 . . . σg 5p (πu 5p)4 πg 5p (σu 5p) ,3 ,3 ,3 ,3 ,1 2u

1u

0u

0u

u

Table 5.28 Comparison of frequency shifts ∆ν (I CN) (cm−1 ) and hardness η (eV). Donor atom

η

Lewis base

∆ν

Cl Br I

4.70 4.24 3.70

i-PrCl i-PrBr i-PrI

8.8 11.4 18.0

O S Se

6.08 4.12 3.86

Et2 O Et2 S Et2 Se

19.4 59.0 75.4

Lewis base

Ph3 PO Ph3 PS Ph3 PSe

∆ν

29.9 51.6 70.2


307

Table 5.29 Correlations between thermodynamic and spectroscopic scales of soft basicity and ∆ν(I CN). Scales µ SP Ds

r

n

Sample of bases

0.747 0.939 0.831 0.969

18 8 41 37

Oxygen, nitrogen, sulfur bases Aniline, nitriles, thiocarbonyl, sulfide C, O, S, N, P, As bases NH3 , Et3 P, NEt3 and piperidine excluded

A higher excited state is + 4 . . . σg 5p (πu 5p)4 πg 5p (σu 5p) ,1 u

The visible band of diiodine at 520 nm in the solvent heptane is assigned [282] +occurring 3 π → 5p → σu 5p transition, which may borrow the intensity to a forbidden 1 + g g 0u + + 1 1 from the allowed g → u transition as a result of the strong interatomic spin–orbit coupling in the heavy iodine atoms. When a Lewis base is added to a diiodine solution, two pronounced changes in the spectrum occur. They can be schematically described by the approximate MO diagram of Figure 5.15. The visible transition (arrow 1 in free diiodine, arrow 2 in the complex) undergoes a blue (hypsochromic) shift and a new band arises in the ultraviolet region that is due to a charge-transfer transition (arrow 3). In medium and weak complexes, the visible band of the complex is overlapped by that of free diiodine. Methods have been developed [38, 39] for measuring accurately and easily the blue shift, hereafter called ∆ν(π g → σ u ). The results of these measurements [38, 39, 65] under standard conditions of solvent (heptane) and temperature (15 ◦ C, in order to increase the intensity of the band of the complex) are presented in the last column of Table 5.26 for π bases, pyridines, nitriles, ethers, carbonyls and sulfur bases. 80 Et3P

70

NH3

Ds / cm

-1

60 50 40 30

NEt3

20 10 0 0

20

40

60

80

100

∆ν (I-CN) / cm-1

Figure 5.14 Correlation between the Ds scale and the spectroscopic scale of soft affinity, ∆ν(I CN): Ds = 0.53∆ν(I CN) + 8 (n = 37; NH3 , Et3 P, NEt3 and piperidine excluded).

308


σu (2)

(1) Energy

πg (3)

n

Diiodine

Complex Lewis base

Figure 5.15 Some MO energy levels in a complex of diiodine with an n donor; n refers to the donor orbital of the base. The MOs in the complex have a strong σ u or n character because of the weak interaction energy.

A rigorous evaluation of the correlation between ∆ν(π g → σ u ) and diiodine affinity, with the accurate spectroscopic data of Table 5.26 and the selected thermodynamic data of Table 5.25, indicates that only 83% of the variance of 45 enthalpies measured in CCl4 and 82% of the variance of 72 enthalpies measured in heptane are explained by ∆ν(π g → σ u ). The quantitative family-independent correlation claimed by a number of workers [282] is denied by this careful evaluation. Much better correlations are found between ∆ν(π g → σ u ) (Table 5.26) and diiodine basicity (Tables 5.6–5.22), but they are family-dependent. The statistical parameters for the equations pK BI2 = a∆ν(π g → σ u ) + b are given in Table 5.30. These correlations will

Table 5.30 Family-dependent correlations between diiodine basicity and the blue shift of the visible diiodine band. Family Aromatic π bases 3- and 4-substituted pyridines Nitriles Ethers Carbonyl bases Single-bonded sulfur basesa Thiophosphoryl bases a

a × 104

b

r

s

n

12.87 ± 1.16 12.78 ± 0.22 10.14 ± 0.79 12.91 ± 0.69 9.99 ± 0.34 15.82 ± 1.41 18.71 ± 2.74

−1.17 ± 0.09 −3.56 ± 0.10 −1.68 ± 0.14 −2.56 ± 0.14 −1.64 ± 0.06 −3.66 ± 0.49 −3.64 ± 0.91

0.984 0.998 0.982 0.986 0.967 0.977 0.989

0.06 0.07 0.08 0.08 0.14 0.16 0.25

6 17 8 12 61 8 3

Ph2 S was excluded from the correlation owing to its polyfunctionality.


309

be useful for checking uncertain diiodine basicity values, for predicting experimentally inaccessible pK BI2 values and, consequently, for making a comprehensive and accurate scale of diiodine basicity.

5.6

Conclusion

The measurement of thermodynamic and spectroscopic parameters of halogen-bonded complexes provides a number of quantitative scales for measuring halogen-bond basicity (affinity). The diiodine basicity scale pK BI2 is rather general for inorganic halogen-bond basicity (e.g. ICl basicity), but is limited by the type of Lewis bases for organic halogenbond basicity (e.g. ICN basicity). The diiodine affinity scale can be useful for selecting the best computational methods to calculate the energy of the halogen bond. The spectroscopic scales ∆ν(I I), ∆ν(I Cl), ∆ν(I CN) and ∆ν(π g → σ u ) give family-dependent relationships with diiodine basicity. The only (approximate) family-independent correlation is found between ∆ν(I CN) and diiodine affinity. ∆ν(I CN) is therefore proposed as a spectroscopic scale of soft affinity.

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6 Gas-Phase Cation Affinity and Basicity Scales 6.1

Cations as Lewis Acids in the Gas Phase

The study of the formation of complexes between cations and chemical species in the gas phase offers the unique possibility of measuring the intrinsic (i.e. free of solvent and counter ion influences) Lewis basicity of these species. Although the charge-transfer component of the interaction energy in cation adducts is normally smaller than the electrostatic and polarization components, gas-phase cations are generally considered as archetypal Lewis acids because the full positive charge of the naked cation allows strong binding with classical Lewis bases. In fact, almost any entity bearing some electron density, such as the rare gases, can form complexes with cations that may be observable in the gas phase. There is a large array of cations that have been studied in the gas phase for their interaction with Lewis bases (generally neutral organic molecules) and are still being actively investigated today. The most studied is the proton [1–6]. This special cation will not be dealt with in this chapter (however, see Section 1.3 in Chapter 1), since the proton and its transfer serve to define the Brönsted acids and bases, as opposed to Lewis acids and bases, which do not exchange protons. Moreover, although the proton and a cation with an empty orbital may be seen as identical from the point of view of the Lewis theory, the special ‘electronic structure’ of the proton (in fact the absence of any electron) makes the proton affinity and basicity scales completely different, qualitatively and quantitatively, from any other cation scales. A significant part of the studies of the interaction of gas-phase metal cations with organic molecules was devoted to their chemical reactivity: bond insertion, bond cleavage, elimination and so on. This aspect is not considered here, although some of these reactions may be used as indirect routes for the formation of cation/molecule adducts. Apart from the proton, experimental studies have focused primarily on the binding of alkali metal cations and transition metal monocations. Few studies have considered alkaline Lewis Basicity and Affinity Scales: Data and Measurement C 2010 John Wiley & Sons, Ltd


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earth metal cations and, among main group metal cations, only the aluminium monocation appears to have been quantitatively studied against an extended variety of Lewis bases. Among non-metal cations, it is worth mentioning that the halogen cations Cl+ and I+ could be investigated as gas-phase Lewis acids, in spite of their strong electrophilic reactivity. In addition to these simple monatomic cations, a few polyatomic cations containing a metal as acceptor centre, such as a metal cation coordinated to a cyclopentadienyl (e.g. CpNi+ ) or halogen ligand (e.g. BrFe+ ), were probed as Lewis acids. Cations corresponding to protonated bases, BH+ , may function as Brönsted or Lewis acids. In the latter case, they may form adducts by sharing the proton through an ionic hydrogen bond BH+ · · ·B with a second base B . Protonated water, H3 O+ , and protonated ammonia, NH4 + , are typical cases, with the possibility of adding several ligands to form gas-phase ‘solvates’, but organic cations, such as the pyridinium cation, C5 H5 NH+ , were also considered. However, B must be a stronger Brönsted base than B to avoid the full proton transfer from B to B . Other typical organic cations that can form adducts are carbenium ions X3 C+ (X = H, alkyl, halogen, etc.). Note that CF3 + is isoelectronic with the typical Lewis acid BF3 . Organoelement analogues of the carbenium ion (Me3 Si+ , Me3 Ge+ , Me3 Sn+ ) have also been studied as Lewis acids. The research field is, in principle, not restricted to monocharged cations so the thermochemistry of multicharged cations, resulting from the ‘solvation’ (addition of several molecules) of alkaline earth dications (Mg2+ , Ca2+ , etc.), has been studied [7]. Nevertheless, such studies are severely limited by competitive electron transfer, since the second ionization energy of a metal is usually much greater than that of molecules, and the charge exchange reaction M2+ + L → M+ + L+

(6.1)

where M is the metal and L the ligand, may occur, instead of adduct formation. A dication adduct can be observed only when it is bonded to a sufficient number of ligands (as in solution) or to a polydentate ligand, for example a peptide [8]. In general, only the basic properties of neutral ligands will be considered, although, in a few cases, the interaction between a cation and a negatively charged ligand may be mentioned. In principle, Lewis basicity scales are defined from the formation of complexes of 1 : 1 stoichiometry, that is, in the case of cation basicity scales, for the coordination of one cation M+ with one ligand L. Gibbs energy and enthalpy changes have also been obtained for two-ligand complexations of FeBr+ , Cu+ , Co+ , Ni+ and Ag+ , that is for the reaction M+ + 2 L L2 M+

(6.2)

These results may be valuable for coordination chemists, even though they provide less direct information on Lewis basicity than scales based on the formation of the simple LM+ adduct. The above list of cations is summarized in Table 6.1. The samplings of Lewis bases complexed with the cations, together with representative references, are also given. Not all the literature data on the thermochemistry of ion/molecule complexation in the gas phase can be usefully translated into absolute cation Lewis basicity (affinity) scales suitable for broad purposes. Only the data that (i) are absolute or can be translated into absolute values, (ii) refer to the same temperature or can be referred to the same temperature, (iii)

Gas-Phase Cation Affinity and Basicity Scales

325

Table 6.1 Some typical gas-phase cationic Lewis affinity and basicity scales that can be established from literature data. Only leading references are given. The stoichiometry column refers to the type of adduct (one- or two-ligand adduct) that was studied. Cationic Lewis acid Alkali metals Li+ Na+ K+ Cs+ Transition metals Mn+ BrFe+ Co+ Ni+ CpNi+ Ni+ Cu+ Cu+ Ag+ Main group metal Al+ Group 14 cations CH3 +

Stoichiometry

Breadth and variety of the scale

Ref.

1:1 1:1 1:1 1:1

Large (more than 200 ligands) Large Large N and O bases, rare gases and a few anions

This chapter This chapter This chapter [9]

1:1 1:2 1:2 1:1 1:1 1:2 1:1 1:2 1:2

Various families of bases 24 ligands in various families of bases Medium, mostly C, O, S bases Diverse, relatively small ligands Various families of bases Large (C, N, O, S, halogen bases, etc.) Diverse, relatively small ligands Large (C, N, O, S, halogen bases, etc.) 16 ligands (C, N, O, S and Br bases)

This chapter [10] [11] [12, 13] This chapter [12, 14] [13, 15, 16] This chapter [17]

1:1

Large, N and O bases

This chapter

1:1

[18, 19]

(CH3 )3 Si+ (CH3 )3 Ge+

1:1 1:1

(CH3 )3 Sn+ Ionic HB CH3 NH3 + Halogens Cl+ I+

1:1

Diverse ligands (C, N, O, S, halogen bases, etc.) Large, N and O bases 17 acetophenones and 4 nitrogen bases 27 ligands (C, N and O bases)

1:1

Large; various families of bases

This chapter

1:1 1:1

Substituted pyridines Mostly substituted pyridines

[23] [24]

[20] [21] [22]

are available for a sufficient number of bases and (iv) correspond to a sufficient variety of coordination centres and functional groups, are recommended to be used for the construction of profitable cation Lewis basicity (affinity) scales. For these reasons, we have chosen to present in this chapter basicity (affinity) scales relative to the cations Li+ , Na+ , K+ , Mn+ , CpNi+ , Cu+ (two-ligand), Al+ and CH3 NH3 + as indicated in Table 6.1. The covalent part of the interaction of the alkali metal and Al+ cations corresponds mainly to sigma bonding, while Mn+ , Cu+ and CpNi+ are additionally capable of π back-donation from the occupied d orbitals of the metal towards antibonding empty orbitals of the ligand. The latter cation provides a scale of ionic hydrogen-bond basicity. In this chapter, we will follow the definition of affinity and basicity used in the field of gas-phase cation adducts. The affinity and basicity of the ligand are defined as, respectively,

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the enthalpy and Gibbs energy of the dissociation reaction LM+ L + M+

(6.3)

The symbols of affinity and basicity for the cation M+ , where M is the symbol of the element, will be written as MCA and MCB, respectively. To avoid confusion, the full symbol, and not the first letter of the symbol or of the name, will be used. For example, LiCB is preferred to ‘LCB’ and KCB to ‘PCB’.

6.2

Structure of Cation/molecule Adducts

When interpreting Lewis basicity measurements, it is expected that the adduct structure, in particular the site of interaction, can be identified experimentally or otherwise. This question, which seems trivial for simple monofunctional ligands, may be more difficult to answer when the molecular complexity increases, since several potential coordination centres are created. In condensed phases, spectroscopic studies often help to resolve the issue. However, classical spectroscopies are difficult to apply in the gas phase, due to the highly dilute state of the ions [25, 26]. Hence the structure of the ion/molecule adducts, which are just a special case of gas-phase ions, has mostly been conjectured from mass spectrometric results (fragmentation, reactivity, exact mass when available, etc.) and energy-related arguments [27]. Modern radiation sources, in particular the synchrotron radiation in the form of ‘free electron laser’, are emerging as a means for recording vibrational spectra of gas-phase ions [28–30]. Nevertheless, obtaining a detailed structural knowledge of gas-phase ions by this advanced technique remains challenging. Fortunately, today quantum chemical calculations can provide reliable structures for a large number of ion/molecule adducts, even of relatively large size [31]. The following discussion on the structure of representative cation adducts is mostly based on the results of such calculations. A general description of the structures of cation adducts is outside the scope of this section. Only representative structures and structural features peculiar to gas-phase ion/molecule interactions will be presented. Among these, the chelate effect can strongly affect cation basicity scales, as compared with scales established towards neutral Lewis acids. Indeed, when a cation and electron-rich centres are initially in a favourable position, or when flexible chains allow a favourable arrangement, the formation of a chelate structure increases the stability of the adduct. In the following, the structures are discussed according to the nature of the coordination centre. Carbon bases. In the case of π bases (aromatic rings, alkenes, alkynes, etc.), the cation is located above the π-electron cloud, in a direction perpendicular to the plane of the sp2 carbon skeleton from (or close to) the centre of the double or triple bond (C2v symmetry for ethylene and acetylene), or on the C6 axis of the benzene ring. This has been established for the adducts of benzene with some transition metal cations (including Cu+ and Mn+ ) of different charge states [32–34], for Al+ [35–37] [including an infrared (IR) spectroscopic study] [38] and for alkali metal cations [39]. The distortion of the sp2 skeleton is limited


327

Figure 6.1 Geometry of the Li+ /n-butylbenzene adduct. Reprinted from [40] by permission of Wiley–VCH. Copyright 2003.

and, basically, the π bonds tend to stretch (decrease of the bond order). In this context, the aluminium cation was categorized as intermediate between alkali and transition metal cations. Other less symmetric structures of very close energy were proposed for the adduct of Cu+ with benzene [15]. The long alkyl chains on substituted benzenes have been shown [40] to enhance the bonding of Li+ . When the cation forms a bond with the benzene ring, it retains most of its charge, and polarizable groups are attracted towards Li+ , producing a kind of chelate. This effect was called the ‘scorpion effect’ owing to the coiling of the alkyl ‘tail’, illustrated in Figure 6.1. When a second phenyl is appended at the end of the alkyl chain, the interaction with this added aromatic ring also produces an enhancement by a ‘pincer effect’ that sandwiches the lithium cation [41] (Figure 6.2). In heterosubstituted benzenes, such as halobenzenes, aniline or phenol, and in aromatic heterocycles, such as furan and thiophene, cations can bind to the aromatic ring (π complex) and/or to the lone pair(s) on the heteroatom (n complex). Halobenzenes behave mostly like π bases with alkali metal cations and Cu+ . In the series of fluorobenzene adducts with alkali metal cations, the calculated structures are in favour of the π complexes for Li+ , Na+ , K+ , as the stability difference with the n complexes on

Figure 6.2 Structure of the most stable conformer of the Li+ /1,3-diphenylpropane adduct. Reprinted from [41] by permission of Wiley–VCH. Copyright 2006.

328


a

b

Figure 6.3 Most stable complex of Na+ with aniline (a) and closest stable structure (b). Reproduced from [43], copyright 2003, Elsevier.

fluorine decreases with increasing size of the cation [42]. The trend is reversed (only by a small difference in energy) for Rb+ and Cs+ . The copper monocation also forms more stable π complexes with the four halobenzenes [15]. Iodobenzene is perhaps an exception, for which calculations favour a ‘mixed’ π–n complex, with Cu+ above the plane of the ring, but near the iodine atom, with the angle C I Cu+ close to 90◦ . In aniline, the electron-donating resonance effect of the amino group increases the ring π-electron density whereas the electron-withdrawing inductive and resonance effects of the phenyl group decrease the nitrogen lone-pair electron density. Indeed, for the five alkali metal cations, calculations show that the preferred site for binding is the benzene ring, in a position close to the centre [43]. Even so, another conformer with the cation situated close to the nitrogen atom, and still above the mean plane of the molecule, is very close in energy, and supposed to be present in room temperature experiments (Figure 6.3). The geometry of the copper cation adduct of aniline was calculated at DFT and correlated ab initio levels [44]. The authors observed that all methods always favour a π complex, but with different positions of the cation relative to the ring carbons. Although the Cr+ cation does not pertain to the series of reference cations selected in Section 6.1, it is worth mentioning that the ring was shown to be the preferred site of complexation of aniline from IR spectroscopic experiments [45]. The comparison of this experimental result with computational data shows that the MPW1PW91 functional better predicts the energy difference of π and n complexes than B3LYP, in the Cr+ /aniline system [45]. The same conclusion was derived when comparing the B3LYP and MPW1PW91 structures of the complexes of phenol with cations of sodium, magnesium, aluminium and first-row transition metals [46]. Calculations for the ring versus O complexation show: (i) similar binding energies for Na+ (ii) a small differential in favour of the ring (8–12 kJ mol−1 ) for Mg+ and (iii) the ring significantly favoured by more than 20–25 kJ mol−1 for Al+ and transition metal monocations, such as Mn+ , Ni+ and Cu+ . For the aromatic heterocycles furan and pyrrole, the available structures indicate that cations seem to bind preferentially by the whole π system rather than by the electrons of the heteroatom [36, 47–50]. Nitrogen bases. For simple nitrogen ligands, when the axes of the putative sp, sp2 or 3 sp nitrogen lone pairs coincide with the direction of the dipole moment, as in acetonitrile, pyridine or ammonia, the cations bind to the coordinating centre along this axis [51].


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Substitutions that remove this coincidence may affect the orientation of the cation, as shown in the alkali metal cation adducts of the three isomers of aminopyridines [52] and of methylpyridines [53]. It is noteworthy that the 2-amino substituent acts as a second basic site, enhancing the bonding by a sort of chelate effect. In 2-methylpyridine the cations are slightly ‘attracted’ toward the methyl group. There are varied examples of polynitrogen bases that form bidentate cyclic structures with metal cations. The most clear-cut cases of such chelate effects are encountered with polyamines, among which ethylenediamine is the best-studied case. The main studies, concerning the cations of Li, Na, K, Mg, Al and Cu, are cited in ref. [54]. The combined spectroscopic and ab initio study of copper complexes shows bidentate cyclic structures for the four N-methyl-substituted derivatives of ethylenediamine, but a hydrogen bond stabilized monodentate structure for the Cu+ /ethylenediamine is also possible [54]. In azines and azoles, when there are two adjacent nitrogen atoms in the cycle(s) as in pyridazine or 1,8-naphthyridine, the alkali metal cations form a ‘bridge’ between them, to maximize the electrostatic interaction [55–57]. However, Hartree–Fock calculations indicate that Al+ is attached at only one nitrogen, similarly to the open structure of the H+ adduct. This was explained by the more covalent character of the Al+ and H+ bonding [58]. The chelate structures of the Li+ adducts of histamine and 2-(β-aminoethyl)pyridine have also been calculated [59]. The flexible chain (two methylene groups) allows the proper conformation for positioning the imino sp2 and amino sp3 donor sites appropriately. Adenine can chelate alkali metal cations between the 6-amino substituent and the N7 nitrogen. This was anticipated from experimental entropies of dissociation [60] and confirmed by MP2/6–31G∗ calculations [61]. It is worth mentioning that chelation is at the cost of some resonance effects, because the NH2 group is twisted out of the plane of the adenine molecule to produce a chelating structure (Figure 6.4). In the case of the Cu+ /adenine adduct, the calculations do not lead to a straightforward conclusion about the preferential site of complexation, but N3 and N7 are clearly favoured over N1 [62]. To complete the picture of chelation by nitrogen bases, it is worth describing the complex of aminoacetonitrile with Ni+ . This bifunctional ligand chelates Ni+ by using the π electrons of the nitrile function and the n electrons of the amino group [63] (Figure 6.5).

Figure 6.4

Most stable calculated structure of Na+ /adenine (N7 complex).

330


Figure 6.5 Geometry of the most stable Ni+ /aminoacetonitrile adduct. Reprinted from [63] by permission of Wiley–Blackwell.

In amidines, the push–pull effect increases the electron density on the imino nitrogen at the expense of the amino nitrogen. In the complexes of formamidine with Li, Na, Mg and Al cations, the structures in which M+ bridges both the amino and imino nitrogen atoms are true minima on the PES but lie about 20 kJ mol−1 above the imino monodentate complex [64]. In the complex of Cu+ with guanidine, the imino nitrogen is also the preferred binding site [65]. Single-bonded oxygen bases. In the M+ complexes of water, alcohols and ethers, the main question concerns the position of the cation relative to the two electron pairs of the oxygen atom. Does the cation interact with the two lone pairs by forming a trigonal complex with an R O M+ angle of 120◦ or with only one sp3 lone pair by making a tetrahedral complex with an R O M+ angle of 109◦ ? In the most studied system, H2 O M+ , nearly all the reported structures present a trigonal geometry around the oxygen and a C2v symmetry [51]. Cations also bind along the twofold axis of dimethyl ether. The formation of non-classical chelates enhances the bonding of Li+ with a number of alcohols (CF3 CH2 OH [66], CCl3 CH2 OH [67], n-BuOH [68], c-C6 H12 CH2 OH [67], C6 H5 CH2 OH [67], etc.), as illustrated in Figure 6.6. True chelation is expected for polyols or polyethers. A simple polyether, 1,2dimethoxyethane, was the target of a number of theoretical studies with alkali metal cations

a

b

Figure 6.6 Geometry of (a) Li+ /CF3 CH2 OH, (b) Li+ /n-BuOH. (a) reproduced from [66] by permission of the Centre National de la Recherche Scientifique (CNRS). (b) reproduced from [68] with permission from the American Institute of Physics. Copyright 2001.


331

Figure 6.7 Geometry of Cu+ /MeOCH2 CH2 OMe. Reproduced from [73], copyright 2001, Elsevier.

[69–72] and Cu+ [73]. At the B3LYP/6–311G∗ level, the copper cation adduct exhibits C2 symmetry, as shown in Figure 6.7. The results for Cu[MeO(CH2 )n OMe]+ with n > 3 also support the chelation, but comparison with Li+ shows that LiCB is less sensitive than CuCB to the number of methylene groups [74]. This is attributed [74] to the sdσ hybridization of Cu+ [75] that is optimal for a linear O Cu O geometry. Ab initio calculations on the structure of alkali metal cation adducts of the smallest crown ether, 12-crown-4, give a C4 symmetry in which the cation is tetracoordinated and positioned above the median plane of the oxygen atoms, in a direction close to the dipole of each C O C moiety [72] (Figure 6.8). Polyol ligands (ethylene glycol, propane-1,2-diol, propane-1,3-diol and propane-1,2,3triol) were used as models for the interaction of carbohydrates with Li+ , Na+ and K+ . Multidentate effects are the rule with these ligands, but the existence of intramolecular hydrogen bonds is a complicating factor in the determination of structures, as compared with polyethers [76]. Carbonyl bases. Formaldehyde is a prototypical model of the carbonyl group of aldehydes, ketones, esters or amides and has been the most studied. Table 6.2 summarizes two ideal geometries expected for H2 COM+ complexes. Whereas protonated formaldehyde has a bent geometry, many cations seem to form linear complexes. This is true of Li+ , Na+ , Mg2+ [77], Al+ [78], Mg+ , Ca+ and Zn+ [79]. However, in the case of Al+ , a slightly bent structure is very close in energy to the linear structure, whereas a bent structure is calculated [80] for Cu+ at the MP2/6–311+G(2f,2d,2p) level. In contrast, the B3LYP geometry of Me2 CO Cu+ is linear [16].

Figure 6.8

Side-view of the Na+ /12-crown-4 complex.

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Table 6.2 Geometries of H2 COM+ complexes. Geometry

Symmetry

Bent

Cs

Linear

C2v

a

Directionality Axis of a putative sp2 lone pair Axis of dipole moment

C O M+ angle (◦ )

Main interactiona

120

Covalent

180

Ion/dipole

Governing the geometry.

The Al+ /acetaldehyde complex is bent by 30◦ toward the hydrogen side of the carbonyl [81] (Figure 6.9). On the other hand, the Li+ and Na+ adducts exhibit a slight bending (8◦ and 5◦ , respectively) in the opposite direction, with the cation towards the methyl side. In the case of acetophenones with Cu+ (one and two ligands), the complex is also bent, with an angle close to 135◦ , the metal being on the side of the methyl group [82]. Carbonyl groups are constitutive of the nucleobases uracil, thymine (5-methyluracil), cytosine and guanine, in the form of amide and urea functional groups. These functions exhibit the largest MCAs in carbonyl compounds. Uracil and thymine are complexed on O4 by Li+ , Na+ and K+ with a (quasi)linear C O M+ arrangement [61]. Methylation(s) of uracil has a relatively small effect on this angle [83]. In contrast, halogenation in position 5 allows a chelation, resulting in a tilt toward the halogen atom with a decrease in the C O M+ angle, from 150◦ for the weakest interaction (K+ iodine) down to 112◦ for the strongest interaction (Li+ fluorine) [84]. When one of the two carbonyl groups is transformed into a thiocarbonyl, monothiouracils remain oxygen bases [85]. The graphical results of B3P86/6–311+G(d, p) calculations on Al+ /uracil indicate a preferred site of complexation and a linear geometry similar to those found with alkali metal cations [86]. In the case of Cu+ , the graphical representation also indicates an O4 complexation, but with a bent geometry [87]. Cytosine and guanine have only one carbonyl group but also bear amidine functions that may coordinate cations by their imino nitrogen(s). Actually, Na+ is chelated in the cytosine adduct by O2 and N3 [88] (Figure 6.10a). This is also the case for Cu+ and Ag+ [89]. Intriguingly, the Al+ adduct presents only a relatively small bending of the C O Al+ angle toward N3 [89, 90] (Figure 6.10b). A similar situation occurs for the bending toward N7 in the Al+ /guanine adduct [91].

Figure 6.9 MP2/6–311++G(2d,2p) geometry of Me(H)C O Al+ . Reprinted from [81] with permission from the American Institute of Physics, copyright 2002.


a

333

b

Figure 6.10 Most stable calculated geometries of (a) Na+ /cytosine and (b) Al+ /cytosine. (a) reprinted from [88] by permission of Wiley–Blackwell.

Amino acids. In addition to the COOH and NH2 groups, many other basic sites may be provided by heteroatoms or π rings of the side-chain R of amino acids H2 NCH(R)COOH. Three complexation motifs are of particular interest [92], as shown in Scheme 6.1. These are the charge-solvated motif CS1, the charge-solvated motif CS2 and the salt-bridge motif SB. M+ is chelated by the COOH and NH2 groups in CS1, the two oxygens of the COOH group in CS2 and the two oxygens of the carboxylate group of the zwitterionic amino acid in SB. When the side-chain takes part in the complexation, it may be directly involved in a tridentate CS structure, or it may indirectly stabilize the zwitterion through an intramolecular hydrogen bond [80, 93, 94]. Structures obtained using IR multiphoton dissociation spectrometry in conjunction with quantum chemical calculations seem the most reliable. They are available for serine [95], threonine [96] and tryptophan [97] with alkali metal cations (Li+ to Cs+ ), phenylalanine [98] with Ag+ and Zn2+ , glycine and proline [99] with Na+ , and also for oligoglycines and oligoalanines [100] with Na+ . Phosphoryl bases. The phosphoryl bond is more polar than a carbonyl group, and is expected to give rise to a strong electrostatic interaction with a cation. Although the strict application of an ion/dipole model of bonding predicts a linear P O M+ system, the calculations give an angle of 139–153◦ in the trialkylphosphate adducts of alkali metal cations, depending on the cation, with no simple relationship with the size [101, 102]. It is

H

M H H

N

H

C C

R

O

O

O H

CS1

H H

N

C C

R

O H

CS2

H

M

H H

N

C C

R

M

O

O H

SB

Scheme 6.1 Motifs for the complexes of amino acids with metal cations.

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noteworthy that the cation is chelated by the phosphoryl oxygen and an oxygen atom of the alkoxy groups.

6.3

Experimental Techniques for Measuring Gas-Phase Cation Affinities and Basicities

The study of the interaction between an ion (cation or anion) and neutral species under very low pressure (isolated in a ‘vacuum’) is usually performed using mass spectrometric techniques. Specialized instrumentation has been designed for this purpose, but it is also possible to use commercial spectrometers for the quantitative determination of thermodynamic quantities related to ion/molecule interactions. Reviews on the experimental techniques relevant to gas-phase ion thermochemistry [103] and on the methods pertaining to the determination of cation/ligand binding energies [104] are recommended to the reader interested in mass spectrometric methodologies. Typical methods are briefly described in the following. Metal cations may be generated in vacuo by several means compatible with low pressures, but in certain instances, atmospheric pressure generation of metal ions, or their adducts, is interfaced to research instruments. The problem of interfacing atmospheric ion generation is counterbalanced by the new possibilities of studying thermally labile or non-volatile molecules, such as those of biological importance, and also cations difficult to generate in the gas phase (refractory metals, multiply-charged cations, etc.). The need to analyse fragile and high molecular weight biomolecules led to the generalization of atmospheric pressure ion sources on commercial mass spectrometers. The methods commonly used for ion generation are mentioned for each specific technique described below. Similarly, the kind of molecule that can be studied is dependent on its properties (size, volatility, stability) and on the technique. These specificities are noted for each case. 6.3.1

High-Pressure Mass Spectrometry (HPMS)

High-pressure mass spectrometry (HPMS) has largely been applied to the determination of bonding energetics in adducts (also called clusters when several ligands are bonded to one cation) of alkali metal cations with simple molecules [105, 106]. The HPMS technique consists in reacting ions with neutral molecules, in the presence of an inert bath gas, at pressures in the range 102 –103 Pa. Note that the name of the technique comes from the relatively ‘high pressure’ in the part of the instrument where ion/molecule reactions occur, as compared with the much lower pressure (10−3 –10−6 Pa range) in the usual mass spectrometers. At HPMS pressures, adduct ions become in thermal equilibrium with the bath gas (and the walls of the ion source) within a few milliseconds through collisions. The equilibrium constants K for reaction 6.3 can be obtained from the ion abundances and the partial pressure of the neutral ligand. After measuring K at several temperatures, a van’t Hoff plot (lnK versus 1/T) can be drawn: the linear relationship gives the enthalpy and entropy of reaction 6.3 from the slope and the intercept, respectively. The enthalpy is the M+ cation affinity (MCA) of L. The precision of these MCA values is roughly in the range 1–6 kJ mol−1 , and 2–15 J K−1 mol−1 for entropies. It is worth


335

recalling (see Section 1.7 in Chapter 1) that such data treatment produces enthalpies and entropies with correlated uncertainties: too small a slope generates too large an intercept and vice versa. The complexes must be sufficiently dissociated to be measured in the HPMS temperature range (rarely below room temperature and up to about 600 K, the limit of stability of many organic compounds). Therefore, too strongly bound complexes cannot be measured and, conversely, too weak adducts are excessively dissociated at the lowest convenient temperatures. Most reported HPMS values for absolute cation affinities fall in the range 10–150 kJ mol−1 , but relative affinities or basicities may also be measured by ligand exchange. As for all mass spectrometric techniques, limitations arise in part from the possibility of generating gas-phase ions. Usually, molecules are vaporized and allowed to react with metal ions. In early studies, alkali metal cations were produced by thermionic emission from a filament coated with an appropriate melt of alkali metal oxide (or carbonate) with silica and alumina. Most recently, electrospray ionization (ESI), one of the atmospheric pressure ion sources mentioned previously, has been used to form gas-phase adducts directly, avoiding some problems associated with volatility, thermal stability, and so on [105].

6.3.2

Collision-Induced Dissociation Threshold (CIDT)

One of the most fruitful methods for the determination of cation/ligand absolute bond dissociation energies is the collision-induced dissociation threshold (CIDT), also called ‘threshold CID’ [13, 107, 108]. Collision-induced dissociation (CID, also called MS/MS) is frequently used in analytical mass spectrometry to remove interference using the different dissociation pathways of ions of the same mass but of different structures. The energy threshold for the dissociation can provide information about the weakest bonds in the ion. Upon dissociation, an adduct between a cation and a ligand usually gives back these initial components, although other bond cleavages may be induced by the coordination with the cation. Threshold measurements must be performed under carefully controlled conditions, in particular the pressure of the collision gas and the kinetic energy of the ions. After fragmentation of adducts (corresponding to the dissociation reaction 6.3) by collision with a rare gas, the intensities of the fragment ions (in this case, only M+ is expected) are measured as a function of the kinetic energy of LM+ . The intensity versus collision energy plot is fitted to a model that relies partly on vibrational frequencies for the complex, which are generally obtained from quantum chemical calculations, and ultimately leads to the dissociation energy at 0 K. Applying thermal corrections, obtained from DFT or ab initio calculations, gives absolute bond dissociation enthalpies at 298 K, accurate to about ±5–15 kJ mol−1 , and also Gibbs energies after entropy calculations. In the original method designed by Armentrout, adducts are formed in a ‘flow tube’ in which metal ions, generated by sputtering by an electric discharge with a cathode containing the pure metal or metal salt, react with ligand vapours diluted in a helium flow. The variety of adducts that can be studied by CIDT was extended by replacing the flow tube with an electrospray ion source [109, 110]. The aim was to generate cation adducts with labile or non-volatile molecules (e.g. amino acids), higher molecular mass systems (peptides), or even multiply-charged complexes, such as Ca(H2 O)n 2+ .

336

6.3.3


Ligand-Exchange Equilibrium Measurements in Trapping Devices

Quadrupole ion trap [111] (QIT) and ion cyclotron resonance (ICR) mass spectrometers [112] have been used for the determination of Gibbs free energies (∆G) of ligand exchange (relative basicity), as illustrated by reaction 6.4 for a metal cation M+ : [L1 M]+ + L2 [L2 M]+ + L1

(6.4)

Adducts are trapped in the presence of a known pressure of the ligands L1 and L2 . When a steady state is observed, usually after a few seconds when working in the 10−5 –10−4 Pa range, an equilibrium constant K can be calculated from the pressure ratio of the ligands and the ratio of ion intensities: K =

P(L1 ) I (L2 M+ ) P(L2 ) I (L1 M+ )

(6.5)

The standard thermodynamic relationship ∆G◦ = −RT lnK gives the difference between the cation basicity of L1 and L2 or the relative MCB. These relative values are fairly accurate, with uncertainties of about 0.5–2 kJ mol−1 , if pressures are well controlled and there are negligible secondary reactions. When these relative values can be linked, they form a ‘ladder’ that may cover ranges of hundreds of kJ mol−1 . This procedure, which adds up numerous individual data, produces an increasing uncertainty as the number of added values grows. The total uncertainty may be evaluated by the ‘propagation of errors’, as 1 the square root of the sum of squared standard errors [Σ(si )2 ] 2 . For example, adding 10 values, each bearing an uncertainty (standard deviation) of 0.8 kJ mol−1 , produces a total uncertainty of 2.53 kJ mol−1 on the sum, corresponding to the total range of the ladder. Relative basicities may be converted into relative affinities by estimating relative entropies. The instrumental temperature range and control are too limited to take advantage of a van’t Hoff plot, and usually the relative MCBs (relative Gibbs energies) are given at a fixed temperature. When very similar ligands appear on each side of the equilibrium 6.4, entropy may be considered negligible and ∆MCA = ∆MCB. More generally, the entropy may be calculated as the sum of translation, rotation and vibration contributions. The relative values may be anchored to some absolute cation affinity or basicity value(s) to obtain absolute MCA or MCB scales. The absolute accuracy of such anchored scales depends mostly on the anchoring data since, in general, absolute values are less accurate than relative values. When an equilibrium state cannot be achieved, bracketing techniques may be used: limits can be placed on the basicity of M+ , derived from fast or slow kinetics of cation exchange. In reaction 6.6: [L1 M]+ + Li → [Li M]+ + L1

(6.6)

the adduct [L1 M]+ is confronted with different ligands Li of increasing basicity until a fast reaction occurs. An absent (or very slow) reaction indicates that Li is a weaker base than L1 , with the underlying postulate that there is no (Gibbs) energy barrier. Conversely, a fast reaction shows that Li is a stronger base than L1 . Bracketing gives uncertainties in the range 5–10 kJ mol−1 . In early studies, alkali metal cations were obtained by thermionic emission, as for HPMS, and later, a large number of metal cations were formed by laser vaporization/ionization of


337

metals or their salts. Although ESI sources may help in generating gas-phase adducts of specific molecules, gas-phase equilibrium studies require that the two ligands L1 and L2 , Equation 6.4, are sufficiently volatile. In the case of bracketing, Equation 6.6, only one gasphase ligand (Li ) is necessary if the reactant adduct ion [L1 M]+ can be generated without the need to vaporize L1 . 6.3.4

Selected Ion Flow Tube (SIFT)

In flow tube techniques, the reacting ion is mixed with a fast stream of helium in which the neutral molecule is diluted. The speed of the flowing bath gas and the length of the tube establish the contact time between the ion and the neutral molecule. The charged product ion(s) are analysed downstream using a mass analyser. Thanks to the setup of innovative ion sources, Bohme and co-workers were able to study the reactivity of transition metal and main group cations with small molecules, using the selected ion flow tube (SIFT) technique [113–115]. The extent of the reaction depends on the pressure of the neutral molecule and the contact time with the ion. SIFT is a technique essentially intended for kinetic studies but, when equilibrium can be reached, an accurate value of the absolute Gibbs free energy for the cation attachment (at a single temperature) can be deduced. Hitherto, the technique has given basicity values for weak interactions (MCBs in the range 10–50 kJ mol−1 ) with various cations [116–118], with a precision of about 1 kJ mol−1 . 6.3.5

Kinetic Method

The kinetic method is probably one of the most commonly used techniques for the determination of relative proton and metal cation affinities, in addition to other cation affinities, gas-phase acidities, electron affinities and so on, because of its ease of implementation using standard mass spectrometers [119, 120]. The kinetic method is based on the dissociation of 1 : 2 adducts, in which two different ligands are bound to a common metal ion, as in the example shown for a metal cation M+ (Scheme 6.2). The weakest metal–ligand bonds easily break down by metastable dissociation or collisional activation into two 1 : 1 adducts, [L1 M]+ and [L2 M]+ . When L1 and L2 have similar structures, the natural logarithm of the ratio of the two unimolecular rate constants ln(k1 /k2 ), with (k1 /k2 ) equal to the ratio of ion intensities, is directly related to the difference in cation affinity or basicity of L1 and L2 , ∆MCA and ∆MCB, respectively, through an equation involving an effective (ad hoc, non-thermodynamic) temperature T eff : ln(k1 /k2 ) ≈ ∆MCA/RTeff ≈ ∆MCB/RTeff

1

k1

[L1-M]+ + L2

k2

[L2-M]+ + L1

2 +

[L -M-L ]

Scheme 6.2 Dissociation of a 1 : 2 adduct into neutral species and 1 : 1 adducts.

(6.7)

338


This derivation has its roots in the transition-state theory and involves some assumptions about the dissociation mechanism. The effective temperature remains unknown unless a series of standards (molecules of known ∆MCA or ∆MCB) are measured under the same conditions. This means that the method necessitates a calibration using known affinities or basicities. It should be noted that the temperature of the resulting ∆MCA or ∆MCB is not T eff but corresponds instead to the temperature at which the standards were measured. As evidenced by the equality between the last two terms in Equation 6.7, entropy effects are neglected in this simplified treatment. This limitation has been addressed for systems where large variations in the degrees of freedom are associated with adduct formation, and a broadening of the scope of the kinetic method in the direction of entropy determinations (the so-called ‘extended kinetic method’) has been proposed [121, 122]. The kinetic method can be used with mass spectrometers [123, 124] that allow CID: electric/magnetic sector (spontaneous unimolecular dissociation of metastable ions in sector instruments may also be exploited), triple quadrupole, quadrupole ion traps, Fourier transform ion cyclotron resonance (FTICR) and several other mass spectrometers in which different techniques are combined. In early studies, most cation-bound dimers were generated by fast atom bombardment [124], but currently ESI [124] is generally used for their production. The wide accessibility of ESI sources, which may be used for the formation of ions of polar and non-volatile molecules, renders the kinetic method broadly applicable.

6.3.6

Radiative Association Kinetics (RAK)

The kinetics of ion/molecule adduct formation, under the conditions of low-pressure [ion cyclotron resonance (ICR)], is controlled by radiative relaxation: the energy released by the complexation can only relax by emission of infrared photons. The link between the strength of the incipient bond and the association rate constants relies on models based on infrared emission for the relaxation of the energy of association [125, 126]. To date, the radiative association kinetics (RAK) technique has been applied to bond formation between metal cations and relatively simple molecules. The consistency of the results observed from different models suggests uncertainties of about ±10–15 kJ mol−1 .

6.3.7

Blackbody Infrared Radiative Dissociation (BIRD)

The blackbody infrared radiative dissociation (BIRD) technique can be considered as the reverse process of radiative association. At very low pressure, trapped ions (essentially in an ICR cell) are slowly dissociated by the blackbody radiation emitted by the walls of the ICR cell [127]. Many different gas-phase ion dissociations have been studied, including cation adducts with biomolecules. The models, on which the thermochemical interpretation of the data relies, depend on the size of the systems. In particular, it is easier to extract data from BIRD of large ion clusters, and most studies report data on the detachment of a small ligand from a large ion or cluster. The precision is similar to, or better than, that of RAK.


[CsF]s

from ∆ fH

∆ subH [CsF]g

Fg + Csg EA

∆H = CsCA

339

IE

F-g + Cs+g

Scheme 6.3 Illustration of the determination of the enthalpy of dissociation of caesium fluoride (equivalent to the CsCA of F− ) using a thermochemical cycle [9].

6.3.8

Vaporization and Lattice Energies

Before the advent of mass spectrometric methods, thermochemical data on isolated (gasphase) ions were obtained from thermodynamic cycles involving lattice energies (or enthalpies), enthalpies of formation, ionization energies and/or electron affinities [2]. The lattice energy (or lattice enthalpy) is the amount of energy (enthalpy) necessary to separate a salt into its gaseous ionic constituents (anions A− and cations M+ ) at infinite distance. It is clear that, if the lattice enthalpy and the vaporization enthalpy of the solid salt [A− M+ ] are known, application of the Hess’s law gives the enthalpy of reaction 6.3, with L = A− , that is, the MCA value of A− . This historical approach is still of interest if the intrinsic anion/cation energy is being searched for because there is currently no direct experimental technique for determining this interaction energy. Scheme 6.3 illustrates how to estimate the enthalpy of the heterolytic dissociation of caesium fluoride. By combining the enthalpy of vaporization ∆vap H (or sublimation ∆sub H) of the salt with the enthalpies of formation ∆f H of the solid salt and of the gaseous monatomic element, the metal ionization energy (IE) and the electron affinity (EA) of the halogen atom, the CsCA of the fluoride anion is obtained. Spectrometric data (EA, IE) on atoms and small molecules are generally fairly accurate. The thermochemical data suitable for the calculation of affinities (or basicities) using Hess’s law may be found already combined in thermodynamic tables. They are tabulated in the form of standard enthalpies (and Gibbs energies) of formation of the gaseous species of interest, for example, the gas-phase enthalpy of formation, ∆f H(g), of salts, metal cations and anions, already evaluated through some particular component(s) of the cycle shown in Scheme 6.3. The most challenging step of the thermodynamic cycle applied to inorganic, non-volatile compounds is the experimental determination of ∆vap H or ∆sub H. Such data are scarce and sometimes of uncertain accuracy [9].

6.4

Ion Thermochemistry Conventions

There are two different conventions used in ion thermochemistry, the so-called ‘electron convention’ and the ‘ion convention’ [2, 128]. They differ in the treatment of the electron captured or liberated during ion formation. In the electron convention, the electron is treated as an element of enthalpy: ∆f H (e− ) = 0

HT − H0 =

5 RT (ideal gas) 2

(6.8)

340


In the ion convention, it is considered as a species without any heat capacity and consequently ∆f H (e− ) = 0

HT − H0 = 0 (at all T )

(6.9) −1

At 298 K, the difference between the two conventions is 6.2 kJ mol per electron. For example, the enthalpy of formation at 298 K of Cs+ found in two reference compilations is 458.0 kJ mol−1 under the ion convention [129] and 451.8 kJ mol−1 under the electron convention [128]. A consistent use of either of the two conventions leads to the same result in terms of reaction thermochemistry (that is, when the electron is not involved in either side of the reaction), provided that there is no mixing of the two. The mass spectrometrist community normally uses the ion convention. A complete description of the two conventions, including details of their origin and their consequences, is given in the ‘Gas-Phase Ion Thermochemistry’ chapter of the NIST Chemistry WebBook and the early NIST data compilation of gas-phase ion and neutral thermochemistry [128]. We also note that the standard pressure used is 1 bar rather than 1 atm, although such a change in the pressure standard state has little effect on the results with regard to the usual level of precision.

6.5

Lithium, Sodium, Potassium, Aluminium, Manganese, Cyclopentadienylnickel, Copper and Methylammonium Cations Affinity and Basicity Scales

The selection of reference cations has been explained in Section 6.1. The present section describes how the experimental scales are compiled, for example if they are either listed directly as in the original articles or rescaled according to new knowledge. Uncertainties were evaluated as far as possible from the original work. When data are calculated by combining two or more values, the total uncertainties U tot are evaluated from individual uncertainties U i (as stated in the original publication) and the classical propagation of 1 absolute standard errors: Utot = (ΣUi2 ) 2 . 6.5.1

Lithium

The earliest scale (1975), determined for 30 ligands from ICR ligand-exchange equilibrium measurements at room temperature, was published in graphic mode [130]. Only 10 out of the 30 values were published later in a numerical form [131]. After anchoring to an experimental and theoretical LiCA value for H2 O, these values were considered by the authors as absolute affinities (LiCAs). The scale needs to be revisited for three reasons: (i) the values are actually relative basicities (∆LiCBs) anchored to an LiCA value, (ii) recent calculations and experiments show that the anchoring datum is inaccurate and (iii) the actual temperature might have been somewhat higher than room temperature, due to the method of production of Li+ by thermoionization. Therefore, the ICR data (taken partly from the figure in the 1975 communication) are re-anchored to an updated absolute LiCB value for H2 O [132] at 298 K, evaluated from CIDT measurements. They are listed in Table 6.3. CIDT values (absolute LiCA at 0 and 298 K and LiCB at 298 K) are also included in this table.

15 16 17 18

1 2 3 4 5 6 7 8 9 10 11 12 13 14

No.

Carbon bases Propene Cyclohexane 2-Methylpropene Propyne Fluorobenzeneb Benzene Pyrroleb 1-Methylpyrroleb Toluene Naphthalene Phenolb Anisoleb Anilineb Indoleb Amines Ammonia Methylamine n-Propylamine Trimethylamine

Lewis base

NH3 MeNH2 n-PrNH2 Me3 N

a

C6 H5 OH C6 H5 OMe C6 H5 NH2

a

MeCH CH2 c-C6 H12 Me2 C CH2 MeC CH C6 H5 F C6 H6 c-(CH)4 NH c-(CH)4 NMe C6 H5 Me

Formula

197.8 ± 6.0

146.9 ± 20.1 161.1 ± 13.5 177.4 ± 16.6 186.2 ± 16.8 183.1 ± 16.0 187.2 ± 15.4 178.5 ± 16.1 184.4 ± 18.3 191.5 ± 22.4 204.5 ± 8.7

LiCA (0 K)

149.6 ± 20.3 164.4 ± 13.6 181.1 ± 16.7 189.8 ± 16.8 186.4 ± 16.3 189.8 ± 15.6 181.6 ± 16.3 187.2 ± 18.5 194.5 ± 22.5 207.7 ± 8.7

LiCA (298 K)

116.7 ± 21.4 132.2 ± 14.3 147.1 ± 17.4 154.8 ± 17.6 150.1 ± 17.5 156.9 ± 17.1 150.8 ± 17.5 153.8 ± 19.8 163.7 ± 23.4 173.6 ± 8.7

LiCB (298 K) (CIDT)

[132]

[42] [39, 132] [48] [48] [149] [150] [132, 151] [152] [43] [153]

Ref.

(Continued)

139.6

132.9 137.9

122.8

67.0 71.0 88.0 89.7

LiCB (298 K) (ICR)

LiCA (0 K): absolute CIDT measurements LiCA (298 K): LiCA (0 K) corrected at 298 K LiCB (298 K) (CIDT): LiCA (298 K) – T∆S (calculated) LiCB (298 K) (ICR): ∆G (298 K) + (112.8 ± 14.3), that is, relative ICR measurements at 298 K [130, 131] anchored to a reference value for H2 O [132]

Table 6.3 Lithium cation affinity (LiCA) and basicity (LiCB) scales (kJ mol −1 ).

41 42 43 44

39 40

32 33 34 35 36 37 38

21 22 23 24 25 26 27 28 29 30 31

Dimethylamine Me2 NH 2-Aminoethanolc H2 NCH2 CH2 OH Six-membered aromatic N-heterocycles 1,3,5-Triazine 1,3,5-N3 C3 H3 Pyrazine 1,4-N2 C4 H4 Pyrimidine 1,3-N2 C4 H4 Pyridine C5 H5 N 2-Methylpyridine 2-MeC5 H4 N 3-Methylpyridine 3-MeC5 H4 N 4-Methylpyridine 4-MeC5 H4 N 3-Aminopyridine 3-H2 NC5 H4 N 4-Aminopyridine 4-H2 NC5 H4 N Pyridazinec 1,2-N2 C4 H4 2-Aminopyridinec 2-H2 NC5 H4 N Five-membered aromatic N-heterocycles a 2-H-1,2,3-Triazole a 2-H-Tetrazole a Pyrazole a 1-H-1,2,4-Triazole a 1-Methylpyrazole a Imidazole a 1-Methylimidazole Nitriles Hydrogen cyanide HC N Acetonitrile MeC N Water, alcohols Water H2 O Methanol MeOH Ethanol EtOH n-Butanol n-BuOH

19 20

Formula

Lewis base

No.


133.1 ± 13.5 155.0 ± 8.5 163.5 ± 6.5 168.6 ± 8.2

137.2 ± 13.7 156.8 ± 8.5 165.5 ± 6.5 171.3 ± 8.2

109.2 ± 14.3 127.7 ± 8.6 136.4 ± 6.7 138.7 ± 8.4

[132] [132, 154] [132, 154] [154]

[56] [56] [48] [56] [48] [49, 132] [49]

138.2 153.4 189.2 ± 16.2 193.7 209.4 ± 18.5 213.2 ± 9.7 244.6 ± 20.3

136.2 ± 7.3 151.3 ± 6.6 187.1 ± 16.1 191.3 ± 7.8 207.2 ± 18.4 210.8 ± 9.5 242.3 ± 20.2

110.5 125.8 159.5 ± 16.9 165.6 178.8 ± 19.1 185.2 ± 10.2 213.7 ± 20.9

[57] [57] [57] [132] [53] [53] [53] [52] [52] [57] [52, 132]

129.4 ± 12.6 151.1 ± 13.2 156.3 ± 10.5 183.0 ± 14.5 196.2 ± 6.7 198.7 ± 14.8 198.2 ± 13.5 203.4 ± 10.2 217.4 ± 20.3 237.2 ± 10.6 241.1 ± 21.2

Ref.

127.4 ± 12.6 149.1 ± 13.2 154.3 ± 10.5 181.0 ± 14.5 194.3 ± 6.5 196.6 ± 14.7 196.2 ± 13.4 201.7 ± 10.0 216.9 ± 20.2 234.4 ± 10.6 237.8 ± 21.1

101.4 ± 12.6 123.1 ± 13.2 128.3 ± 10.5 155.1 ± 14.5 168.4 ± 7.6 169.1 ± 15.2 170.3 ± 13.9 175.6 ± 10.9 192.9 ± 20.8 207.7 ± 10.6 211.8 ± 21.6

LiCB (298 K) (CIDT) [132]

LiCA (298 K)

289.5 ± 9.0

LiCA (0 K)

112.8 125.4

118.7 151.1

154.9

141.3

LiCB (298 K) (ICR)

63 64 65 66 67

62

61

60

53 54 55 56 57 58 59

50 51 52

45 46 47 48 49

2-Methyl-1-propanol n-Propanol 2-Propanol 2-Butanol tert-Butanol Ethers Dimethyl ether 1,2-Dimethoxyethanec 12-Crown-4d Carbonyl compounds Formaldehyde Acetaldehyde Methyl formate Methyl acetate Acetone Butan-2-one N,N-Dimethylformamidee Nitroalkane Nitromethane Phosphoryl compound Trimethyl phosphate Sulfide Methyl sulfide Haloalkanes Chloromethane Difluoromethane Dichloromethane 2-Iodopropane 2-Chloropropane CH3 Cl CH2 F2 CH2 Cl2 CH3 CHICH3 CH3 CHClCH3

Me2 S

MeNO2

H2 CO MeCHO HCOOMe MeCOOMe Me2 CO MeCOEt HCONMe2

a

Me2 O (MeOCH2 )2

i-PrCH2 OH n-PrOH i-PrOH EtCH(OH)Me t-BuOH

280.8 ± 14.5

282.1 ± 14.5

192.2 ± 7.0

166.8 ± 10.7 245.1 ± 18.3 378 ± 51

165.0 ± 10.6 241.2 ± 18.3 371.5 ± 51

190.1 ± 7.0

171.1 ± 7.6 172.8 ± 8.6 174.9 ± 7.5 176.2 ± 8.9 180.1 ± 10.2

168.8 ± 7.6 170.3 ± 8.6 172.8 ± 7.5 174.3 ± 8.9 178.2 ± 10.2

256.2 ± 15.0

161.4 ± 9.5

138.9 ± 11.8 209.1 ± 19.4

139.6 ± 7.8 141.1 ± 8.8 144.9 ± 7.7 146.3 ± 9.1 150.8 ± 10.3

[102]

[132]

[132] [69, 132] [69, 132]

[154] [132, 154] [132, 154] [154] [154]

(Continued)

75.4 81.7 91.5 94.6 96.4

104.4

133.8

179.9

116.6 144.1 145.1 153.7 156.7

129.5

a a a

Uracilf

Adenineg

Cytosineh Guanineh Amino acids Glycinei Azetidine-2-carboxylic acidj Pipecolic acidj Prolinej

73

74

75 76

77 78 79 80

72

a

a

H2 NCH2 COOH

a

a

a

CH3 CHBrCH3 CH3 F CH3 CH2 F CH3 CHFCH3

2-Bromopropane Fluoromethane Fluoroethane 2-Fluoropropane Nucleobases Thyminef

68 69 70 71

Formula

Lewis base

No.


220.0 ± 8.0 250 ± 14 272 ± 16 278.8 ± 9.7

226.1 ± 6.1

211.5 ± 6.1

210.1 ± 7.0

LiCA (0 K)

253 ± 14 274 ± 16 282 ± 10

212.4 ± 7.0 215 213.8 ± 6.1 211 230.0 ± 6.1 226 232 239

LiCA (298 K)

223 ± 15 244 ± 17 248 ± 11

195.9 ± 6.1

186.2 ± 6.1

184.8 ± 7.0

LiCB (298 K) (CIDT)

[132] [155] [155] [132, 155]

[61] [60] [61] [60] [61, 132] [60] [60] [60]

Ref.

98.1 98.5 111.2 118.6

LiCB (298 K) (ICR)

N H

4

O

R

72 73

14

R: Me R: H

N

32

N

N

b

N

N

74

NH

H

Miscellaneous Argon Carbon monoxide Nitrogen monoxide Phosphine Tetraphosphorus

π complex or π and n complexes. c Bidentate behaviour. d Tetradentate chelate. e Extrapolated value. f Mainly O4 complex. g Bidentate chelate (NH2 /N7). h Uncertain site or chelation. i Charge-solvated complex. j Salt-bridge complex.

O

N

10

Formulae:

H

a

81 82 83 84 85

H

N 7

N

33

N

N

O

N

N

H

2

3

75

H

N

NH

R

N

H

HN

N

N

34 36

Ar CO NO PH3 P4

76

N

6

O

N H

N 7

R: H R: Me N

COOH

H

78

N H

35

N R

N

N

32.8 ± 13.5 55.0 ± 12.5 59.8 ± 10

N H

79

N

80

N H

R: H R: Me

COOH

37 38

57.0 ± 13

O

O

52

COOH

O

O

[132] [132] [132] [156] [156]

79.3 ± 13.8 79.3 ± 12.3

346


By far the most comprehensive list of LiCB values was published in 2000 [133]. The LiCBs were determined at 373 K from FTICR ligand-exchange equilibrium measurements. These relative values were anchored to LiCB(H2 O, 373 K) derived from experimental LiCA and an ab initio-calculated entropy. The anchor value used here (103.6 kJ mol−1 ) is corrected for a small rounding error of 0.3 kJ mol−1 in the original work [133]. Subsequent measurements were performed by the equilibrium or by the kinetic method, and added to the list, after adjustment for consistency if necessary. LiCB values obtained for dimethoxyalkanes, MeO(CH2 )n OMe, n = 2–9, at a nominal temperature of 343 K [74] were linked here to the absolute LiCB for 1,2-dimethoxyethane at 373 K. Considering that the temperature effect (∆T = 30 K) in such a small range of relative LiCBs is certainly small and negligible and that the relative values were anchored to a value at 373 K, the absolute values in ref. [74] correspond in fact to the 373 K scale in ref. [133]. All these values are listed in Table 6.4. Additional LiCBs (373 K) may be calculated from the very good correlation between the LiCBs at 373 and 298 K common to both Tables 6.3 and 6.4 LiCB(373 K) = 1.0145 (±0.024) LiCB(298 K) −9.84 (±3.25) n = 18, r = 0.9957, s = 1.8 kJ mol−1 .

(6.10)

By these means, 16 new LiCB values were added to Table 6.4 for the compounds: cyclohexane, propene, 2-methylpropene, propyne, nitromethane, fluoromethane, fluoroethane, 2-fluoropropane, difluoromethane, chloromethane, 2-chloropropane, dichloromethane, 2bromopropane, 2-iodopropane, tetraphosphorus and phosphine. This equation is worthy of comment. The slope is close to unity. This means that variable entropy effects are minor within the temperature range 373–298 K (∆T = 75 K) and that the intercept (−9.84 kJ mol−1 ) can be assigned to a quasi-constant entropy term ∆T × ∆S. From this hypothesis, a rough entropy value of (9.84 ± 3.25) × 103 /75 = 131 (±43) J K−1 mol−1 is obtained, in approximate agreement with calculated entropies for Li+ adduct dissociation in the case of simple monofunctional ligands (see, for example, ref. [132]). 6.5.2

Sodium

As in the case of the lithium cation, there are roughly two major sources of sodium cation affinities and basicities: r Absolute NaCA measurements by CIDT, with possible entropy calculations to estimate NaCB at 298 K, to which a few HPMS values may be added. r Relative NaCBs obtained from equilibrium measurements, in general at room temperature, using FTICR mass spectrometry. In a few cases, the kinetic method was also utilized, particularly in the case of amino acids and their derivatives. As indicated in the main contribution to this scale [106], uncertainties on relative values are about 0.4 kJ mol−1 at the lower part of the scale, with an upper limit of 2.5 kJ mol−1 for the largest values. On the basis of a standard deviation of 0.4 kJ mol−1 per step, and 53 steps, a total of 2.9 kJ mol−1 is obtained, in reasonable agreement with authors’ estimation. The uncertainty on the absolute value for methylamine (81.6 kJ mol−1 ) is 0.8 kJ mol−1 .


347

Table 6.4 Lithium cation basicity (LiCB) scale (kJ mol−1 ) from FTICR ligand-exchange equilibria at 373 K, anchored to the experimental and ab initio value for H2 O (see text). No. 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43

Lewis base

Formula

Carbon bases Propene MeCH CH2 Cyclohexane c-C6 H12 3,3,3-Trifluoro-1-propyne CF3 C CH 2-Methylpropene Me2 C CH2 Propyne MeC CH Chlorobenzene C6 H5 Cl Bromobenzene C6 H5 Br Thiophenol C6 H5 SH Benzene C6 H6 Phenol C6 H5 OH Styrene C6 H5 CH CH2 Toluene C6 H5 CH3 a Naphthalene Ethylbenzene C6 H5 CH2 CH3 n-Butylbenzene C6 H5 (CH2 )3 CH3 a Phenanthrene a Anthracene a Azulene n-Heptylbenzene C6 H5 (CH2 )6 CH3 1,1-Diphenylethane C6 H5 CH(CH3 )C6 H5 1,2-Diphenylethane C6 H5 (CH2 )2 C6 H5 1,7-Diphenylheptane C6 H5 (CH2 )7 C6 H5 1,3-Diphenylpropane C6 H5 (CH2 )3 C6 H5 Amines Perfluoro-tert-butylamine (CF3 )3 CNH2 Ammonia NH3 Methylamine MeNH2 Trimethylamine Me3 N Dimethylamine Me2 NH 4,4,4-Trifluorobutylaminec CF3 (CH2 )3 NH2 Six-membered aromatic N-heterocycles Pentafluoropyridine C5 F5 N Pyrazine 1,4-N2 C4 H4 4-Trifluoromethylpyridine 4-CF3 C5 H4 N Pyrimidine 1,3-N2 C4 H4 3-Chloropyridine 3-ClC5 H4 N 2,6-Difluoropyridine 2,6-F2 C5 H3 N Pyridine C5 H5 N 2-Fluoropyridine 2-FC5 H4 N 3-Methylpyridine 3-MeC5 H4 N 3-Dimethylaminopyridine 3-Me2 NC5 H4 N Pyridazinec 1,2-N2 C4 H4 4-Dimethylaminopyridine 4-Me2 NC5 H4 N a 1,8-Naphthyridinec 2-(2-Aminoethyl)pyridinec 2-NH2 CH2 CH2 C5 H4 N

LiCBb

Ref.

58.1 62.2 74.9 79.4 81.2 99.4 101.8 112.3 112.8 117.8 122.8 124.5 127.9 130.4 136.6 141.3 141.7 145.3 150.3 155.4 175.8 184.3 185.2

g g

[133] g g

[157] [157] [157] [133] [157] [157] [157] [158] [40] [40] [158] [158] [158] [40] [41] [41] [41] [41]

99.7 126.4 131.2 134.1 134.5 155.0

[133] [133] [133] [133] [133] [133]

93.4 119.9 123.6 124.9 132.4 139.0 146.7 147.2 152.9 161.9 173.3 175.9 181.8 197.9

[133] [133] [133] [133] [133] [133] [133] [133] [133] [133] [133] [133] [133] [59] (Continued)

348


Table 6.4 (Continued) No. 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60 61 62 63 64 65 66 67 68 69 70 71 72 73 74 75 76 77 78 79 80 81 82 83 84 85 86 87

Lewis base

Formula

Five-membered aromatic N-heterocycles a 1,2,3-Triazole a 1,2,4-Triazole a Isoxazole a Tetrazolec a Thiazole a Pyrazole a 1-Methylpyrazole a 3(5)-Methylpyrazole a 4-Methylpyrazole a 1,4-Dimethylpyrazole a 1,5-Dimethylpyrazole a Imidazole a 1,3,5-Trimethylpyrazole a 3,4,5-Trimethylpyrazole a 1,3,4,5-Tetramethylpyrazole a 1-Methylimidazole a 1,2-Dimethylimidazole a 2,4,5-Trimethylimidazole a Histaminec Guanidine Tetramethylguanidine Me2 NC(NMe2 ) NH Nitriles Trifluoroacetonitrile CF3 C N Hydrogen cyanide HC N Fluoroacetonitrile FCH2 C N Malononitrile CH2 (C N)2 Trichloroacetonitrile CCl3 C N Dichloroacetonitrile Cl2 CHC N Chloroacetonitrile ClCH2 C N Cyanogen bromide BrC N Methoxyacetonitriled MeOCH2 C N Acetonitrile CH3 C N (Methylthio)acetonitrile MeSCH2 C N Benzyl cyanide C6 H5 CH2 C N Propionitrile EtC N Butyronitrile n-PrC N Benzonitrile C6 H5 C N Isobutyronitrile i-PrC N Valeronitrile n-BuC N Trimethylacetonitrile t-BuC N n-Heptyl cyanide n-HeptC N n-Octyl cyanide n-OctC N 1-Adamantyl cyanide 1-AdamC N Dimethylcyanamide Me2 NC N Water, alcohols Perfluoro-tert-butanol (CF3 )3 COH 1,1,1,3,3,3-Hexafluoro-2-propanol (CF3 )2 CHOH

LiCBb

Ref.

134.4 136.9 137.7 139.4 140.0 140.8 143.7 147.2 149.6 154.9 157.1 159.6 160.6 162.1 163.4 167.8 174.5 178.5 205.0

[133] [133] [133] [133] [133] [133] [133] [133] [133] [133] [133] [133] [133] [133] [133] [133] [133] [133] [59]

177.5

[133]

89.2 108.6 109.7 110.2 112.2 116.0 123.1 123.2 137.4 142.1 143.7 147.1 147.8 148.3 148.6 149.6 150.0 152.1 154.0 157.0 159.6 162.9

[133] [133] [133] [133] [133] [133] [133] [133] [133] [133] [133] [133] [133] [133] [133] [133] [133] [133] [133] [133] [133] [133]

85.1 99.7

[133] [133]


349


Lewis base

Formula

LiCBb

Ref.

88 89 90 91 92 93 94 95 96 97 98 99 100 101 102

Water 2,2,2-Trifluoroethanolc Methanol Trichloroethanolc Ethanol n-Propanol 2-Propanol Isobutanol n-Butanol 2,2-Dimethyl-1-propanol tert-Butanol 2-Butanol Cyclohexanemethanol Benzyl alcohol 2-Methoxyethanolc Ethers 1,1,1,3,3,3-Hexafluoro-2-propyl methyl ether 2,2,2-Trifluoroethyl vinyl ether Bis(2,2,2-trifluoroethyl) ether Dimethyl ether 2,2,2-Trifluoroethyl methyl ether Anisole 1,4-Dioxane Tetrahydrofuran Diethyl ether tert-Butyl methyl ether 2-Methyltetrahydrofuran Methyl trimethylsilyl ether Di-n-propyl ether 2,5-Dimethyltetrahydrofuran tert-Butyl ethyl ether Diisopropyl ether Di-n-butyl ether 1,2-Dimethoxyethanec 1,3-Dimethoxypropanec 1,4-Dimethoxybutanec 1,5-Dimethoxypentanec 1,6-Dimethoxyhexanec 1,7-Dimethoxyheptanec 1,8-Dimethoxyoctanec 1,9-Dimethoxynonanec Carbonyl compounds Ketones Carbonyl difluoride Hexafluoroacetone

H2 O CF3 CH2 OH MeOH CCl3 CH2 OH EtOH n-PrOH i-PrOH i-BuOH n-BuOH neo-C5 H9 OH t-BuOH s-BuOH c-C6 H11 CH2 OH PhCH2 OH MeOCH2 CH2 OH

103.6 111.1 119.5 127.4 127.4 131.6 135.4 136.2 137.4 138.7 139.5 139.5 143.7 150.0 178.6

[133] [133] [133] [133] [133] [133] [133] [133] [133] [133] [133] [133] [133] [133] [133]

(CF3 )2 CHOMe

109.8

[133]

CF3 CH2 OCH CH2 (CF3 CH2 )2 O Me2 O CF3 CH2 OMe PhOMe

114.7 122.3 123.6 124.0 126.6 127.0 137.2 139.7 143.3 143.7 144.9 145.8 146.7 148.3 148.7 152.9 187.9 191.7 193.3 195.8 199.2 202.1 200.4 200.8

[133] [133] [133] [133] [133] [133] [133] [133] [133] [133] [133] [133] [133] [133] [133] [133] [133] [74] [74] [74] [74] [74] [74] [74]

77.1 80.0

[133] [133]

103 104 105 106 107 108 109 110 111 112 113 114 115 116 117 118 119 120 121 122 123 124 125 126 127

128 129

a

c-(CH2 )4 O Et2 O t-BuOMe a

Me3 SiOMe n-Pr2 O

a

t-BuOEt i-Pr2 O n-Bu2 O (CH2 OMe)2 CH3 O(CH2 )3 OCH3 CH3 O(CH2 )4 OCH3 CH3 O(CH2 )5 OCH3 CH3 O(CH2 )6 OCH3 CH3 O(CH2 )7 OCH3 CH3 O(CH2 )8 OCH3 CH3 O(CH2 )9 OCH3

F2 CO (CF3 )2 CO

(Continued)

350



Lewis base

Formula

LiCBb

Ref.

130 131 132 133 134 135 136 137 138 139 140 141 142 143

Perfluorodiisopropyl ketone 1,1,3,3-Tetrafluoroacetone 1,1,1-Trifluoroacetone 1,1,1,5,5,5-Hexafluoro-2,4-pentanedione 1,1,1-Trifluoro-2,4-pentanedione Acetone Butan-2-one Pentan-3-one Methyl cyclopropyl ketone 2,4-Dimethylpentan-3-one 4-Methylacetophenone Dicyclopropyl ketone Isophorone Acetylacetonec,e Aldehydes Trifluoroacetaldehyde Formaldehyde Trichloroacetaldehyde Acetaldehyde Propionaldehyde Butyraldehyde Valeraldehyde Hexaldehyde Heptaldehyde Benzaldehyde Carboxylic acid, esters, carbonates 2,2,2-Trifluoroethyl trifluoroacetate Methyl trifluoroacetate Methyl chloroformate Ethyl trifluoroacetate Methyl formate Acetic acid Ethyl formate n-Propyl formate n-Butyl formate Ethyl perfluoropivalatec Methyl acetate Ethyl acetate Methyl propionate Dimethyl terephthalate Methyl benzoate Dimethyl carbonate Dimethyl isophthalate Ethyl 2,2-dimethylpropanoate Dimethyl phthalatec Thioesters S-Methyl trifluorothioacetate S-Methyl thioacetate

[(CF3 )2 CF]2 CO (CHF2 )2 CO CF3 COMe (CF3 CO)2 CH2 CF3 COCH2 COMe Me2 CO MeCOEt Et2 CO c-PrCOMe i-Pr2 CO 4-MeC6 H4 COMe c-Pr2 CO MeC(OH) CHCOMe

91.8 103.1 113.2 114.4 147.9 147.9 150.8 153.8 156.7 157.1 159.6 160.9 173.9 180.5

[133] [133] [133] [133] [133] [133] [133] [133] [133] [133] [133] [133] [133] [133]

CF3 CHO H2 CO CCl3 CHO MeCHO EtCHO n-PrCHO n-BuCHO C5 H11 CHO C6 H13 CHO PhCHO

91.3 106.5 114.0 133.3 137.4 139.5 141.6 144.0 144.8 158.0

[133] [133] [133] [133] [133] [133] [133] [133] [133] [133]

CF3 COOCH2 CF3 CF3 COOMe ClCOOMe CF3 COOEt HCOOMe MeCOOH HCOOEt HCOO-n-Pr HCOO-n-Bu (CF3 )3 CCOOEt MeCOOMe MeCOOEt EtCOOMe C6 H4 -1,4-(COOCH3 )2 C6 H5 COOMe (MeO)2 CO C6 H4 -1,3-(COOCH3 )2 t-BuCOOEt C6 H4 -1,2-(COOCH3 )2

107.7 121.0 121.1 128.2 135.8 137.0 142.1 143.7 143.7 144.4 147.4 150.8 151.9 152.1 154.8 155.0 157.2 162.8 196.7

[133] [133] [133] [133] [133] [133] [133] [133] [133] [133] [133] [133] [133] [159] [159] [133] [159] [133] [159]

CF3 COSMe MeCOSMe

125.2 141.5

[133] [133]

144 145 146 147 148 149 150 151 152 153 154 155 156 157 158 159 160 161 162 163 164 165 166 167 168 169 170 171 172 173 174

a


351

Table 6.4 (Continued) No. 175 176 177 178 179 180 181 182 183 184 185

186 187 188 189 190 191 192 193 194 195 196 197 198 199 200 201 202 203 204 205 206 207 208 209 210 211 212

Lewis base

Formula

Amides, carbamates Trifluoroacetamide CF3 CONH2 Dimethylcarbamoyl cyanide NCCONMe2 Formamide HCONH2 N-Methylformamide HCONHMe N,N-Dimethyl-2,2,2CF3 CONMe2 trifluoroacetamide Acetamide MeCONH2 N,N-Dimethylmethylcarbamate MeOCONMe2 N-Methylacetamide MeCONHMe N,N-Dimethylformamide HCONMe2 N,N-Dimethylacetamide MeCONMe2 Imide 2,2,2-Trifluoro-N(CF3 CO)2 NH (trifluoroacetyl)acetamidef Nitroalkane Nitromethane MeNO2 Sulfur dioxide, sulfinyl and sulfonyl compounds Sulfur dioxide SO2 a Glycol sulfate Dimethyl sulfate (MeO)2 SO2 a Glycol sulfite Methyl 4-nitrophenyl sulfone 4-NO2 C6 H4 SO2 Me Methyl methanesulfonate MeOSO2 Me S-Methyl methanethiosulfonate (MeS)SO2 Me a 1,3-Propanesultone Dimethyl sulfone Me2 SO2 Phenyl methylsulfonate C6 H5 OSO2 Me Tetramethylene sulfone c-(CH2 )4 SO2 a 1,3-Propanesultine Methyl phenyl sulfone PhSO2 Me Methyl 4-tolyl sulfone 4-MeC6 H4 SO2 Me Diphenyl sulfone Ph2 SO2 Dimethyl sulfoxide Me2 SO Methyl phenyl sulfoxide PhSOMe Tetramethylene sulfoxide c-(CH2 )4 SO Diphenyl sulfoxide Ph2 SO Phosphoryl compounds Phosphoryl chloride Cl3 PO Dimethyl phosphonate (MeO)2 POH Trimethyl phosphate (MeO)3 PO Dimethyl methylphosphonate (MeO)2 POMe Diisopropyl phosphonate (i-PrO)2 POH Diethyl chloromethylphosphonate CH2 ClPO(OEt)2 4-(Trifluoromethyl)phenyl 4-CF3 C6 H4 OPOPh2 diphenylphosphinate

LiCBb

Ref.

141.9 144.5 157.4 165.9 166.2

[133] [133] [133] [133] [133]

166.6 167.2 173.4 173.7 179.1

[133] [133] [133] [133] [133]

147.5

[133]

125.9

g

76.3 137.7 142.5 148.7 150.3 151.5 152.5 153.2 155.4 157.5 163.4 163.7 164.5 168.4 169.5 175.1 179.3 180.0 183.3

[133] [133] [133] [133] [133] [133] [133] [133] [133] [133] [133] [133] [133] [133] [133] [133] [133] [133] [133]

145.2 177.6 183.1 183.7 184.2 184.7 185.0

[133] [133] [133] [133] [133] [133] [133] (Continued)

352



Lewis base

Formula

LiCBb

Ref.

213 214 215 216 217 218 219 220 221 222

Diethyl methylphosphonate Triethyl phosphate Methyl methylphenylphosphinate Triphenyl phosphate 4-Fluorophenyl diphenylphosphinate Trimethylphosphine oxide Triethylphosphine oxide Phenyl diphenylphosphinate Hexamethylphosphoric triamide Triphenylphosphine oxide Thiols Methanethiol Ethanethiol 2-Propanethiol 1-Propanethiol 2-Methyl-1-propanethiol 2-Methyl-2-propanethiol 1-Butanethiol Thioethers Trifluoromethyl disulfide Methyl sulfide Ethyl methyl sulfide Tetrahydrothiophene Pentamethylene sulfide Ethyl sulfide Isobutyl methyl sulfide n-Propyl sulfide Isopropyl sulfide n-Butyl sulfide tert-Butyl sulfide Haloalkanes Chloromethane Difluoromethane Dichloromethane 2-Iodopropane 2-Chloropropane 2-Bromopropane Fluoromethane Fluoroethane 2-Fluoropropane Miscellaneous Tetraphosphorus Phosphine Glycine

(EtO)2 POMe (EtO)3 PO (MeO)MePOPh (C6 H5 O)3 PO (4-FC6 H4 O)POPh2 Me3 PO Et3 PO C6 H5 OPOPh2 (Me2 N)3 PO Ph3 PO

187.9 188.4 188.4 188.8 190.4 190.9 195.0 195.9 198.4 198.4

[133] [133] [133] [133] [133] [133] [133] [133] [133] [133]

MeSH EtSH i-PrSH n-PrSH i-BuSH t-BuSH n-BuSH

85.1 89.8 93.9 94.4 99.4 99.8 100.6

[133] [133] [133] [133] [133] [133] [133]

(CF3 S)2 Me2 S EtSMe c-(CH2 )4 S c-(CH2 )5 S Et2 S i-BuSMe n-Pr2 S i-Pr2 S n-Bu2 S t-Bu2 S

80.5 98.1 104.8 108.2 108.6 110.7 114.7 121.1 121.1 128.2 130.3

[133] [133] [133] [133] [133] [133] [133] [133] [133] [133] [133]

CH3 Cl CH2 F2 CH2 Cl2 CH3 CHICH3 CH3 CHClCH3 CH3 CHBrCH3 CH3 F CH3 CH2 F CH3 CHFCH3

66.7 73.0 83.0 86.1 88.0 89.7 90.1 103.0 110.5

g

P4 PH3 H2 NCH2 COOH

70.6 (±12.3) 70.6 (±13.8) 171.7

h

223 224 225 226 227 228 229 230 231 232 233 234 235 236 237 238 239 240 241 242 243 244 245 246 247 248 249 250 251 252

g g g g g g g g

h i


353

Table 6.4 (Continued) a

Formulae: N

N N

N

13

16

N

N

17

18 49 50 51 52 53

N N

N H

S

47

48 R4

5

R

N

N R1

R2

R1 H Me H H Me

55 59 60 61

R1 H Me Me H

R2 H H Me Me

R4 H H H Me

42 R3 H H Me H H

R4 H H H Me Me

R5 H H H Me

N

44 R5 H H H H H

54 56 57 58

N

N H

N H

N

O

45 R1 Me Me H Me

R3 H Me Me Me

O

O

46 R4 H H Me Me

R5 Me Me Me Me

NH N O

N H

62

109

113

O

O

O O

O

116

142

188

O

O

O

S

S

S O

O

190

O

O

194

O S O

198

b Uncertainties on relative LiCBs are in general less than 0.5 kJ mol−1 if the values are not too distant, and increase with the difference in LiCB, due to the accumulation of errors. The experimental error on the absolute LiCB anchor point is 8.8 kJ mol−1 . c Bidentate behaviour: Li+ bridges two heteroatoms. d Uncertain site (N and/or O). e In the adduct the ligand may be in the diketone tautomeric form. f Oxygen base in the enol form. g Calculated from Equation 6.10 and data from refs [130, 131]. h Calculated from Equation 6.10 and data from ref. [156]. i Glycine was measured against HCONMe2 . The value is estimated from entropy data at 298 K in ref. [160] assumed to be valid at 373 K and the LiCB (HCONMe2 ) in this table.

The absolute NaCA and NaCB values are reported in Table 6.5 and the relative FTICR NaCBs, anchored to an accurate HPMS value for MeNH2 , are given in Table 6.6. The two NaCB scales for the compounds common to Tables 6.5 and 6.6 are in fair agreement, within the combined uncertainties, the largest discrepancy being for H2 O (8.7 kJ mol−1 ).

6.5.3

Potassium

The majority of KCA and also KCB data are absolute values obtained from CIDT. In a few instances, HPMS and the kinetic method were also used. The K+ data are reported in Table 6.7. As most KCA and KCB values come from absolute measurements, the uncertainties are fairly large and may obscure trends with closely related compounds. This was particularly observed within the series of amides, for which the trend was really indistinct (see, for example, ref. [134]) until relative KCAs were established [135].

354

6.5.4


Aluminium

The Al+ basicities, AlCBs, were mainly measured as relative values using the ICR equilibrium method in 1982 [136] and 1984 [137]. The correlation between the two sets of data (∆AlCB relative to formaldehyde) is very good: ∆AlCB (1984) = (1.0999 ± 0.0082) ∆AlCB (1982) −(3.61 ± 0.48) n = 130, r 2 = 0.9984, s = 0.58 kJ mol−1 .

(6.11)

However the slope is about 10% greater than expected for a perfect accord. This may be attributed to a systematic difference of about 30 K between the true temperatures of each ICR spectrometer. Yet such a large difference seems unlikely, hence the difference between the 1982 and 1984 AlCB scales is rather to be ascribed, at least in part, to the absence of corrections on pressure reading in the 1982 ICR measurements. The more extended 1984 measurements were therefore preferred. Two values from the 1982 scales were added after correcting them with Equation 6.11. The choice for anchoring the relative AlCB scale at an experimental absolute value is limited to formaldehyde [78] and pyridine [138]. However, the relative value of AlCB (pyridine) is uncertain (±8 kJ mol−1 ) due to competitive proton transfer [137]. Moreover, the experimental absolute AlCB value of formaldehyde has a large uncertainty of about 10 kJ mol−1 and its corresponding van’t Hoff plot is centred at 523 K, far from 298 K. Hence high-level ab initio calculations have been performed [139]. The comparison between G3X computed results and available absolute AlCA and AlCB experimental data appears satisfactory [139]. In particular, the G3X AlCB value of formaldehyde (93.6 kJ mol−1 ) and the experimental value (90.3 ± 10.5 kJ mol−1 ) compare fairly well. Consequently, the ab initio AlCB value of formaldehyde was chosen as anchor. The AlCB scale anchored at H2 CO is given in Table 6.8. The few available absolute AlCA values are collected in Table 6.9.

6.5.5

Manganese

As for Al+ , there are two methods of determination: ICR for measuring relative manganese cation basicities, ∆MnCBs (discussed in terms of ‘ligand binding energies’ in ref. [140] but actually Gibbs energies), and CIDT for absolute measurements of manganese cation affinities, MnCAs. The values obtained by the latter method are listed in Table 6.10. Unfortunately, there is no compound for which an experimentally derived absolute MnCB value could be used for anchoring the relative values. The most accurate experimental MnCA value at 298 K corresponds to ammonia, but no MnCB was derived. With its small size, the adduct Mn+ NH3 is amenable to high-level calculations of bonding energetics. Such calculations, at the CCSD(T)/aug-cc-pVTZ level, leading to MnCA and MnCB, were performed for NH3 [139]. The MnCA(NH3 ) value compared favourably with the experimental value [139], therefore the calculated MnCB(NH3 ) was chosen as the anchor value for the relative ICR scale. The anchored MnCB data are given in Table 6.11. The

Methylamine

Propylamine CH3 CH2 CH2 NH2 2-Aminoethanolc H2 NCH2 CH2 OH Six-membered aromatic N-heterocycles 1,3,5-Triazine 1,3,5-N3 C3 H3 Pyrimidine 1,3-N2 C4 H4 Pyrazine 1,4-N2 C4 H4 Pyridine C5 H5 N 2-Methylpyridine 2-MeC5 H4 N 3-Methylpyridine 3-MeC5 H4 N

13

14 15

16 17 18 19 20 21

12 MeNH2

NH3

c-(CH)4 NMe

1-Methylpyrrole Amines Ammonia

11

c-(CH)4 NH

CH4 CO H2 C CH2 C6 H5 F C6 H5 Cl C6 H5 Br C6 H5 I C6 H6 C6 H5 OH

Carbon bases Methane Carbon monoxide Ethylene Fluorobenzene Chlorobenzene Bromobenzene Iodobenzene Benzene Phenolb

Pyrrole

Formula

Lewis base

10

1 2 3 4 5 6 7 8 9

No. 12.6 1.1 ± 7.7 19.6 ± 4.4 42.7 ± 7.9 45.1 ± 4.1 52.0 ± 5.0 57.2 ± 5.7 59.3 ± 4.3 69.7 ± 3.4 75.0 ± 7.9 71.6 ± 7.7 69.2 ± 7.4 80.8 ± 6.6 81.4 ± 5.4 79.9 ± 2.6 83.4 ± 2.6 81.8 ± 2.6 90 ± 6 145 ± 7 61.5 ± 3.0 76.0 ± 3.9 80.7 ± 3.5 100.0 ± .9 99.7 ± 6.4 105.0 ± 6.2

106.2 ± 5.4 107.1 ± 0.8 110.0 ± 0.8 110.0 ± 0.8 124 ± 6 177 ± 7 89.0 ± 3.0 103.5 ± 3.9 108.2 ± 3.5 127.5 ± 2.9 129.2 ± 4.7 134.2 ± 4.4

NaCB

30.1 33.3 ± 7.7 44.6 ± 4.4 70.7 ± 4.1 74.7 ± 3.5 81.5 ± 3.3 87.0 ± 4.4 90.0 ± 4.3 100.3 ± 3.4 103.6 ± 3.9 103.1 ± 7.7 103.7 ± 4.8 113.1 ± 3.4

NaCA

(Continued)

[57] [57] [57] [57] [53] [53]

[162] [164] [162],g [164] [165] [165]

[161],f [162] [162] [42] [163] [163] [163] [162] [162] [151] [49] [48] [48]

Ref.

Table 6.5 Sodium cation affinity (NaCA) and basicity (NaCB) scales (kJ mol−1 ) at 298 K from CIDT measurements (unless indicated otherwise).

n-Butanol 2-Propanol tert-Butanol 2-Butanol Ethers Dimethyl ether 1,2-Dimethoxyethanec 12-crown-4d

36 37 38

39 40 41 42

43 44 45 a

Me2 O (MeOCH2 )2

n-BuOH i-PrOH t-BuOH s-BuOH

EtOH i-BuOH n-PrOH

Ethanol 2-Methyl-1-propanol n-Propanol

34 35

27 28 29 30 31 32 33

4-Methylpyridine 4-MeC5 H4 N 3-Aminopyridine 3-H2 NC5 H4 N 4-Aminopyridine 4-H2 NC5 H4 N 2-Aminopyridinec 2-H2 NC5 H4 N Pyridazinec 1,2-N2 C4 H4 Five-membered aromatic N-heterocycles a 2H-1,2,3-Triazole a 2H-Tetrazole a 1H-1,2,4-Triazole a Pyrazole a 1-Methylpyrazole a Imidazole a 1-Methylimidazole Water, alcohols Water H2 O Methanol MeOH

22 23 24 25 26

Formula

Lewis base

No.


69.7 89.2 97.0 101.1 ± 10.3 103.6 ± 6.4 113.3 ± 5.2 132.6 ± 7.6 74.4 ± 7.5 65.7 ± 5.7 74.8 ± 2.6 75.3 ± 3.7 76.8 ± 5.7 80.4 ± 4.1 85 ± 4 82.2 ± 4.7 85.5 ± 4.3 89.0 ± 4.1 88.4 ± 5.1 66.4 ± 4.8 128.7 ± 3.9

98.1 ± 7.5 93.2 ± 5.7 100.4 ± 0.8 103.3 ± 3.7 106.4 ± 5.7 109.3 ± 4.1 114 ± 4 110.3 ± 4.7 114.4 ± 4.3 117.6 ± 4.1 118.4 ± 5.1 92.6 ± 4.8 159.9 ± 3.9 254 ± 13

106.9 ± 6.0 109.2 ± 6.0 119.7 ± 6.6 120.8 ± 7.4 131.1 ± 3.2

134.4 ± 4.0 136.7 ± 4.0 146.8 ± 4.8 148.0 ± 6.0 160.4 ± 3.2 96.6 115.4 124.6 129.0 ± 8.5 132.4 ± 2.9 140.9 ± 5.2 161.8 ± 5.0

NaCB

NaCA

[162] [162] [70]

[162] [162] [164] [162] [166] [162] [165] [166] [162] [166] [166]

[56],h [56],h [56],h [48] [48] [48, 162] [48]

[53] [52] [52] [52] [57]

Ref.

Butan-2-one Formamide N-Methylformamide Acetamide N,N-Dimethylformamide

N-Methylacetamide N,N-Dimethylacetamide Succinic diamidec Rare gas Argon Amino acidse Glycine Alanine Valine Leucine Cysteine Isoleucine Serine Proline Threonine Phenylalanine Tyrosine Aspartic acid

52 53 54 55 56

57 58 59

61 62 63 64 65 66 67 68 69 70 71 72

60

49 50 51

15-crown-5d 18-crown-6d Dibenzo-18-crown-6d Carbonyl compounds Acetaldehyde Propionic acid Acetone

46 47 48

a

MeCHOHCH(NH2 )COOH PhCH2 CH(NH2 )COOH 4-HOC6 H4 CH2 CH(NH2 )COOH HOOCCH2 CH(NH2 )COOH

HCH(NH2 )COOH MeCH(NH2 )COOH i-PrCH(NH2 )COOH i-PrCH2 CH(NH2 )COOH HSCH2 CH(NH2 )COOH EtC(Me)HCH(NH2 )COOH HOCH2 CH(NH2 )COOH

Ar

MeCONHMe MeCONMe2 (H2 NCOCH2 )2

MeCOEt HCONH2 HCONHMe MeCONH2 HCONMe2

MeCHO EtCOOH Me2 CO

a

a

a

−0.9 ± 8.7

17.3 ± 8.7 161 ± 8 167 ± 8 173 ± 8 175 ± 8 175 ± 8 176 ± 8 192 ± 8 196 ± 8 197 ± 8 198 ± 8 201 ± 8 203 ± 8

(Continued)

[169] [169] [169] [169] [169] [169] [169] [169] [169] [169] [169] [169]

[162]

[162] [165] [162] [164] [165] [135],j [135],j [135],j [162] [135] [135],j [135],j [134]

88.7 ± 3.4 92 ± 6 105.7 ± 4.1 102.7 ± 6.6 105 ± 7 110.5 ± 2.7 118.7 ± 2.7 119.0 ± 2.6 130.6 ± 3.7 125.9 ± 2.5 126.2 ± 2.6 132.8 ± 2.7

114.4 ± 3.4 119 ± 6 131.3 ± 4.1 128.9 ± 2.1 132 ± 7 142.4 ± 3.8 150.6 ± 3.8 150.9 ± 3.7 157.8 ± 3.7 158.1 ± 3.7 164.7 ± 3.8 186.2

[167] [167] [168],i

298 ± 18 300 ± 19 299 ± 26

46

O

O

O

O

47

O

N

O

O

N

H

O O

48

O

O

O

O

N R

N

N

H

29

N

N

68

H

N

31

30

COOH

74

R: Me

R: H N

H

N

NH

CH CHCOOH

R

N

HN C(NH2 )NH(CH2 )3 CH(NH2 )COOH

a

H2 NCOCH2 CH(NH2 )COOH H2 NCOCH2 CH2 CH(NH2 )COOH

HOOCCH2 CH2 CH(NH2 )COOH

a

Formula

N

77

H

N

33

32

NH

CH CHCOOH

R: Me

R: H

204 ± 8 210 ± 8 217 ± 8 222 ± 8 228 ± 8 242 ± 8

NaCA

45

O

O

O

O

NaCB

[169] [169] [94] [94] [94] [94]

Ref.

b

O and π base. c Bidentate behaviour. d Polydentate complexes. e See Section 6.2 for structures. f HPMS value. g HPMS NaCA = 110.9 kJ mol−1 reported in ref. [162] is actually 110.0 kJ mol−1 . The value of NaCB is corrected accordingly. h Although not explicitly given, uncertainties are expected to be in the 4–8 kJ mol−1 range of uncertainties of NaCA (0 K) values. i Obtained by bracketing using CIDT values as references. j Kinetic method data linked to NaCA (HCONMe2 ) = 157.8 kJ mol−1 . Relative NaCAs and NaCBs are considered equal for this series of amides (i.e. constant entropy is assumed). For consistency with Table 6.6, NaCB (amides) are linked to the FTICR value of HCONMe2 = 125.9 kJ mol−1 [106].

O

O

O

N

O

H

28

N

N

27

N

N

Formulae:

Glutamic acid Tryptophan Asparagine Glutamine Histidine Arginine

73 74 75 76 77 78

a

Lewis base

No.



359

Table 6.6 Sodium cation basicity (NaCB) scale (kJ mol−1 at 298 K), from FTICR ligand-exchange equilibria [106] unless noted otherwise, anchored to the HPMS and ab initio values for methylamine. No. 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37

Lewis base Carbon π bases 2-Methylpropene Butadiene Furan Chlorobenzene Bromobenzene Benzene Phenolc Amines Ammonia Trimethylamine Methylamine Dimethylamine Ethylamine Diethylamine Nitriles Acetonitrile Propionitrile Water, alcohols Water Methanol Ethanol n-Propanol n-Butanol 2-Methyl-1-propanol 2-Propanol 2-Butanol tert-Butanol 2-Fluoroethanold Ethers Dimethyl ether Diethyl ether Diethoxymethane 1,2-Dimethoxyethaned Carbonyl compounds Methyl acetate Acetone β-Butyrolactone Pentan-2-one Cyclopenten-2-one 3,3-Dimethylacrolein Formamide N-Methylformamide

Formula (CH3 )2 C CH2 H2 C CHCH CH2

NaCBb

C6 H5 Cl C6 H5 Br C6 H6 C6 H5 OH

41.8 43.1 49.0 51.9 53.6 65.7 69.9

NH3 (CH3 )3 N CH3 NH2 (CH3 )2 NH CH3 CH2 NH2 (CH3 CH2 )2 NH

77.8 79.5 81.6 82.0 87.0 91.6

CH3 C N CH3 CH2 C N

98.7 102.9

a

H2 O CH3 OH CH3 CH2 OH CH3 CH2 CH2 OH CH3 CH2 CH2 CH2 OH (CH3 )2 CHCH2 OH (CH3 )2 CHOH CH3 CH2 CHOHCH3 (CH3 )3 COH FCH2 CH2 OH

65.7 72.4 79.5 81.6 82.4 82.4 85.4 87.4 89.5 98.7

(CH3 )2 O (C2 H5 )2 O C2 H5 OCH2 OC2 H5 CH3 OCH2 CH2 OCH3

73.6 89.1 115.5 133.1

CH3 COOCH3 CH3 COCH3

97.5 100.8 100.8 104.2 107.9 108.8 110.5e 118.7e

a

CH3 COCH2 CH2 CH3

a

HCOCH C(CH3 )2 HCONH2 HCONHCH3

(Continued)

360



Lewis base

Formula

NaCBb

38 39 40 41 42

Acetamide N,N-Dimethylformamide N-Methylacetamide N,N-Dimethylacetamide Acetic anhydrided Sulfinyl compound Dimethyl sulfoxide Chloroalkanes Chloroethane 1-Chloropropane 1-Chlorobutane 2-Chloropropane 1-Chloro-2-methylpropane 2-Chlorobutane 2-Chloro-2-methylpropane Bromoalkanes 1-Bromobutane 2-Bromo-2-methylpropane Sulfide Methyl sulfide

CH3 CONH2 HCON(CH3 )2 CH3 CONHCH3 CH3 CON(CH3 )2 CH3 COOCOCH3

119.0e 125.9 126.2e 132.8e 130.1e

(CH3 )2 SO

129.7

43 44 45 46 47 48 49 50 51 52 53 a

CH3 CH2 Cl CH3 CH2 CH2 Cl CH3 CH2 CH2 CH2 Cl CH3 CHClCH3 (CH3 )2 CHCH2 Cl CH3 CH2 CHClCH3 (CH3 )3 CCl

45.2 48.5 50.6 51.5 51.5 54.4 57.3

CH3 CH2 CH2 CH2 Br (CH3 )3 CBr

51.0 57.3

(CH3 )2 S

59.4

Formulae O O

3

CH3

32

O

34

b Uncertainties on relative values are about 0.4 kJ mol−1 at the lower part of the scale, with an upper limit of 2.5 kJ mol−1 for the largest values (see text); the uncertainty on the absolute value for methylamine (81.6 kJ mol−1 ) is 0.8 kJ mol−1 . c Polyfunctional base (O and π ). d Na+ bridges two basic sites. e Kinetic method values from ref. [135]. NaCB values for these amides are linked to the FTICR value NaCB (HCONMe2 ) = 125.9 kJ mol−1 .

uncertainties stated in the original work are ±0.8 kJ mol−1 on relative values in a restricted range, hence it can be estimated that they may add up to about ±4.5 kJ mol−1 for the total range.

6.5.6

Cyclopentadienylnickel

The η5 -cyclopentadienylnickel cation (CpNi+ ) basicity scale, CpNiCB, was established by ICR mass spectrometry [141]. Although the equilibrium data were discussed by the authors as relative affinities, by assuming negligible variations in entropies (as in the case of Al+ and Mn+ ), these values are really relative basicities. To anchor them experimentally, our

Gas-Phase Cation Affinity and Basicity Scales Table 6.7

No. 1 2 3 4 5

Potassium cation affinity (KCA) and basicity (KCB) scales (kJ mol −1 ) at 298 K.

Lewis base

Formula

Carbon π bases Fluorobenzene Chlorobenzene Bromobenzene Iodobenzene Benzene

C6 H5 F C6 H5 Cl C6 H5 Br C6 H5 I C6 H6

6 7 8 9 10

Phenolb Anisolec Toluene Naphthalene Pyrroled

C6 H5 OH C6 H5 OMe C6 H5 Me

11 12 13

1-Methylpyrroled Anilined Indolee Amines Ammonia

c-(CH)4 NMe C6 H5 NH2

14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38

361

a

c-(CH)4 NH

a

NH3

Methylamine MeNH2 Dimethylamine Me2 NH Trimethylamine Me3 N n-Propylamine n-PrNH2 Ethylenediaminef H2 NCH2 CH2 NH2 Six-membered aromatic N-heterocycles 1,3,5-Triazine 1,3,5-N3 C3 H3 Pyrazine 1,4-N2 C4 H4 Pyrimidine 1,3-N2 C4 H4 Pyridine C5 H5 N 2-Methylpyridine 2-MeC5 H4 N 4-Methylpyridine 4-MeC5 H4 N 3-Methylpyridine 3-MeC5 H4 N 3-Aminopyridine 3-H2 NC5 H4 N 2-Aminopyridine 2-H2 NC5 H4 N 4-Aminopyridine 4-H2 NC5 H4 N g Pyridazine 1,2-N2 C4 H4 Five-membered N-heterocycles 2-H-1,2,3-Triazoleh a a 1-H-Pyrazoleh 1-H-1,2,4-Triazoleh 2-H-Tetrazoleh 1-Methylpyrazole 1-H-Imidazoleh 1-Methylimidazole Nitrile Acetonitrile

a a a a a

MeC N

KCA

KCB

Ref.

55.5 ± 3.6 56.6 ± 3.3 66.2 ± 3.0 67.5 ± 3.2 74.2 ± 4.1 76.6 74.6 ± 3.8 79.4 ± 3.3 80.8 ± 5.3 81.2 ± 5.3 85.1 ± 3.6 79.9 ± 3.9 89.1 ± 6.3 95.4 ± 4.2 100.6 ± 3.9

29.4 ± 7.8 28.7 ± 5.9 39.1 ± 5.2 39.4 ± 5.2 45.9 ± 7.1 48.5 47.8 ± 8.1 50.2 ± 8.4 48.9 ± 9.4 52.3 ± 9.4 54.6 ± 6.8 49.4 ± 3.9 57.9 ± 8.6 66.1 ± 4.2 70.7 ± 4.0

[42] [163] [163] [163] [39] [170] [151] [152] [149] [150] [48] [49] [48] [171] [153]

82 ± 8 84 ± 8 75 ± 8 79.9 ± 4.2 81.6 ± 4.2 83.7 ± 4.2 91.2 ± 4.2 107.5 ± 4.2

55 ± 11 55 ± 9 49 ± 8 53.1 ± 4.2 54.8 ± 4.2 54.4 ± 4.2 59.4 ± 4.2 79.5 ± 4.2

[172] [172, 173] [171, 172] [171] [171] [171] [171] [171]

55.6 ± 3.0 67.6 ± 3.6 69.7 ± 4.3 90.6 ± 3.9 86.6 ± 4.2 98.5 ± 3.5 99.0 ± 4.0 100.1 ± 3.5 101.3 ± 3.4 104.6 ± 3.8 108.8 ± 3.4 130.9 ± 2.6

29.6 ± 3.0 41.6 ± 3.6 43.7 ± 4.3 64.6 ± 3.9 63.6 ± 4.2 69.6 ± 5.7 73.0 ± 6.0 72.5 ± 5.7 75.4 ± 5.9 76.2 ± 5.8 82.8 ± 5.9 102.2 ± 2.6

[57] [57] [57] [57] [171] [53] [53] [53] [52] [52] [52] [57]

55.6 ± 5.5 79.9 ± 3.9 84.2 ± 3.3 87.5 ± 4.5 89.6 ± 4.6 94.8 ± 3.6 109.7 ± 5.7 117.7 ± 2.7

30.4 ± 5.5 49.4 ± 3.9 58.0 ± 6.7 61.5 ± 4.5 64.0 ± 4.6 67.7 ± 6.9 83.5 ± 7.3 90.5 ± 6.4

[56] [49] [48] [56] [56] [48] [56] [48]

102.1 ± 1.7

75.3 ± 2.1

[174] (Continued)

362


Table 6.7

(Continued)

No. Lewis base 39 40

Water, alcohols Water

Formula H2 O

KCA

KCB

Ref.

74.9 72.8 ± 2.1 91.6 ± 2.1

47.7 46.9 ± 1.1 52.3 ± 1.3

[175],k [176] [177],l

74.0 ± 4.0 87.0 ± 4.2 93.3 ± 4.2 120.0 ± 4.0 191 ± 11 205 ± 15 221 ± 19 235 ± 13

43.0 ± 4.0 56.1 ± 4.2 62.6 ± 4.2 86.6 ± 4.0

[178],m [171] [171] [178],n [178] [13] [168] [13]

Methanol Ethers Dimethyl ether

MeOH

Et2 O (MeOCH2 )2

48

Diethyl ether 1,2-Dimethoxyethaneg 12-Crown-4i 15-Crown-5i Dibenzo-18-crown-6i 18-Crown-6i Carbonyl compounds Acetone

49 50 51

Formamide N-Methylformamide Acetamide

HCONH2 HCONHMe MeCONH2

52

N-Methylacetamide

MeCONHMe

53

N,N-Dimethylformamide HCONMe2

54

N,N-Dimethylacetamide MeCONMe2

55

Sulfonyl compound Dimethyl sulfoxide

Me2 SO

130.1 146.4 ± 12.6

104.6 ± 12.6

[134],o [179]

56 57

Phosphoryl compounds Trimethyl phosphate Triethyl phosphate

(MeO)3 PO (EtO)3 PO

135.3 ± 3.9 135.3 ± 5.8

108.6 ± 4.9 108.0 ± 8.9

[102] [101]

58

Amino acidsj Glycine

HCH(NH2 )COOH

59 60 61 62

Alanine Valine Leucine Isoleucine

63

41 42 43 44 45 46 47

Me2 O

a a

a a

Me2 CO

102.1 108.8 111.1 117.2 124.3 117.8 127.2 123.8 123.4 129.7 121.3 129.7 128.7

79.5

96.2 100.4 ± 8.4

[134],o [179],p [135],q [135],q [134],o [135],q [134],o [135],q [134],o [179],p [134],o [179] [135],q

125.5 119.3 MeCH(NH2 )COOH 123.6 i-PrCH(NH2 )COOH 128.0 i-PrCH2 CH(NH2 )COOH 129.3 EtC(Me)HCH(NH2 )COOH 129.9

[134],o [180] [180] [180] [180] [180]

Nucleobasesj Uracil

a

64

Thymine

a

65

Adenine

a

66 67

Cytosine Guanine

a

[60] [61] [60] [61] [60] [61] [60] [60]

a

101 105.0 ± 2.8 102 104.6 ± 3.8 106 97.4 ± 3.2 110 117

Gas-Phase Cation Affinity and Basicity Scales Table 6.7

(Continued)

No.

Lewis base

Formula

68 69 70

Miscellaneous Argon Carbon monoxide Carbon dioxide

Ar CO CO2

a

KCA

KCB

15.4 ± 7.0 19.0 ± 5.0 35.6

363

Ref. [161],r [161, 181] [161]

15.9

Formulae: N N N H

9

N

13

N

36 37

N

H

32 35

R: H R: Me

R

O O

O

O O

45

H

R

N

O

O

O

O H

O

O

47

NH2

N H

O

O

46

H2N

N

N

N

N

N

H

H

N H

O

N

N

O

O

N

N

O

O

O

NH2

R: H R: Me

34

O

O

O

H

N

H

O

O

N N

N

33

44

63 64

N N

N R

31

N

O

N

R: H R: Me

65

66

67

The favoured sites are the π ring from calculated KCA and oxygen in terms of KCB. c π complex slightly favoured. d Structure of complex not studied. e π base on the six-membered ring. f Chelate. g Bidentate behaviour. h The more likely complexed tautomer is deduced from calculations. i Polydentate complex. j See Section 6.2 for structure of complex. k Mean of two values. l Value measured relative to water. m KCB estimated from an HPMS entropy of 104 J K mol−1 . n KCB estimated from an HPMS entropy of 112 J K mol−1 . o Uncertainties of about 10 kJ mol−1 . p Uncertainties not given; values rounded to the nearest kcal mol−1 indicate uncertainties of about 4 kJ mol−1 . q Relative affinities obtained by the kinetic method (calibrated on theoretical values) and anchored to the KCA value of HCONMe2 (123.4 kJ mol−1 ). r Value at 0 K, corrected to 298 K in ref. [180]. b

knowledge about the absolute energy of CpNi+ adducts is limited to a dissociation energy of the CpNiNO+ adduct obtained as follows. A mass spectrometric study [142] yielded the ionization energy of the neutral adduct CpNiNO (8.5 eV) and the appearance energy of the CpNi+ fragment (10.5 eV). Combined together, these data give D0 (CpNi+ -NO) = 2 eV = 192.5 kJ mol−1 . Errors of ±20.9 kJ mol−1 [141], then of ±8.4 kJ mol−1 [143],

364


Table 6.8 Aluminium cation basicity (AlCB) scale (kJ mol−1 ), from ICR ligand-exchange equilibria at 298 K (ref. [137] unless noted otherwise) anchored at the ab initio value for formaldehyde. No. 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26

27 28 29 30 31 32 33 34 35 36 37 38 39

Lewis base Six-membered N-heteroaromatics Pyrimidine Pyridine Nitriles Malononitrile Acetonitrile Acrylonitrile Propionitrile Butyronitrile Isobutyronitrile Benzonitrile Dimethylcyanamide Alcohols Methanol Ethanol 1-Propanol 2-Propanol 1-Butanol 2-Methyl-1-propanol, isobutanol 2-Butanol Cyclopentanol tert-Butanol Ethers Dimethyl ether Diethyl ether Tetrahydrofuran tert-Butyl methyl ether Diisopropyl ether 2-Methyltetrahydrofuran 2,2-Dimethyltetrahydrofuran Carbonyl compounds Aldehydes Formaldehyde Acetaldehyde Propanal 2-Methyl-1-propanal Ketones 1,1,1-Trifluoroacetone Cyclobutanone Acetone Butanone 2-Pentanone Cyclopropyl methyl ketone 2,2,4,4-Tetramethyl-3-pentanone 4-Methylacetophenone Acetylacetone

Formula

AlCBb

1,3-N2 C4 H4 C5 H5 N

129.8 ± 9.3c 166.6 ± 8d

N CCH2 C N CH3 C N CH2 CHC N CH3 CH2 C N CH3 CH2 CH2 C N (CH3 )2 CHC N C6 H5 C N (CH3 )2 NC N

91.9 132.1 134.2 140.9 145.5 146.7 153.0 184.0

CH3 OH CH3 CH2 OH CH3 CH2 CH2 OH (CH3 )2 CHOH CH3 CH2 CH2 CH2 OH (CH3 )2 CHCH2 OH CH3 CH2 CHOHCH3

120.4 133.3 140.5 144.2 145.1 145.1 149.7 150.9e 153.9 ± 4e

a

(CH3 )3 COH CH3 OCH3 CH3 CH2 OCH2 CH3 c-(CH2 )4 O (CH3 )3 COCH3 (CH3 )2 CHOCH(CH3 )2 a a

139.2 155.1 162.2 166.8 168.9 169.3 177.3

H2 CO HCOCH3 HCOCH2 CH3 HCOCH(CH3 )2

93.6f 130.0 140.0 145.9

CF3 COCH3 c-(CH2 )3 CO CH3 COCH3 CH3 CH2 COCH3 CH3 CH2 CH2 COCH3 c-(CH2 )3 COCH3 (CH3 )3 CCOC(CH3 )3 4-CH3 C6 H4 COCH3 CH3 COCH2 COCH3

99.0 149.7 157.6 165.1 171.8 183.1 185.6 190.3 191.5


365


Lewis base

Formula

AlCBb

40 41

Dicyclopropyl ketone Isophorone Acid, esters, carbonate Trifluoroacetic acid Methyl trifluoroacetate Ethyl trifluoroacetate Methyl formate Ethyl formate n-Propyl formate n-Butyl formate Methyl acetate Methyl propionate Ethyl acetate Dimethyl carbonate Ethyl propionate Ethyl butyrate Amides, carbamate N,N-Dimethyltrifluoroacetamide Methyl dimethylcarbamate N,N-Dimethylformamide

c-(CH2 )3 COc-(CH2 )3 a

195.7 204.5

CF3 COOH CF3 COOCH3 CF3 COOCH2 CH3 HCOOCH3 HCOOCH2 CH3 HCOOCH2 CH2 CH3 HCOOCH2 CH2 CH2 CH3 CH3 COOCH3 CH3 CH2 COOCH3 CH3 COOCH2 CH3 CH3 OCOOCH3 CH3 CH2 COOCH2 CH3 CH3 CH2 CH2 COOCH2 CH3

90.7 111.6 122.1 140.5 150.1 154.7 156.8 164.3 170.6 172.7 174.8 179.0 182.7

CF3 CON(CH3 )2 CH3 OCON(CH3 )2 HCON(CH3 )2

189.8 204.5 209.5

42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 a

Formulae: O

OH

18

O

O

25

26

41

Uncertainties of ±1 kJ mol−1 on relative values, unless noted otherwise. c CIDT data from Ref. [182]. d In good agreement with the value of 161.1 ± 10.3 kJ mol−1 in ref. [138]. e Calculated from ref. [136] using Equation (6.11). f G3X AlCB used as anchor point, see text. b

were stated without obvious justification. If the thermal energy is neglected, this value can be taken for the CpNiCA of nitrogen monoxide. However, the entropy needed to obtain the corresponding CpNiCB value still remains to be evaluated. It seems safer to anchor the scale to a theoretical value. CCSD(T) calculations [139] on the thermodynamics of dissociation of the CpNiNH3 + adduct yield CpNiCB(NH3 ) = 172.1 kJ mol−1 and CpNiCA(NH3 ) = 212.3 kJ mol−1 . The latter value compares well with the value of 219.3 kJ mol−1 estimated from D0 (CpNi+ NO) = 192.5 kJ mol−1 (assumed to be equal to CpNiCA) and the difference in basicity (26.8 kJ mol−1 ) between NH3 and NO (assumed to be equal to the affinity difference). The CpNiCB scale anchored to the ab initio value for NH3 is given in Table 6.12. The alkyl substituent effect on the CpNiCA of 12 alkylpyridines has been measured by the kinetic method [144]. Unfortunately, the CpNiCB of unsubstituted pyridine has not yet

366


Table 6.9 Aluminium cation affinity (AlCA) scale (kJ mol−1 ) at 298 K unless noted otherwise. No.

Lewis base

Formula

1 2 3 4 5 6 7 8

Water Formaldehyde Benzene Phenol Pyrimidine Pyrrole Indole Pyridine

H2 O H2 CO C6 H6 C6 H5 OH 1,3-N2 C4 H4 c-(CH)4 NH

a

a

C5 H5 N

AlCA

Ref.

104.2 ± 15.1 115.1 ± 10.5 146.0 ± 8.4 154.8 ± 16.7 160.1 ± 6.1 172.0 ± 8.4 188.3 ± 16.7 191.5 ± 10.3

[183],b [78] [184, 185] [185],c [182] [47],d [185],c [138]

Formula:

N H

7 b c d

At 0 K. Uncertain temperature. Estimated from the difference ∆AlCA = 25.9 kJ mol−1 with benzene value.

been determined. Therefore, the alkylpyridine data cannot be included in the CpNiCB scale of Table 6.12 at present.

6.5.7

Copper

The scales relative to the copper monocation correspond to the total dissociation of the double adduct L2 Cu+ into Cu+ and the two ligands L. When studying ligand exchange equilibria L1 Cu+ + L2 L2 Cu+ + L1

(6.12)

for measuring copper cation basicity (CuCB), observing and measuring the simple equilibrium 6.12 is not possible. The reaction of L1 Cu+ with L2 yields mostly L1 Cu+ L2 because of the special stability of the two-ligand complex. If ligands are exchanged from L1 Cu+ L1 , it is not possible to stop the exchange to reaction 6.13: L1 Cu+ L1 + L2 L1 Cu+ L2 + L1

(6.13)

L1 Cu+ L2 + L2 L2 Cu+ L2 + L1

(6.14)

because reaction (6.14):


367

Table 6.10 Manganese cation affinity (MnCA) scale (kJ mol−1 ) at 0 K unless noted otherwise. No. 1 2 3 4 5 6 7 8 9 10 11 a

Lewis base

Formula

Carbon monoxide Carbon monosulfide Ethylene Water Benzene

CO CS H2 C CH2 H2 O C6 H6

Ammonia Acetone Pyrimidine Pyrrole Pyridine Adenine

NH3 Me2 CO 1,3-N2 C4 H4 c-(CH)4 NH C5 H5 N

a

MnCA

Ref.

25.1 ± 9.7 80.1 ± 21.2 90.7 ± 11.6 121.8 ± 5.9 133.2 ± 8.7 144.7 ± 9.6 147 ± 8 159.2 ± 14.5 159.9 ± 9.7 177.4 182.6 ± 8.6 219.8 ± 8.1

[186] [186] [187] [188],b [189] [190] [191],b [192] [182],b [47],c,d [138],b [62],b

Formula: NH2 N

N

N H

N

11 b

Value at 298 K. c π complex, ref. [47]. d Approximate value.

occurs almost simultaneously. For experimental convenience, reactions 6.13 and 6.14 are combined and the double exchange reaction 6.15 is considered: L1 Cu+ L1 + 2 L2 L2 Cu+ L2 + 2 L1

(6.15)

The Gibbs energy and enthalpy of reaction 6.15 provide the two-ligand copper cation affinity and basicity of L2 relative to L1 . The two-ligand affinity and basicity scales are labelled CuCA(2L) and CuCB(2L), respectively. These scales are considered very valuable even though they are not expected to be simply related to the corresponding one-ligand scales. The comparison of CuCA(2L) with CuCA(1L) for a sample of 14 ligands is shown in Figure 6.11. For seven small ligands (CO, H2 O, C2 H4 , Me2 O, Me2 CO, MeCN and NH3 ), the second binding enthalpy is nearly identical with the first one. Accordingly, CuCA(1L) explains almost 99% of the variance of CuCA(2L) and the regression line of Equation 6.16: CuCA(2L) = 1.86 (±0.09) CuCA(1L) +34 (±17) n = 7, r 2 = 0.989, s = 7.7 kJ mol−1 .

(6.16)

has a slope close to two and an intercept hardly significant. However larger and/or bidentate ligands (pyridine, 4,4 -bipyridine, benzene, pyrrole, 1,2-dimethoxyethane, 2,2 -bipyridine and 1,10-phenanthroline) have a significantly weaker second binding enthalpy than the first one, possibly because of repulsive interactions between the two ligands. Consequently, the

368


Table 6.11 Manganese cation basicity (MnCB) scale (kJ mol−1 ) from ICR ligand-exchange equilibria at 298 K (ref. [140] unless noted otherwise), anchored at the ab initio value for ammonia. No. 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18

19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34

Lewis base

Formula

Carbon π base Benzene C6 H6 Amines Ammonia NH3 Trimethylamine (CH3 )3 N Six-membered aromatic N-heterocycles Pyrimidine 1,3-N2 C4 H4 Pyridine C5 H5 N Nitriles Hydrogen cyanide HC N Acetonitrile CH3 C N Propionitrile CH3 CH2 C N Alcohols Methanol CH3 OH Ethanol CH3 CH2 OH n-Propanol CH3 CH2 CH2 OH 2-Methyl-1-propanol CH3 CH(CH3 )CH2 OH tert-Butanol (CH3 )3 COH Ethers Oxirane c-(CH2 )2 O Dimethyl ether (CH3 )2 O 1,4-Dioxane C4 H8 O2 Diethyl ether (CH3 CH2 )2 O Tetrahydrofuran c-(CH2 )4 O Carbonyl compounds Aldehydes Formaldehyde H2 CO Acetaldehyde HCOCH3 Propanal HCOCH2 CH3 Butanal HCOCH2 CH2 CH3 2-Methylpropanal HCOCH(CH3 )2 2,2-Dimethyl propanal HCOC(CH3 )3 Ketones Acetone (CH3 )2 CO 2-Butanone CH3 CH2 COCH3 Esters Methyl formate HCOOCH3 Ethyl formate HCOOCH2 CH3 n-Butyl formate HCOO(CH2 )3 CH3 Methyl acetate CH3 COOCH3 Thiols, sulfides Methanethiol CH3 SH Ethanethiol CH3 CH2 SH 1-Propanethiol CH3 CH2 CH2 SH Methyl sulfide (CH3 )2 S

MnCB 117.9b 118.9c 137.5 130.0 ± 12.0d 152.4 ± 8.6e 99.8 129.4 136.2 109.0 115.6 120.4 124.0 131.1 111.4 116.4 122.8b 133.5 133.5

98.8 117.1 122.6 125.3 125.7 129.8 136.4 140.7 122.2 129.2 133.4 137.0 95.1 101.1 103.9 109.7 (Continued)


369


Lewis base Isocyanate Methyl isocyanate Nucleobase Adenine

35 36 a

Formula

MnCB

CH3 N C O

112.6

a

180.1 ± 9.4f

Formula: NH2 N

N

N H

N

36 b

The symmetry correction (–RTln2) applied in the original article [140] is removed c Ab initio anchoring value [139]. d Ref. [182]. e Ref. [138]. f Ref. [62].

variance of CuCA(2L) explained by CuCA(1L) falls to 93% when the whole sample is considered. The first relative CuCB(2L) scale was built in 1982 from ICR measurements for relatively weak ligands [145]. The upper part of the scale was explored in 1998 by the HPMS method [146]. Two specific series, ring-substituted acetophenones [82] and pyridines [147], were studied in 2007 by the FTICR method.

700 14 13

CuCA(2L) / kJ mol-1

600 7

500

6

10

5 400

12 11

9

4 8 2

300

200 100

1

3

200

300

CuCA(1L) /kJ mol

400

-1

Figure 6.11 Plot of CuCA(2L) versus CuCA(1L). The drawn line is for the relation CuCA(2L) = 2CuCA(1L). Points 1–7 correspond to small ligands and 8–14 correspond to benzene, pyrrole, 1,2-dimethoxyethane, 4,4 -bipyridine, pyridine, 2,2 -bipyridine and 1,10-phenanthroline.

370


Table 6.12 Cyclopentadienylnickel cation basicity (CpNiCB) scale (kJ mol−1 ) from ICR ligand-exchange equilibria at 298 K [141], anchored at the theoretical value for ammonia (see text). No. 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30

Lewis base Carbon base (isonitrile) Methyl isocyanide Amines Ammonia Methylamine Trimethylamine Dimethylamine Phosphines, arsine Phosphine Trimethylarsine Trimethylphosphine Nitriles Hydrogen cyanide Acetonitrile Alcohols Methanol Ethanol 2-Propanol tert-Butanol Ethers Oxirane Dimethyl ether 1,4-Dioxane Tetrahydropyran Tetrahydrofuran Diethyl ether Carbonyl compounds Acetaldehyde Propanal 2-Methylpropanal 2-Propenal 2,2-Dimethylpropanal Acetone Methyl acetate Thiols, sulfides Methanethiol Methyl sulfide Miscellaneous Nitrogen monoxide

Formula

CpNiCBa

MeN C

194.3

NH3 CH3 NH2 (CH3 )3 N (CH3 )2 NH

172.1b 184.2 189.2 190.9

PH3 (CH3 )3 As (CH3 )3 P

143.6 192.6 194.3

HC N CH3 C N

152.4 175.8

CH3 OH CH3 CH2 OH (CH3 )2 CHOH (CH3 )3 COH

144.0 152.0 160.4 165.8

c-(CH2 )2 O (CH3 )2 O C4 H8 O2 c-(CH2 )5 O c-(CH2 )4 O (CH3 CH2 )2 O

144.5 149.5 151.2 c 165.4 165.8 166.2

HCOCH3 HCOCH2 CH3 HCOCH(CH3 )2 HCOCH CH2 HCOC(CH3 )3 (CH3 )2 CO CH3 COOCH3

149.9 154.5 158.7 159.5 160.8 165.4 166.2

CH3 SH (CH3 )2 S

150.3 167.1

NO

145.3

The uncertainties stated in ref. [141] on relative values are ±0.8 kJ mol−1 . See text for a general comment. Anchoring value (see text). c The symmetry correction applied in the original article is removed. a

b


371

A number of absolute affinities have been determined by the CIDT method from the two dissociation steps of LCu+ L: L–Cu+ –L → L–Cu+ + L L–Cu+ → Cu+ + L

(6.17) (6.18)

The affinities (and deduced basicities) corresponding to dissociations 6.17 and 6.18 can be added to give two-bond absolute values. These values may be considered for anchoring the above relative scale (and enlarging the database). Among them, the values CuCA(2L) = 411.8 ± 7.8 and CuCB(2L) = 347.9 ± 8.3 kJ mol−1 for acetone are selected as anchors for their good accuracy and the position of acetone in the relative scales, relatively high in the 1982 ICR scale and relatively low in the 1998 HPMS scale. The values CuCA(2L) = 368.0 ± 4.4 and CuCB(2L) = 297.49 ± 4.5 kJ mol−1 for methanol are even more precise. However, these values have the drawback of being in the lower part of the scale and their choice as anchor would require a large number of steps to cover the whole scale, with an associated increase in uncertainty. In any case, the choice of methanol as anchor would shift the CuCA(2L) and CuCB(2L) values upwards by only 1.3 and 2.6 kJ mol−1 , respectively. The relative CuCB(2L) values anchored to acetone of refs [82, 145–147] are collected in Table 6.13. This table also contains the HPMS CuCA(2L) values of ref. [146] recalculated using their corresponding entropies and the CuCB(2L) of acetone as the starting point. Basicities are subject to the uncertainties on the anchoring value (±8.3 kJ mol−1 ) plus uncertainties on relative values (±0.8 kJ mol−1 for the values of ref. [145]), as explained in Section 6.3. It should be noted that HPMS measurements correspond to a mean temperature of 393 K, and are subject to 298 K adjustment corrections using calculated entropies, adding about 4.2 kJ mol−1 of uncertainty [146]. In a few cases, entropies estimated from compounds of close structural similarity are used to complete the CuCB(2L) from ref. [146]. The values of substituted acetophenones and pyridines have been converted into absolute CuCB(2L) or approximate CuCA(2L) as indicated in the footnotes of Table 6.13. The conversion utilizes ab initio entropy values, which are supposed to be constant within a closely related series. All other absolute CuCA(2L) and CuCB(2L) values, obtained by CIDT, are presented separately in Table 6.14. 6.5.8

Methylammonium

The association between a protonated hydrogen-bond donor and a neutral hydrogen-bond acceptor represents a special kind of hydrogen bond, the so-called ionic hydrogen bond [148]. The maximum number of data were found for the methylammonium cation (MAC), CH3 NH3 + , as the reference hydrogen-bond donor. Enthalpies (MACA) and entropies were obtained by HPMS, from van’t Hoff plots, in a temperature range depending on the bond strength. These parameters are considered valid at 298 K and the corresponding basicities, MACB, are calculated using this hypothesis. Uncertainties cited in most articles are in the range 4–8 kJ mol−1 for MACA and 8–16 J K−1 mol−1 for entropies, leading to combined uncertainties on MACB in the range 5–9 kJ mol−1 . The MACA and MACB values are given in Table 6.15.

7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27

3 4 5 6

1 2

No.

Carbon π bases Propene 2-Methylpropene Amines Ammonia n-Propylamine n-Butylamine Tri-n-Butylamine Six-membered aromatic N-heterocycles 4-Cyanopyridine 3-Cyanopyridine 4-Trifluoromethylpyridine 3-Trifluoromethylpyridine 2-Fluoropyridine 3-Fluoropyridine 2-Cyanopyridine 3-Chloropyridine 2-Chloropyridine 3-Acetylpyridine Pyridine 2-Methylpyridine 3-Methylpyridine 4-Methylpyridine 4-Ethylpyridine 2,6-Dimethylpyridine 4-Methoxypyridine 3,5-Dimethylpyridine 3,4-Dimethylpyridine 4-Aminopyridine 4-N,N-Dimethylaminopyridine

Lewis base

4-N CC5 H4 N 3-N CC5 H4 N 4-CF3 C5 H4 N 3-CF3 C5 H4 N 2-FC5 H4 N 3-FC5 H4 N 2-N CC5 H4 N 3-ClC5 H4 N 2-ClC5 H4 N 3-CH3 COC5 H4 N C5 H5 N 2-CH3 C5 H4 N 3-CH3 C5 H4 N 4-CH3 C5 H4 N 4-CH3 CH2 C5 H4 N 2,6-(CH3 )2 C5 H3 N 4-CH3 OC5 H4 N 3,5-(CH3 )2 C5 H3 N 3,4-(CH3 )2 C5 H3 N 4-H2 NC5 H4 N 4-(CH3 )2 NC5 H4 N

NH3 n-PrNH2 CH3 (CH2 )3 NH2 [CH3 (CH2 )3 ]3 N

CH3 CH CH2 (CH3 )2 C CH2

Formula

531.2

510.7

524.0

457.1 500.6

CuCA(2L)b

390.2 401.9 403.5 405.6 419.0 419.9 420.3 421.5 424.0 427.4 442.0 452.9 458.8 461.3 462.5 465.5 471.7 473.0 475.5 497.3 511.5

385.1 428.7 436.0 452.1

302.3 321.2

CuCB(2L)b

[147],e [147],e [147],e [147],e [147],e [147],e [147],e [147],e [147],e [147],e [146, 147] [147],e [147],e [147],e [146, 147]f [147],e [147],e [147],e [147],e [147],e [147],e

[146] [146] [147],d [146]

[145] [145],c

Ref.

Table 6.13 Copper monocation two-ligand affinity [CuCA(2L)] and basicity [CuCB(2L)] scales (kJ mol−1 ) at 298 K, from ligand-exchange equilibria anchored at the experimental values for acetone.

49 50 51 52

48

42 43 44 45 46 47

39 40 41

32 33 34 35 36 37 38

30 31

29

28

Butan-2-one Pentan-3-one Heptan-4-one 2,4-Dimethylpentan-3-one

1,2-Dimethylimidazole Nitriles Hydrogen cyanide Acetonitrile Water, alcohols Water Methanol Ethanol n-Propanol n-Butanol 2-Propanol tert-Butanol Ethers Oxirane Diethyl ether Tetrahydrofuran Carbonyl compounds Aldehydes Formaldehyde Acetaldehyde Propanal Butanal 2-Methylpropanal 2,2-Dimethylpropanal Aliphatic ketones Acetone

Five-membered N-heterocycles 1-Methylimidazole

CH3 COCH2 CH3 CH3 CH2 COCH2 CH3 CH3 (CH2 )2 CO(CH2 )2 CH3 (CH3 )2 CHCOCH(CH3 )2

(CH3 )2 CO

H2 CO HCOCH3 HCOCH2 CH3 HCOCH2 CH2 CH3 HCOCH(CH3 )2 HCOC(CH3 )3

c-(CH2 )2 O (CH3 CH2 )2 O c-(CH2 )4 O

H2 O MeOH CH3 CH2 OH CH3 CH2 CH2 OH CH3 (CH2 )3 OH (CH3 )2 CHOH (CH3 )3 COH

HC N MeC N

a

a

425.3 433.3

414.4 ± 7.8

416.1

334.1 366.7 385.1

470.5

582.6

347.9 ± 8.3 352.5 358.8 367.0 368.0

263.8 311.9 318.2 323.6 326.6 331.2

295.2 340.4 344.6

264.2 294.8 310.2 315.7 320.3 324.5 339.5

322.0 396.4

503.1 494.1 515.9

(Continued)

[145] [145] [146] [145]

[16],g

[145] [145] [145] [145] [145] [145]

[145] [145] [145]

[145, 146] [145, 146] [145, 146] [145] [145] [145] [145]

[145] [146]

[146] [147],d [147],d

77

75 76

69 70 71 72 73 74

53 54 55 56 57 58 59 60 61 62 63 64 65 66 67 68

No.

Ring-substituted acetophenones 4-Trifluoromethylacetophenone 3-Trifluoromethylacetophenone 3-Fluoroacetophenone 3-Chloroacetophenone 4-Chloroacetophenone 4-Fluoroacetophenone 1,4-Diacetylbenzene Acetophenone 3-Methylacetophenone 3-Methoxyacetophenone 3-Chloro-4-methoxyacetophenone 4-Methylacetophenone 3,5-Dimethylacetophenone 4-(Methylthio)acetophenone 3,4-Dimethylacetophenone 4-Methoxyacetophenone Esters Methyl formate n-Butyl formate Methyl acetate Ethyl propionate Ethyl butyrate n-Propyl butyrate Amides N-Methylacetamide N,N-Dimethylacetamide Nitroalkane Nitromethane

Lewis base


CH3 NO2

CH3 CONHCH3 CH3 CON(CH3 )2

HCOOMe HCOO(CH2 )3 CH3 CH3 COOCH3 CH3 CH2 COOCH2 CH3 CH3 CH2 CH2 COOCH2 CH3 CH3 CH2 CH2 COOCH2 CH2 CH3

4-CH3 C6 H4 COCH3 3,5-(CH3 )2 C6 H3 COCH3 4-CH3 SC6 H4 COCH3 3,4-(CH3 )2 C6 H3 COCH3 4-CH3 OC6 H4 COCH3

a

4-CF3 C6 H4 COCH3 3-CF3 C6 H4 COCH3 3-FC6 H4 COCH3 3-ClC6 H4 COCH3 4-ClC6 H4 COCH3 4-FC6 H4 COCH3 4-H3 CCOC6 H4 COCH3 C6 H5 COCH3 3-CH3 C6 H4 COCH3 3-CH3 OC6 H4 COCH3

Formula

480.5 501.5

419.0

CuCA(2L)b

294.8

411.9 431.2

314.0 332.4 341.2 355.9 361.3 364.2

336.8 339.6 348.8 351.9 355.4 356.9 358.8 368.1 376.2 377.0 383.1 383.4 387.6 388.6 389.7 398.0

CuCB(2L)b

[145]

[146] [146]

[145] [145] [145, 146] [145] [145] [145]

[82],h [82],h [82],h [82],h [82],h [82],h [82],h [82],h [82],h [82],h [82],h [82],h [82],h [82],h [82],h [82],h

Ref.

28

Me

N

N

Formulae:

N

29

Me

N

Me

63

OMe

COCH

Cl

Sulfinyl compound Dimethyl sulfoxide Isocyanate Methyl isocyanate Thiols, sulfide Methanethiol Ethanethiol Methyl sulfide Isothiocyanate Methyl isothiocyanate Haloalkanes Chloroethane 1-Chloropropane 1-Chloro-2-methylpropane 2-Chloropropane Bromoethane 1-Bromopropane 1-Bromo-2-methylpropane 2-Bromopropane 2-Bromobutane CH3 CH2 Cl CH3 (CH2 )2 Cl (CH3 )2 CHCH2 Cl (CH3 )2 CHCl CH3 CH2 Br CH3 (CH2 )2 Br (CH3 )2 CHCH2 Br (CH3 )2 CHBr CH3 CH2 CHBrCH3

CH3 N C S

CH3 SH CH3 CH2 SH (CH3 )2 S

CH3 N C O

(CH3 )2 SO

354.2

336.2

317.8

440.4

401.0

488.1

261.7 268.4 272.6 278.0 276.8 284.7 290.6 295.2 301.9

360.0

334.1 350.0 373.8

274.7

416.9

[145, 146] [145] [145] [145] [145, 146] [145] [145] [145, 146] [145]

[145]

[145, 146] [145] [146]

[145]

[146]

Values are subject to the uncertainty on the anchoring value plus other errors on relative values and on temperature correction for the HPMS data, see text. c Symmetry correction of ref. [145] removed. d Evaluated in ref. [147] from a correlation with data of ref. [146]. e Obtained by adding the substituent effect on CuCB(2L) measured at 343 K to the value for pyridine (442.0 kJ mol−1 ). This is reasonable since the entropy and thermal corrections are minimal in the narrow temperature interval 343–298 K. The values obtained in this way are actually linked to the absolute value of acetone. f Entropy estimated by similarity. g This value was made consistent with CuCA(2L) of ref. [146] listed in this table, by using the calculated entropy of this reference rather than the value from ref. [16]. h Obtained by adding the substituent effect on CuCB(2L) measured at 343 K to the value of acetophenone (368.1 kJ mol−1 ). See footnote e.

b

a

84 85 86 87 88 89 90 91 92

83

80 81 82

79

78

376


Table 6.14 Copper monocation two-ligand affinity [CuCA(2L)] and basicity [CuCB(2L)] scales (kJ mol−1 ) at 298 K (unless noted otherwise), from CIDT measurements. No.

Lewis base Carbon π bases Ethylene Benzene Pyrrole Amine Ammonia

1 2 3 4

Formula

CuCA(2L) 349.3 ± 18.4 373.5 ± 15.1 432.0

NH3

485.0 ± 18.0 490.4 ± 23.7

422.0 ± 24.1

[191] [16],d

495.9 ± 12.7 499.1 ± 11.6 608.8 ± 16.1 628.7 ± 13.2

417.4 ± 16.2 421.1 ± 13.7 520.2 ± 18.3 540.1 ± 15.7

[193] [193] [193],e [193],e

478.1 ± 9.2

408.8 ± 10.5

[194]

10 11 12

Methanol Dimethyl ether

MeOH (CH3 )2 O

13

(CH3 OCH2 )2

14

1,2-Dimethoxyethane Ketone Acetone

(CH3 )2 CO

411.1 ± 10.8 411.8 ± 7.8

15

Miscellaneous Carbon monoxide

CO

321.0 ± 7.6

9

a

Ref.

H2 C CH2 C6 H6 c-(CH)4 NH

Six-membered aromatic N-heterocycles 4,4 -Bipyridine (C5 H4 N)2 Pyridine C5 H5 N 2,2 -Bipyridine (C5 H4 N)2 a 1,10-Phenanthroline Nitrile Acetonitrile MeC N Water, alcohols, ethers Water H2 O

5 6 7 8

CuCB(2L)

327.1 ± 10.3 331.0 ± 10.1 332.0 ± 14.6 368.0 ± 4.4 378.3 ± 13.9 375.6 ± 19.4 444.2 ± 9.8

[187],b [189],b [47],c

272.7 ± 15.4 297.4 ± 4.5 313.9 ± 23.7 369.1 ± 11.8 346.4 ± 11.5 347.9 ± 8.3

[195],b [188] f

[196] [195],b f

[73],e [16],f [16] [197],b

Formula:

1 N

N

10

8 b

At 0 K. c Temperature not stated. d These values appear too large, as discussed in ref. [16]. e Bidentate behaviour. f Estimated from 393 K data in ref. [16].

6.6 6.6.1

Significance and Comparison of Gas-Phase Cation Scales Properties of Cations and Significance of MCB and MCA Scales

Properties of cations (including H+ for comparison) relevant to the binding energy, ∆Eel , and therefore contributing to MCA and MCB are given in Table 6.16 and discussed in the following.


377

Table 6.15 Methylammonium cation affinity (MACA) and basicity (MACB) scales (kJ mol−1 ) at 298 K, from HPMS measurements. No.

Lewis base

Formula

4

Carbon π bases Cyclohexene Benzene Pyrrole Carbon bases (isonitriles) Methyl isocyanide

5

Ethyl isocyanide

CH3 CH2 N C

6 7

Amines Ammonia Methylamine

NH3 CH3 NH2

1 2 3

8 9 10 11 12 13 14 15

Dimethylamine Trimethylamine Nitriles Hydrogen cyanide Cyanoacetylene Acetonitrile Butyronitrile Benzonitrile Water, alcohols Water

MACAa

MACBa

c-C6 H10 C6 H6 c-(CH)4 NH

48.5 78.7 77.8

27.4 47.4 51.6

[198] [198] [198]

CH3 N C

99.7

71.4

99.5

69.2

[148, 199] [148, 199]

89.5 98.7 106.0 90.8 115.1 136.0

57.1 66.9 72.0 61.4 84.0

[200] [148] [201] [200] [200] [200],b

87.0 100.4 109.6 102.5 117.6 123.0

58.4 69.2 77.5 70.3 83.9 84.1

[202] [202] [202] [199] [202] [202] [148] [203] [204] [205],c [206] [206] [206] [206] [206] [206]

(CH3 )2 NH (CH3 )3 N HC N HC CC N CH3 C N CH3 CH2 CH2 C N C6 H5 C N H2 O

74.5 77.0 78.7

MeOH CF3 CH2 OH CH3 CH2 OH CH3 CH2 CH2 OH (CH3 )3 COH CH3 (CH2 )3 OH

79.5 79.9 89.1 92.0 95.8 98.3

22 23

Methanol 2,2,2-Trifluoroethanol Ethanol n-Propanol tert-Butanol n-Butanol Ethers Dimethyl ether Diethyl ether

43.7 45.9 44.8 49.4 44.4 57.9 60.1 63.4 65.9

(CH3 )2 O (CH3 CH2 )2 O

90.0 92.0

53.3 60.8

24 25

Di-n-propyl ether Di-n-butyl ether

(CH3 CH2 CH2 )2 O [CH3 (CH2 )3 ]2 O

100.4 104.6

67.1 69.7

26 27 28

Di-n-hexyl ether 1,2-Dimethoxyethane 1,3-Dimethoxypropane

[CH3 (CH2 )5 ]2 O (CH3 OCH2 )2 CH3 O(CH2 )3 OCH3

113.8 125.9 130.5

74.6 88.4 90.6

16 17 18 19 20 21

Ref.

[206] [206, 207] [206] [206, 207] [206] [207] [206] (Continued)

378



29 30 31 32 33 34 35 36 37 38 39 40 41 42

43 44 45 46 47 48 49 50 51 52

Lewis base Carbonyl compounds Ketones Acetone Butan-2-one Pentan-3-one Hexan-3-one Acids, esters, amide Formic acid Ethyl trifluoroacetate Methyl formate Acetic acid Methyl acetate n-Butyl formate n-Propyl pentanoate Isopropyl acetate Formamide N-Acetylalanine methyl ester Nitroalkane Nitromethane Thiols, sulfide Hydrogen sulfide Methanethiol Ethanethiol Propanethiol Ethyl methyl sulfide Haloalkanes Chloromethane Bromomethane Fluoromethane Miscellaneous Carbon dioxide

MACAa

MACBa

(CH3 )2 CO CH3 COCH2 CH3 CH3 CH2 COCH2 CH3 CH3 CH2 COCH2 CH2 CH3

100.4 105.4 108.4 113.0

71.4 74.2 76.0 79.3

[206] [206] [206] [206]

HCOOH CF3 COOCH2 CH3 HCOOCH3 CH3 COOH CH3 COOCH3 HCOO(CH2 )3 CH3 CH3 (CH2 )3 COOCH2 CH2 CH3 CH3 COOCH(CH3 )2 HCONH2

79.5 89.5 89.5 92.0 98.3 102.5 125.5 125.5 125.5

49.4 57.1 59.6 61.7 67.4 70.1 82.1 81.6 88.1

CH3 CONHCH(CH3 )COOCH3

167.8

124.0

[206] [206] [206] [206] [206] [206] [206] [206] [206, 208] [208]

CH3 NO2

85.8

57.1

[206]

H2 S CH3 SH CH3 CH2 SH CH3 CH2 CH2 SH CH3 SCH2 CH3

45.2 56.1 61.1 73.2 82.8

20.2 28.5 36.7 42.1 51.6

[209] [209] [209] [161] [161]

CH3 Cl CH3 Br CH3 F

44.8 46.9 49.4

19.1 20.7 20.3

[206] [206] [206]

CO2

55.2

28.5

[210]

Formula

Ref.

a

See text for uncertainty. Empirical estimate. c Measured by using an atmospheric ionization source. Stated error ±0.8 kJ mol−1 . b

The electronic structure of the cation determines the orbital able to accept electron density from the ligand. These are empty or half-filled AOs, the LUMO of CpNi+ with strong Ni 4s character and a natural bond orbital σ ∗ (N H) of CH3 NH3 + . However repulsion between ligand electrons and electrons on M+ , such as the 1s2 core of Li+ , the 3s2 electrons on Al+ or the electrons in the half-filled 4s and 3d orbitals of Mn+ , should severely limit this charge transfer, and the L M+ bond is expected to be largely ionic in character. Nevertheless some covalent bonding can also occur by delocalization of 3d electrons of transition metals into suitable unoccupied ligand orbitals. Unsaturated ligands such as benzene, CO, nitriles or isonitriles would be able to accept metal electrons by using empty π ∗ MOs. This so-called π back-bonding is considered an important feature of soft Lewis acids [211]. Stronger

(∞) 930.0 853.6

0.441

0.977 0.76 0.60 35.12 183.0

3 1s2 1 S 2s

Li+

0.90 0.95 21.08 127.5 107.1

11 2p6 1 S 3s

Na+

0.98 1.33 13.64 90.6 75

19 3p6 1 S 4s

K+

0.72 Hard 191.5

13 3s2 1 S 3p

Al+

b

Natural charge on M given by a natural population analysis at B3LYP/6–311+G(2d,2p)//B3LYP/6–311G. Charge on M taken from MP2 calculations. c ˚ A. d eV. e MCA in kJ mol−1 . f 14 outer electrons. g MO with strong Ni 4s character. h MO with strong Ni 3d character.

a

1

Atomic number Electronic configuration Term Electron-accepting orbital Electron-donating orbital Q (M+ – pyridine)a Q (M+ – N7-adenine)b Ionic radiusc Hardnessd Acidity (towards Pyr)e Acidity (towards NH3 )e

1s

H+

Property

Table 6.16 Properties of cations.

0.75 6.28 182.6 147

25 3d5 4s1 7 S 4s, 4p, 3d 3d

Mn+ 29 3d10 1 S 4s, 4p 3d 0.862 0.65 0.96 Soft 264.0 241.6

Cu+

Hard 89.5

212.3

σ ∗ (N H)

CH3 NH3 +

soft

4sg 3dh

f

CpNi+

380


back-bonding is expected for Cu+ than for Mn+ since 3d orbitals on Mn+ have only one electron each whereas for Cu+ they are full. The covalency can be estimated from the charge retained on the cation in the complex, Q(M+ ). The combination of results on the complexes of pyridine [147] and adenine [62] indicates that the covalent character of the N M+ bond decreases in the order H+ > Cu+ > Li+ > Na+ > K+ ≈ 0 and goes from a significantly covalent N H+ bond to a largely ionic N K+ bond. The trend in electrostatic attraction follows the trend in cation size, estimated in Table 6.16 by the Pauling univalent radius [212]. Indeed, the smaller the cation radius, the shorter is the bond length and the stronger the electrostatic energy, as shown by the expression of the ion/permanent dipole energy: U (ion/permanent dipole) = −eµ/4π ε0 R 2

(6.19)

where e is the proton charge, µ the ligand dipole, ε 0 the permittivity of vacuum and R can be approximated by the ion–ligand distance. A small cation radius also implies a strong polarizing power, and hence a large dipole induced by the field of the unit charge into the polarizable ligand and a large polarization energy, as shown by the expression U (ion/induced dipole) = −e∆µ/4π ε0 R 2 = −e2 α/2 (4π ε0 )2 R 4

(6.20)

where ∆µ is the induced dipole and α the ligand polarizability. Calculations on the example ˚ yields of the H2 O–Li+ complex [213], where µ = 1.854 D, ∆µ = 1.01 D and R = 1.85 A, U = −157 and U = −85 kJ mol−1 . Compared with the total binding energy of −150 kJ mol−1 , this crude decomposition analysis reveals that the water/Li+ attraction is dominated by electrostatic (ES) and polarizability (POL) energy components. The importance of POL requires a new substituent constant, σ α [214], to be added to structure–cation basicity relationships. In combination with the field/inductive and resonance constants, σ F and σ R [215], respectively, the polarizability constant σ α enables the effects of molecular structure on cation basicities [67] to be described, empirically but successfully. The polarizability of the cation is another property contributing to POL and DISP (dispersion). For alkali metal cations, the polarizability becomes significant only from K+ [in a0 3 : α(Li+ ) = 0.061, α(Na+ ) = 0.346 and α(K+ ) = 4.063]. The soft, borderline or hard character of the MCA and MCB scales is given by the absolute hardness of the cation, η. When η is unknown, the cation is classified soft (or hard) when the basicity increases (or decreases) in descending a group (15, 16 and/or 17) of the periodic table. Thus CH3 NH3 + is hard because MACB (H2 O) > MACB (H2 S). Cu+ is a classical example of a soft cation because CuCB(2MeSH) > CuCB(2MeOH), Cu+ can back-donate 3d electrons, and some covalency of the pyridine Cu+ bond is manifested by the partial neutralization of the unit charge on copper. Finally, the acid strength of studied cations (and H+ for comparison) is given in Table 6.16 towards pyridine, a borderline base and ammonia, a hard base. A good agreement is found between the MCA(C5 H5 N) order: H+ Cu+ (1L) > Al+ ≈ Li+ ≈ Mn+ > Na+ > K+


381

and the MCA(NH3 ) order: + H+ Cu+ (1L) ≈ CpNi+ > Mn+ > Na+ > CH3 NH+ 3 ≈K

6.6.2

Relationship of MCA with MCB

It was shown in Chapters 1 and 4 that there was no general relationship between affinity and basicity when a series of Lewis bases react with a neutral Lewis acid such as BMe3 or a hydrogen-bond donor. Only relationships limited to series of closely related bases (so-called family-dependent relationships) could be established (see Figure 4.10). The situation is different in ion/molecule reactions. When MCA is plotted against MCB, for a given cation, rather general relationships are observed. They are the consequence of the extrathermodynamic enthalpy–entropy relationship [216] 6.21: ∆H ◦ = β∆S ◦ + ∆G ◦ (at T = β) = β∆S ◦ + ∆G ◦β

(6.21)

often called the compensation relationship, since the contributions of ∆H ◦ and ∆S◦ to the resulting ∆G◦ are opposite. The introduction of relationship 6.21 into the thermodynamic relationship 6.22: ∆G ◦T = ∆H ◦ − T ∆S ◦

(6.22)

∆H ◦ = γ ∆G ◦T + (1 − γ )∆G ◦β

(6.23)

yields [217] Equation (6.23):

where γ = 1/(1 − T /β) or, in terms of MCA and MCB, MCA = γ MCB + constant

(6.24)

The statistics of relationship 6.24 are summarized in Table 6.17. The plot of NaCA versus NaCB is illustrated in Figure 6.12. The value of the slope γ is close to 1 for Mn+ , Li+ , Na+ and K+ . The value γ = 1 corresponds to β = ∞ , that is, to isoentropic reactions. Indeed the assumption that entropy changes are small has been postulated to explain the Gibbs energy of complexation of Al+ [136], Mn+ [140], Ni+ [14], Co+ [11], Cu+ [145], CpNi+ [141] and BrFe+ [10] in terms of electronic binding energy. It has also been observed [164] that the entropy changes for the addition of Na+ to Lewis bases are nearly all within the range −92 ± 8 J K−1 mol−1 and that the entropies of complexation of Cs+ are nearly constant for monodentate ligands [9]. In the reaction M+ + L LM+ , three translational degrees of freedom are lost, three new vibrational modes are created (the L M+ stretching and two bending modes), and the moments of inertia of LM+ increase compared with L. Hence the translation entropy is negative whereas the vibration and rotation entropies are positive. By far the greatest part of the total negative entropy change is accounted for by the translation entropy change. For example, the total, translational, rotational and vibrational entropies for the addition of Al+ to dimethyl ether, calculated [78] at the MP2/6–31G∗ level, are −96.8, −144.1, +14.1 and +33.1 J K−1 mol−1 respectively. This may explain why the total entropy change is not very sensitive to the chemical structure of the ligands. However, two contributions to ∆S, symmetry number changes and chelate formation, could provoke a departure from isoentropic behaviour.

382


Table 6.17 Relationships between experimental gas-phase cation affinity and basicity at 298 K. Cation

Data in Table:

Cu+ (2L)

6.13 (CIDT) 6.12 (ICR) 6.9 6.2 6.4 6.6 6.14

Mn+ Li+ Na+ K+ CH3 NH3 + a

γ

Intercepta

n

r

sa

1.11 ± 0.02 1.04 ± 0.02 1.04 ± 0.12 1.00 ± 0.01 1.04 ± 0.01 1.05 ± 0.02 1.16 ± 0.02

31 ± 7 54 ± 6 22 ± 16 30 ± 2 25 ± 1 25 ± 1 22 ± 1

12 24 5 47 58 56 55

0.999 0.998 0.981 0.997 0.996 0.990 0.992

4.7 5.0 3.4 2.7 2.9 3.3 3.2

kJ mol−1 .

The MACA versus MACB relationship shows a slope significantly greater than unity. This means that the formation of the ionic hydrogen bond cannot be considered as isoentropic. However, the compensation effect allows the existence of a relationship between MACA and MACB. The same compensation effect is also observed for the calculated [139] enthalpies and entropies of complexation of Al+ (Figure 6.13). It must be noted that these compensations are very moderate. For example, whereas AlCA increases by ∼86 kJ mol−1 on going from the weakest to the strongest complex, the corresponding T∆S◦ increase is only ∼5 kJ mol−1 at 298 K. 6.6.3

The Computation of MCB and MCA Scales

The performance of theoretical methods to compute gas-phase cation basicity and/or affinity scales can be evaluated using Equation 6.25: MCA (MCB) (calc.) = a MCA (MCB) (exp .) + b

250

NaCA / kJ mol

-1

200 150 100 50 0 -10

40

90

140

190

-1

NaCB / kJ mol

Figure 6.12

Plot of sodium cation affinity versus sodium cation basicity.

(6.25)


8

-1

200

o

-∆H / kJ mol

383

7

150

6

4

5 2 1

100 80

3

90

100 o

-1

110

-1

-∆S / J K mol

Figure 6.13 Calculated [78, 139] enthalpy ∆H◦ versus entropy ∆S◦ for the Al+ binding to a series of Lewis bases: 1 H2 CO [78], 2 H2 CO [139], 3 H2 O, 4 MeCN, 5 NH3 , 6 Me2 O, 7 Me3 N and 8 C5 H5 N.

where calc. and exp. stand for experimental and calculated, respectively. A good agreement between the calculated and experimental properties is reflected by a slope (a) close to unity, an intercept (b) close to zero, a squared correlation coefficient (r2 ) approaching unity and a standard deviation from the regression line (s) comparable to the experimental error. Random errors are indicated by r2 and s, while a slope different from unity and a non-zero intercept imply systematic errors (of calculation if the experiment is assumed to be without systematic error). Five comprehensive sets of data computed by the same authors at the same level of theory are available in the literature. They relate to the lithium, sodium and potassium cations. The results of comparisons are summarized in Table 6.18. The agreement between theory and experiment is good for Na+ and K+ with both modest and advanced computational effort. In contrast, B3LYP and even G2 calculations do not perform well for Li+ . A critical evaluation [132] of the multiple levels of theory available for determining LiCA reveals that core correlation on the lithium cation is needed to describe the binding energy accurately. An adequate level of theory is provided by correlation consistent polarized core/valence basis sets. They permit 1s electrons of lithium to polarize away from the ligand and to correlate with the ligand electrons. There are no such systematic theoretical studies concerning a comprehensive series of ligands for the other cations selected in this chapter. The interested reader is referred to the publications listed for these cations in Section 6.5. These papers often combine experimental and theoretical determinations. 6.6.4

MCA and MCB Scales and the Concept of a Cation/π Interaction

The LiCA scale (Table 6.3) shows that benzene, a ligand without a permanent dipole, binds Li+ more tightly than water, a dipolar ligand. This indicates that the interaction between cations and the face of a π system is a potentially strong binding force. This binding can be explained by the contribution of attractive forces (mainly ion/quadrupole, ion/induced dipole and dispersion and, less importantly, charge transfer) well known for a long time. However, it has been given a new, specific name: the cation/π interaction. The concept of a

a 1.34 ± 0.03 1.17 ± 0.04 1.04 ± 0.04 1.07 ± 0.01 1.005 ± 0.029 0.997 ± 0.30

Computational level

B3LYP/6–311+G∗∗ G2 MP2(full)/aug-cc-pVTZ(Li-C)// MP2(full)/cc-pVDZ(Li-C)b MP2(full)/6–311+G(2d,2p)// MP2(full)/6–31G∗ CBS-Q//MP2(full)/6–31G∗ B3LYP/6–31+G(3df,2p)// B3LYP/6–31G(d)

b

a

kJ mol−1 . Li-C stands for lithium core correlation. c Phenylalanine, cytosine and guanine are excluded.

LiCA (0 K) NaCB (298 K) NaCA (0 K) KCA (298 K)

LiCB (373 K)

Property

−24.6 ± 4.7 −15.4 ± 4.9 −4.9 ± 6.9 −5.6 ± 1.2 1±3 −1±3

ba

62 37 25 38 17 47c

n

0.962 0.968 0.970 0.994 0.988 0.960

r2

7.5 5.0 10.6 2.2 4.5 5.5

sa

[133] [133] [132] [106] [162] [180]

Ref.

Table 6.18 Linear regression analysis of the relationship between the calculated and experimental gas-phase cation basicities and affinities.


385

Pyridine affinity / kJ mol-1

300 250

Cu Al

200 150

Na

100

Mn

Li

K

50 0 0

50

100

150

200

250

-1

Benzene affinity / kJ mol

Figure 6.14 Plot of pyridine affinity versus benzene affinity for neutral Lewis acids (•) (BF3 , IBr, I2 and 4-FC6 H4 OH) and cations (+). The line with slope of unity is drawn for comparison.

cation/π interaction has become very popular since it was recognized that this interaction makes significant contributions to crystal packing, to the secondary structure of proteins and to molecular recognition in both biological and synthetic receptors and in both gas and condensed phases [218, 219]. MCA and MCB scales provide a simple demonstration of the strength and specificity of the cation/π interaction. Compare the affinity of benzene, a π ligand, to that of pyridine, an n ligand of similar structure, for the class of cationic Lewis acids and that of non-cationic ones, such as BF3 , IBr, I2 and 4-FC6 H4 OH. In the comparison plot in Figure 6.14, it appears that different comparison lines are required for each class. Their slopes show that the interaction of cations with a π ligand becomes fairly large compared with that of neutral Lewis acids. The strong cation basicity of π ligands is also shown by the preference of cations for the π site of heteroaromatics and heterosubstituted benzenes (see Section 6.2). In contrast, neutral Lewis acids are not so selective and form both π and n complexes (see Chapters 4 and 5). Lastly, it is noteworthy that the cation/π interaction yields an Li+ /1,3-diphenylpropane complex (see Figure 6.2) which is 10 kJ mol−1 more stable than the Li+ complex of dimethyl sulfoxide, a strong ligand with a high dipole moment (4 D). MCA and MCB scales (Tables 6.3–6.15) are the source of numerous structure–basicity relationships in the field of cation/π interactions. The following sequences of basicity (affinity) are observed in the comparison of benzene with: r condensed aromatics: azulene > anthracene ≈ phenanthrene > naphthalene > benzene r alkenes and alkynes: benzene > propyne > propene

386


r heteroaromatics: indole > 1-methylpyrrole > pyrrole > benzene > furan r heterosubstituted benzenes: fluorobenzene < chlorobenzene < bromobenzene < iodobenzene < thiophenol ≈ benzene < phenol < anisole < toluene < ethylbenzene < aniline The last sequence can be made quantitative by means of the field, resonance and polarizability constants of the ring substituents, σ F , σR◦ and σ α [215], respectively: LiCB = 113.0(±1.7) −85.1(±5.6) σF − 72.9(±8.6) σR◦ − 18.6(±3.3) σα

(6.26)

n = 8 (Cl, Br, SH, H, OH, CH = CH2 , Me, Et), r = 0.995, s = 1.8 kJ mol−1 . Regression coefficients of expression 6.26 show that the lithium cation/π interaction is affected almost equally by the field and resonance effect of the ring substituent and that the polarizability effect is smaller, although significant. However the fundamental features of the MCA and MCB scales of π ligands are best revealed by quantum mechanical calculations. Notable studies concern the π complexes of Li+ [220], Na+ [220, 221], K+ [221], Al+ [35], Cu+ [15] and MeNH3 + [222]. 6.6.5

Conventional Versus Ionic Hydrogen-Bond Basicity and Affinity Scales

The MACB and MACA scales occur in the chapter on metal cation scales because the techniques of determination (Section 6.3) are the same for both inorganic and organic cations. Moreover, numerous similarities have been pointed out in the interactions of ammonium and potassium cations [223]. Nonetheless, the percentages of variance of MACB and MACA explained by KCB and KCA are only 57% (n = 11) and 29% (n = 13), respectively. It would probably be best to retain MACB and MACA in the category of hydrogen-bond scales, largely because the geometry of the ionic hydrogen bond MeNH3 + · · · L is similar to that of the conventional O H · · · L hydrogen bond. Indeed, both are two-centre and linear and have similar directionalities (see Section 4.1 in Chapter 4). It is true that the percentages of variance of MACB and MACA explained by 4-fluorophenol basicity and affinity scales still remain low (54 and 43%, respectively), but they gain significance because they refer to a larger number of ligands (n = 31 and 18). The differences between ionic and conventional hydrogen bonds originate from the charge on the NH groups, which amplifies the electrostatic attraction and the polarization effects. Consequently, ligand properties such as dipole moment and polarizability gain importance in the binding energy of ionic hydrogen bonds. Indeed, if the dipole moment µ and a crude parameter of polarizability are added as explanatory variables, MACB is fairly well related to the conventional HB basicity through the three-parameter Equation 6.27: MACB = 22.6(±3.2) + 2.78(±0.24) ( 4-FC6 H4 OH basicity) + 5.98(±0.90) µ + 3.41(±0.52)N n = 29, r = 0.951, s = 5.0 kJ mol−1 where N is the number of polarizable heavy atoms (C, O and N) of the ligand.

(6.27)


387

Table 6.19 Statistics of the linear regressions MCB = aMCB + b. MCB

MCB

Sample

n

r2

sa

NaCB KCB AlCB

LiCB LiCB LiCB

AlCB

NaCB

MnCB

AlCB

CpNiCB

MnCB

All bases All bases All bases Alcohols Nitriles Carbonyls All bases Oxygen bases All bases Oxygen bases All bases Oxygen bases

27 30 45 8 7 21 16 14 22 18 20 14

0.965 0.831b 0.876 0.973 0.934 0.982 0.673 0.938 0.761 0.950 0.425 0.947

3.9 9.2 9.6 1.9 7.7 4.5 12.0 5.4 5.8 2.5 8.9 2.0

a

kJ mol−1 . r rises to 0.932 when three outliers (water, 1,2-dimethoxyethane and 1,2,3-triazole) are excluded.

b 2

6.6.6

Comparison of Cation Basicity Scales

It seems obvious that the nine MCA and nine MCB scales in this chapter are interdependent to some extent. A systematic exploration of the similarities and differences between two given scales requires 72 correlations (36 for basicity and 36 for affinity) to be performed. If two given scales are found to be not, or poorly, correlated, further statistical work is necessary to understand the chemistry underlying the differences, for example by splitting up the sample of bases into families. Here, only MCB scales will be compared since MCA is related to MCB. Moreover, the correlation analysis will be limited to the comparison of alkali metals, the comparison of aluminium with alkali metals and with transition metals, and the comparison of transition metals. The results are summarized in Table 6.19. The comparison of alkali metals shows that NaCB is well correlated with LiCB. Such a good correlation indicates that the blend of interactions contributing to LiCB and NaCB is similar for the two alkali metals within the wide range of studied ligands. As expected, the similarity decreases on going from Na+ to K+ . The comparison of aluminium with alkali metals is characterized by a family-dependent behaviour. For the whole sample of bases, the percentage of variance of AlCB explained by LiCB is not satisfactory (87.6%) and even falls with NaCB (67.3%). Good correlations are obtained only when nitriles and oxygen bases are treated separately. Moreover, comparison reveals three separate lines for oxygen bases, corresponding to ethers, alcohols and carbonyl compounds. Figure 6.15 illustrates the family-dependent similarity of Al+ with the lithium cation. MnCB is plotted versus AlCB in Figure 6.16. The data points for the oxygen bases are well correlated (r2 = 0.950). Those for the nitriles and pyridines are clearly off the correlation line. Their Mn+ basicities are about 13 kJ mol−1 stronger than expected from the line of oxygen bases. This can be attributed [140] to π back-bonding of Mn 3d electrons into empty π ∗ orbitals of nitriles and pyridines.

388

Lewis Basicity and Affinity Scales 220 200

AlCB / kJ mol

-1

180 160 140 120 100 80 100

120

140

160

LiCB / kJ mol

180

-1

Figure 6.15 Comparison of AlCB and LiCB. The following families can be distinguished from left to right: ethers ( ), alcohols ( ), carbonyls (except acetylacetone) ( ) and nitriles ( ).

The comparison of CpNi+ basicity with Mn+ basicity, Figure 6.17, also shows a good linear correlation for oxygen ligands. Several bases fall above the line of oxygen ligands: MeSH, Me2 S, HCN, MeCN, NH3 and NMe3 . The deviation of nitriles towards stronger CpNi+ basicity could be explained in terms of π back-bonding. Indeed, stronger backbonding is expected for CpNi+ compared with Mn+ since 3d orbitals on Mn+ have only one electron each whereas for CpNi+ they are nearly full. A greater contribution of the POL

160

MnCB / kJ mol

-1

140

120

100

80 80

100

120

140

160

180

AlCB / kJ mol-1

Figure 6.16 Comparison of AlCB and MnCB. The line is a least-squares fit to the data for oxygen bases (alcohols and ethers , carbonyls ) omitting the data points for nitriles ( ) and pyridines ( ).


389

CpNiCB / kJ mol-1

190

170

150

130 90

100

110

120

MnCB / kJ mol

130

140

150

-1

Figure 6.17 Comparison of CpNiCB and MnCB. The line is a least-squares fit to the data for oxygen bases (alcohols and ethers , carbonyls ) omitting the data points for nitriles ( ), amines ( ) and single-bonded sulfur bases ( ).

and charge-transfer terms to the binding energy could account for the deviation of sulfur bases and amines, respectively, since sulfur bases are more polarizable and amines have lower ionization energies than the corresponding oxygen bases. This tentative analysis of deviations needs to be confirmed with the help of ab initio calculations and binding energy decomposition schemes.

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160. Hoyau, S. and Ohanessian, G. (1998) Interaction of alkali metal cations (Li+ –Cs+ ) with glycine in the gas phase: a theoretical study. Chem. Eur. J., 4, 1561–1569. 161. Keesee, R.G. and Castleman, A.W. (1986) Thermochemical data on gas-phase ion/molecule association and clustering reactions. J. Phys. Chem. Ref. Data, 15, 1011–1071. 162. Armentrout, P.B. and Rodgers, M.T. (2000) An absolute sodium cation affinity scale: threshold collision-induced dissociation experiments and ab initio theory. J. Phys. Chem. A, 104, 2238–2247. 163. Hallowita, N., Udonkang, E., Ruan, C. et al. (2009) Inductive effects on cation/π interactions: structures and bond dissociation energies of alkali metal cation/halobenzene complexes. Int. J. Mass Spectrom., 283, 35–47. 164. Hoyau, S., Norrman, K., McMahon, T.B. and Ohanessian, G. (1999) A quantitative basis for a scale of Na+ affinities of organic and small biological molecules in the gas phase. J. Am. Chem. Soc., 121, 8864–8875. 165. Moision, R.M. and Armentrout, P.B. (2002) Experimental and theoretical dissection of sodium cation/glycine interactions. J. Phys. Chem. A, 106, 10350–10362. 166. Rodgers, M.T. and Armentrout, P.B. (1999) Absolute binding energies of sodium ions to short chain alcohols, Cn H2n+2 O, n = 1–4, determined by threshold collision-induced dissociation experiments and ab initio theory. J. Phys. Chem. A, 103, 4955–4963. 167. More, M.B., Ray, D. and Armentrout, P.B. (1999) Intrinsic affinities of alkali cations for 15crown-5 and 18-crown-6: bond dissociation energies of gas-phase M+ –crown ether complexes. J. Am. Chem. Soc., 121, 417–423. 168. Anderson, J.D., Paulsen, E.S. and Dearden, D.V. (2003) Alkali metal binding energies of dibenzo-18-crown-6: experimental and computational results. Int. J. Mass Spectrom. Ion Processes, 227, 63–76. 169. Kish, M.M., Ohanessian, G. and Wesdemiotis, C. (2003) The Na+ affinities of α-amino acids: side-chain substituent effects. Int. J. Mass Spectrom., 227, 509–524. 170. Sunner, J., Nishizawa, K. and Kebarle, P. (1981) Ion/solvent molecule interactions in the gas phase. The potassium ion and benzene. J. Phys. Chem., 85, 1814–1820. 171. Davidson, W.R. and Kebarle, P. (1976) Binding energies and stabilities of potassium ion complexes from studies of the gas phase ion equilibriums K+ + M = K+ M. J. Am. Chem. Soc., 98, 6133–6138. 172. Iceman, C. and Armentrout, P.B. (2003) Collision-induced dissociation and theoretical studies of K+ complexes with ammonia: a test of theory for potassium ions. Int. J. Mass Spectrom., 222, 329–349. 173. Castleman, A.W. (1978) The properties of clusters in the gas phase: ammonia about Bi+ , Rb+ , and K+ . Chem. Phys. Lett., 53, 560–564. 174. Davidson, W.R. and Kebarle, P. (1976) Ionic solvation by aprotic solvents. Gas phase solvation of the alkali ions by acetonitrile. J. Am. Chem. Soc., 98, 6125–6133. 175. Dzidic, I. and Kebarle, P. (1970) Hydration of the alkali ions in the gas phase. Enthalpies and entropies of reactions M+ (H2 O)n–1 + H2 O = M+ (H2 O)n . J. Phys. Chem. A, 74, 1466– 1474. 176. Sunner, J. and Kebarle, P. (1981) Unimolecular dissociation of ions. Effect on massspectrometric measurements of ion/molecule association equilibriums. J. Phys. Chem., 85, 327–335. 177. Evans, D.H., Keesee, R.G. and Castleman, A.W. (1991) Thermodynamics of gas-phase mixedsolvent cluster ions: water and methanol on K+ and Cl− and comparison to liquid solutions. J. Phys. Chem. A, 95, 3558–3564. 178. More, M.B., Ray, D. and Armentrout, P.B. (1997) Cation–ether complexes in the gas phase: bond dissociation energies of K+ (dimethyl ether)x , x = 1–4; K+ (1,2-dimethoxyethane)x , x = 1 and 2; and K+ (12-crown-4). J. Phys. Chem. A, 101, 4254–4262. 179. Sunner, J. and Kebarle, P. (1984) Ion–solvent molecule interactions in the gas phase. The potassium ion and dimethyl sulfoxide, DMA, DMF, and acetone. J. Am. Chem. Soc., 106, 6135–6139. 180. Lau, J.K.-C., Wong, C.H.S., Ng, P.S. et al. (2003) Absolute potassium cation affinities (PCAs) in the gas phase. Chem. Eur. J., 9, 3383–3396.

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181. Walter, D., Sievers, M.R. and Armentrout, P.B. (1998) Alkali ion carbonyls: sequential bond energies of Li+ (CO)x (x = 1–3), Na+ (CO)x (x = 1, 2) and K+ (CO). Int. J. Mass Spectrom. Ion Processes, 175, 93–106. 182. Amunugama, R. and Rodgers, M.T. (2001) Periodic trends in the binding of metal ions to pyrimidine studied by threshold collision-induced dissociation and density functional theory. J. Phys. Chem. A, 105, 9883–9892. 183. Dalleska, N.F., Tjelta, B.L. and Armentrout, P.B. (1994) Sequential bond energies of water to Na+ (3s0 ), Mg+ (3s1 ), and Al+ (3s2 ). J. Phys. Chem. A, 98, 4191–4195. 184. Dunbar, R.C., Klippenstein, S.J., Hrusak, J. et al. (1996) Binding energy of Al(C6 H6 )+ from analysis of radiative association kinetics. J. Am. Chem. Soc., 118, 5277–5283. 185. Ryzhov, V. and Dunbar, R.C. (1999) Interactions of phenol and indole with metal ions in the gas phase: models for Tyr and Trp side-chain binding. J. Am. Chem. Soc., 121, 2259–2268. 186. Rue, C., Armentrout, P.B., Kretzschmar, I. et al. (2001) Guided ion beam studies of the statespecific reactions of Cr+ and Mn+ with CS2 and COS. Int. J. Mass Spectrom., 210/211, 283–301. 187. Sievers, M.R., Jarvis, L.M. and Armentrout, P.B. (1998) Transition-metal ethene bonds: thermochemistry of M+ (C2 H4 )n (M = Ti–Cu, n = 1 and 2) complexes. J. Am. Chem. Soc., 120, 1891–1899. 188. Dalleska, N.F., Honma, K., Sunderlin, L.S. and Armentrout, P.B. (1994) Solvation of transition metal ions by water. Sequential binding energies of M+ (H2 O)x (x = 1–4) for M = Ti to Cu determined by collision-induced dissociation. J. Am. Chem. Soc., 116, 3519–3528. 189. Meyer, F., Khan, F.A. and Armentrout, P.B. (1995) Thermochemistry of transition metal benzene complexes: binding energies of M(C6 H6 )x + (x = 1, 2) for M = Ti to Cu. J. Am. Chem. Soc., 117, 9740–9748. 190. Lin, C.-Y., Chen, Q., Chen, H. and Freiser, B.S. (1997) Observing unimolecular dissociation of metastable ions in FTICR: a novel application of the continuous ejection technique. J. Phys. Chem. A, 101, 6023–6029. 191. Walter, D. and Armentrout, P.B. (1998) Sequential bond dissociation energies of M+ (NH3 )x (x = 1–4) for M = Ti–Cu. J. Am. Chem. Soc., 120, 3176–3187. 192. Lin, C.-Y., Chen, Q., Chen, H. and Freiser, B.S. (1997) Bond dissociation energy determinations for MOC(CH3 ) 2 + and MOC(CD3 ) 2 + , (M = Cr, Mn), using the continuous ejection and radiative association methods. Int. J. Mass Spectrom. Ion Processes, 167/168, 713–722. 193. Rannulu, N.S. and Rodgers, M.T. (2007) Noncovalent interactions of Cu+ with N-donor ligands (pyridine, 4,4 -dipyridyl, 2,2 -dipyridyl, and 1,10-phenanthroline): collision-induced dissociation and theoretical studies. J. Phys. Chem. A, 111, 3465–3479. 194. Vitale, G., Valina, A.B., Huang, H. et al. (2001) Solvation of copper ions by acetonitrile. Structures and sequential binding energies of Cu+ (CH3 CN)x , x = 1–5, from collision-induced dissociation and theoretical studies. J. Phys. Chem. A, 105, 11351–11364. 195. Koizumi, H., Zhang, X.-G. and Armentrout, P.B. (2001) Collision-induced dissociation and theoretical studies of Cu+ –dimethyl ether complexes. J. Phys. Chem. A, 105, 2444–2452. 196. Yang, Z., Rannulu, N.S., Chu, Y. and Rodgers, M.T. (2008) Bond dissociation energies and equilibrium structures of Cu+ (MeOH)x , x = 1–6, in the gas phase: competition between solvation of the metal ion and hydrogen-bonding interactions. J. Phys. Chem. A, 112, 388–401. 197. Meyer, F., Chen, Y.-M. and Armentrout, P.B. (1995) Sequential bond energies of Cu(CO)x + and Ag(CO)x + (x = 1–4). J. Am. Chem. Soc., 117, 4071–4081. 198. Meot-Ner, M. and Deakyne, C.A. (1985) Unconventional ionic hydrogen bonds. 2. NH+ · · · π . Complexes of onium ions with olefins and benzene derivatives. J. Am. Chem. Soc., 107, 474–479. 199. Meot-Ner, M., Sieck, L.W., Koretke, K.K. and Deakyne, C.A. (1997) The carbon lone pair as electron donor. Ionic hydrogen bonds in isocyanides. J. Am. Chem. Soc., 119, 10430–10438. 200. Yamdagni, R. and Kebarle, P. (1973) Gas-phase basicities of amines. Hydrogen bonding in proton-bound amine dimers and proton-induced cyclization of α,ω-diamines. J. Am. Chem. Soc., 95, 3504–3510. 201. Meot-Ner, M. (1992) Intermolecular forces in organic clusters. J. Am. Chem. Soc., 114, 3312–3322. 202. Speller, C.V. and Meot-Ner, M. (1985) The ionic hydrogen bond and ion solvation. 3. Bonds involving cyanides. Correlations with proton affinities. J. Phys. Chem., 89, 5217–5222.


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7 The Measurement of Lewis Basicity and Affinity in the Laboratory This chapter is intended to provide detailed examples of the spectroscopic and thermodynamic determination of most Lewis basicity scales presented in the previous chapters, namely the BF3 affinity scale, 4-fluorophenol basicity and affinity scales, the methanol infrared (IR) shift scale, the 4-nitrophenol solvatochromic shift scale, diiodine basicity and affinity scales, the iodine cyanide IR shift scale, the diiodine blue shift scale and the lithium cation basicity scale. With these examples, it is hoped that professional chemists, and also students of physicochemical sciences, will be able to supplement the scales for the molecules in which they are interested. A basic knowledge of calorimetric and Fourier transform infrared (FTIR) and ultraviolet (UV) spectrometric techniques, and of laboratory work, is required before starting with the experiments in solution. Gas-phase measurements of BF3 affinity and lithium cation basicity must be carried out by chemists experienced in gas/liquid calorimetric techniques and mass spectrometric techniques, respectively. All users are asked to follow the safety instructions for laboratory work carefully and, especially, for the chemical substances they deal with. For safety reasons, and also because water may react with or catalyse reactions involving Lewis acids and bases, and act as a competitive Lewis acid, it is strongly recommended that chemicals be handled in a thoroughly dried glove-box.

7.1

7.1.1

Calorimetric Determination of the BF3 Affinity of Pyridine by Gas/Liquid Reaction Introduction: Principles and Difficulties in the Calorimetric Measurement of the Enthalpy of a Gas/Liquid Reaction

The measurement is based on recording the heat evolved when a precisely known quantity of gaseous BF3 is reacted with a dilute solution of a Lewis base (LB) (pyridine in our case). Lewis Basicity and Affinity Scales: Data and Measurement C 2010 John Wiley & Sons, Ltd


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When a complete reaction occurs for the formation of an adduct between BF3 and the LB, the molar enthalpy of reaction is simply calculated from the evolved heat, Q, divided by the number of moles of BF3 , n. This simple relationship follows from the fact that, for a 1 : 1 adduct not appreciably dissociated under the experimental conditions, the number of moles of gas is equal to the number of moles of adduct formed, as shown by Equation 7.1 : BF3 (g) + LB(soln) LB−BF3 (soln)

(7.1)

The case of an incomplete reaction, that is when the complex is significantly dissociated, is outside the scope of this experiment, since the knowledge of the quantity of reacted BF3 requires the knowledge of the complexation constant. This case is treated in ref. [1] for the BF3 complexes of weakly basic carbonyl compounds. In the case of standardized BF3 affinity measurements (Chapter 3), the gaseous Lewis acid is introduced by small increments (5–10 additions, as in a stepwise titration) in a dichloromethane solution of the Lewis base, until saturation. This saturation corresponds to the formation of the 1 : 1 complex for a monofunctional base. If the adduct is not dissociated, the values of the ratio Q/n for the consecutive 5–10 injections are practically constant, and their mean value can be taken as the molar enthalpy of complexation. There are two devices to be set up and coupled for these experiments: a gas pressure measurement device and a calorimeter. The gas is manipulated in, and transferred from, a vacuum line. The amount of gas introduced into the calorimeter is obtained by pressure, volume and temperature (PVT) measurements before and after each injection. PVT measurements, the coupling of the vacuum line to the calorimeter cell, and the calorimeter are described below [2]. The volume V of gas is measured in a glass bulb of about 100–200 ml, calibrated by weighing water at a known temperature and using density tables. For safety reasons, the quantity of pure gas introduced into the calibrated bulb is such that the pressure is well below atmospheric pressure. The temperature T is the temperature of the room thermoregulated to ±0.1 ◦ C (a thermostatic bath, or an air-bath, limited to the calibrated volume may also be used). This room also provides a stable environment for the calorimeter. In the original apparatus, the pressure P was measured using a mercury manometer and a cathetometer [2]. Nowadays, for reasons of health hazards (avoiding mercury) and simplicity, a high-precision manometer may conveniently replace the reading of mercury levels. The interface between the vacuum line and the calorimeter is crucial to transfer the gas back and forth quantitatively. It is a one-way valve, made of a glass tube terminated by a sintered-glass disk (glass frit) of fine porosity, immersed in mercury [3]. For the transfer of boron trifluoride from the glass bulb to the calorimetric cell, standard techniques of gas manipulation in a vacuum line are used. The mercury cannot penetrate the pore of the sintered glass (if the porosity is appropriate), even under vacuum, and makes a vacuumtight valve. The gas is condensed in a small volume (about 1 ml), using liquid nitrogen and then gently warmed. When the pressure in the tube becomes greater than the pressure in the calorimeter (which is close to atmospheric pressure), the gas pushes the mercury and reaches the solution placed above the mercury pool. The gas flow should be controlled in order to avoid ‘jumps’ of the mercury pool. When enough gas is injected into the cell,

The Measurement of Lewis Basicity and Affinity in the Laboratory

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the remaining gas is condensed and returned to the calibrated bulb. The new pressure is measured after temperature stabilization. A Tian–Calvet heat flux calorimeter was used in the measurements described in ref. [2]. This type of calorimeter is also called ‘isothermal’ [4, 5], in contrast to other kinds of calorimeter. A tutorial [6] on heat-conduction calorimetry gives a good account of the technique. Peak integration of the heat flux against time may be performed by a numerical integration method, such as Simpson’s method, on a personal computer interfaced to the calorimeter [7]. This experiment describes the measurement of the BF3 affinity of pyridine in CH2 Cl2 solution, namely the enthalpy of reaction 7.2: BF3 (g) + C5 H5 N(soln) C5 H5 N−BF3 (soln)

(7.2)

7.1.2 Reagents and Equipment r Electronic-grade boron trifluoride gas (99.99+%), anhydrous pyridine (99.8%), anhydrous dichloromethane (99.8+%). It is essential to use dry solvents and reactants (including the mercury placed at the bottom of the calorimetric cell, to be dried under vacuum) because traces of humidity (and also other impurities) tend to induce an additional heat of reaction. Moreover, boron trifluoride releases hydrogen fluoride by slow hydrolysis, resulting in etching of the glass parts of the system. r Glassware and analytical balance for preparing the solutions. r Glass high-vacuum line. A detailed scheme of the actual apparatus is given in ref. [2], but the short description hereafter is sufficient to understand its principle and to design a device tailored to the student laboratory. The vacuum line comprises a high-vacuum pumping system (mechanical plus diffusion or turbomolecular pumps), fitted with a liquid nitrogen trap and vacuum manometers, a gas introduction and purification stage, a gas reservoir, a calibrated bulb fitted with a high-precision manometer for gas measurement and a vacuum-tight connection to the calorimetric cell. Bulbs and traps are designed with provision for freezing the gas using liquid nitrogen. High vacuum TeflonTM valves and joints (grease-less) are recommended. The mercury manometer should be replaced by a mercury-free device. For example, MKS (USA) makes corrosion-resistant capacitance R ) with a resolution of 10−5 in the 0.1–1 bar range, with an absolute manometers (Baratron accuracy of 0.05–0.1%. r Calorimetric cell (glass [2] or corrosion-resistant stainless steel) shown in Figure 7.1a. It is dried at a temperature above 100 ◦ C and flushed with a dry gas such as nitrogen or argon. The blanket of inert gas also avoids possible oxidation of some compounds. r Calibration cell (shown in Figure 7.1b). An electrical calibration may be performed using a stabilized power supply (current generator), a high-accuracy resistor, a precision voltmeter and an electronic chronometer. Very high absolute tolerance resistors (0.01–0.001%) with low temperature coefficient are available, for example from VishaySFERNICE, France. r Any good-quality heat-flux calorimeter or isothermal titration calorimeter, such as those marketed by TA Instruments (USA) and the associated company Thermometric (Sweden), CSC (USA), Microcal (USA) or Setaram (France). r Personal computer.

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Lewis Basicity and Affinity Scales To the vacuum line

To external manometer

Heat flux sensors of the calorimeter Frit a

To the power supply and voltmeter

Solution

Solvent

Mercury

Mercury High accuracy resistor b

Figure 7.1 (a) Gas/liquid calorimetric cell showing the solution of the Lewis base and the glass frit (sintered glass) plunged into mercury, making the one-way valve. The gas is introduced through the glass frit when a pressure greater than atmospheric is produced in the vacuum line. A pressure increase (read on the external manometer) indicates unreacted boron trifluoride. (b) Calibration cell (Joule effect). The end of the introduction tube is replaced by a highaccuracy resistor. This calibration device is plunged into the same volumes of mercury and solvent as used for measurements.

7.1.3

Experiment

The calorimeter can be calibrated by the Joule effect. The resistor is placed in a cell containing mercury and a liquid representative of the solution, in amounts identical with those for the measurements (Figure 7.1b). Another possibility is to replace this absolute calibration by the measurement of a known enthalpy of reaction with BF3 , which may be used as an indirect calibration. Enough mercury is put at the bottom of the calorimetric cell to cover the glass frit (sintered glass), the system that constitutes the one-way valve. The pyridine solution must not touch the frit. A known volume (3 ml in our case) of this solution, about 0.5 mol l−1 , is placed in the cell and the temperature of the system is allowed to settle. The volume and concentration of the solution give the number of moles of pyridine in the calorimeter, for a future evaluation of the stoichiometry. After measuring the pressure in the calibrated bulb, BF3 is condensed in the small bulb then slowly warmed until the pressure increase allows the gas to pass through the frit/mercury valve and to react with pyridine. Heat is instantaneously released and the gas injection is stopped after a few seconds by refreezing the gas with liquid nitrogen. The quantity of injected BF3 should be sufficient for a precise measurement of the gas quantity and of the heat, but not too large, to allow for several (at least five) Q/n measurements on the same solution. With a 100 ml bulb, this corresponds to about 40 mbar (4 × 103 Pa) of gas injected at each measurement The rapid heat evolution and the return to the baseline following an exponential law indicates a simple and rapid reaction, as


405

expected for the adduct formation between BF3 and pyridine. A secondary, slower reaction (polymerization, crystallization, isomerization) induces a shift of the baseline, usually in the exothermic direction. This partition between fast and slow events by heat-flux calorimeters is particularly useful. Conversely, a pressure increase over the solution (as observed on the external manometer) may indicate an incomplete reaction, solvent evaporation or an excess of boron trifluoride, and shows up as an endothermic shift of the baseline. This phenomenon happens at the saturation of the solution, and the approximate stoichiometry of the complex may be inferred from the comparison between the number of moles of base and of BF3 . 7.1.4

Results

These are summarized in Table 7.1, in which AQ (area) is the result of the peak integration, Q (in joules) is AQ multiplied by the calibration constant (1.0685 × 10−2 J per unit area), ∆P is the pressure difference (before and after each injection, in Pa = 10−5 bar or 10−2 mbar) measured in the calibrated bulb, n(BF3 ) is equal to (∆P)V/RT in moles, and Q/n is in kJ mol−1 . Remember that energy (heat) is leaving the system (exothermic reaction) so Q should be negative, like the enthalpy change. Six digits were used for intermediate calculations, although the expected final precision is about 0.5–1%. The value of the bulb calibration factor V/RT used in Table 7.1 (6.183 53 × 10−8 mol Pa−1 ) was established from the temperature (22.0 ◦ C) and the bulb volume (151.74 cm3 ). It is found that the ratio −Q/n, that is, the enthalpy for each injection, is reasonably constant, to within better than 1%. As often observed, the first injection leads to a −Q/n value that is not very consistent with the others; this is attributed to the presence of impurities. The mean of the remaining seven values (that is −∆H ◦ of complexation) is 127.69 kJ mol−1 . The repeatability (same sample, same conditions, same analyst) is ±0.64 kJ mol−1 at the 95% confidence level. The published value [8], −∆H ◦ = 128.08 ± 0.50 kJ mol−1 (95% confidence level, 16 individual values), corresponds to the combination of two Table 7.1 Calorimetric data for the reaction of gaseous BF3 with a 0.401 mol l−1 solution of pyridine in CH2 Cl2 a . Injection No. AQ −Q ∆P n(BF3 ) × 104 −Q/n n(BF3 ) × 104 1 1261 2 1260 3 1180 4 1035 5 2212 6 1143 7 1023 8 1790 Saturation 2794 ∆P = 22822.4 Pa

a b

13.4738 1739.53 13.4631 1719.62 12.6083 1585.50 11.0590 1414.20 23.6352 2977.12 12.2130 1539.02 10.9308 1384.99 19.1262 2420.74 29.8539 8040.36 n(BF3 ) = 1.411× 10−3 mol

1.07 564 1.07 564 125.263 1.06 333 2.13 897 126.613 0.980 399 3.11 937 128.604 0.874 475 3.99 384 126.464 1.84 091 5.83 475 128.389 0.951 658 6.78 641 128.334 0.856 413 7.64 283 127.635 1.49 687 9.13 970 127.774 4.97 178 14.1122 — Mean = 127.69 kJ mol−1 Standard deviation = 0.86 kJ mol−1 CIb = 0.79 kJ mol−1

Units of AQ , Q, ∆ P, n and Q/n are: arbitrary, J, Pa, mol and kJ mol−1 , respectively. 95% confidence interval.

406


-ΣQ/J mol

-1

120

80 40 0 0

5

10

15

20

Σn*10

4

Figure 7.2 Plot of heat evolved as a function of the number of moles of BF3 injected. The break in the curve may be used as an estimate of the number of BF3 moles reacted, assuming that the heat is nil after formation of the 1 : 1 complex.

series of experiments. See ref. [2] for a discussion about the reproducibility. The amount of pyridine in the calorimetric cell was 1.203 × 10−3 mol and the total amount of boron trifluoride injected was 1.411 × 10−3 mol, corresponding to an excess of about 17% BF3 for a 1 : 1 complex. After correction for the unreacted BF3 (overpressure in the calorimetric cell indicated on the external manometer), the molar ratio [BF3 ]/[pyridine] is estimated to be 0.96 ± 0.04, close to the expected 1 : 1 stoichiometry. Another way to treat and visualize the data is to plot ΣQ as a function of Σn(BF3 ), as in Figure 7.2. The proportionality between Q and n translates into a line of slope −Q/n (or −∆H ◦ ) until saturation of the solution, then Q/n (or −∆H ◦ ) ≈ 0.

7.2

7.2.1

Calorimetric Determination of the BF3 Affinity of Pyridine by Liquid/Liquid Reaction Introduction: Measuring Relative BF3 Affinity by Ligand Exchange in Solution

Considering the complexity of the gas measurement system and its coupling to the calorimeter, a simpler method related to titration calorimetry is also proposed. Here, the source of BF3 is the addition compound with diethyl ether, a liquid easier to handle than gaseous BF3 . The reaction studied is now (C2 H5 )2 O−BF3 (soln) + C5 H5 N(soln) C5 H5 N−BF3 (soln) + (C2 H5 )2 O(soln)

(7.3)

where (C2 H5 )2 O is displaced from its BF3 complex by pyridine, because this compound is a much stronger base than the ether. We have followed Brown and Horowitz [9] in choosing nitrobenzene as the solvent. Indeed, nitrobenzene is much less volatile than CH2 Cl2 . Moreover, a very good relationship has been established [8] between the BF3 affinities in dichloromethane and in nitrobenzene.


407

The experiment proposed corresponds to a variant of titration calorimetry. The titration is done by stepwise (discontinuous) injections of small increments of a pyridine solution into a given amount of the BF3 complex in solution, until a complete reaction. The amount of complex is calculated to give 5–10 injections before complete displacement. The enthalpy of dilution should be taken into account by measuring a ‘blank’ injection of pyridine in nitrobenzene. 7.2.2 Reagents and Equipment r Boron trifluoride diethyl etherate purified by distillation, anhydrous pyridine (99.8%), ˚ molecular sieves. As in experiment 7.1, it is essential nitrobenzene (>99.5%) dried on 4 A to use dry solvents and reactants. r Glassware, analytical balance. r Commercial titration calorimeter equipped with an injection device for the study of liquid/liquid reactions, such as those marketed by TA Instruments (USA) and the associated company Thermometric (Sweden), CSC (USA), Microcal (USA) or Setaram (France). For a student experiment, simpler systems based on Dewar flasks, such as those used by Brown et al. [9, 10], Arnett et al. [11] and Sherry and Purcell [12], may also be used. The calorimetric cell should be thoroughly dried by flushing with a dry gas such as nitrogen or argon. The blanket of inert gas also avoids contact with atmospheric humidity and oxygen. 7.2.3

Experiment

The calorimeter may be calibrated by the Joule effect (see experiment 7.1) or using a standard reaction, chosen among those suggested by an IUPAC technical report [13]. A 5 ml volume of the 0.5 mol l−1 solution of BF3 etherate in nitrobenzene is placed in the calorimetric cell under dry nitrogen or argon. The syringe pusher is filled with a sufficient volume of 2.5 mol l−1 pyridine solution. Using these values, 0.100–0.200 ml injection steps allow for 5–10 additions. For an isothermal titration calorimeter, the time interval between each injection should be sufficient for the signal to return to the baseline (for a Dewar calorimeter, allowance is made for temperature stabilization after each injection). When the temperature of the system is equilibrated, the data acquisition and the injection programme are started. A preliminary experiment should be carried out to measure the blank value, corresponding to the heat of dilution of pyridine solution in the pure solvent. 7.2.4

Results

The results corresponding to each injection are presented in a table, which is similar to Table 7.1 but with an added column corresponding to the correction for the heat of dilution of pyridine solution. Thus, the column −Q in Table 7.1 becomes corrected −Q, −Q(corr). Also, n becomes the number of moles of pyridine, equal to the molar concentration multiplied by the volume (in litres) of pyridine solution. As in experiment 7.1, the ratio −Q(corr)/n should be constant, since it is equal to the enthalpy of the displacement reaction. This enthalpy is also equal to the slope of the regression line of −ΣQ(corr) versus Σn, the number of moles of pyridine injected.

408


Table 7.2 BF3 affinities (kJ mol−1 ) of pyridine in C6 H5 NO2 : comparison of the results of the displacement and of the direct methods for two ethers. Displacement BF3 affinity BF3 affinity of pyridine enthalpy/kJ mol–1 of ether (displacement method)

Ether Et2 O c-(CH2 )5 O

−58.99 −51.04

81.35 86.45

BF3 affinity of pyridine (direct method)

140.34 137.49

137.86 137.86

The result found by Brown and Horowitz [9] for the enthalpy of the displacement reaction 7.3 is −58.99 kJ mol−1 . Combining this value with the enthalpy of reaction 7.4 in nitrobenzene: BF3 (g) + Et2 O(soln) Et2 O−BF3 (soln)

(7.4)

which is equal to −81.35 kJ mol−1 according to ref. [8], we arrive at the BF3 affinity of pyridine in nitrobenzene as −∆H (7.3) − ∆H (7.4) = 140.34 kJ mol−1 This affinity in nitrobenzene can be converted into the affinity in CH2 Cl2 by the relationship 7.5 established in ref. [8]: BF3 affinity in CH2 Cl2 = 0.958 × (BF3 affinity in nitrobenzene) − 0.31

(7.5)

We finally obtain 134.14 kJ mol−1 for the BF3 affinity of pyridine in CH2 Cl2 . It can be concluded that the displacement method yields a value in reasonable agreement with the direct method described in experiment 7.1, considering the various sources of error: r evaporation from the solution of the ethyl ether formed in the displacement reaction; r uncertainties in the enthalpies ∆H (7.3) and ∆H (7.4); r standard deviation (2.3 kJ mol−1 ) of relationship 7.5. Because of the evaporation of liberated Et2 O, Brown and Horowitz recommended the use of the addition compound of tetrahydropyran with BF3 (we did not follow this recommendation because the tetrahydropyran complex does not appear to be commercially available. However, it can be readily prepared and purified by distillation under reduced pressure [9]). Indeed, Table 7.2 shows that the displacement of tetrahydropyran yields a result in better agreement with the direct method than the displacement of Et2 O.

7.3

7.3.1

Determination by FTIR Spectrometry of the Complexation Constants of 4-Fluorophenol with Isopropyl Methyl Ketone and Progesterone Introduction: Recognition of Progesterone by its Receptor

˚ crystal structure of the progesterone 1–human receptor complex provides a The 1.8 A molecular basis for the recognition of this hormonal steroid by its receptor. Concerning hydrogen-bond interactions contributing to the stability of the complex, the carbonyl C3 O group participates in a specific and rigid network of hydrogen bonds, while the C20 O


409

O O 20

3 O 1

2

Scheme 7.1 Progesterone 1 and isopropyl methyl ketone 2.

group does not appear to accept any hydrogen bond. To explain this hydrogen-bond selectivity, the hydrogen-bond basicity of each carbonyl group of progesterone (Scheme 7.1) is studied in this experiment, by measuring the complexation constant of progesterone with 4-fluorophenol in CCl4 . Since progesterone has two potential hydrogen-bond acceptor sites, the measured constant, K t , corresponds to the sum of the constants of complexation to O3 and O20: K t = K (O3) + K (O20) To obtain the hydrogen-bond basicity of each carbonyl group separately, the 20-keto oxygen is modelled by the oxygen of 3-methylbutan-2-one (isopropyl methyl ketone) 2 (Scheme 7.1), and the basicity of the 3-keto oxygen is deduced by the difference K t – K(i-PrCOMe). The measurement of the complexation constants of 4-fluorophenol with isopropyl methyl ketone and progesterone is presented below and the experimental results are explained by the electronic structure of progesterone. A detailed study of the hydrogen bonding of progesterone can be found in ref. [14]. 7.3.2 Reagents and Equipment r Spectrometric-grade CCl4 dried on 4 A ˚ molecular sieves, 4-fluorophenol (99%) vacuum r r r r

sublimed over P2 O5 , progesterone (99+%), and isopropyl methyl ketone (99%) passed through a column of freshly activated basic aluminium oxide. Volumetric flasks, spatula, pipettes. Analytical balance. FTIR spectrometer flushed with dry air. A 1 cm pathlength Infrasil quartz cell placed in a cell holder thermoregulated at 25 ◦ C.

7.3.3 Experiment r Prepare 20 cm3 of stock solutions of ∼8 × 10−2 mol l-1 4-fluorophenol, 0.3 mol l−1 isopropyl methyl ketone and 0.1 mol l−1 progesterone in CCl4 . Two stock solutions of each ketone should be prepared. r Prepare four sample solutions of 4-fluorophenol and ketone in 10 cm3 volumetric flasks by mixing aliquots of stock solutions and CCl4 . The concentration of 4-fluorophenol must

410


be ∼4 × 10−3 mol l−1 to avoid self-association. The ketone concentrations are varied from 8 × 10−3 to 2.5 × 10−2 mol l−1 in order to complex 30–60% of 4-fluorophenol and to maintain the ketone in excess to avoid the formation of 2(4-fluorophenol) : 1(ketone) complexes. r Prepare four reference solutions of ketone in 10 cm3 volumetric flasks, so that the ketone concentration in the reference solution is exactly the same as that in the corresponding sample solution. r Record the FTIR spectra of reference solutions and then of sample solutions at a resolution of 1 cm−1 . r Display the spectra of the solute (a mixture of free and hydrogen-bonded 4-fluorophenol) in absorbance units.

7.3.4

Results and Discussion

The resulting spectra are shown in Figures 7.3 (isopropyl methyl ketone) and 7.4 (progesterone) in the region 3700–3100 cm−1 , where the sharp OH band of free 4-fluorophenol appears at 3614 cm−1 and the broad OH band of hydrogen-bonded 4-fluorophenol at lower wavenumbers. The characteristic shape of the band of the i-PrCOMe complex is attributed to the presence of two-component bands, corresponding to two stereoisomeric complexes

1

Absorbance

0.8

0.6

0.4 E D C 0.2 B

0 3700

A 3600

3500

3400

3300

3200

3100

Wavenumber / cm-1

Figure 7.3 Influence of the addition of increasing quantities of isopropyl methyl ketone on the OH band of 4-fluorophenol. Ketone concentrations are (A) 0, (B) 0.032, (C) 0.063, (D) 0.068 and (E) 0.115 mol l−1 .


411

1

Absorbance

0.8

0.6

E D

0.4

C B

0.2

A 0 3700

3600

3500

3400 Wavenumber / cm

3300

3200

3100

-1

Figure 7.4 Influence of the addition of increasing quantities of progesterone on the OH band of 4-fluorophenol. Progesterone concentrations are (A) 0, (B) 0.008, (C) 0.012, (D) 0.016 and (E) 0.024 mol l−1 .

[15]. The OH band of the complex with progesterone is still broader and more asymmetric, because of the existence of two 1 : 1 complexes. The absorbance A of the band at 3614 cm−1 (absorbance at maximum minus absorbance of baseline at 3800 cm−1 ) yields the equilibrium concentration Ca of free 4-fluorophenol from Beer’s law: Ca = A/εl

(7.6)

where l is the cell length in centimetres and ε the molar absorption coefficient. The value 233.9 l mol−1 cm−1 is found from spectra A (Ca0 = 4.097 mmol l−1 in the i-PrCOMe experiment, and 3.997 mmol l−1 in the progesterone experiment; l = 1 cm). The concentrations of complex, Cc , and base (ketone), Cb , at equilibrium are deduced from the initial concentrations, Ca0 and Cb0 , as follows: Cc = Ca0 − Ca Cb =

Cb0

− Cc =

Cb0

− Ca0 − Ca

(7.7) (7.8)

and the equilibrium constant K c is calculated by the expression Kc =

C 0 − Ca 0 a Ca Cb − Ca0 − Ca

(7.9)

412


Table 7.3 Calculation of the equilibrium constant (l mol−1 ) for the complex of 4-fluorophenol with isopropyl methyl ketone in CCl4 at 25 ◦ C. Molar mass: 112.1 g mol−1 Molar mass: 86.14 g mol−1 Molar mass: 153.82 g mol−1

Acid: 4-fluorophenol Base: isopropyl methyl ketone Solvent: carbon tetrachloride Sample 1 Absorbance 0.6341 Density 1.5 975 208 Ca0 × 103 4.057 657 Cb0 × 102 3.1 607 294 Ca × 103 2.710 638 Cb × 102 3.0 260 274 Cc × 103 1.347 019 Kc 16.42 % H-bonded 33.2 4-FC6 H4 OH K¯ c = 16.3 ± 0.2 (95% CI)

Sample 2

Sample 3

Sample 4

0.4707 1.5 969 415 4.036 969 6.3 438 622 2.012 139 6.1 413 792 2.024 829 16.39 50.2

0.4563 1.594 045 4.040 075 6.8 243 067 1.950 582 6.6 153 574 2.089 493 16.19 51.7

0.32 406 1.5 936 588 3.936 366 11.5 405 365 1.385 285 11.2 854 284 2.551 081 16.32 64.8

The calculations are easily performed in a spreadsheet as illustrated in Table 7.3 for isopropyl methyl ketone and Table 7.4 for progesterone. The concentrations Ca0 and Cb0 depend on the temperature at which the solutions were prepared. Consequently, they must be corrected for density changes. The constant K t = 77 l mol−1 of progesterone can thus be separated into two constants, K(O20) = 16 l mol−1 and K(O3) = 77 – 16 = 61 l mol−1 . The latter value compares well with the value of 56 l mol−1 found for isophorone [15]. Indeed, isophorone 3 is a good model of the cyclohexenone moiety of progesterone 1 (Scheme 7.2). Table 7.4 Calculation of the overall equilibrium constant (l mol−1 ) for the two 1 : 1 complexes of 4-fluorophenol with progesterone in CCl4 at 25 ◦ C. Molar mass: 112.1 g mol−1 Molar mass: 314.46 g mol−1 Molar mass: 153.82 g mol−1

Acid: 4-fluorophenol Base: progesterone Solvent: carbon tetrachloride Sample 1 Absorbance 0.6204 Density 1.5 988 725 Ca0 × 103 4.04843 Cb0 × 103 8.206 366 Ca × 103 2.652 617 Cb × 103 6.810 553 Cc × 103 1.395 813 Kc 77.26 % H-bonded 34.5 4-FC6 H4 OH K¯ c = 76.6 ± 1.4 (95% CI)

Sample 2

Sample 3

Sample 4

0.4487 1.5 981 001 4.012 952 16.427 029 1.918 487 14.332 564 2.094 465 76.17 52.2

0.5194 1.5 965 553 4.012 612 12.225 017 2.220 776 10.433 181 1.791 836 77.34 44.7

0.3666 1.594 045 4.073 905 23.662 727 1.567 456 21.156 277 2.50 645 75.58 61.5


413

3

O

O 1

3

Scheme 7.2 Similarity of isophorone 3 with the cyclohexenone moiety of progesterone 1.

These results show that oxygen O3 is more basic by 0.58 pK unit, that is, 3.3 kJ mol−1 on the Gibbs energy scale, than oxygen O20. This greater hydrogen-bond basicity explains, in part, the preference of the human receptor for forming hydrogen bonds with O3 rather than with O20. The greater hydrogen-bond basicity of O3 can be rationalized from two quantumchemical descriptors of the oxygen lone pairs. First, the ionization energy of the oxygen lone pair is lower for O3 (+10.66 eV) than for O20 (+10.77 eV) (in Koopman’s approximation); consequently, the charge-transfer component of the hydrogen-bond energy is greater for O3 than for O20. Second, the electrostatic potential is more negative around O3 (−201.7 kJ mol−1 ) than around O20 (−169.0 kJ mol−1 ); consequently, the electrostatic component of the hydrogen-bond energy also favours a binding to O3. In terms of the resonance concept, the basicity difference results from a greater contribution of the canonical form 1b to the electronic structure of 1 than of 2b to the structure of 2, that is, to an additional negative charge given to O3 by the conjugation of the double bond with the carbonyl group (Scheme 7.3).

7.4

Determination by FTIR Spectrometry of the Complexation Enthalpy and Entropy of 4-Fluorophenol with Cyclopropylamine

7.4.1

Introduction

In this experiment, the enthalpy and entropy of complexation of 4-fluorophenol with cyclopropylamine is measured from the temperature variation of the complexation constant on the mole fraction scale, Kx . The experimental result is compared with high-level calculations [MP2/aug-cc-pVTZ//B3LYP/6–31+G(d,p)] of the enthalpy in vacuo. For this reason,

O

O 3 O

O

1a

1b

2a

2b

Scheme 7.3 Canonical forms 1a and 1b of progesterone and 2a and 2b of isopropyl methyl ketone.

414


cyclohexane is chosen as the solvent because it is expected that the solvation term will be smaller for cyclohexane than for the usual solvent CCl4 . A new procedure is implemented for an accurate and fast determination of the hydrogenbond enthalpy. In this method, a single solution of cyclopropylamine and 4-fluorophenol is studied by measuring its absorbance change as a function of temperature. 7.4.2 Reagents and Equipment r Spectrometric-grade cyclohexane and cyclopropylamine (98%) dried on 4 A ˚ molecular sieves, 4-fluorophenol as in experiment 7.3. r Device for regulating temperature from 10 to 70 ◦ C. r Thermocouple. The temperature is taken in the cell centre, just after the spectral record of the reference solution. r The spectrometer and the cell are the same as in experiment 7.3, except for the cell pathlength, which is reduced to 0.2 cm because of the lower IR transparency of cyclohexane compared with CCl4 . 7.4.3 Experiment r Determination of the temperature dependence of the absorption coefficient of the OH band of 4-fluorophenol. Prepare a sample solution of 4-fluorophenol in cyclohexane with a 4-fluorophenol mole fraction of 3.6 × 10−4 . Ensure that the cell has equilibrated at the desired temperature. Record the reference spectra (neat cyclohexane) at 10, 25, 40, 55 and 65 ◦ C, then the sample spectra at the same temperatures. r Prepare a single sample solution containing 3.6 × 10-4 mole fraction of 4-fluorophenol and 2.3 × 10−3 mole fraction of cyclopropylamine in cyclohexane, and a reference solution containing only cyclopropylamine at the same mole fraction as the sample. Record the reference spectra, then the sample spectra at the previous temperatures. r Display the spectra of the solute (a mixture of free and hydrogen-bonded 4-fluorophenol) in absorbance units. 7.4.4

Results

The spectra at various temperatures are shown in Figure 7.5. The sharp band near 3620 cm−1 and the broad band near 3140 cm−1 are attributed to the OH stretching of free and hydrogenbonded 4-fluorophenol, respectively. The ν as and ν s (NH2 ) bands of cyclopropylamine cannot be well compensated in the sample minus reference spectra subtraction because their frequencies and intensities vary on complexation. Hence they appear near 3367 cm−1 . The variation with temperature of the absorbance of the band near 3620 cm−1 in the spectra of a 4-fluorophenol solution enables the absorption coefficient at several temperatures to be calculated from Beer’s law: ε = A/xl

(7.10)

where ε is in cm−1 , x in mole fraction and l = 0.2 cm. The temperature dependence of the absorption coefficient can be deduced using Equation 7.11, where t is the temperature


415

0.12

0.1

Absorbance

0.08

0.06

10 25

0.04

40 55 65

0.02

0 3640

3540

3440

3340

3240

3140

3040

-1

Wavenumber / cm

Figure 7.5 FTIR spectra of a single solution of 3.6 × 10−4 mole fraction 4-fluorophenol and 2.3 × 10−3 mole fraction cyclopropylamine in cyclohexane at various temperatures in a 0.2 cm cell. The absorbance of the sharp OH band of free 4-fluorophenol near 3620 cm−1 decreases with decrease in temperature (+65 to 10 ◦ C) with a concomitant absorbance increase in the broad OH band of the hydrogen-bonded complex near 3140 cm−1 . The missing parts of the spectra correspond to energy losses caused by the cyclohexane absorption.

in ◦ C: ε(t) = ε(25) − 18.969 (t − 25) + 0.0505 (t 2 − 252 )

(7.11)

The absorbance of the free OH band in the spectra of the single solution of 4-fluorophenol and cyclopropylamine in cyclohexane allows the calculation of the equilibrium mole fraction of 4-fluorophenol, xa , cyclopropylamine, xb , and the complex, xc , and the equilibrium constant Kx at each studied temperature, from the initial mole fractions xa0 and xb0 and Equations 7.12–7.15: xa = A/εl

(7.12)

xc = xa0 − xa − xc =

xb = Kx =

xa0 − xa xa xb0 − xa0 − xa

xb0

−

xb0

xa0

− xa

(7.13) (7.14) (7.15)

416


Table 7.5 Calculations of xa0 , xb0 , xa , xb, xc , K x , ∆H◦ and ∆Sx◦ for the complex of cyclopropylamine with 4-fluorophenol in cyclohexane. Molar mass: 112.1 g mol−1 Molar mass: 57.09 g mol−1 Molar mass: 84.16 g mol−1 dd/dt = 0.00 101 xa0 = 3.62 423 × 10−4 xb0 = 22.81 139 × 10−4

Acid: 4-fluorophenol Base: cyclopropylamine Solvent: cyclohexane Density (25 ◦ C) = 0.7735 Pathlength = 0.2 cm Temperature (◦ C)

9.5

25.2

40.2

55.9

65.7

Absorbance Absorption coefficient xa × 104 xb × 104 xc × 104 H-bonded 4-FC6 H4 OH Kx lnKx 1/T × 103

0.0586 2797.7

0.0839 2527.4

0.0994 2292.4

0.1112 2070.7

0.1142 1945.0

1.0478 20.2399 2.5773 71%

1.6592 20.8502 1.9657 54%

2.16 829 21.35 824 1.45 647 40%

2.6846 21.8736 0.93 998 25%

2.9360 22.1245 0.68 844 18%

1215.26 7.10 3.53 794

568.20 6.34 3.35 177

314.50 5.75 3.19 132

160.08 5.08 3.03 905

105.98 4.66 2.95 116

Least-squares treatment of the van’t Hoff plot Slope 4119.8835 Intercept Standard deviation of slope 90.8772 Standard deviation of intercept Squared correlation 0.9985 Standard deviation of estimate coefficient ∆H◦ = −34.25 kJ mol−1 ∆Sx◦ = −62.0 J K−1 mol−1 Kx (25 ◦ C) = 580

−7.4553 0.2927 0.0430

The complexation enthalpy, ∆H ◦ , is obtained from the slope of lnKx versus 1/T, the so-called van’t Hoff plot: ∆H ◦ = −R d(ln K x ) /d(1/T )

(7.16)

and the complexation entropy relative to mole fraction, ∆Sx◦ , from Equations 7.17 and 7.18: ∆G ◦x = −RT ln K x ∆Sx◦ = ∆H ◦ − ∆G ◦x /T

(7.17) (7.18)

The calculations are detailed in the spreadsheet of Table 7.5. The same kind of experiment can be repeated with CCl4 as the solvent. The result, ∆H ◦ = −31.88 kJ mol−1 , shows that the transfer of the reagents and product from cyclohexane to CCl4 is accompanied by a complexation enthalpy change of +2.37 kJ mol−1 . Since the complexation entropy change from cyclohexane to CCl4 is also positive (+7.1 J K−1 mol−1 ), on the whole the hydrogen-bond basicity of cyclopropylamine is less sensitive to the solvent change than its affinity. The comparison of the hydrogen-bond basicity of cyclopropylamine, pK BHX = 1.74, with that of a primary amine with the same carbon content, isopropylamine, pK BHX = 2.22 (see Table 4.4), shows that the cyclopropyl substituent decreases the basicity of the NH2 group significantly. In contrast, in ketones, the cyclopropyl substituent increases the


417

Figure 7.6 B3LYP/6–31+G(d,p) structure of the hydrogen-bonded complex of 4-fluorophenol with cyclopropylamine.

basicity of the carbonyl group: pK BHX increases by 0.12 pK unit on going from i-PrCOMe to c-PrCOMe (see Table 4.16). The versatility of the electronic effects of this substituent is attributed to the unsaturated character of the cyclopropane ring. 7.4.5

Comparison with Theoretical Calculations

Theoretical calculations were carried out using Gaussian 03 software. The geometries of cyclopropylamine, 4-fluorophenol and the complex were fully optimized at the B3LYP level using the 6–31+G(d, p) basis set, a split valence 6–31G basis augmented by polarization functions on all atoms and diffuse functions on non-hydrogen atoms. Three stationary points were detected in the potential energy surface (PES) of the complex and confirmed as true minima via vibrational frequency calculations. The structure of the deepest minimum is displayed in Figure 7.6. ◦ , is given by Equation 7.19 (explained in Chapter 1): The binding enthalpy, ∆H298 ◦ ∆H298 = ∆E el + BSSE + ∆ZPVE + ∆E vib,therm + ∆E rot + ∆E trans + ∆n RT = −47.19+5.29+5.92 + 11.24−3.72 − 3.72 − 2.48 = −34.66 kJ mol−1 (7.19)

Because ∆Eel is the dominant term in the enthalpy expression, single-point electronic energy calculations were performed at higher levels than the level of geometry optimization. As expected and shown in Table 7.6, the negative enthalpy increases with the level of computation. Table 7.6 Variation of the computed enthalpy with the level of calculation and of the experimental enthalpy with the medium. ◦ ∆H298 /kJ mol−1

Computational level

Computed

Experimental

B3LYP/6–31+G(d,p)//B3LYP/6–31+G(d,p) MP2/6–31+G(d,p)//B3LYP/6–31+G(d,p) MP2/6–31+G(3df,2p)//B3LYP/6–31+G(d,p) MP2/aug-cc-pVTZ//B3LYP/6–31+G(d,p) Medium

−27.81 −30.20 −33.54 −34.66 Gas

−34.25 Cyclohexane

−31.88 CCl4

418


A satisfactory convergence to a common gas-phase calculated value of −34.7 kJ mol−1 is observed in Table 7.6 when both the computational level and the ‘gas-phase-like’ behaviour of the medium increases.

7.5

7.5.1

FTIR Determination of the OH Shift of Methanol Hydrogen Bonded to Pyridine, Mesitylene and N-Methylmorpholine Introduction

In the following experiment, the IR spectra of three hydrogen-bonded complexes of methanol are measured in the range 4000–3000 cm−1 , where the only fundamental absorption is attributed to the stretching of the O H bond. In the gas phase, the wavenumber of the band maximum is at 3681 cm−1 . In CCl4 solution, intermolecular interactions shift the band to 3644 cm−1 . In the following, this band will be referred to as the OH band of ‘free’ methanol, even if the O H group is (very) weakly hydrogen bonded to the chlorines of CCl4 . This band is sharp (half-band width of 20 cm−1 ), its integrated intensity is 25 500 m mol−1 and its absorption coefficient is 67 ± 1 l mol−1 cm−1 (mean of 28 determinations). When a Lewis base is added to a CCl4 solution of methanol, a broader and more intense band is observed at lower wavenumbers. This is attributed to the OH band of the hydrogenbonded complex of methanol with the base. The ‘spectroscopic basicity’ towards methanol is calculated as the difference ∆ν = ν(free OH) − ν(hydrogen-bonded OH)

(7.20)

This scale has been discussed and tabulated for ∼800 organic bases in Chapter 4. Here ∆ν(OH) is measured for three bases: pyridine, mesitylene (1,3,5-trimethylbenzene) and N-methylmorpholine. In the first case, one hydrogen-bonded OH band is observed with a nearly Gausso-Lorentzian shape. In the second case, the wavenumber shift is not large because mesitylene is a weak base, so the free band and the band of the complex overlap each other. The spectrum can be resolved by either (i) a mathematical separation into GaussoLorentzian component bands using a program included in the software treating the FTIR spectra, or (ii) subtracting the absorbance due to the free OH band. N-methylmorpholine is a polyfunctional base with an ether and an amine function; it forms two 1 : 1 complexes, when the base is in excess over methanol, and two bands are observed, one for the O complex and the other for the N complex.

7.5.2 Reagents and Equipment r Spectroscopic-grade methanol, dried on 3 A ˚ molecular sieves, and CCl4 , anhydrous-grade pyridine, mesitylene and N-methylmorpholine purified by distillation and carefully dried ˚ molecular sieves. on 4 A r Two 10 cm3 volumetric flasks, pipettes. r FTIR spectrometer flushed with dry air operating at 1 cm−1 resolution. r A 1 cm pathlength Infrasil cell placed in a cell holder thermoregulated at 25 ◦ C.


7.5.3

419

Experiment

Good spectra correspond to a significant absorbance of the OH band of methanol hydrogen bonded to the base and to the absence of the OH band(s) of self-associated methanol. Hence the base and/or methanol concentrations must be high enough to displace the equilibrium towards the formation of the complex. However, the methanol concentration must be kept below the limit of significant self-association (i.e. below ∼7 × 10-3 mol l-1 ). Moreover, the addition of base to CCl4 must not change the properties of the medium significantly. The experiment is described for pyridine. r Prepare a stock solution of ∼0.3 mol l−1 of pyridine in CCl4 (0.25 cm3 of pyridine in 10 cm3 of CCl4 ). Record its spectrum (reference spectrum). r Add a trace of methanol to the pyridine solution. Record the spectrum of this sample solution (sample spectrum). r Display the absorbance spectrum of the solute (a mixture of free methanol and of methanol hydrogen bonded to pyridine). If the absorbance of the band at 3644 cm−1 (free methanol) is higher than 0.3, dilute the sample solution with the stock solution. If the absorbance is lower than 0.2, add another trace of methanol. The experiment is repeated with mesitylene and N-methylmorpholine by preparing stock solutions of 1.2 mol l−1 mesitylene and 0.2 mol l−1 N-methylmorpholine. 7.5.4


The spectra of the three complexes are shown in Figure 7.7. The spectroscopic basicities towards methanol are as follows:

mesitylene pyridine N-methylmorpholine

∆ν ∆ν ∆ν ∆ν

= 3644 − 3594 = 50 cm−1 = 3644 − 3358 = 286 cm−1 = 3644 − 3502 = 142 cm−1 (O complex) = 3644 − 3261 = 383 cm−1 (N complex)

The results may vary from laboratory to laboratory because the shifts depend on the concentration of base. In the case of pyridine, the concentration gradient d(∆ν OH)/dCb is ∼13 cm−1 l mol−1 in the range 0–1 mol l−1 pyridine concentration, Cb . The shift of the hydrogen-bonded OH band to low wavenumbers is caused by a clustering of polar pyridine molecules around the polar methanol–pyridine complex. The dependence of ∆ν(OH) on base concentration can be corrected by extrapolation to infinite dilution. This tedious procedure can be avoided by choosing as low a concentration of base as allowed by the accurate measurement of the maximum of the hydrogen-bonded OH band. In order to facilitate the choice of the base concentration, Table 7.7 gathers a number of complexation constants and of absorption coefficients of the hydrogen-bonded OH band. The absorbance of this band can be estimated from these data, using Beer’s law and the expression of the complexation constant.

420

Lewis Basicity and Affinity Scales 0,4 a

b

c

d

e

Absorbance

0,3

0,2

0,1

0 3800

3100 -1

Wavenumbers / cm

Figure 7.7 OH stretching band in the IR spectra of methanol, free (a) and hydrogen bonded to mesitylene (b), N-methylmorpholine (c, O-complex; e, N-complex) and pyridine (d).

7.6

7.6.1

Solvatochromic Shifts of 4-Nitrophenol upon Hydrogen Bonding to Nitriles Introduction

The longest wavelength absorption band of the UV spectrum of 4-nitrophenol is at 263.3 nm (37 980 cm−1 ) in the gas phase. This band is attributed to a π → π ∗ transition with intramolecular charge-transfer character. In acetonitrile, the band is red shifted to 306.7 nm Table 7.7 Complexation constants, Kc (l mol−1 ), of methanol with selected organic bases in CCl4 at 25 ◦ C. Absorption coefficients, ε c (l mol−1 cm−1 ), of the corresponding hydrogen-bonded OH band a . Base

Kc

εc

Base

Kc

εc

Mesitylene Fluorocyclohexane Chlorocyclohexane

0.34 0.38 0.32

102 126 71

8.87 11.29 43.85

192 132 182

Acetonitrile Tetrahydrofuran Tetrahydrothiophene

1.31 1.35 0.45

143 149 75

Trimethyl phosphate Dimethyl sulfoxide Hexamethylphosphoric triamide Pyridine N-oxide Pyridine N-Methylmorpholine

Methyl acetate

0.98

a b

65; 99b

M. Lucon, University of Nantes, personal communication. Two stereoisomeric complexes.

Triethylamine

12.20 138 2.81 133 1.0 (O complex) 2.8 (N complex) 4.97 118


421

(32 600 cm−1 ). The solvatochromic shift of 5380 cm−1 can be considered as the sum of two terms, one due to the formation of the hydrogen bond 4-NO2 C6 H4 OH· · ·NCCH3 , and the other due to other interactions. The latter contribution can be evaluated by using 4-nitroanisole, which resembles 4-nitrophenol in all ways except in the ability to hydrogen bond. The hydrogen-bond shift can be calculated by a solvatochromic comparison method as explained in Chapter 4. This method yields Equation 7.21: ∆ν˜ (OH − OMe) = [1.0381ν(OMe) ˜ − 384] − ν˜ (OH)

(7.21)

where ∆˜ν (OH − OMe) is the hydrogen-bond shift and ν˜ (OMe) and ν˜ (OH) are the wavenumbers of the longest wavelength transition of 4-NO2 C6 H4 OMe and 4-NO2 C6 H4 OH, respectively. Here, ∆˜ν (OH − OMe) is determined for three nitriles: chloroacetonitrile, acetonitrile and dimethylcyanamide. The results are analysed in terms of substituent effects on the basicity of the nitrile group and compared with other hydrogen-bond basicity scales, in order to test the significance of electronic spectral shifts as basicity parameters. 7.6.2 Reagents and Equipment r 4-Nitrophenol, recrystallized repeatedly (m.p. 113 ◦ C) from ethanol, and 4-nitroanisole, fractionally distilled. r Spectrometric-grade acetonitrile. Chloroacetonitrile and dimethylcyanamide purified by distillation and carefully dried on molecular sieves. r Micro-syringe, spatula, pipettes. r Any good-quality computer-assisted UV–visible spectrometer. r A 1 mm pathlength Suprasil quartz cell, placed in a cell holder thermoregulated at 25 ◦ C. For less transparent nitriles (e.g. benzonitrile), a sealed 30 µm cell is recommended. A 1 cm cell is needed for the nitriles (e.g. pentadecanenitrile) in which 4-nitrophenol and/or 4-nitroanisole are only slightly soluble. 7.6.3

Experiment

The experiment is the same for each nitrile: (a) Prepare the two following solutions. r Solution 1. The nitrile is pipetted into the cell and a small crystal of 4-nitrophenol is dissolved directly in the cell. The concentration is adjusted to between 0.5 and 1.0 absorbance units at the band maximum. r Solution 2. A drop of 4-nitroanisole is added to the nitrile in the cell and the absorbance value of the solution is adjusted in the same range. (b) Record the absorbance spectra, S1 and S2 , of each solution, and also the spectrum of the pure nitrile, S0 , against air in the reference beam, with a 2 nm slit-width, over the range 400–250 nm. (c) Do the following subtractions: r S1 minus S0 : spectrum of the solute 4-nitrophenol in the nitrile solution. r S2 minus S0 : spectrum of 4-nitroanisole. (d) Locate the overall band maximum with the appropriate algorithm of the spectrometer software package.

422

Lewis Basicity and Affinity Scales 1,2 1

Absorbance

0,8 0,6 0,4 0,2 0

285.2

250

306.7 Wavelength / nm

380

Figure 7.8 UV spectra (normalized to absorbance 1) of 4-nitrophenol in cyclohexane (red) and acetonitrile (green) at 25 ◦ C. The 2460 cm−1 (21.5 nm) bathochromic shift decomposes into a 785 cm−1 shift upon hydrogen bonding to the nitrile group and a 2460 − 785 = 1675 cm−1 shift caused by stronger van der Waals interactions in acetonitrile (relative permittivity: 36.0) than in cyclohexane (relative permittivity: 2.02).

7.6.4


Representative spectra are shown in Figure 7.8. The values of ν˜ (OH) and ν˜ (OMe) measured on spectra and ∆˜ν (OH − OMe) calculated by Equation 7.21 are given in Table 7.8. The solvatochromic comparison graph is shown in Figure 7.9. The quality of the results depends on: r A good calibration of the spectrometer. Calibration can be made on the prominent 360.9 nm band of a holmium oxide glass. r A good solvatochromic comparison line. The line equation proposed here has been calculated from the data collected in Table 7.9, by a regression minimizing deviations in the direction normal to the regression line, since the errors on the two axes are the same (about ±30 cm−1 ). Table 7.8 Wavenumbers (cm−1 ) of the longest wavelength π → π ∗ transition of 4-nitrophenol, ν˜ (OH), and 4-nitroanisole, ν˜ (OMe), dissolved in chloroacetonitrile, acetonitrile and dimethylcyanamide. Corresponding hydrogen-bond shifts, ∆˜ν (OH − OMe) (cm−1 ). Comparison with other hydrogen-bond basicity scales. Solvent

ν˜ (OH)

ν˜ (OMe)

∆˜ν (OH− OMe)

∆˜ν (NH2 − NMe2 ) a

∆ν(OH)b

4-FC6 H4 OH affinityc

Chloroacetonitrile Acetonitrile Dimethylcyanamide

32 340 32 600 32 050

31 890 32 530 32 310

381 785 1107

480 1030 1370

48 75 117

16.25 19.29 22.36

Hydrogen-bond shift of 4-nitroaniline in cm−1 . Infrared OH shift of methanol in cm−1 . c kJ mol−1 . a

b


423

Wavenumber (OH) / kcm

-1

39

37

35

33 3 1

2

31 31

33

35

37

39

Wavenumber (OMe) / kcm-1

Figure 7.9 Solvatochromic comparison of 4-nitrophenol and 4-nitroanisole. Chloroacetonitrile 1, dimethylcyanamide 2 and acetonitrile 3 are displaced from the regression line of non-hydrogen-bonding solvents (◦). Data from Tables 7.8 and 7.9.

Table 7.9 Wavenumbers (kcm−1 ) of the longest π → π ∗ transition of 4-nitrophenol, ν˜ (OH), and nitroanisole, ν˜ (OMe), in the gas phase and 32 non- or very weak hydrogen-bond acceptor and hydrogen-bond donor solvents at 25 ◦ C. Medium

ν˜ (OH) ν˜ (OMe) Medium

ν(OH) ˜ ν˜ (OMe)

Gas phase (105 ◦ C) Perfluorohexane FC-75a Perfluoro(methylcyclohexane) Perfluorodecalin 2-Methylbutane Pentane Tetramethylsilane 1,1,1-Trichlorotrifluoroethane Hexane Hexafluorobenzene Trichlorofluoromethane Heptane Pentafluorobenzene Methyltrichlorosilane Dodecane

37.98 36.26 36.06 36.24 36.09 35.49 35.40 35.35 35.00 35.35 34.23 34.82 35.26 34.06 34.26 35.03

35.06 33.69 33.21 33.38 34.21 34.54 33.30 33.94 34.72 33.15 34.39 33.46 33.89 33.15 33.29 32.46

a

37.08 35.26 35.11 35.26 35.07 34.49 34.47 34.41 34.09 34.38 33.47 33.89 34.27 33.17 33.28 34.15

Cyclohexane 1,4-Difluorobenzene 1,2-Dichloroethane cis-1,2-Dichloroethylene trans-1,2-Dichloroethylene Carbon tetrachloride 1,1,2-Trichloroethane Trichloroethylene cis-Decalin 1,2,3-Trichloropropane Tetrachloroethylene 1,3-Dichlorobenzene Hexachloropropene 1,2-Dichlorobenzene 1,3-Dibromobenzene 1,2-Dibromobenzene

A mixture of the positional isomers of perfluoro-n-butyltetrahydrofuran (C8 F16 O).

34.12 32.95 32.37 32.40 33.34 33.62 32.25 32.96 33.91 32.24 33.52 32.55 33.18 32.27 32.27 31.96

424


The variation in the hydrogen-bond shifts of the three nitriles follows the electronic effects of the substituents. Compared with acetonitrile, the inductive electron-withdrawing effect of the ClCH2 substituent decreases the basicity, whereas the resonance electron-donating effect of the Me2 N substituent increases the basicity. These results are also consistent with other spectroscopic and thermodynamic scales of hydrogen-bond basicity, as shown in Table 7.8. Finally, the total red shift of 5380 cm−1 of the solvatochromic band of 4-nitrophenol, on going from the gas phase to acetonitrile solution, has been resolved into a 785 cm−1 hydrogen-bond shift (15% of the total shift) and a 4595 cm−1 shift caused by nonspecific (van der Waals) interactions. With triethylamine as solvent, the hydrogen-bond shift amounts to 42% of the total shift. These examples show the importance of quantifying the Lewis basicity of solvents for a quantitative description of solvent effects in chemistry.

7.7

7.7.1

Determination of the Complexation Constant of Diiodine with Iodocyclohexane by Visible Spectrometry Introduction: Measuring the Weak Diiodine Basicity of Haloalkanes

Haloalkanes are weak Lewis bases towards all Lewis acids. For this reason, they have attracted little attention in the construction of Lewis basicity scales. In this experiment, their basicity towards diiodine is studied by measuring the complexation constant of diiodine with iodocyclohexane in cyclohexane and comparing the result with those for other haloalkanes. Visible spectrometry is the method of choice for measuring diiodine basicity because the visible band of diiodine and the blue-shifted band of its complexes have rather high absorption coefficients (∼1000 l mol−1 cm−1 ). Consequently, there are significant variations in absorbance on complexation, even when small quantities of complex are formed, and accurate complexation constants can be obtained for weak Lewis bases. A disadvantage of electronic spectrometry is the width of the absorption bands. Hence, even significant blue shifts on complexation cannot prevent the overlap of the visible bands of free and complexed diiodine. Therefore, there are two unknowns, the complexation constant and the absorption coefficient of the complex, in the equation relating the equilibrium constant to the absorbance and the initial concentrations of diiodine and base. A wide variety of methods has been devised for the treatment of experimental data. The simple and rigorous method of Rose–Drago is used here.

7.7.2

The Rose–Drago Method

For a 1 : 1 complex (C) formed between a Lewis acid (A = I2 ) and a Lewis base (B = C6 H11 I) according to the reaction A + B C, the expression of the equilibrium constant in molar concentration units is K =

(C 0a

Cc − Cc )(Cb0 − Cc )

(7.22)


425

where Cc is the concentration of the complex and Ca0 and Cb0 are the initial concentrations of acid and base, respectively. Equation 7.22 may be rearranged to C 0C 0 −1 0 0 (7.23) K = Cc − Ca + Cb + a b Cc The concentration of the complex is obtained from the visible band of diiodine as follows. A highly diluted cyclohexane solution (S0 ) of diiodine at concentration Ca0 obeys Beer’s law at any wavelength λ: A0 = εa Ca0l

(7.24) −1

−1

where A is the absorbance, εa the absorption coefficient of diiodine in l mol cm at λ and l the cell pathlength in centimetres. A solution (S1 ) containing diiodine at initial concentration Ca0 and the base at initial 0 concentration Cb , hence the complex at concentration Cc , shows the absorptions of I2 , B and C. The total absorbance (A), at wavelength λ, is (7.25) A = εa Ca0 − Cc l + εb Cb0 − Cc l + εc Ccl 0

where ε b and εc are the absorption coefficients of the base and complex, respectively. When the base does not absorb at the studied wavelength, Equation 7.25 becomes (7.26) A = εa Ca0 − Cc l + εc Ccl Substituting εa = A0 /Ca0l into Equation 7.26 and rearranging allows the concentration of the complex to be expressed as A − A0 Ca0 /Ca0 (7.27) Cc = (εc − εa ) l Substituting this expression for Cc into Equation (7.23) yields A − A0 Ca0 /Ca0 Ca0 Cb0 (εc − εa ) l −1 0 0 K = − Ca + Cb + (εc − εa ) l A − A0 Ca 0 /Ca0

(7.28)

The general Equation 7.28 contains two unknowns, K −1 and εc . If a second equation like 7.28 is set up for solution S2 containing a concentration of diiodine and/or base different from solution S1 , an analytical solution can be derived specifically for εc as follows. Equation 7.28 can be written in the form α (7.29) y = +β +γx x where y = K −1

(7.30)

x = (εc − εa ) l α = A − A0 Ca0 /Ca0 β = − Ca0 + Cb0

(7.31) (7.32) (7.33)

426


and

γ =

Ca0 Cb0 A − A0 Ca 0 /Ca0

(7.34)

The equality 7.35 of the right-hand terms of Equation 7.29 written for solutions S1 and S2 yields the quadratic Equation 7.36: α2 α1 (7.35) + β1 + γ1 x = + β2 + γ2 x y= x x (γ2 − γ1 )x 2 + (β2 − β1 )x + (α2 − α1 ) = 0 which provides two roots: x=

−(β2 − β1 ) ±

(β2 − β1 )2 − 4(γ2 − γ1 )(α2 − α1 ) 2 (γ2 − γ1 )

(7.36)

(7.37)

When the indices one and two are chosen such that Cb0 (S1 ) < Cb0 (S2 ), only the root with the + sign in front of the square root has a physical significance. Finally, the substitution of x into Equation 7.29 yields K −1 . If the experiment is repeated with solutions S3 and S4 , the corresponding root, x34 , should ideally be equal to the root obtained with solutions S1 and S2 , x12 , as there is only one εc for a given system. In practice, experimental errors on absorbances and concentrations, and also model errors, such as the influence of the base concentration on the spectral properties of the system, yield different roots for different pairs of data. If n solutions have been studied, there are n(n − 1)/2 roots for all possible pairs of data, and hence n(n − 1)/2 values of K −1 . If each of these roots is given equal weight, the mean of K values can be used as the best value, and the standard error of the mean as an estimate of the uncertainty. To determine equilibrium constants accurately, two conditions must be met. First, the base concentration must vary as much as allowed by other constraints (solubility, ideality, change of the medium, etc.). Second, the wavelength λ of the measurements should be chosen such that the A – A0 term of Equation 7.28 is the largest possible. This often occurs near the maximum of the band of either free diiodine or the complex. In any case, it is recommended that the variation of K and its uncertainty with wavelength is studied. 7.7.3 Reagents and Equipment r Spectrometric-grade cyclohexane dried on 4 A ˚ molecular sieves, bisublimed diiodine, iodocyclohexane (98%) passed through a column of freshly activated basic aluminium oxide and copper, further protected from light. r One 100 cm3 , one 50 cm3 and four 10 cm3 volumetric flasks, spatula and pipettes. r Analytical balance. r Any good-quality UV–visible computer-assisted spectrometer. r A 1 cm pathlength Suprasil quartz cell placed in a cell holder thermoregulated at 25 ◦ C. 7.7.4 Experiment r Prepare 100 cm3 of a stock solution of ∼9 × 10−3 mol l−1 diiodine. r Prepare 35 cm3 of a stock solution of ∼0.65 mol l−1 iodocyclohexane.


427

2

Absorbance

1,5

1

0,5

0 400

450

500

550

600

650

Wavelength / nm

Figure 7.10 Visible spectra of free diiodine (red) and mixtures of diiodine and iodocyclohexane in cyclohexane at 25 ◦ C.

r From the diiodine stock solution, prepare a solution S0 of ∼1.75 × 10−3 mol l−1 diiodine. r From the diiodine and iodocyclohexane stock solutions, prepare 10 cm3 of solutions of ∼1.75 × 10−3 mol l−1 diiodine and ∼0.14 (S1 ), ∼0.28 (S2 ), ∼0.42 (S3 ) and 0.∼56 (S4 ) mol l−1 iodocyclohexane. r Record the spectra of solutions S –S against cyclohexane from 350 to 650 nm. 0 4 r Measure the absorbance at 470 nm.

7.7.5

Results and Discussion: Illustration of the HSAB Principle

The spectra are shown in Figure 7.10. The visible band of free diiodine at 520 nm is increasingly blue shifted as the equilibrium is displaced towards the complex. The observed bands are the sum of overlapping bands of free and complexed diiodine. A quasi-isosbestic point is observed because only two absorbing species are present and the sum of their concentrations is quasi-constant Ca + Cc = Ca0 . The calculations are shown in Table 7.10. The results are as follows: εc − εa = 944 ± 37 l mol−1 cm−1 at 470 nm K¯ c = 1.33 ± 0.08 l mol−1 (95% CI) Table 7.11 summarizes the results obtained for the diiodine basicity of haloalkanes (Table 5.21 in Chapter 5) and compares the diiodine basicity with the 4-fluorophenol basicity in cyclohexane. The sequence of diiodine basicity RI RBr > RCl RF lies in the reverse order of the 4-fluorophenol basicity, RF RCl > RBr > RI. Similar

428


Table 7.10 Calculation of the complexation constant of diiodine with iodocyclohexane in cyclohexane at 25 ◦ Ca . Molar mass = 253.81 g mol-1 Molar mass = 210.06 g mol-1 Molar mass = 84.16 g mol-1 dd/dt = 0.00101 Temperature = 16.1 ◦ C

Acid: diiodine Base: iodocyclohexane Solvent: cyclohexane Density (25 ◦ C) = 0.7735 Pathlength = 1 cm Wavelength (nm) = 470.15 nm Stock solution of acid Acid weight = 0.2218 Solution weight = 77.5447

Stock solution of base Base weight = 4.7773 Solution weight = 27.3742

10 ml flasks

S0

S1

S2

S3

S4

Acid solution weight Base solution weight Ca0 (S0 ), Ca0 (S1 to S4 ) Cb0 Absorbance εa α β γ

1.5746 0 0.0 017 541 0 0.5743 327.42

1.5703 1.7124 0.0 017 493 0.14 063 0.8322

1.5753 3.4052 0.0 017 549 0.27 965 1.0229

1.5752 5.1095 0.0 017 548 0.41 962 1.1633

1.5756 6.8340 0.001 753 0.56 125 1.2809

0.2594 −0.1424 0.0009

0.4483 −0.2814 0.0011

0.5888 −0.4214 0.0013

0.7069 −0.563 0.0014

εc − εa S1 S2

Solutions of Equation 7.28 S2 S3 S4 948.9 921.7 947.6 896.2 947.0

S3 Kc S1 S2

S2 1.3197

S3 a

S3 1.3661 1.4283

1003.2 S4 1.3217 1.3232 1.1991

Mean = 944.1 Standard deviation = 35.6 95% CI = 37.4 Mean = 1.33 Standard deviation = 0.075 95% CI = 0.08

Units for weight, d, C, ε and K are g, g cm−3 , mol l−1 , l mol−1 cm−1 and l mol−1 , respectively.

sequences have been observed for the basicity of halide ions towards metallic cations (see Table 1.15): I− > Br− > Cl− F− towards soft Mn+ F− > Cl− > Br− I− towards hard Mn+ All four sequences obey the Hard and Soft Acid and Base (HSAB) principle. Indeed, diiodine is a soft Lewis acid, 4-fluorophenol a hard Lewis acid, and the softness of the electron donor atom increases (that is hardness decreases) in descending group 17 of the periodic table (see Table 1.16 in Chapter 1).


429

Table 7.11 Lewis basicity of haloalkanes towards I2 (pK BI2 ) and 4-FC6 H4 OH (logKc ) in cyclohexane.

Kc (l mol−1 ) pK BI2 Kc (l mol−1 ) LogKc

7.8

7.8.1

C6 H11 I

C6 H11 Br

(CH3 )2 CHCl

C6 H11 F

1.33 0.12 C5 H11 I 0.60 −0.22

0.38 −0.42 C5 H11 Br 0.72 −0.14

0.11 −0.96 C5 H11 Cl 0.78 −0.11

Not measurable C5 H11 F 1.41 0.15

Determination of the Complexation Enthalpy and Entropy of Diiodine with Dimethyl Sulfoxide by Visible Spectrometry Introduction

The best conditions for obtaining accurate diiodine affinities by visible spectrometry from the temperature variation of the complexation constant are studied in this experiment using the example of the complexation of diiodine with dimethyl sulfoxide (DMSO) in CCl4 . The result will be compared with those of two previous studies by Drago et al. [16] and Klæboe [17]. The importance of the solvation term in the diiodine affinity measured in CCl4 will be evaluated by comparison with a value measured in cyclohexane, a less solvating medium. 7.8.2 Reagents and Equipment r Spectrometric-grade CCl4 dried on 4 A ˚ molecular sieves and anhydrous DMSO passed through a column of freshly activated basic aluminium oxide. Diiodine as in experiment 7.7. r Balance, glassware, cell and UV–visible spectrometer as in experiment 7.6. Thermocouple as in experiment 7.4. 7.8.3

Experiment

The equilibrium constant K c is determined by the Rose–Drago method, as described in experiment 7.7, at five temperatures from −5 to +55 ◦ C. At each temperature, the initial diiodine concentration is ∼0.0015 mol l−1 and five initial DMSO concentrations are studied from ∼0.01 to ∼0.09 mol l−1 . The compositions of the five solutions studied at 25 ◦ C are given in Table 7.12. The concentrations at the various temperatures studied are corrected for volume changes. Table 7.12 Initial concentrations (mmol l−1 ) at 25 ◦ C of the five solutions of diiodine (C a0 ) and DMSO (Cb0 ) in CCl4 . Solution Ca0 Cb0

1

2

3

4

5

1.5006 11.2051

1.4868 27.6342

1.4967 44.7080

1.4990 69.2691

1.4920 92.2853

430

Lewis Basicity and Affinity Scales 1.4

1.2

5 4

1

Absorbance

3 0.8 2 0.6 1 0.4

0.2

0 350

400

450

500

550

600

650

Wavelength / nm

Figure 7.11

Visible spectra of the I2 –DMSO solutions S1 –S5 (Table 7.12) in CCl4 at 25.1 ◦ C.

The spectra of the solvent and of the five solutions are recorded from 350 to 650 nm at each temperature (measured inside the solvent cell) against air. Then the spectrum of the solvent is subtracted from that of the solution. The absorbance can be read at any wavelength. However, it is better to calculate the diiodine concentrations at the wavelength where the A – A0 term of Equation 7.28 is the largest. This extremum is easily located by computing the curve A – A0 versus λ. 7.8.4

Results

Figure 7.11 shows the visible spectra of the five I2 –DMSO solutions in CCl4 at 25.1 ◦ C. When the DMSO concentration increases, it appears that the absorbance of the free diiodine band at 516.7 nm decreases, to the benefit of the band of complexed diiodine near 442 nm. Figure 7.12 shows the visible spectra of solution 2 at five temperatures. Now the absorbance decrease of the 516.7 nm band and the absorbance increase of the 442 nm band are caused by a temperature decrease. Figures 7.1 and 7.12 illustrate the shift of the equilibrium 7.38: (CH3 )2 SO + I2 (CH3 )2 SO · · · I−I

(7.38)

towards the formation of the complex when the DMSO concentration increases or the temperature decreases. The 1 : 1 stoichiometry of the complex has been determined by Klæboe [17] by the method of continuous variation of the I2 /DMSO system in CCl4 at 20 ◦ C.


431

1.4

1.2

1

Absorbance

-4.6 0.8

9.7

0.6

25.1

0.4

39.7 53.1

0.2

0 350

400

450

500

550

600

650

Wavelength / nm

Figure 7.12 Visible spectra of the I2 –DMSO solution S2 in CCl4 at the five indicated temperatures (◦ C).

On the A – A0 versus λ curve (not shown), the wavelength at which the absorbance variations upon the shift of equilibrium 7.38 are the greatest is located at 442 nm. A second extremum is found near 522 nm. The equilibrium constant values calculated by the Rose–Drago method, as explained in experiment 7.7, are summarized in Table 7.13. A least-squares treatment of lnK c versus 1/T yields Equation 7.39: ln K c = 2403.40 (±22.65)

1 − 5.56 (±0.08) T

(7.39)

n = 5, r2 = 0.9997, s = 0.012 Table 7.13 Complexation constants Kc (l mol−1 ) and absorption coefficients (ε c − εa ) (l mol−1 cm−1 ) for the I2 –DMSO complex in CCl4 at 442 nm. t/◦ C −4.6 9.7 25.1 39.4 53.9

Kc

εc − εa

29.72 ± 1.47 18.75 ± 1.04 12.14 ± 0.92 8.26 ± 1.21 6.04 ± 1.36

1278 ± 33 1270 ± 47 1271 ± 78 1301 ± 194 1326 ± 401

432


corresponding to K c (25 ◦ C) = 12.16 l mol−1 . The standard thermodynamic functions are calculated as follows: ∆H ◦ = −R [d lnK c /d (1/T )] − α RT 2 ∆H ◦ = −20.9 (±0.2) kJ mol−1 ∆G ◦298 = −RT lnK c (25 ◦ C) ∆G ◦298 = −6.19 kJ mol−1 ◦ ∆S298 = −5.56 (±0.08) × 8.3145 ◦ ∆S298 = −46.2 (±0.7) J K−1 mol−1

(7.40)

(7.41)

(7.42)

Gibbs energies and entropies refer to molar concentrations.

7.8.5 7.8.5.1

Discussion Effect of Wavelength on the Complexation Enthalpy Measured

The calculation of ∆H ◦ at 522 nm, the wavelength corresponding to the second extremum of the A – A0 versus λ curve, and at 550 nm, the wavelength selected by Drago et al. [16], yields ∆H ◦ = −21.2 (±1.3) kJ mol−1 (at 522 nm) ∆H ◦ = −20.8 (±0.2) kJ mol−1 (at 550 nm) The enthalpies calculated at 522 and 550 nm are in agreement with the value obtained at 442 nm. This does not mean that the calculated enthalpy does not vary with the wavelength but that concordant results can be obtained if the wavelength of measurement is judiciously chosen. In fact, there is a blue shift in the wavelength of the band of the complex (with a concomitant loss of the isosbestic point near 485 nm) when the DMSO concentration increases and deviations in ∆H ◦ should occur if the calculations are carried out on the side of a shifting band. Shifting of the band of the complex has a negligible effect on the A – A0 values at 442 and 522 nm because the A – A0 versus λ curve is very flat near these extrema. An extensive study of the wavelength variation of diiodine complexation constants can be found in ref. [18].

7.8.5.2

Effect of DMSO Concentration on the Complexation Enthalpy Measured

Table 7.13 shows that there is no significant change in the absorption coefficient term εc − εa with temperature differences in the range −5 to 54 ◦ C at 442 nm. This allows the calculation of K c from Equation 7.28 at any temperature from a single mixture of I2 and DMSO in CCl4 . Consequently, the calculation of the complexation enthalpy can be carried out for each DMSO concentration, and the possible variation of ∆H ◦ with DMSO concentration can be studied in the range 0.01–0.09 mol l−1 . The results are summarized in Table 7.14.


433

Table 7.14 Calculation of the complexation enthalpy for a given DMSO concentration assuming the temperature independence of εc − εa from –5 to +54◦ C at 442 nm.a DMSO concentration/mol l−1 −∆H◦ /kJ mol−1 a

0.011

0.028

0.044

0.069

0.092

21.09 ± 0.10

21.00 ± 0.15

21.35 ± 0.12

21.37 ± 0.16

21.41 ± 0.16

The value at 25 ◦ C, 1271 l mol−1 cm−1 , is chosen for the calculation.

The results in Table 7.14 show that the complexation enthalpy can be assumed to be independent of DMSO concentration in the studied range. This indicates that it is not necessary to make any activity coefficient correction in the thermodynamic study of the system I2 /DMSO in CCl4 , at the dilute concentrations employed. 7.8.5.3

Comparison with Literature Results

The least-squares treatment of the K c data of Drago et al. [16], determined at 550 nm and 23.4, 28.9, 35.4, 40.6 and 46.1 ◦ C, yields ln K c = 2193.09 (±41.71)

1 − 4.89(±0.14) T

(7.43)

n = 5, r2 = 0.9989, s = 0.008 −∆H ◦ = 19.1 (±0.3) kJ mol−1 The data of Klæboe [17], determined at 440 nm and 15, 25, 35 and 45 ◦ C, yield ln K c = 1920.03 (±81.20)

1 − 4.03 (±0.27) T

(7.44)

n = 4, r2 = 0.996, s = 0.020 −∆H ◦ = 16.9 (±0.7) kJ mol−1 We believe that the results of our experiment are to be preferred to these literature results for the simple reason that our measurements use a much wider temperature range (58.5 ◦ C) than those of Drago et al. [16] (23 ◦ C) and Klæboe [17] (30 ◦ C). 7.8.5.4

Solvent Effect on the Complexation Enthalpy

Although cyclohexane dissolves much smaller amounts of DMSO than CCl4 does, it is interesting to carry out the measurement of the I2 –DMSO system in this more ‘inert’ solvent, to evaluate the solvent effect on the complexation enthalpy. The reader will check the results given in Table 7.15, obtained under the following conditions. The absorbance is read at 442 nm, the temperature is varied from 10 to 59 ◦ C, the initial diiodine concentration is ∼0.0015 mol l−1 and the initial DMSO concentration

434

Lewis Basicity and Affinity Scales Table 7.15 Complexation constants Kc (l mol−1 ) and absorption coefficients (εc − εa ) (l mol−1 cm−1 ) for the I2 –DMSO complex in cyclohexane at 442 nm. t/◦ C

Kc

εc − εa

9.7 25.1 39.3 58.7

79.60 ± 5.66 47.20 ± 2.76 28.40 ± 3.08 17.50 ± 2.21

1189 ± 45 1135 ± 41 1117 ± 96 1002 ± 116

ranges from ∼0.009 to ∼0.017 mol l−1 . These data yield 1 − 6.00 (±0.31) T n = 4, r2 = 0.9979, s = 0.037 ln K c = 2935.28 (±94.97)

(7.45)

∆H ◦ = −25.3 (±0.8) kJ mol−1 K c (25 ◦ C) = 46.54 l mol−1 ∆G ◦298 = −9.52 kJ mol−1 ◦ ∆S298 = −49.9( ± 2.6) J K−1 mol−1

A significant difference of 4.4 ± 1.0 kJ mol−1 (that is, a 21% relative difference) is found between the enthalpies measured in CCl4 and in cyclohexane. It appears that the diiodine affinities measured in alkanes must not be mixed with those measured in CCl4 , as was unfortunately done in the Drago EC analysis of Lewis affinity. Moreover, the experimental value to be compared with the diiodine affinity of DMSO computed in vacuo by quantum chemical methods is the value measured in cyclohexane. Taking into account the solvent effect of cyclohexane, a good calculated value should be at least −25.3 (±0.8) kJ mol−1 .

7.9

7.9.1

FTIR Determination of the Shift of the I C Stretching of Iodine Cyanide upon Halogen Bonding to Phosphine Chalcogenides Introduction

The linear triatomic molecule ICN has three fundamental frequencies, ν 1 (I C stretching), ν 2 (doubly degenerate I CN bending), and ν 3 (CN stretching), at 486, 320 and 2168 cm−1 , respectively, in chloroform solution. The decrease in the I C stretching frequency as this molecule forms halogen-bonded complexes with Lewis bases can be used to construct a spectroscopic scale of halogen-bond affinity (see Chapter 5). In the HSAB classification, ICN is a soft Lewis acid. However, thermodynamically, ICN is significantly harder than diiodine, as expected by the ICN and I2 dipole moments of 3.74 and 0 D, respectively. Nevertheless, the frequency shift ∆ν(I CN) appears very sensitive to the softness of the Lewis base (see Chapter 5), possibly because of the significant contribution of the n(π ) → σ ∗ (C I) charge transfer to the decrease in the force constant of the C I bond. Consequently, the ∆ν(I CN) scale has been proposed as a spectroscopic scale of soft affinity [19].


435

In this experiment, the soft character of the ∆ν(I CN) scale is studied by measuring this scale for three phosphine chalcogenides, (C6 H5 )3 PO, (C6 H5 )3 PS and (C6 H5 )3 PSe. In this series of Lewis bases, the prominent structural feature is the increase in softness of the electron-donor atom, in the order O < S < Se. Dichloromethane is chosen as the solvent because it dissolves ICN better than alkanes and CCl4 and its transparency allows the use of a cell of 1 mm pathlength. While it is true that CH2 Cl2 interacts weakly with ICN and Lewis bases, it is satisfactory that the ν(I C) frequencies in CH2 Cl2 (485 cm−1 ) and in cyclohexane (486.5 cm−1 ) are almost equal. 7.9.2 Reagents and Equipment r Spectroscopic-grade dichloromethane, dried on 4 A ˚ molecular sieves. r Iodine cyanide recrystallized from CHCl . 3 r Triphenylphosphine oxide, triphenylphosphine sulfide and triphenylphosphine selenide recrystallized from light petroleum, benzene and ethanol, respectively. r Volumetric flasks, spatula, pipettes. r Analytical balance. r FTIR spectrometer flushed with dry air operating at a resolution of 1 cm−1 . r A 1 mm pathlength KBr cell placed in a cell holder thermoregulated at 15 ◦ C (the temperature is lowered in order to increase the complexation constants and the concentration of complexes). 7.9.3 Experiment r Prepare a stock solution of 0.2 mol l−1 ICN. r Prepare three stock solutions of bases: 0.3 mol l−1 (C6 H5 )3 PO, 0.2 mol l−1 (C6 H5 )3 PS and 0.15 mol l−1 (C6 H5 )3 PSe. r In 10 cm3 volumetric flasks, prepare three corresponding sample solutions by adding 5 cm3 of stock solution of base to 5 cm3 of stock solution of ICN. r In 10 cm3 volumetric flasks, prepare three reference solutions by adding 5 cm3 of CH Cl 2 2 to 5 cm3 of each of the stock solutions of bases (dilution by 1 : 2). r Record the three reference spectra, then the three sample spectra. r Display the absorbance spectra of the solute (an equilibrium mixture of ‘free’ ICN and ICN halogen-bonded to phosphine chalcogenide). 7.9.4

Results and Discussion: ∆ν(ICN) as a Spectroscopic Scale of Halogen-Bond Affinity

IR spectra of the three halogen-bonded complexes of iodine cyanide with Ph3 PO, Ph3 PS and Ph3 PSe are shown in Figure 7.13. They show two bands in the range 500–360 cm−1 . The first, at 485 cm−1 , is sharp with a shoulder at low wavenumbers (attributed to a hot band); it is assigned to ν(I C) of ‘free’ ICN. The second one is broader and situated at 455, 433 or 415 cm−1 ; it is attributed to ν(I C) of ICN halogen bonded to (C6 H5 )3 PO, (C6 H5 )3 PS or (C6 H5 )3 PSe, respectively. In Table 7.16, the spectroscopic results are compared with the diiodine affinity for the three phosphine chalcogenides. There is good agreement between the ICN spectroscopic scale and the diiodine affinity scale. Both follow the softness order Ph3 PO < Ph3 PS < Ph3 PSe. Moreover, since diiodine is the archetype of soft Lewis acids,

436


1

b

c

d

Absorbance

a

0 500

390 -1

Wavenumber / cm

Figure 7.13 FTIR spectra, in the ν(I CN) region, of free ICN (a) and ICN complexed to Ph3 PO (b), Ph3 PS (c) and Ph3 PSe (d) in CH2 Cl2 at 15 ◦ C in a 1 mm KBr cell.

the ∆ν(I CN)–diiodine affinity relationship confirms the soft character of the ∆ν(I CN) spectroscopic scale of halogen-bond affinity.

7.10

7.10.1

Blue Shift of the Visible Diiodine Transition Upon Halogen Bonding to Pyridines Introduction

Among the uses of the correlations that have been established between spectroscopic and thermodynamic scales of basicity, the calculation of new thermodynamic values appears of primary importance in the context of the quantitative measurement of Lewis basicity. In fact, it is generally easier and faster to measure a spectroscopic shift than an equilibrium constant. This experiment illustrates the calculation of pK BI2 values for three pyridines, Table 7.16 Changes in the I C stretch (cm−1 ) of iodine cyanide on complexation with phosphine chalcogenides. Comparison with diiodine affinity (kJ mol −1 ). Phosphine chalcogenide Ph3 PO Ph3 PS Ph3 PSe a

From Table 5.25.

ν(I CN) free

ν(I CN) complex

∆ν(I CN)

I2 affinitya

485.1 485.1 485.1

455.2 433.5 414.9

29.9 51.6 70.2

18.8 30.5 38.5


437

3-iodopyridine, 3-(trifluoromethyl)pyridine and 4-(trifluoromethyl)pyridine, from the correlation between the diiodine basicity and the blue shift of the visible diiodine transition, established in Chapter 5 for 3- and 4-substituted pyridines (Equation 7.46): pK BI2 = 12.78 (±0.22) ×10−4 ∆ν(πg → σu ) −3.56 (±0.10)

(7.46)

n = 17, r = 0.998, s = 0.07. The standard error of the estimate, s, is very close to the experimental uncertainty on pK BI2 (about ±0.05). Thus, the diiodine basicity of any 3- and 4-substituted pyridine can be reliably obtained from the determination of a spectroscopic shift. 7.10.2 Reagents and Equipment r Spectroscopic-grade heptane. r 3-Iodopyridine, recrystallized from light petroleum, 3-(trifluoromethyl)pyridine and 4(trifluoromethyl)pyridine passed through a column of freshly activated basic aluminium oxide. r Same glassware, UV–visible spectrometer and quartz cell as in experiment 7.7. The cell holder is thermoregulated at 15 ◦ C. 7.10.3

Experiment

To obtain a band of complexed diiodine of absorbance 0.2 in a 1 cm cell, assuming an absorption coefficient of 1000 l mol-1 cm−1 , a concentration of complex of about 2 × 10−4 mol l−1 is required. With an initial concentration of 8 × 10−4 mol l−1 in diiodine (yielding an absorbance of 0.76 for the band of free diiodine at 521.4 nm), the initial concentration of base, Cb0 , may be adjusted from the expression for the equilibrium constant K c : 2 × 10−4 8 × 10−4 − 2 × 10−4 Cb0 − 2 × 10−4

Kc =

(7.47)

The complex concentration can be neglected compared with the base concentration in excess. The relationship between K c (in l mol−1 ) and Cb0 becomes Cb0 ≈ (3K c )−1

(7.48)

From the K c values given in Table 5.8 in Chapter 5, a K c value is assumed for the studied pyridine in order to calculate Cb0 . For example, the K c value of 3-iodopyridine is assumed close to the value of 32 l mol−1 of 3-fluoropyridine. Thus, the initial concentration of 3-iodopyridine is chosen to be about 10−2 mol l−1 . Hence the preparation of solutions is as follows for 3-iodopyridine: r Prepare a stock solution of about 10−2 mol l−1 3-iodopyridine. r Prepare a stock solution of 20 mg of I in 10 cm3 of heptane (stir for 12 hours for complete 2 dissolution). r Prepare a solution 0 by adding 1 cm3 of diiodine stock solution to a 10 cm3 flask and filling to the mark with heptane. r Prepare a solution 1 by adding 1 cm3 of heptane to a 10 cm3 flask and filling to the mark with the 3-iodopyridine stock solution.

438

Lewis Basicity and Affinity Scales 1 a

Absorbance

0,8

0,6 b 0,4

0,2

0 350

400

450

500

550

600

650

700

Wavelength / nm

Figure 7.14 Visible spectrum of a mixture of diiodine and 3-iodopyridine in heptane showing the free diiodine visible band (a) at 521.4 nm and the complexed diiodine band (b) blue shifted to 431.4 nm.

r Prepare a solution 2 by adding 1 cm3 of diiodine stock solution to a 10 cm3 flask and filling to the mark with the 3-iodopyridine stock solution. The spectra S0 of solution 0, S1 of solution 1 and S2 of solution 2 are recorded in the wavelength range from 700 to 350 nm. Then the spectrum (S2 – S1 ) – f S0 is computed. The factor f is iterated until the free diiodine band at 521.4 nm disappears. The resulting spectrum corresponds to the visible band of complexed diiodine. Repeat the experiment with the two other pyridines, with stock solutions of about 0.02 and 0.025 mol l−1 4-(trifluoromethyl)pyridine and 3-(trifluoromethyl)pyridine, respectively. 7.10.4

Results and Discussion: Substituent Effects

A representative spectrum is shown in Figure 7.14 and the results of the blue shift and of pK BI2 calculated from Equation 7.46 are reported in Table 7.17. Table 7.17 Changes in the π g → σ u diiodine transition wavenumber (cm−1 ) on complexation with pyridines in heptane. Prediction of pK B I 2 from Equation (7.46). Pyridine

ν free I2

ν complex

∆ν(π g → σ u )

pK BI2

3-IC5 H4 N 4-CF3 C5 H4 N 3-CF3 C5 H4 N

19 180 19 180 19 180

23 180 22 960 22 850

4000 3780 3670

1.55 1.27 1.13

The Measurement of Lewis Basicity and Affinity in the Laboratory 2,4

439

H

pKBI2

1,9 1,4

COOMe COMe CF3

0,9 0,4 -0,1

CN NO2

0,1

0,3

0,5

0,7

σF

Figure 7.15 pK B I 2 of 3-substituted pyridines (data from Tables 5.8 and 7.17) versus the field-inductive substituent constant σ F [20].

The most important factor limiting the accuracy of ∆ν(π g → σ u ) is the influence of the medium. Even at dilute concentrations of pyridines in heptane, there is a clustering of polar pyridine molecules around the polar complexes, causing a shift of the band of the complex to short wavelengths. The concentration dependence of ∆ν(π g → σ u ) can be minimized by using as low a pyridine concentration as allowed by an accurate measurement of the visible band maximum. The calculated pK BI2 values depend on the substituents as expected from the electronic theories. The 3-iodo substituent behaves similarly to other 3-halo substituents: 3-F (1.51), 3-Br (1.40) and 3-Cl (1.38). The CF3 substituent decreases the basicity through the same electron-withdrawing inductive effect as the COOMe, COMe, CN and NO2 substituents. The pK BI2 values follow the order of the σ F field-inductive substituent constant as shown by Equations 7.49 and 7.50 and as illustrated in Figure 7.15. pK BI2 = −2.21 (±0.13) σF + 2.27 (±0.06)

(7.49)

n = 5 para substituents (H, COMe, CF3 , CN, NO2 ), r = 0.995, s = 0.07 pK BI2 = −2.27 (±0.19) σF + 2.17 (±0.04)

(7.50)

n = 6 meta substituents (H, COOMe, COMe, CF3 , CN, NO2 ), r = 0.997, s = 0.05. From a σ F database [20], many new pK BI2 values can now be estimated by means of Equations 7.49 and 7.50.

7.11

7.11.1

Mass Spectrometric Determination of the Gas-Phase Lithium Cation Basicity of Dimethyl Sulfoxide and Methyl Phenyl Sulfoxide by the Kinetic Method The Kinetic Method

As explained in Section 6.3, Fourier transform ion cyclotron resonance mass spectrometry (FTICR-MS) and high-pressure mass spectrometry (HPMS) are the methods of choice for

440


the determination of equilibrium constants of ligand exchange about a given cation. Unfortunately, the application of these mass spectrometric techniques is hampered in most laboratories by the unavailability of such complex, expensive and specialized instrumentation. An experiment using a more generally available mass spectrometric method is chosen instead. A mass spectrometer offering the possibility of collision-induced dissociation (CID), also called MS/MS or tandem MS, will allow measurements of relative gas-phase affinity/basicity, provided that a ‘cation-bound dimer’ [L1 M+ L2 ] (L = ligand or base; M = H or metal) may be formed in the ion source. The CID technique of ion fragmentation is widespread in mass spectrometry laboratories, with the aim of obtaining structural information as well as a more selective detection and quantitation of analytes. The basic principles of the kinetic method are briefly recalled here. In the first step, the cation-bound dimer [L1 M+ L2 ] is selected from the various ions formed in the source. Then, this ion is accelerated to a given kinetic energy and a part of the added kinetic energy is transformed into internal energy, owing to the collisions with gas atoms (He, Ar, N2 , etc.). Some fragmentation of the excited ion occurs depending on the bond strength and the amount of internal energy acquired from the collision. Under well-controlled conditions, fragmentation of this ion into [L1 M+ ] and [L2 M+ ], treated by means of the so-called ‘kinetic method’ (or Cooks method) [21, 22], leads to the relative basicity, as described in Chapter 6. In the case of M+ = Li+ (or other alkali metal cations), using electrospray ionization (ESI) is the easiest approach to obtain a range of Li+ adducts. The basic assumptions of the kinetic method are better fulfilled when measuring compounds of similar structures and functions. The mass spectrometer may be a triple-quadrupole or an ion-trap instrument. Other kinds of mass spectrometers [23, 24] such as linear ion traps, electric/magnetic sector instruments, or more sophisticated hybrid instruments, for example a quadrupole/time-of-flight (QTOF) apparatus, are also used for CID experiments. Note that, for some instruments, the low m/z (massto-charge ratio) limit may be a limitation for the choice of the molecule to be measured. The proposed experiment consists in determining the LiCB of dimethyl sulfoxide (DMSO) and methyl phenyl sulfoxide (MPSO) from the measurement of the intensity ratios of the ions [L1 Li+ ] and [L2 Li+ ], I[L1 Li+ ]/I[L2 Li+ ], involving five ligands Li . Among them, three amides [N-methylformamide (MFA), N,N-dimethylformamide (DMF) and N,N-dimethylacetamide (DMA)] are considered as references with known LiCBs. The unknown LiCBs of the two sulfoxides are determined from a plot of ln(I[L1 Li+ ]/I[L2 Li+ ]) against known thermodynamic LiCB values. As explained in Section 6.3, the ratio of ion intensities is equal to the ratio of the unimolecular rate constants of the dissociation reactions of [L1 M+ L2 ] into [L1 Li+ ] and [L2 Li+ ]. Hence ln(I[L1 Li+ ]/I[L2 Li+ ]) may be called a relative lithium cation ‘kinetic basicity’. Remember that, under specified assumptions, (7.51) ln I L1 − Li+ /I L2 − Li+ ≈ ∆LiCA/RTeff ≈ ∆LiCB/RTeff where T eff is the effective temperature (see Chapter 6). 7.11.2 Reagents and Equipment r Lithium bromide, MFA, DMF, DMA, DMSO, MPSO (purity is not critical) and methanol (LC/MS grade). r Standard volumetric glassware and analytical balance.


441

r Mass spectrometer with MS/MS (CID) possibilities and a source for generating the Li+ clusters. The experiments described here were carried out on a Thermo Finnigan Model TSQ 7000 triple-quadrupole mass spectrometer, using argon as collision gas. 7.11.3

Experiment

Stock solutions of lithium bromide and of the five Lewis bases are prepared in the concentration range 0.05–0.1 g ml−1 in methanol. The working ESI solution is obtained by dilution of 2 to 8 µl of LiBr stock solution and two competing ligand stock solutions in 1 ml of methanol, to obtain respective concentrations in the range (1–5) × 10−3 mol l−1 . The final concentration is not critical but may be adjusted according to the instrument and the ESI source conditions so as to obtain an optimum intensity of the cation bound dimer [L1 M+ L2 ]. In the experiment described, six solutions containing the competing ligands are prepared: DMF with DMA; DMSO with MFA, DMF and DMA; and MPSO with DMF and DMA. More combinations may be tried since there is a maximum of 10 possible combinations, but a minimum of four is necessary to link the five ligands (Figure 7.16). The instrument is set up according to the standard MS/MS conditions, taking into account the aim of the experiment, which is dissociating the lithium cation-bound dimers. The most important parameters are flow rate of solution, spray voltage and temperature of the desolvation capillary. As far as possible, the same experimental conditions should be used for all the experiments, in particular the pressure of the collision gas (argon) and the range of collision energies. The ESI solution containing two ligands and the lithium salt is introduced into the source using the ‘infusion’ mode (continuous injection using a syringe pump).

L5 L4 L3

L2

L1

Figure 7.16 Schematic representation of the possible links between relative basicities for five ligands in terms of ln(I[L1 Li+ ]/I[L2 Li+ ]). The minimum number of links (or values) necessary for establishing the scale is four, either as direct links (L1 ↔ L2 ↔ L3 ↔ L4 ↔ L5 ) or indirect (for example L1 ↔ L2 ↔ L3 , L2 ↔ L4 , L2 ↔ L5 ). When more than four links are established (maximum = 10), some averaging can be made and the self-consistency of the scale may be evaluated.

442


Table 7.18 Results of the kinetic method obtained for the six pairs of ligands. m/z (L1 Li+ )

L1 DMA DMSO DMSO DMSO DMF DMF

94 85 85 85 80 80

L2 MPSO DMA DMF MFA DMA MPSO

m/z (L2 Li+ )

ln(I[L1 Li+ ]/I[L2 Li+ ]

147 94 80 66 94 147

−0.152 −1.350 0.646 2.797 −2.036 −2.158

For the kinetic method, typical ranges for working conditions of the TSQ 7000 instrument [25], with the solutions prepared above, are as follows: r r r r r

Syringe flow: 2–5 µl min−1 . ESI spray voltage: 4–5 kV. Heated capillary temperature: 150–200 ◦ C. Scan range: m/z 50–250. Pressure of the collision gas (argon): 133–200 Pa [(1–1.5) × 10−3 Torr]. The pressure is not very critical, but it is convenient to keep it constant for the complete set of measurements. r Collision energy: several values between 4 and 12 eV, by 1–2 eV steps. In the present study, the MS/MS experiment on [L1 Li+ L2 ] is conducted at seven collision energies in this range. The results are not very sensitive to this parameter within this energy range and can be averaged.

The main fragment ions should be [L1 Li+ ] and [L2 Li+ ], because the weakest bonds are actually the L Li+ bonds. Other fragments may be observed, mainly at the higher collision energies, resulting from other fragmentation pathways.

7.11.4

Data Treatment

The relative intensities I[L1 Li+ ]/I[L2 Li+ ] are measured and mean values of ln(I[L1 Li+ ]/I[L2 Li+ ]) (relative kinetic basicities) are calculated for the ligand pairs given above. Table 7.18 shows the results. A positive value of the logarithm implies that L1 is a stronger base than L2 (and the reverse for a negative logarithm). The uncertainty on each value is in the range ±0.1–0.2. Based on Figure 7.16 and Table 7.18, a kinetic Li+ basicity scale, limited to ligands of close basicity, is established relative to N-methylformamide in Figure 7.17 by adding pair values. The kinetic basicity of amides, reported in Table 7.19, is related to their thermodynamic basicity by means of Equation 7.52: ln(I [LLi+ ] / I [MFALi+ ]) = 0.313(±0.027) LiCB(373 K) − 51.99(±4.67) n = 3, r 2 = 0.993, s = 0.25 kJ mol−1 .

(7.52)


Ligand

443

ln (I [LLi+] / I [MFALi+])

ln(I[L1Li+] / I [L2Li+])

4.304

MPSO 0.152

DMA 1.350

4.167

2.158 2.036 2.797

DMSO 0.646

2.151

DMF 2.797

MFA

(0) +

Figure 7.17 (Not to exact scale). Relative kinetic Li basicity ladder built using data in Table 7.18. The final ln(I[LLi+ ]/I[MFALi+ ]) scale for the five ligands is given relative to MFA.

The LiCBs of sulfoxides are calculated by Equation 7.53: LiCB(373 K) = 3.1729(±0.273) ln(I [LLi+ ] /I [MFALi+ ]) + 166.22(±0.74)

(7.53)

n = 3, r 2 = 0.993, s = 0.8 kJ mol−1 . The results are LiCB(DMSO) = 175.1 and LiCB(MPSO) = 179.9 kJ mol–1 , in good agreement with the literature data. By assuming a constant entropy of Li+ complexation (equivalent to a negligible entropy of ligand exchange), LiCA values can also be calculated. The kinetic data are then converted to thermodynamic values from the theoretical LiCAs at 298 K of Table 7.19 and Equation 7.54: LiCA(298 K) = 4.8475(±0.0075) ln(I [LLi+ ] /I [MFALi+ ]) + 211.89(±0.02)

(7.54)

n = 3, r 2 = 0.9999, s = 0.02 kJ mol−1 . Table 7.19 Kinetic Li+ basicity and the corresponding experimental LiCB and ab initio LiCA (kJ mol–1 ) of the three reference amides. Ligand

Ln(I[LLi+ ]/I[MFALi+ ])

LiCB (373 K)a

LiCA (298 K)b

MFA DMF DMA

0.000 2.151 4.167

165.9 173.7 179.1

211.9 222.3 232.1

a

From Table 6.4. From ref. [26] Calculations at the G3 (GCP) level, with full-core electron correlation and inclusion of geometric effects in BSSE correction. b

444


Table 7.20 Influence of a methyl/phenyl substitution on neutral and cationic basicity and affinity scales (kJ mol−1 ), and relevant substituent constants. Ligand Scale BF3 affinityc 4-FC6 H4 OH basicityd I2 basicitye Li+ basicity (LiCB)f Li+ affinity (LiCA)f

CH3 SOCH3

C6 H5 SOCH3

C6 H5 SOC6 H5

∆1 a

∆2 b

105.34 14.50 8.88 175.1 225.4

97.37 12.79 7.01 179.9 232.8

90.34 11.64 6.29 183.3 —

−7.97 −1.71 −1.87 +4.8 +7.4

−7.03 −1.15 −0.72 +3.4

Substituent of the sulfoxide function

Field effect (σ F )g Polarizability effect (σ α )g

Methyl

Phenyl

0 −0.35

+0.10 −0.81

a

Variation in basicity (affinity) upon a first Me/Ph substitution. Variation in basicity (affinity) upon a second Me/Ph substitution. c Chapter 3. d Chapter 4. e Chapter 5. f This experiment and Table 6.9. g Ref. [20]. b

The results are LiCA(DMSO) = 225.4 and LiCA(MPSO) = 232.8 kJ mol–1 . The conversion Equations 7.53 and 7.54 may be used for converting additional Li+ relative kinetic basicity, provided that the expected LiCBs and LiCAs are within (or very close to) the LiCB (or LiCA) calibration range. 7.11.5

Discussion: Substituent Effects

In the LiCB (LiCA) scale of sulfoxides, a Me/Ph substitution (that is, the change from CH3 SOCH3 to C6 H5 SOCH3 ) yields an increase in basicity (affinity) of +4.8 (+7.4) kJ mol–1 . In contrast, the same substitution in the sulfoxide basicity (affinity) scales constructed towards neutral Lewis acids as references (BF3 , 4-FC6 H4 OH, I2 ) gives a decrease in basicity (affinity), as shown in Table 7.20. The decrease in basicity (affinity) can be explained by the electron-withdrawing field/inductive effect of Ph compared with Me, as measured by the corresponding substituent constants [20] reported in Table 7.20, whereas the LiCB (LiCA) increase can be explained by the greater polarizability effect [20] of the phenyl group as compared with methyl. The increased influence of the polarizability effect on LiCB is the consequence of the charge on lithium, which creates an increased induced dipole in the ligand and raises the polarization component of the binding energy. As expected, a second Me/Ph substitution, that is, the change from C6 H5 SOCH3 to C6 H5 SOC6 H5 , confirms the trends observed with the first substitution.


445

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Index References to figures are given in italic type. References to tables are given in bold type. acetamidines 137 acetic acid 2 acetone 80 Brönsted basicity 5 ECM parameters 49 halogen-bond spectroscopic basicities 298 acetonitrile 4-fluorophenol hydrogen-bond enthalpy 173 antimony pentachloride complexation 72 diiodine basicity 255 potassium cation affinity and basicity 361 acetophenone Brönsted basicity 5 diiodine complexation constants 241 4-fluorophenol hydrogen-bond enthalpies and entropies 179 acetophenones boron trifluoride affinity 94 copper monocation basicity and affinity 374 4-fluorophenol hydrogen-bond basicity 147–148 methanol IR hydrogen-bonding shift 199 acid halides 258 acidic sites 19 activity equilibrium constant 22 affinity see Lewis affinity alcohols 111 aluminium cation basicity 364 antimony pentachloride affinity 75 copper monocation affinity and basicity 373, 376 cyclopentadienylnickel basicity 370 4-fluorophenol hydrogen-bond basicity 142 lithium cation affinity and basicity 343, 348–349 manganese cation basicity 368 methylammonium cation affinity and basicity 377 sodium cation affinity and basicity 356, 359 Lewis Basicity and Affinity Scales: Data and Measurement C 2010 John Wiley & Sons, Ltd

solvatochromic comparison 216 aldehydes aluminium cation basicity 364 boron trifluoride affinity 93 diiodine basicity 258 4-fluorophenol hydrogen-bond basicity 146 4-fluorophenol hydrogen-bond enthalpies and entropies 179 lithium cation affinity and basicity 350 manganese cation basicity 368 methanol IR hydrogen-bonding shift 198 aliphatic ethers see ethers, aliphatic alkanones see ketones alkenes diiodine complexation enthalpy 295 4-fluorophenol hydrogen-bond basicity 122 4-fluorophenol hydrogen-bond enthalpies and entropies 175 halogen-bond basicity 295 halogen-bond spectroscopic basicities 295 methanol IR hydrogen-bonding shift 190 solvatochromic comparison 214 alkylbenzenes diiodine basicity 244 diiodine complexation enthalpy 287, 295 halogen-bond basicity 295 see also carbon bases alkynes diiodine complexation enthalpy 295 4-fluorophenol hydrogen-bond basicity 122 4-fluorophenol hydrogen-bond enthalpies and entropies 175 halogen-bond basicity 295 methanol IR hydrogen-bonding shift 190 aluminium 46 cation properties 379 gas-phase basicity scale 354, 364–365 Christian Laurence and Jean-François Gal

448

Index

amides 4-fluorophenol hydrogen-bond affinity 180 aluminium cation basicity 364–365 boron trifluoride affinity 97–98 copper monocation basicity and affinity 374 diiodine basicity 260, 261 4-fluorophenol hydrogen-bond basicity 153–154 lithium cation affinity and basicity 351 methylammonium cation affinity and basicity 378 primary, hydrogen-bond affinity 111 secondary, hydrogen-bond affinity 111 solvatochromic comparison 217–218 amidines boron trifluoride affinity 92 4-fluorophenol hydrogen-bond basicity 137–138 methanol IR hydrogen-bonding shift 195–196 amine oxides antimony pentachloride affinity 77 diiodine basicity 265 4-fluorophenol hydrogen-bond basicity 159 4-fluorophenol hydrogen-bond enthalpies and entropies 180 methanol IR hydrogen-bonding shift 203 amines 4-fluorophenol basicity, solvent effects 34 antimony pentachloride affinity 75 boron trifluoride affinity 91 complexation constants 40 copper monocation basicity and affinity 372 cyclopentadienylnickel basicity 370 diiodine complexation enthalpy 287–288, 295–296 halogen-bond basicity 295–296 halogen-bond complexation, bond geometry 235 lithium affinity and basicity 341–342, 347 manganese cation basicity 368 methylammonium affinity and basicity 377–378 potassium cation affinity and basicity 361 primary diiodine basicity 247 4-fluorophenol hydrogen-bond basicity 124 4-fluorophenol hydrogen-bond enthalpies and entropies 175 hydrogen-bond affinity 111 methanol IR hydrogen-bonding shift 190–191 secondary diiodine basicity 247–248

4-fluorophenol hydrogen-bond basicity 125–126 4-fluorophenol hydrogen-bond enthalpies and entropies 175–176 hydrogen-bond affinity 111 methanol IR hydrogen-bonding shift 191 sodium cation affinity and basicity 355, 359 solvatochromic comparison 214 tertiary diiodine basicity 248–249 4-fluorophenol hydrogen-bond basicity 127–128 4-fluorophenol hydrogen-bond enthalpies and entropies 176 methanol IR hydrogen-bonding shift 192 see also nitrogen bases amino acids cationic adduct structure 333 lithium cation affinity and basicity 344 potassium cation affinity and basicity 362 sodium cation affinity and basicity 357, 357–358 aminoacetonitrile, nickel adduct 329–330, 330 aminoboranes, diiodine basicity 249 ammonia 3 4-fluorophenol hydrogen-bond basicity 124 4-fluorophenol hydrogen-bond enthalpies and entropies 175 anharmonicity 14 aniline 5, 328 antimony pentachloride 59 complexes 71–73, 72 antimony pentachloride affinity 26, 75–78, 80–81 comparison with boron trichloride scale 103 experimental determination 73 arenes diiodine basicity 245 diiodine complexation solvent effects 242 4-fluorophenol hydrogen-bond basicity 122 4-fluorophenol hydrogen-bond enthalpies and entropies 175 methanol IR hydrogen-bonding shift 190 solvatochromic comparison 214 aromatic ethers 4-fluorophenol hydrogen-bond basicity 145 methanol IR hydrogen-bonding shift 198 solvatochromic comparison 216 aromatic N-heterocycles aluminium cation basicity 364 boron trifluoride affinity 91–92 complexation constants 40 copper monocation affinity and basicity 372, 376 diiodine basicity 250–253

Index 4-fluorophenol hydrogen-bond basicity 131–135 4-fluorophenol hydrogen-bond enthalpies and entropies 177–178 lithium cation affinity and basicity 342, 347–348 methanol IR hydrogen-bonding shift 193, 194 potassium cation affinity and basicity 361 sodium cation affinity and basicity 355–356 solvatochromic comparison 215 5-membered rings diiodine basicity 254 diiodine complexation enthalpy 289 4-fluorophenol hydrogen-bond basicity 135 halogen-bond spectroscopic basicities 297 see also nitrogen bases arsine oxides 4-fluorophenol hydrogen-bond basicity 159 methanol IR hydrogen-bonding shift 204 arsines, halogen complex bond geometry 235 arylacetylenes, 4-fluorophenol hydrogen-bond basicity 122 arylamines, 4-fluorophenol hydrogen-bond basicity 129–130 atoms in molecules (AIM) 18–20 azines, cationic adduct structure 329 azoles, cationic adduct structure 329 basic sites 19 cationic adducts 326–335 multiple 38–42 basicity-dependent property (BDP) plots 52–56 hydrogen-bond complexes 111–113 basin 19 bathochromic shift 211–213 see also solvatochromic comparison BDP plots 52–53 Principal Component Analysis 53–56 benzaldehyde, Brönsted basicity 5 benzamidines, 4-fluorophenol hydrogen-bond basicity 136 benzene 80 diiodine complexation constant 241 gas-phase cationic adduct structure 327 benzonitrile antimony pentachloride affinity 79 antimony pentachloride thermodynamic equilibrium data 78 benzophenones 4-fluorophenol hydrogen-bond basicity 148–149 methanol IR hydrogen-bonding shift 199

449

benzoyl chloride, antimony pentachloride complexation 72 beta and xi equation 52–53 blackbody infrared dissociation (BIRD) 338 bond critical point 19 boron 46 boron trifluoride 7, 49, 59 complex formation 85–86 boron trifluoride affinity calorimetry 402–406 comparison with antimony pentachloride scale 103 complex structures 86–88, 87 computation 104–105 liquid/liquid reactions 406–408 acetophenones 94 amides 97–98 amines 91 camphors 93–94 carboxamidates 98–99 cyclohexenones 95 esters 97 ethers 92–93 ketones 93, 93–94, 97 nitriles 92 nitro compounds 99 sulfinyl compounds 99 sulfonyl compounds 99 thioethers 100 thiophosphoryl compounds 100 bromine 46 bromoalkanes 4-fluorophenol hydrogen-bond basicity 163 4-fluorophenol hydrogen-bond enthalpies and entropies 182 solvatochromic comparison 219 3-bromocamphor, antimony pentachloride complexation 72 bromophenyl ethanone, antimony pentachloride complexation 72 2-bromopyridine 5 Brönsted acids 3 Brönsted basicity xiii definition 2–3 gas phase 6 in solution 3–4 butyl ethanoate 80 butyrophenones 94–95 calorimetric cell 404 calorimetry 168–170, 402–406 antimony pentachloride affinity 73 boron trifluoride affinity 89 camphors 93–94

450

Index

carbamates 4-fluorophenol hydrogen-bond basicity 153–154 lithium cation affinity and basicity 351 carbimazole 41 carbon 46 carbon bases cation-π interaction 383–386 copper monocation affinity and basicity 376 copper monocation basicity and affinity 372 cyclopentadienylnickel basicity 370 ECW parameters 49 gas-phase basicity 326–327 lithium 341 lithium cation affinity and basicity 347 manganese cation basicity 368 potassium cation affinity and basicity 361 sodium cation affinity and basicity 355, 359 carbon monoxide 345 carbon tetrachloride 25, 79, 167–168, 168–169 4-fluorophenol complexation constants in 32 diiodine complexation constants in 33, 242 carbonates aluminium cation basicity 365 boron trifluoride affinity 97 4-fluorophenol hydrogen-bond basicity 151–152 methanol IR hydrogen-bonding shift 200–201 solvatochromic comparison 217 carbonyl compounds 4-fluorophenol basicity, solvent effects 34 antimony pentachloride affinity 75–76 cationic adduct structure 331–332 copper monocation basicity and affinity 373 cyclopentadienylnickel basicity 370 diiodine basicity 258–261, 308 diiodine complexation enthalpy 290–291 diiodine complexation solvent effects 242 ECW parameters 49 halogen-bond spectroscopic basicities 298–300, 357 lithium cation affinity and basicity and affinity 342–343 manganese cation basicity 368, 368–369 methylammonium cation affinity and basicity 378 potassium cation affinity and basicity 362 sodium cation affinity and basicity 357, 359–360 see also oxygen bases carboxamidates boron trifluoride affinity 98–99 4-fluorophenol hydrogen-bond basicity 155 methanol IR hydrogen-bonding shift 202

carboxylic acids hydrogen-bond affinity 111 lithium cation affinity and basicity 350 cation/π interaction 383–384 cationic adducts 326–334 amino acids 332–333 carbon 326–328 carbonyl bases 331–332 nitrogen bases 328–330 oxygen bases 330–331 phosphoryl bases 333–334 cationic basicity and affinity see gas-phase basicity and affinity cations 323–324 charge-controlled reactions 11 chelate effects, gas-phase cationic adducts 329–331 chemometrics 52–56 chlorine 46 4-chloroacetophenone, diiodine complexation constants 241 chloroalkanes 4-fluorophenol hydrogen-bond basicity 163–164 4-fluorophenol hydrogen-bond enthalpies and entropies 182 methanol IR hydrogen-bonding shift 205 solvatochromic comparison 218 chlorobenzene 4-fluorophenol complexation constants in 32 diiodine complexation constants in 33 chloroform 25, 49 hydrogen-bond enthalpy 183 10-chlorophenoxarsine oxide 72 3-chloroquinuclidine 121 chlorotrimethylstannane 49 chromium 46 cobalt 46 collision-induced dissociation threshold (CIDT) 335 complexation constants 171–172 antimony pentachloride complexes 78 diiodine 241 hydrogen-bond 113, 120–162 measurement 21–22, 408–413, 424–428 solvent choice and 32–33 complexation enthalpy see enthalpy copper 46 copper monocation affinity and basicity 366–371, 379 acetophenone derivatives 374 amines 372 aromatic N-heterocycles 372 carbon bases 372 comparison to manganese basicity scale 389

Index esters 374 two-ligand 376 cotinine 42, 121, 165 critical point (of electron density gradient) 19 crown ethers 257 cyanoacetylene 183, 184 hydrogen-bond enthalpy 183 cyanogen 72 cycloalkanones boron trifluoride affinity 93–94 4-fluorophenol hydrogen-bond basicity 147 4-fluorophenol hydrogen-bond enthalpies and entropies 179 methanol IR hydrogen-bonding shift 199 solvatochromic comparison 217 cyclohexane 25 cyclohexanol 183 cyclohexanone 95 cyclohexenones 95 cyclohexylmethyl metal amines 41 cyclopentadienylnickel basicity 360, 366, 370, 379 cyclopropylamine 413–418 cytosine 332–333, 344 density functional theory (DFT) 13 di-n-butylamine, Brönsted basicity 5 di-n-propyl ether, antimony pentachloride affinity 79 dibutyl ether, diiodine complexation constants 241 1,4-dichlorobenzene, 4-fluorophenol complexation constants in 32 dichloroethane 4-fluorophenol complexation constants in 32 diiodine complexation constants in 33 1,2-dichloroethane 25 dichloromethane 25, 103, 167 4-fluorophenol complexation constants in 33 boron trifluoride affinity measurements in 90, 102–103 diiodine complexation constants in 33 3,5-dichloropyridine, Brönsted basicity 5 dicycloproyl ketone, Brönsted basicity 5 diethyl ether, Brönsted basicity 5 diethyl ketone, Brönsted basicity 5 diethyl sulfide 50 diethylamine, Brönsted basicity 5 diethylphosphoryl ethane 49 diiodine 424–428 affinity scale 285–286 basicity scale 237, 302 alkenes 246 alkylbenzenes 244–245 arenes 245

451

definition 237–239 drugs 281–282 experimental technique 429 haloalkanes 280–282 IR spectroscopic shifts 304 isothiocyanates 281 methods 238–239 oxides 266 solvent effects 239–240 sulfides 273–276 temperature correction 239 thiocarbonyl compounds 267 thiols 273 thiophosphoryl compounds 276 visible-band blueshift 306–309 complexation constants and solution 33 complexation enthalpy alkylbenzenes 287, 295 amines 287–288 ethers 289 five-membered N-heterocycles 289 phosphoryl compounds 291 pyridines 288 selenium bases 293–294 thiophenes 287 thiophosphoryl compounds 293 ECW parameters 49 dimethyl ether 50 boron trifluoride affinity 103 Brönsted basicity 5 dimethyl selenide 51 dimethyl sulfoxide 49, 429–434, 439–444 Brönsted basicity 5 ECW parameters 49 N,N-dimethylacetamide Brönsted basicity 5 diiodine complexation constants 241 dimethylamine, Brönsted basicity 5 N-N-dimethylaniline, Brönsted basicity 5 5,5-dimethyl-2-cyclohexen-1-ones, boron trifluoride affinity 95–96 dimethylformaldedyde, diiodine complexation constants 33 N-N-dimethylthioacetamide 5 2,6-dimethylpyridine 5 diphenyl sulfide 241 diphenyl sulfoxide 72 displacement reactions 9 disulfides 162 donor number see antimony pentachloride affinity drugs, diiodine basicity 281–282 ECT model 50–51 ECW model 47–52

452

Index

electron convention 339 electron density function 19 electronegativity 44–45 enthalpy 26 4-fluorophenol complexes 170–174, 182–185 complexation constant and 30 diiodine complexes 287–295 FTIR spectrometry 413–418 hydrogen-bond dependent scales 113 measurement 25–28, 172–174 see also calorimetry equilibrium constant see complexation constants esters aluminium basicity 365 boron trifluoride affinity 97 copper monocation basicity and affinity 374 diiodine basicity 258 4-fluorophenol hydrogen-bond basicity 151, 151–152 4-fluorophenol hydrogen-bond enthalpies and entropies 179 lithium cation affinity and basicity 350 manganese cation basicity 368 methanol IR hydrogen-bonding shift 200–201 methylammonium cation affinity and basicity 378 solvatochromic comparison 217 ethanol 5 ethers aliphatic, diiodine basicity 256 aluminium cation basicity 364 antimony pentachloride affinity 75 aromatic 4-fluorophenol hydrogen-bond basicity 145 methanol IR hydrogen-bonding shift 198 solvatochromic comparison 216 boron trifluoride affinity 92 copper monocation affinity and basicity 373, 376 crown 257 cyclopentadienylnickel basicity 370 diiodine basicity 256–257 visible-band blueshift 308 diiodine complexation enthalpy 289 diiodine complexation solvent effects 242 ECW parameters 49 4-fluorophenol hydrogen-bond basicity 143 halogen-bond spectroscopic basicities 297–300 lithium cation affinity and basicity 342–343, 349

manganese cation basicity 368 methanol IR hydrogen-bonding shift 197–198 methylammonium cation affinity and basicity 377–378 potassium basicity and affinity 362 sodium cation affinity and basicity 356 solvatochromic comparison 216 see also oxygen bases ethyl acetate 241 ethyl chloroacetate, antimony pentachloride affinity 79 ethyl dichloroacetate 78 ethyl dimethylcarbamate, antimony pentachloride affinity 79 ethyl sulfide, Brönsted basicity 5 ethyl trichloroacetate 78 ethylamine 5 ethylene diamine 329 fluorine 43, 46 fluoroalkanes 4-fluorophenol hydrogen-bond basicity 163 methanol IR hydrogen-bonding shift 205 3-fluorophenol, comparison with other hydrogen-bond affinities 186 4-fluorophenol 59 comparison of basicity/affinity measures 185–188 comparison with other hydrogen-bond affinities 186 ECW parameters 49 effects 32 FTIR spectrocopy 408–413 hydrogen-bond affinity 170–174, 183–185 early studies 168–170 measurement 408–413 acetophenones 147–148 alcohols 142 aldehydes 146 aliphatic ketones 146–147 alkenes 123 alkynes 123 amides 153–154 amine oxides 159 arenes 122 aromatic N-heterocycles 131–134 5-membered 135 arsine oxides 159 arylacetylenes 123 arylamines 129–130 benzophenones 148–149 bromoalkanes 163 carbamates 154 carbonates 151, 152

Index carboxamidates 155 cycloalkanones 147 disulfides 162 esters 151, 152 ethers 143 haloalkanes 163–164 hindered pyridines 122 imides 165 imines 139 isothiocyanates 160 lactones 151, 152 nitro compounds 156 oximes 139 phenols 142 phosphoroso compounds 159 polybasic compounds 165 primary amines 125 secondary amines 125 seleninyl compounds 158 steric effects 185–186 styrenes 123 sulfides 162 sulfinyl compounds 158 sulfonyl compounds 157 temperature effects 171–172, 172 tertiary amines 127–128 thiocarbonyls 160–161 thiols 162 thiophosphoroso compounds 165 ureas 154 water 142 hydrogen-bond enthalpies 173 aldehydes 179 alkenes 175 amides 180 amidines 178 amines 175–177 arenes 175 aromatic N-heterocycles 177–178 disulfides 181 ethers 178–179 haloalkanes 181–182 imines 178 nitriles 178 phosphoroso compounds 180–181 selenoxides 180 sulfides 181 sulfites 180 sulfoxides 180 thiocarbonyl compounds 181 thiols 181 ureas 180 formaldehyde 331 formamidines, 4-fluorophenol hydrogen-bond basicity 136

453

Fourier transform infrared (FTIR) spectroscopy 119–120, 167, 172–173, 408–413, 413–418, 418–419 furan, see also ethers gas-phase basicity and affinity 323–326 boron trifluoride 103 Brönsted 6 computation 57 experimental determination 439–444 hydrogen-bond basicity scales and 386 measurement 21–22, 334–339 relationship of basicity to affinity 381–386 relationship between scales 387–389 solution basicity and 30–31 gas-phase basicity scales aluminium 354, 364–365 copper 367–371 lithium 341–345, 347–352 manganese 354, 367, 368–369 methylammonium 371, 377–378 potassium 353, 361–363 sodium 355–360 general interaction properties function (GIPF) 56–57 geometry optimization 13–14 Gibbs energy, complexation enthalpy and 26, 27 guanidines 4-fluorophenol hydrogen-bond basicity 137 methanol IR hydrogen-bonding shift 195 solvatochromic comparison 215 guanine 332, 344 haloalkanes copper monocation basicity and affinity 375 copper monocationic affinity and basicity 375 diiodine basicity 280–282 4-fluorophenol hydrogen-bond basicity 163 4-fluorophenol hydrogen-bond enthalpies and entropies 181–182 lithium cation affinity and basicity 343–344, 352 methanol IR hydrogen-bonding shift 205–206 solvatochromic comparison 218–219 halobenzenes 327–328 halogen bonding 229–230 bond geometry 231–236 stretching-force constants 233–234 X-ray crystallography 234–236 halogen-bond basicities 295–302, 303–307 experimental determination 436–439

454

Index

halogen-bond basicities (Cont.) alkenes 295 amines 295–296 carbonyls 298–300 ethers 297–298 experimental determination 434–437 family relationships 302 hydrogen-bond scales and 230 phosphoryl compounds 300–301 pyridine derivatives 296–297 quinoline 296 relation between different halogen basicities 284–285 selenium bases 301–302 sulfides 301 sulfinyl compounds 300 sulfonyl compounds 300 thiocarbonyl compounds 301 thiophene 295 thiophosphoryl compounds 301 hard acids and bases 43, 44, 102–103 hydrogen-bond donors 111–112 heptane 25, 33 hexafluoro-2-propanol, hydrogen-bond enthalpy 183 hexamethylbenzene 5, 241 hexamethylphosphoric triamide 5 high-dilution calorimetry 168–169 high-pressure mass spectrometry (HPMS) 334–335 histamine 329 HSAB principle 43–47, 427–428 boron trifluoride 102–103 hydrazine dithione 72 hydrogen 46, 379 hydrogen halides 111 hydrogen-bond acceptors 114–115 polyfunctional 120–121 hydrogen-bond affinity scales 112, 113, 117–119, 136–138 4-fluorophenol 119–120, 121–145 4-fluorophenol see 4-fluorophenol BDP plot principal components analysis 54–55 cationic hydrogen-bond scales and 386–387 comparison of measures 185–188 experimental methods 418–424 solvatochromic comparison see solvatochromic comparison validity range 167–168 halogen-bond scales and 230 methanol see methanol hydrogen-bond complexes 42 structure 113–117 hydrogen-bond donors 111, 174

bond geometry 114–115, 116 enthalpies 183 imidazole, see also aromatic N-heterocycles imides 4-fluorophenol hydrogen-bond basicity 165 lithium cation affinity and basicity 351 imines 4-fluorophenol hydrogen-bond basicity 139 methanol IR hydrogen-bonding shift 195 indium 43 infrared (IR) spectroscopy 112, 113, 239 halogen-bond complexes 286, 302–306 hydrogen-bond donoro affinity scales 117–118 see also Fourier transform infrared spectroscopy iodine bromide, ECW parameters 49 iodine complexes see diiodine iodine cyanide basicity 303–304, 434–437 iodine monochloride basicity 302, 303 see also halogen-bond spectroscopic basicities iodoalkanes 4-fluorophenol hydrogen-bond basicity 163 4-fluorophenol hydrogen-bond enthalpies and entropies 182 solvatochomic comparison 219 iodocyclohexane 424–428 ion convention 340 ion cyclotron resonance (ICR) mass spectrometry 336 ion pair compounds 4-fluorophenol hydrogen-bond basicity 165–166 methanol IR hydrogen-bonding shift 206 iron 46 isobutyl methyl ketone, 4-fluorophenol basicity 408–413 isocyanates 375 isocyanic acids 183 isothiocyanates copper monocation basicity and affinity 375 diiodine basicity 281 4-fluorophenol hydrogen-bond basicity 160–161 hydrogen-bond affinity 181 methanol IR hydrogen-bonding shift 204 ketones aluminium cation basicity 364–365 boron trifluoride affinity 93, 97 copper monocation affinity and basicity 376 diiodine basicity 259

Index 4-fluorophenol hydrogen-bond basicity 146–147 copper monocation basicity and affinity 373 methanol IR hydrogen-bonding shift 198 solvatochromic comparison 217 lithium cation affinity and basicity 349–350 manganese cation basicity 368 methanol IR hydrogen-bonding shift 199–200 see also cycloalkanones lactams 4-fluorophenol hydrogen-bond basicity 153–154 methanol hydrogen-bond affinity 201 lactones diiodine basicity 258 4-fluorophenol hydrogen-bond basicity 151–152 methanol IR hydrogen-bonding shift 200–201 solvatochromic comparison 217 lattice energy 339 Lewis acids 7 Lewis affinity measurement 24–29 boron trifluoride 88–89 equipment 403–405 Lewis bases 8 Lewis basicity definition 6–9 quanitative 42–47 measurement 20–22 boron trifluoride 89 gas-phase 21–22 in solution 22–24 Lewis, G.N. xiii linear combination of atomic orbitals (LCAO-MO) 11, 12, 13 liquid/liquid reactions 406–408 lithium 46 adduct formation 329 cation properties 379 gas-phase basicity scales 325, 340, 341–345, 346, 347–352 compared to aluminium 388 manganese cation properties 379 gas-phase basicity scales 325, 354, 365–366, 367, 368–369 compared to aluminium 388 compared to copper 389 mass spectrometry 334–339, 439–444

455

matrices xiv mercury, Lewis basicity 43 mesistylene 418–419 methanol 5, 59–60 antimony pentachloride affinity 75 FTIR spectroscopy 418–424 hydrogen-bond enthalpy 183 IR stretching wavenumber on bonding to Lewis bases 1,10-phenanthrolines 194 acetophenones 199 alcohols 195, 197 alkenes 190 alkynes 190 amidines 195 amine oxides 203 amines 190–192 arenes 190 aromatic heterocyclics 192–193, 194 arsine oxides 204 benzophenones 199, 200 carbonates 200–201 carboxamidates 202 esters 200–201 ethers 197–198 aldehydes 198 aromatic 198 formamidines 194–195 haloalkanes 205–206 imines 195 ion pair compounds 206 isothiocyanates 204 ketones 199, 199–200 aliphatic 198–199 lactones 200–201 nitriles 196–197 peroxides 198 phosphoroso compounds 204 relation to basicity 208–210 seleninyl compounds 203 selenophosphoroso compounds 206 sulfinyl compounds 203 sulfonyl compounds 203 thiocarbonyl compounds 204 thioethers 205 thiophene 190 thiophosphoroso compounds 206 methyl acetate 5 methyl dithiovalerate 33 methyl phenyl sulfoxide 173, 439–444 methyl propionate 5 methyl sulfide 5 4-methylacetophenone, diiodine complexation constants 241 methylamine, Brönsted basicity 5

456

Index

methylammonium 371–381, 377–378 cation properties 379 methylimidazole, diiodine complexation constants 33 N-methylmorpholine 419 methylnitrile, antimony pentachloride complexation 72 2-methylpyridine, 241 3-methylpyridine-N-oxide, 72 4-methylpyridine 5 molecular orbital models 10–12 molybdenum 46 morpholine 209 N-oxides see amine oxides naming conventions 1–2 natural bond orbital (NBO) theory 17–18 NBO theory 17–18 neutron diffraction crystallography 115 nickel 46 see also cyclopentadienylnickel nicotine, 4-fluorophenol hydrogen-bond basicity 165 niobium 46 nitriles 289 antimony pentachloride affinity 75 boron trifluoride affinity 92 copper monocation affinity and basicity 376 copper monocation basicity and affinity 373 cyclopentadienylnickel basicity 370 diiodine complexation solvent effects 242 4-fluorophenol hydrogen-bond basicity 140–141 halogen-bond spectroscopic basicities 297 lithium cation affinity and basicity 342, 348 manganese cation basicity 368 methanol IR hydrogen-bonding shift 196–197 potassium cation affinity and basicity 361–362 sodium cation affinity and basicity 359 solvatochromic comparison 215 see also nitrogen bases nitro compounds antimony pentachloride affinity 76 boron trifluoride affinity 99 4-fluorophenol hydrogen-bond basicity 156 4-fluorophenol hydrogen-bond enthalpies and entropies 180 methanol IR hydrogen-bonding shift 203 solvatochromic comparison 218 4-nitroaniline 210–211 see also solvatochromic comparison nitrobenzene antimony pentachloride affinity 79

antimony pentachloride thermodynamic equilibrium data 78 Brönsted basicity 5 nitrogen bases ECW parameters 49 gas-phase cationic adduct structure 328–330 nitromethane, lithium cation affinity and basicity 351 4-nitrophenol 210–213 see also solvatochromic comparison nitroso compounds methanol IR hydrogen-bonding shift 203 nuclear magnetic resonance (NMR) spectroscopy 112, 239 nucleobases 344 manganese cation basicity 369 potassium cation affinity and basicity 362 orbital-controlled reactions 11 oxides diiodine basicity 266 group 15 organoderivatives 159 oximes 139 oxygen 46 oxygen bases diiodine complexation enthalpy 291–292 halogen-bond spectroscopic basicities 303–305 palladium 46 pentachlorophenol, hydrogen-bond complexation 115 pentan-2-one, antimony pentachloride affinity 79 perfluoro-tert-2-butanol, hydrogen-bond enthalpy 183 peroxides 4-fluorophenol hydrogen-bond basicity 144 methanol IR hydrogen-bonding shift 198 1, 10-phenanthrolines, methanol IR hydrogen-bonding shift 194 1,10-phenanthrolines, diiodine basicity 253 phenol ECW parameters 49 hydrogen-bond enthalpy 183 solvatochromic comparison 215 phenols 111, 210–211 4-fluorophenol hydrogen-bond basicity 142 5-phenyl-2-adamantone, antimony pentachloride complexation 72 phosophoroso compounds 4-fluorophenol hydrogen-bond enthalpies and entropies 180–181 methanol IR hydrogen-bonding shift 204

Index phosphine boron trifluoride affinity 100 boron trifluoride complexation 88 phosphines 434–436 antimony pentachloride affinity 77 cyclopentadienylnickel basicity 370 halogen complex bond geometry 235 phosphoramides, antimony pentachloride affinity 74 phosphoroso compounds, 4-fluorophenol hydrogen-bond basicity 159 phosphorus 46 phosphoryl chloride antimony pentachloride complexation 72 antimony pentachloride thermodynamic equilibrium data 78 phosphoryl compounds antimony pentachloride affinity 77 boron trifluoride affinity 99–100, 99–100 cationic adducts 333–334 diiodine basicity 264 diiodine complexation, solvent effects 242 diiodine complexation enthalpy 291 ECW parameters 49 halogen-bond spectroscopic basicities 300–301 lithium cation affinity and basicity 351–352 potassium cation affinity and basicity 362 see also oxygen bases polybasic compounds 38–42, 165 boron trifluoride affinity 104 polybasic compounds halogen-bond affinity 305 potassium 46 cation properties 379 gas-phase basicity and affinity scales 325, 353, 361–362 principal components analysis (PCA) 52–56 progesterone 165, 408–412 propiophenones 94 proprionyl chloride 78 propylene carbonate 79 proton affinity 5 proton exchange 2–3 pure-base calorimetry 169 pure-solvent calorimetry 169–170 pyrazine 121, 241, 252 pyridine 50 antimony pentachloride affinity 75 boron trifluoride affinity 402–406 liquid/liquid reaction measurement 406–407 Brönsted basicity 5

457

methanol IR wavenumber shift 191 methanol spectroscopic basicity 419 pyridines 51, 436–439 4-fluorophenol basicity 34, 122 Brönsted basicity 5 diiodine basicity 250–251, 252–253, 265 diiodine complexation enthalpy 288 diiodine complexation solvent effects 242 ECW parameters 49 halogen complex bond geometry 235 halogen-bond spectroscopic basicities 296–297 methanol IR hydrogen-bonding shift 190, 208–209 solvatochromic comparison 214 pyrrole 49, 183 quantum chemistry atoms in molecules 18–20 basicity scale descriptors 56–58 boron trifluoride affinity 104–105 model complexes 57 natural bond orbital (NBO) theory 17–18 perturbation molecular orbital theory 10–12 valence-band models 10 variational supermolecular methods 12–17 quinoline 5, 241, 296 quinones 149 quinuclidine 5, 51 see also aromatic N-heterocycles radiative association kinetics (RAK) 338–339 reference acids xiii relativity 15 rhodium 46 Rose-Drago method 424–427 rotational spectroscopy 232–234 rubidium 46 selected ion flow tube (SIFT) method 337 seleninyl compounds antimony pentachloride affinity 76 4-fluorophenol hydrogen-bond basicity 158 selenium 46 selenium bases diiodine complexation 277–279, 293–294 halogen-bond spectroscopic basicities 301–302 selenium oxychloride 72 selenoanisole 241, 277 selenoethers 235 selenophosphoroso compounds 165, 182 selenoxides 115–116 SIFT (selected ion flow tube) 337 silicon 46

458

Index

sodium cation affinity and basicity 346, 352 alcohols 356, 359 amino acids 357, 357–358 carbonyl compounds 357 cation properties 379 electronegativity 46 ethers 356 gas-phase basicity scales 325 sulfinyl compounds 360 soft acids and bases 43, 44 solution basicity computation 57–58 gas-phase basicity and 30–31 measurement 22–23 solvent effects 29–34 solvatochromic comparison 210–213, 220–221, 420–424 acetophenone 217 alcohols 216 aliphatic ketones 217 alkenes 214 amides 217–218 amines 214–215 arenes 214 aromatic N-heterocycles 215 carbonates 217 cycloalkanones 217 disulfides 218 esters 217 ethers 216 guanidines 215 haloalkanes 218–219 lactones 217 nitriles 215 nitro compounds 218 phosophoro compounds 218 pyridine 214 sulfinyl compounds 218 sulfonyl compounds 218 water 216 solvents 24, 25 antimony pentachloride affinity and 79 basicity measurement and 29–34 diiodine basicity scale 239–240 effect on antimony tetrachloride affinity 79 effect on boron trifluoride affinity 102–103 hydrogen-bond measurements 167–168 standardisation for measurement 32–33 spectroscopic basicity measurement 34–37, 117–118 hydrogen-bond donors 112 methanol scale 188–189, 190–206 see also Fourier transform infrared spectroscopy; visible spectroscopy

standard concentration equlibrium constant 21–22 standard pressure equilibrium constant 21 statistical errors 28 steric hindrance 187–188, 189 styrenes 122 sulfides cyclopentadienylnickel basicity 370 diiodine basicity 273–276 4-fluorophenol hydrogen-bond basicity 162 halogen-bond spectroscopic basicities 301 lithium cation affinity and basicity 343 manganese cation basicity 368 sulfinyl compounds antimony pentachloride affinity 76 boron trifluoride affinity 99 4-fluorophenol hydrogen-bond basicity 158 halogen-bond spectroscopic basicities 300 lithium cation affinity and basicity 351 methanol IR hydrogen-bonding shift 203 sodium cation affinity and basicity 360 solvatochromic comparison 218 sulfonyl compounds 4-fluorophenol basicity 34, 157 antimony pentachloride affinity 76 boron trifluoride affinity 99 4-fluorophenol hydrogen-bond enthalpies and entropies 180 halogen-bond spectroscopic basicities 300 lithium cation affinity and basicity 351 phosphorus cation affinity and basicity 362 potassium cation affinity and basicity 362 solvatochromic comparison 218 sulfoxide compounds 180 sulfur 46 sulfur bases, ECW parameters 49 sulfur dioxide 49 supermolecules 12–13 anharmonicity 14–15 atomic orbital basis sets 13 energy decomposition schemes 16–17 experimental validation 15 geometry optimization 13–14 see also halogen-bond complexes; hydrogen-bond complexes temperature aluminium cation affinity and 366 correction 239 equilibrium constants and 26, 27 tetrachloroethylene 25 tetrachloroethylene carbonate 72 tetrahydrofuran antimony pentachloride solvent influence 79 boron trifluoride affinity 103

Index Brönsted basicity 5 diiodine complexation constant 33 tetrahydropyran 103 tetramethylbenzoyl chloride 72 tetramethylurea 5 tetranitrogen tetrasulfide 72 thiazole 92 thiazolidine 125 thioanisole 241 thiocarbonyl compounds diiodine basicity 267 diiodine complexation solvent effects 242 halogen-bond complex geometry 235 halogen-bond spectroscopic basicities 301 methanol IR hydrogen-bonding shift 204 thioesters, ECW parameters 49 thioethers boron trifluoride affinity 100 diiodine complexation solvent effects 242 lithium cation affinity and basicity 351, 352 thiols copper monocation basicity and affinity 375 cyclopentadienylnickel basicity 370 diiodine basicity 273 4-fluorophenol hydrogen-bond basicity 162 hydrogen-bond affinity 181 lithium cation affinity and basicity 352 manganese cation basicity 368 thionyl chloride, Brönsted basicity 78 thionyl compounds ECW parameters 49 see also oxygen bases thionyls diiodine basicity 263 diiodine complexation solvent effects 242 thiooethers, halogen-bond complex geometry 235 thiophene 4-fluorophenol hydrogen-bond enthalpies and entropies 175 halogen-bond spectroscopic basicities 295 thiophenes diiodine basicity 245–246 diiodine complexation enthalpy 287 thiophosphoroso compounds 4-fluorophenol hydrogen-bond basicity 165 4-fluorophenol hydrogen-bond enthalpies and entropies 182 methanol IR hydrogen-bonding shift 206 thiophosphoryl compounds diiodine basicity 276, 308 diiodine complexation enthalpy 293

459

halogen-bond spectroscopic basicities 301 thiourea-diiodine complexes 232 thioureas 269–271 tri-n-butylammonium 183 trichloromethane, diiodine complexation constants in 33 triethoxyaminophosphine sulfide, diiodine complexation constants 241 triethoxyphosphine sulfide, diiodine complexation constants 241 triethylamine 5, 74 triethylgallane 49 trifluoroethanol 49 hydrogen-bond enthalpy 183 3-trifluoromethylphenol 186 trimethylalumane 49 trimethylamine boron trifluoride affinity 103 boron trifluoride complexation 88 Brönsted basicity 5 trimethylbenzoyl chloride, antimony pentachloride complexation 72 trimethylborane, ECW parameters 49 trimethylgermanium, gas-phase basicity scales 325 trimethylindane 49 trimethylphosphine oxide, antimony pentachloride complexation 72 2,4,6-trimethylpyridine, Brönsted basicity 5 trimethylsilyl, gas-phase basicity scale 325 trimethyltin 325 triphenylamine 5 triphenylphosphine, Brönsted basicity 5 triphenylphosphine oxide antimony pentachloride complexation 72 Brönsted basicity 5 diiodine complexation constants 33 triphenylphosphine selenide 33 triphenylphosphine sulfide 33, 50 ultraviolet (UV) spectroscopy 112 diiodine complexation constants 238 ureas 4-fluorophenol hydrogen-bond affinity 180 boron trifluoride affinity 97–98 diiodine basicity 262 4-fluorophenol hydrogen-bond basicity 154 urethanes see carbamates valence-band modeling 10 vanandium, electronegativity 46 van’t Hoff equation 26 vaporization energy 339

460

Index

visible-band spectroscopy 306–309, 424–428, 429–434 water antimony pentachloride complexation 72 4-fluorophenol hydrogen-bond basicity 142 lithium cation affinity and basicity 342 methanol hydrogen-bond affinity 196

parameters for organic base protonation 5 sodium cation affinity and basicity 356 X-ray structures 72, 73 halogen-bonded complexes 234–236 zinc 43 zirconium 46

Plate 1 – see Figure 2.1




(a)

(b)


a

b

c

d Plate 6 – see Figure 5.2

a

b

c


a

b

c




a

b Plate 11 – see Figure 6.3



a





a


1

Absorbance

0.8

0.6

0.4 E D C 0.2 B

0 3700

A 3600

3500

3400

3300 -1

Wavenumber / cm


3200

3100

1

Absorbance

0.8

0.6

E D

0.4

C B

0.2

A 0 3700

3600

3500

3400

3300

Wavenumber / cm

3200

3100

-1


0.12

0.1

Absorbance

0.08

0.06

10 25

0.04

40 55 65

0.02

0 3640

3540

3440

3340

3240 -1

Wavenumber / cm


3140

3040


0,4 a

b

c

d

e

Absorbance

0,3

0,2

0,1

0 3800

3100 -1

Wavenumbers / cm


1,2 1

Absorbance

0,8 0,6 0,4 0,2 0

285.2

250

306.7 Wavelength / nm

380


2

Absorbance

1,5

1

0,5

0 400

450

500

550

Wavelength / nm


600

650

1.4

1.2

5 4

1

Absorbance

3 0.8 2 0.6 1 0.4

0.2

0 350

400

450

500

550

600

550

600

650

Wavelength / nm

Plate 26 – see Figure 7.11 1.4

1.2

1

Absorbance

-4.6 0.8

9.7

0.6

25.1

0.4

39.7 53.1

0.2

0 350

400

450

500 Wavelength / nm


650

1

b

c

d

Absorbance

a

0 500

390 -1

Wavenumber / cm


1 a

Absorbance

0,8

0,6 b 0,4

0,2

0 350

400

450

500

550

Wavelength / nm


600

650

700

Lewis Basicity and Affinity Scales: Data and Measurement

Acidity and Basicity

Affinity

Measurement Data Modeling and Parameter Estimation

Affinity

Affinity

Affinity

Affinity

Scales and a Tail

Scales and Arpeggios

affinity chromatography principles and methods

Affinity Chromatography Methods and Protocols

Affinity Chromatography: Methods and Protocols

Scales and a Tail

Scales and a Tail

Scales and a Tail

Scales and a Tail

Tails and Scales

Scales and a Tail

Affinity Chromatography Methods and Protocols

Diatonic Major and Minor Scales

Financial Planning and Counseling Scales

Affinity Chromatography

Quantum Measurement and Control

Tree and Forest Measurement

Diatonic Major and Minor Scales

Data Collection: Planning for and Collecting All Types of Data (Measurement and Evaluation Series)

Quantum Measurement and Control

Biomedical Sensors and Measurement

Experimentation and Measurement

Measurement systems and sensors

Lewis Basicity and Affinity Scales: Data and Measurement

Acidity and Basicity

Affinity

Measurement Data Modeling and Parameter Estimation

Affinity

Affinity

Affinity

Affinity

Scales and a Tail

Scales and Arpeggios

affinity chromatography principles and methods

Affinity Chromatography Methods and Protocols

Affinity Chromatography: Methods and Protocols

Scales and a Tail

Scales and a Tail

Scales and a Tail

Scales and a Tail

Tails and Scales

Scales and a Tail

Affinity Chromatography Methods and Protocols

Diatonic Major and Minor Scales

Financial Planning and Counseling Scales

Affinity Chromatography

Quantum Measurement and Control

Tree and Forest Measurement

Diatonic Major and Minor Scales

Data Collection: Planning for and Collecting All Types of Data (Measurement and Evaluation Series)

Quantum Measurement and Control

Biomedical Sensors and Measurement

Experimentation and Measurement

Measurement systems and sensors

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