Buffers for pH and Metal Ion Control Buffers for pH and Metal Ion Control D. D. Perrin John Curtin School of Medical Research Australian National University Canberra Boyd Dempsey Faculty of Military Studies University of New South Wales Royal Military College Duntroon London CHAPMAN AND HALL First published in 1974 by Chapman and Hall Ltd 11 New Fetter Lane, London EC4P 4EE First issued as a Science Paperback 1979 © 1974 D.D. Perrin and Boyd Dempsey Typeset by Santype Ltd (Coldtype Division) Salisbury, Wiltshire ISBN-J3: 978-0-412-21890-3 e-ISBN-13: 978-94-009-5874-6 DOl: /0./007/978-94-009-5874-6 This paperback edition is sold subject to the condition that it shall not, by way of trade or otherwise, be lent, re-sold, hired out, or otherwise circulated without the publisher's prior consent in any form of binding or cover other than that in which it is published and without a similar condition including this condition being imposed on the subsequent purchaser All rights reserved. No part of this book may be reprinted, or reproduced or utilized in any form or by any electronic, mechanical or other means, now known or hereafter invented, including photocopying and recording, or in any information storage or retrieval system, without the permission in writing from the publisher. Distributed in the US.A. by Halsted Press, a Division of John Wiley & Sons, Inc, New York Contents Preface page viii 1. Introduction 1 1.1 The concept of buffer action 1 1.2 Why are buffers needed? 2 1.3 Some naturally occurring buffers 3 2. The Theory of Buffer Action 4 2.1 Equilibrium aspects 4 2.2 Activity effects 6 2.3 Effect of dilution 7 2.4 Salt effects 8 2.5 Ampholytes and zwitterions 10 2.6 Buffer capacity 10 2.6.1 Buffer capacity of a polybasic acid 12 2.7 Pseudo buffers 15 2.8 Self buffers 15 2.9 Mixtures of buffers 17 2.10 Temperature dependence 18 2.11 Effect of pressure on buffers 18 2.12 Further reading 19 3. Applications of pH Buffers 24 3.1 Factors governing the choice of a buffer 24 3.2 Measurement of pH 25 3.3 Biochemistry and biology 27 3.4 Spectroscopy 32 3.5 Buffers for special applications 33 3.5.1 Volatile buffers 33 3.5.2 Buffers for electrophoresis 34 3.5.3 Buffers for complexometric titrations 34 3.5.4 Buffers for chromatography 35 vi . Contents 3.5.5 Buffers for polarography 35 3.5.6 Buffers for proton magnetic resonance studies 36 3.5.7 Buffers for solvent extraction 36 3.5.8 Isotonic phannaceutical buffers 37 3.5.9 Miscellaneous 38 4. Practical Limitations in the Use of Buffers 55 4.1 Chemical problems 55 4.2 Biological effects 58 4.3 Influence on chemical reactions 60 5. New pH-Buffer Tables and Systems 62 5.1 On calculating buffer composition tables 62 5.1.1 Buffers of constant ionic strength. No added electrolyte 62 5.1.2 Constant ionic strength buffers with added electrolyte 64 5.1.2.1 Preparation of amine buffers of constant ionic strength 65 5.1.3 Buffers by direct titration of weak bases or acids with strong acids or bases 65 5.2 On designing a new pH-buffer system 69 6. Buffers for use in Partially Aqueous and Non-Aqueous Solvents and Heavy Water 77 6.1 pH* Scales 78 6.2 pH* Buffers 78 6.3 The measurement of pH* 79 6.4 A universal pH scale 80 6.5 The pD scale and the measurement of pD 81 6.6 The use of pH* and pD buffers 82 6.6.1 The determination of dissociation constants of acids 82 6.6.2 Rate studies in heavy water 82 6.7 Surfactants 83 7. Metal-ion Buffers 94 7.1 The concept of pM 94 7.2 Uses of metal-ion buffers 95 7.3 Calculation of pM 96 7.4 pH-Independent metal-ion buffers 99 7.5 Effects of pH buffer substances on pM 101 7.6 Anion buffers 102 7.7 Redox buffering 103 8. Purification of Substances Used in Buffers 109 9. Preparation of Buffer Solutions 117 Contents· vii 10. Appendices 123 Appendix I. Tables for constructing buffer tables 123 Appendix II. Composition-pH tables of some commonly used buffers 128 Appendix III. Thermodynamic acid dissociation constants of prospective buffer substances 157 Appendix IV. The Henderson-Hasselbalch equation 164 References 167 Index 173 Preface This book is intended as a practical manual for chemists, biologists and others whose work requires the use of pH or metal-ion buffers. Much information on buffers is scattered throughout the literature and it has been our endeavour to select data and instructions likely to be helpful in the choice of suitable buffer substances and for the preparation of appropriate solutions. For details of pH measurement and the preparation of standard acid and alkali solutions the reader is referred to a companion volume, A. Albert and E. P. Serjeant's The Determination of Ionization Constants (1971). Although the aims of the book are essentially practical, it also deals in some detail with those theoretical aspects considered most helpful to an understanding of buffer applications. We have cast our net widely to include pH buffers for particular purposes and for measurements in non-aqueous and mixed solvent systems. In recent years there has been a significant expansion in the range of available buffers, particularly for biological studies, largely in conse quence of the development of many zwiUerionic buffers by Good et al. (1966). These are described in Chapter 3. However, there are very many substances that could be, or have been, of use as buffers, and Appendix III lists some of these. Chapter 5 shows how new pH-buffer tables can be constructed from the thermodynamic pKa values, and some simple computer programmes are included to facilitate the necessary calculations. Tables and worked examples are given for use if a computer is not available. In view of the importance of metal ion concentrations, particularly in biological work, we have considered it appropriate to include a section on metal-ion buffers. These buffers may also be useful in preparing convenient standards for ion-selective electrodes. Canberra, D. D. Perrin August, 1973 Boyd Dempsey Chapter One Introduction 1.1 The concept of buffer action When partly neutralized weak acids or bases are present in aqueous solution the addition of small amounts of strong acid or strong base causes little change in pH. This resistance to change in the free hydrogen ion concentration of a solution was described by Fernbach and Hubert (1900) as 'buffering'. In studies of the enzyme, amylase, they found that a partly neutralized solution of phosphoric acid acted as a 'protection against abrupt changes in acidity or alkalinity: the phosphates behave as a sort of buffer'. Following this observation use was soon made of mixtures of monohydro gen phosphate/dihydrogen phosphate, ammonia/ammonium chloride, acetate/acetic acid, phthalate/phthalic acid, and p-nitrophenolate/p-nitrophenol to obtain solutions which were 'practically unaffected by the presence of traces of (acidic or basic) impurities in the water or salts used, which is far from being the case with very dilute solutions of strong acids and bases' (Fels, 1904). In Lowry-Bronsted terms, the common feature of these mixtures is the presence of an acid and its conjugate base. This acid-base pair, together with appropriate counter ions, constitutes a buffer substance. A convenient definition (Van Slyke, 1922) of a buffer is a 'substance which by its presence in solution increases the amount of acid or alkali that must be added to cause unit change in pH.' Addition of 1 ml of 1M Hel to a litre of distilled water, pH 7, lowers the pH to 3. Alternatively, addition of I ml of 1 M NaOH raises the pH to 11. Much less change (only about 0.02 pH units) is observed if the same amount of acid or alkali is added to a litre of solution, also pH 7, that is 0.05M in imidazolinium hydrochloride and 0.047M in imidazole. Thus imidazole/imidazolinium ion is a good buffer at pH 7. 2 . Buffers for pH and Metal Ion Control Another useful property of a buffer, at least within the range pH 4 to 10, is that its pH remains substantially unchanged upon dilution of the solution. The effectiveness of a buffer depends on its buffer capacity (resistance to pH change on addition of acid or alkali), the pH change on dilution, and the effects of adding neutral salts or changing the temperature. These are discussed in Chapter 2. 1.2 Why are buffers needed? Many biological and chemical systems involve acid-base equilibria and therefore depend critically on the pH of the solution. An example is the extent to which the viability and growth of organisms and tissues depends on the pH of the cell fluids and of the media in which the cells grow (Albert, 1968). The effectiveness of many chemical separations and the rates of many chemical reactions are governed by the pH of the solution. Buffer solutions offer advantages for controlling reaction conditions and yields in organic syntheses. The objection that this adds material which must later be removed is less serious if volatile buffers are used. A particularly promising field in this respect is biomimetic chemistry (Breslow, 1972) which attempts to imitate natural reactions and enzymic processes as a way to improve the power of organic chemistry. In analytical and industrial chemistry, adequate pH control may be essential in determining the courses of precipitation reactions and of the electrodeposition of metals. Physico chemical studies of reaction kinetics and chemical equilibria often require solutions to be maintained at a definite pH value. Buffers are needed for pH standardization and control in the research laboratory, the factory and the medical clinic. For kinetic, equilibrium and physiological studies it is often desirable to make measurements over a controlled range of pH values while, at the same time, maintaining constant ionic strength in the medium.