Intracellular pH and Its Measurement Authors Arnost Kotyk, Ph.D., D.Sc. and Jan Slavik, Ph.D. Institute of Physiology Czechoslovak Academy of Sciences Prague, Czechoslovakia CRC Press, Inc. Boca Raton, Florida Library of Congress Cataloging-in-Publication Data Kotyk, Amost. Intracellular pH and its measurement. Bibliography: p. Includes index. 1. Cytochemistry-Methodology. 2. Hydrogen-ion concentration—M easurement. I. Slavik, Jan II. Title. QH611.K67 1989 574.87' 6042 88-19253 ISBN 0-8493-4916-8 This book represents information obtained from authentic and highly regarded sources. Reprinted material is quoted with permission, and sources are indicated. A wide variety of references are listed. Every reasonable effort has been made to give reliable data and information, but the author and the publisher cannot assume responsibility for the validity of all materials or for the consequences of their use. All rights reserved. This book, or any parts thereof, may not be reproduced in any form without written consent from the publisher. Direct all inquiries to CRC Press, Inc., 2000 Corporate Blvd., N.W., Boca Raton, Florida, 33431. © 1989 by CRC Press, Inc. International Standard Book Number 0-8493-4916-8 Library of Congress Card Number 88-19253 Printed in the United States PREFACE At the very beginning of my scientific career more than 30 years ago, I was interested in shifts of various phosphate fractions of yeast under different physiological conditions; the measurements included the effects of various ions on these fractions. In several experiments where a Na+-containing medium was compared with a K+-containing one there was a strikingly higher level of inorganic orthophosphate in the yeast with K+ than there was with Na+. We invented several working hypotheses and set out to test them. When I checked the bottles of Na- and K-phosphates we had used, sure enough, one was Na2HP04 and the other KH2P04, with the pH of the somewhat acid-buffered medium being 8.8 and 6.1, respectively. The observed effect was solely due to the difference in pH. This was my first practical lesson on the importance of pH in biology. Others followed later after it had become clear that the intracellular (as well as the extracellular) pH is a quantity that affects the kinetics and the energetics of many transport systems and, finally, when the local pH— whatever its interpretation may be — was found to differ quite generally from the bulk pH. The realization that H+ ions not only play a role in enzyme (and transport) catalysis but also are involved in the energization of some transport processes, has led to a burst of emphasis on accurate measurements of their concentrations, a task seemingly quite straight­ forward but, in practice, hampered by both conceptual and technical problems. We attempted to summarize in this book the present state-of-the-art of pH measurement, with a special accent on intracellular pH, proceeding from the rather extensive experience with several methods of the senior author (A.K.), and the intensive experience with the dual fluorescence method invented by the junior author (J.S.), the method being the one of choice in general laboratory practice where local pH values are sought. We had the opportunity to draw on the know-how of many colleagues, some of whom have read and corrected parts of the book; thus, we wish to acknowledge the following people: Prof. J. Koryta of the J. Heyrovsky Institute of Physical Chemistry and Electro­ chemistry in Prague (Chapters 1 and 2), Dr. J. Budesmsky of the Institute of Organic Chemistry and Biochemistry in Prague, Dr. P. Sedmera of the Institute of Microbiology in Prague, Dr. I. Goljer of the Slovak Technical College in Bratislava (Chapter 4E), and Dr. J. Plasek of the Charles University of Prague (Chapters 4C and 4D). We also wish to thank the 75 authors, quoted at the appropriate places of the book, who sent their reprints for our perusal. Arnost Kotyk Jan Slavik THE AUTHORS Arnost Kotyk, Ph.D., D.Sc., is a senior scientist at the Institute of Physiology, Czech­ oslovak Academy of Sciences in Prague, Czechoslovakia. Dr. Kotyk received his various academic degrees at the University of California in Berke­ ley, at the Charles University in Prague, and at the Institute for Organic Chemistry and Biochemistry also in Prague, between the years 1950 and 1958. Dr. Kotyk is the chairman of the State Research Board on Special Cell Biology, lectures at the J. E. Purkyne University in Brno and at the Charles University in Prague, and holds important positions in the International Union of Biochemistry and in the UNESCO Global Program on Molecular and Cellular Biology. Dr. Kotyk has published 240 original papers, has written six monographs from the field of membrane transport and enzyme kinetics, and has lectured in 26 countries in five lan­ guages. Jan Slavik, Ph.D., is a Research Scientist at the Institute of Physiology, Czechoslovak Academy of Sciences in Prague. He majored in biophysics and received his degrees at the Charles University in Prague. Dr. Slavik is the leader of a grpup researching fluorescence techniques at the Institute of Physiology of the Czechoslovak Academy of Sciences and has published over 25 original papers and several reviews for international journals. Dr. Slavik’s current interests include localization and interactions of membrane transport proteins. TABLE OF CONTENTS Chapter 1 The Concept of pH, Hydrolysis, and Buffers........................................................................... 1 A. Kotyk Chapter 2 pH in Biological Systems............................................................................................................15 A. Kotyk Chapter 3 Methods of Determining Intracellular pH with Electrodes.....................................................29 A. Kotyk Chapter 4A Assays of Intracellular pH Using Chemical Probes: Principles of pH Indicator Response.......................................................................................................................37 J. Slavik Chapter 4B Assays of Intracellular pH Using Chemical Probes: Distribution Techniques......................51 J. Slavik and A. Kotyk Chapter 4C Assays of Intracellular pH Using Chemical Probes: Absorption Spectroscopy..................69 J. Slavik Chapter 4D Assays of Intracellular pH Using Chemical Probes: Fluorescence Spectroscopy................87 J. Slavik Chapter 4E Assays of Intracellular pH Using Chemical Probes: Nuclear Magnetic Resonance Spectroscopy.............................................................................................................127 J. Slavik Chapter 5 The Choice of the Assay Method...................................................... J. Slavik Chapter 6 Heterogeneity of Intracellular pH.............................................................................................169 A. Kotyk and J. Slavik Index 177 1 Chapter 1 THE CONCEPT OF pH, HYDROLYSIS, AND BUFFERS Arnost Kotyk TABLE OF CONTENTS I. Definition of pH..................................................................................................................2 II. Ionization of Acids and Bases..........................................................................................3 III. Hydrolysis of Salts...............................................................................................................8 IV. Buffering Power................................................................................................................10 V. The Activity Coefficient...................................................................................................13 References.......................................................................................................................................14 2 Intracellular pH and Its Measurement I. DEFINITION OF pH Some 80 years ago, Sorensen introduced the symbol pH for the negative logarithm of concentration (moles or equivalents per cubic decimeter) of H+ ions.1 The letter “p” derives either from the Latin pondus (weight or significance) or from potentia (might or potential) and has since been used to denote the negative logarithm of concentration of various other ions; e.g., log cNa+ = -pNa, which is correct or at least consistent. It has also been used for the negative logarithms of various kinetic quantities having dimensions of concentration, such as the dissociation constant K3, where -log Ka = pATa, which is certainly incorrect in the original sense of “p” but is useful in cases where sums or differences, such as pA" - pH, occur in formulae. In Sorensen’s original version pH = -log10 cH+ (1) but it is now defined in terms of activities a. Thus (with 7 being the activity coefficient), pH = — l og10 (cH+ yH+) = — l og qh+ (2) which, particularly in more concentrated solutions, represents an appreciable difference, as will be seen below. It should be observed that in rigorous derivations of these equations both c and a (and, hence, 7) are dimensionless quantities that can be understood as concentrations or activities relative to a unit reference. Their values are thus numerically identical with true concentra­ tions or activities expressed in moles per cubic decimeter. Water, which is the principal milieu of all subcellular and cellular systems, is an extremely weak electrolyte; but it is still dissociated into ions, according to H20 H+ + OH" (3) As the H+ ion cannot exist as free proton in aqueous media, the equation should be properly written with an oxonium ion in place of H+. Thus, 2H20 ^ H30 + + OH" (4) Even in an acid solution of pH = 0 the concentration of free protons was shown to be about IQ-13° !2 However, this way of expressing the hydrogen ion concentration does not reflect its actual solvation degree in water which, in fact, shows highest stability for H30 +*3H20 and for 0H -3H20. From Equation 4, then, the dissociation constant of water is = aH30+ * a0H~/aH20 = CH30 + 7h30 + * COH"7oH_/CH2o7h20 (^) The dissociation of water being negligible we may consider the concentration of H20 as constant and the activity coefficients equal to unity. The concentration of water itself is 55.35 mol dm-3 at 25°C and it is generally included in what is termed the ionization constant (ion product) of water, or (6) K w — 0h3O+ * aOH~ 3 T/*C FIGURE 1. Temperature dependence of the ionization con­ stant (ion product) of water. The value of Kw is almost exactly 10_ 14 mol dm ~ 3 at 25°C and depends on temperature as shown in Figure 1. The degree of dissociation of water a (= cH30+/cHl0 = c0H-/cH20) at 25°C is, thus, 1.4M0-9. Because of the temperature dependence, neutrality (identical number of plus and minus charges in solution) is achieved at 6.81 at human body temperature. For the sake of comparison, it should be noted that the self-ionization (autoprotolysis) constants of protogenic solvents (strong acids) are high (2.69 • 10-4M for sulfuric acid), but much like water for weaker acids or alcohols (3.16 • 10~ 15M for acetic acid and 2.51 * 10_17M for ethanol), and are extremely low for protophilic solvents (bases), e.g., 10_29M for ammonia. II. IONIZATION OF ACIDS AND BASES Although water is by far the most abundant component of all living systems, its dissociation into oxonium ions is so weak that the pH of a salt solution, either extra- or intracellular, is determined by the presence of components that readily dissociate or readily bind an oxonium ion, i.e., acids and bases, respectively. Throughout modem electrochemistry, three theories of acids and bases came into prom­ inence. The first theory, that of Arrhenius3 dates back to 1887 when he postulated the universal existence of dissociation of electrolytes in solution, supporting his views by conductometric measurements. He calculated the degree of dissociation a from the ratios of equivalent conductivities at a given and at an infinite dilution. Thus, a = A/A^ It was Arrhenius who defined acids and bases in a simple way, stating that an acid (HA) is characterized by dissociation of hydrogen ions HA H+ + A- (7a) while a base (BOH) is recognized as a substance dissociating hydroxide ions BOH B+ + OH- (7b)