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 New York
`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
`Reprinted 1979, 1983 twice
`
`Published in the USA by
`Chapman and Hall,
`733 Third Avenue, New York NY 10017
`
`© 1974 D.D. Perrin and Boyd Dempsey
`
`Printed in Great Britain by
`J. W. Arrowsmith Ltd, Bristol
`
`ISBN O 412 21890 9
`
`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 permission in writing
`from the publisher
`
`Contents
`
`Preface
`1.
`Introduction
`1.1 The concept of buffer action
`1.2 Why are buffers needed?
`1.3 Some naturally occurring buffers
`2.
`The Theory of Buffer Action
`2.1 Equilibrium aspects
`2.2 Activity effects
`2.3 Effect of dilution
`2.4 Salt effects
`2.5 Arnpholytes and zwitterions
`2.6 Buffer capacity
`2.6.1
`Buffer capacity of a polybasic acid
`2.7 Pseudo buffers
`2.8 Self buffers
`2.9 Mixtures of buffers
`2.10 Temperature dependence
`2.11 Effect of pressure on buffers
`2.12 Further reading
`3.
`Applications of pH Buffers
`3.1 Factors governing the choice of a buffer
`3.2 Measurement of pH
`3.3 Biochemistry and biology
`3.4 Spectroscopy
`3.5 Buffers for special applications
`3.5.1
`Volatile buffers
`3.5.2 Buffers for electrophoresis
`3.5.3 Buffers for complexometric titrations
`3.5.4 Buffers for chromatography
`
`page viii
`1
`1
`2
`3
`4
`4
`6
`7
`8
`1-0
`10
`12
`IS
`IS
`17
`18
`18
`19
`24
`24
`25
`27
`32
`33
`33
`34
`34
`35
`
`

`

`16 · Buffers for pH and Metal Ion Control
`
`convenient as pH standards, particularly as the solutions are
`rather insensitive to variations in concentration if the pH lies
`between about 3 and 11.
`A much wider range of self buffers is encompassed when
`partially neutralized salts of substances with two or more
`pKa values are involved, Phthalic acid is dibasic, with pKa
`values at 25°C of 2.95 and 5.41; potassium hydrogen
`phthalate solutions have a pH roughly midway between these
`values (slightly lower because of ionic strength effects) and
`the buffer capacity is as expected for a solution 1.2 pH units
`removed from the pKa. This is the basis of the use of 0.05M
`potassium hydrogen phthalate as the reference standard in
`pH measurement. Another example is potassium hydrogen
`tartrate (pKa values of parent acid, 3.04 and 4.37) which has
`a better buffer capacity but is not very soluble. Sodium
`hydrogen malate (pKa values of parent acid, 3.40 and 5.13)
`and potassium dihydrogen citrate (pKa values of 3 .13 and
`4.76) could also be useful. The pH of a 0.05m KH2 citrate
`solution is 3.68 at 25°C. A 0.2M solution of sodium
`hydrogen diglycollate (pKa values 2.96 and 4.43) has a
`pH= 3.40 ± 0.02 over the temperature range 10-35°C, has a
`slight ~ pHy. and has been proposed as a pH reference buff er
`(Keyworth and Hahn, 1958). On the other hand, the pKa
`separations are too great for a salt such as potassium
`dihydrogen phosphate (pKa values of phosphoric acid are
`2~16, 7.21, 12.33) to be useful as a self buffer.
`Mention has already been made of the self buffering by a
`salt of a weak acid and a weak base, ~ch as ammonium
`acetate (pKa values of 4.76 and 9.25). Irrthis example the pH
`of a solution is near to 7 (R=< ½(PKHo Ac+ pKNH 3 ), but the
`buffer capacity is small because the pH is more than 2 units
`away from either pKa value. (In analytical chemistry this
`disadvantage is partially overcome by using high concentra(cid:173)
`tions}. Ammonium bicarbonate (pKa values of 6.35, 9.25) or
`diammonium hydrogen phosphate (pKa values 7.20, 9.25)
`giving pH values of approximately 7.4 and 8.0, respectively,
`for 0.05M solutions, have much better buffer capacities.
`
`The Theory of Buffer Action · 17
`
`Piperazine phosphate monohydrate, C4 H1 2 N2 HPO4 · H2 0,
`is a very good self buffer (PKa values of piperazine 5.333,
`9.731; relevant pKa value of phosphoric acid 7.198 at 25°C)
`and it has been proposed for use as a pH standard. Measured
`pH values for an 0.02m solution from Oto 50°C are given in
`Table 2.5. For an 0.05m solution the pH values should be
`increased by 0.009.
`Ampholytes (zwitterions) include an extensive range of
`substances that could be used as self buffers. Table 2.6 gives a
`list of isoionic ampholytes proposed for use as buffers in
`protein fractionation in a natural pH gradient (Svensson,
`1962). Solutions of these substances in water have pH values
`close to the listed value of pl and when pl - pK1 is less than
`1.5 they can be considered to be self buffers.
`
`2.9 Mixtures of buffers
`
`The effective buff er range for a weak acid or base is
`approximately from pH= pKa + I to pH = pKa -
`l. When
`two or more buffers are present, the effects are additive so
`that the buffering ability is spread over a wider pH range.
`Examples are Mcllvaine's (1921) citric acid-phosphate mix(cid:173)
`tures for pH 2.6-8 (Table 10.45) and Smith and Smith's
`(1949) piperazine-glycylglycine mixtures for pH 4.4-10.8
`(Table 10.46). The pKa of ethanolamine (9.5) falls con(cid:173)
`veniently between two of the pKa values of phosphoric acid
`(7 .2, 12.3) so that ethanolamine-phosphate mixtures provide
`almost uniform buffer capacity between pH 6.7 and pH 12.8
`(Thies and Kallinich, 1953).
`If a buffer system has several successive pKa values which
`differ by about 2 pH units, approximately linear buffer
`capacity results. This property has been exploited in 'univer(cid:173)
`sal' buffers having high buffer capacities over a wide pH
`range. Britton and Robinson ( 1931) used equirnolar mixtures
`based on seven pKa values of citric, phosphoric, boric and
`diethylbarbituric acids to cover the pH range 2.6-12, as
`listed in Tables 10.47 and 10.48. Coch Frugoni (1957) gave
`
`

`

`'t:~·-~-,;:,,~ .,~~,;~i;f;,:it;ft }~-rz,r·,7; .
`18 • Bu'f.fen ft,r. :~y.i,h.ilM'll1i:itlh eontrdl .
`ta~ies '~~ii>:ifi~;:br~~ ~~~tiA' ,;fal~io1{i~t .ionic s,trengths
`0.00S and ().02 by varying the amount of water added, and
`to I= 0.1; 0.S and 1.0 by addition of sodium chloride.
`2.10 Temperature dependence
`The effect of temperature on the pH of a buff er solution
`depends on the temperature dei,endence of the activity
`coefficient terins and of the pK8 of the buffer species. The
`latter is usually much the more important. For syste~ of the
`type, B/Blr', the effect of temperature on the pKa is given,
`to a first approximation, by
`-d(pK8 ) pK8 - 0.9
`--,-=--=·= ------~-
`(2.22)
`dT
`T
`where Tis in °K (Perrin, 1964). For dications, the equation is.
`-d(pK8 ) pKa
`-=--==-
`(2.23}
`dT
`T
`Around 25° C, for a base such as piperidine (PK a = 11.12),
`the pK9 decreases by 0.034 ~ts per degree. For carboxylic
`acids around ambient temperatures, on the other hand,
`changes in . pKa values are much smaller. Th,e temperature
`coefficients for -phenols are also smaller than for bases having
`comparable i:>Ka values.
`
`2. 11 Effect of pressui'e on buffen
`High pressur~ increases the ionizatjon ..Q!.J.Veak electrolytes,
`by enhancing the solvation of the ions, but the effect "I not
`very great. At 25° C and 3000 atmospheres, p.Ka values are
`decreased as follows: formic acid, by 0.38, acetic acid, by
`0.50, and propjonic acid, by 0.55 (Hamann and Strau§,
`1955). The logarithm of the b!!Sic dissociation constant of
`ammonia (into ammonium ion and. hydroxyl ion) is incteased
`by · 1.14 fQr a ·· press~ increase of 3000 at~ospheres
`(Buch~an and Ham~, 1953) and ~Y 2.72 for a ~~ure
`
`

`

`132 • Buffers for pH and Metal Ion Control
`Table 10.13 Formic acid, sodium formate buffer(0°C, 25°C}*
`Contains x ml of M formic acid and y ml M NaOH, diluted to 100 ml
`
`I= 0.05
`
`I= 0.1
`
`[=0.2
`
`0°C 25°C
`X
`X
`
`pH
`
`0°C
`
`25°C
`
`0°C
`
`25°C
`
`y
`
`X
`
`X
`
`y
`
`X
`
`X
`
`y
`
`---2
`
`0.0
`20.0
`20.0
`20.0
`20.0
`20.0
`20.0
`20.0
`20.0
`
`-
`-
`-
`-
`-
`2.6 74.3 68.4 5.0
`2.8 47.8 44.2 5.0 45.9t 84.4 10.0 -
`-
`-
`-
`56.6 10.0
`3.0 31.7 29.4 5.0 61.3
`39.3 10.0 41.6t 77.3
`3.2 21.7 20.3 5.0 42.2
`56.1
`28.4 10.0 59.8
`3.4 15.5 14.6 5.0 30.3
`21.6 10.0 45.1 · 42.8
`3.6 11.6 11.0 5.0 22.8
`34.3
`17.3 10.0 35.8
`8.8 5.0 18.0
`9.1
`3.8
`29.1
`14.6 10.0 30.0
`7.4 5.0 15. l
`7.6
`4.0
`25.7
`12.9 10.0 26.3
`6.5 5.0 13.2
`6.6
`4.0
`23.6
`11.8 10.0 24.0
`6.0 5.0 12.0
`6.0
`4.4
`22.3
`11.1 10.0 22.5
`5.6 5.0 11.3
`5.7
`4.6
`21.4
`10.7 10.0 21.6
`5.4 5.0 10.8
`5.4
`4.8
`
`t 2M solution
`
`*Long (196 l)
`Table 10.14 Citric acid, sodium citrate buffer
`(23°C)*
`Contains x ml 0.lM citric acid (21.01 g C6 H8 O7 -
`) and (50-x) ml 0.lM Na3 citrate (29.41
`H2O1- 1
`g C6 H5 O7 Na 3··2H2O 1-1 ), diluted to 100 ml
`
`pH
`
`3.0
`3.2
`3.4
`3.6
`3.8
`4.0
`4.2
`4.4
`4.6
`
`X
`
`46.5
`43.7
`40.0
`37.0
`35 .0
`33.0
`31.5
`28.0
`25.5
`
`pH
`
`4.8
`5.0
`5.2
`5.4
`5.6
`5.8
`6.0
`6.2
`
`X
`
`23.0
`20.5
`18.0
`16.0
`13.7
`11.8
`9.5
`7.2
`
`*Gomori (1955)
`For disodium hydrogen citrate, NaOH, HCI buff(cid:173)
`ers covering the pH range 2.2-6.8, see Sorensen
`(1909, 1912)
`
`

`

`Appendix. I I • 153
`
`Table 10.45 Wide range buffer(pH 2.6-8.0) (21°C)*
`Contains x ml 0.1 M citric acidt and (100 - x) ml
`0.2M Na2HPO4t.
`.
`
`ti
`
`13
`
`--0
`
`.061
`0.066
`0.066
`0.061
`0.0S6
`0.0S0
`0.047
`0.046
`0.046
`0.044
`0.042
`0.039
`0.036
`0.036
`0.036
`0.036
`0.036
`0.037
`0.040
`0.048
`0.060
`0.076
`0.092
`0.101
`0.100
`0.088
`0.071
`0.053
`
`/§
`
`-
`-
`0.044
`0.062
`0.081
`0.100
`0.12
`0.135
`0.1S
`0.17
`0.19
`0.205
`0.22
`0.24
`0.25
`0.27
`0.28
`0.30
`0.31
`0.32
`0.34
`0.35
`0.36
`0.37
`0.40
`0.43
`0.46
`0.49
`0.52
`o.ss
`
`pH
`
`2.2
`2.4
`2.6
`2.8
`3.0
`3.2
`3.4
`3.6
`3.8
`4.0
`4.2
`4.4
`4.6
`4.8
`5.0
`5.2
`5.4
`S.6
`S.8
`6.0
`6.2
`6.4
`6.6
`6.8
`7.0
`7:2
`7.4
`7.6
`7.8
`8.0
`
`X
`
`98.0
`94.9
`90.3
`85.8
`81.1
`76.6
`72.4
`68.7
`6S.2
`61.9
`59.0
`5(,.3
`53.8
`51.4
`49.0
`47.0
`44.8
`42.6
`40.2
`37.5
`34.6
`31.1
`27.1
`22.8
`17.8
`13.0
`9.4
`6.5
`4.2
`2.8
`
`*Mcilvaine (1921); Whiting (1966).
`t21.0l g C6HsO7·H2O1- 1 .
`t35.6 lg Na2 HPO4 ·2H2O or 28.40 g Na2 HPO4 1-1 _
`§ To adjust these buffers to a constant ionic strength of 0.5
`to 1.0, by adding KCl, see Elving et al., (1956).
`
`