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Buffers: Principles and Practice1
Vincent S. Stoll and John S. Blanchard Contents
1. Introduction 2. Theory 3. Buffer Selection 4. Buffer Preparation 5. Volatile Buffers 6. Broad-Range Buffers 7. Recipes for Buffer Stock Solutions References 43 44 45 48 49 50 50 56

1. Introduction
The necessity for maintaining a stable pH when studying enzymes is well established(Good and Izawa, this series; Johnson and Metzler, this series). Biochemical processes can be severely affected by minute changes in hydrogen ion concentrations. At the same time, many protons may be consumed or released during an enzymatic reaction. It has become increasingly important to find buffers to stabilize hydrogen ion concentrations while not interfering with the function of the enzymebeing studied. The development of a series of N-substituted taurine and glycine buffers by Good et al. (1966) has provided buffers in the physiologically relevant range (6.1–10.4) of most enzymes, which have limited side effects with most enzymes (Good et al., 1966). It has been found that these buffers are nontoxic to cells at 50 mM concentrations and in some cases much higher (Ferguson et al.,1980).

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Department of Biochemistry, Albert Einstein College of Medicine, Bronx, New York Reprinted from Methods in Enzymology, Volume 182 (Academic Press, 1990) Copyright # 1990 by Academic Press All rights reserved.

Methods in Enzymology, Volume 463 ISSN 0076-6879, DOI: 10.1016/S0076-6879(09)63006-8

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Vincent S. Stoll and John S. Blanchard

2. Theory
The observation thatpartially neutralized solutions of weak acids or weak bases are resistant to pH changes on the addition of small amounts of strong acid or strong base leads to the concept of ‘‘buffering’’ (Perrin and Dempsey, 1974). Buffers consist of an acid and its conjugate base, such as carbonate and bicarbonate, or acetate and acetic acid. The quality of a buffer is dependent on its buffering capacity(resistance to change in pH by addition of strong acid or base), and its ability to maintain a stable pH upon dilution or addition of neutral salts. Because of the following equilibria, additions of small amounts of strong acid or strong base result in the removal of only small amounts of the weakly acidic or basic species; therefore, there is little change in the pH: HAðacidÞ Ð Hþ þ AÀ ðconjugate baseÞBðbaseÞ þ Hþ Ð BHþ ðconjugate acidÞ ð6:1Þ ð6:2Þ

The pH of a solution of a weak acid or base may be calculated from the Henderson–Hasselbalch equation: pH ¼ pKa0 þ log ½basic speciesŠ ½acidic speciesŠ ð6:3Þ

The pKa of a buffer is that pH where the concentrations of basic and acidic species are equal, and in this basic form the equation is accurate between the pH range of 3–11. Below pH 3 andabove pH 11 the concentrations of the ionic species of water must be included in the equation (Perrin and Dempsey, 1974). Since the pH range of interest here is generally in the pH 3–11 range, this will be ignored. From the Henderson–Hasselbalch equation an expression for buffer capacity may be deduced. If at some concentration of buffer, c, the sum [AÀ] þ [HA] is constant, then the amount ofstrong acid or base needed to cause a small change in pH is given by the relationship: ( ) d½BŠ ½Ka0 c½Hþ Š Kw ð6:4Þ þ ½Hþ Š þ þ ¼ 2:303 H dpH ðKa0 þ ½Hþ ŠÞ2 In this equation, Kw refers to the ionic product of water, and the second and third terms are only significant below pH 3 or above pH 11. In the pH range of interest (pH 3–11), this equation yields the following expression: 2:303c ¼ 0:576c; ð6:5Þ4 which represents a maximum value for d[B]/dpH when pH ¼ pKa. The buffer capacity of any buffer is dependent on the concentration, c, and may bmax ¼

Buffers: Principles and Practice

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0.025

b 0.015 0.005 −1.0

−0.5

0.0 ΔpH

0.5

1.0

Figure 6.1 Buffer capacity (b) versus DpH over the range Æ 1 pH unit of the pKa for HEPES (0.05 M). Points calculated using Eq. (6.5),...