Buffer solutions

Introduction

The pH of a solution is defined as the logarithm to base 10 of 1 divided by the concentration of the free hydrogen ions in solution (i.e. pH = log ₁₀ 1/[A ⁺ ] = −log ₁₀ [H ⁺ ]). A neutral solution is defined as pH = −log ₁₀ [10 ⁻⁷ ] = 7. The pH may greatly affect many chemical and immunohistochemical reactions, and consequently it is frequently important to minimize large changes in free hydrogen ion content, i.e. stabilize the pH.

Buffers are typically solutions in which the addition of small quantities of acids or bases causes little or no change in the pH of the solution. In other words, the solution ‘buffers’ against a change in pH. This is accomplished by solutions of inorganic and organic acids or bases plus salts, which together absorb free hydrogen or free hydroxyl ions to prevent major changes in pH. Several major buffer systems are used in histochemical and/or immunohistochemical staining. Buffer systems include citric acid, sodium citrate, acetic acid-sodium acetate and mixtures of sodium or potassium phosphates. One frequently used system is based on the use of Tris(hydroxylmethyl)aminomethane, often called ‘Tris’. Tris buffer systems include Tris-maleic acid. Tris buffers are susceptible to temperature changes, so pH values specified at multiple temperatures are shown. The following buffer tables are the primary buffers referred to in this edition. For any buffers required and not listed here, the reader is referred to and , or to suitable biochemical texts.

General notes regarding buffer solutions

The salts and acids used in the preparation of buffers should be of at least laboratory reagent grade. When preparing buffers, the molecular weight given on the reagent bottle should be checked, as many chemicals are available in a number of states of hydration.