buffer

A solution that resists change in pH when small amounts of an acid or alkali are added over a cer­tain range or when the solution is di­luted. Acidic buffers consist of a weak acid with a salt of the acid. The salt provides the negative ion A-, which is the conjugate base of the acid HA. An example is carbonic acid and sodium hydrogencarbonate. Basic buffers have a weak base and a salt of the base (to provide the conjugate acid). An example is am­monia solution with ammonium chloride.

In an acidic buffer, for example, molecules HA and ions A- are pre­sent. When acid is added most of the extra protons are removed by the base:

A^– + H⁺ → HA

When base is added, most of the extra hydroxide ions are removed by reaction with undissociated acid:

OH^– + HA → A^– + H₂O

Thus, the addition of acid or base changes the pH very little. The hy­drogen-ion concentration in a buffer is given by the expression

K_a= [H⁺] = [A^–]/[HA]

i.e. it depends on the ratio of conju­gate base to acid. As this is not al­tered by dilution, the hydrogen-ion concentration for a buffer does not change much during dilution.

In the laboratory, buffers are used to prepare solutions of known stable pH. Natural buffers occur in living or­ganisms, where the biochemical re­actions are very sensitive to change in pH. The main natural buffers are H₂CO₃/HCO₃^– and H₂PO₄^–/HPO₄^–. Buffer solutions are also used in medicine (e.g. in intravenous injec­tions), in agriculture, and in many industrial processes (e.g. dyeing, fermentation processes, and the

food industry).