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
Ka= [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).
