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By definition, a buffer is a substance (usually a weak acid and its conjugate base) which is added to a sample to avoid changes in its pH. An example of a buffer would be an amino acid as it has a carboxyl group and an amine group. The name for this type of buffer with both a positive ion and negative ion in its molecule is called a zwitterion. A buffer must contain the chemical species for “neutralizing” added amounts of acid or base. For example, if a buffer was a solution of a weak acid and its conjugate base it would contain acetic acid and sodium acetate or a weak base and conjugate acid it would contain ammonia and ammonium chloride. If there is a lot of acid or alkali added then the buffer will not be able to cope with such a change and will no longer able to maintain the pH.

Within the body, there are three major buffer systems; the carbonic acid-bicarbonate buffer system, the phosphate buffer system and the protein buffer system.

Buffers are most effective in the range PH = pK’a ± 1. Outside the range the concentration of either the acid or the conjugate base is too small to effectively resist the effect of added hydrogen or hydroxide ions.

Examples of buffers that are found in the blood[1].


  1. Lawrie Ryan and Roger Norris. Cambridge International AS and A Level Chemistry Coursebook. 2nd ed. In: Chapter 21: Further aspects of equilibria, page 315. United Kingdom: University Printing House, Cambridge CB2 8BS.
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