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 zwitter ion. 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.
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
- Hydrogencarbonate ions, HCO3-
- Haemoglobin and plasma proteins
- Dihydrogenphosphate (H2PO4-) and hydrogenphosphate (HPO42-) ion.
- ↑ SlideShare. Buffer in the blood. 2013 [cited 24/11/2018]; Available from: https://www.slideshare.net/tohapras/buffer-in-the-blood