There are two useful rules of thumb for selecting buffer mixtures:
A good buffer mixture should have about equal concentrations of both of its components.
A buffer solution has generally lost its usefulness when one component of the buffer pair is less than about 10% of the other. The graph below shows how pH changes for an acetic acid-acetate ion buffer as base is added. The initial pH is 4.74. A change of 1 pH unit occurs when the acetic acid concentration is reduced to 11% of the acetate ion concentration.
Change in pH as an increasing amount of a 0.10-M NaOH solution is added to 100 mL of a buffer solution in which, initially, [CH3CO2H] = 0.10 M and [CH3CO2–] = 0.10 M. [CH3CO2– ]= 0.10 M.
Note the greatly diminished buffering action occurring after the buffer capacity has been reached, resulting in drastic rises in pH on adding more strong base.
Weak acids and their salts are better as buffers for pHs less than 7; weak bases and their salts are better as buffers for pHs greater than 7.
Blood is an important example of a buffered solution, with the major acid and ion responsible for the buffering action being carbonic acid, H2CO3, and the bicarbonate ion, $HCO_3^-$.
When a hydronium ion is introduced to the blood stream, it is removed primarily by the reaction:
$$H_3O^+(aq)+HCO_3^-(aq)⟶H_2CO_3(aq)+H_2O(l)$$
An added hydroxide ion is removed by the reaction:
$$OH^-(aq)+H_2CO_3(aq)⟶HCO_3^-(aq)+H_2O(l)$$
The added strong acid or base is thus effectively converted to the much weaker acid or base of the buffer pair (H3O+ is converted to H2CO3 and OH– is converted to HCO3–). The pH of human blood thus remains very near the value determined by the buffer pairs pKa, in this case, 7.35. Normal variations in blood pH are usually less than 0.1, and pH changes of 0.4 or greater are likely to be fatal.