How do buffer solutions maintain the pH of blood?

The most important buffer for maintaining the blood's acid-base balance is the carbonic acid - bicarbonate buffer.

#H_((aq))^(+) + HCO_(3(aq))^(-) rightleftharpoons H_2CO_(3(aq)) rightleftharpoons H_2O_((l)) + CO_(2(g))#

SInce pH is determined by the concentration of #H^(+)#, let's try and determine a relationship between the concentrations of all the species involved in this reaction. The two ractions that take place are

#H_2CO_(3(aq)) + H_2O_((l)) rightleftharpoons HCO_(3(aq))^(-) + H_3O_((aq))^(+)# - (1) an acid-base reaction, has an equilibrium constant #K_1#;

#H_2CO_(3(aq)) + H_2O_((l)) rightleftharpoons CO_(2(g)) + 2 H_2O_((l))# - (2) carbonic acid dissociates rapidly to produce water and #CO_2# - equilibrium constant #K_2#

For the first reaction, carbonic acid (#H_2CO_3#) is the weak acid and the bicarbonate ion (#HCO_3^(-)#) is its conjguate base.

Using the Henderson-Hasselbach equation, and without going through the entire derivation, the pH can be written as

#pH = pK - log(([CO_2])/([HCO_3^-]))#, where #K = K_1/K_2#.

So, the blood's pH depends on the ratio between the amount of #CO_2# present in the blood and the amount of #HCO_3^(-)# present in the blood. Since the concentrations of both buffer components are very large, the pH will remain unchanged when #H^(+)# is added to the blood.

When #H^(+)# is added to the blood as a result of a metabolic process, the amount of #HCO_3^(-)# decreases (relative to the amount of #CO_2#); however, this change is small compared to the amount of #HCO_3^(-)# present in the blood. Optimal buffering takes place when the pH is between 5.1 and 7.1.

When too much protons are added to the blood, the buffer system gets a little help from the lungs and the kidneys:

• The lungs remove excess #CO_2# from the blood #-># this increases the pH;
• The kidneys remove excess #HCO_3^(-)# from the body #-># this lowers the pH.

  Here's a nice video detailing the carbonic acid - bicarabonate ion buffer system: