How do buffer solutions maintain the pH of blood?

1 Answer
Dec 26, 2014

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

H+(aq)+HCO3(aq)H2CO3(aq)H2O(l)+CO2(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

H2CO3(aq)+H2O(l)HCO3(aq)+H3O+(aq) - (1) an acid-base reaction, has an equilibrium constant K1;

H2CO3(aq)+H2O(l)CO2(g)+2H2O(l) - (2) carbonic acid dissociates rapidly to produce water and CO2 - equilibrium constant K2

For the first reaction, carbonic acid (H2CO3) is the weak acid and the bicarbonate ion (HCO3) is its conjguate base.

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

pH=pKlog([CO2][HCO3]), where K=K1K2.

So, the blood's pH depends on the ratio between the amount of CO2 present in the blood and the amount of HCO3 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 HCO3 decreases (relative to the amount of CO2); however, this change is small compared to the amount of HCO3 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 CO2 from the blood this increases the pH;
  • The kidneys remove excess HCO3 from the body this lowers the pH.

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