The most important buffer for maintaining the blood's acid-base balance is the carbonic acid - bicarbonate buffer.
H+(aq)+HCO−3(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)⇌HCO−3(aq)+H3O+(aq) - (1) an acid-base reaction, has an equilibrium constantK1;
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 (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([CO2][HCO−3]), 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 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 CO2); 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 CO2 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: