# How do buffer solutions maintain the pH of blood?

Dec 26, 2014

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

${H}_{\left(a q\right)}^{+} + H C {O}_{3 \left(a q\right)}^{-} r i g h t \le f t h a r p \infty n s {H}_{2} C {O}_{3 \left(a q\right)} r i g h t \le f t h a r p \infty n s {H}_{2} {O}_{\left(l\right)} + C {O}_{2 \left(g\right)}$

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

${H}_{2} C {O}_{3 \left(a q\right)} + {H}_{2} {O}_{\left(l\right)} r i g h t \le f t h a r p \infty n s H C {O}_{3 \left(a q\right)}^{-} + {H}_{3} {O}_{\left(a q\right)}^{+}$ - (1) an acid-base reaction, has an equilibrium constant ${K}_{1}$;

${H}_{2} C {O}_{3 \left(a q\right)} + {H}_{2} {O}_{\left(l\right)} r i g h t \le f t h a r p \infty n s C {O}_{2 \left(g\right)} + 2 {H}_{2} {O}_{\left(l\right)}$ - (2) carbonic acid dissociates rapidly to produce water and $C {O}_{2}$ - equilibrium constant ${K}_{2}$

For the first reaction, carbonic acid (${H}_{2} C {O}_{3}$) is the weak acid and the bicarbonate ion ($H C {O}_{3}^{-}$) is its conjguate base.

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

$p H = p K - \log \left(\frac{\left[C {O}_{2}\right]}{\left[H C {O}_{3}^{-}\right]}\right)$, where $K = {K}_{1} / {K}_{2}$.

So, the blood's pH depends on the ratio between the amount of $C {O}_{2}$ present in the blood and the amount of $H C {O}_{3}^{-}$ 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 $H C {O}_{3}^{-}$ decreases (relative to the amount of $C {O}_{2}$); however, this change is small compared to the amount of $H C {O}_{3}^{-}$ 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 $C {O}_{2}$ from the blood $\to$ this increases the pH;
• The kidneys remove excess $H C {O}_{3}^{-}$ from the body $\to$ this lowers the pH.

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