How do buffer solutions maintain the pH of blood?

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How do buffer solutions maintain the pH of blood?

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Answer 1 · 2014-12-26T12:28:26

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

$H_{(a q)}^{+} + H C O_{3 (a q)}^{-} ⇌ H_{2} C O_{3 (a q)} ⇌ H_{2} O_{(l)} + C O_{2 (g)}$ $H_((aq))^(+) + HCO_(3(aq))^(-) rightleftharpoons H_2CO_(3(aq)) rightleftharpoons H_2O_((l)) + CO_(2(g))$

SInce pH is determined by the concentration of $H^{+}$ $H^(+)$ , 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

$H_{2} C O_{3 (a q)} + H_{2} O_{(l)} ⇌ H C O_{3 (a q)}^{-} + H_{3} O_{(a q)}^{+}$ $H_2CO_(3(aq)) + H_2O_((l)) rightleftharpoons HCO_(3(aq))^(-) + H_3O_((aq))^(+)$ - (1) an acid-base reaction, has an equilibrium constant $K_{1}$ $K_1$ ;

$H_{2} C O_{3 (a q)} + H_{2} O_{(l)} ⇌ C O_{2 (g)} + 2 H_{2} O_{(l)}$ $H_2CO_(3(aq)) + H_2O_((l)) rightleftharpoons CO_(2(g)) + 2 H_2O_((l))$ - (2) carbonic acid dissociates rapidly to produce water and $C O_{2}$ $CO_2$ - equilibrium constant $K_{2}$ $K_2$

For the first reaction, carbonic acid ( $H_{2} C O_{3}$ $H_2CO_3$ ) is the weak acid and the bicarbonate ion ( $H C O_{3}^{-}$ $HCO_3^(-)$ ) is its conjguate base.

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

$p H = p K - log (\frac{[C O_{2}]}{[H C O_{3}^{-}]})$ $pH = pK - log(([CO_2])/([HCO_3^-]))$ , where $K = \frac{K_{1}}{K_{2}}$ $K = K_1/K_2$ .

So, the blood's pH depends on the ratio between the amount of $C O_{2}$ $CO_2$ present in the blood and the amount of $H C O_{3}^{-}$ $HCO_3^(-)$ present in the blood. Since the concentrations of both buffer components are very large, the pH will remain unchanged when $H^{+}$ $H^(+)$ is added to the blood.

When $H^{+}$ $H^(+)$ is added to the blood as a result of a metabolic process, the amount of $H C O_{3}^{-}$ $HCO_3^(-)$ decreases (relative to the amount of $C O_{2}$ $CO_2$ ); however, this change is small compared to the amount of $H C O_{3}^{-}$ $HCO_3^(-)$ 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 $C O_{2}$ $CO_2$ from the blood $\to$ $->$ this increases the pH;
The kidneys remove excess $H C O_{3}^{-}$ $HCO_3^(-)$ from the body $\to$ $->$ this lowers the pH.

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

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How do buffer solutions maintain the pH of blood?

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