# Bicarbonate buffering system

The bicarbonate buffering system is an important buffer system in the acid-base homeostasis of living things, including humans. As a buffer, it tends to maintain a relatively constant plasma pH and counteract any force that would alter. In this system, carbon dioxide (CO2) combines with water (H2O) to form carbonic acid (H2CO3), which in turn rapidly dissociates to form hydrogen ions (H+) and bicarbonate (HCO3- ) as shown in the reactions below.

$\rm CO_2 + H_2O \rightleftarrows H_2CO_3 \rightleftarrows HCO_3^- + H^+$

The carbon dioxide - carbonic acid equilibrium is catalyzed by the enzyme carbonic anhydrase; the carbonic acid - bicarbonate equilibrium is simple proton dissociation/association and needs no catalyst.

Any disturbance of the system will be compensated by a shift in the chemical equilibrium according to Le Chatelier's principle. For example, if one attempted to acidify the blood by dumping in an excess of hydrogen ions (acidemia), some of those hydrogen ions will associate with bicarbonate, forming carbonic acid, resulting in a smaller net increase of acidity than otherwise. This buffering system becomes an even more powerful regulator of acidity when it is coupled with the body's capacity for respiratory compensation, in which breathing is altered to modify the amount of CO2 in circulation. In the above example, increased ventilation would increase the loss of CO2 to the atmosphere, driving the equilibria above to the left. The process could continue until the excess acid is all exhaled.

This process is extremely important in the physiology of blood-having animals. It manages the many acid and base imbalances that can be produced by both normal and abnormal physiology. It also affects the handling of carbon dioxide, the constantly produced waste product of cellular respiration.

## Henderson–Hasselbalch equation

A modified version of the Henderson–Hasselbalch equation can be used to relate the pH of blood to constituents of the bicarbonate buffering system:[1]

$pH = pK_{a~H_2CO_3}+ \log \left ( \frac{[HCO_3^-]}{[H_2CO_3]} \right )$

, where:

This is useful in arterial blood gas, but these usually state pCO2, that is, the partial pressure of carbon dioxide, rather than H2CO3. However, these are related by the equation:[1]

$[H_2CO_3] = k_{\rm H~CO_2}\, \times pCO_2$

, where:

Taken together, the following equation can be used to relate the pH of blood to the concentration of bicarbonate and the partial pressure of carbon dioxide:[1]

$pH = 6.1 + \log \left ( \frac{[HCO_3^-]}{0.03 \times pCO_2} \right )$

, where:

• pH is the acidity in the blood
• [HCO3-] is the concentration of bicarbonate in the blood
• pCO2 is the partial pressure of carbon dioxide in the blood

## References

1. ^ a b c page 556, section "Estimating plasma pH" in: Bray, John J. (1999). Lecture notes on human physiolog. Malden, Mass.: Blackwell Science. ISBN 978-0-86542-775-4.