Bicarbonate buffering system

File:2325 Carbon Dioxide Transport.jpg

Carbon dioxide, a product of metabolic function, is dissolved in the blood, where it is taken up by red blood cells and converted to carbonic acid by carbonic anhydrase. Most of the carbonic acid then dissociates to bicarbonate and hydrogen ions.

The bicarbonate buffering system is an acid-base homeostatic mechanism involving the balance of carbonic acid (H₂CO₃), bicarbonate ion (HCO₃^-), and carbon dioxide (CO₂) in order to maintain pH in the blood and duodenum, among other tissues, to support proper metabolic function. Catalyzed by carbonic anhydrase, carbon dioxide (CO₂) reacts with water (H₂O) to form carbonic acid (H₂CO₃), which in turn rapidly dissociates to form a hydrogen ion (present in solution as hydronium ion, H₃O⁺) and bicarbonate ion (HCO₃⁻ ) as shown in the following reaction:

$\rm CO_2 + H_2O \rightleftarrows H_2CO_3 \rightleftarrows HCO_3^- + H_3O^+$

As with any buffer system, the pH is balanced by the presence of both weak acid (i.e. H₂CO₃) and conjugate base (i.e. HCO₃^-) so that any excess acid or base introduced to the system is neutralized.

Failure of this system to function properly results in acid-base imbalance such as acidemia (pH<7.35) and alkalemia (pH>7.45) in the blood. In the duodenum, failure to protect epithelial cells with bicarbonate ion results in ulcers, as seen in H. Pylori infection.

Physiological Mechanism

In tissue, cellular respiration produces carbon dioxide as a waste product; as one of the primary roles of the cardiovascular system, most of this CO₂ is rapidly removed from the tissues by its hydration to bicarbonate ion. The bicarbonate ion present in the blood plasma is transported to the lungs, where it is dehydrated back into CO₂ and released during exhalation. These hydration and dehydration conversions of CO₂ and H₂CO₃, which are normally very slow, are facilitated by carbonic anhydrase in both the blood and duodenum. While in the blood, bicarbonate ion serves to neutralize acid introduced to the blood through other metabolic processes (e.g. lactic acid, ketone bodies); likewise, any bases (e.g. urea from the catabolism of proteins) are neutralized by carbonic acid.

The bicarbonate buffering system plays a vital role in other tissues as well. In the human stomach and duodenum, the bicarbonate buffering system serves to both neutralize gastric acid and stabilize the intracellular pH of epithelial cells via the secretion of bicarbonate ion into the gastric mucosa.

Regulation

As calculated by the Henderson-Hasselbalch equation, in order to maintain a normal pH of 7.4 in the blood (whereby the pKa of carbonic acid is 6.1 at physiological temperature), a 20:1 bicarbonate to carbonic acid must constantly be maintained; this homeostasis is mainly mediated by the respiratory and renal systems. In the blood of most animals, the bicarbonate buffering system is coupled to the lungs via respiratory compensation, the process by which the rate of breathing changes to compensate for changes in the blood concentration of CO₂. By Le Chȃtlier’s Principle, the release of CO₂ from the lungs pushes the reaction above to the left, causing carbonic anhydrase to form CO₂ until all excess acid is removed. Bicarbonate concentration is also further regulated by renal compensation, the process by which the kidneys regulate the concentration of bicarbonate ion by filtering out excess or retaining it when in low concentrations.

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:

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

, where:

• pK_{a H2CO3} is the negative logarithm (base 10) of the acid dissociation constant of carbonic acid. It is equal to 6.1.
• [HCO₃⁻] is the concentration of bicarbonate in the blood
• [H₂CO₃] is the concentration of carbonic acid in the blood

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:

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

, where:

• [H₂CO₃] is the concentration of carbonic acid in the blood
• k_{H CO2} is a constant including the solubility of carbon dioxide in blood. k_{H CO2} is approximately 0.03 (mmol/L)/mmHg
• pCO₂ is the partial pressure of carbon dioxide in the blood

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:

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

, where:

• pH is the acidity in the blood
• [HCO₃⁻] is the concentration of bicarbonate in the blood, in mmol/L
• pCO₂ is the partial pressure of carbon dioxide in the blood, in mmHg

Derivation of The Kassirer-Bleich Approximation

The Henderson Equation, which is derived from the Law of Mass Action, can be modified with respect to the bicarbonate buffering system to yield a simpler equation that provides a quick approximation of the H⁺ or HCO₃^- concentration without the need to calculate logarithms:

$\rm K_{a,H_2CO_3} = \frac{[HCO_3^-][H_3O^+]}{[H_2CO_3]}$

Since the partial pressure of carbon dioxide is much easier to obtain from measurement than carbonic acid, the Henry’s Law solubility constant – which relates the partial pressure of a gas to its solubility – for CO₂ in plasma is used in lieu of the carbonic acid concentration. After rearranging the equation and applying Henry’s Law, the equation becomes:

$\rm [H^+] = \frac{K'\cdot0.03P_{CO_2}}{[HCO_3^-]}$

Where K’ is the dissociation constant from the pK_a of carbonic acid, 6.1, which is equal to 800nmol/L (since K’ = 10^-pKa = 10^-(6.1) ≈ 8.00X10⁻⁰⁷mol/L = 800nmol/L).

By multiplying K’ (expressed as nmol/L) and 0.03 (800 X 0.03 = 24) and rearranging with respect to HCO₃^-, the equation is simplified to:

$\rm [HCO_3^-] = 24\frac{P_{CO_2}}{[H^+]}$

External links

• Physiology: 7/7ch12/7ch12p17 - Essentials of Human Physiology