Bicarbonate buffer

The "bicarbonate buffer system" (also "bicarbonate") is the most important and effective buffering system in the body. Especially in the blood, where it accounts for up to 53% buffering capacity. Its importance lies in its good ability to maintain a stable pH mainly due to the fact that the concentration of both components can change independently - CO_{2</ sub>} by breathing, HCO₃^- by kidney and liver activity. Therefore, the bicarbonate buffer in the body is referred to as an open buffering system.

Bicarbonate buffer titration curve

The reaction proceeds as follows: CO₂ + H₂O $\rightleftharpoons$ HCO₃^- + H< sup>+.

Buffers composed of weak acids and their salts (or weak bases and their salts) with the same substance concentration have the greatest buffering capacity, i.e., more precisely, for which pH = pK_A'. The optimal blood pH value is 7.4' ± 0.04. The pK_A value for bicarbonate buffer is 6.1. Thus, it appears that this buffer will not buffer pH fluctuations very well. But the opposite is true. For a better idea, here is an example:

To Henderson-Hasselbalch equation

$pH=pK_{A}+\log {\frac {[HCO_{3}^{-}]}{[CO_{2}]}}$

we substitute physiological concentrations

HCO₃^- = 24 mmol/l and CO₂ 1.2 mmol/l.

(Ratio of base to acid is therefore 20:1.)

$pH=6.1+\log {\frac {24mmol/l}{1.2mmol/l}}$     Resulting pH = 7.4 .

In the case of a closed system, after the addition of H⁺, conjugated acid CO₂ is formed, which cannot escape from the system, and thus its concentration rises. An increase in the concentration of CO₂ by 2 mmol/l is reciprocally balanced by a decrease in the concentration of HCO₃^-.

$pH=6.1+\log {\frac {22mmol/l}{3.2mmol/l}}$     Resulting pH = 6 ,93.   (In this case, the buffering capacity of the buffer is very small, because the value of 6.93 is quite far from 7.4.)

However, if the resulting CO₂ is removed (exhaled) from the system, as is the case with the bicarbonate open system, only the concentration changes with the addition of H⁺ HCO₃^-. The ratio of HCO₃^- to CO₂, and thus the pH value, will shift much less.

$pH=6.1+\log {\frac {22mmol/l}{1.2mmol/l}}$     Resulting pH = 7 ,36.

Summary: An increase in H⁺ in the blood leads to the production of CO₂, which is soon exhaled in the lungs, allowing a constant pCO_{2< /sub>, i.e. a concentration of 1.2 mmol/l.}

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Related Articles

• Buffers
• Henderson-Hasselbalch equation
• pH
• Protein Buffer System
• Buffer systems

References

• LEDVINA, M. Biochemie pro studující medicíny II. 2. edition. Nakladatelství Karolinum, 2009. 281 pp. ISBN 978-80-246-1415-1.