Which expression represents the Henderson-Hasselbalch equation as applied to blood pH?

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Multiple Choice

Which expression represents the Henderson-Hasselbalch equation as applied to blood pH?

Explanation:
The key idea is that blood pH is set by the carbonic acid–bicarbonate buffering system and follows Henderson–Hasselbalch: pH = pKa + log([base]/[acid]). For this system, the base is bicarbonate (HCO3−) and the acid is carbonic acid (H2CO3). In plasma, the amount of H2CO3 is approximated by the dissolved CO2, which equals 0.03 × PCO2 (with PCO2 in mmHg). The pKa for carbonic acid at body temperature is about 6.1. Putting these together yields pH = 6.1 + log([HCO3−]/(0.03 × PCO2)). This is why the expression with HCO3− in the numerator and 0.03 × pCO2 in the denominator is correct. Using just pCO2 in the denominator would ignore the conversion from partial pressure to the dissolved CO2 concentration, which is essential for the buffering relationship. Reversing the ratio or omitting the 0.03 factor would also misrepresent the balance between bicarbonate and dissolved CO2, leading to an incorrect pH prediction.

The key idea is that blood pH is set by the carbonic acid–bicarbonate buffering system and follows Henderson–Hasselbalch: pH = pKa + log([base]/[acid]). For this system, the base is bicarbonate (HCO3−) and the acid is carbonic acid (H2CO3). In plasma, the amount of H2CO3 is approximated by the dissolved CO2, which equals 0.03 × PCO2 (with PCO2 in mmHg). The pKa for carbonic acid at body temperature is about 6.1. Putting these together yields pH = 6.1 + log([HCO3−]/(0.03 × PCO2)). This is why the expression with HCO3− in the numerator and 0.03 × pCO2 in the denominator is correct.

Using just pCO2 in the denominator would ignore the conversion from partial pressure to the dissolved CO2 concentration, which is essential for the buffering relationship. Reversing the ratio or omitting the 0.03 factor would also misrepresent the balance between bicarbonate and dissolved CO2, leading to an incorrect pH prediction.

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