In the Henderson-Hasselbalch equation for blood pH, what is the pKa value for the HCO₃⁻/H₂CO₃ buffer system?

Explanation

## Henderson-Hasselbalch Equation and the Bicarbonate Buffer System ### The Equation $$pH = pKa + \log \frac{[HCO_3^-]}{[H_2CO_3]}$$ For the bicarbonate buffer system: $$pH = 6.1 + \log \frac{[HCO_3^-]}{0.03 \times PaCO_2}$$ **Key Point:** The pKa of the HCO₃⁻/H₂CO₃ buffer system is **6.1**. This is a fixed constant and must be memorized. ### Why pKa = 6.1? The pKa is the negative logarithm of the acid dissociation constant (Ka): $$pKa = -\log(Ka)$$ For the first dissociation of carbonic acid (H₂CO₃ ⇌ H⁺ + HCO₃⁻), the Ka ≈ 7.9 × 10⁻⁷, giving pKa ≈ 6.1. ### Clinical Application At normal blood pH (7.4), the ratio of [HCO₃⁻] to [H₂CO₃] is approximately 20:1, which satisfies the Henderson-Hasselbalch equation: $$7.4 = 6.1 + \log(20) = 6.1 + 1.3$$ **High-Yield:** This 20:1 ratio is the hallmark of normal acid-base balance and is frequently tested. The pKa of 6.1 is a fundamental constant used in all acid-base calculations. **Mnemonic:** "**6.1 is the pKa, 20:1 is the ratio, 7.4 is the pH**" — remember these three numbers for all acid-base problems. ### Why Other Options Are Wrong | Option | Value | Why Wrong | |--------|-------|----------| | 7.4 | Normal blood pH | This is the pH, not the pKa | | 7.35 | Lower normal pH | This is a pH value, not the pKa constant | | 8.3 | pKa of phosphate buffer | This is the pKa for the H₂PO₄⁻/HPO₄²⁻ system, not bicarbonate | **Clinical Pearl:** The bicarbonate buffer system is the most important buffer in blood because both components (HCO₃⁻ and CO₂) are independently regulated — the kidneys control HCO₃⁻ and the lungs control CO₂. This makes it superior to other buffers like phosphate.