pKa and pKb from Constants

pK_a and pK_b constants

Take a conjugate acid–base pair sitting in water at 25 °C. Its dissociation constants K_a and K_b obey K_a·K_b = K_w = 10⁻¹⁴, which is where pK_a + pK_b = 14 comes from. A smaller pK_a means a stronger acid, and a smaller pK_b means a stronger base. Some reference values to anchor this: HCl pK_a ≈ −7 (very strong acid), HF pK_a = 3.17, acetic acid (CH₃COOH) pK_a = 4.76, HCN pK_a = 9.21, water pK_a = 15.7. As an example, acetate is the conjugate base of acetic acid, so its pK_b = 14 − 4.76 = 9.24, which tells you acetate is a very weak base. Because of this relationship you can work out one constant from the other without having to measure both.

Applications

This is the standard reference table analytical chemists reach for when ranking acid/base strength. It also drives buffer design, where you pick a system whose pK_a lands within ±1 of the target pH. In pharmaceutical drug delivery the drug's pK_a sets how much of it ionizes at gastric or intestinal pH, which in turn shapes absorption. And it helps you predict titration end-points and pick the right indicator.

FAQ

Does pK_a + pK_b = 14 always hold? Only in water at 25 °C. K_w shifts with temperature (it's about 10⁻¹³·⁶ at 50 °C), so the sum drifts a little.

Why is the pK_a of HCl negative? In water HCl dissociates completely, so K_a is enormous (≈ 10⁷) and taking its log lands you on a big negative number. For strong acids and bases the pK_a values you see are estimated or extrapolated.

How do I choose a buffer? Go for a weak acid whose pK_a sits close to the pH you want, within ±1. Buffering capacity peaks right at pH = pK_a.