Henderson–Hasselbalch equation

In chemistry and biochemistry, the Henderson–Hasselbalch equation

relates the pH of a chemical solution of a weak acid to the numerical value of the acid dissociation constant, K_a, of acid and the ratio of the concentrations, ${\displaystyle {\frac {[{\ce {Base}}]}{[{\ce {Acid}}]}}}$ of the acid and its conjugate base in an equilibrium.

${\displaystyle \mathrm {{\underset {(acid)}{HA}}\leftrightharpoons {\underset {(base)}{A^{-}}}+H^{+}} }$

For example, the acid may be acetic acid

${\displaystyle \mathrm {CH_{3}CO_{2}H\leftrightharpoons CH_{3}CO_{2}^{-}+H^{+}} }$

The Henderson–Hasselbalch equation can be used to estimate the pH of a buffer solution by approximating the actual concentration ratio as the ratio of the analytical concentrations of the acid and of a salt, MA.

The equation can also be applied to bases by specifying the protonated form of the base as the acid. For example, with an amine, ${\displaystyle \mathrm {RNH_{2}} }$

${\displaystyle \mathrm {RNH_{3}^{+}\leftrightharpoons RNH_{2}+H^{+}} }$