Henderson–Hasselbalch equation

In chemistry and biochemistry, the pH of weakly acidic chemical solutions can be estimated using the Henderson-Hasselbalch Equation: $${\displaystyle {\ce {pH}}={\ce {p}}K_{{\ce {a}}}+\log _{10}\left({\frac {[{\ce {Base}}]}{[{\ce {Acid}}]}}\right)}$$

The equation relates the pH of the weak acid to the numerical value of the acid dissociation constant, K_a, of the acid, and the ratio of the concentrations of the acid and its conjugate base.

Acid-base Equilibrium Reaction

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

The Henderson-Hasselbalch equation is often used for estimating the pH of buffer solutions by approximating the actual concentration ratio as the ratio of the analytical concentrations of the acid and of a salt, MA. It is also useful for determining the volumes of the reagents needed before preparing buffer solutions, which prevents unncessary waste of chemical reagents that may need to be further neutralized by even more reagents before they are safe to expose.

For example, the acid may be carbonic acid

$${\displaystyle {\ce {HCO3-}}+\mathrm {H^{+}} \rightleftharpoons {\ce {H2CO3}}\rightleftharpoons {\ce {CO2}}+{\ce {H2O}}}$$

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

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

The Henderson–Hasselbalch buffer system also has many natural and biological applications, from physiological processes (e.g., metabolic acidosis) to geological phenomena.