AS & A-Level Chemistry 13 — Quantitative equilibria and kinetics

Calculating the pH of strong acids requires assuming complete dissociation, so $[\mathrm{H^+}]$ equals the concentration of a monoprotic acid. Strong alkalis also dissociate fully; $[\mathrm{H^+}]$ is found by dividing $K_\mathrm{w}$ by $[\mathrm{OH^-}]$. For weak acids, partial dissociation means $[\mathrm{H^+}]$ must be calculated using $K_\mathrm{a}$, often with the approximation that equilibrium $[\mathrm{HA}]$ equals its initial concentration, yielding $[\mathrm{H^+}] = \sqrt{K_\mathrm{a} [\mathrm{HA}]}$.

A buffer solution is a mixture that resists changes to pH upon the addition of small amounts of strong acid or strong alkali. Acidic buffers are commonly made from a weak acid and its conjugate base (e.g., sodium ethanoate). When acid is added, the conjugate base reacts with the incoming $\mathrm{H^+}$; when alkali is added, the weak acid neutralises the incoming $\mathrm{OH^-}$. An important blood buffer near pH 7.4 is the $\mathrm{H_2CO_3} / \mathrm{HCO_3^-}$ buffer system, operating alongside physiological regulation. Buffer pH calculations use the $K_\mathrm{a}$ expression directly, applying the standard assumption that initial concentrations of acid and salt reliably approximate their equilibrium values. The strong-acid and alkali shortcuts require their ion concentrations to dominate water auto-ionisation. For an isolated weak acid, check that the calculated dissociation is small compared with its initial concentration and that water contributes negligible hydrogen ions. After a buffer reacts with added strong acid or alkali, calculate the remaining acid and conjugate$- \text{base}$ amounts by stoichiometry before applying the buffer expression.