Strong acids/bases ionise completely; weak ones only partially, setting up an equilibrium. The extent is measured by Ka (acid): larger Ka (smaller pKa) = stronger acid. Acid dissociation constant \ce{HA <=> H+ + A^-}\qquad K_a=\frac{[\ce{H+}][\ce{A^-}]}{[\ce{HA}]} pKa = −log Ka pH of a weak acid (approximation) [\ce{H+}]=\sqrt{K_a\,c} Valid when dissociation is small (< ~5%): assumes [HA] ≈ initial c and [H⁺] ≈ [A⁻]. Ka × Kb relationship K_a\times K_b=K_w Lets you find Kb of the conjugate base from Ka. Worked example pH of a weak acid Find the pH of 0.10 mol L⁻¹ ethanoic acid, Ka = 1.8×10⁻⁵. [H⁺] = √(Ka × c) = √(1.8×10⁻⁵ × 0.10) = √(1.8×10⁻⁶). [H⁺] = 1.34×10⁻³ mol L⁻¹. pH = −log(1.34×10⁻³) = 2.87. Check dissociation: 1.34×10⁻³ / 0.10 = 1.3% (< 5%, so the approximation is valid). ✓ pH = 2.87. ✕
Weak ≠ dilute
A weak acid is one that partially ionises (small Ka); a dilute acid just has a low concentration. A concentrated weak acid and a dilute strong acid can have the same pH for very different reasons.