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Acid strength is quantified by Ka (acid dissociation constant); base strength by Kb (base dissociation constant). Larger Ka or Kb indicates stronger acid or base. Strong acids and bases ionize completely; weak acids and bases establish equilibrium. Conjugate acid-base pairs are related by Ka × Kb = Kw = 1.0 × 10⁻¹⁴ at 25°C.
Acid strength is not a binary property — it exists on a continuous spectrum captured by the acid dissociation constant Ka. When a weak acid HA dissolves in water, it partially ionizes: HA ⇌ H⁺ + A⁻. The Ka is the equilibrium constant for this reaction: Ka = [H⁺][A⁻] / [HA]. A large Ka means the equilibrium lies far to the right — most of the acid has donated its proton and the acid is strong. A small Ka means the equilibrium lies left — most of the acid remains intact, and only a small fraction has ionized. Strong acids like HCl and HNO₃ have Ka values so large that ionization is essentially complete; weak acids like acetic acid (Ka ≈ 1.8 × 10⁻⁵) ionize only partially.
Working with Ka numerically usually means using logarithms, since Ka values span many orders of magnitude. The pKa = −log(Ka) compresses this range into a more convenient scale: a lower pKa corresponds to a stronger acid (more ionization). For example, acetic acid has pKa ≈ 4.74, while hydrofluoric acid has pKa ≈ 3.17, confirming HF is the stronger acid of the two. When calculating the pH of a weak acid solution, you set up an ICE table (Initial, Change, Equilibrium) and solve the equilibrium expression — often using the approximation that x ≪ initial concentration when Ka is small.
The conjugate base relationship is a critical organizing principle. Every acid HA has a conjugate base A⁻ formed when it donates its proton. The Ka of the acid and the Kb of its conjugate base are linked by Ka × Kb = Kw = 1.0 × 10⁻¹⁴ at 25°C. This means a strong acid (large Ka) always has a weak conjugate base (small Kb), and vice versa. Acetic acid's conjugate base, acetate, has Kb ≈ 5.6 × 10⁻¹⁰ — a weak base, but not negligible. This is why sodium acetate solutions are slightly basic: acetate slowly picks up protons from water.
A common misconception is that Ka directly tells you the pH of a solution without considering concentration. Ka measures ionization tendency, not the resulting H⁺ concentration in a specific solution. A 0.001 M weak acid will have a higher pH than a 1.0 M solution of the same acid even though Ka is identical. The pH depends on both Ka and the initial concentration, which is why the ICE table approach accounts for both.