Chemical Equilibrium, Acids, and Bases
Equilibrium
At equilibrium, forward and reverse rates are equal, concentrations are constant, and the reaction is dynamic. The equilibrium constant K expresses products/reactants raised to stoichiometric coefficients; pure solids and liquids are omitted. K > 1 favors products. Le Chatelier's principle predicts shifts: adding reactant shifts right; increasing pressure shifts toward fewer gas moles; heating shifts in the endothermic direction; a catalyst speeds both directions without changing K.
Acids and bases
Arrhenius acids produce H+; Brønsted–Lowry acids donate protons. Strong acids and bases ionize completely; weak ones partially ionize with Ka or Kb. pH = −log[H3O+], pOH = −log[OH−], and pH + pOH = 14 at 25 °C. Buffers resist pH change via a weak acid/base conjugate pair; the Henderson–Hasselbalch equation is pH = pKa + log([A−]/[HA]). Titration equivalence depends on stoichiometry, not pH 7 for weak acids.
Worked example
Find the pH of 0.010 M HCl. HCl is strong, so [H3O+] = 0.010 M = 1.0 × 10^−2 M. pH = −log(1.0 × 10^−2) = 2.00. Then pOH = 12.00 and [OH−] = 1.0 × 10^−12 M.
Traps
Treating weak acids as fully ionized; using pH 7 as the equivalence point for every titration; forgetting that K changes with temperature; and omitting solids from equilibrium expressions.