Equilibrium is the state in which the rates of the forward and reverse processes become equal so that measurable properties stay constant. It is dynamic, possible only in a closed system, and applies to both physical and chemical processes.
Physical equilibria (solid-liquid, liquid-vapour, solid-vapour, dissolution) each have a constant parameter at fixed temperature (e.g. constant vapour pressure, melting point, solubility).
Law of chemical equilibrium / law of mass action: for aA + bB forming cC + dD, Kc = [C]^c[D]^d / [A]^a[B]^b. Kc for the reverse reaction is 1/Kc; multiplying the equation by n raises Kc to the power n. For gaseous reactions Kp = Kc(RT)^dng where dng = (moles of gaseous products) - (moles of gaseous reactants). Pure solids and liquids are omitted in heterogeneous equilibria. A large K (greater than 10^3) means products predominate; a small K (less than 10^-3) means reactants predominate. The reaction quotient Q predicts direction: Q less than K (forward), Q greater than K (reverse), Q = K (equilibrium). K relates to thermodynamics by dG = -RT ln K.
Le Chatelier's principle: a system at equilibrium shifts to counteract any change. Adding a reactant or removing a product shifts forward; increasing pressure favours the side with fewer moles of gas; raising temperature favours the endothermic direction (decreasing K for exothermic reactions). Inert gas at constant volume and catalysts do not shift the equilibrium; only temperature changes the value of K.
Ionic equilibrium: acids and bases are classified by Arrhenius (H+ / OH- producers), Bronsted-Lowry (proton donor / acceptor, with conjugate acid-base pairs differing by one proton), and Lewis (electron-pair acceptor / donor). Strong electrolytes ionize almost completely, weak ones only partially with ionization constants Ka and Kb (Ka = c alpha^2 / (1 - alpha)). Water self-ionizes with Kw = [H+][OH-] = 10^-14 at 298 K. pH = -log[H+]; pH + pOH = 14. For a conjugate pair, Ka x Kb = Kw and pKa + pKb = 14.
The common ion effect (a Le Chatelier consequence) suppresses ionization. Buffer solutions resist pH change; the Henderson-Hasselbalch equation gives pH = pKa + log([salt]/[acid]), so pH = pKa when [salt] = [acid]. Salt hydrolysis: salts of strong acid + strong base are neutral, weak acid + strong base are basic, strong acid + weak base are acidic, and weak acid + weak base give pH = 7 + (1/2)(pKa - pKb). For sparingly soluble salts, Ksp = product of ion concentrations raised to their coefficients; precipitation occurs when Qsp greater than Ksp.