Chapter 6: EquilibriumClass 11 Chemistry — summary, notes, extra questions & MCQ quiz

Summary

This unit develops the idea of equilibrium for both physical and chemical systems. Physical equilibria — solid-liquid, liquid-vapour, solid-vapour and the dissolution of solids and gases in liquids (Henry's law) — are dynamic, with opposing processes proceeding at equal rates in a closed system so that measurable properties stay constant. Chemical equilibrium is likewise dynamic, demonstrated by the Haber synthesis of ammonia and by isotope-scrambling experiments. The law of chemical equilibrium (the law of mass action of Guldberg and Waage) gives the equilibrium constant \(K_c\) as the ratio of product to reactant concentrations each raised to its stoichiometric coefficient. For gaseous reactions \(K_p\) is used, with \(K_p=K_c(RT)^{\Delta n}\); pure solids and liquids are omitted from heterogeneous equilibrium constants. The magnitude of \(K\) indicates the extent of a reaction, while the reaction quotient \(Q\) predicts the direction of net change. Equilibrium connects to thermodynamics through \(\Delta G^{\circ}=-RT\ln K\). Le Chatelier's principle explains how concentration, pressure, temperature and catalysts affect equilibrium, a catalyst speeding both directions equally without shifting the position. The unit then turns to ionic equilibrium: strong and weak electrolytes; the Arrhenius, Brønsted-Lowry (conjugate acid-base pairs) and Lewis concepts of acids and bases; ionisation of acids and bases; the ionic product of water \(K_w=1\times10^{-14}\), the pH scale, buffer solutions and solubility product.

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