Topic overview

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Le Chatelier’s principle describes how a disturbed equilibrium shifts in a direction that opposes the change. For concentration changes, comparing the reaction quotient, Q, with the equilibrium constant, K, helps identify the direction of reaction. Pressure changes in gaseous systems depend on the number of gas moles on each side, while inert-gas effects depend on whether volume or pressure is held constant. Temperature is distinct because it changes K; the direction of that change depends on whether the reaction is exothermic or endothermic. The lesson also connects K to temperature through the van’t Hoff equation and to standard Gibbs energy through ΔG° = −RT ln K. Catalysts speed the approach to equilibrium without changing K or its position. In non-ideal electrolyte solutions, activities and activity coefficients provide a more accurate basis for describing equilibrium.