Chapter 6 — Equilibrium

Chemical equilibrium is a dynamic state in a closed system where forward and reverse reaction rates are equal, resulting in constant concentrations of reactants and products; it involves physical processes (like phase transformations) and chemical reactions, and is characterized by a constant equilibrium constant (Kc or Kp) that depends on temperature.

Chapter 6 on Equilibrium covers the fundamental concept that chemical equilibria are dynamic processes occurring in closed systems where opposing reactions proceed at equal rates. The chapter explains equilibrium in physical processes (solid-liquid, liquid-vapour, solid-vapour, and dissolution equilibria), introduces the Law of Chemical Equilibrium and equilibrium constants (Kc for concentrations and Kp for partial pressures), and details how equilibrium can be reached from either direction. Key relationships include Kp = Kc(RT)^Δn for gaseous systems, and the chapter discusses homogeneous and heterogeneous equilibria with practical examples from reactions like the Haber process (N₂ + 3H₂ ⇌ 2NH₃) and H₂ + I₂ ⇌ 2HI.

1. 01Equilibrium is dynamic—both forward and reverse reactions continue simultaneously, though concentrations remain constant
2. 02Equilibrium constant expression Kc = [products]^coefficients / [reactants]^coefficients relates concentration of species at equilibrium for a given reaction
3. 03For gaseous reactions, Kp = Kc(RT)^Δn where Δn = (moles of gaseous products) – (moles of gaseous reactants)
4. 04The equilibrium constant for a reverse reaction K'c = 1/Kc; multiplying stoichiometric coefficients by n gives equilibrium constant K^n
5. 05Physical equilibria (melting point, boiling point, vapour pressure, solubility) are characterized by constant values at given temperature and pressure
6. 06Equilibrium can be established from either reactants or products; identical equilibrium mixtures result regardless of starting direction