Class 11 Chemistry Equilibrium Notes

Class 11 Chemistry Equilibrium Notes

6.1 Equilibrium in Physical Processes

  • Equilibrium: State in which forward and reverse processes occur at the same rate.
  • Examples:
    • Liquid-vapor equilibrium in a closed container.
    • Dissolution of solids in liquids.
  • Characteristics:
    • Macroscopic properties remain constant.
    • Dynamic equilibrium occurs at microscopic level.

6.2 Equilibrium in Chemical Processes – Dynamic Equilibrium

  • Dynamic Equilibrium: Forward and backward reactions occur at the same rate.
  • Represented as:

A+B⇌C+DA + B \rightleftharpoons C + DA+B⇌C+D

  • Key points:
    • Concentrations of reactants and products remain constant.
    • Occurs in closed systems.

6.3 Law of Chemical Equilibrium and Equilibrium Constant

  • Law of Chemical Equilibrium: At equilibrium, the ratio of product concentrations to reactant concentrations, raised to their stoichiometric coefficients, is constant.

Kc=[C]c[D]d[A]a[B]bK_c = \frac{[C]^c [D]^d}{[A]^a [B]^b}Kc​=[A]a[B]b[C]c[D]d​

  • Equilibrium constant (Kc or Kp): Measures extent of reaction.
    • K >> 1 → Products favored
    • K << 1 → Reactants favored

6.4 Homogeneous Equilibria

  • Equilibria where all reactants and products are in the same phase.
  • Example:

N2(g)+3H2(g)⇌2NH3(g)N_2(g) + 3H_2(g) \rightleftharpoons 2NH_3(g)N2​(g)+3H2​(g)⇌2NH3​(g)

6.5 Heterogeneous Equilibria

  • Equilibria where reactants and products exist in different phases.
  • Example:

CaCO3(s)⇌CaO(s)+CO2(g)CaCO_3(s) \rightleftharpoons CaO(s) + CO_2(g)CaCO3​(s)⇌CaO(s)+CO2​(g)

  • Only gaseous and aqueous species are included in the equilibrium expression.

6.6 Applications of Equilibrium Constants

  • Predict reaction direction.
  • Calculate equilibrium concentrations.
  • Design industrial processes (e.g., Haber process, Contact process).

6.7 Relationship between Equilibrium Constant (K), Reaction Quotient (Q) and Gibbs Energy (G)

  • Reaction Quotient (Q): Same formula as K, but not at equilibrium.
  • Spontaneity:
    • Q < K → Reaction proceeds forward.
    • Q > K → Reaction proceeds backward.
  • Gibbs Energy Relation:

ΔG=ΔG∘+RTln⁡Q\Delta G = \Delta G^\circ + RT \ln QΔG=ΔG∘+RTlnQ

  • At equilibrium: ∆G = 0 and Q = K

6.8 Factors Affecting Equilibria

  • Concentration: Change shifts equilibrium (Le Chatelier’s Principle).
  • Pressure: Affects equilibria involving gases.
  • Temperature: Changes K; exothermic/endothermic reaction behavior differs.

6.9 Ionic Equilibrium in Solution

  • Equilibria involving ions in aqueous solutions.
  • Common examples: acid-base reactions, salt hydrolysis.

6.10 Acids, Bases, and Salts

  • Acid: Proton donor (Arrhenius/Brønsted-Lowry).
  • Base: Proton acceptor.
  • Salt: Product of acid-base neutralization.

6.11 Ionization of Acids and Bases

  • Degree of ionization (α): Fraction of molecules ionized.
  • Strong acids/bases: Almost fully ionized.
  • Weak acids/bases: Partially ionized.
  • Ionization constant (Ka/Kb):

Ka=[H+][A−][HA]Ka = \frac{[H^+][A^-]}{[HA]}Ka=[HA][H+][A−]​ Kb=[OH−][B+][BOH]Kb = \frac{[OH^-][B^+]}{[BOH]}Kb=[BOH][OH−][B+]​

6.12 Buffer Solutions

  • Definition: Solution resisting change in pH upon addition of acid or base.
  • Types:
    • Acidic buffer (weak acid + its salt)
    • Basic buffer (weak base + its salt)
  • Example: CH₃COOH + CH₃COONa → Maintains pH.

6.13 Solubility Equilibria of Sparingly Soluble Salts

  • Sparingly soluble salts: Salts with very low solubility.
  • Solubility Product (Ksp):

Ksp=[A+]m[B−]nK_{sp} = [A^+]^m [B^-]^nKsp​=[A+]m[B−]n

  • Used to predict precipitation and calculate ion concentrations in solution.

Key Points to Remember

  • Equilibrium is dynamic, not static.
  • Kc, Kp, and Q help predict reaction direction.
  • Le Chatelier’s Principle explains effects of concentration, pressure, and temperature.
  • Ionic equilibrium governs acidity, basicity, buffer solutions, and solubility.
  • Gibbs free energy links spontaneity to equilibrium constants.