Equilibrium, Acids and Redox Reactions
What you must understand
Equilibrium is dynamic: forward and reverse rates are equal and concentrations remain constant, not equal. Write and use Kc and Kp expressions, apply Le Chatelier's principle qualitatively, and use reaction quotient Q to predict direction. Understand the effect of temperature on K (concentration and pressure changes do not alter K). For acids and bases, use Brønsted–Lowry definitions, conjugate pairs, Kw, pH, pOH, Ka and pKa, buffers, and titration curves including equivalence point and indicator choice. For redox, assign oxidation numbers, identify oxidant and reductant, balance half-equations in acidic and basic media, and apply electrochemical cells, standard electrode potentials and Faraday calculations.
Traps in questions
- Adding a catalyst or changing pressure by adding inert gas at constant volume does not shift equilibrium position.
- Only temperature changes the value of K; a catalyst changes how fast equilibrium is reached.
- Strong acid does not mean concentrated; distinguish strength from concentration.
- In titration curves, the equivalence point pH depends on the salt formed: weak acid–strong base gives pH above 7.
- Oxidation numbers: oxygen is usually −2 but −1 in peroxides; hydrogen is +1 except in metal hydrides.
- When balancing redox in base, balance as if acidic then add OH⁻ to neutralise H⁺.
Rule of thumb and worked example
Rule: compare Q with K to predict direction; if Q < K the reaction proceeds forward.
Example: For N₂ + 3H₂ ⇌ 2NH₃, suppose at a given moment [N₂] = 0.50, [H₂] = 0.30, [NH₃] = 0.20 mol L⁻¹ and Kc = 4.0 at that temperature. Q = (0.20)² / (0.50 × 0.30³) = 0.040 / 0.0135 ≈ 2.96. Since Q < K, the forward reaction is favoured until Q reaches K. If the temperature is then raised and the reaction is exothermic, K decreases and the equilibrium shifts left.