Understand equilibrium constant temperature dependence, common ion effect in solubility, and salt hydrolysis calculations.
Equilibrium is one of the foundational conceptual chapters in Class 11 Physical Chemistry. It bridges thermodynamics with quantitative chemical reactivity and ion behavior in aqueous solutions.
Chemical Equilibrium Principles: The Equilibrium Constant (Kc and Kp) is constant at a given temperature and is independent of initial concentrations, volume, or catalyst presence. Temperature dependence follows Van't Hoff equation: log (K2 / K1) = (Delta H° / 2.303 R) * [(T2 - T1) / (T1 T2)].
Ionic Equilibria & Salt Hydrolysis: Salts of weak acids and strong bases (e.g., CH3COONa) undergo anion hydrolysis producing basic solutions (pH = 7 + 1/2 pKa + 1/2 log C). Salts of strong acids and weak bases (e.g., NH4Cl) undergo cation hydrolysis producing acidic solutions (pH = 7 - 1/2 pKb - 1/2 log C).
Common Ion Effect & Buffer Action: The addition of a common ion suppresses the ionization of a weak electrolyte. This principle governs qualitative salt analysis precipitation (Group II vs Group IV sulfides) and buffer solution stabilization. Dr. Aarzoo Saini guides students through equilibrium numericals at We-Gyaan Classes Roorkee.
Key Takeaways for Students
- Remember that Equilibrium Constant K varies exclusively with temperature.
- Calculate salt hydrolysis pH using derived pKa and pKb relationships.
- Apply common ion effect to explain selective sulfide precipitation in qualitative analysis.
- Solve Henderson-Hasselbalch equations for acidic and basic buffer solutions.
Authored by Dr. Aarzoo Saini
Founder & Lead Educator at We-Gyaan Classes Roorkee, with over 15 years of teaching excellence in Science and Chemistry for Board Exams, NEET, JEE, and CUET.