What is Ionic Equilibrium?
Ionic equilibrium deals with the equilibrium established between ions and molecules in an aqueous solution. It is one of the most high-weightage chapters in NEET Chemistry, typically contributing 2–3 questions every year.
Strong vs Weak Electrolytes
Strong electrolytes (HCl, NaOH, NaCl) dissociate completely in water. Weak electrolytes (CH₃COOH, NH₄OH) partially dissociate and establish an equilibrium described by the dissociation constant Ka (for acids) or Kb (for bases).
pH Scale
pH = −log[H⁺]. At 25°C, neutral water has pH = 7. For strong acid HCl at 0.01 M: pH = −log(0.01) = 2.
For weak acid HA with concentration C and Ka: [H⁺] ≈ √(Ka × C), so pH ≈ ½(pKa − log C).
Buffer Solutions
A buffer resists change in pH. The Henderson–Hasselbalch equation:
pH = pKa + log([A⁻]/[HA])
Common buffers in NEET: CH₃COONa/CH₃COOH (acidic buffer) and NH₄Cl/NH₃ (basic buffer).
Common-Ion Effect
Adding a common ion suppresses ionisation. Adding NaCl to CH₃COOH decreases [CH₃COO⁻] and shifts equilibrium left, reducing ionisation.
Solubility Product (Ksp)
For AgCl ⇌ Ag⁺ + Cl⁻: Ksp = [Ag⁺][Cl⁻]. If ionic product Q > Ksp, precipitation occurs.
Solubility (s) from Ksp: For AB type, s = √Ksp. For AB₂ type (e.g., PbCl₂), Ksp = 4s³, so s = (Ksp/4)^(1/3).
Salt Hydrolysis
Salts of weak acid + strong base (e.g., CH₃COONa) give basic solution. Salts of strong acid + weak base (e.g., NH₄Cl) give acidic solution. pH for hydrolysed salt: pH = 7 + ½(pKa − pKb).
NEET Previous Year Trends
- pH of buffer with given concentrations (almost every year)
- Identifying whether precipitation occurs using Ksp
- Degree of hydrolysis and pH of salt solutions