Quick Summary
Equilibrium is one of the most NEET-critical chapters (3–5 questions per year). It covers dynamic equilibrium, equilibrium constants Kp and Kc, Le Chatelier's principle, ionic equilibrium (acids, bases, salts), pH calculations, and buffers. Board exam tests 5-mark derivations and numerical problems.
Key Concepts and Definitions
- Dynamic equilibrium: Forward and reverse reactions occur at equal rates; no net change in concentrations
- Kc: Equilibrium constant in terms of molar concentrations
- Kp: Equilibrium constant in terms of partial pressures
- Reaction quotient (Q): Same expression as K but using non-equilibrium concentrations; Q < K → forward, Q > K → reverse
- Degree of dissociation (α): Fraction of moles dissociated at equilibrium
- pH: −log[H⁺]; pH + pOH = 14 at 25°C
- Ka and Kb: Dissociation constants of weak acids and bases
- Kw: Ionic product of water = [H⁺][OH⁻] = 10⁻¹⁴ at 25°C
- Common ion effect: Suppression of ionisation by addition of a common ion
Formulae
| Concept | Formula |
|---|---|
| Equilibrium expression | Kc = [products]^stoi / [reactants]^stoi |
| Kp–Kc relation | Kp = Kc(RT)^Δng |
| pH | pH = −log[H⁺] |
| Weak acid pH | pH = ½(pKa − log C) |
| Henderson-Hasselbalch | pH = pKa + log([A⁻]/[HA]) |
| Kw | Kw = [H⁺][OH⁻] = 10⁻¹⁴ |
| pKa + pKb | = pKw = 14 |
| Degree of hydrolysis | h = √(Kw / Ka · C) |
Le Chatelier's Principle — Effect Summary
| Change Applied | Equilibrium Shift |
|---|---|
| Increase concentration of reactant | Forward |
| Increase pressure | Towards fewer moles of gas |
| Increase temperature | Towards endothermic direction |
| Addition of catalyst | No shift (only speeds up both directions equally) |
| Addition of inert gas at const V | No shift |
Ionic Equilibrium
Strong acids/bases: Fully dissociate — [H⁺] = concentration of acid Weak acids: Partially dissociate — use Ka and degree of dissociation α Salt hydrolysis:
- Salt of strong acid + weak base → acidic solution
- Salt of weak acid + strong base → basic solution
- Salt of strong acid + strong base → neutral solution
Solubility product (Ksp): For AₘBₙ ⇌ mA + nB; Ksp = [A]^m [B]^n
Important Questions
One Mark (5)
- Define dynamic equilibrium.
- What is the pH of 0.01 M HCl?
- State Le Chatelier's principle.
- Define buffer solution.
- What is Kw at 25°C?
Short Answer 3-mark (3)
- For N₂ + 3H₂ ⇌ 2NH₃, write Kc and Kp expressions. Derive the Kp–Kc relation.
- Using Le Chatelier's principle, explain the effect of increasing pressure on N₂ + 3H₂ ⇌ 2NH₃.
- Calculate the pH of 0.1 M CH₃COOH (Ka = 1.8 × 10⁻⁵).
Long Answer 5-mark (2)
- Explain ionic equilibrium. Discuss the common ion effect with an example.
- State Le Chatelier's principle and apply it to the Haber process for NH₃ synthesis to explain the choice of conditions.
Common Mistakes
- Including solids/pure liquids in the equilibrium expression (they have activity = 1).
- Forgetting that adding a catalyst does NOT shift equilibrium — only speeds it up.
- Confusing Q < K (forward shift) with Q > K (reverse shift).
- Using log instead of −log for pH.
Board Exam Tips
- Kp–Kc derivation is a standard 5-mark question.
- Le Chatelier applied to Haber process or Contact process is very common.
- pH numericals (weak acid, buffer) appear every year as 3-mark questions.
- NEET: Buffer pH using Henderson-Hasselbalch and common ion effect are high-frequency.
Revision Checklist
- [ ] Kc and Kp expressions written for 3 different reactions
- [ ] Le Chatelier's table completed for all changes
- [ ] pH of strong acid, weak acid, and buffer calculated
- [ ] Kp = Kc(RT)^Δng derived from scratch
- [ ] Salt hydrolysis type identified for 5 examples
Frequently Asked Questions
At constant temperature, the rate of a reaction is proportional to the product of the molar concentrations of the reactants, each raised to the power of its stoichiometric coefficient.