ChemistryCHEM 03
Thermodynamics & Chemical Equilibrium
ΔH, ΔG, ΔS, Hess's law, Kc, Kp, Le Chatelier's principle, pH, Ka, Ksp
Concept Core
Essential theory — everything NCERT tests on Thermodynamics & Chemical Equilibrium
THERMODYNAMIC FUNCTIONS
Internal energy: U. Enthalpy: H = U + PV. ΔH = qp (at constant P). Entropy: S = disorder. ΔS > 0: spontaneous if ΔH favourable.
Gibbs energy: ΔG = ΔH − TΔS. ΔG < 0: spontaneous. ΔG = 0: equilibrium. ΔG > 0: non-spontaneous.
HESS'S LAW & THERMOCHEMISTRY
ΔH for a reaction is same regardless of pathway (state function).
ΔH°formation: 1 mole of compound from elements. Elements in standard state: ΔH°f = 0. Bond energies: ΔH = Σ(bonds broken) − Σ(bonds formed).
Resonance energy = ΔH°exp − ΔH°calc. Benzene resonance ≈ −36 kcal/mol.
CHEMICAL EQUILIBRIUM
Kc = [products]/[reactants] (molar conc.). Kp = Kc(RT)^Δn. Δn = moles of gaseous products − reactants.
Q < K: reaction goes forward. Q > K: backward. Q = K: equilibrium.
Le Chatelier's principle: system opposes any change to equilibrium. Pressure increase: shifts toward fewer moles of gas. Temp increase: shifts toward endothermic side. Catalyst: no shift, only speeds up reaching equilibrium.
IONIC EQUILIBRIUM
Ka: acid dissociation constant. Stronger acid = larger Ka = smaller pKa. pH = −log[H⁺].
Buffer: Henderson-Hasselbalch: pH = pKa + log([A⁻]/[HA]).
Ksp: solubility product. AgCl: Ksp = [Ag⁺][Cl⁻]. Solubility s = √Ksp for 1:1 salt. Common ion effect reduces solubility.
Fact & Formula Vault
High-yield facts, numbers, and formulas
Thermodynamic Criteria
ΔG < 0: spontaneous
ΔG = 0: equilibrium
ΔG > 0: non-spontaneous
ΔG = ΔH − TΔS
Equilibrium Relations
Kp = Kc(RT)^Δn
Q < K: forward reaction
Le Chatelier: opposes change
Catalyst: no effect on K
Worked Examples
NEET-style questions solved step-by-step
EASYAt equilibrium for N₂ + 3H₂ ⇌ 2NH₃, if pressure is increased:▾
At equilibrium for N₂ + 3H₂ ⇌ 2NH₃, if pressure is increased:
Δn = 2 − (1+3) = −2. Le Chatelier: increased pressure shifts toward fewer moles of gas = towards products (NH₃ side). This is the basis of Haber process.
MEDIUMΔH = −400 kJ, ΔS = −200 J/K at 1000K. Is the reaction spontaneous?▾
ΔH = −400 kJ, ΔS = −200 J/K at 1000K. Is the reaction spontaneous?
ΔG = ΔH − TΔS = −400,000 − (1000)(−200) = −400,000 + 200,000 = −200,000 J = −200 kJ. ΔG < 0 → spontaneous.
HARDKp = Kc(RT)^Δn. For PCl₅ ⇌ PCl₃ + Cl₂, Δn = ?▾
Kp = Kc(RT)^Δn. For PCl₅ ⇌ PCl₃ + Cl₂, Δn = ?
PCl₅(g) → PCl₃(g) + Cl₂(g). Gaseous moles: products = 2, reactants = 1. Δn = 2 − 1 = 1. So Kp = Kc(RT)¹ = Kc × RT.
Mistake DNA
Common NEET traps for this chapter
⚠ Catalyst and K
A catalyst does NOT change the equilibrium constant K or the equilibrium position. It only changes the rate of reaching equilibrium.
✓ Fix: Catalyst → faster equilibrium, same K value
⚠ ΔG and K relationship
ΔG° = −RT ln K. If K > 1: ΔG° < 0 (products favoured). If K < 1: ΔG° > 0 (reactants favoured).
✓ Fix: K > 1 → products favoured → ΔG° < 0
⚠ Ksp and solubility
For AB type salt: s = √Ksp. For A₂B: Ksp = 4s³. Common ion decreases solubility.
✓ Fix: Don't mix up Ksp formula for different salt types
Chapter Intelligence
Exam data and last-minute strategy
NEET Frequency
3–4 Q/year. ΔG = ΔH−TΔS, Le Chatelier effect of pressure/temperature, Kp vs Kc, pH buffer, Ksp.
High-Yield
ΔG < 0 = spontaneous. Kp = Kc(RT)^Δn. Le Chatelier: P↑ → fewer gas moles side. Catalyst: no effect on K. pH = −log[H⁺].
Strategy
ΔG problems: be careful with units (kJ vs J). Equilibrium: identify Δn first for Kp/Kc conversion. Le Chatelier: always state what changes and which side has fewer moles.
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