Chemical Thermodynamics

Why This Chapter Matters

Thermodynamics in Chemistry overlaps with Physics but adds Hess's Law, bond enthalpies, and Gibbs free energy criteria for spontaneity — 6-8 marks in JEE every year.

Core Concepts

1. System, Surroundings, Types

Open system: exchanges matter + energy with surroundings

Closed system: only energy exchange

Isolated system: no exchange

State functions (path-independent): T, P, V, U, H, G, S

Path functions (path-dependent): Q (heat), W (work)

2. Internal Energy and Enthalpy

ΔU = q + w (first law, q = heat added, w = work done ON system)

w = -P_ext ΔV (work of expansion against external pressure)

Constant pressure: ΔH = ΔU + PΔV = q_p (enthalpy = heat at constant P)

Constant volume: ΔU = q_v (heat at constant V)

Relation: ΔH = ΔU + Δn_g RT (Δn_g = change in moles of gas)

3. Hess's Law

Enthalpy change is path-independent (state function).

ΔH of reaction = sum of ΔH of steps (can add/subtract thermochemical equations)

Bond enthalpy: ΔH_rxn = ΣBE(reactants broken) - ΣBE(products formed)

Standard enthalpy of formation ΔHf°: enthalpy for forming 1 mol from elements in standard states.

ΔH°rxn = ΣΔHf°(products) - ΣΔHf°(reactants)

Important enthalpy types:

ΔH_combustion: burning in excess O₂. Always negative (exothermic).

ΔH_neutralisation: strong acid + strong base → always -57.3 kJ/mol (in water)

ΔH_vaporisation: liquid → gas

ΔH_sublimation: solid → gas

4. Entropy (S)

Measure of disorder/randomness.

ΔS > 0: disorder increases (spontaneous tendency)

Gases have higher S than liquids than solids.

ΔS°rxn = ΣS°(products) - ΣS°(reactants)

Second law: In spontaneous processes, total entropy (system + surroundings) increases.

For isolated system: ΔS_total > 0 (spontaneous), = 0 (equilibrium)

5. Gibbs Free Energy (G) — Most Important for JEE!

G = H - TS

ΔG = ΔH - TΔS (at constant T, P)

ΔG < 0: spontaneous | ΔG = 0: equilibrium | ΔG > 0: non-spontaneous

|---|---|---|

Relation to equilibrium: ΔG° = -RT ln K (K = equilibrium constant)

Also: ΔG = ΔG° + RT ln Q

PYQs

2024: ΔG for a reaction at 300K: ΔH = -40 kJ/mol, ΔS = -100 J/mol·K. Spontaneous?

ΔG = -40000 - 300×(-100) = -40000 + 30000 = -10000 J = -10 kJ. ΔG<0 → Spontaneous.

2023: At what temperature does a reaction with ΔH=+20 kJ and ΔS=+50 J/K become spontaneous?

ΔG < 0: ΔH - TΔS < 0 → T > ΔH/ΔS = 20000/50 = 400 K.

2022: Δn_g for: N₂(g) + 3H₂(g) → 2NH₃(g). Relation between ΔH and ΔU?

Δn_g = 2-4 = -2. ΔH = ΔU + Δn_g RT = ΔU + (-2)RT = ΔU - 2RT.

Revision Notes

FIRST LAW: ΔU = q + w  (q in, w on system)
ΔH = q_p (constant pressure)
ΔH = ΔU + Δn_g RT  (Δn_g = moles gas products - moles gas reactants)

HESS'S LAW: ΔH path-independent (state function)
ΔH_rxn = ΣBE(broken) - ΣBE(formed)
ΔH°rxn = ΣΔHf°(P) - ΣΔHf°(R)

ENTROPY: disorder. Gas > liquid > solid. ΔS_total ≥ 0 (2nd law)

GIBBS: ΔG = ΔH - TΔS
ΔG<0: spontaneous | ΔG=0: equilibrium | ΔG>0: non-spontaneous
ΔG° = -RT ln K (links thermodynamics to equilibrium)