Chemical Thermodynamics

Concept overview

Chemical thermodynamics asks a single big question: can this reaction happen, and if so, how much energy does it release or absorb? The answer comes from three state functions — internal energy U, enthalpy H, and Gibbs free energy G — along with the first and second laws that connect them.

The first law states ΔU = q + w — internal energy changes only through heat and work. For a chemical reaction at constant pressure, the enthalpy change ΔH = ΔU + PΔV captures both the energy exchange and any pressure-volume work. Hess's law generalises this: enthalpy is a state function, so the total enthalpy change for a multi-step reaction is the sum of the individual step enthalpies. This is the workhorse identity for computing ΔH of complicated reactions from tabulated formation enthalpies.

The second law introduces entropy S — a measure of disorder or the number of microscopic states consistent with a given macroscopic condition. Spontaneous processes always increase the total entropy of the universe. Combining ΔH and ΔS gives the Gibbs free energy ΔG = ΔH − TΔS, and a reaction is spontaneous at a given temperature exactly when ΔG < 0.

The chapter has a small formula set and a large logic surface. Most JEE difficulty comes from setting up the problem — identifying whether you're at constant P or V, whether a phase change is involved, and whether ΔH_f or ΔH_c values are the right table to use.

Key formulas & concepts

• First law: ΔU = q + w
• Enthalpy definition: H = U + PV
• ΔH at constant P: ΔH = q_P
• Hess's law: ΔH_reaction = Σ (n × ΔH_f, products) − Σ (n × ΔH_f, reactants)
• Kirchhoff's law: ΔH(T₂) = ΔH(T₁) + ΔC_p × (T₂ − T₁)
• Entropy change: ΔS = q_rev / T
• Gibbs free energy: ΔG = ΔH − TΔS
• Spontaneity criterion: ΔG < 0 (at constant T, P)
• ΔG standard vs K: ΔG° = −RT ln K

Common mistakes to avoid

• Confusing ΔH_f (formation) with ΔH_c (combustion) — they have opposite signs and different tabulated values.
• Applying ΔH = ΔU + PΔV to a solid-liquid reaction where ΔV ≈ 0 and getting a small correction wrong; for gas-phase reactions ΔH − ΔU = Δn_gas × RT.
• Using Δn without checking whether every species is a gas — solids and liquids contribute zero.
• Assuming ΔS positive means spontaneous — spontaneity depends on ΔG, which weighs ΔS against ΔH/T.

JEE Main relevance

JEE Main asks 2–3 thermodynamics problems every paper. Expect one direct Hess's-law numerical, one ΔG-based spontaneity question, and one that tests whether you correctly compute ΔH − ΔU using Δn_gas × RT. Formulas are few; the trap is in setup.

JEE Advanced relevance

Advanced thermodynamics problems combine multiple concepts — Hess's law + Kirchhoff's law (temperature dependence of ΔH), or ΔG connecting to K_eq through ΔG° = −RT ln K, then linking to a chemical equilibrium problem. Bomb calorimetry setups (constant V, adiabatic) are common in Advanced.

Previous-year questions (PYQ)

Thermodynamics PYQs cluster into three archetypes — Hess's law, ΔG-spontaneity and enthalpy-of-solution. Filter by 'Chemical Thermodynamics' in the PYQ bank and drill archetype-by-archetype.