Take timed practice tests on Chemical Thermodynamics and Thermochemistry for JEE Main and JEE Advanced with session-wise drills, score review, and explanation-led revision.
Six 20-question timed sessions plus a 60-question chapter module. Each item is original and reframed for copyright safety.
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1. The change in enthalpy ΔH is defined as:
Explanation: ΔH = Qp (heat at constant pressure). ΔU = Qv (heat at constant volume). Relation: ΔH = ΔU + ΔngRT where Δng = change in moles of gas.
2. The first law of thermodynamics is:
Explanation: First Law: Energy is conserved. ΔU = Q − W (IUPAC: W = work done by system). Or ΔU = Q + W (W = work done on system). Both conventions are used.
3. An exothermic reaction has:
Explanation: Exothermic: heat is released to surroundings → ΔH 0.
4. Hess's law states that the enthalpy change for a reaction is:
Explanation: Hess's Law: ΔH for an overall reaction = sum of ΔH values for each step (because H is a state function). It allows calculation of ΔH for reactions that cannot be measured directly.
5. ΔH_reaction using bond enthalpies = :
Explanation: ΔH_rxn = Σ(bond energies broken in reactants) − Σ(bond energies formed in products). Breaking bonds is endothermic (+), forming bonds is exothermic (−).
6. Entropy (S) is a measure of:
Explanation: Entropy (S) quantifies the disorder or randomness of a system. Boltzmann's equation: S = k ln W, where W = number of microstates. Spontaneous processes generally increase entropy of the universe.
7. The second law of thermodynamics states:
Explanation: Second Law: Entropy of the universe (system + surroundings) increases for spontaneous processes (ΔS_universe > 0) and remains constant at equilibrium (ΔS_universe = 0).
8. The Gibbs free energy G = H − TS. A process is spontaneous when:
Explanation: ΔG = ΔH − TΔS. Spontaneous: ΔG 0. At constant T and P, ΔG
9. A reaction with ΔH 0 is:
Explanation: ΔG = ΔH − TΔS. If ΔH 0: ΔG = (−) − T(+) = always negative (−) at any T > 0. Spontaneous at all temperatures.
10. For the combustion of 1 mol CH₄ (g): ΔH = −890 kJ/mol. If ΔH = ΔU + ΔngRT at 25°C (R = 8.314 J/mol K), and Δng = −1, then ΔU is:
Explanation: CH₄ + 2O₂ → CO₂ + 2H₂O(g): Δng = 3−3 = 0. Wait: Δng = (1+2) − (1+2) = 0. ΔU = ΔH. Actually if H₂O is liquid: Δng = 1−3 = −2. ΔU = ΔH − ΔngRT = −890 − (−2)(8.314×10⁻³)(298) = −890 + 4.95 ≈ −885 kJ/mol. For Δng = −1: ΔU = −890 − (−1)(8.314×10⁻³)(298) = −890 + 2.48 ≈ −887.5 kJ/mol.
11. ΔG° = −RT ln K. If K = 1, then ΔG° is:
Explanation: ΔG° = −RT ln K. If K = 1: ΔG° = −RT ln 1 = −RT × 0 = 0. The standard Gibbs energy is zero when K = 1.
12. For a reaction, ΔH = −20 kJ and ΔS = −100 J/K. Above what temperature does the reaction become non-spontaneous?
Explanation: ΔG = ΔH − TΔS = −20000 − T(−100) = −20000 + 100T. Non-spontaneous when ΔG > 0: −20000 + 100T > 0 → T > 200 K. Above 200 K, ΔG > 0, so non-spontaneous.
13. Enthalpy of formation of an element in its standard state is:
Explanation: By convention, the standard enthalpy of formation (ΔHf°) of any element in its standard state (most stable allotrope) is defined as zero. E.g., ΔHf°(O₂ gas) = 0.
14. Resonance energy of benzene is determined experimentally from:
Explanation: Resonance energy = (ΔHcombustion theoretical for Kekulé structure) − (ΔHcombustion actual). Benzene burns with less energy than three isolated double bonds → difference is resonance stabilisation energy (~150 kJ/mol).
15. The standard free energy ΔG° is related to equilibrium constant K by:
Explanation: ΔG° = −RT ln K (from thermodynamics). For electrochemical cells: ΔG° = −nFE°cell. Both expressions are valid and can be equated: RT ln K = nFE°cell.