Questions
5 questions per paper
Difficulty
Medium-Hard
Importance
High yield for JEE Advanced and NEET
Overview
Thermodynamics in Physical Chemistry focuses on energy transformations and the criteria for chemical spontaneity using state functions like Enthalpy, Entropy, and Gibbs Free Energy. It is a high-yield topic in competitive exams because it bridges the gap between molecular interactions and macroscopic equilibrium. Mastery requires a deep understanding of sign conventions and the application of Hess's Law in complex multi-step reactions.
Enthalpy and Hess's Law
Enthalpy (H) accounts for the total heat content of a system at constant pressure, while Hess's Law provides a powerful tool for calculating enthalpy changes in complex pathways. Since enthalpy is a state function, the heat of reaction remains independent of the route taken, allowing us to combine thermochemical equations algebraically.
- ΔH = ΔU + PΔV (for constant pressure)
- ΔH = ΣH(products) - ΣH(reactants)
- Hess's Law: ΔH_total = ΔH1 + ΔH2 + ...
- ΔH = ΔU + Δn(g)RT
- Standard enthalpy of combustion vs formation definitions
Entropy and the Second Law
Entropy (S) serves as a measure of the randomness or disorder within a system, acting as the primary indicator for the direction of spontaneous processes. The Second Law of Thermodynamics dictates that for any spontaneous process, the total entropy of the universe must increase, providing a definitive boundary for physical and chemical changes.
- ΔS_univ = ΔS_sys + ΔS_surr > 0 for spontaneity
- ΔS = q(rev) / T
- ΔS_surr = -ΔH_sys / T
- S(gas) > S(liquid) > S(solid)
- Entropy increases with temperature and volume
Gibbs Energy and Spontaneity
Gibbs Free Energy (G) combines enthalpy and entropy into a single criterion to predict spontaneity at constant temperature and pressure. By monitoring the sign of ΔG, aspirants can determine if a reaction is feasible, at equilibrium, or non-spontaneous without calculating the entropy of the surroundings.
- ΔG = ΔH - TΔS
- ΔG < 0 (spontaneous), ΔG = 0 (equilibrium), ΔG > 0 (non-spontaneous)
- ΔG° = -RT ln K
- ΔG = ΔG° + RT ln Q
- Temperature dependence of spontaneity based on ΔH and ΔS signs
Bond Enthalpies
Bond dissociation enthalpy is the energy required to break one mole of gaseous bonds into gaseous atoms. This concept is vital for estimating enthalpy changes of reactions where experimental data is missing, especially in gas-phase organic synthesis.
- ΔH_rxn = Σ(Bond energy of reactants) - Σ(Bond energy of products)
- Applies strictly to gaseous systems
- Mean bond enthalpy used for polyatomic molecules
- Directly proportional to bond strength and inversely to bond length
Formula Sheet
ΔH = ΔU + PΔV
ΔH = ΔU + Δn(g)RT
ΔS_univ = ΔS_sys + ΔS_surr
ΔS = q(rev)/T
ΔG = ΔH - TΔS
ΔG° = -RT ln K
ΔG = ΔG° + RT ln Q
ΔH_rxn = ΣBE(reactants) - ΣBE(products)
Exam Tip
Always check the units of ΔH (kJ/mol) and ΔS (J/mol·K) before plugging them into the ΔG = ΔH - TΔS equation; the most common trap is the unit mismatch.
Common Mistakes
- Ignoring the Δn(g) term when converting between ΔH and ΔU, often failing to account for only gaseous moles.
- Confusing the signs of ΔS_surr with ΔS_sys, leading to incorrect calculations for ΔS_total.
- Forgetting to convert Temperature to Kelvin in the Gibbs energy equation, causing significant arithmetic errors.
More Revision Notes
Ready to test yourself?
Play topic-wise Thermodynamics (Physical Chem) questions in Aspirant Arcade — gamified MCQ practice.
Download Free