Chemical Energetics progresses from enthalpy calculations to Born-Haber cycles, entropy, and Gibbs free energy — providing the thermodynamic tools to predict whether reactions occur spontaneously.
Standard enthalpy: measured at 298 K, 100 kPa. ΔH°f (formation): forming 1 mol from elements. ΔH°c (combustion): complete combustion of 1 mol. ΔH°neut (neutralisation). Hess\'s law: enthalpy change depends only on initial and final states, not the route. Cycle method: construct alternative routes using known values. Bond energy method: ΔH ≈ Σ(bonds broken) − Σ(bonds formed). Calorimetry: q = mcΔT, then ΔH = -q/n.
For ionic compounds: construct energy cycle linking ΔH°f to: atomisation energies, ionisation energies, electron affinities, and lattice energy. Lattice energy: energy released when gaseous ions form 1 mol of lattice (exothermic). Factors: smaller ions and higher charges → larger lattice energy. Compare theoretical (perfect ionic model) with experimental — large discrepancy suggests significant covalent character (e.g., AgI).
Entropy S: measure of disorder (J/K/mol). Gases > liquids > solids. ΔS > 0: increase in disorder (more gas molecules produced, dissolving). ΔS < 0: decrease in disorder (precipitation, fewer gas molecules). Gibbs free energy: ΔG = ΔH − TΔS. Reaction spontaneous when ΔG < 0. Four combinations of ΔH and ΔS signs determine temperature dependence of spontaneity.
Yes — spontaneity depends on ΔG = ΔH − TΔS, not just ΔH. An endothermic reaction (ΔH > 0) can be spontaneous if the entropy increase (TΔS) is large enough to make ΔG < 0. This requires a large positive ΔS and/or high temperature. Example: dissolving NH₄NO₃ in water is endothermic (the solution cools) but spontaneous because the ions become more disordered in solution (large ΔS > 0). At room temperature, TΔS > ΔH, so ΔG < 0. Many dissolving processes and decompositions are entropy-driven endothermic spontaneous reactions.
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