Units 6 and 9 cover energy changes in reactions (enthalpy) and the thermodynamic criteria for spontaneity (entropy and Gibbs free energy) — fundamental to predicting whether reactions will occur.
ΔH < 0: exothermic (releases heat). ΔH > 0: endothermic (absorbs heat). Hess\'s law: ΔH depends only on initial and final states. ΔH°rxn = Σ ΔH°f(products) - Σ ΔH°f(reactants). Bond enthalpy method: ΔH ≈ Σ(bonds broken) - Σ(bonds formed). Calorimetry: q = mcΔT, then ΔH = -q/n.
Entropy S: measure of energy dispersal/disorder (J/mol·K). ΔS > 0: more disorder (more gas molecules, dissolving, higher temperature). Second law: total entropy of universe always increases for spontaneous processes. Gibbs: ΔG = ΔH - TΔS. Spontaneous when ΔG < 0. ΔG° = Σ ΔG°f(products) - Σ ΔG°f(reactants). Four cases: (-H, +S always spontaneous), (+H, -S never), (-H, -S at low T), (+H, +S at high T).
Yes — spontaneity is determined by ΔG = ΔH - TΔS, not ΔH alone. An endothermic reaction (ΔH > 0) is spontaneous if TΔS > ΔH, making ΔG negative. This requires a large positive ΔS (increase in disorder) and sufficiently high temperature. Example: dissolving NH₄NO₃ in water is endothermic (solution gets cold) but spontaneous because the ions become disordered in solution (large +ΔS). Similarly, decomposition reactions at high temperatures are often endothermic but spontaneous because ΔS is very positive (solid → gases).
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