Topic 5 covers energy changes in chemical reactions: measuring enthalpy changes, applying Hess\'s law, and analysing bond energies. HL extends to lattice enthalpy, Born-Haber cycles, entropy, and spontaneity (Gibbs free energy).
ΔH < 0: exothermic (releases heat, temperature rises). ΔH > 0: endothermic (absorbs heat, temperature falls). Standard conditions: 298 K, 100 kPa. Standard enthalpy of formation ΔHf°: enthalpy change forming 1 mol of compound from elements in standard states. Standard enthalpy of combustion ΔHc°: complete combustion of 1 mol of substance.
Q = mcΔT where m = mass of solution (water), c = 4.18 J/g/K. ΔH = −Q/n (negative for exothermic). Hess\'s law: total enthalpy change is independent of route. ΔH_rxn = ΣΔHf°(products) − ΣΔHf°(reactants). Also calculated via combustion data or bond enthalpies.
Energy required to break one mole of bonds in gaseous molecules. ΔH ≈ Σ(bonds broken) − Σ(bonds formed). Average bond enthalpies give approximate ΔH (exact only for gaseous reactions). Stronger bonds have higher bond enthalpy (triple > double > single for same pair of atoms).
Entropy S: measure of disorder. ΔS > 0 when gas produced, solid → liquid, more particles. Gibbs free energy: ΔG = ΔH − TΔS. Spontaneous when ΔG < 0. Four cases based on signs of ΔH and ΔS: both negative (spontaneous at low T), both positive (spontaneous at high T), ΔH < 0 and ΔS > 0 (always spontaneous), ΔH > 0 and ΔS < 0 (never spontaneous).
Bond enthalpies used in IB are average values measured across many different molecules. The actual energy of a C-H bond depends on its molecular environment — C-H in CH₄ is slightly different from C-H in C₂H₆. Also, bond enthalpy calculations assume all species are gaseous, which may not be true. For precise values, use enthalpy of formation data instead.
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