Chemistry › Thermodynamics
Thermodynamics
Unit C08 · 3 lessons.
Whether a reaction can happen at all is a question of free energy, not of enthalpy alone. Born–Haber cycles find a lattice enthalpy that cannot be measured directly, enthalpies of solution and hydration explain why some ionic solids dissolve, and entropy and the Gibbs free-energy change combine enthalpy and disorder into one test of feasibility.
Whole answers are lost on units. Entropy is tabulated in J K⁻¹ mol⁻¹ and enthalpy in kJ mol⁻¹, so ΔS is divided by 1000 before ΔG = ΔH − TΔS is used, and the temperature is in kelvin. State which convention a lattice enthalpy uses, formation or dissociation, before giving its sign, because the two differ only in that sign.
On the written papers it is set on AQA Papers 1 and 3, OCR A Papers 1 and 3, Edexcel Papers 1, 2 and 3 and Cambridge Paper 4.
Specification sections: AQA 3.1.8.1, 3.1.8.2 · OCR A 5.2.1, 5.2.2 · Edexcel 2.14 to 2.21, 4.6, 13.1 to 13.22 · CAIE 23.1, 27.1, 23.2, 23.3, 23.4
Notes
Illustrated notes for each topic in this unit, with worked examples and a checklist of objectives.
- Lattice enthalpy and Born–Haber cycles
Lattice enthalpy of formation and of dissociation and their signs, the enthalpies of atomisation, first and second ionisation energy, first and second electron affinity and formation, Born–Haber cycles for NaCl, KBr, MgCl₂ and MgO drawn as energy levels and used to calculate a missing value, the factors and trends in first electron affinity in Groups 6 and 7, the perfect ionic model against experiment as evidence of polarisation and covalent character, polarising power and polarisability, and the thermal stability of the Group 1 and 2 carbonates and nitrates explained by polarisation of the anion.
- Enthalpies of solution and hydration
Enthalpy of solution and enthalpy of hydration defined, the polar water molecule orientated round cations and anions, how charge and ionic radius set the hydration enthalpy, the cycle linking enthalpy of solution to lattice enthalpy of dissociation and hydration enthalpies, why some salts dissolve endothermically, and predicting solubility trends such as the Group 2 sulfates and hydroxides and the silver halides, the calcium chloride hand warmer and the temperature rise it gives, and the entropy change on dissolving.
- Entropy and Gibbs free-energy change
Entropy as disorder and its units, how it changes on melting, boiling and with the number of moles of gas, calculating the entropy change of a reaction from standard entropies, the Gibbs free-energy change ΔG = ΔH − TΔS, the condition ΔG ≤ 0 for feasibility, the temperature T = ΔH/ΔS at which a reaction becomes feasible, the four sign cases, the graph of ΔG against temperature, the entropy change of the surroundings, −ΔH/T, and the total entropy change, ΔG = −RT ln K and what the sign of ΔG says about the equilibrium constant, and why a feasible reaction need not be fast.
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