AQA A Level Chemistry 3.1.8 Thermodynamics Exam Questions

About AQA A Level Chemistry 3.1.8 Thermodynamics

Thermodynamics is the A Level extension of energetics. On the AQA course you construct Born-Haber cycles to calculate lattice enthalpies, compare experimental values with those from a perfect ionic model to judge covalent character, and use cycles of enthalpies of solution and hydration. You also meet entropy as a measure of disorder, calculate entropy changes, and use the Gibbs free energy equation ΔG = ΔH - TΔS to decide whether a reaction is feasible and to find the temperature at which it becomes feasible. Examiners test precise definitions, multi-step cycle calculations, unit conversion between J and kJ, and graphs of ΔG against temperature.

What you need to know

• Key definitions Enthalpy of lattice formation: the enthalpy change when one mole of a solid ionic compound is formed from its gaseous ions (exothermic). Lattice dissociation is the reverse (endothermic). Enthalpy of atomisation: the enthalpy change when one mole of gaseous atoms is formed from the element in its standard state.
• Electron affinity First electron affinity is the enthalpy change when one mole of gaseous atoms each gain one electron to form one mole of gaseous 1- ions, and is exothermic. The second electron affinity is endothermic because an electron is added to a negative ion, which repels it.
• Born-Haber cycles The enthalpy of formation equals the sum of atomisation enthalpies, ionisation energies, electron affinities and the lattice formation enthalpy. Rearrange the cycle to find any unknown term.
• Perfect ionic model Theoretical lattice enthalpies assume spherical ions with only ionic bonding. If the Born-Haber value is more exothermic than the theoretical one, the compound has some covalent character caused by polarisation of the anion.
• Solution and hydration Enthalpy of solution = lattice dissociation enthalpy + sum of the enthalpies of hydration of the ions. Hydration is more exothermic for smaller, more highly charged ions.
• Entropy Entropy, S, measures disorder and has units of J K-1 mol-1. Gases have much higher entropy than liquids and solids. ΔS = sum of S(products) minus sum of S(reactants).
• Gibbs free energy ΔG = ΔH - TΔS. A reaction is feasible when ΔG is less than or equal to zero. Setting ΔG = 0 gives the temperature at which it becomes feasible, T = ΔH / ΔS, with ΔS converted to kJ K-1 mol-1.

Common mistakes in 3.1.8 Thermodynamics

• Mixing units in ΔG = ΔH - TΔS. Entropy is usually given in J K-1 mol-1 and must be divided by 1000 before combining with ΔH in kJ mol-1.
• Giving the second electron affinity a negative sign. It is endothermic because the incoming electron is repelled by the negative ion.
• Using the bond enthalpy of Cl2 for the atomisation of chlorine. Atomisation forms one mole of Cl atoms, which is half the bond enthalpy.
• Assuming a feasible reaction will happen quickly. A large activation energy can make a feasible reaction too slow to observe.
• Explaining the difference between theoretical and experimental lattice enthalpies without mentioning polarisation and covalent character.

Try three 3.1.8 Thermodynamics exam-style questions