Thermodynamics (A Level) flashcards

Topic 8: Thermodynamics

Enthalpy of atomisation, ΔatH⦵

The enthalpy change when 1 mole of gaseous atoms is formed from the element in its standard state, e.g. ½Cl₂(g) → Cl(g). Always positive.

First and second electron affinity of oxygen: equations and signs

First: O(g) + e⁻ → O⁻(g), exothermic. Second: O⁻(g) + e⁻ → O²⁻(g), endothermic (adding an electron to a negative ion).

Lattice enthalpy of formation vs lattice enthalpy of dissociation

Formation: 1 mole of solid ionic compound formed from its gaseous ions (negative). Dissociation: 1 mole of solid separated into its gaseous ions (positive). Same size, opposite sign.

Enthalpy of solution from lattice and hydration enthalpies

ΔsolH = lattice enthalpy of dissociation + Σ ΔhydH of all the ions (count each ion).

Born–Haber value more exothermic than the perfect ionic model value: what does it show?

Covalent character. The cation polarises the anion’s electron cloud, giving extra attraction. Greatest with a small, highly charged cation and a large anion.

When is ΔS positive?

When a solid becomes liquid or gas, a liquid becomes gas, or the number of moles of gas increases.

Gibbs free-energy equation and the feasibility condition

ΔG = ΔH − TΔS (T in K, ΔS in kJ K⁻¹ mol⁻¹ to match ΔH). A reaction is feasible when ΔG ≤ 0.

Temperature at which a reaction becomes feasible

Set ΔG = 0, so T = ΔH ÷ ΔS. With ΔH and ΔS both positive it is feasible above this T; both negative, below it.

Graph of ΔG against T: gradient and intercept

Gradient = −ΔS. Intercept at T = 0 = ΔH. It crosses ΔG = 0 where the reaction becomes feasible.

Classic unit trap in ΔG = ΔH − TΔS

Entropies are in J K⁻¹ mol⁻¹ but ΔH is in kJ mol⁻¹: divide ΔS by 1000 first.