AS & A-Level Chemistry 12 — Thermodynamics and electrochemistry

Thermochemistry extends to ionic lattices using theoretical energy cycles. The standard enthalpy change of atomisation, $\Delta H^\ominus_\mathrm{at}$, is the enthalpy change when one mole of gaseous atoms is formed from the element in its standard state. The first electron affinity $(E_{A})$ is the enthalpy change when one mole of electrons is added to one mole of gaseous atoms to form one mole of gaseous 1- ions. Lattice energy, $\Delta H^\ominus_\mathrm{latt}$, is defined by Cambridge as the enthalpy change when one mole of an ionic solid is formed from its gaseous ions; it is inherently an exothermic quantity. A Born-Haber cycle applies Hess's Law to link the standard enthalpy change of formation of an ionic compound to its constituent atomisation enthalpies, ionisation energies, electron affinities, and lattice energy. Electron affinity is influenced by nuclear charge, atomic radius, and shielding. Down Groups 16 and 17, first electron affinity generally becomes less exothermic due to increasing atomic radius and shielding, which outweigh the increased nuclear charge. However, the first electron affinities of oxygen and fluorine are anomalously less exothermic than those of sulfur and chlorine. Their small atomic radii lead to significant electron-electron repulsion in the compact $2 p$ subshell, making the addition of an incoming electron less energetically favourable.