Ionic solids and Lattice Enthalpy

Ionic solids are crystalline arrays of cations and anions held together purely by coulombic (electrostatic) attraction, with every cation surrounded by a fixed number of anions and vice versa in a regular, repeating three-dimensional lattice. Forming an ionic solid from separate gaseous atoms happens in stages: first ionisation, $M(g)\rightarrow M^{\oplus}(g)+e^-$ (always an ENDOTHERMIC step, since removing an electron from a neutral atom always costs energy), then electron gain, $X(g)+e^-\rightarrow X^-(g)$ (which can be exothermic or endothermic depending on the element), and finally condensation of the gaseous ion pair into the ordered solid lattice — a strongly exothermic step. For NaCl, ionising Na costs $495.8$ kJ mol⁻¹ and chlorine's electron gain releases $348.7$ kJ mol⁻¹, so forming gaseous Na⊕Cl⊖ ion pairs from separate gaseous atoms already releases a net $147.7$ kJ mol⁻¹ — but condensing those gaseous ion pairs further into the solid NaCl lattice releases a much larger $788$ kJ mol⁻¹. This shows that an ionic compound's real stability comes overwhelmingly from the energy released on forming its crystal lattice, not merely from each ion completing its octet. Lattice Enthalpy is accordingly DEFINED as the energy required to completely separate one mole of a solid ionic compound into its gaseous ions — e.g. NaCl's lattice enthalpy of $788$ kJ mol⁻¹ is the energy needed to pull one mole of solid NaCl apart into gaseous Na⊕(g) and Cl⊖(g) ions at infinite separation. Comparing compounds (Table 5.1), lattice entha …