Lattice Energy and Coulomb's Law

An ionic lattice is held together by the electrostatic pull between cations and anions, and that pull follows Coulomb's law: the attraction is proportional to (charge of cation x charge of anion) divided by the distance between their centers. Lattice energy, the energy released when separate gaseous ions come together into a solid crystal, tracks that same ratio. Bigger charges or shorter distances mean a stronger pull, more energy released, and a lattice that is harder to break apart with heat.

Charge and radius both matter, but charge usually wins because it enters as a product, not a sum. Going from a 1+/1- pair to a 2+/2- pair roughly quadruples the charge product, while swapping in an ion one row down the periodic table only stretches the distance by a modest amount. That is why MgO, with 2+ and 2- ions packed close together, has a far larger lattice energy than NaCl, even though both have similar-sized ions. It also explains why a small, highly charged pair like MgO beats a large, singly charged pair like CsI despite CsI's bigger radius sum.

This ties directly to melting point: a stronger lattice needs more thermal energy to pull ions apart, so high lattice energy means high melting point, which is why MgO melts near 2800 C while NaCl melts at 801 C. The same logic carries into solubility. Dissolving an ionic solid means breaking the lattice and replacing those ion-ion attractions with ion-water attractions, so a very high lattice energy can make a salt stubbornly insoluble even when water is eager to surround the ions.