Polarizability and Fajans' Rules

Melting points: LiCl 605 °C, NaCl 801 °C, KCl 770 °C. A purely ionic (lattice-energy) model would predict LiCl to have the highest melting point of the three, since Li⁺ is the smallest cation and should pack most tightly against Cl⁻. Instead, LiCl's melting point is markedly lower than NaCl's — exactly what Fajans' rules predict, because Li⁺'s small size gives it high charge density and strong polarising power, pulling covalent character into the Li–Cl bond and weakening the simple ionic-lattice picture. LiCl is also the only one of the three that dissolves appreciably in organic solvents like ethanol.

Ag⁺ ([Kr]4d¹⁰, ~115 pm) and Na⁺ ([Ne], ~102 pm) are not hugely different in size or charge, so rules 1–3 alone would predict only a small difference. Yet AgCl is far less soluble in water than NaCl, has a distinctly off-white/cream colour, and is light-sensitive — all signs of substantial covalent character. The difference is rule 4: Ag⁺'s filled 4d¹⁰ shell shields the nucleus poorly, giving Ag⁺ much greater polarising power than a true noble-gas-configuration ion of similar size and charge like Na⁺.

Keeping the cation fixed at Ag⁺ isolates the effect of anion size. Solubility and covalent character follow the anion's polarizability: AgF is largely ionic and freely soluble in water; solubility falls sharply through AgCl and AgBr; AgI is the most covalent and least soluble, and visibly coloured (pale yellow), because I⁻ is the largest, most easily distorted halide ion in the series.

BeCl₂ and AlCl₃ are both markedly covalent — BeCl₂ and AlCl₃ (as its dimer Al₂Cl₆) both sublime, dissolve in organic solvents and act as Lewis acids — because Be²⁺ and Al³⁺ combine small ionic radius with high charge, giving very high charge density. MgCl₂ and CaCl₂, with lower charge density cations, stay much closer to the ionic end of the spectrum. (This is the same reasoning behind the Be–Al diagonal relationship — see the related article below.)