The Diagonal Relationship in the Periodic Table

Approximate six-coordinate ionic radii: Li⁺ ≈ 76 pm, Na⁺ ≈ 102 pm, Mg²⁺ ≈ 72 pm.
Li⁺'s radius (76 pm) is far closer to Mg²⁺'s radius (72 pm) than to Na⁺'s radius (102 pm) — and Li⁺ and Mg²⁺ are close enough in size that Mg²⁺'s doubled charge gives it a charge density in the same range as small, singly-charged Li⁺. Na⁺, despite being in the same group as Li⁺, is both larger and lower in charge density. This is exactly why Li behaves more like Mg than like its own group neighbour Na.

Three genuine chemical similarities, none shared with the other alkali metals:
• Li₂CO₃ decomposes on heating to Li₂O + CO₂, just as MgCO₃ decomposes to MgO + CO₂ — the carbonates of Na, K, Rb and Cs are thermally stable and do not decompose this way.
• Both Li and Mg react directly with N₂ to form a nitride (Li₃N, Mg₃N₂) on heating in air — the other alkali metals do not form nitrides directly.
• LiF is only sparingly soluble in water, matching MgF₂'s low solubility — the other alkali metal fluorides (NaF, KF, etc.) are freely soluble.

Be²⁺ and Al³⁺ both have unusually high charge density for their groups, giving three clear parallels:
• Both oxides are amphoteric — BeO and Al₂O₃ dissolve in both acids and bases. Every other Group 2 oxide (MgO, CaO, etc.) is purely basic.
• Both chlorides, BeCl₂ and AlCl₃, are covalent, exist as dimers in the vapour phase, and act as Lewis acids — unlike the ionic chlorides of the rest of Group 2.
• Both metals dissolve in NaOH solution releasing H₂ gas, a reaction the other Group 2 metals do not undergo.

• Both form acidic oxides: B₂O₃ and SiO₂, both used together in borosilicate glass.
• Both form covalent halides that hydrolyse completely in water to give the corresponding oxoacid plus HX — for example, BCl₃ + 3 H₂O → H₃BO₃ + 3 HCl, and SiCl₄ hydrolyses analogously.
• Both form hydrides (boranes and silanes) that are unstable and readily hydrolysed — unlike the far more stable hydrocarbons formed by carbon, boron's own period-2 neighbour.