Structure: Models of Bonding and Structure

Pure ionic bonding (complete electron transfer) and pure covalent bonding (perfectly equal sharing) are idealised extremes. Most real bonds — Al–Cl, for instance — sit somewhere in between, and HL specifically expects you to justify where on that spectrum a given bond falls. Structure 2.1–2.3 build each model; Structure 2.4 uses them together to explain real material properties like melting point, solubility and brittleness.

An ionic lattice is a giant 3-D array where every ion attracts every oppositely-charged neighbour non-directionally, which is why lattice enthalpies are so large. Coordination number — the number of nearest oppositely-charged neighbours — is fixed by the radius ratio $r_{+}/r_{-}$: a small cation relative to its anion fits fewer neighbours around it. Rock-salt structures give CN 6:6 (FCC anions, cations in octahedral holes); caesium chloride gives CN 8:8 because the larger Cs${}^{+}$ fits comfortably against 8 corner Cl${}^{-}$ ions. Shearing an ionic lattice brings like-charge ions adjacent, and the repulsion shatters the crystal — hence hard but brittle, never malleable.

Covalent bonding shares electrons rather than transferring them. VSEPR predicts molecular shape from electron domains around a central atom, with lone pairs repelling more strongly than bonding pairs and compressing bond angles below the ideal. Hybridisation (HL) links the number of electron domains to orbital mixing — 4 domains means sp³, for example. A sigma bond (head-on overlap, always present, allows rotation) differs from a pi bond (HL, sideways overlap of unhybridised p-orbitals, restricts rotation, found in double/triple bonds). Delocalisation in systems like benzene spreads pi electron density over several atoms, making bonds intermediate in length and unusually stable.

A metal is a lattice of positive ion cores sitting in a sea of delocalised valence electrons that belong to the whole structure, not any single atom. Because this bonding is non-directional, planes of cations can slide past each other without breaking bonds, explaining malleability and ductility. The same free-moving electrons conduct electricity in both the solid and molten states — unlike ionic solids, metals don't need to melt first to conduct.

The recurring exam move is to give lattice data — charges, radii, or structure type — and ask which factor controls melting point, solubility or brittleness. NaCl versus MgO melting points, for example, come down entirely to the higher charges in MgO increasing lattice enthalpy through the $q_{+}{q}_{-}/(r_{+}+r_{-})$ relationship — molar mass is irrelevant because you're breaking electrostatic bonds, not comparing molecular weight. Brittleness is explained by like-charge repulsion after shearing, not by 'weak bonds', since ionic bonds are actually very strong, just non-directional.