Structure: Models of Bonding and Structure

Ionic, covalent and metallic models are deliberate simplifications: real electron density never sits neatly in one category, but each model lets you predict melting point, conductivity, shape and mechanical behaviour with a handful of rules. The model you choose has to match the electronegativity data, not the other way round. Almost every exam question is really asking which structural feature — coordination number, lone pair count, or delocalized electrons — explains the observed property.

An ionic bond is the electrostatic attraction between oppositely charged ions in a 3-D lattice — there is no single bond, just the sum of every ion–ion attraction acting in all directions. Coordination number is set by the radius ratio $r_{+}/r_{-}$, not by charge: NaCl is 6:6 because Na+ is too small relative to Cl– to fit 8 neighbours, while CsCl is 8:8 because Cs+ is large enough. Lattice enthalpy — and therefore melting point — scales with $q_{+}\times q_{-}$ over interionic distance, so charge usually dominates over radius when both change. Ionic solids only conduct once molten or dissolved, because in the solid the ions vibrate about fixed lattice points and cannot migrate.

A covalent bond is the electrostatic attraction between a shared electron pair and the nuclei of both bonded atoms; more shared pairs between the same two atoms means a shorter, stronger bond (single < double < triple). A sigma bond forms by head-on overlap along the bond axis and allows free rotation; a pi bond forms by sideways overlap of parallel p-orbitals and locks atoms into a planar, non-rotating arrangement, which is why cis/trans isomers exist around C=C but not C–C. A coordinate (dative) covalent bond has both electrons donated by the same atom — classic exam examples are NH4+, H3O+ and [Ag(NH3)2]+ — but once formed it behaves identically to any other covalent bond.

VSEPR treats every bonding pair and lone pair around the central atom as an electron domain that repels every other domain, pushing them as far apart as possible. Repulsion strength follows lone pair–lone pair > lone pair–bonding pair > bonding pair–bonding pair, so each extra lone pair compresses the bond angle a little further. The named shape only describes where the atoms sit, but lone pairs still take up room.

Covalent network (giant covalent) structures repeat strong covalent bonding throughout the entire crystal — diamond, SiO2, SiC — giving very high melting points and usually poor electrical conductivity, since there are no ions or delocalized charge carriers. Graphite is the exam-favourite exception: each carbon forms 3 sigma bonds plus delocalized pi electrons across a layer, and the layers are held together only by weak London forces. That structure lets graphite conduct along its layers while flaking and lubricating between them.