Metallic bonding

Metals are described by a single, unifying model: a giant three-dimensional lattice of positive metal ions held together by a sea of delocalised electrons.

The metal ions form because each metal atom releases its outer-shell electrons into a shared pool. These electrons are not bound to any single ion and are free to move throughout the structure.

Key Definition Metallic bonding itself is the strong electrostatic attraction between these positive ions and the delocalised electrons.

This attraction acts in every direction across the lattice, so the bonding extends throughout the entire solid rather than between specific pairs of atoms. The strength of the bond increases with higher ionic charge and a greater number of delocalised electrons released per atom.

The model directly explains the key physical properties of metals required at this level. Electrical conductivity arises because the delocalised electrons can drift through the lattice when a potential difference is applied, carrying charge in both the solid and liquid (molten) states. High melting temperatures arise because the many strong electrostatic attractions throughout the lattice require large amounts of energy to overcome. Differences in melting temperature between metals can therefore be explained by differences in ionic charge and in the number of delocalised electrons released per atom.

The bigger picture is that metallic bonding is one of the three main bonding types, alongside ionic and covalent bonding. Each is rooted in electrostatic attraction but differs in how the electrons are arranged.

Exam Tip In exam answers, full marks depend on naming both the positive metal ions and the delocalised electrons, and on identifying the attraction between them as electrostatic. Vague references to "metallic bonds" without describing the underlying particles consistently lose marks.