Metallic bonding

The Metallic Lattice Model

Metals exist as giant lattices of positive metal ions surrounded by a sea of delocalised electrons, and this regular three-dimensional arrangement defines all metallic structures.

When metal atoms pack together, each atom releases its outer-shell electrons into a shared pool. These electrons are no longer attached to any one atom and are described as delocalised [free to move throughout the entire lattice].

The atoms left behind become positive metal ions (cations) because they have lost their outer electrons. These cations sit in fixed positions, arranged in a regular, repeating pattern that extends in all directions.

The structure is described as a giant lattice because the same pattern of ions and delocalised electrons repeats continuously throughout the whole solid. There are no individual molecules in a metal.

Defining Metallic Bonding

Metallic bonding is the strong electrostatic attraction between positive metal ions and the sea of delocalised electrons that surrounds them.

This attraction acts in all directions throughout the lattice because the delocalised electrons are spread evenly between every metal ion. As a result, the bonding holds the entire structure together as one rigid solid.

The strength of the metallic bond depends on two factors. The first is the charge on the metal ion — ions with higher charges (e.g. Mg²⁺) attract the delocalised electrons more strongly than ions with lower charges (e.g. Na⁺). The second is the number of delocalised electrons per ion — more delocalised electrons per atom produce a stronger attraction.

MisconceptionMetallic bonding is not the “sharing” of electrons between two atoms like a covalent bond. It is the electrostatic attraction between many positive ions and a shared sea of mobile electrons spread across the whole lattice.

Exam TipAlways state “electrostatic attraction between positive metal ions and delocalised electrons” — partial definitions lose the mark.

Properties Explained by the Model

The metallic bonding model explains electrical conductivity and high melting temperatures by linking each property directly to the behaviour of delocalised electrons and the strength of the lattice.

Metals conduct electricity in both the solid and liquid state because the delocalised electrons are free to move throughout the lattice. When a potential difference is applied, these electrons drift towards the positive terminal, carrying charge through the metal.

Metals have high melting temperatures because a large amount of energy is required to overcome the strong electrostatic attraction between the positive metal ions and the delocalised electrons. Breaking down the giant lattice means breaking many of these attractions at once, so a high temperature is needed.

The melting temperature varies between different metals. Metals with smaller ions and higher charges (e.g. magnesium) have stronger metallic bonding and therefore higher melting points than metals with larger ions and lower charges (e.g. sodium).

Examiner InsightWhen asked to explain a property, name the specific feature of the model that causes it. For melting points, state “strong electrostatic attraction between metal ions and delocalised electrons” — vague phrases like “strong metallic bonds” may not earn full credit.

Exam TipLink the property to the specific particles involved.