Metallic bonding
A lattice of positive ions in a sea of delocalised electrons, and why metals conduct and are malleable.
Review these first
Learning objectives
What you need to be able to do
Teacher-mapped phrasing — check against the official Cambridge syllabus for exact wording.
- 2.7.1Describe metallic bonding as a lattice of positive ions in a sea of delocalised electrons.Supplement
- 2.7.2Explain electrical conductivity and malleability in terms of metallic structure.Supplement
6 minute read
Metallic bonding
Metallic bonding is the strong electrostatic attraction between a giant lattice of positive metal ions and a "sea" of delocalised electrons that move freely throughout the whole structure.
Each metal atom loses its outer-shell electrons into this shared "sea", becoming a positive ion. The ions are held in a regular lattice by their attraction to the surrounding delocalised electrons, which belong to the structure as a whole rather than to any single ion.
Why metals conduct electricity
Because the delocalised electrons are not fixed to any one ion, they are free to move throughout the lattice, carrying charge — this is what allows metals to conduct electricity (and heat) so well, in both the solid and molten state.
Why metals are malleable
The layers of positive ions can slide over one another when a force is applied, without breaking the metallic bonding — because the sea of delocalised electrons moves with them and continues to hold the structure together from any new position. This is why metals can be hammered or bent into shape (malleable) and drawn into wires (ductile), unlike ionic compounds, whose rigid lattice shatters if the layers are forced to shift.
Why metals have high melting points
The electrostatic attraction between the positive ions and the delocalised electrons is strong and extends throughout the entire lattice, so a large amount of energy is needed to overcome it — giving metals generally high melting and boiling points.
Think of it like this
A metal lattice is like a tray of marbles (the positive ions) glued together not by touching each other, but by being immersed in a pool of glue (the delocalised electrons) that surrounds and holds them all. Tilt the tray, and the marbles can slide into new positions while the glue simply flows and re-sets around them — the structure never breaks apart the way it would if the marbles were rigidly fixed.
Common misconceptions
- Confusing metallic bonding with ionic bonding, since both involve positive ions. In metallic bonding the ions are held by delocalised electrons shared throughout the structure; in ionic bonding, ions are held by direct attraction to oppositely charged ions.
- Thinking metals are malleable because the bonds between ions are weak. The bonding is actually strong (giving high melting points) — malleability comes from the *non-directional* nature of the bonding, allowing layers to slide without breaking it.
- Believing only the outer electrons of some atoms are delocalised. In metallic bonding, the outer electrons of every atom in the lattice contribute to the shared sea.
In the exam
- The word "delocalised" is essential in any answer about metallic bonding or conductivity — without it, an answer describing "free electrons" can be marked as incomplete.
- When explaining malleability, mention explicitly that the delocalised electrons move with the shifting layers, which is why the bonding is not broken — this is the step most answers miss.