Metallic bonding
Metals make up most of the periodic table and dominate engineering. Their properties — strength, conductivity, malleability — all flow from a single bonding model.
The metallic bonding model
In a metal, atoms pack together in a regular giant lattice. Each atom donates its outer-shell electrons into a "sea" of delocalised electrons that flow freely throughout the structure.
What's left of each atom is a positive ion (cation). The lattice is held together by the electrostatic attraction between the positive ions and the sea of negative electrons.
This is the picture you must learn:
- Regular rows/layers of positive metal ions.
- A surrounding "sea" of delocalised electrons (one or more per atom, depending on group).
- Strong, non-directional bonds throughout the whole lattice.
Why metals conduct electricity
Delocalised electrons are free to move through the lattice. When a potential difference is applied, electrons drift, carrying charge. This works in solid and molten metals, unlike ionic substances which only conduct when molten/aqueous.
Why metals conduct heat
Delocalised electrons gain kinetic energy at the hot end and quickly transfer it through the structure by collisions. The closely packed lattice also allows efficient vibration transfer.
Why metals are malleable and ductile
The layers of positive ions can slide over each other when force is applied — the delocalised electrons keep holding the structure together. So metals can be hammered into shape (malleable) or drawn into wires (ductile) without shattering, unlike ionic crystals.
High melting and boiling points
Metallic bonds are strong; the electrostatic attraction between many positive ions and the electron sea requires a lot of energy to overcome. Mercury is the famous exception (liquid at room temperature) because of weak bonding between its full-shell-like electron configuration.
Alloys — why they're harder than pure metals
A pure metal has identical atoms in identical positions. Layers slide easily, so pure metals are soft. An alloy is a mixture of a metal with one or more other elements (often metals).
The added atoms have a different size, which distorts the regular layers. Distorted layers cannot slide as easily over each other, so the alloy is harder and stronger than the pure metal.
Examples:
- Bronze (Cu + Sn) — harder than copper.
- Brass (Cu + Zn) — used for fittings, instruments.
- Steel (Fe + C, sometimes other metals) — harder than iron.
- Stainless steel (Fe + Cr + Ni) — corrosion resistant.
- Gold alloys (e.g. 18-carat gold = 75% Au + Cu/Ag) — harder, more durable than pure gold.
✦Worked example— Worked example — explaining a property
"Why is copper a good electrical conductor?"
- Copper has metallic bonding B1
- It contains delocalised electrons B1
- Free electrons can move through the lattice and carry charge when a p.d. is applied B1
⚠Common mistakes
- Saying metals share electrons in covalent-style pairs. They donate them to a delocalised sea — non-directional.
- Saying alloys are compounds. Alloys are mixtures — atoms aren't chemically combined in fixed proportions.
- Saying alloys are weaker than pure metals. Opposite — they're stronger because layers can't slide easily.
- Confusing thermal and electrical conduction. Both rely on delocalised electrons, but heat also uses ion vibrations.
Links
Builds on C2.6 (linking properties to bonding) and C10.8 (specific alloys). Compare with C2.2 (ionic) and C2.3 (covalent).
AI-generated · claude-opus-4-7 · v3-deep-chemistry