Metallic bonding
Section: 3 Chemical Bonding | Syllabus: Cambridge AS Level Chemistry (9701)
Metallic Bonding
Metallic Bonding The electrostatic attraction between positively charged metal ions (cations) and a sea of delocalised electrons.When metal atoms pack together in a solid, their outer (valence) electrons leave their parent atoms and become free to move throughout the entire metallic lattice. This creates:
- A regular lattice of positive metal ions (cations) - the metal atoms having lost their valence electrons.
- A "sea" of delocalised electrons - free electrons that are no longer associated with any particular atom and can move throughout the structure.
- Strong electrostatic attraction between the positive ions and the delocalised electrons holds the lattice together.
Figure 3.14: Metallic Bonding
(Regular array of positive metal ions (circles with + symbol) in a lattice. Between and around them, randomly scattered small dots representing delocalised electrons. Label: "sea of delocalised electrons" and "positive metal ions". Arrows showing electrostatic attraction between ions and electrons.)
Factors Affecting Strength of Metallic Bonding
- Number of delocalised electrons per atom: More valence electrons donated → stronger metallic bonding. E.g. Al (3 delocalised e⁻) has stronger bonding than Na (1 delocalised e⁻).
- Ionic radius: Smaller ions → delocalised electrons are closer to the positive nucleus → stronger attraction → stronger bonding.
- Ionic charge: Higher charge on metal ion → greater electrostatic attraction → stronger metallic bond.
| Metal | Delocalised e⁻ per atom | Melting point / °C | Trend |
|---|---|---|---|
| Na | 1 | 98 | Melting point increases across Period 3 as metallic bonding strengthens |
| Mg | 2 | 650 | |
| Al | 3 | 660 |
Properties Explained by Metallic Bonding
- High melting and boiling points: Strong electrostatic attractions between positive ions and delocalised electrons require a lot of energy to overcome.
- Good electrical conductivity: Delocalised electrons can move freely and carry charge through the metal.
- Good thermal conductivity: Delocalised electrons transfer kinetic energy rapidly through the lattice.
- Malleability and ductility: Layers of ions can slide over each other without disrupting the sea of delocalised electrons - the bonding is non-directional.
- Lustre (shiny appearance): Delocalised electrons absorb and re-emit light.
Metallic bonding is non-directional - the electrostatic attraction acts equally in all directions. This is why metals are malleable: layers of ions can slip without the bond breaking completely, unlike ionic or covalent lattices.
Relative Strength of Bonding Types
In general, the three primary bonding types are all significantly stronger than intermolecular forces:
| Bonding Type | Typical strength | Example |
|---|---|---|
| Covalent (triple bond) | ~200–1000 kJ mol⁻¹ | N≡N (940 kJ mol⁻¹) |
| Ionic | ~600–3000 kJ mol⁻¹ (lattice) | MgO |
| Metallic | Moderate to very strong | Al, Fe, W |
| Hydrogen bonding | ~10–40 kJ mol⁻¹ | H₂O···H₂O |
| Permanent dipole–dipole | ~2–20 kJ mol⁻¹ | HCl···HCl |
| id–id (London dispersion) | ~0.1–10 kJ mol⁻¹ | Noble gases, alkanes |
The syllabus requires you to state that ionic, covalent and metallic bonding are in general stronger than intermolecular forces. This explains why molecular substances (e.g. Cl₂, HCl) have much lower melting/boiling points than ionic or metallic solids.
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