Factors Affecting Bond Strength
Why some metals melt at 63°C and others at 3400°C.
The Rule of Strength
Strength ∝ (Number of Delocalized e⁻ × Charge) / Ionic Radius
Smaller, highly charged ions with more delocalized electrons form stronger metallic bonds (Charge Density).
Trend 1: Across Period 3 (Na → Mg → Al)
| Metal | Charge | e⁻ in Sea | Ionic Radius (pm) | MP (°C) |
|---|---|---|---|---|
| Sodium (Na) | +1 | 1 | 102 | 98 |
| Magnesium (Mg) | +2 | 2 | 72 | 650 |
| Aluminium (Al) | +3 | 3 | 54 | 660 |
- Charge increases: stronger attraction between nucleus and the sea.
- Delocalized electrons increase: more "gluing" agents per cation.
- Radius decreases: electrons closer to nucleus → stronger electrostatic attraction.
Trend 2: Down Group 1 (Li → Cs)
Melting point decreases down the group.
↓ Radius increases, Strength decreases
All Group 1 metals form +1 ions with 1 delocalized electron. As cations get larger (more shells), the distance between the nucleus and the delocalized electrons increases, weakening the electrostatic force.
Think About It
Aluminium has a melting point (660°C) that is only slightly higher than magnesium (650°C), despite having one extra delocalized electron and a smaller ionic radius. Why isn't the difference larger?
The Al³⁺ ion is very small with a very high charge density. This actually distorts (polarises) the electron cloud so much that Al begins to show some covalent character in its bonding. The metallic model breaks down slightly — the simple "more electrons = higher MP" trend isn't perfectly linear when charge density becomes extreme.