Introduction
Metals make up most of the periodic table, and their special properties — shiny surfaces, good conductivity, malleability — all come from one type of chemical bond: the metallic bond. In this lesson we'll build a model of metallic bonding, see what controls its strength, and connect it to alloys like bronze and brass.
The electron sea model
In a metal, atoms are packed closely together in a regular lattice. Each metal atom loses its outer (valence) electrons, which become delocalised — free to move throughout the whole structure rather than belonging to one atom. This leaves behind a lattice of positive metal cations surrounded by a "sea" of mobile electrons.
The metallic bond is the electrostatic attraction between the positive metal cations and the delocalised electrons surrounding them. This explains many typical metal properties:
- Electrical conductivity – delocalised electrons can flow and carry charge.
- Malleability and ductility – layers of cations can slide past each other without breaking bonds, because the electron sea adjusts and keeps attracting the ions.
- Thermal conductivity – mobile electrons transfer kinetic energy quickly.
- Lustre – free electrons interact with light, reflecting it.
What controls the strength of a metallic bond?
Two main factors determine how strong the metallic bond is (Talbot, p.202):
- Number of delocalised electrons per atom – more valence electrons per ion means a higher positive charge on the ion and stronger attraction to the electron sea.
- Size of the metal ion (ionic radius) – smaller ions have their charge concentrated into a smaller volume, giving a higher charge density (charge density = ion charge ÷ ion volume). Higher charge density means stronger attraction to the delocalised electrons.
So overall:
strength of metallic bond ∝ charge density ∝ (number of valence electrons) / (ionic radius)
Worked example 1
Compare Na, Mg, and Al (all from Period 3).
- Na⁺ has 1 delocalised electron and a relatively large radius.
- Mg²⁺ has 2 delocalised electrons and a smaller radius than Na⁺.
- Al³⁺ has 3 delocalised electrons and an even smaller radius.
Since charge increases and radius decreases across the row, charge density increases Na⁺ < Mg²⁺ < Al³⁺. Therefore metallic bond strength (and melting point) follows the same order: Na < Mg < Al.
Worked example 2
Why do metals have much higher boiling points than melting points? Melting only needs to loosen the rigid lattice arrangement enough for ions to move past each other — most metallic bonding survives in the liquid. Boiling, however, requires completely separating atoms into a gas, meaning all the metallic bonding must be broken (Talbot, p.202). Since bond-breaking is endothermic and requires a lot of energy here, boiling points are far higher than melting points.
Alloys: modifying metallic properties
An alloy is a mixture of a metal with one or more other elements (often metals), designed to improve properties. Because delocalised electrons still hold the structure together, alloys remain metallic in bonding, but different-sized atoms disrupt the regular lattice, often making alloys harder and more corrosion-resistant than the pure metal (Talbot, p.214). For example, stainless steel (iron, nickel, chromium) resists rusting far better than pure iron, and bronze (copper and tin) resists corrosion better than copper alone. Alloying can also change colour — brass (copper + zinc) is yellow, unlike its red and silver-white components.
Key takeaways
- Metallic bonding is the attraction between a lattice of positive metal cations and a surrounding sea of delocalised electrons.
- Bond strength increases with higher ion charge and smaller ionic radius (greater charge density).
- Delocalised electrons explain conductivity, malleability, ductility, and lustre.
- Boiling points of metals are much higher than melting points because boiling requires breaking all metallic bonds, while melting only loosens the lattice.
- Alloys keep metallic bonding but often have improved hardness, corrosion resistance, and altered colour compared to pure metals.