Metals conduct electricity. They are malleable and ductile. They have a luster. These properties come from metallic bonding. In a metal, the atoms release their valence electrons into a shared sea of delocalized electrons. The positively charged metal ions are held together by their attraction to this electron sea. The electrons are free to move, which is why metals conduct.

The metallic bond is not directional. Each ion is attracted to the electron cloud as a whole, not to specific neighbors. That is why metals can be deformed without breaking. The layers of ions can slide past each other, and the electron sea adjusts. Ionic crystals shatter when deformed because like charges repel. Metals bend.

The strength of a metallic bond depends on the number of delocalized electrons and the size of the ions. Magnesium has two valence electrons and a small ion, so it has a strong bond and a high melting point. Sodium has one valence electron and a large ion, so it has a weaker bond and a low melting point. Alloys are mixtures of metals that often have stronger bonds and better properties than pure metals.

Metallic bonding explains why metals are good conductors of heat and electricity, why they are opaque and reflective, and why they can be hammered into sheets or drawn into wires. It is a model, but a useful one.