* Metallic Bonding: Metals form metallic bonds, not covalent bonds. Metallic bonding involves a "sea" of delocalized electrons shared by all the metal atoms in the structure. These electrons are free to move throughout the metal, giving metals their excellent conductivity.

* Covalent Bonding: Covalent bonding involves the sharing of electrons between atoms to form strong bonds. These electrons are localized between the atoms and not free to move.

Here's a breakdown of the properties of metals and giant covalent structures:

| Property | Metals | Giant Covalent Structures |

|---|---|---|

| Bonding Type | Metallic | Covalent |

| Electrical Conductivity | Excellent | Poor (except for graphite) |

| Malleability | High | Generally brittle |

| Ductility | High | Generally brittle |

| Melting Point | Generally high | Generally high (except for graphite) |

Examples:

* Metals: Iron, copper, gold, sodium

* Giant Covalent Structures: Diamond, silicon dioxide (quartz)

In summary: Metals are excellent conductors of electricity due to their metallic bonding, which allows for free movement of electrons. Giant covalent structures, on the other hand, have localized electrons and are generally poor conductors.