Metallic bonds explain the conductivity of metals through the “sea of electrons” model. Here’s how:
- Delocalized electrons – In metals, the outermost (valence) electrons are not bound tightly to any one atom. Instead, they become free to move throughout the entire structure, forming a “sea of electrons.”
- Positive ion lattice – The metal atoms lose these valence electrons and become positively charged ions arranged in a regular lattice.
- Electrons move freely – Since the electrons are delocalized, they can move easily when an electric potential (voltage) is applied. This movement of electrons constitutes an electric current.
- Conductivity in all directions – The electrons are mobile in all directions, so metals can conduct electricity equally well in any direction.
- Thermal conductivity – The same free electrons also carry heat energy quickly, explaining why metals are good conductors of heat too.
- In short: Metallic bonds allow a sea of delocalized electrons to flow freely, making metals excellent conductors of electricity and heat.