Metallic bonding — common mistakes

A very common mistake is to talk about 'molecules' or 'intermolecular forces' when discussing metals. Metals are giant structures with metallic bonding throughout. Avoid using terms like 'atoms' when explaining properties like conductivity; it is the positive ions in a sea of electrons.

The high density of free electrons on the surface of a metal can absorb photons of light and become excited to higher energy levels. They immediately fall back down, re-emitting the light at the same frequency. Our eyes perceive this re-emitted light as a characteristic metallic lustre or shininess.

Yes, they do. In the liquid state, the positive ions are no longer in fixed positions, but the delocalised electrons are still present and mobile. Therefore, molten metals can still conduct electricity very well.

This is an excellent and subtle point. While the charge and number of delocalised electrons increase from Mg to Al, aluminium has a more complex crystal structure than magnesium. This structure is slightly less efficient at packing, which can slightly weaken the overall bonding compared to what would be expected from charge and radius alone. For A-Level purposes, focusing on the trend of increasing charge and electron number is sufficient to explain the general increase in melting point across the period.