9.07.1 Metallic Bonding
Syllabus
Metals consist of giant structures of atoms arranged in a regular pattern. The electrons in the outer shell of metal atoms are delocalised and so are free to move through the whole structure.
The sharing of delocalised electrons gives rise to strong metallic bonds. The bonding in metals may be represented in the following form:
What does this mean?
Metallic bonding and electrical conductivity.
Metals and non metals form ionic bonds by exchanging electrons to form positive and negative ions that then attract each other.
Non-metals overlap their outer shells to form covalent bonds.
But metals must also bond strongly to each other or they wouldn't have such high melting points.
All metals have 1,2 or 3 electrons in their outer shells.
So metal atoms lose their outer electrons to become ions.
These positive ions are then attracted to the "sea" of (negative) electrons that surround them.
The electrons are no longer held in place within a shell, so they can move around.
We say they are "delocalised".
Because the electrons can move they can be attracted to the positive terminal of a battery - an electric current.
Brittle or Malleable?
Ionic compounds have alternating positive and negative ions.
Disrupting the pattern is likely to push positive ions closer together (and negative ions closer together)
This causes repulsion and breaks the lattice.
Ionic compounds are Brittle.
But in a metallic crystal all the ions are positive and they are attracted to electrons that are constantly changing place.
So, moving one layer of ions over another by hitting them with a hammer is unlikely to add any more repulsions and the electrons simply find a new arrangement.
