Carbon structures

In diamond, each carbon atom forms four covalent bonds with other carbon atoms in a rigid giant covalent structure.

• Very hard — the hardest natural substance — so it is used in cutting tools.
• Very high melting point, because many strong covalent bonds must be broken.
• Does not conduct electricity, because all four outer electrons are used in bonds, so there are no free electrons.

In graphite, each carbon atom forms three covalent bonds, making layers of hexagonal rings.

• There are no covalent bonds between the layers, only weak forces, so the layers slide over each other. Graphite is soft and slippery, useful as a lubricant and in pencils.
• It has a very high melting point, because the covalent bonds within the layers are strong.
• Each carbon has one delocalised electron, which can move along the layers. So graphite conducts electricity and heat, like a metal.

Graphene is a single layer of graphite — just one atom thick.

It is very strong (because of its covalent bonds), very light, and conducts electricity (because of its delocalised electrons). This makes it useful in electronics and in composites, where adding a little graphene makes a material much stronger.

Fullerenes are molecules of carbon atoms with hollow shapes. They are based on hexagonal rings of carbon atoms, but may also have rings of five or seven carbon atoms.

The first one discovered was Buckminsterfullerene (C₆₀), which is spherical, like a football. Fullerenes can be used to deliver drugs into the body, as lubricants and as catalysts.

Carbon nanotubes are cylindrical fullerenes: very long and very thin, with a very high length to diameter ratio. They are very strong for their weight and conduct electricity, so they are used in electronics, nanotechnology and to reinforce materials such as tennis rackets.