Diamond
Has a very high melting point (almost 4000°C).
Very strong carbon-carbon covalent bonds have to be broken throughout the structure before melting occurs.
Is very hard.
This is again due to the need to break very strong covalent bonds operating in 3-dimensions.
Doesn’t conduct electricity.
All the electrons are held tightly between the atoms, and aren’t free to move.
Is insoluble in water and organic solvents.
There are no possible attractions which could occur between solvent molecules and carbon atoms which could outweigh the attractions between the covalently bound carbon atoms.
The giant covalent structure of graphite
Graphite has a layer structure which is quite difficult to draw convincingly in three dimensions.
The diagram below shows the arrangement of the atoms in each layer, and the way the layers are spaced.
The bonding in graphite
Each carbon atom uses three of its electrons to form simple bonds to its three close neighbours. That leaves a fourth electron in the bonding level. These “spare” electrons in each carbon atom become delocalized over the whole of the sheet of atoms in one layer.
They are no longer associated directly with any particular atom or pair of atoms, but are free to wander throughout the whole sheet. The important thing is that the delocalized electrons are free to move anywhere within the sheet – each electron is no longer fixed to a particular carbon atom.
There is, however, no direct contact between the delocalized electrons in one sheet and those in the neighbouring sheets. The atoms within a sheet are held together by strong covalent bonds – stronger, in fact, than in diamond because of the additional bonding caused by the delocalized electrons.
So what holds the sheets together? In graphite you have the ultimate example of van der Waals dispersion forces. As the delocalized electrons move around in the sheet, very large temporary dipoles can be set up which will induce opposite dipoles in the sheets above and below – and so on throughout the whole graphite crystal.