Graphite and diamond
Graphite and diamond differ in one thing only: how many neighbours each carbon atom bonds to.
Figure to be added
Graphite
In graphite each carbon atom forms single bonds to three others. That produces flat sheets, which stack on top of one another. Carbon has four valence electrons, so one is left over on each atom, and those spare electrons form only weak attractions between the sheets.
Strong bonds within a sheet, weak attractions between sheets. So one layer slides over another easily, and that is why graphite feels slippery and works as a lubricant.
Diamond
In diamond each carbon atom forms single bonds to four others, in three dimensions. There are no layers and no leftover electrons — every valence electron is used in a bond, and the network extends rigidly in all directions.
Nothing can slide, because there are no separate pieces to slide. Diamond is the hardest substance found in nature.
| Graphite | Diamond | |
|---|---|---|
| Bonds per carbon atom | 3 | 4 |
| Arrangement | Flat layers, stacked | Three-dimensional network |
| Spare valence electron | One per atom | None |
| Behaviour | Soft, slippery, a lubricant | The hardest natural substance |
The exception to remember
Covalent compounds usually have low melting and boiling points. Graphite and diamond do not: melting them means breaking covalent bonds throughout a giant lattice, not just separating small molecules. Both have very high melting and boiling points, and this is the exception that section 10.5 will need.