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Diamond and graphite are both allotropes of carbon, which means they share the same chemical structure. But they differ a lot in their physical properties. Graphite is soft and slippery, while diamonds are extremely hard. This difference has to do with how the carbon atoms rearrange themselves.
Diamonds have a giant covalent structure, where each carbon atom is covalently bonded to four other atoms in a rigid, tetrahedral order. This creates strong bonds that are equally strong in all directions, and it makes diamonds very hard.
In graphite, the carbon atoms bond together in layers. Each layer is held together by weak intermolecular forces called Van der Waals dispersion forces. The delocalized electrons of the different layers can set up large temporary dipoles, and these can induce opposite dipoles in the neighboring sheets. This causes the sheets to be attracted to each other, but the forces are not as strong as the covalent bonds in a diamond.
As a result, the carbon atoms in graphite layers can slide over each other easily, and this makes graphite very soft. Graphite is used in pencils, and it is also an important ingredient in lubricants. It even conducts electricity, thanks to the’spare’ electrons that are delocalized between the layers of carbon atoms.
Another way that the structures of diamonds and graphite differ is that they can be cleaved along certain planes but not others. For example, a diamond can be cleaved along its x-axis but not its y-axis.
