How does the structure of diamond differ from graphite?

Diamond and graphite differ in structure; diamond has a tetrahedral structure while graphite has a layered structure.

Diamond and graphite are both allotropes of carbon, meaning they are made up of carbon atoms arranged in different ways. In diamond, each carbon atom is covalently bonded to four other carbon atoms in a tetrahedral structure. This results in a three-dimensional network of carbon atoms, making diamond extremely hard and giving it its characteristic brilliance. The strong covalent bonds require a lot of energy to break, making diamond the hardest known natural substance. It's also a poor conductor of electricity as there are no free electrons available.

On the other hand, in graphite, each carbon atom is covalently bonded to three other carbon atoms, forming layers of hexagonal rings. These layers are held together by weak Van der Waals forces, which allow the layers to slide over each other easily. This gives graphite its characteristic slippery feel and makes it useful as a lubricant and in pencils. Unlike diamond, graphite is a good conductor of electricity. This is because one electron from each carbon atom is delocalised, allowing it to move freely between the layers and conduct electricity.

The differences in the structures of diamond and graphite lead to their distinct physical properties. Diamond's tetrahedral structure makes it incredibly hard and a poor conductor of electricity, while graphite's layered structure makes it soft, slippery, and a good conductor of electricity. Understanding these structural differences is crucial in appreciating why diamond and graphite, despite both being forms of carbon, have such different characteristics.