How is the structure of diamond different from graphite?

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

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 that makes diamond extremely hard and gives it its characteristic sparkling appearance. The strong covalent bonds require a lot of energy to break, making diamond the hardest known natural substance. It's also an excellent thermal conductor because the vibrations of the atoms can travel easily through the rigid lattice.

On the other hand, in graphite, each carbon atom is covalently bonded to three other carbon atoms, forming layers of hexagonal structures. The layers are held together by weak Van der Waals forces, which allow them to slide over each other easily. This gives graphite its 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, meaning it can move freely within the layers, carrying an electric charge.

In summary, the key difference between diamond and graphite lies in their structure. Diamond's tetrahedral structure makes it hard and a good thermal conductor, while graphite's layered hexagonal structures make it slippery and a good electrical conductor. Understanding these differences is crucial in appreciating why these two allotropes of carbon have such different physical properties and uses.