How is graphite different from diamond?

Graphite and diamond differ in their structure, hardness, conductivity, and appearance due to their different atomic arrangements.

Graphite and diamond are both forms of carbon, known as allotropes, but they have different physical and chemical properties due to the way their carbon atoms are arranged. In diamond, each carbon atom is bonded to four other carbon atoms in a tetrahedral structure, forming a rigid three-dimensional network. This structure makes diamond extremely hard, the hardest known natural substance in fact, and gives it a high melting point. It also means that diamond does not conduct electricity, as all its electrons are involved in bonding and none are free to move.

On the other hand, in graphite, each carbon atom is bonded to three other carbon atoms, forming layers of hexagonal rings. These layers can slide over each other, which makes graphite soft and slippery – it's this property that makes it useful in pencils. The spare electron from each carbon atom is free to move, which means graphite can conduct electricity.

The different structures also give graphite and diamond different appearances. Diamond is transparent and sparkles because its rigid structure refracts light. Graphite, however, is black and has a metallic lustre because its loosely bonded layers absorb light.

So, while graphite and diamond are both made of carbon, their different atomic arrangements result in very different properties. This is a great example of how the arrangement of atoms in a substance can dramatically affect its properties.