Why does graphite conduct electricity but diamond does not?

Graphite conducts electricity due to its free electrons in its structure, while diamond doesn't have any free electrons.

Graphite and diamond are both forms of carbon, known as allotropes. However, they have different structures and properties due to the way the carbon atoms are arranged. In graphite, each carbon atom is bonded to three other carbon atoms, forming layers of hexagonal structures. This leaves one electron from each carbon atom free to move. These free electrons are delocalised, meaning they can move throughout the layers of graphite. When a voltage is applied, these free electrons move, carrying charge from one place to another, and thus conducting electricity.

On the other hand, in diamond, each carbon atom is bonded to four other carbon atoms, forming a rigid tetrahedral structure. This means that all the electrons in diamond are involved in bonding and none are free to move. Therefore, diamond does not have any free electrons to carry charge and so it does not conduct electricity.

The difference in electrical conductivity between graphite and diamond is a great example of how the arrangement of atoms in a substance can significantly affect its properties. Even though graphite and diamond are both made of the same type of atom, their different structures result in very different properties. This is a key concept in chemistry, and it's particularly important when studying the properties of different materials.