Compare the structures of Diamond and Graphite, making references to the bonding, the shape of the structures, and location of the electrons within the structures. Account for the fact that graphite conducts electricity and diamond does not.

In diamond, each carbon atom forms 4 strong covalent sigma bonds with other carbons, resulting in a tetrahedral 3D arrangement of atoms where all the electrons are fixed in place or localised. In graphite, each carbon forms 3 strong covalent sigma bonds with other carbons, resulting in lots of 2D layers with a hexagonal (or trigonal planar) arrangement of carbon atoms and weak forces of interaction between the two layers. As each carbon only forms 3 bonds, only 3 of the 4 outer shell electrons are fixed in place, which means one of the electrons from each carbon is delocalised. These delocalised electrons are free to move around and create a current, which is why graphite conducts electricity and diamond does not.