Key concepts in chemistry Study Guide

Diamond and graphite are both allotropes of carbon, meaning they are different structural forms of the same element. Their contrasting physical properties arise directly from their distinct arrangements of carbon atoms. Diamond has a giant covalent structure where each carbon atom is covalently bonded to four other carbon atoms in a tetrahedral arrangement. This forms a rigid, three-dimensional network, making diamond exceptionally hard and giving it a very high melting point. As all valence electrons are involved in bonding, diamond does not conduct electricity. Graphite, on the other hand, also has a giant covalent structure but consists of layers of carbon atoms. Within each layer, each carbon atom is covalently bonded to three other carbon atoms, forming hexagonal rings. These layers are held together by weak intermolecular forces. The remaining unbonded valence electron from each carbon atom becomes delocalised within its layer. This layered structure allows graphite to be soft and slippery, as the layers can slide over each other. The delocalised electrons enable graphite to conduct electricity and heat efficiently along its layers.

Graphite and diamond are allotropes of carbon, meaning they are different structural forms of the same element. Their vastly different properties arise from their distinct atomic structures and bonding. In diamond, each carbon atom is covalently bonded to four other carbon atoms in a tetrahedral arrangement, forming a giant covalent lattice. This strong, rigid 3D structure makes diamond exceptionally hard, explaining its use in cutting tools. All valence electrons are involved in these strong covalent bonds, so diamond does not conduct electricity. In contrast, graphite consists of carbon atoms arranged in hexagonal layers. Within each layer, each carbon atom is covalently bonded to three other carbon atoms, leaving one delocalised electron per atom. These delocalised electrons are free to move between the layers, making graphite an electrical conductor, suitable for electrodes. The forces between the layers are weak intermolecular forces (van der Waals forces), allowing the layers to slide easily over each other. This enables graphite's use as a lubricant.