Graphite — AQA GCSE Chemistry

Graphite is a giant covalent structure of carbon. Each carbon atom bonds covalently to three other carbon atoms, leaving one outer electron per atom not used in bonding. The bonded atoms arrange into flat layers of hexagonal rings, like a honeycomb. Within a layer, each carbon has three single covalent bonds, so the layer is a continuous two-dimensional network. Between layers there are no covalent bonds; the layers are held together only by weak intermolecular forces. This explains why graphite is soft and slippery: the layers can slide over one another because the forces between them are weak, even though the covalent bonds within each layer are strong. For example, this sliding allows graphite to be used as a lubricant and in pencil leads.

In graphite each carbon atom forms three covalent bonds, using three of its four outer electrons. The remaining outer electron from each carbon atom becomes delocalised. These delocalised electrons are free to move throughout each layer of hexagonal rings, so graphite conducts electricity. The layers themselves are held together by weak intermolecular forces, allowing them to slide. This is why graphite is used for electrodes in electrolysis and in batteries. For example, a graphite rod conducts electricity because the delocalised electrons can carry charge through the structure. The number of delocalised electrons is one per carbon atom, so the ratio of carbon atoms to delocalised electrons is 1:1.

Graphite is a giant covalent structure in which each carbon atom forms three covalent bonds to three other carbon atoms, producing layers of hexagonal rings with no covalent bonds between the layers. Within a layer the covalent bonds are strong, so graphite has a very high melting point because many strong covalent bonds must be broken. Between layers there are only weak forces, so the layers can slide over one another, making graphite soft and slippery and useful as a lubricant and in pencil leads. One electron from each carbon atom is delocalised, so graphite conducts electricity and is used for electrodes. For example, graphite conducts electricity but diamond does not, because in diamond all outer electrons are localised in bonds.

Graphite is a giant covalent structure of carbon atoms, each bonded to three others in flat hexagonal layers. Each carbon uses three of its four outer electrons in covalent bonds, so one electron per atom is left free to move throughout the layer. These delocalised electrons are the key similarity to metals, because metals also have delocalised outer electrons that are not tied to a particular atom. In both materials, the mobile electrons can carry charge and transfer energy, so graphite conducts electricity and conducts heat along its layers. Unlike metals, graphite is not malleable in the same way, because its layers are held by weak forces and can slide, and it is a non-metal with covalent bonding within layers. A useful check is to compare graphite with copper: both conduct, but only graphite has strong covalent bonds inside its layers.