Graphite 1: introduction

Graphite is one of the allotropes of carbon. Surprisingly, it is the most stable form of carbon. The only reason diamond doesn't spontaneously transform into graphite is because the activation energy is so high. Its properties are remarkably different to those of diamond: unlike diamond it is very soft, an electrical conductor and opaque. These differences can be accounted for by its structure and bonding. The unit cell is hexagonal (diamond's is cubic) and each carbon is strongly bonded in a planar network of tesselated hexagons. Only van der Waals' forces operate between the layers. There is electron delocalisation within the planes, which explains the electrical conductivity. The softness of graphite accounts for its use in pencils - indeed its name comes from the Greek "to draw". Graphite is an excellent lubricant, though recent research has shown that this is due to fluids between the layers such as water, which are adsorbed from the environment.

Even though graphite is soft, related "carbon fibre" is the high-strength component in many extremely strong composite materials. Graphite sheets are rolled up into fibres and the fibres twisted into threads, which are very strong owing to the strong carbon-carbon bond within the graphite sheets. Graphite is produced from the metamorphism of organic material in rocks. Indeed it may be considered the highest grade of coal, although it is not usually used as a fuel as it is difficult to ignite.

A fragment of the bulk structure of graphite is shown to the left. Each carbon atom in the bulk is bonded to three nearest neighbours within the planar sheets. It is evident that all the carbon atoms occupy equivalent positions in the lattice.

Go to page 2 to look at the unit cell within the bulk structure of graphite.

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