States of Matter: Question 9

Worked solution

Part (a): Structure and bonding in diamond

In diamond, every carbon atom uses all four of its outer-shell electrons to form four strong covalent bonds to four neighbouring carbon atoms, arranged tetrahedrally around it. This pattern of bonding is repeated throughout the entire crystal, producing a rigid, three-dimensional giant covalent network in which there are no separate molecules. The whole crystal is, in effect, one enormous covalently bonded structure.

Part (b): Structure and bonding in graphite

In graphite, each carbon atom forms strong covalent bonds to only three neighbouring carbon atoms, arranged in flat, two-dimensional hexagonal (honeycomb) layers. This uses three of each carbon atom’s four outer-shell electrons; the fourth electron from each atom is delocalised, free to move within the plane of its layer. The layers themselves are held together only by weak van der Waals forces of attraction, with a relatively large separation between adjacent layers compared with the short covalent bond length within a layer.

Part (c): Why graphite conducts but diamond does not

Electrical conduction in a solid requires charge carriers (such as electrons) that are free to move through the structure. In graphite, each carbon atom contributes one delocalised electron that is free to move along the plane of its hexagonal layer, so graphite can conduct electricity along the layers. In diamond, however, every one of each carbon atom’s four outer-shell electrons is used in a localised covalent bond to a specific neighbouring atom, there are no delocalised or mobile electrons anywhere in the structure, so diamond cannot conduct electricity.

Part (d): Why graphite is much softer than diamond

Diamond’s hardness comes from its rigid, three-dimensional network of strong covalent bonds extending equally in every direction: breaking or distorting this network (for example, by scratching it) would require breaking many strong covalent bonds at once, which needs a very large amount of energy. Graphite, in contrast, consists of strong covalent bonds only within each flat layer; the separate layers are held together by much weaker van der Waals forces. Applying a force to graphite allows the layers to slide over one another with relatively little energy needed, since only the weak forces between layers (not the strong covalent bonds within a layer) need to be overcome. This is why graphite is soft, slippery, and useful as a lubricant, despite containing covalent bonds just as strong as diamond’s within each layer.

Final answers

• (a) Diamond: each carbon atom covalently bonded to four others in a rigid, three-dimensional tetrahedral network
• (b) Graphite: each carbon atom covalently bonded to three others within flat hexagonal layers, with weak van der Waals forces between layers
• (c) Graphite has one delocalised electron per carbon atom, free to move along each layer and carry charge; diamond has no delocalised electrons because all four outer-shell electrons per atom are used in localised covalent bonds
• (d) Graphite’s layers, held together only by weak van der Waals forces, can slide over one another; diamond’s rigid three-dimensional network of strong covalent bonds in all directions cannot