Diamond 1: introduction

To go directly to the unit cell structure, click the link to page 5 below.

Diamond is an allotrope of carbon and a remarkable mineral. Its name comes from the Greek for 'invincible'. Nearly half of all diamonds come from central and southern Africa, where they are mined from volcanic pipes deep in the Earth: carbon-based materials need to be exposed to high temperatures and pressures for diamond crystals to form. It is not only the hardest mineral, but is four times harder than the second hardest, corundum (aluminium oxide), varieties of which are ruby and sapphire. Its hardness can be accounted for by its short and strong carbon-carbon covalent bonds. Indeed it needs to be heated to an extremely high temperature before it sublimes (about 4000 K). Diamond has the highest lattice density of any substance in terms of number of atoms per unit volume. The combination of the strong bonds, high lattice density and high symmetry means that diamond is also the best conductor of heat of any substance, five times better than silver, the best metallic conductor - despite the fact that diamond is an electrical insulator! Heat energy, carried by vibrations of its atoms, is dispersed very rapidly and efficiently through its lattice structure.

Despite its hardness, diamond will split if it receives a sharp blow in the direction of one of its four cleavage planes, and so gemstones tend to be mounted so that these planes are not in a position to be struck. The high refractive index of diamond gives crystals their sparkle. This, combined with its hardness, gives gemstones their great value. Crystals often possess cubic characteristics - not surprising, given its cubic unit cell. Diamond is transparent from the ultraviolet all the way through to the far infrared - by far the largest range of any substance. Diamonds often appear coloured though, due to impurities and crystal defects. A common impurity is nitrogen, which imparts a yellow colouration. When not being used as gemstones diamonds (usually synthetic ones) are used as abrasives on account of their hardness. It is amazing how different diamond is compared to carbon's other allotropes, graphite and Buckminsterfullerene (see those pages).

To the left is the structure of diamond. It is a giant covalent network, with each carbon connected to four neighbours, each with a single covalent bond. These four neighbours are in a tetrahedral arrangement.

In fact, orbitals on the atoms overlap to form bonds so the structure isn't actually the most realistic possible representation. It is however, a useful representation for seeing inside the structure. A more realistic representation is shown on the next page.

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