Giant Covalent Structures Giant Covalent Structures: Diamond, Graphite, Graphene and Silica

A giant covalent (macromolecular) structure has no separate molecules. Millions of atoms are joined by strong covalent bonds into one continuous lattice, so the whole solid is essentially one molecule. Breaking it down means breaking many strong bonds, which needs a lot of energy. That is why these solids have very high melting points, are hard, and do not dissolve in water. Carbon and silicon are the usual elements.

In diamond each carbon atom forms 4 covalent bonds to 4 other carbons in a tetrahedral, three-dimensional network. There are no weak points, so diamond is the hardest natural substance and melts above 3500 degrees C. It does not conduct electricity because all 4 outer electrons are locked in bonds, leaving none free to move. It is used in cutting tools and drill tips.

In graphite each carbon bonds to only 3 others, forming flat hexagonal layers. The spare 4th electron from each atom becomes delocalised, so graphite conducts electricity. The layers are held together by weak forces and slide over each other, making graphite soft and a good lubricant. Graphene is just one of those layers, one atom thick: it is the strongest material known and an excellent conductor.

Silica (SiO2) is found in sand, quartz and glass. Its formula is just a ratio, not a small molecule: each silicon bonds to 4 oxygens and each oxygen bridges 2 silicons, building a giant covalent lattice like diamond. So silica is hard, has a high melting point (about 1700 degrees C), is insoluble and does not conduct. It is used to make glass and optical fibres.