Giant covalent structures — AQA GCSE Chemistry

A giant covalent structure is a network of atoms joined by covalent bonds throughout the whole substance. Examples include diamond, silicon dioxide and graphite. Because the covalent bonds are strong and extend in three dimensions, melting the substance requires breaking many strong covalent bonds, not just overcoming intermolecular forces. This means giant covalent substances are solids with very high melting points. For example, diamond melts at over 3500 °C and silicon dioxide melts at about 1700 °C. Graphite also has a very high melting point because its carbon atoms are held in layers by strong covalent bonds. The high melting point is therefore a consequence of the strong covalent bonding throughout the giant structure, not of intermolecular forces between separate molecules.

A giant covalent structure is a continuous three-dimensional network in which every atom is joined to neighbouring atoms by covalent bonds. Each bond is a shared pair of electrons, so the network has no separate molecules and no overall fixed formula unit. In diamond, each carbon atom forms four single covalent bonds to four other carbon atoms in a tetrahedral arrangement, giving a rigid lattice. In graphite, each carbon atom bonds to three others in hexagonal layers, leaving one delocalised electron per atom. In silicon dioxide, each silicon atom bonds to four oxygen atoms and each oxygen atom bonds to two silicon atoms, so the ratio is 1:2 but the network is continuous. Because bonding extends throughout the structure, melting involves breaking many strong covalent bonds rather than separating small molecules.

Melting or boiling a giant covalent substance requires the covalent bonds in the network to be broken, not just weakened or separated. In a simple molecular substance, intermolecular forces between molecules are overcome and the covalent bonds inside each molecule stay intact, so melting points are low. In a giant covalent structure there are no separate molecules, so the strong covalent bonds themselves must be overcome throughout the lattice. This needs a large amount of energy, which is why diamond melts at over 3500 °C, silicon dioxide melts at about 1700 °C and graphite sublimes at very high temperature. The number and strength of covalent bonds broken per atom, and the rigidity of the network, determine how much energy is needed.

Diamond, graphite and silicon dioxide are the three named examples of giant covalent structures. Diamond and graphite are both made only of carbon atoms but have different arrangements, so they are allotropes with different properties. In diamond, each carbon atom forms four single covalent bonds in a rigid tetrahedral network, making diamond very hard and giving it a very high melting point. In graphite, each carbon atom forms three covalent bonds in hexagonal layers, with delocalised electrons between layers, so graphite conducts electricity and layers can slide. Silicon dioxide contains silicon and oxygen in a continuous network in which each silicon atom bonds to four oxygen atoms and each oxygen atom bonds to two silicon atoms. These examples show how the same type of bonding can produce different properties when the arrangement differs.

Recognising a giant covalent structure from a diagram means identifying a continuous network of atoms joined by covalent bonds, with no separate molecules and no finite repeating unit. In a diagram of diamond, each carbon atom is shown with four bonds extending in three dimensions to other carbon atoms. In graphite, each carbon atom is shown with three bonds within a flat hexagonal layer, and the layers are stacked with delocalised electrons between them. In silicon dioxide, silicon and oxygen atoms alternate in a continuous network, with each silicon atom bonded to four oxygen atoms and each oxygen atom bonded to two silicon atoms. A diagram of a simple molecular substance, by contrast, shows small groups of atoms with weak forces between the groups.