Giant covalent structures contain very many atoms, each joined to adjacent atoms by covalent bonds. The atoms are usually arranged into giant regular lattices – extremely strong structures because of the many bonds involved.
The graphic shows the molecular structure of graphite and diamond (two allotropes of carbon).
Graphite is a form of carbon in which the carbon atoms form covalent bonds with three other carbon atoms. This means that each carbon atom has a ‘spare’ electron (as carbon has four outer electrons) which is delocalised between layers of carbon atoms. These layers can slide over each other, so graphite is much softer than diamond. It is used in pencils, and as a lubricant. Graphite conducts electricity due to the ‘spare’ electrons being delocalised between the layers. This conductivity makes graphite useful as electrodes for electrolysis.
However, graphite still has a very high melting and boiling point because the strong covalent bonds that hold the carbon atoms together in the layers require a lot of heat energy to break.
Diamond is a form of carbon in which each carbon atom is joined to four other carbon atoms, forming a giant covalent structure. As a result, diamond is very hard and has a high melting point. This explains why it is used in cutting tools. It does not conduct electricity as there are no delocalised electrons in the structure.
Silica (or silicon dioxide), which is found in sand, has a similar structure to diamond, so its properties are similar to diamond. It is hard and has a high melting point, but contains silicon and oxygen atoms, instead of carbon atoms.