Giant Covalent Structures

Introduction to Covalent Bonds

Formation of Covalent Compounds: Carbon

Carbon, a non-metal, forms covalent compounds by sharing its four valence electrons.
Carbon can form several types of giant covalent structures including diamond, graphite, and fullerenes.

Giant Covalent Structures: Diamond

Each carbon atom in diamond is covalently bonded to four other carbon atoms in a tetrahedral arrangement.
Diamonds are extremely strong and hard due to the presence of these strong covalent bonds in all directions.
Diamond does not conduct electricity as all of the electrons are held in localised bonds and are thus not free to move and carry charge.

Giant Covalent Structures: Graphite

In graphite, each carbon atom is bonded to three others, forming layers of hexagonal structures.
Graphite can conduct electricity along its layers. While three of each carbon atom’s electrons form covalent bonds, the fourth is delocalised and is free to move and carry charge. Thus, graphite is a good conductor of electricity.
The layers of graphite can slide over each other because the forces between the layers are much weaker than the covalent bonds within the layers. These properties make graphite soft and slippery, ideal for applications such as pencil “lead” and lubricants.

Giant Covalent Structures: Fullerenes (including Buckminsterfullerene and Graphene)

Fullerenes are giant covalent structures that include hexagonal, pentagonal, or heptagonal rings of carbon atoms.
The most famous fullerene is buckminsterfullerene (C60), which has a structure shaped like a football made up of twenty hexagonal and twelve pentagonal rings. It can be considered a “molecular crystal”.
Graphene is a single layer of graphite forming a lattice of hexagonal rings. It is noteworthy for its exceptional strength and conductivity.

General Properties of Giant Covalent Structures

Giant covalent structures have high melting and boiling points because all the atoms are linked by strong covalent bonds, which take a large amount of energy to break.
They do not generally conduct electricity (with the notable exception of graphite and graphene) because there are usually no free electrons or ions to carry charge.

Limitations of Dot-and-Cross Diagrams

Just like ionic structures, dot-and-cross diagrams are used to display the covalent bonding in compounds by showcasing the shared pair of electrons between atoms.
Despite their usage, these diagrams fail to illustrate the actual arrangement of atoms in a giant covalent structure.
These diagrams also don’t signify the high melting and boiling points which result from the large number of strong covalent bonds present in the structure.