Covalent bonds between adjacent atoms throughout the crystal. They are also called giant molecules.
Covalent bonds are strong and directional in nature, therefore atoms are held very strongly at their positions. Such solids are very hard and brittle.
They have extremely high melting points and may even decompose before melting.
They are insulators and do not conduct electricity.
Graphite, diamond and silicon carbide are typical examples of such solids.
Graphite is soft and a conductor of electricity.
Its exceptional properties are due to its typical structure. Carbon atoms are arranged in different layers and each atom is covalently bonded to three of its neighbouring atoms in the same layer. The fourth valence electron of each atom is present between different layers and is free to move about. These free electrons make graphite a good conductor of electricity.
Different layers can slide one over the other. This makes graphite a soft solid and a good solid lubricant.
Carbon has an electronic configuration of 2,4. In diamond, each carbon shares electrons with four other carbon atoms forming four single bonds.
This is a giant covalent structure. It continues on and on in three dimensions. It is not a molecule, because the number of atoms joined up in a real diamond is completely variable depending on the size of the crystal.
Diamond has a very high melting point (almost 4000°C). Very strong carbon-carbon covalent bonds have to be broken throughout the structure before melting occurs.
Diamond is very hard. This is again due to the need to break very strong covalent bonds operating in three dimensions.
Diamond doesn’t conduct electricity. All the electrons are held tightly between the atoms, and aren’t free to move.
Diamond is insoluble in water and organic solvents. There are no possible attractions which could occur between solvent molecules and carbon atoms which could outweigh the attractions between the covalently bound carbon atoms.
Buckminster fullerene or fullerene:
In 1985 a new allotrope of carbon (C60) was discovered in carbon soot. Sixty carbon atoms form the shape of a ball like a football with a carbon atom at each corner of the 20 hexagons and 12 pentagons.
Each carbon atom (shown as a sphere) has three bonds.
The size of the molecule is almost exactly 1 nm in diameter.
These are not called giant molecules because there are only sixty atoms.
A large number of these molecules can fit together to form a transparent yellow solid called fullerite.
This form of carbon was named after the American architect Buckminster Fuller, who was famous for designing a large dome which looked similar (sort of) to the molecular structure of C60.
Many other balls of carbon called fullerenes, have since been made, including C70, C76, and C84. These molecules have become known as “buckyballs”.
Fullerenes are used as catalysts and lubricants. They are also used in nanotubes for strengthening materials (for example sports equipment) and are sometimes used as a way of delivering drugs into the body.