Giant Covalent Structures — Covalent Bonding | IGCSE Chemistry

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Dot-and-cross diagrams and comparison of giant covalent vs simple molecular structures appear regularly.

A giant covalent structure (also called a macromolecule) contains a vast number of non-metal atoms joined by strong covalent bonds in one continuous lattice
Two classic examples, both built from carbon only:
- Diamond
- Graphite
Diamond and graphite are allotropes of carbon: different forms of the same element with different bonding arrangements

There are many strong covalent bonds running throughout the lattice
Melting or vaporising the solid requires breaking very large numbers of these bonds
That takes a great deal of thermal energy, so the melting point is high

Each carbon atom forms 4 single covalent bonds to four other carbon atoms
The result is a rigid 3D tetrahedral network extending throughout the crystal
Every outer electron on every carbon is locked into a covalent bond → no free electrons

Properties of diamond:

Very hard: the rigid 3D network of strong covalent bonds resists deformation. Diamond is used in cutting tools, drill bits and saw edges where extreme hardness matters
High melting point: the many strong covalent bonds running through the lattice need a large amount of energy to break
Does not conduct electricity: every outer electron is locked in a covalent bond, so there are no free electrons or ions to move and carry charge

Each carbon atom forms only 3 single covalent bonds with three other carbon atoms
These bonded atoms arrange themselves into flat layers of hexagons (sometimes called sheets)
Each carbon has 4 outer electrons in total; 3 are used in bonds and the 4th is a delocalised electron that can move freely along the layer
The hexagonal layers stack on top of each other, held together only by weak intermolecular forces between layers, not by covalent bonds

Properties of graphite:

Soft and slippery: the weak forces between layers let the layers slide easily over each other. Graphite is used as a dry lubricant and as the writing "lead" in pencils
Conducts electricity (and heat): the delocalised electrons can move along the layers and carry charge. Graphite is one of the very few non-metals that conducts electricity well
High melting point: within each layer the covalent bonds are strong, and a great deal of energy is needed to break enough of them for the structure to fall apart

An electric current is a flow of mobile charged particles
In diamond, every carbon uses all 4 of its outer electrons in covalent bonds. There are no free electrons available to move, so diamond is an electrical insulator
In graphite, every carbon uses only 3 of its 4 outer electrons in bonds. The 4th is delocalised and can move freely along the layer, carrying charge along with it. Graphite therefore conducts electricity