Giant Covalent Structure – IGCSE Chemistry Definition

Giant covalent structures are three-dimensional lattices where every atom is bonded to its neighbours by strong covalent bonds. The 0620 syllabus focuses on three examples — diamond, graphite, and silicon dioxide — and expects you to explain their properties in terms of their structure and bonding. This is a high-value topic: comparison questions between diamond and graphite appear frequently on Papers 2 and 4.

The 0620 definition

A giant covalent structure is a lattice of atoms held together by strong covalent bonds extending throughout the structure. There are no individual molecules — the entire structure is one giant network.

The three key examples

Diamond (carbon)

• Each carbon atom is covalently bonded to four other carbon atoms in a tetrahedral arrangement
• Very hard — strong covalent bonds in all directions
• Very high melting point — many strong covalent bonds to break
• Does not conduct electricity — all four outer electrons are used in bonding, none are free to move

Graphite (carbon)

• Each carbon atom is bonded to three other carbon atoms in flat hexagonal layers
• The fourth outer electron from each carbon is delocalised between the layers
• Layers are held together by weak intermolecular forces
• Slippery/lubricant — layers slide over each other
• Conducts electricity — delocalised electrons can move along the layers
• High melting point — strong covalent bonds within the layers

Silicon dioxide (SiO₂)

• Each silicon atom is bonded to four oxygen atoms and each oxygen is bonded to two silicon atoms
• Very high melting point — giant network of strong covalent bonds
• Hard — strong bonds in all directions
• Does not conduct electricity

Diamond vs graphite comparison

Property	Diamond	Graphite
Bonds per carbon	4 covalent bonds	3 covalent bonds (+ 1 delocalised electron)
Structure	Tetrahedral 3D lattice	Flat hexagonal layers
Hardness	Very hard	Soft, slippery
Electrical conductivity	Does not conduct	Conducts (delocalised electrons)
Melting point	Very high	Very high
Use	Cutting tools, jewellery	Lubricant, electrodes, pencils

Worked exam question

Explain why graphite conducts electricity but diamond does not. (3)

Mark scheme

In graphite, each carbon is bonded to only three other carbons [1]; the fourth outer electron is delocalised / free to move along the layers [1]; in diamond, each carbon is bonded to four other carbons and all outer electrons are used in covalent bonds / no free electrons [1]

Common exam mistakes

• Saying diamond has “no bonds” because it does not conduct electricity. Diamond has many strong covalent bonds — it does not conduct because there are no free or delocalised electrons.
• Forgetting to explain why graphite is slippery. State that the layers are held together by weak intermolecular forces, so layers can slide over each other.
• Confusing giant covalent with simple molecular. Simple molecular substances (like H₂O) have weak forces between molecules and low melting points. Giant covalent substances have strong bonds throughout and very high melting points.