Giant Covalent StructuresCambridge IGCSE Chemistry: Revision notes

Giant covalent structures (also called macromolecular structures) are networks of atoms bonded together by covalent bonds in a continuous three-dimensional lattice. Each atom is bonded to several other atoms through strong covalent bonds, extending throughout the entire structure. This is different from simple molecular structures, where only a few atoms are bonded together.

• Unlike ionic or metallic structures, there are no free ions or delocalised electrons to move throughout the structure
• The entire compound acts as one giant molecule
• Giant covalent structures have very high melting points because breaking them down requires breaking many strong covalent bonds
• Common examples include diamond, graphite, and silicon(IV) oxide (SiO₂)

Diamond is a giant covalent structure of carbon where each carbon atom forms four covalent bonds with other carbon atoms in a tetrahedral arrangement.

Structure of diamond:

• Each carbon atom is bonded to four other carbon atoms at the corners of a tetrahedron
• The bonding extends throughout the entire structure in a continuous network
• All bonds are strong covalent C–C bonds
• The structure is extremely rigid and compact

Properties related to structure:

• Very high melting point (~3823 K) – requires breaking thousands of strong covalent bonds
• Extremely hard – the tetrahedral arrangement and continuous bonding create a very strong, rigid structure
• Does not conduct electricity – there are no free electrons or ions to carry charge
• Insoluble in water – the covalent bonds cannot be broken by polar molecules
• Low density – despite being hard, the tetrahedral bonding is relatively spacious

Uses of diamond:

• Cutting tools and drill bits – hardness allows it to cut through other materials
• Abrasives – used in grinding and polishing applications
• Jewellery – brilliant appearance due to light refraction through the structure

Graphite is a giant covalent structure of carbon with a very different arrangement from diamond. Carbon atoms form layers in which each atom bonds to only three others.

Structure of graphite:

• Each carbon atom is bonded to three other carbon atoms in the same layer, forming a hexagonal pattern
• Layers are held together by weak van der Waals forces (not covalent bonds)
• Within layers, strong covalent C–C bonds create extended planes
• Layers can slide over one another easily due to weak intermolecular forces between them
• Each carbon atom has a delocalised electron that moves freely within the layer

Properties related to structure:

Uses of graphite:

• Lubricant – layers slide easily over each other, reducing friction between surfaces
• Electrodes – delocalised electrons can conduct electricity and allow reactions to occur at the electrode surface
• Pencil 'lead' – soft, slippery layers deposit easily on paper

Both diamond and graphite are allotropes of carbon (different forms of the same element), but their structures produce very different properties.

Why the differences?

• Hardness: Diamond's tetrahedral structure creates an interconnected 3D lattice with no weak points. Graphite's layers slide past each other easily.
• Electrical conductivity: Graphite has delocalised electrons within layers (one per atom from incomplete bonding); diamond has all electrons localised in covalent bonds.
• Lubrication: Only graphite lubricates because its layers slide; diamond is too hard.

Silicon(IV) oxide is a giant covalent structure where silicon and oxygen atoms are bonded together in a continuous 3D network. It is commonly found as quartz or sand.

Structure of SiO₂:

• Each silicon atom forms four covalent bonds to oxygen atoms (Si is in group 14, like C)
• Each oxygen atom forms two covalent bonds to silicon atoms
• The ratio is always Si:O = 1:2 (one silicon bonded to two oxygens, accounting for the formula SiO₂)
• The bonding extends throughout the entire structure in a continuous network
• No free electrons or ions exist

Structure representation:

• Can be represented as Si–O–Si–O–Si–O (continuous chain/network)
• Each Si is at the centre of a tetrahedron with four O atoms at the corners
• Each O bridges between two Si atoms

Properties of SiO₂:

• Very high melting point (~1973 K) – requires breaking strong Si–O covalent bonds throughout the structure
• Very hard and brittle – rigid continuous 3D network
• Insoluble in water – covalent bonds cannot be broken by polar molecules
• Non-conductor – no free electrons or ions
• Transparent/translucent – light can pass through the continuous structure without obstruction

Diamond and silicon(IV) oxide have very similar properties despite being different compounds. This is because their structures are fundamentally similar.

Structural similarities:

• Both are giant covalent structures extending continuously in 3D
• Both have each atom bonded tetrahedrally (four bonds per atom)
• Both have strong covalent bonding throughout the entire structure
• Neither has free electrons or ions

Property similarities:

Why are they similar? Both materials have the same type of bonding structure: strong covalent bonds throughout a rigid 3D lattice. Although diamond contains only carbon and SiO₂ contains silicon and oxygen, the tetrahedral bonding arrangement and continuous network create analogous properties.