Giant Covalent Structures

This lesson covers giant covalent structures, as required by the AQA GCSE Chemistry specification (4.2.2). You need to understand the structures and properties of diamond, graphite, graphene, and fullerenes (including buckminsterfullerene and carbon nanotubes). These substances have very different properties from simple molecular substances because their covalent bonds extend throughout the entire structure.

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What Are Giant Covalent Structures?

A giant covalent structure (also called a macromolecular structure) is a three-dimensional network of atoms connected by strong covalent bonds in all directions. Unlike simple molecules, there are no individual molecules — the entire structure is one giant molecule.

Because all the bonds are strong covalent bonds, a huge amount of energy is needed to break them, giving these substances very high melting and boiling points.

Exam Tip: The key phrase for giant covalent structures is "many strong covalent bonds that require a large amount of energy to break." This directly explains the very high melting points. Always link the property to the bonding.

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Diamond

Diamond is an allotrope of carbon. An allotrope is a different structural form of the same element in the same physical state.

Structure of Diamond

• Each carbon atom forms four covalent bonds with four other carbon atoms.
• The bonds are arranged in a tetrahedral shape (109.5 degree bond angles).
• This creates a rigid, three-dimensional giant covalent lattice.
• There are no free electrons — all four outer electrons of each carbon atom are used in bonding.
• There are no weak intermolecular forces — only strong covalent bonds throughout.

Properties of Diamond

Property	Value / Description	Explanation
Melting point	3550 degrees C	Many strong covalent bonds must be broken, requiring a very large amount of energy.
Hardness	Very hard (hardest natural substance)	Rigid 3D network of strong covalent bonds in all directions.
Electrical conductivity	Does NOT conduct	All four outer electrons are involved in bonding — no free electrons or ions to carry charge.
Solubility	Insoluble in water	The strong covalent bonds cannot be overcome by water molecules.

graph TD
    A["Carbon Allotropes"] --> B["Diamond"]
    A --> C["Graphite"]
    A --> D["Graphene"]
    A --> E["Fullerenes"]
    B --> F["4 bonds per C<br/>Tetrahedral<br/>Very hard<br/>Does NOT conduct"]
    C --> G["3 bonds per C<br/>Layers<br/>Soft/slippery<br/>DOES conduct"]
    D --> H["3 bonds per C<br/>Single layer<br/>Strong and light<br/>DOES conduct"]
    E --> I["Hollow molecules<br/>or tubes<br/>e.g. C60, nanotubes"]

    style A fill:#2c3e50,color:#fff
    style B fill:#3498db,color:#fff
    style C fill:#27ae60,color:#fff
    style D fill:#e67e22,color:#fff
    style E fill:#8e44ad,color:#fff

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Graphite

Graphite is another allotrope of carbon. It has very different properties from diamond because its structure is different.

Structure of Graphite

• Each carbon atom forms three covalent bonds with three other carbon atoms.
• This creates flat layers (sheets) of hexagonally arranged carbon atoms.
• The fourth outer electron of each carbon atom is delocalised — it is free to move along the layers.
• The layers are held together by weak intermolecular forces (van der Waals forces).

Properties of Graphite

Property	Value / Description	Explanation
Melting point	3652 degrees C (sublimes at ~3642 under normal pressure)	Many strong covalent bonds within the layers must be broken.
Hardness	Soft and slippery	The weak intermolecular forces between layers allow them to slide over each other easily.
Electrical conductivity	DOES conduct electricity	Each carbon atom has one delocalised electron that is free to move along the layers, carrying charge.
Uses	Pencil lead, lubricant, electrodes	Softness (pencils, lubricants); conductivity (electrodes).

Exam Tip: Graphite is the ONLY giant covalent structure you are likely to encounter that conducts electricity. The reason is the delocalised electrons — one electron per carbon atom is free to move and carry charge. Make sure you explain this using the term "delocalised electrons," not just "free electrons."

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Graphene

Graphene is a single layer of graphite — a one-atom-thick sheet of carbon atoms arranged in hexagons. It was first isolated in 2004 by Andre Geim and Konstantin Novoselov at the University of Manchester.

Properties of Graphene

Property	Description	Explanation
Strength	Extremely strong (strongest material ever tested)	Strong covalent bonds between carbon atoms in the hexagonal layer.
Thickness	One atom thick	It is literally a single layer of atoms.
Electrical conductivity	Excellent conductor	Delocalised electrons are free to move across the sheet.
Flexibility	Very flexible	The single layer can bend without breaking.
Transparency	Transparent	Only one atom thick, so it allows light to pass through.

Uses of Graphene

Graphene has potential applications in:

• Electronics — transistors, touchscreens, flexible displays.
• Composites — making materials stronger and lighter.
• Sensors — detecting individual molecules.
• Energy — improving batteries and solar cells.
• Medical — drug delivery and biosensors.

Exam Tip: Graphene is described as a "wonder material" because of its extraordinary combination of properties: it is the thinnest, strongest, lightest, and most conductive material known. If asked about graphene, always mention that it is a single layer of graphite with strong covalent bonds and delocalised electrons.

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Fullerenes

Fullerenes are molecules of carbon shaped as hollow balls, tubes, or other shapes. They are NOT giant covalent structures — they are large but have a defined number of atoms.

Buckminsterfullerene (C60)

• Contains 60 carbon atoms arranged in a sphere of pentagons and hexagons (like a football / soccer ball).
• Named after architect Buckminster Fuller, who designed geodesic domes with a similar shape.
• Each carbon atom forms three covalent bonds.
• Has delocalised electrons on its surface.

Uses of Fullerenes