Allotropes of Carbon: GCSE Chemistry

What is an allotrope?

An allotrope is a different structural form of the same element. Allotropes of carbon all contain only carbon atoms, but those atoms are bonded together in different ways. Because the physical and chemical properties of a substance depend on its structure and bonding, each allotrope has a completely different set of properties — even though the chemical composition is identical (100% carbon).

This makes carbon uniquely versatile and is why understanding allotropes is a central GCSE chemistry topic covering giant covalent structures, simple molecular structures, and delocalised electrons.

What is the structure of diamond and why is it so hard?

Diamond has a giant covalent structure. Each carbon atom forms four covalent bonds to four neighbouring carbon atoms in a tetrahedral arrangement, creating a continuous three-dimensional lattice with no free edges or layers.

Key properties that follow from this structure:

• Extremely hard — every atom is held in place by four strong covalent bonds; there are no weak points where layers could slip.
• Very high melting point (≈3550 °C) — enormous energy is required to break the strong covalent bonds throughout the lattice.
• Does not conduct electricity — all four outer electrons of every carbon atom are used in covalent bonds; there are no free (delocalised) electrons to carry charge.
• Transparent — no free electrons to absorb visible light.

Uses: gemstones, industrial cutting and grinding tools (drill bits, saw blades), polishing compounds.

What is the structure of graphite and why is it unusual?

Graphite also has a giant covalent structure, but a very different one. Each carbon atom forms three covalent bonds in a flat, hexagonal arrangement. These hexagonal rings join to form flat sheets called graphene layers. The layers are held together by weak intermolecular (van der Waals) forces, and each carbon has one electron left over — these electrons are delocalised between the layers.

Key properties:

• Soft and slippery — the weak inter-layer forces allow layers to slide over each other easily (used as a dry lubricant).
• Very high melting point — strong covalent bonds within each layer require a lot of energy to break.
• Conducts electricity — the delocalised electrons between layers are free to move and carry charge. (Graphite is the only non-metal that conducts electricity well under normal conditions.)
• Opaque and black — free electrons absorb all visible light.

Uses: pencil "leads" (graphite mixed with clay), lubricants, electrodes in electrolysis, electric motor brushes.

What is graphene and why does it matter?

Graphene is a single atomic layer of graphite — one sheet of carbon atoms arranged in a hexagonal lattice, exactly one atom thick. It was first isolated in 2004 by scientists at the University of Manchester (Geim and Novoselov received the Nobel Prize in Physics 2010 for this).

Properties of graphene:

• Extremely strong — despite being one atom thick, graphene is stronger than steel for its mass due to its in-plane covalent bonds.
• Excellent electrical conductor — delocalised electrons move freely across the sheet.
• Transparent — absorbs only about 2% of visible light.
• Very light — it is essentially two-dimensional.

Potential uses: flexible electronics, ultrastrong composites, faster transistors, biomedical sensors, water filtration membranes.

What is buckminsterfullerene (C₆₀)?

Buckminsterfullerene contains 60 carbon atoms arranged in a hollow sphere made of 12 pentagons and 20 hexagons — the same pattern as a football. It was discovered in 1985. Unlike diamond and graphite, C₆₀ has a simple molecular structure (not giant covalent), so the molecules are held together only by weak van der Waals forces.

Key properties:

• Low melting point (≈600 °C) compared to diamond or graphite — only weak intermolecular forces need to be overcome, not covalent bonds.
• Poor electrical conductor under normal conditions (electrons are not delocalised between molecules), but under certain conditions modified fullerenes can be made to conduct.
• Soluble in organic solvents such as benzene (unlike diamond or graphite).

Uses: potential drug-delivery vehicles (hollow interior can carry molecules), lubricants, superconductors (when doped with alkali metals), nanotechnology research.

Other fullerene shapes include carbon nanotubes (cylindrical sheets of graphene), which combine the strength and conductivity of graphene in a tubular form.

Comparison table of the four carbon allotropes

How should you answer allotrope exam questions?

Exam questions on this topic almost always ask you to link structure to property. The key chain is:

1. Identify the type of structure (giant covalent or simple molecular).
2. State the bonding within and between particles.
3. Explain the property that results (e.g. high melting point because many strong covalent bonds must be broken; conducts electricity because delocalised electrons can move).

Worked example: "Graphite conducts electricity but diamond does not. Explain this in terms of their structures."

• In graphite, each carbon atom forms three covalent bonds, leaving one electron per atom delocalised between the layers. These delocalised electrons can carry charge through the material.
• In diamond, each carbon forms four covalent bonds; all outer electrons are used in bonding, so there are no delocalised electrons and no charge carriers — diamond does not conduct.

Frequently asked questions

Why does graphite have a higher melting point than C₆₀ if both are forms of carbon?

Graphite has a giant covalent structure, meaning that to melt it you must break the very strong covalent bonds within each graphene layer — this requires enormous energy (~3650 °C). C₆₀ is a simple molecular substance; melting only requires overcoming the weak van der Waals forces between whole C₆₀ molecules, leaving the bonds inside each molecule intact. The intramolecular bonds do not need to be broken to melt a simple molecular substance, which is why C₆₀ melts at only around 600 °C.

Why can graphite be used as a lubricant but diamond cannot?

In graphite, the flat layers are held together only by weak van der Waals forces. When a shear force is applied — for example, in a bearing or between sliding surfaces — the layers slide over each other easily, reducing friction. In diamond, every atom is locked into a rigid three-dimensional lattice by four strong covalent bonds in all directions. There are no planes along which the structure can slip, so diamond is rigid and cannot act as a lubricant.

What makes graphene different from a sheet of graphite?

Graphene is a single isolated layer of carbon atoms; graphite is a stack of many graphene layers held together by van der Waals forces. In graphite, the presence of the other layers slightly damps graphene's properties — for example, electron mobility is slightly lower in bulk graphite. An isolated graphene sheet has exceptional strength, near-perfect electron mobility, and optical transparency that bulk graphite does not exhibit to the same degree. The challenge of isolating a single layer (originally done with sticky tape) was the technical breakthrough that won the 2010 Nobel Prize.

Is C₆₀ the only fullerene?

No. C₆₀ (buckminsterfullerene) is the most famous, but fullerenes are a family of molecules. C₇₀ (70 carbon atoms in an elongated sphere) was discovered at the same time. Carbon nanotubes — cylinders of rolled graphene — are also classed as fullerenes. Researchers have made fullerenes ranging from C₂₀ to C₅₄₀ and beyond. C₆₀ remains the most studied because it can be produced in relatively large quantities and its spherical symmetry makes it useful as a model compound in chemistry research.

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Professor Curie at aitutors.me can walk you through each allotrope's particle model, drill you on structure-to-property chains, and help you write flawless six-mark answers on carbon allotropes.