Allotropes of Carbon: Diamond, Graphite, Graphene and Fullerenes

Carbon can exist in several different structures, all made only of carbon atoms. These different structural forms of one element are called allotropes. Diamond and graphite are the classic pair: both are pure carbon, yet one is the hardest natural material and the other is soft enough to leave a mark on paper.

The difference is entirely down to how the carbon atoms are bonded and arranged. This is the key idea that runs through the whole topic: structure determines properties, and properties determine uses.

Allotrope	Bonds per carbon	Structure	Key property
Diamond	4	Giant covalent, 3D	Very hard, no conduction
Graphite	3	Layers of hexagonal rings	Soft, slippery, conducts
Graphene	3	Single layer, one atom thick	Strong, conducts
Fullerenes (C60, nanotubes)	3	Hollow molecules of rings	Cage/tube shapes, many uses

Diamond and graphite are both pure carbon. Any difference in their properties must come from the difference in their structure and bonding, never from a difference in the atoms themselves.

Work through each form in turn, linking structure to property at every step.

In diamond, each carbon atom forms four covalent bonds to four other carbon atoms. This produces a giant covalent structure: a rigid three-dimensional network in which every atom is locked in place by four strong bonds.

That structure explains diamond's properties directly:

• Very hard. Every atom is held by four strong covalent bonds in a rigid lattice, so the structure resists being deformed or scratched. This is why diamond is used in cutting tools.
• Very high melting point. Melting means breaking a huge number of strong covalent bonds, which needs an enormous amount of energy.
• Does not conduct electricity. All four outer electrons on each carbon are used in covalent bonds, so there are no free (delocalised) electrons to carry charge.

Diamond does not conduct electricity because it has no delocalised electrons. Every outer electron on every carbon is held in a covalent bond.

The single sentence "each carbon forms four covalent bonds in a giant covalent structure" is the root of every diamond property, so learn it first and derive the rest from it.

In graphite, each carbon atom forms only three covalent bonds to three other carbon atoms. This arranges the atoms into flat layers of hexagonal rings (six-carbon rings). There are no covalent bonds between the layers — the layers are held together only by weak forces, so they can slide.

Because each carbon uses only three of its four outer electrons for bonding, the fourth electron on every atom becomes delocalised (free to move). Each carbon contributes one delocalised electron.

These two features explain graphite's properties:

• Soft and slippery. There are no bonds between the layers, so the layers slide over each other easily. This makes graphite useful as a lubricant and as pencil "lead".
• Conducts electricity and thermal energy. The one delocalised electron per atom is free to move through the layers and carry charge, similar to the delocalised electrons in a metal.
• High melting point. Melting still requires breaking the strong covalent bonds within the layers, which needs a large amount of energy.

Feature of graphite	Property it causes
No bonds between layers	Soft, slippery, good lubricant
One delocalised electron per atom	Conducts electricity and heat
Strong covalent bonds in layers	High melting point

Graphite conducts electricity because of its delocalised electrons, just like a metal, even though it is a non-metal. The exam often contrasts this directly with diamond, which has none.

Graphene is a single layer of graphite — just one atom thick. It is a sheet of carbon atoms arranged in the same hexagonal rings, with each carbon bonded to three others and one delocalised electron per atom.

Because it keeps graphite's bonding but is only one atom thick, graphene combines useful properties:

• It is very strong for its weight, as every atom is held by strong covalent bonds within the sheet.
• It conducts electricity, because of its delocalised electrons.
• It is almost transparent and extremely thin.

These properties make graphene useful in electronics (for example in fast, thin electronic devices) and in composites, where it is added to other materials to make them stronger. You need to be able to recognise graphene from a diagram and give a use.

Graphene = one layer of graphite. If you can describe graphite's bonding, you already know graphene's structure; it is simply a single sheet of it.

Fullerenes are molecules of carbon with hollow shapes. Their structures are based on hexagonal rings of carbon atoms, but they may also contain rings of five or seven carbon atoms, which allow the sheet to curve into a closed cage.

The first fullerene to be discovered was Buckminsterfullerene, which has the formula C60. It is a spherical molecule, shaped like a hollow ball (its shape resembles a football). Because fullerenes are individual molecules, they behave very differently from the giant structures diamond and graphite.

Carbon nanotubes are cylindrical fullerenes — long, hollow tubes of carbon. They have a very high length-to-diameter ratio, meaning they are extremely long compared with how wide they are. This shape, together with their bonding, gives them useful properties:

• They have a high tensile strength, so they can be used to reinforce materials (for example in strong, lightweight composites for tennis rackets or aircraft parts).
• They conduct electricity, so they are useful in electronics and nanotechnology.

Fullerene	Shape	Example use
Buckminsterfullerene (C60)	Hollow sphere	Delivering drugs, lubricants, catalysts
Carbon nanotube	Long hollow cylinder	Reinforcing materials, electronics

Uses of fullerenes you can quote include drug delivery, lubricants and catalysts (for C60), and reinforcing materials and electronics (for nanotubes). Naming a hollow, molecular cage or tube structure is what marks fullerenes out from the giant structures.

The whole topic reduces to one habit: state the structure and bonding, then read off the property. This table gathers the comparisons the exam draws most often.

Form	Bonds per C	Delocalised electrons?	Conducts electricity?	Hardness / feel
Diamond	4	No	No	Very hard
Graphite	3	Yes (1 per atom)	Yes	Soft, slippery
Graphene	3	Yes	Yes	Very strong, thin
Fullerenes	3	Yes	Yes (nanotubes)	Molecular cages/tubes

The single most common comparison is diamond versus graphite conduction:

• Diamond: 4 bonds per carbon → all outer electrons used in bonds → no delocalised electrons → does not conduct.
• Graphite: 3 bonds per carbon → one electron left over per atom → delocalised electrons → conducts.

Learn that contrast well and most exam questions on this topic fall into place.

1. Saying diamond conducts electricity

Diamond does not conduct. Every carbon forms four covalent bonds, so all its outer electrons are held in bonds and none are free to move. Only graphite, graphene and nanotubes conduct, because they have delocalised electrons.

2. Forgetting to say why graphite is soft

Graphite is soft because there are no covalent bonds between the layers, so the layers slide over each other. Do not say the covalent bonds are weak; the bonds within each layer are strong, and it is the gaps between layers that matter.

3. Confusing graphene and graphite

Graphene is a single layer, one atom thick. Graphite is many such layers stacked together. Describing graphene as "layers" loses the key point.

4. Getting the fullerene formula or shape wrong

Buckminsterfullerene is C60 and is a hollow sphere. Carbon nanotubes are hollow cylinders with a high length-to-diameter ratio. Do not mix up the sphere and the tube.

5. Treating fullerenes as giant structures

Fullerenes are molecules with hollow shapes, not giant covalent networks like diamond and graphite. That molecular, cage-like structure is exactly what the question is testing when it names fullerenes.