Structures of Carbon
Carbon is special
Carbon atoms can form four covalent bonds each, and can bond to other carbon atoms to build huge structures. This gives several different forms (allotropes) of carbon with very different properties, plus important carbon-based materials.
Diamond
- Each carbon atom is bonded to 4 others in a rigid giant covalent structure.
- Very hard and has a very high melting point (many strong covalent bonds to break).
- Does not conduct electricity — all four outer electrons are used in bonding, so there are no free electrons.
- Uses: cutting tools, drill tips.
Graphite
- Each carbon atom bonds to only 3 others, forming layers of hexagons.
- The layers are held together by weak forces, so they can slide — graphite is soft and slippery (used as a lubricant and in pencils).
- Each carbon has one delocalised electron, so graphite conducts electricity — unusual for a non-metal.
- Also has a high melting point (strong covalent bonds within layers).
Graphene
- A single layer of graphite — one atom thick.
- Extremely strong, light, and an excellent conductor of electricity.
- Promising for electronics and composite materials.
Fullerenes and nanotubes
- Fullerenes are carbon molecules shaped like hollow balls or tubes (e.g. buckminsterfullerene, C₆₀, a football-shaped molecule).
- Carbon nanotubes are cylinders of graphene — very strong with a high length-to-diameter ratio; good conductors.
- Uses: drug delivery, catalysts, strong lightweight materials, electronics.
Comparison table
| Form | Bonds per carbon | Conducts? | Key property |
|---|---|---|---|
| Diamond | 4 | No | Very hard |
| Graphite | 3 (+ layers) | Yes | Soft, slippery |
| Graphene | 3 (single layer) | Yes | Strong, thin |
| Fullerenes/nanotubes | 3 | Yes | Hollow, high surface area/strength |
Worked example
Explain why graphite conducts electricity but diamond does not.
- In graphite, each carbon bonds to only 3 others, leaving one delocalised electron per atom free to move and carry charge. In diamond, all 4 outer electrons are used in bonding, so there are no free electrons. ✓
Common mistakes
- Saying diamond conducts — it doesn't (no free electrons).
- Forgetting graphite conducts because of delocalised electrons (one per atom).
- Confusing graphene (one layer) with graphite (many layers).
Exam tips
- Always link the property to the structure and bonding (e.g. slippery ↔ sliding layers; conducts ↔ delocalised electrons).
- Learn the number of bonds per carbon for diamond (4) and graphite (3).
- Know one use for each form.
Key facts to remember
- Diamond: 4 bonds/carbon, hard, high mp, doesn't conduct.
- Graphite: 3 bonds/carbon in sliding layers, soft, conducts (delocalised electron).
- Graphene = single graphite layer (strong, conductive); fullerenes/nanotubes = hollow carbon structures with many uses.