What are allotropes?
Allotropes are different forms of the same element in the same physical state. The atoms are the same; only the way they are joined is different.
Carbon is in Group 14 (old Group 4). It has 4 outer electrons, so each atom can share electrons to make up to 4 covalent bonds. A covalent bond is a shared pair of electrons, and it is very strong.
Because carbon can join in different patterns, we get diamond, graphite, graphene and fullerenes. Their properties are very different, even though all are pure carbon.
Diamond: a giant covalent structure
In diamond each carbon atom is joined to 4 other carbon atoms by strong covalent bonds. This makes a huge 3D network called a giant covalent structure.
- Very hard: to scratch it you must break many strong bonds. It is the hardest natural substance.
- Very high melting point (above 3500 °C): lots of energy is needed to break so many bonds.
- Does not conduct electricity: all 4 outer electrons are used in bonds, so there are no free electrons or ions to move.
Uses: cutting tools, drill tips, glass cutters and jewellery (it sparkles because it bends light strongly).
Graphite: layers that slide
In graphite each carbon atom bonds to only 3 others. The atoms form flat layers of hexagons (six-sided rings).
- Bonds inside a layer are strong covalent bonds, so graphite also has a very high melting point.
- Forces between layers are weak, so layers slide over each other. Graphite is soft and slippery.
- Each atom has one spare electron that is not in a bond. These delocalised electrons can move along the layers, so graphite conducts electricity and heat.
Uses: pencil leads, electrodes in electrolysis and batteries, and a dry lubricant for machines.
Graphene, fullerenes and nanotubes
Graphene is a single layer of graphite, just one atom thick. It is very strong for its weight, almost see-through, and an excellent conductor. Uses being developed: flexible electronics, better batteries, and stronger, lighter materials (composites).
Fullerenes are molecules of carbon shaped like hollow balls or tubes. They are made mainly of hexagons, with some pentagons (five-sided rings) or heptagons to make them curve. The first one found was buckminsterfullerene, C₆₀, which looks like a football.
Carbon nanotubes are fullerenes rolled into long thin cylinders. They have a very high length-to-width ratio, high tensile strength and good conductivity. Uses: strengthening sports gear, electronics and nanotechnology.
Fullerenes can carry drug molecules inside their cage (drug delivery), act as lubricants, and act as catalysts thanks to their large surface area.
Chemical properties of carbon
All allotropes are carbon, so they react in the same way. At room temperature carbon is quite unreactive.
- Burning in plenty of oxygen: C + O₂ → CO₂ (carbon dioxide). Even diamond burns if heated strongly in oxygen!
- Burning in little oxygen: 2C + O₂ → 2CO (carbon monoxide, a poisonous gas).
- As a reducing agent: hot carbon takes oxygen from some metal oxides, e.g. 2CuO + C → 2Cu + CO₂. This is how iron is made from iron ore in a blast furnace.
Burning a diamond and a lump of graphite gives the same gas, CO₂. This proves both are made of carbon only.
Key formulas and definitions
- Allotropes = forms of one element with different arrangements of atoms
- Diamond: 4 bonds per atom, giant 3D network → hard, high melting point, no conduction
- Graphite: 3 bonds per atom, hexagon layers + delocalised electrons → soft, conducts
- Graphene: one layer of graphite; fullerenes: hollow cages (C₆₀); nanotubes: rolled cylinders
- C + O₂ → CO₂ (plenty of air); 2C + O₂ → 2CO (little air)
- 2CuO + C → 2Cu + CO₂ (carbon as a reducing agent)
Worked examples
1. Both diamond and graphite have very high melting points, but only graphite conducts. Explain.
Both have giant structures held by many strong covalent bonds, which need a lot of energy to break, so both melt at very high temperatures. In diamond all 4 outer electrons are in bonds, so no charge can move. In graphite each atom uses only 3 electrons in bonds; the 4th is delocalised and moves along the layers, carrying current.
2. Why is graphite used as a lubricant but diamond is not?
Graphite's layers are held together only by weak forces, so they slide easily over each other and reduce friction. Diamond's atoms are locked in every direction by strong bonds, so nothing slides.
3. How many bonds are there in one C₆₀ molecule if each atom bonds to 3 others?
Each of the 60 atoms has 3 bonds: 60 × 3 = 180 bond ends. Each bond joins 2 atoms, so it is counted twice. Bonds = 180 ÷ 2 = 90.
Common mistakes
- Saying graphite conducts because of ions. Graphite has no ions; it conducts because of delocalised electrons.
- Saying layers in graphite slide because covalent bonds are weak. The bonds inside layers are strong; only the forces between layers are weak.
- Calling C₆₀ a giant structure. A fullerene is a molecule with a fixed number of atoms.
- Thinking allotropes are different elements or compounds. All carbon allotropes are pure carbon and burn to give only CO₂.