Allotropes of carbon: diamond, graphite, graphene and fullerenes
A pencil "lead" and the stone in an engagement ring are both made of pure carbon. One smears on paper and the other can cut glass. How can the same element behave so differently? The answer lies in how the atoms are joined together, and we will come back to it at the end of the page.
What is an allotrope?
Carbon has several allotropes: diamond, graphite, graphene and the fullerenes (including buckminsterfullerene, $\mathrm{C_{60}}$, and carbon nanotubes). Because the bonding is different, their physical properties differ greatly. Their chemical formula is the same: just C, with every atom in oxidation state 0.
Each carbon atom has four outer electrons. What it does with them decides the allotrope:
| Allotrope | Bonds per C atom | Hybridisation | Delocalised electrons per atom |
|---|---|---|---|
| Diamond | 4 sigma | $sp^3$ | 0 |
| Graphite | 3 sigma (+ 1 delocalised) | $sp^2$ | 1 |
| Graphene | 3 sigma (+ 1 delocalised) | $sp^2$ | 1 |
| Fullerene $\mathrm{C_{60}}$ | 3 sigma (+ 1 delocalised) | $sp^2$ (approx.) | 1 |
a) Define the term allotrope. b) Name four allotropes of carbon. c) State the type of bonding between atoms within each layer of graphite.
- a) Define the term allotrope.
- b) Name four allotropes of carbon.
- c) State the type of bonding between atoms within each layer of graphite.
Diamond and graphite
Diamond
Each carbon atom is bonded to four others by strong covalent bonds, pointing to the corners of a tetrahedron (bond angle 109.5°). The carbon atoms are $sp^3$ hybridised. This builds a continuous giant covalent lattice in all three directions.
Consequences:
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Very hard and very high melting point: a huge number of strong bonds must be broken.
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Does not conduct electricity: all four outer electrons are held in localised bonds, so there are no mobile charge carriers.
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Excellent thermal conductor and transparent.
Uses: cutting and drilling tools (drill-bit tips for rock), jewellery.
Graphite
Each carbon atom is bonded to three others by strong covalent bonds in flat layers of hexagons ($sp^2$ hybridised, bond angle 120°). The fourth electron on each atom is delocalised across the layer. The layers are held together only by weak London dispersion forces.
Consequences:
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Conducts electricity (along the layers): the delocalised electrons are free to move. This makes graphite a good electrode material, and it is also unreactive and has a high melting point.
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Soft and slippery: the layers slide over each other easily because the forces between them are weak. This makes it useful as a dry lubricant and in pencils.
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Very high melting point: melting would require breaking the strong covalent bonds within the layers, not just the weak forces between them.
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Less dense than diamond: the layers are relatively far apart (held by weak forces), so the structure is more open.
Diamond and graphite are both pure carbon. a) State the number of covalent bonds formed by each carbon atom in diamond and the hybridisation of the carbon atoms. b)…
Diamond and graphite are both pure carbon.
- a) State the number of covalent bonds formed by each carbon atom in diamond and the hybridisation of the carbon atoms.
- b) State the number of covalent bonds formed by each carbon atom in graphite.
- c) Explain why graphite conducts electricity but diamond does not.
Graphite is used as a dry lubricant and as electrodes in electrolysis. a) Describe the structure of graphite, including the forces between layers. b) Explain why…
Graphite is used as a dry lubricant and as electrodes in electrolysis.
- a) Describe the structure of graphite, including the forces between layers.
- b) Explain why graphite is a good lubricant.
- c) Explain why graphite is a good electrode material.
- d) Explain why graphite has a very high melting point even though it is soft.
Graphene
Graphene is a single layer of graphite: one atom thick, with carbon atoms in a hexagonal "honeycomb" lattice. Each C is $sp^2$ hybridised, with three sigma bonds and one delocalised electron.
Properties and uses:
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Excellent electrical conductor (delocalised electrons): electronics, flexible transparent touchscreens.
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Extremely strong for its mass and very light: composites for sports equipment.
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Almost transparent and flexible.
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Good thermal conductor.
Match each use to the most suitable allotrope (diamond, graphite, graphene, fullerene), giving a reason in each case. a) The tip of a drill bit for cutting rock. b) The…
Match each use to the most suitable allotrope (diamond, graphite, graphene, fullerene), giving a reason in each case.
- a) The tip of a drill bit for cutting rock.
- b) The "lead" in a pencil.
- c) A flexible, transparent conducting layer for a touchscreen.
- d) A hollow cage molecule investigated for drug delivery.
Graphene is a single layer of carbon atoms. A sheet of graphene has a mass of $0.77\,\text{mg}\,\text{m}^{-2}$. a) Describe the structure of graphene. b) Calculate the…
Graphene is a single layer of carbon atoms. A sheet of graphene has a mass of $0.77\,\text{mg}\,\text{m}^{-2}$.
- a) Describe the structure of graphene.
- b) Calculate the number of carbon atoms in $1.00\,\text{m}^2$ of graphene ($A_r$ C = 12, $L = 6.02\times10^{23}\,\text{mol}^{-1}$).
- c) Suggest, with reasons, two properties that make graphene useful in electronics or composites.
Fullerenes and nanotubes
Fullerenes are hollow cage-like molecules of carbon. The best known is buckminsterfullerene, $\mathrm{C_{60}}$: 60 carbon atoms forming a sphere of 12 pentagons and 20 hexagons, like a football. Each atom has three sigma bonds and one delocalised electron.
Carbon nanotubes can be pictured as a sheet of graphene rolled into a cylinder.
Key difference: $\mathrm{C_{60}}$ is a simple molecular substance. Its molecules are held together by weak London dispersion forces, so it is soft, has a much lower melting point than diamond (it sublimes at a far lower temperature), and does not conduct well in the solid state. Its relative molecular mass is $60 \times 12 = 720$.
Nanotubes:
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Very high tensile strength along the axis: the strong covalent bonds of the hexagonal lattice are all aligned along the tube, and there are no weak layers to slip.
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Conduct electricity because of the delocalised electrons.
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Uses: strong lightweight composites, electronics, and research into filtration membranes. Fullerenes are investigated for drug delivery because their hollow cage can carry molecules.
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Concern: nanomaterials released into the environment may be toxic or accumulate in living organisms, and their long-term effects are not well known.
The buckminsterfullerene molecule has formula $\mathrm{C_{60}}$. a) Calculate its relative molecular mass ($A_r$ C = 12). b) Describe its shape and the arrangement of…
The buckminsterfullerene molecule has formula $\mathrm{C_{60}}$.
- a) Calculate its relative molecular mass ($A_r$ C = 12).
- b) Describe its shape and the arrangement of the carbon atoms.
- c) Explain why $\mathrm{C_{60}}$ is soft and has a much lower melting point than diamond.
- d) Calculate the number of molecules in $1.00\,\text{g}$ of $\mathrm{C_{60}}$ ($L = 6.02\times10^{23}\,\text{mol}^{-1}$).
Carbon nanotubes can be viewed as rolled-up sheets of graphene. a) Explain why a nanotube has very high tensile strength along its axis. b) Explain why nanotubes conduct…
Carbon nanotubes can be viewed as rolled-up sheets of graphene.
- a) Explain why a nanotube has very high tensile strength along its axis.
- b) Explain why nanotubes conduct electricity.
- c) A research group proposes using nanotubes in a water-filtration membrane. Suggest one advantage of this and one concern about releasing nanomaterials into the environment.
- d) Explain why $\mathrm{C_{60}}$, graphene and nanotubes are all described as having carbon in the same oxidation state, 0, and yet different properties.
Stability, density and structure
Graphite is less dense than diamond because its layers are spaced widely apart, leaving more empty space. Diamond's tetrahedral network packs atoms more closely.
At room temperature graphite is the more stable allotrope, but diamond does not turn into it. The change would require breaking very strong covalent bonds, so the activation energy is enormous and the rate is effectively zero. Diamond is stable only in the sense of being kinetically unreactive.
The densities of diamond and graphite are $3.51\,\text{g}\,\text{cm}^{-3}$ and $2.27\,\text{g}\,\text{cm}^{-3}$ respectively. a) Explain, in terms of structure, why…
The densities of diamond and graphite are $3.51\,\text{g}\,\text{cm}^{-3}$ and $2.27\,\text{g}\,\text{cm}^{-3}$ respectively.
- a) Explain, in terms of structure, why graphite is less dense than diamond.
- b) Calculate the volume occupied by $12.0\,\text{g}$ of each allotrope.
- c) The conversion of graphite to diamond requires very high pressure. Use the densities to explain why high pressure favours diamond.
- d) Explain why diamond does not spontaneously convert to graphite at room temperature even though graphite is more stable.
Compare the structure and properties of diamond, graphite and graphene in a table or in prose, addressing: a) hybridisation of carbon and number of delocalised electrons…
Compare the structure and properties of diamond, graphite and graphene in a table or in prose, addressing:
- a) hybridisation of carbon and number of delocalised electrons per atom;
- b) electrical conductivity;
- c) hardness and the reason for it;
- d) one application for each, justified by a property.