Structure, Properties and Carbon Allotropes
Properties of Ionic Compounds
Ionic compounds form giant ionic lattices. The strong electrostatic forces between the ions act in all directions, which gives ionic compounds a characteristic set of properties:
- High melting and boiling points — a large amount of energy is needed to overcome the many strong bonds acting throughout the lattice.
- Conduct electricity when molten or dissolved in water — the ions are then free to move and carry charge. They do not conduct when solid because the ions are fixed in position.
- Brittle — if an external force shifts one layer of ions, identical charges align and repel each other strongly, shattering the crystal.
Ionic compounds: high melting points (strong forces in all directions), conduct when molten or dissolved (ions free to move), do not conduct when solid (ions fixed), brittle (layers shift → repulsion).
Properties of Simple Molecular Substances
Simple molecules (e.g. \(H_2O\), \(CO_2\), \(O_2\)) have strong covalent bonds within the molecule but weak intermolecular forces between molecules. This gives them:
- Low melting and boiling points — little energy is needed to overcome the weak intermolecular forces (not the covalent bonds themselves).
- Do not conduct electricity — no free electrons or ions to carry charge.
As molecules get larger, their intermolecular forces become stronger, so boiling points increase with molecular size. This is why large hydrocarbons are liquid or solid at room temperature while small ones (e.g. methane) are gases.
When a simple molecular substance melts or boils, the weak intermolecular forces are overcome — the strong covalent bonds within molecules remain intact. This explains why simple molecules have low melting points despite having strong internal bonds.
Giant Covalent Structures
Giant covalent structures have millions of atoms connected by covalent bonds in a continuous network. They have very high melting points because many strong covalent bonds must be broken to separate the atoms. Examples include diamond, graphite and silicon dioxide.
In diamond, each carbon atom forms four covalent bonds arranged in a rigid tetrahedral pattern. This gives diamond:
- Extreme hardness — the rigid lattice resists deformation.
- A very high melting point — all four bonds per atom must be broken.
- No electrical conductivity — all four outer electrons are used in bonding, leaving none free to carry charge.
In graphite, each carbon atom bonds to three others in flat hexagonal layers. The fourth outer electron from each carbon becomes delocalised — free to move along the layers, allowing graphite to conduct electricity. The layers are held together only by weak intermolecular forces, so they slide easily over one another, making graphite soft and slippery — useful as a lubricant and in pencil leads.
Graphene, Fullerenes and Carbon Nanotubes
Graphene is a single layer of graphite: a one-atom-thick sheet of hexagonally arranged carbon atoms. Scientists Andre Geim and Konstantin Novoselov isolated it at the University of Manchester in 2004, work for which they received the Nobel Prize in Physics in 2010. Graphene is extraordinarily strong relative to its mass and is an excellent conductor of electricity and heat.
Fullerenes are carbon molecules shaped as hollow spheres or tubes. Buckminsterfullerene (\(C_{60}\)) — 60 carbon atoms arranged like a football — was first made in 1985 by Harold Kroto, Robert Curl and Richard Smalley, who were awarded the Nobel Prize in Chemistry in 1996.
Carbon nanotubes are cylindrical fullerenes with remarkable tensile strength and electrical conductivity. They are used in electronics and as reinforcing agents in composite materials.
Diamond: 4 bonds per carbon, rigid, hard, no conductivity. Graphite: 3 bonds per carbon + 1 delocalised electron, layers slide, conducts electricity. Graphene: single graphite layer — isolated 2004, strongest known material per unit mass.
Polymers are very large molecules made from many repeating units called monomers, joined by covalent bonds. Examples include poly(ethene) and poly(propene). The long tangled chains experience stronger intermolecular forces than simple molecules, so polymers are solid at room temperature. They do not conduct electricity and generally have moderate melting points depending on chain length and degree of cross-linking.
When comparing allotropes in the exam, always link structure to property: graphite conducts because delocalised electrons move along the layers, whereas diamond does not because all four outer electrons are covalently bonded. Fullerenes and graphene are studied for nanotechnology applications — fullerene cages can carry drug molecules, while graphene's exceptional strength comes from the hexagonal arrangement of carbon atoms in a single layer.