Atoms, Elements and Compounds – Giant Covalent Structures
Learning Objective
Describe the giant covalent (network) structures of diamond and graphite, explain how their bonding and atomic arrangement give rise to their physical properties, and give examples of their common uses (Cambridge IGCSE 0620).
What is a Giant Covalent Structure?
- A solid in which the atoms are linked together by covalent bonds throughout an extended three‑dimensional lattice.
- There are no discrete molecules; the covalent network extends indefinitely in one, two or three dimensions.
- Breaking the solid into separate atoms requires breaking a very large number of strong covalent bonds, giving such materials high melting points and characteristic physical properties.
1. Diamond – a three‑dimensional giant covalent solid
- Hybridisation: each carbon atom is sp³ hybridised.
- Bonding pattern: the four sp³ orbitals form four strong σ‑bonds with four neighbouring carbon atoms.
- Structure: a continuous three‑dimensional tetrahedral network; every carbon is equivalent and the lattice repeats indefinitely.
- Bond length & angles: C–C bond length ≈ 1.54 pm; tetrahedral angle ≈ 109.5°.
Why the structure matters (structure → properties)
- Extreme hardness (Mohs 10): a very large number of strong σ‑bonds must be broken or deformed.
- Very high melting / sublimation point (≈ 3550 °C): simultaneous rupture of many σ‑bonds.
- Electrical insulator: no free or delocalised electrons are present in the lattice.
- Transparency & high refractive index (≈ 2.42): a regular, defect‑free lattice allows light to pass with little scattering.
Common uses (as listed in the syllabus)
- Jewellery – valued for brilliance and hardness.
- Industrial abrasives and cutting tools – exploit the hardness and high melting point.
- Extension (enrichment): heat‑conducting elements in electronics (high thermal conductivity while electrically insulating).
2. Graphite – a two‑dimensional giant covalent solid
- Hybridisation: each carbon atom is sp² hybridised.
- Bonding pattern: three sp² orbitals form three σ‑bonds with three neighbouring carbons, giving a planar hexagonal (honey‑comb) sheet.
- π‑electron system: the remaining un‑hybridised p‑orbital contains one electron that is delocalised over the whole sheet, creating a continuous π‑bonding network.
- Layered structure: the σ‑bonded sheets (graphene layers) are stacked; weak van der Waals forces hold the layers together, allowing them to slide over one another.
- Bond length & angles: C–C bond length ≈ 1.42 pm; bond angle ≈ 120° within the sheet.
Why the structure matters (structure → properties)
- Soft, slippery feel: weak inter‑layer forces permit easy shear of the layers (lubricating property).
- Good electrical conductor (parallel to layers): delocalised π‑electrons move freely within each sheet.
- Very high melting point (≈ 3600 °C): breaking the strong σ‑bonds within a sheet requires a great amount of energy.
- Opaque, black, greasy texture: result of the layered structure and the presence of mobile electrons.
Common uses (as listed in the syllabus)
- Pencil “lead” – sheets detach and leave a mark on paper.
- Lubricants and dry‑powder lubricating agents – layers slide with minimal friction.
- Electrodes in batteries, fuel cells and electric arc furnaces – conductivity within the planes.
- Refractory material for high‑temperature crucibles and furnace linings.
3. Comparison of Diamond and Graphite
| Feature |
Diamond |
Graphite |
| Hybridisation of C atoms |
sp³ |
sp² + delocalised p‑electron |
| Bonding pattern |
Four σ‑bonds in a 3‑D tetrahedral network |
Three σ‑bonds in 2‑D sheets + π‑bond network |
| Structure |
Continuous three‑dimensional network (giant covalent solid) |
Stacked planar sheets held together by weak van der Waals forces (giant covalent solid) |
| Bond length |
≈ 1.54 pm |
≈ 1.42 pm |
| Hardness |
Very hard (Mohs 10) |
Soft, slippery |
| Electrical conductivity |
Insulator (no free electrons) |
Conductor parallel to layers (delocalised π‑electrons) |
| Melting / sublimation point |
Very high – ≈ 3550 °C (sublimes) |
Very high – ≈ 3600 °C (decomposes) |
| Typical uses (syllabus list) |
Jewellery; abrasives & cutting tools |
Pencil “lead”; lubricants; electrodes; refractory linings |
4. Extended Content – Silicon(IV) oxide (SiO₂)
Relevance: In the extended IGCSE syllabus, students are also required to describe the giant covalent structure of silicon(IV) oxide.
- Hybridisation: silicon is sp³ hybridised; each oxygen is sp³ hybridised.
- Bonding pattern: each Si atom forms four σ‑bonds to four O atoms; each O atom forms two σ‑bonds to two Si atoms, giving a three‑dimensional tetrahedral network (SiO₄ tetrahedra share corners).
- Structure: continuous 3‑D network; no discrete SiO₂ molecules.
- Properties arising from the structure: very high melting point (~ 1710 °C), hardness, electrical insulating behaviour, and transparency (used for glass).
- Common uses: glass manufacturing, ceramics, refractory linings, and as a component of concrete.
5. Suggested Diagrams (to be drawn by the learner)
- Diamond – a three‑dimensional tetrahedral network showing each carbon atom bonded to four neighbours.
- Graphite – a single graphene sheet (hexagonal pattern) with arrows indicating the delocalised π‑electrons, plus a side view showing stacked sheets held together by weak van der Waals forces.
- Silicon(IV) oxide (extension) – a 3‑D network of SiO₄ tetrahedra sharing corners.