Diamond & Graphite
Section: 2. Atoms, Elements & Compounds | Syllabus: Cambridge IGCSE Chemistry (0620)
Giant Covalent Structures
Diamond and graphite are both allotropes of carbon - different structural forms of the same element in the same physical state. Both are giant covalent structures: a continuous network of atoms joined by strong covalent bonds throughout, rather than separate molecules. Despite being made of identical carbon atoms, their very different bonding arrangements give them almost opposite properties.
Diamond's 3D network of 4 bonds per atom versus graphite's stacked layers of 3 bonds per atom explains why the two forms of carbon behave so differently
Diamond
- Each carbon atom forms 4 strong covalent bonds to 4 other carbon atoms in a rigid, tetrahedral 3D network
- All 4 outer electrons of every carbon atom are used in bonding, so there are no free electrons
| Property | Diamond | Explanation |
|---|---|---|
| Melting point | Very high (3550°C) | Huge numbers of strong covalent bonds must be broken throughout the structure |
| Hardness | Extremely hard | Rigid 3D network of strong bonds in every direction |
| Electrical conductivity | Does not conduct | No free electrons - all 4 outer electrons per atom are used in bonding |
Use: cutting tools and drill bits, because of its extreme hardness.
Graphite
- Each carbon atom forms only 3 covalent bonds, to 3 neighbours within a flat hexagonal layer
- Strong covalent bonds hold atoms together within each layer, but only weak forces hold separate layers together, allowing layers to slide over each other
- Each carbon atom's 4th outer electron is not used in bonding - it becomes delocalized, free to move along the layers
| Property | Graphite | Explanation |
|---|---|---|
| Melting point | Very high (3652°C) | Strong covalent bonds within each layer must still be broken |
| Hardness | Soft, slippery | Weak forces between layers let them slide over each other easily |
| Electrical conductivity | Does conduct | Delocalized electrons are free to move along the layers and carry charge |
Uses: as a lubricant and in pencils (layers slide/rub off easily), and as an electrode (it conducts electricity and has a very high melting point).
Worked Example: Uses Linked to Structure
- Question: explain why graphite, but not diamond, can be used to make an electrode. [2]
- Answer: in graphite, each carbon atom uses only 3 of its 4 outer electrons in bonding, leaving one delocalized electron per atom free to move and carry charge, so graphite conducts electricity [1]; in diamond, all 4 outer electrons per atom are used in covalent bonding, so there are no free electrons and diamond does not conduct [1]
Common Mistakes
- ❌ Thinking graphite conducts electricity because of metallic bonding
- ✓ graphite is a non-metal held together entirely by covalent bonds - it conducts because one delocalized electron per carbon atom is free to move, which is a different reason from the "sea of electrons" in metallic bonding
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