Graphite
Summary: Giant covalent allotrope of carbon with hexagonal layers — soft, slippery, and an electrical conductor thanks to delocalised electrons; used as a lubricant and as electrodes. Tags: igcse chemistry Created: 2026-07-18
Graphite is a giant covalent structure of the element carbon, and one of carbon’s allotropes alongside Diamond. Each carbon atom is covalently bonded to only three other carbon atoms, forming flat, parallel layers of hexagonal rings that extend throughout the whole crystal — a giant structure rather than separate molecules. Because carbon has four outer electrons but uses only three in covalent bonds, the fourth electron from every atom becomes a Delocalised Electron, free to move along the layers, which makes graphite a rare non-metal conductor of electricity. The layers themselves are held together only by weak forces of attraction (not covalent bonds), so they slide over one another easily, making graphite soft and slippery. This distinctive combination of properties explains its everyday uses as a lubricant, as pencil “lead”, and as inert electrodes in Electrolysis.
Structure
- A giant covalent structure: strong covalent bonding extends throughout each layer — there are no individual molecules.
- Each carbon atom is bonded to 3 other carbon atoms by strong covalent bonds.
- The atoms are arranged in flat layers of hexagonal rings (bond angle 120°).
- The fourth outer electron of each carbon atom is delocalised between the layers and is free to move.
- Between the layers there are only weak forces of attraction — no covalent bonds — so the layers can slide over each other.
Contrast with Diamond: same element, but in diamond each carbon bonds to 4 others in a rigid tetrahedral network, so diamond is extremely hard and does not conduct.
Structure–Property–Use Table
| Property | Explanation from structure | Related use |
|---|---|---|
| Soft and slippery | Layers are held together only by weak forces, so they slide over each other easily. | Lubricant for machinery; pencil “lead” (layers rub off onto paper). |
| Conducts electricity | Each carbon atom contributes one delocalised electron that moves freely along the layers, carrying charge. | Electrodes in Electrolysis; brushes in electric motors. |
| Very high melting point (~3600 °C) | The covalent bonds within each layer are very strong and require enormous energy to break. | Electrodes can operate at the high temperatures of molten electrolytes without melting. |
| Insoluble in water | Strong covalent bonds within layers cannot be disrupted by water. | Suitable as an inert, unreactive electrode material. |
Why the Uses Follow from the Structure
Lubricant — a good lubricant must let surfaces move past each other with little friction. Graphite’s layers slide because only weak forces (not covalent bonds) act between them, so graphite powder reduces friction between moving metal parts without the mess of oil.
Electrode — an electrode must (1) conduct electricity and (2) survive harsh conditions. Graphite conducts because of its delocalised electrons, and its very high melting point (strong covalent bonds within layers) means it does not melt even in molten-salt Electrolysis such as the extraction of aluminium. It is also fairly unreactive, so it does not usually interfere with the products.
Note the key exam point: the softness and the high melting point are caused by different bonds — weak forces between layers (softness) versus strong covalent bonds within layers (high melting point).
Comparison with Diamond
| Feature | Graphite | Diamond |
|---|---|---|
| Bonds per carbon atom | 3 | 4 |
| Arrangement | Flat hexagonal layers | Rigid 3D tetrahedral network |
| Hardness | Soft, slippery | Hardest natural substance |
| Electrical conductivity | Conducts (delocalised electrons) | Does not conduct (no free electrons) |
| Melting point | Very high | Very high |
| Uses | Lubricant, electrodes, pencils | Cutting tools, drill bits |
Both are allotropes of carbon — see Giant Structures for the full comparison, including Silicon(IV) Oxide.
Graphite as a Lubricant
Graphite is used as a solid lubricant because its layers can slide over one another easily. In graphite, each carbon atom is bonded to only 3 other carbon atoms in flat hexagonal layers. While the covalent bonds within each layer are strong, the forces between the layers are only weak intermolecular forces. These weak interlayer forces allow the layers to slide past each other when a small force is applied, making graphite slippery and an effective dry lubricant. This is particularly useful in high-temperature environments where oil-based lubricants would decompose.
Exam point: Graphite is a lubricant because of weak forces between layers → layers slide over each other.
Sources
- Cambridge IGCSE Chemistry 0620 Syllabus (2023-2025) — Section 2.6 Giant covalent structures, Core points 1–2, Cambridge Assessment International Education
Common Misconceptions
| Misconception | Reality |
|---|---|
| ”Graphite is soft, so it must have a low melting point” | Softness comes from weak forces between layers; the melting point is very high because the covalent bonds within layers are strong and must be broken to melt it. |
| ”The layers in graphite are held together by covalent bonds” | Only weak forces of attraction act between layers — that is exactly why they slide and graphite is slippery. Covalent bonds exist only within each layer. |
| ”Graphite conducts because it is a metal” | Graphite is a non-metal. It conducts because each carbon atom donates one delocalised electron that can move along the layers — not because of Metallic Bonding. |
| ”Graphite and diamond are different elements” | Both are carbon — allotropes of the same element. Their different properties come entirely from different bonding arrangements. |
| ”All four of carbon’s outer electrons are bonded in graphite” | Only three are used in covalent bonds; the fourth is delocalised, which is the source of graphite’s conductivity. |