Structure and Properties of Graphite Revision Notes | GCSE AQA Higher Chemistry

1. What Makes Graphite Special?

You know diamond is the hardest substance on Earth. But here's the twist - graphite is made of exactly the same atoms as diamond, yet it's soft enough to write with. The secret lies in how those carbon atoms are arranged.

Graphite has a unique structure that gives it properties more like metals than like other non-metals. It conducts electricity, something most carbon compounds can't do. Understanding why requires looking at the atomic level.

2. The Structure of Graphite

Picture a honeycomb pattern - that's what each layer of graphite looks like. Each carbon atom forms exactly three covalent bonds with three neighbouring carbon atoms, creating hexagonal rings.

Here's the crucial part: these hexagonal layers stack on top of each other, but there are no covalent bonds between the layers. The layers are held together only by weak intermolecular forces.

Graphite has a layered giant covalent structure. Each carbon atom bonds to three others in hexagonal sheets, and weak forces between sheets allow layers to slide over each other.

This layered structure explains why graphite feels slippery and why pencils work - the layers can slide over each other easily, leaving marks on paper.

3. The Fourth Electron - Key to Conductivity

Carbon atoms have four electrons in their outer shell. In graphite, each carbon uses three of these electrons to form covalent bonds. But what happens to the fourth electron?

This fourth electron becomes delocalised. Instead of being tied to one specific carbon atom, it can move freely throughout the layer. Think of it like a sea of mobile electrons floating above and below each layer of carbon atoms.

These delocalised electrons are what make graphite conduct electricity. When you apply a voltage, the electrons can flow through the structure, carrying electrical current.

4. Graphite vs Metals - Surprising Similarities

Graphite behaves like metals in one key way - it conducts electricity. This is unusual for a non-metal element. The reason is the same in both cases: delocalised electrons.

In metals, electrons form a 'sea' around positive metal ions. In graphite, electrons form a 'sea' above and below each layer of carbon atoms. Both structures allow electrons to move freely when a voltage is applied.

However, there's an important difference. Metals conduct equally well in all directions. Graphite conducts well along the layers but poorly between layers, because the delocalised electrons can't easily jump from one layer to another.

5. Properties Explained by Structure

Now we can explain graphite's key properties using its structure and bonding. This structure-property relationship is exactly what examiners want you to understand.

Worked example: Explaining Graphite's Electrical Conductivity

Explain why graphite conducts electricity. [3 marks]

1. Each carbon atom forms three covalent bonds — [1 mark] - Must mention the specific number of bonds
2. The fourth electron from each carbon is delocalised — [1 mark] - Key concept that enables conductivity
3. These delocalised electrons can move and carry electrical current — [1 mark] - Must explain how electrons enable current flow

Answer: Graphite conducts electricity because each carbon forms three covalent bonds, leaving the fourth electron delocalised and free to move, carrying electrical current.

The layered structure also explains why graphite is soft and slippery. The weak forces between layers mean they can slide over each other easily. This is completely different from diamond, where every carbon is locked in place by four strong covalent bonds.

6. Exam Success - Getting Full Marks

Graphite questions often ask you to explain properties in terms of structure. The key is linking the atomic-level structure to the observable property through a clear chain of reasoning.

Common mistakes include confusing graphite with diamond (four bonds vs three), forgetting to mention delocalised electrons when explaining conductivity, and not linking structure to properties clearly.

Exam precision note: In Structure and Properties of Graphite - Graphite - Structure and bonding of carbon, always identify the key chemical idea first, then link it to evidence, data, or an equation. A high-scoring GCSE answer names the particles involved, states what changes, and justifies the conclusion using correct scientific language. Where calculations are used, show each step with units and keep significant figures sensible. Finally, check that terminology matches AQA wording so your answer is unambiguous and mark-scheme friendly.

Graphite and diamond are both forms of carbon, but their structures lead to different properties. In graphite, each carbon forms three covalent bonds and leaves one delocalised electron, so graphite conducts electricity. The layers are held together by weak forces and can slide over each other, making graphite soft and slippery. In diamond, each carbon forms four covalent bonds in a giant 3D lattice, so it is very hard and does not conduct.

Graphite is used for electrodes because it conducts electricity and has a very high melting point. Its delocalised electrons can move through the structure, while the strong covalent bonds mean it stays solid even when it gets hot. This is a better GCSE application than discussing how graphite conductivity varies with temperature.