Diamond and graphite: structure and uses
Explain why two forms of carbon have different hardness and electrical behaviour.
Section 2.6 · Giant covalent structures
Exam scope and assessment notes
Core and Extended · Chemistry 0971, 2026–2028. Cambridge states that IGCSE 0971 has the same subject content as IGCSE 0620, with 9–1 grading. This StudyForA pathway maps the 2026–2028 syllabus only. This 9–1 graded qualification is available only in a limited number of administrative zones. Confirm that your centre can enter you for this code and exam series with Cambridge International. Cambridge has published a separate syllabus for 2029. StudyForA has not reconciled that version yet; this page supports exams in 2026, 2027 and 2028 only.
Before you begin
- Distinguish a giant structure from separate simple molecules.
- Know that carbon supplies four outer electrons.
Work on paper. Try each question before revealing help, and explain why your method works.
Quick readiness check
What is a giant covalent structure?
What you will learn
- Compare the giant covalent structures of diamond and graphite.
- Explain diamond cutting tools and graphite lubricant and electrode uses.
Diamond is a three-dimensional network
In diamond, each carbon atom forms four covalent bonds to other carbon atoms in a three-dimensional network. There are no small diamond molecules. Strong covalent bonds hold the network together in every direction, explaining its hardness and high melting point. All four outer electrons are involved in bonding, so diamond has no freely mobile electrons to carry current in this model.
A cutting tool needs a material that resists being scratched or deformed by the workpiece. Diamond's rigid bonded network makes it suitable for that role. Hardness is different from electrical conductivity: the bonds explain why diamond is hard, while the absence of mobile charge carriers explains its poor conductivity.
Graphite has strong layers and weaker attractions between them
In graphite, each carbon bonds to three others in flat hexagonal layers. Strong covalent bonds act within each layer, while weaker attractions act between layers. Layers can slide over one another, so graphite can work as a lubricant. The bonds within the sheets do not all break when the sheets slide.
One outer electron from each carbon contributes to delocalised electrons that can move along the layers. Graphite therefore conducts electricity and can be used as an electrode. Both graphite and diamond are giant covalent structures with high melting points, despite their different mechanical and electrical properties.
Pause and explain
Which feature directly explains diamond's hardness?
Worked example
Explain why graphite can be an electrode but diamond is generally an electrical insulator.
Show the worked solution
- Graphite's carbons bond to three neighbours, leaving delocalised electrons.
- These electrons can move along the layers and carry charge.
- Diamond's carbons use all four outer electrons in covalent bonds, leaving no mobile electrons in this model.
Answer Graphite has mobile delocalised electrons; diamond does not.
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Write your answer and reasoning first. Use a hint only if you are stuck. The model solution is for self-checking; your written work is not automatically marked.
Compare the number of bonded carbon neighbours in diamond and graphite, then link the structures to hardness or sliding.
Give me a hint
Count neighbours and identify the direction of strong bonding.
Compare my reasoning
- Diamond has four bonded neighbours per carbon in a three-dimensional network.
- Graphite has three bonded neighbours per carbon within layers.
- The rigid network makes diamond hard; weaker between-layer attractions allow graphite layers to slide.
Diamond: four neighbours and a rigid network. Graphite: three neighbours in sliding layers.
Look for these in your work
- Both neighbour counts are correct.
- A structural feature explains each mechanical property.
Explain two reasons graphite is useful as an electrode in a suitable electrolytic cell.
Give me a hint
An electrode must conduct, and graphite also tolerates high temperatures.
Compare my reasoning
- Delocalised electrons carry charge along graphite's layers.
- Its giant covalent structure has strong bonds within the sheets.
- These bonds contribute to a high melting point, although chemical suitability still depends on the cell.
It conducts via delocalised electrons and has a high melting point from strong covalent bonding.
Look for these in your work
- Conduction uses electrons rather than moving carbon ions.
- Strong sheet bonds are distinguished from weaker interlayer attractions.
A carbon sample conducts electricity and rubs into slippery flakes. Which structure fits, and why is 'all its bonds are weak' an incorrect explanation?
Give me a hint
Use both observations and distinguish within-layer from between-layer bonding.
Compare my reasoning
- Conduction and slipping flakes fit graphite.
- Weak attractions between layers allow sliding.
- Strong covalent bonds within layers still explain its giant network and high melting point.
Graphite fits; only the interlayer attractions are weak, not all its covalent bonds.
Look for these in your work
- Both observations support the identification.
- The explanation distinguishes the two types of interaction.
Common mistakes
- Saying all bonds in graphite are weak.
- Calling diamond a molecule of four carbon atoms.
What can you explain now?
Why can graphite act as a lubricant?
Compare with the explanation
Its layers can slide over one another.
Weak attractions between layers allow sliding, while strong bonds remain within each layer.
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Make it stick
Sketch and annotate diamond and graphite using neighbour counts, mobile electrons and directions of strong bonding. Connect each feature to one use.
If you needed the solution, close it and solve the problem again from a blank page. A correct answer is stronger when you can explain the reason for each step.