Graphite and diamond are both giant covalent structures made entirely of carbon atoms, yet they have very different physical properties. Explain why graphite is soft and can be used as a lubricant, whereas diamond is the hardest natural substance.
Written & reviewed by James Millett — Biology (Imperial College London), PGCE Science (University of Cambridge).
Model answer (4 marks)
Graphite consists of layers of hexagonal rings of carbon atoms. The layers are held together only by weak van der Waals forces, so they can slide over one another easily, making graphite soft and useful as a lubricant. In diamond each carbon atom is covalently bonded to four other carbons in a tetrahedral arrangement, forming a rigid three‑dimensional network of strong covalent bonds. This gives diamond its extreme hardness.
Examiner tips
- Use the word ‘layers’ and ‘van der Waals forces’ for graphite; mention ‘tetrahedral’ and ‘3‑D network’ for diamond.
- Show the contrast between weak interlayer forces and strong covalent bonds to justify the properties.
Common mistakes
- Confusing graphite’s layers with a 3‑D network; describing diamond as having weak bonds; omitting the term ‘tetrahedral’ or ‘van der Waals’.
Mark scheme (4 marks)
- In graphite, carbon atoms are arranged in layers (of hexagonal rings)
- The layers in graphite are held together by weak (intermolecular/van der Waals) forces, allowing them to slide over one another easily
- In diamond, each carbon atom is covalently bonded to four other carbon atoms (tetrahedral arrangement)
- Diamond has a rigid three-dimensional network of strong covalent bonds throughout the structure, making it very hard
Key terms in this question
giant covalent structure · lubricant
Related
- All Cambridge International IGCSE Chemistry (0620) revision notes →
- How to answer a "Explain" question →
- Decode the mark scheme abbreviations →