Why are diamond and graphite so different even though both are made only of carbon?
Describe giant covalent structures including diamond and graphite, and relate their bonding and structure to their very different properties.
A focused answer to WJEC GCSE Chemistry topic 2.1, covering giant covalent structures, the bonding and structure of diamond and graphite, and how these explain their hardness, melting points and electrical conductivity.
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What this topic is asking
WJEC topic 2.1 wants you to describe giant covalent structures and use diamond and graphite to show how the same element (carbon) can have very different properties because of different structures. You must relate the bonding and structure to properties such as hardness, melting point and conductivity.
Giant covalent structures
Because the strong covalent bonds run through the entire structure, giant covalent substances have very high melting and boiling points: a huge amount of energy is needed to break the many covalent bonds. The two most important examples are both made only of carbon: diamond and graphite.
Diamond
In diamond, every carbon atom forms four covalent bonds to four other carbon atoms, arranged in a rigid 3D tetrahedral lattice.
Its hardness makes diamond useful for cutting tools and drill tips.
Graphite
In graphite, every carbon atom forms only three covalent bonds, producing flat layers of carbon atoms arranged in hexagons. The layers are held together only by weak intermolecular forces.
Its slipperiness makes graphite useful as a lubricant and in pencils, and its conductivity makes it useful for electrodes.
Why structure explains the difference
Diamond and graphite are both carbon, so the difference in their properties comes entirely from how the atoms are bonded and arranged. Four bonds per carbon give the rigid, non-conducting diamond; three bonds per carbon give the layered, conducting graphite. This is a clear example of how structure determines properties.
Another giant covalent substance you should recognise is silicon dioxide (), found in sand and quartz. Like diamond it has a rigid 3D network of strong covalent bonds, so it is hard, has a very high melting point and does not conduct electricity. It is used to make glass and to line furnaces, because it can withstand high temperatures without melting. Whenever you meet a substance described as very hard with a very high melting point and no conductivity, a giant covalent structure is the most likely answer.
Exam-style practice questions
Practice questions written in the style of WJEC exam questions on this dot point, with worked answer explainers. The year tag is the paper they imitate, not the source.
WJEC sample4 marksDiamond and graphite are both forms of carbon. Explain why diamond is very hard but graphite is soft and slippery, in terms of their structures.Show worked answer →
A Unit 2.1 compare-and-explain question. Reward: in diamond each carbon atom forms four strong covalent bonds in a rigid 3D giant lattice, so it is very hard. In graphite each carbon forms only three covalent bonds, making layers of carbon atoms; there are only weak forces between the layers, so the layers can slide over each other, making graphite soft and slippery. Markers credit four bonds and a rigid 3D network for diamond, and layers with weak forces between them that allow sliding for graphite. A common error is to mention intermolecular forces without saying the layers slide.
WJEC sample3 marksExplain why graphite conducts electricity but diamond does not.Show worked answer →
A Unit 2.1 explanation question. Reward: in graphite each carbon forms only three covalent bonds, so each carbon has one delocalised (free) electron; these electrons can move through the layers and carry charge, so graphite conducts. In diamond each carbon forms four bonds, using all its outer electrons, so there are no free electrons to carry charge and diamond does not conduct. Markers credit graphite having delocalised electrons free to move, and diamond having all electrons used in bonding with none free. A common slip is to say diamond conducts because it is a giant structure.
Related dot points
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- Use the properties of a substance (melting point, conductivity, state) to deduce its type of bonding and structure.
A focused answer to WJEC GCSE Chemistry topic 2.1, bringing ionic, simple molecular, giant covalent and metallic structures together and showing how to deduce the bonding in a substance from its melting point, electrical conductivity and state.
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Sources & how we know this
- WJEC GCSE Chemistry specification (from 2016) — WJEC (2016)