Graphite — AQA GCSE Combined Science
Test yourself on Graphite with AQA GCSE practice questions.
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Graphite explained
Graphite is a giant covalent structure of carbon, but each carbon atom bonds covalently to only three others, producing flat layers of hexagonal rings.
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The fourth outer electron per atom is delocalised and can move along the layer, so graphite conducts electricity. Within a layer, the strong covalent bonds need a great deal of energy to break, giving graphite a very high melting point. Between layers there are no covalent bonds, only weak forces, so layers can slide over one another; this makes graphite soft and slippery, suitable for pencil leads and as a lubricant. A useful method is to separate the two length scales: strong covalent bonding within each layer, weak forces between layers.
In graphite, one electron from each carbon atom is delocalised.
Carbon has four outer-shell electrons. In graphite, each carbon atom forms three covalent bonds to three other carbon atoms in a hexagonal layer, using three of those electrons. The fourth electron from each carbon atom is not used in bonding and becomes delocalised: it is free to move throughout the layer rather than staying with one atom. Because every carbon contributes one delocalised electron, the layer contains a sea of moving electrons. This explains why graphite conducts electricity along its layers, unlike diamond where all four outer electrons are localised in covalent bonds. A useful check is to count bonds: three sigma bonds per carbon means one electron remains, so the ratio of delocalised electrons to carbon atoms is 1:1.
Students should be able to explain the properties of graphite in terms of its structure and bonding.
Graphite is a giant covalent structure of carbon atoms arranged in hexagonal layers. Within each layer, every carbon atom forms three strong covalent bonds, so three outer electrons are localised in bonds and the fourth is delocalised. The many strong covalent bonds make graphite very high melting, while the delocalised electrons allow it to conduct electricity along the layers. Between layers there are only weak forces, so layers can slide over one another, making graphite soft and slippery and useful as a lubricant and in pencils. The weak forces also mean the layers are far apart, giving graphite a lower density than diamond. To explain a property, always name the structural feature responsible and then link it to the property.
Students should know that graphite is similar to metals in that it has delocalised electrons.
Graphite is a giant covalent structure of carbon atoms arranged in hexagonal layers. Each carbon atom forms three covalent bonds, leaving one electron per atom delocalised within the layers. These delocalised electrons are free to move through the structure, so graphite conducts electricity and heat like a metal. Unlike metals, graphite is not malleable in the same way, because its layers are held by weak forces and can slide, and it has no sea of positive ions. A useful method is to compare graphite with a metal such as copper: both conduct because charged particles can move, but in copper the moving charges are electrons in a metallic lattice, while in graphite they are delocalised electrons within covalent layers. This explains why graphite is used for electrodes and feels slippery.
Your focus
- Describe the structure of graphite as layers of hexagonal rings in which each carbon atom forms three covalent bonds.
- Explain why graphite has a high melting point but is soft and slippery, referring to bonding within and forces between layers.
- Explain why graphite conducts electricity in terms of delocalised electrons.
Show all 12 objectives
- Describe the bonding in a graphite layer, including the number of covalent bonds formed by each carbon atom.
- Identify that one electron per carbon atom is delocalised in graphite.
- Relate delocalised electrons in graphite to its ability to conduct electricity.
- Describe the structure of graphite, including its hexagonal layers and the bonding within and between layers.
- Explain how the strength of covalent bonds and the weakness of forces between layers account for the melting point, softness and density of graphite.
- Explain how delocalised electrons in graphite allow it to conduct electricity.
- Describe the bonding in graphite, including that each carbon atom forms three covalent bonds.
- Explain how delocalised electrons allow graphite to conduct electricity.
- Compare the conduction mechanism in graphite with that in metals.
Graphite exam tips
Marking Points
- Each carbon atom in graphite forms three covalent bonds with three other carbon atoms, leaving one outer electron delocalised.
- The bonded atoms form layers of hexagonal rings, and each layer is a giant covalent structure.
- There are no covalent bonds between the layers; only weak forces act between them, so the layers can slide over one another.
- The strong covalent bonds within a layer require much energy to break, so graphite has a very high melting point.
- The delocalised electrons can move along the layers, so graphite conducts electricity.
- The weak forces between layers explain why graphite is soft and slippery, unlike diamond.
- Each carbon atom in graphite forms three covalent bonds to neighbouring carbon atoms in the same layer.
- Three of the four outer-shell electrons per carbon atom are used in these covalent bonds.
- The remaining fourth outer-shell electron per carbon atom becomes delocalised.
- Delocalised electrons are free to move within the layer, so graphite conducts electricity.
- The ratio of delocalised electrons to carbon atoms is 1:1, because every carbon atom contributes one delocalised electron.
- Delocalisation occurs within layers; movement between layers is limited, which helps explain why conduction is mainly along the layers.
- Graphite has a giant covalent structure in which each carbon atom is covalently bonded to three others in a hexagonal layer.
- Strong covalent bonds within layers require a large amount of energy to break, giving graphite a very high melting point.
- One delocalised electron per carbon atom is free to move within a layer, so graphite conducts electricity.
- Weak forces between layers allow layers to slide over one another, making graphite soft and slippery.
- The weak forces between layers also mean the layers are separated by relatively large distances, so graphite has a lower density than diamond.
- Graphite is insoluble in water because the covalent bonds within layers are too strong to be broken by water molecules.
- Graphite is a giant covalent structure in which each carbon atom is covalently bonded to three other carbon atoms.
- Each carbon atom contributes one electron that becomes delocalised, so delocalised electrons are present within the layers.
- The delocalised electrons are free to move, allowing graphite to conduct electricity and thermal energy.
- This is similar to metals, where delocalised electrons also carry charge through the structure.
- Graphite differs from metals because it has a layered covalent lattice rather than a lattice of positive ions.
- The weak forces between layers allow them to slide, which explains why graphite is slippery and useful as a lubricant.
Examiner Tips
- 💡Always distinguish within-layer covalent bonding from between-layer forces, since this contrast explains both the high melting point and the softness.
- 💡When explaining conduction, name delocalised electrons explicitly and say they can move along the layers.
- 💡Use a labelled diagram of hexagonal layers if it helps, but keep the written explanation of bonding and forces clear.
- 💡Link the number of bonds per carbon atom directly to the number of delocalised electrons in your answer.
- 💡Use the phrase 'delocalised electrons' rather than 'free electrons' when explaining conduction in graphite.
- 💡If asked to compare graphite and diamond, state the bonding count for each before drawing the conductivity conclusion.
- 💡For each property, name the structural feature first and then state the consequence, for example 'weak forces between layers, so layers slide'.
- 💡Use comparative language when contrasting graphite with diamond, such as 'graphite has three bonds per carbon atom whereas diamond has four'.
- 💡Check that every property you mention is linked to a specific structural feature rather than given as a standalone fact.
- 💡Link the structure of graphite directly to its property of conduction in one sentence, for example: three bonds per carbon atom leave delocalised electrons that can move.
- 💡When comparing graphite with metals, name the similarity (delocalised electrons) and one difference (layered covalent structure versus positive ion lattice).
- 💡Use the phrase delocalised electrons rather than free electrons if the question asks for the reason graphite conducts, and avoid saying electrons are shared between layers.
Common Mistakes
- Saying graphite is a simple molecular substance: correct this by describing it as a giant covalent structure made of layers of hexagonal rings.
- Stating that graphite melts easily because layers slide: correct this by distinguishing weak forces between layers from the strong covalent bonds within layers, which give a very high melting point.
- Claiming graphite conducts because of covalent bonds: correct this by explaining that delocalised electrons, not the bonds themselves, carry charge.
- Saying all four electrons are delocalised: correct this by counting three covalent bonds per carbon atom, leaving only one electron delocalised.
- Confusing graphite with diamond and stating that all outer electrons are localised in bonds: correct this by noting that diamond has four bonds per carbon and no delocalised electrons.
- Describing delocalised electrons as 'free atoms' or 'ions': correct this by stating they are electrons that are not associated with any single atom and can move through the layer.
- Saying graphite melts easily because layers slide: correct this by distinguishing weak forces between layers from strong covalent bonds within layers, which give a very high melting point.
- Stating that graphite conducts because it contains ions: correct this by explaining that delocalised electrons, not ions, carry charge in graphite.
- Claiming graphite is soft because its covalent bonds are weak: correct this by stating the covalent bonds are strong but the forces between layers are weak.
- Thinking graphite conducts because its ions move: correct this by stating that the moving charges are delocalised electrons, not ions.
- Believing each carbon atom in graphite forms four covalent bonds: correct this by noting that only three bonds form, leaving one delocalised electron per carbon atom.
- Confusing graphite with diamond and saying it does not conduct: correct this by explaining that diamond has no delocalised electrons, whereas graphite does.