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Graphene is a single layer of graphite, with carbon atoms in a hexagonal lattice each bonded to three others. Free notes and 14 questions answered.
Discovering Graphene and Fullerenes is part of Structure and bonding of carbon in GCSE AQA Higher Chemistry.
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GCSE AQA Higher
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Graphene is a single layer of graphite, with carbon atoms in a hexagonal lattice each bonded to three others. Free notes and 14 questions answered.
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Use this lesson to stay inside Structure and bonding of carbon while connecting the detail back to the wider Topic 2: Bonding, structure, and the properties of matter topic.
Read the core explanations from Discovering Graphene and Fullerenes before testing yourself from memory.
In 2004, two scientists at Manchester University used sticky tape to peel layers off graphite until they had just one atom thick. This simple experiment won them the Nobel Prize and launched the 'wonder material' revolution. What makes a single layer of carbon so special?
Carbon is the ultimate shape-shifter. The same element that makes soft graphite pencil lead can also form the hardest natural substance - diamond. Now we've discovered even more exotic forms: graphene and fullerenes. These aren't just laboratory curiosities - they're reshaping technology.
Think of carbon allotropes like different buildings made from identical bricks. Same carbon atoms, completely different structures and properties. Diamond has each carbon bonded to four others in a 3D network - making it incredibly hard. Graphite has carbon atoms in layers that slide over each other - making it slippery and soft.
But what happens when you take just one layer from graphite? You get something extraordinary: graphene. And when you roll carbon into football-shaped cages? You get fullerenes. Let's explore these carbon superstars.
Picture chicken wire made from carbon atoms. That's graphene - a flat sheet just one atom thick. Each carbon atom forms three covalent bonds with its neighbours, creating a hexagonal pattern that stretches in all directions. The fourth electron from each carbon isn't used in bonding - these electrons are delocalised and can move freely across the sheet.
This structure gives graphene incredible properties. It's 200 times stronger than steel, yet flexible like rubber. It conducts electricity better than copper and heat better than diamond. You could stretch a sheet of graphene over a coffee cup, and it would support the weight of a car - if you could see it!
Imagine folding a flat sheet of graphene into a football shape. That's essentially what fullerenes are - hollow carbon cages. The most famous is buckminsterfullerene (C₆₀), which looks exactly like a football with 60 carbon atoms arranged in pentagons and hexagons.
Unlike graphene's flat structure, fullerenes curve by including five-membered rings alongside the six-membered rings. This curvature closes the structure into a cage. Each carbon still forms three covalent bonds, but the geometry creates a hollow space inside - perfect for trapping other molecules.
Fullerenes come in many sizes: C₆₀ (football-shaped), C₇₀ (rugby ball-shaped), and even tube-shaped versions called carbon nanotubes. The hollow interior makes them useful for drug delivery - imagine a molecular cage that can carry medicine directly to cancer cells.
Why is graphene so strong while graphite is soft? Why do fullerenes form cages while diamond forms crystals? The answer lies in how the same carbon atoms bond differently in each structure. Let's decode the structure-property relationships.
Worked example: Explaining Graphene's Electrical Conductivity
Explain why graphene conducts electricity, using its structure and bonding. [4 marks]
Answer: Graphene conducts electricity because it has delocalised electrons that can move freely and carry current.
The key insight is that structure determines properties. Graphene's flat structure with delocalised electrons makes it conductive and strong. Fullerenes' curved cage structure makes them perfect for encapsulation. Diamond's 3D network makes it the hardest natural material.
These aren't just laboratory curiosities. Graphene's unique properties are revolutionising technology. Its combination of strength, conductivity, and transparency makes it perfect for flexible electronics - imagine a phone screen you can fold like paper.
In composites, adding just 1% graphene to plastic can double its strength while making it conductive. Tennis rackets with graphene are lighter yet more powerful. Car bodies with graphene composites are stronger yet use less material.
Fullerenes are finding uses in medicine and materials science. Their cage structure can trap drug molecules for targeted delivery. Some fullerenes show promise as superconductors - materials that conduct electricity with zero resistance.
Graphene and fullerenes questions test your ability to link structure to properties. Examiners love asking you to explain why these materials have specific properties - always trace it back to the bonding and electron arrangement.
Common exam mistakes include confusing graphene with graphite, forgetting that fullerenes are hollow, and not explaining why delocalised electrons enable conductivity. Practice drawing the structures - hexagonal sheets for graphene, cage shapes for fullerenes.
For evaluation questions about applications, consider both advantages (unique properties) and limitations (cost, manufacturing challenges). The key is showing you understand how atomic structure determines macroscopic properties.
In fullerenes, each carbon atom forms three covalent bonds. Delocalised electrons are still present, but the curved cage structure means fullerenes do not conduct electricity as effectively as graphene, where electrons can move across a flat layer.
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Move to the related lessons or back to the topic page when you need the wider sequence before exam-style practice.
Graphene is a single layer of graphite, made of carbon atoms arranged in a hexagonal lattice with each carbon bonded to three others by covalent bonds. It is a flat sheet just one atom thick.
Fullerenes are hollow carbon structures in which carbon atoms are arranged in rings of five or six atoms, forming cage-like molecules. The most famous is buckminsterfullerene, C₆₀, which has 60 carbon atoms arranged like a football.
An allotrope is one of the different structural forms of the same element, where the atoms are arranged in different ways to give different properties. Diamond, graphite, graphene and fullerenes are all allotropes of carbon.
Graphene is used in flexible electronics because of its strength, conductivity and transparency, and in composites, where adding just 1% graphene to plastic can double its strength while making it conductive. Tennis rackets with graphene are lighter yet more powerful.
Graphene conducts electricity because each carbon atom forms 3 covalent bonds, using 3 of its 4 outer electrons and leaving 1 electron per carbon atom. These electrons are delocalised and can move freely across the sheet, carrying electrical current.
Fullerenes still have delocalised electrons because each carbon atom forms three covalent bonds, but they do not conduct electricity as effectively as graphene. The curved cage structure is the reason, whereas in graphene the electrons can move across a flat layer.
Graphene is a single layer of graphite, a flat sheet of carbon just one atom thick. Graphite is made of many such layers that slide over each other, making it soft, while a single graphene sheet is very strong and conducts electricity across the sheet.
Graphene is added to composites because adding just 1% graphene to plastic can double its strength while making it conductive. Tennis rackets with graphene are lighter yet more powerful, and car bodies with graphene composites are stronger yet use less material.
Graphene was discovered in 2004 by two scientists at Manchester University, who used sticky tape to peel layers off graphite until they had a layer just one atom thick. This simple experiment won them the Nobel Prize and launched the 'wonder material' revolution.
Graphene is 200 times stronger than steel, yet flexible like rubber. A sheet of graphene stretched over a coffee cup could support the weight of a car, even though it is a flat sheet of carbon atoms just one atom thick.
Fullerenes come in many sizes and shapes: C₆₀ is football-shaped, C₇₀ is rugby ball-shaped, and tube-shaped versions are called carbon nanotubes. All are hollow carbon structures with atoms arranged in rings of five or six atoms.
Graphene is used in flexible electronics because it combines strength, conductivity and transparency in a sheet that is flexible like rubber. It could be used to make a phone screen that folds like paper.
Graphene conducts heat better than diamond and conducts electricity better than copper. These properties come from its structure: a flat hexagonal sheet of carbon atoms, each bonded to three others, with delocalised electrons that can move freely across the sheet.
Each carbon atom in graphene forms three covalent bonds with its neighbours, creating a hexagonal pattern. This uses 3 of carbon's 4 outer electrons, so the fourth electron from each carbon is delocalised and can move freely across the sheet.