Bonding & Structure In Forms Of Carbon (AQA GCSE Combined Science: Synergy: Physical Sciences): Revision Note

Diamond

• Diamond and graphite are allotropes of carbon

• Both substances contain only carbon atoms but due to the differences in bonding arrangements they are physically completely different

• In diamond, each carbon atom bonds with four other carbons, forming a tetrahedron

• All the covalent bonds are identical and very strong

  • All four outer electrons of each carbon atom are held in covalent bonds, leaving no free electrons

Properties of diamond

• Diamond does not conduct electricity

  • All the outer shell electrons in carbon are held in the four covalent bonds around each carbon atom

  • As a result, there are no freely moving particles to carry a charge

• Diamond has a very high melting point

  • Diamond has a giant covalent structure

  • There are strong covalent bonds between the carbon atoms

  • These need lots of energy to break 

• It is extremely hard and dense

  • It has strong covalent bonds and each carbon atom is bonded to four other carbon atoms 

  • Diamond's hardness makes it very useful in cutting tools like drills 

Graphite

• Each carbon atom in graphite forms three covalent bonds to other carbon atoms forming layers of hexagons, leaving one free electron per carbon atom

• These free electrons migrate along the layers and are free to move and carry charge, hence graphite can conduct electricity

• Freely moving electrons are called delocalised electrons - that is, the opposite of localised - they are not confined to any particular place

• The covalent bonds within the layers are very strong

• The layers are attracted to each other by weak intermolecular forces, so the layers can slide over each other making graphite soft and slippery

Properties of graphite

• Graphite conducts electricity 

  • Each carbon atom is bonded to three others leaving one free electron per carbon atom 

  • These free (delocalised) electrons exist in between the layers

  • They are free to move through the structure and carry charge

• Graphite has a high melting point

  • Graphite has a giant covalent structure

  • There are strong covalent bonds between the carbon atoms

  • These need lots of energy to break 

• Graphite is slippery 

  • Graphite is arranged in layers

  • Although the atoms within the layers are joined by strong covalent bonds, the layers have only weak intermolecular forces between them

  • As a result the layers can slide over each other

  • This property allows graphite to be used in pencils and as an industrial lubricant

• Graphite can be used to make inert electrodes for electrolysis, which is particularly important in the extraction of metals such as aluminium

Graphene

• The structure of graphene consists of a single layer of graphite which is a sheet of carbon atoms covalently bonded forming a continuous hexagonal layer

• It is essentially a 2D molecule since it is only one atom thick

• It has very unusual properties that make it useful in:

  • Composite materials

  • Electronics

Properties of graphene

• Graphene has the following properties:

  • It is extremely strong but also amazingly light

  • It conducts heat and electricity

  • It is flexible

  • It is transparent

Strength

• It would take an elephant with excellent balance to break through a sheet of graphene

  • It is very strong due to its unbroken pattern and the strong covalent bonds between the carbon atoms. Even when patches of graphene are stitched together, it remains the strongest material out there

Conductivity

• It has delocalised electrons which can move along its surface allowing it to conduct electricity

  • It is known to move electrons 200 times faster than silicon

  • It is also an excellent conductor of heat

Flexibility

• Those strong bonds between graphene’s carbon atoms are also very flexible. They can be twisted, pulled and curved to a certain extent without breaking, which means graphene is bendable and stretchable

Transparent

• Graphene absorbs 2.3 percent of the visible light that hits it, which means you can see through it without having to deal with any glare

  • This gives it the potential to be used for making computer screens of the future

Fullerenes

• Fullerenes are a group of carbon allotropes which consist of molecules that form hollow tubes or spheres

• The molecules are made of interlocking hexagonal rings, but they can also be rings of five or seven carbons atoms

• Fullerenes can be used to trap other molecules by forming around the target molecule and capturing it, making them useful for targeted drug delivery systems

• They also have a huge surface area and are useful for trapping catalyst molecules onto their surfaces making them easily accessible to reactants so catalysis can take place

• Some fullerenes are excellent lubricants and are starting to be used in many industrial processes

• The first fullerene to be discovered was Buckminsterfullerene which is affectionately referred to as a “Buckyball”

• In this fullerene, 60 carbon atoms are joined together forming 20 hexagons and 12 pentagons which produce a hollow sphere that is the exact shape of a football

Buckminsterfullerene

Carbon nanotubes

• Graphene can also be rolled into a cylinder to produce an interesting type of fullerene called a nanotube

• These have high tensile strength and are resistant to breaking or stretching

• As in graphene, nanotubes can also conduct electricity which makes them useful in composites and specialised materials, electronics and nanotechnology

Carbon nanotubes