Structures+-+Elise

Allotropes of carbon

**Diamond** **The structure of diamond**

Diamond is a form of pure carbon. Each carbon shares electrons with four other carbon atoms - forming single bonds This is a giant covalent structure. It is not a molecule, because the number of atoms joined up in a real diamond is variable, depending on the size of the crystal.



**The physical properties of diamond** Diamond has very high melting point. Very strong carbon-carbon covalent bonds have to be broken throughout the structure before melting occurs.  Diamond is very hard. This is because of the very strong carbon-carbon covalent bonds which extend throughout the whole crystal in three dimensions.  Diamonds doesn’t conduct electricity. All the electrons are held tightly between the atoms, and aren’t free to move around.  Diamond doesn’t dissolve in water or any other solvent. There are no possible attractions which could occur between solvent molecules and carbon atoms which could outweigh the attractions between the covalently bound carbon atoms.

**The structure of graphite** 

**The bonding in graphite**  The atoms within a layer are held together by strong covalent bonds.  Each carbon atom uses three of its electrons to form simple bonds to its three close neighbours. That leaves a fourth electron in the bonding level. These spare electrons in each carbon atom become delocalised over the whole of the atoms in one layer.  The delocalised electrons are free to move anywhere within the layer. However, there is no direct contact between the delocalised electrons in one layer and those in the neighbouring layers.  As the delocalised electrons move around in the layer, very large temporary dipoles can be set up which will induce opposite dipoles in the layers above and below - and so on throughout the whole graphite crystal.

**The physical properties of graphite** <span style="font: normal normal normal 10px/normal Helvetica; letter-spacing: 0px; margin-bottom: 0px; margin-left: 0px; margin-right: 0px; margin-top: 0px;"> Graphite is a soft material with a slimy feel and is used in pencils. Although the forces holding the atoms together in each later are very strong, the attractions between the layers are much weaker. <span style="font: normal normal normal 10px/normal Helvetica; letter-spacing: 0px; margin-bottom: 0px; margin-left: 0px; margin-right: 0px; margin-top: 0px;"> Graphite is insoluble in any solvents. Attractions between solvent molecules and carbon atoms will never be strong enough to overcome the strong covalent bonds in graphite. <span style="font: normal normal normal 10px/normal Helvetica; letter-spacing: 0px; margin-bottom: 0px; margin-left: 0px; margin-right: 0px; margin-top: 0px;"> Graphite is less dense than diamond, because the layers in graphite are relatively far apart. A graphite crystal contains a lot of wasted space, which isn’t there in a diamond crystal. <span style="font: normal normal normal 10px/normal Helvetica; letter-spacing: 0px; margin-bottom: 0px; margin-left: 0px; margin-right: 0px; margin-top: 0px;"> Graphite conducts electricity. Each carbon atom uses three of its electrons to form these simple covalent bonds. The fourth electron in the outer layer of each atom is free to move around throughout the whole of the layer. The movement of these electrons allows the graphite to conduct electricity.

<span style="font: normal normal normal 10px/normal Helvetica; letter-spacing: 0px; margin-bottom: 0px; margin-left: 0px; margin-right: 0px; margin-top: 0px;">**C60 fullerene** <span style="font: normal normal normal 10px/normal Helvetica; letter-spacing: 0px; margin-bottom: 0px; margin-left: 0px; margin-right: 0px; margin-top: 0px;"> **The structure of fullerene** <span style="font: normal normal normal 10px/normal Helvetica; letter-spacing: 0px; margin-bottom: 0px; margin-left: 0px; margin-right: 0px; margin-top: 0px;"> Fullerenes consist of 20 hexagonal and 12 pentagonal rings as the basis of an icosohedral symmetry closed cage structure, with a carbon atom at the vertices of each polygon and a bond along each polygon edge. In fullerenes each sp² hybridised carbon is bonded by sigma bonds to three other carbons, but because the surface of the sphere is not planar, there is only a little delocalisation of the unpaired bonding electrons.

Electrical conductivity is greater than that of diamond, but much less than that of graphite. Electrons cannot flow easily from one C₆₀ molecule to the next. <span style="font: normal normal normal 10px/normal Helvetica; letter-spacing: 0px; margin-bottom: 0px; margin-left: 0px; margin-right: 0px; margin-top: 0px;"> C₆₀ behaves as an electron deficient molecule readily accepting electrons from reducing agents to form anions with a variety of charges. <span style="font: normal normal normal 10px/normal Helvetica; letter-spacing: 0px; margin-bottom: 0px; margin-left: 0px; margin-right: 0px; margin-top: 0px;"> Fullerenes are molecular, hence they will dissolve in non-polar solvents and have low melting points.