The Structures and Properties of Diamond and Graphite

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The Structure of Graphite
It is a GIANT covalent structure formed in layers stacked on top of each other.
The distance between the layers is about 2.5 times the distance between the atoms within
each layer.
Each carbon atom uses three of its electrons to form simple covalent bonds with other
carbon atoms. The fourth electron in each carbon becomes delocalised ­ they are no longer
associated with any particular atom.
As the delocalised electrons move around the sheets of graphite, very large temporary
dipoles form which will induce other dipoles in the sheets above and below throughout the
whole graphite crystal.
The atoms in each sheet, held together by covalent bonds, are stronger than diamonds
because of the bonding caused by the delocalised electrons.
Properties of Graphite
Has a high melting point as covalent bonds are strong and need to be broken throughout the
whole crystal.
Has a low density because of the space `wasted' in-between the sheets.
Insoluble as attractions between solvent molecules and carbon atoms will never be strong
enough to overcome the covalent bonds.
Conducts electricity as the delocalised electrons are free to carry charge across the sheets.
The Structure of Diamond
Each carbon atom is covalently bonded to four other carbon atoms ­ forming four single
bonds.
It is a giant covalent structure ­ a macromolecular crystal.
Has no delocalised electrons.
Properties of Diamond
Has a very high melting point as there are very strong carbon-carbon covalent bonds
throughout the structure.
It is very hard because of the three dimensional covalent bonds.
Does not conduct electricity as there are no electrons which are free in order to carry the
charge.
Insoluble as there are no possible attractions which could occur between solvent molecules
and carbon atoms which could outweigh the attraction between the covalently bound carbon
atoms.

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