11.12. Explain the difference in properties of diamond and graphite on the basis of their structures.

<p>Diamond has a crystalline lattice where each carbon atom undergoes sp³ hybridisation and linked to four other carbon atoms by using hybridised orbitals in tetrahedral fashion.&nbsp; The C–C bond length is 154 pm. The structure extends in space and produces a rigid three- dimensional network of carbon atoms. It is very difficult to break extended covalent bonding and, therefore, diamond is a the hardest substance on the earth.</p><p>Graphite has layered structure in which the layers are held by van there Waals forces and distance between two layers is 340 pm. Each layer is composed of planar hexagonal rings of carbon atoms. C—C bond length within the layer is 141.5 pm. Each carbon atom in hexagonal ring undergoes sp² hybridisation and makes three sigma bonds with three neighbouring carbon atoms. Fourth electron forms a? bond. The electrons are delocalised over the whole sheet. Electrons are mobile and, therefore, graphite conducts electricity along the sheet. Graphite cleaves easily between the layers and, therefore, it is very soft and slippery. That is why graphite is used as a dry lubricant in machines running at high temperature, where oil cannot be used as a lubricant.</p>

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