• # question_answer 20) Explain the difference in properties of diamond and graphite on the basis of their structures.

In diamond, each carbon is in $s{{p}^{3}}-hybridised$state and linked to four other carbon atoms tetrahedrally giving a ant three tensional polymeric structure. As the atoms are firmly by strong covalent bonds, diamond is hardest substance, possesses very high melting point and is chemically inert. Since, there is no mobile electron present, diamond is non-conductor of electricity. However, it is an excellent thermal conductor. Graphite has two dimensional sheet structure. Each carbon is in $s{{p}^{2}}-hybridised$and is linked to three other carbon atoms m a hexagonal plannar structure. There are $\pi$-bonds in the structure. The $\pi$-electrons are free to move throughout the entire layers, i.e., graphite is a good conductor of electricity. The adjacent layers can easily slide over each other, hence graphite is soft and possess low density. It is chemically more active than diamond.