Like first member of other groups, carbon also differs from rest of the members of its group. It is due to its smaller size, higher electronegativity, higher ionisation enthalpy and unavailability of d orbitals.
In carbon, only s and p orbitals are available for bonding and, therefore, it can accommodate only four pairs of electrons around it.
This would limit the maximum covalence to four whereas other members can expand their covalence due to the presence of d orbitals.
Carbon also has unique ability to form p∏– p∏ multiple bonds with itself and with other atoms of small size and high electronegativity.
Few examples of multiple bonding are: C=C, C≡C, C = O, C = S, and C ≡ N. Heavier elements do not form p– p bonds because their atomic orbitals are too large and diffuse to have effective overlapping.
Carbon atoms have the tendency to link with one another through covalent bonds to form chains and rings. This property is called catenation.
This is because C—C bonds are very strong. Down the group the size increases and electronegativity decreases, and, thereby, tendency to show catenation decreases.
This can be clearly seen from bond enthalpies values. The order of catenation is C > > Si > Ge Sn. Lead does not show catenation.
Due to property of catenation and p∏– p∏ bond formation, carbon is able to show allotropic forms.