Q.Distinguish between sigma and pi bond.

Covalent bonds are classified by the GEOMETRY of the atomic-orbital overlap that forms them. A sigma (σ) bond forms when the overlap lies directly ALONG the internuclear (bond) axis — from two 's' orbitals overlapping head-on (e.g. H2's s-s overlap), from an 's' and a 'p' orbital overlapping end-on (e.g. HF's s-p overlap), or from two 'p' orbitals overlapping end-on along that same axis (e.g. F2's p-p σ overlap, via each atom's half-filled $2p_z$ orbital). A pi (π) bond, by contrast, forms when two 'p' orbitals overlap SIDEWAYS (laterally) rather than end-on, so the resulting bond lies PERPENDICULAR to the internuclear axis, with electron density concentrated in two lobes, above and below that axis, and a node passing through the axis itself. Because end-on overlap is inherently more extensive/direct than sideways overlap, a σ bond is generally STRONGER than a π bond between the same pair of atoms — and, as a direct consequence, in every double or triple bond the FIRST bond formed is always a σ bond, with any additional bond(s) necessarily being π bonds (e.g. a C=C double bond is one σ + one π; a C≡C triple bond is one σ + two mutually perpendicular π bonds). In molecular orbital theory, the same σ/π classification carries over to molecular orbitals themselves: a σ molecular orbital is symmetric all the way around the bond axis, while a π molecular orbital is not, having instead a positive lobe on one side of the axis and a negative lobe on the other.