pi bond

When two atoms want to bond more than once, the second connection cannot be another head-on grip — that spot is taken by the sigma bond. Instead, leftover p orbitals on each atom overlap sideways, above and below the line joining the nuclei, like two parallel logs touching along their lengths rather than end to end. That sideways overlap is a pi bond. It is the 'extra' bond that turns a single bond into a double or triple bond.

More precisely, a pi bond forms when two parallel p orbitals overlap side-by-side, placing the shared electron density in two lobes, one above and one below the sigma-bond axis, with none directly on the axis itself. This sideways overlap is less effective than head-on overlap, so a pi bond is weaker than a sigma bond. But it has two important consequences. First, it forbids rotation: to twist the two atoms relative to each other you would have to break the sideways p-orbital overlap, so a C=C double bond is rigid and locked, which is why cis and trans isomers exist. Second, the pi electrons sit exposed above and below the molecule, loosely held and easy for electrophiles to grab.

Pi bonds matter because they are where the action is. The reactive double bonds of alkenes, the triple bonds of alkynes, the carbonyl C=O, and the delocalized ring of benzene are all built from pi bonds, and their exposed, weakly-held electrons make them the favourite targets of addition and electrophilic reactions. The rule to carry everywhere: a double bond is one sigma plus one pi; a triple bond is one sigma plus two pi bonds. The sigma is the strong, rotatable skeleton; the pi is the reactive, rotation-locking extra.