Multiple bonding

Like with the diatomics multiple bonds are provided by the overlap of pπ orbitals perpendicular to the direction of the bond, in difference to the σ orbitals that point in the bond direction. A simple instance is ethene, C₂H₄. In which the planar structure of the molecule results from sp² bonding with each carbon forming two σ bonds to hydrogens and one to the other carbon. The p orbitals not involved in hybrids are directed perpendicular to the molecule and can overlap to form the π bonding MO displayed, that is occupied by two  electrons. The combination of σ+π  MOs gives a double C=C  bond. The Maximum bonding overlap of the π orbitals relies on the coplanar arrangement of atoms and there is a considerable hurdle to rotation about double bonds, not like the single bonds where groups can rotate fairly freely. Triple bonds (example.  in C₂H₂)  are provided  by the overlap of two  sets of perpendicular pπ orbitals, like in diatomics like N₂ and CO.

In some  examples in which a localized description of σ bonding is possible this is not so for the π bonds. An instance is the carbonate ion CO₃^2- where a resonance picture  is essential in simple models. Figure 2b depicts the planar framework with sp² bonding in the central atom. The pπ AOs of four atoms can overlap together to create a delocalized MO as displayed. Out of the three orbital combinations  feasible for the three oxygen πAOs only one can overlap and bond with carbon in this way. There are two others that is not shown, that remain nonbonding on the oxygen. So one π bonding MO is distributed  over three C-O bonds, with nonbonding  charge density subsequent to two MOs  distributed  over the three oxygen atoms.  This is necessarily identical to the resonance picture.