Resonance and Hybrid Structures

Most molecules and ions are represented by a single Lewis structure. However, in particular compounds, some electrons are delocalized over multiple bonds or atoms rather than localized to a specific bond or atom. These compounds can be represented accurately with multiple Lewis structures.

Consider the Lewis structure for sulfur trioxide. The single bonds between each oxygen and the central sulfur atoms satisfy the octet for the oxygen atoms. However, to reach a full octet for the sulfur, an additional bond must be formed between sulfur and one of the oxygen atoms.

These multiple Lewis structures of sulfur trioxide are called resonance structures or contributing structures. The spatial positions of all of the component atoms remain the same; however, the valence electrons are distributed differently.

These double-headed arrows can be thought of as commas and should not be confused with resonance structures being in equilibrium. In these molecules, the actual structure is the weighted average or a hybrid of its resonance structures.

Resonance structures are identified by ‘electron pushing’, or transforming lone pairs into bonds and vice versa, as denoted with curved arrows. This maps the delocalized areas.

Such delocalization results in resonance stabilization—that is, a molecule with a lower potential energy than that of any theoretical non-delocalized structure.

If a contributing structure is lower in energy than another, it more closely resembles the actual molecular structure. Certain preferences are used to estimate the relative energies of various contributing structures.

Firstly, structures in which all atoms have filled valence shells are more stable. Secondly, structures with a greater number of covalent bonds are more stable. Thirdly, structures that minimize formal charges are more stable.

Finally, when other factors are equal, structures with negative charges on more electronegative atoms and positive charges on less electronegative atoms are more stable.