Resonance

RESONANCE

Valence shell electrons of the atoms in a molecule are either shared or unshared. The shared electrons are found in either σ or π bonds. Unshared electrons are found in AOs (usually hybrid AOs for first-row elements). Lewis struc-tures provide a way to indicate the shared and unshared pairs of electrons in molecules. Sometimes, however, it is possible to indicate the electron distri-bution in molecules by more than one Lewis structure. For example, a car-boxylate anion can be represented by two equivalent but different Lewis structures.

These structures are equivalent because they have the same numbers of bonds, unshared pairs of electrons, and the same charge. They are different because the negative charge is located on different oxygen atoms. Moreover the bonds from carbon to a particular oxygen are double in one structure and single in the other. When more than one correct Lewis structure can be written for a molecule, each structure is a resonance form of the molecule. The actual molecule is a resonance hybrid of the contributing resonance forms, and its properties result from a combination of the properties of the contributing resonance forms. Thus each oxygen atom carries a – 1/2 charge, and the bonds between carbon and each oxygen atom have a bond order of 1.5 and are of the same length.

A very good analogy is a mule. A mule is a hybrid of a horse and a donkey. A mule is neither a horse nor a donkey but it has properties of each. The reso-nance hybrid of the carboxylate anion is a resonance hybrid of the contributing resonance forms and has properties of each.

Another classic example of resonance is the benzene molecule. The localized resonance forms are termed Kekule´ forms (after Friedrich August Kekule,´ who first deduced the structure of benzene) and have alternating single and double bonds between carbon atoms. The actual benzene molecule is a resonance hybrid of the contributing resonance forms as the bond lengths are equal (single and double bonds have different lengths).

The bond order is between one (single) and two (double). The resonance hybrid is often pictured with a circle in the ring to indicate the delocalized electron distribution in the molecule.

Double-headed arrows are used to indicate resonance forms. It is important to note that resonance forms are not in equilibrium, just as a mule is not a horse part of the time and a donkey the rest of the time.

The presence of resonance forms means that the electrons are not localized between two nuclei but are delocalized over more than two nuclei. The result of electron delocalization is that electrons are attracted by a greater number of nuclei, which leads to a lower energy for the molecule and hence greater stability. Simply put, resonance delocalization is a stabilizing feature of molecules.

A molecule for which resonance forms can be written is more stable than any of the contributing resonance forms. Thus the carboxylate ion (a resonance hybrid) is more stable than either of the contributing resonance forms. The dif-ference in energy between the energy of the molecule and the energy of the most stable resonance form is the resonance energy (RE) of the molecule. The reso-nance energy represents the stabilization of the molecule due to the delocalization of electrons.

The amount of resonance energy is related to the relative energies of the con-tributing resonance forms. The greatest resonance stabilization is found when the contributing resonance forms are degenerate (equal) in energy. Thus molecules such as the carboxylate ion, benzene, the allyl anion, and the allyl cation all have significant resonance stabilization because the main resonance contributors are of the same energy.

In contrast, resonance stabilization is less in an amide because the resonance forms A₁ and A₂ given below are very different in energy. Nevertheless, because an amide is a resonance hybrid of A₁ and A₂, it is predicted that there should be some double-bond character in the bond between carbon and nitrogen. This is in fact the case since many amides show restricted rotation around the C–N bond (typical of a π bond). Moreover, the nitrogen atom in amides is nearly planar and not very basic, also indicating that the lone pair is delocalized.

It is also generally true that the greater the number of contributing resonance forms, the greater will be the resonance stabilization. For this reason the enolate of a β-diketone has much more resonance stabilization than the enolate of a simple ketone (three resonance forms versus two). The electrons are delocalized over five atoms in the former versus three atoms in the latter. In addition, the electron density on the carbon atom is less in the diketone enolate than in a simple methyl ketone enolate.

Resonance has a significant influence on the electron distributions and energies of molecules. The delocalization of electrons is described by the contributions of resonance forms, which are themselves localized structures with discrete bonds. Such structures are known as valence bond (VB) structures, and this approach to the description of bonding in molecules is called the valence bond approach. As long as one keeps in mind that resonance forms are limiting VB structures and that the actual molecule is a resonance hybrid of these VB structures, a great deal of insight into the structure and properties of molecules can be gained.