Introduction

When the bonding in a molecule or polyatomic ion cannot be represented by a single Lewis formula, the phenomenon of resonance is used to describe the behavior of the delocalized electrons involved. Several resonance structures can be used to depict a molecule or ion with delocalized electrons. The Lewis nuclear framework The only thing that changes about these resonance structures is the position of the electrons in them. They are extremely important in chemistry and aid in deciphering the structures of several compounds. These bonds are central to the valence bond theory, which attempts to provide a mechanistic account of the chemical interactions between different elements.

History

This idea was initially offered in 1899 by Johannes Thiele in his "Partial Valence Theory" to explain the unexpected stability of benzene, which followed August Kekulé's structure described in 1865, which had consecutive single and double bonds.

Unlike most alkenes, benzene goes through substitution reactions rather than addition reactions. He contended that benzene's carbon-carbon bond is a single and double-bond hybrid.

The resonance concept also partially explained a large number of benzene derivative isomers. If we apply Kekulé's structure to the molecule dibromo benzene, for instance, we get predictions for four different isomers comprising two ortho isomers in which the brominated carbon atoms are connected by either a single or double bond. In actuality, there are just three di-bromobenzene isomers, and only one of them is ortho, which is consistent with the notion that there is only one sort of carbon-carbon bond, which lies between a single and a double bond. 

Mechanism of Resonance

The double bonds in some compounds may be twisted or rearranged, known as resonance or mesomerism. The mesomerism increases the compound's stability and, thus, its reactivity. Covalent chemistry describes the bonds between some chemical compounds, such as benzene, ozone, etc. When the Lewis structure fails to express and explain delocalized electrons sufficiently, it also refers to a strategy for expressing and explaining them.

It occurs when a nonbonding electron is present or when a pi bond is altered. The pi-electron locations or even the positions of nonbonding electrons cause an atom's orbital to change. Each of the resonance structures must possess the same degree of energy. It happens in unsaturated systems due to the delocalization of electrons during this process.

Resonance Energy

Resonance energy is linked to the concept of aromaticity. Resonance energy is used as a first step in determining the stability of a resonating structure. Because the bonding in a molecule cannot be expressed uniquely by a single Lewis structure, the presence of resonance energy can be used to identify the delocalized electrons present in the molecule.

The resonance energy may be calculated by comparing the energies of two compounds, one with a localized Lewis structure and the other with a delocalized real structure.

A mathematical representation of resonance energy follows

∆E = Edelocalized - Elocalized. 

The more the value of resonance energy, the more stable the compound.

Examples of resonance structures

Benzene:

One of the most important hydrocarbons for research in organic chemistry is benzene (C₆H₆). It has a cyclic structure with alternating single and double bonds. Benzene's ring structure allows for two distinct resonance modes, one of which involves delocalizing the pi-electrons. 

The six carbon atoms that make up benzene are all sp2 hybridized, with an unhybridized p orbital running perpendicular to the ring plane. Because of the equivalence of the p orbitals of the 6 carbon atoms, it is inconceivable for them to overlap with only one neighboring p orbital to form three distinct double bonds. The six p orbitals instead overlap cyclically, each with its neighbor. This allows for greater overlap between the p orbitals than would be achieved from the linear 1,3,5-hexatriene analogue since the p orbitals are delocalized into molecular orbitals that stretch around the ring.

The ring must be planar for this to occur, because without it, the p orbitals would not be able to overlap appropriately, and benzene is typically thought of as a planar molecule.