Resonance energy

In chemistry, resonance or mesomerism is a way of describing delocalized electrons within certain molecules or polyatomic ions where the bonding cannot be expressed by one single Lewis structure. A molecule or ion with such delocalized electrons is represented by several contributing structures (or forms) (also variously known as resonance structures (or forms), canonical structures, or, in older works or translations, mesomers), which collectively constitute a resonance hybrid. (The concept of a resonance hybrid is unrelated to orbital hybridization.)

Under the framework of valence bond theory, resonance is an extension of the idea that the bonding in a chemical species can be described by a Lewis structure. For many chemical species, a single Lewis structure, consisting of atoms obeying the octet rule, possibly bearing formal charges, and connected by bonds of positive integer order, is sufficient for describing the chemical bonding and rationalizing experimentally determined molecular properties like bond lengths, angles, and dipole moment. However, in some cases, more than one Lewis structure could be drawn, and experimental properties are inconsistent with any one structure. In order to address this type of situation, several resonance structures are considered together as an average, and the molecule is said to be represented by a resonance hybrid in which several Lewis structures are used collectively to describe its true structure. For instance, in NO₂^–, nitrite anion, the two N–O bond lengths are equal, even though no single Lewis structure has two N–O bonds with the same formal bond order. However, its measured structure is consistent with a description as a resonance hybrid of the two major contributing structures shown above: it has two equal N–O bonds of 125 pm, intermediate in length between a typical N–O single bond (145 pm in hydroxylamine, H₂N–OH) and N–O double bond (115 pm in nitronium ion, [O=N=O]⁺). According to the contributing structures, each N–O bond is an average of a formal single and formal double bond, leading to a true bond order of 1.5. By virtue of this averaging, the Lewis description of the bonding in NO₂^– is reconciled with the experimental fact that the anion has equivalent N–O bonds.

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