Resonance Theory (c. 1945)

There are several pages on the Avogadro Web Site that refer to the delocalisation of electrons within chemical structures. Where this has been observed, some additional stability has been associated with the structures concerned. What follows is a brief insight into Resonance Theory.

Whilst it is common to represent a chemical structure in a consistent and recognisable way, often the structural formula shown does not adequately represent the distribution of electrons within it. Rather, it is necessary to represent the molecule (or ion) by various alternative structures, with the true structure, and its distribution of electrons, lying somewhere between the extremes.

Benzene provides a first example¼

The benzene molecule is often represented as a hexagonal ring of alternate single and double carbon-carbon bonds. But, given one such structure drawn on paper, the single and double bonds can be shown to have changed positions simply by re-distributing the pairs of electrons. These are the two Kekulé structures of benzene.

Move the mouse across the structure to change between the Kekulé forms of benzene.

The true structure of benzene lies somewhere in between the Kekulé forms, and is referred to as a RESONANCE HYBRID of these.

It is essential to appreciate this idea; above all else, it is important to realise the resonance hybrid is not a mixture of the canonical forms, nor that the forms are rapidly alternating from one to another in a dynamic equilibrium.

The different structures contributing to the resonance hybrid are referred to as CANONICAL FORMS of the molecule. The more stable a canonical form, the more it contributes to the resonance hybrid. The resonance hybrid is more stable than any one of its canonical forms, and the difference in energy between the resonance hybrid and the most stable canonical form is known as the RESONANCE ENERGY or DELOCALISATION ENERGY. Generally, the greater the number of canonical forms of about the same energy, the greater the resonance energy.

Modern physical analysis confirms the resonance hybrid nature of benzene. A simple and convenient way of representing the benzene ring is to place a circle inside the hexagonal ring.

Resonance Energy of Benzene

Add a little liquid bromine to an alkene, such as cyclohexene, and an immediate reaction is observed. An electrophilic addition reaction occurs with the decolorisation of the bromine. Add liquid bromine to benzene and no such similar observation is made. If a reaction is taking place then it does so very slowly indeed. Now add some iron filings or iron(III) bromide and the formation of white fumes of hydrogen bromide gas is a sign of a reaction, in which electrophilic substitution is taking place.

This is an illustration of the additional stability afforded to the benzene ring owing to the delocalisation of electrons.

The resonance energy of benzene can be determined by comparing the enthalpies of hydrogenation of benzene and cyclohexene. If benzene did have the hypothetical 'cyclohexatriene' structure (the Kekulé structure), then it would be expected to have an enthalpy of hydrogenation that is three times that of cyclohxene.


The enthalpy of hydrogenation of cyclohxene is -121 kJ mol-1, providing a theoretical value of -363 kJ mol-1 for Kekulé benzene with its three separate carbon-carbon double bonds.

In fact, the actual enthalpy of hydrogenation of benzene is only -209 kJ mol-1. Therefore, 154 kJ mol-1 of energy less than might have been expected is evolved. This is the energy that must have been expended in overcoming the additional stability conferred upon the benzene ring, owing the delocalisation of electrons. It is called the resonance energy or delocalisation energy.

In general, substances in which electrons are fairly evenly distributed over the entire or part of the molecules tend to be more stable and less reactive than those in which electrons are localised. This is because the polarity of chemical bonds is kept to a minimum, making bond breaking more difficult.

Here's another example¼

Ethanoic acid, Phenol, and Ethanol can be considered as weak acids. Their relative strengths as acids are:

Ethanoic acid > Phenol > (Water) > Ethanol

This can be explained (in part) in terms of the stability of their respective conjugate bases; the more stable a base, the greater in the tendency for it to form. The relative base strengths are:

Ethoxide ion > Phenoxide ion > (Hydroxide ion) > Ethanoate ion

Both the Ethanoate ion, CH3COO-, and the Phenoxide ion, C6H5O-, allow for some delocalisation of electrons in their structures. There is no such delocalisation of electrons in the Ethoxide ion, CH3CH2O-.

The diagrams below attempt to illustrate the delocalisation of electrons in the anions where this exists.

The comparative stability of the Ethanoate ion is attributable to the delocalisation of electrons as represented by the two canonical forms. Move the mouse pointer over the canonical forms to represent them as a single resonance hybrid structure. This structure also illustrates the equivalence of both carbon-oxygen bonds and the equal distribution of negative charge between the oxygen atoms.

The Phenoxide ion is also stabilised by resonance. The lone pair of oxygen electrons are considered to interact with, and form part of, the delocalised electron system of the benzene ring. Move the mouse pointer over the canonical forms to represent them as a single resonance hybrid structure.

The ability of alcohols to liberate hydrogen in their reaction with alkali metals illustrates their acidic nature, but their acidic strengths are much weaker than even that of water. Ethanol, for example, is a weaker acid than Phenol which in turn is a weaker acid than Ethanoic acid. There is no delocalisation of electrons in the Ethoxide ion.