Therefore, the real structure of benzene cannot be represented by either. This would mean that two adjacent carbon atoms of the benzene molecule are joined neither by a single nor a double bond. In terms of the theory of resonance propounded by Pauling, the two valence bond structures ‘must be superimposed, fused together, to represent the molecule of benzene, with consideration given also to the stabilizing effect of resonance’. According to this theory, several possible Lewis structures for benzene, such as 8, 9 and others called canonical forms, are drawn, and energy of each is calculated from the wave equation for each structure. The concept of the real structure as the weighted average of these canonical forms is called resonance, and the difference in energy between the actual molecule and the canonical form of the lowest energy is called resonance energy.
In benzene, the resonance energy has been calculated from the heats of combustion, and from the heats of hydrogenation. The values obtained by either approach are close to each other, and work out to about 36 kcal/mol or 150 kJ/mol.
As shown above, the resonance is depicted using a double-headed arrow (<-> ) which is placed between each pair of the contributing forms. It is important to note that the double-headed arrow does not signify the oscillation of molecule from one Lewis structure to another. In fact, the word resonance does not mean the mixing of formal structures. Therefore, a less confusing term, π-electron delocalization, has been coined for this phenomenon.
The true structure of such a species is considered to be a hybrid of the various Lewis structures that can be written for it. The structure which differs only in the arrangement of electrons are called contributing or resonance structures, and we have only an idea of the true structure by considering the properties of two or more Lewis structures, which contain the same π electrons localized in various ways. This way of studying the properties of a molecule is called resonance method.
The concept of resonance stabilization is of particular importance for a qualitative understanding of reaction mechanism. Most of the organic reactions involve the formation of reactive intermediates such as carbocations, carbanions, and free radicals. Any variation in their resonance stabilization often changes the course of the reaction. Addition of bromine to 1,3-butadiene, resulting in the formation of two dibromo products, offers an illustrative example. The formation of the product of 1,4-addition can be rationalized by assuming the intermediacy of a carbocation, in which the positive charge is divided between 2 and 4 carbons. Addition of bromide ion then occurs at either of the two partially charged carbon atoms.
It is significant that initial attack of Br+ takes place at the position 1 in preference to the 2-position of the diene, as in the latter case the intermediate carbocation (CH2=CH-CHBr-CH2+) is unable to gain resonance stabilization.
Rules for the Resonance Method
The following rules for the use of this method enable us to predict when and to what extent resonance is important in determining the properties of various molecules:
(i) Structures contributing towards the hybrid must conform to real Lewis structures, i.e. there can be no structure with pentacovalent carbon and bicovalent hydrogen.
(ii) It involves the movement of only π or n-electrons, and a resonance structure can be derived from another by a series of one or more short electron shifts without changing the relative position of atoms. In other words, if two or more double bonds, or a double bond and an unshared pair of electrons are in conjugation, there is delocalization of π or n-electrons. For example, the cyclopropenyl cation can be represented as:
(iii) All resonance structures must have the same number of paired electrons. Thus, .CH2–.CH2 is not a resonance form of CH2=CH2.
(iv) All these structures do not contribute equally to the hybrid. The relative energies of various contributing structures are estimated by considering bond energies. Since each bond adds about 50-100 kcal/mol or 209-418 kJ/mol to the stability of the system, it is to be expected that structures with more covalent bonds are more stable than those with fewer. The nonpolar form of butadiene (10), for instance, is the most stable of all the resonance structures. The other forms, 11 and 12 makes a minor contribution, if at all, to the resonance hybrid.
(v) When atoms of different electronegativities are involved, the structure which places a negative charge on the most electronegative atom would be more important. Thus, for a ketone, 13 is more important than 14.
(vi) Resonance contribution is the greatest when there are two or more equivalent contributing forms of lower energy. The species involved may be neutral or ionic. For example, resonance in allyl cation, cyclopentadienyl anion, and benzene involves contributing forms of comparable energies and renders them quite stable. It may be noted that ionic species are especially stabilized if resonance is possible in them. This is due to the movement of electrons, there is a dispersal of charge over the whole system rather than being localized at a particular atom.
(vii) If there is only a single contributing structure of lowest energy, the resonance hybrid possesses properties expected for that structure. Ethylene and butadiene are, thus, well represented by their conventional structures, but benzene and carbonate anion is not.
(viii) These structures involving charge separation are less stable than nonpolar structures.
(ix) These structures having like charges on adjacent atoms are of high energy and hence unimportant. Structure 18 plays a minor role in the stabilization of 17.
(x) In order to achieve maximum overlap of p-orbitals, which is necessary for delocalization, it is essential that the skeleton should remain planar in a conjugated system.