There are a number of the molecules and ions in which basic modes of bond formation are fully adequate to explain the properties of the molecules or ions.


Carbon dioxide is represented as O = C = O, the carbon – oxygen bond length in CO_2, \text{is} 1.15 \AA while the expected value for double bond is 1.22 \AA and for triple is 1.10 \AA. It means that only double bond formula is inadequate for carbon dioxide.

Similarly, hydrogen molecule is represented as H – H. if it is so then the theoretical bond energy is quite low in comparison to the experimental value. It means that only covalent bond formula is inadequate for hydrogen molecule.

Therefore, when a molecule cannot be defined completely by a single structure but its characteristic properties can be described by two or more different structures, these structures are called resonance hybrid. These structures are as follows:

O = C = O \leftrightarrow \overset{-}{O}- C \equiv \overset{+}{O} \leftrightarrow \overset{+}{O} \equiv C- \overset{-}{O} \\ H^+H^- \leftrightarrow H- H \leftrightarrow H^-H^+


Hence, we can say that there is resonance between the various possible structures. Such structures are called resonating structures since due to resonance, molecule acquires minimum energy hence molecule has maximum stability. We can say resonance gives extra stability to the molecule. Such structures are shown by (\leftrightarrow) arrow.

Definition: It may be defined as:

(i) Resonance is the description of the electronic structure of molecules by means of several schemes of pairing electrons, with features of each scheme contributing to final description.

(ii) The individual pairing models are called resonance structures. The resonance structures put partial concepts of molecules and have physical meaning only in combination with all other pertinent resonance structures.

(iii) Resonance involves canonical forms in which certain atoms are required to bear unit charges. These charges are termed formal charges and their calculations are based upon the number of electrons in valency shell of the atom (shared electron counts as one half of a unit negative charge).

Principles of Resonance

(i) Whenever a molecule can be represented by two or more structures that differ from one another in the arrangement of electrons and not of atomic nuclei, the molecule is said to involve resonance. For example in CO^{2-}_3 ion C and O atoms are at the same position with different positions of electrons.

(ii) Resonance is particularly important when the contributing structures are of about the same stability.

(iii) The number of unpaired electrons in each contribution structure must be the same.

(iv) The more stable is the contributing structure the greater will be its contribution to the resonance hybrid.

(v) The resonance hybrid is more stable than any of the contributing structures by an amount known as resonance energy.

Resonance energy = [energy of most stable canonical form] – [experimental value of energy]

(vi) The greater is the number of contributing structures that can be written for a molecule, the more stable it would be.

(vii) In resonating structures, atoms should not undergo appreciable shift (Tautomerism).

(viii) Resonating structures must have very high energy do not contribute, lower the energy greater is the contribution.

(a) Structures having less number of bonds having higher energy.


(b)   Structures with unfavorable charge distribution do not coincide.

(c)    If adjacent atoms have the same charge, structures do not contribute


(ix)The canonical form should be of similar energies because a relatively vary high energy form is not expected to contribute much towards resonance.

(x) The relative position of various atoms involved must remain unchanged in all canonical forms.

Limitations of the Electronic Theory of Valency

Although electronic theory of valency provides a general explanation of chemical bonding even then there are many exceptions. For example, formation of BF_3, AlCl_3, PCl_5, SF_6, OSO_4 etc. molecules.

Transition metal ions do not follow the octet rule. The formation of one electron (H^+_2), three electron bond (O_3, NO, NO_2 etc.) metallic bond etc. cannot be explained by this theory.

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