General Organic Chemistry
17.0 Resonance
17.0 Resonance
Many organic compounds are known which are not adequately represented by single valence bond structures. All the properties of such a compound are not explained by a single structure. Thus, the valence bond structure of benzene indicates three each carbon-carbon single and double bonds.
Measurement of bond distances in benzene, however shows that all the carbon-carbon distances are the same and it is $1.4\mathop A\limits^ \circ $ i.e. between a single bond ($1.54\mathop A\limits^ \circ $) and a double bond ($1.34\mathop A\limits^ \circ $).
Further benzene on hydrogenation gives out heat which is less than the heat calculated from its structure.
For such compounds, it is necessary that other structures be devised to explain all their properties. This is the basis of the concept of resonance. Thus, when two or more good Lewis structures can be devised for a compound, resonance is invoked.
The different structures of a compound devised by different methods of pairing electrons in a fixed atomic skeleton are called resonance or canonical structures.
The actual structure of the compound is then a combination of these structures and hence the compound is called a resonance hybrid.
A hybrid is more stable than any one of the contributing structures. The contributing resonance structures are shown by double-headed arrow $\left( \leftrightarrow \right)$ indicating that the real structure involves both ways of pairing electrons.
Each resonance structure represents only partially to the real state of the compound and all of them in combination represent the compound completely. Final structural description of the molecule is then what is obtained on superimposing all the resonance structures on one another. The actual molecule therefore does not vibrate or oscillate from one structure to another but has one and only one structure which is an average of all the structures, The different resonance structures do not exist, they are drawn by different schemes of pairing electrons to consider the extent of delocalization and hence to access the stability of the hybrid.
The $C_1-C_2$ and $C_3-C_4$ bonds are found to be longer than a carbon-carbon double bond and $C_2-C_3$ bond is slightly shorter than a carbon-carbon single bond. To explain this, several other structures may be devised by shifting a pair of $\pi$ electrons.
The structures (B) and (D) show partial single bond character of $C_1-C_2$ and $C_3-C_4$ bonds and structures (C) and (E) show partial double-bond character of $C_2-C_3$ bond. This explains the observed anomalies in the bond distances in the hybrid.
Thus, in the hybrid structure (i.e. real structure) of the compound, lateral overlaps of all the four $p$ atomic orbitals have taken place, i.e., the electrons are delocalized.
As a result, two bonding and two antibonding MOs are formed. The hybrid representation shows that each pair of electrons binds four carbon nuclei instead of two. This gives a net stability to the hybrid over any one of the contributing resonance structures.
Experimental confirmation about the stability of the hybrid is afforded from the qualitative measurements of the heat of combustion or hydrogenation.
[Diagram]
It is seen that the actual compound gives out lesser energy than that calculated for the structure of 1,3-butadiene. Hence, the actual compound is at a lower-energy state, i.e., more stable.
The difference in the experimental and calculated energies is the amount of energy by which the compound is stable. This difference in the energies is known as the resonance or delocalization energy (RE). It is about $4\ kcal/mol$ for 1,3-butadiene and $36\ kcal/mol$ for benzene which is also a resonance hybrid.
