Resonance hybridization, which occurs in molecules or ions with resonance structures, has an effect on bond length. In resonance, multiple Lewis structures, called resonance structures, are used to represent a molecule or ion that cannot be adequately described by a single Lewis structure.
The resonance hybrid is a combination or average of all the resonance structures, and it represents the actual electronic structure of the molecule or ion. The delocalization of electrons in resonance structures leads to the following effects on bond length:
Bond Lengths between Resonance Structures: In resonance structures, the actual bond lengths are intermediate between the bond lengths predicted by the contributing Lewis structures. For example, if a molecule has two resonance structures, one with a double bond and one with a single bond, the actual bond length will be shorter than a typical single bond but longer than a typical double bond.
Delocalization of π-electrons: Resonance allows for the delocalization of π (pi) electrons over multiple atoms. This delocalization leads to electron density being spread out over a larger region of space. Consequently, the electron-electron repulsions are reduced, which results in a slight shortening of the bonds involved in the resonance.
Equivalent Bond Lengths: In some cases, resonance structures contribute equally to the resonance hybrid, and all the bonds between the same atoms are of equal length. For example, in benzene, all the carbon-carbon bonds have the same length due to the equal contribution of resonance structures.
In summary, resonance hybridization affects bond length by leading to intermediate bond lengths between contributing resonance structures. The delocalization of π-electrons and the equal contribution of resonance structures can result in a slight shortening of bonds involved in resonance and the presence of equivalent bond lengths, respectively.