# Resonance

Resonance Structures & Hydbrid Orbitals
15:38 — by Dennis H.

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## Key Questions

Resonance refers to the existence of numerous forms of a compound, and is a component of valence bond theory.

#### Explanation:

Resonance refers to the existence of numerous forms of a compound, and is a component of valence bond theory. Shown below are the two resonance structures of benzene:

These appear to be simple mirror images of each other, but this in fact represents a process that occurs continuously throughout the lifetime of a benzene ring. Observe that benzene is composed of both $\text{C - C}$ single bonds, and $\text{C = C}$ double bonds.

Since a $\text{C = C}$ double bond is stronger than a $\text{C - C}$ single bond, due to the presence of a $\pi$ bond, it follows that the $\text{C = C}$ bonds must have a shorter bond length than the $\text{C - C}$ bonds. The bond length of a $\text{C - C}$ single bond is $154 \text{pm}$, whilst the bond length of a $\text{C = C}$ double bond is $133 \text{pm}$.

Yet the bond length shown in benzene is $139 \text{pm}$. Not only that, but all of the bonds have the same bond length. How can this be? Molecular orbital theory is able to explain this, because resonance is integral to its conception, but valence bond theory must take a different tact if it wishes to explain this.

Recall that $\pi$ bonds result from the extension of half-filled, unhybridised p atomic orbitals from the bonding plane, which can overlap side-on - perpendicular to the $\sigma$ bond. In doing this, the overlapping regions will result in a single $\pi$ bond that contains a pair of bonding electrons. When this happens in benzene, we see an alternating pattern of $\text{C - C}$ single bonds and $\text{C = C}$ double bonds, as shown in the first diagram.

But a p atomic orbital on, say, ${\text{C}}^{1}$, can overlap with, to form a $\pi$ bond, a p orbital on ${\text{C}}^{2}$ or ${\text{C}}^{6}$, can it not? Upon consideration, neither of the two local carbon atoms is any more favorable than the other for this purpose. Because of this, valence bond theory postulates that the structure resonates: it holds in one formation for an instant, before moving onto the next one, then back again, because neither one is more stable than the other.

This actually strengthens the structure, because what appear to be simple $\text{C - C}$ single bonds actually experience double bonded properties instantaneously. The reason for the intermediate bond length in benzene is due to the fact that these bonds are effectively intermediate in identity between single and double bonds.

Resonance plays an important structural role in covalent compounds in accordance with valence bond theory. It can be applied not only to benzene, but also to other compounds such as ozone and carbocations.

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