S2.2.11 Resonance structures (Higher Level Only)

Delocalization of Electrons: Resonance and Its Applications

Imagine you're rearranging furniture in a room, trying to find the perfect layout. You realize there are two equally good options, but neither feels like the "true" setup because the room's functionality depends on a blend of both.

This is similar to what happens in certain molecules. Instead of furniture, we're dealing with electrons, and instead of a room, we're talking about molecular structures.

This phenomenon, where a molecule cannot be fully represented by a single structural formula but instead exists as a hybrid of multiple structures, is known as resonance.

What Is Resonance?

The Basics of Delocalization

In most covalent bonds, electrons are localized between two atoms. However, in some molecules, electrons can be delocalized, meaning they are shared across more than two atoms.

This occurs when there are multiple valid ways to position double or triple bonds while still satisfying the rules of chemical bonding.

When this happens, no single Lewis structure can accurately describe the molecule. Instead, the molecule is represented by resonance structures.

The actual molecule exists as a resonance hybrid, which is a blend of all the resonance structures.

Ozone ($O_3$)

Ozone consists of three oxygen atoms. Its Lewis structures suggest two possible arrangements for the double bond and lone pairs:

1. A double bond between the first and second oxygen atoms, with a single bond between the second and third.
2. A double bond between the second and third oxygen atoms, with a single bond between the first and second.

These two resonance structures can be represented as:

$$\text{O=O-O ↔ O-O=O}$$

Experimental evidence shows that both O–O bonds in ozone are identical in length and strength, intermediate between a single bond and a double bond. This indicates that the electrons in the double bond are delocalized across the entire molecule.

The true structure of ozone is a resonance hybrid, where the double bond is "spread out" over both O–O bonds.

Benzene ($C_6{H}_6$)

Benzene is an aromatic hydrocarbon with a six-membered ring of carbon atoms, each bonded to one hydrogen atom. Its Lewis structures suggest alternating single and double bonds:

However, X-ray diffraction studies reveal that all six carbon–carbon bonds in benzene are identical in length, shorter than a single bond but longer than a double bond.

This uniformity arises from delocalized electrons in the π-bond system, which form a continuous ring of electron density above and below the plane of the molecule.

The delocalized structure of benzene is often represented as a hexagon with a circle inside, symbolizing the shared electron cloud.

This delocalization makes benzene more stable than hypothetical structures with alternating single and double bonds. The extra stability provided by resonance is called resonance energy.

How Does Resonance Affect Bond Properties?

Bond Length and Strength

In molecules with resonance, bonds that appear as single or double bonds in individual resonance structures often have intermediate properties in the resonance hybrid. For example:

• In ozone, both O–O bonds have a bond order of 1.5 (average of a single and double bond).
• In benzene, all C–C bonds have a bond order of 1.5.

Bond length and strength data confirm these intermediate properties. For instance:

• A single O–O bond has a length of 148 pm and a bond enthalpy of 144 ${\mathrm{kJ}\mathrm{mol}}^{-1}$.
• A double O=O bond has a length of 121 pm and a bond enthalpy of 498 ${\mathrm{kJ}\mathrm{mol}}^{-1}$.
• The O–O bonds in ozone have a length of 128 pm and a bond enthalpy of 362 ${\mathrm{kJ}\mathrm{mol}}^{-1}$, consistent with a bond order of 1.5.

Why Does Resonance Matter?

Stability

Resonance increases a molecule's stability by delocalizing its electrons, which lowers the overall energy of the system. For example, benzene is much less reactive than alkenes due to the stability of its delocalized π-electrons, which resist disruption during chemical reactions.

Reactivity

Resonance also shapes how molecules react:

• In benzene, the delocalized π-electrons attract electrophiles, leading to electrophilic substitution reactions instead of addition reactions.
• In ozone, the delocalized electrons make the molecule reactive enough to absorb ultraviolet (UV) radiation, playing a critical role in shielding Earth's surface from harmful UV rays.

Reflection and Practice