Structure 2.2.11 — Resonance Structures

 Resonance structures (also known as resonance forms) are two or more valid Lewis structures that can be drawn for a molecule or ion, where the position of electrons (especially π electrons and lone pairs) varies, but the arrangement of atoms remains the same.

Key Concepts:

• Resonance occurs in species where delocalized electrons are shared between more than two atoms.
• Instead of being fixed in a single position (as in a double bond), the electrons are spread out (delocalized) over multiple atoms.
• The true structure of the molecule is a resonance hybrid – an average of all valid resonance structures, not flipping between them.
• Double-headed arrows (↔) are used to indicate resonance between different Lewis structures.
• Resonance increases stability by spreading out charge or electron density.

Important Notes:

• Only electrons move between resonance forms — atom positions do not change.
• All resonance structures must follow the octet rule where applicable and be valid Lewis structures.
• Resonance is especially common in aromatic compounds, carboxylates, and polyatomic ions such as nitrate or sulfate.

Example:

• \( \text{CO}_3^{2-} \) (carbonate ion): Delocalized π electrons over three equivalent oxygen atoms.

Delocalization:

This refers to the spreading of electrons across multiple atoms. In resonance structures, π electrons are not associated with a single bond or atom, but are instead shared over a region (e.g. entire ring or carboxylate group).

Visual Representation:

For example, for the carbonate ion \( \text{CO}_3^{2-} \):

• Each resonance form has one double bond and two single bonds to oxygen.
• The actual structure has bond lengths between single and double bonds (delocalized).

Bond Order in Resonance Structures:

The bond order is the average number of bonds across all resonance forms.

For carbonate \( (\text{CO}_3^{2-}) \): \( \text{Bond order} = \frac{\text{Total number of bonds}}{\text{Number of positions}} = \frac{4}{3} \approx 1.33 \)