Resonance structures are a way to represent molecules or ions whose electrons cannot be shown accurately with a single Lewis structure. They are especially important when pi electrons or lone pairs are spread over several atoms. Resonance helps explain why some bonds have equal lengths, why charges are distributed, and why certain molecules are more stable than one drawing suggests.
It is a drawing tool for electron delocalization, not a picture of a molecule rapidly switching shapes.
A resonance contributor differs from another contributor only in the placement of electrons, not in the positions of atoms. Curved arrows show how electron pairs move from lone pairs or bonds to form new pi bonds, lone pairs, or charges. The real structure is the resonance hybrid, which combines the electron distribution of all important contributors.
Major contributors usually have full octets, fewer formal charges, and negative charge on more electronegative atoms.
Understanding Chemistry: Resonance Structures
A useful way to build resonance drawings is to start with one correct Lewis structure, then search for a connected set of p orbitals. A p orbital may come from a double bond, a lone pair, an empty orbital, or an atom carrying a positive charge. The atoms in this set need to be next to one another so their orbitals can overlap.
Only electrons within that connected region can be rearranged. Sigma bonds form the basic framework of a molecule and usually stay fixed. This rule prevents a common mistake, which is moving hydrogen atoms or breaking single bonds while drawing contributors.
Formal charge is a bookkeeping tool that helps compare possible contributors. Count the valence electrons assigned to each atom after a structure is drawn. An atom receives all of its nonbonding electrons and one half of the electrons in each bond.
A sensible contributor keeps second row atoms, such as carbon, nitrogen, oxygen, and fluorine, within an octet whenever possible. It avoids unnecessary charge separation.
If charges must appear, negative charge is usually better placed on an atom that attracts electrons strongly, such as oxygen. These patterns do not make a structure impossible or possible by themselves, but they give a reliable ranking for most school level examples.
Some contributors matter far more than others. In the acetate ion, the two oxygen atoms have the same surroundings. The two valid contributors therefore have equal importance.
The negative charge is shared across both oxygens, and both carbon oxygen bonds have the same intermediate character. In a molecule with unlike atoms, the contributors are often unequal. For example, a form with a negative charge on oxygen usually has more importance than one with that charge on carbon.
This unequal sharing affects measurable properties. Bond lengths can fall between typical single and double bond lengths, while charge density can be spread across several atoms.
Resonance has direct consequences in reactions. A charged particle is often less reactive when its charge is spread out, because no one atom carries the full charge concentration. This helps explain why carboxylate ions are relatively stable after a carboxylic acid loses a hydrogen ion.
It also helps explain the basic behavior of amides, where a nitrogen lone pair is shared toward a nearby carbon oxygen group. That lone pair is less available to bond to a hydrogen ion, so amides are weaker bases than many amines.
When studying reaction mechanisms, check whether a new charge can be delocalized after a bond changes. A product or intermediate with more effective delocalization is often lower in energy.
Curved arrow practice becomes clearer when every arrow begins at electrons. Start an arrow at a lone pair or at a bond, never at a positive charge or an atom symbol. End it where the electron pair will be located, either in a new bond or as a lone pair.
After each move, recount bonds, lone pairs, octets, and formal charges. Do not force resonance into every molecule.
A lone pair beside a saturated carbon cannot usually join a delocalized system because there is no suitable p orbital pathway. Careful orbital connection, rather than memorizing pictures, is the main skill.
Key Facts
- Resonance contributors have the same atom positions but different electron placement.
- Curved arrows show electron-pair movement, not atom movement.
- Formal charge = valence electrons - nonbonding electrons - 1/2 bonding electrons.
- The resonance hybrid is more stable than any single contributor.
- Equivalent resonance contributors contribute equally to the hybrid.
- Bond order in a resonance hybrid can be fractional, such as 1.5 for equivalent single and double bond sharing.
Vocabulary
- Resonance contributor
- One valid Lewis structure that helps describe the delocalized electron arrangement in a molecule or ion.
- Resonance hybrid
- The real electron distribution of a molecule or ion represented as a blend of its resonance contributors.
- Delocalized electrons
- Electrons that are spread over three or more atoms instead of being confined between two atoms or on one atom.
- Curved arrow
- A notation that shows the movement of an electron pair from an electron-rich region to a new location.
- Formal charge
- The charge assigned to an atom in a Lewis structure by comparing its assigned electrons with its valence electrons.
Common Mistakes to Avoid
- Moving atoms while drawing resonance structures is wrong because resonance only changes electron placement, not the molecular skeleton.
- Using a double-headed equilibrium arrow between contributors is wrong because resonance contributors are not separate species in equilibrium.
- Breaking the octet rule for second-row atoms like carbon, nitrogen, oxygen, or fluorine is wrong because these atoms cannot have more than eight valence electrons.
- Choosing the contributor with the most separated charges as the best structure is wrong because major contributors usually minimize formal charge and place negative charge on more electronegative atoms.
Practice Questions
- 1 For ozone, O3, draw the two equivalent resonance contributors and calculate the average O-O bond order in the resonance hybrid.
- 2 For carbonate, CO3^2-, draw the three equivalent resonance contributors and calculate the average C-O bond order.
- 3 Explain why the resonance hybrid of carbonate has three identical C-O bonds even though each resonance contributor shows one C=O bond and two C-O bonds.