2.6 Resonance

Resonance & Formal Charge

Work through all three parts of topic 2.6 on 12 species: equivalent contributors and their hybrid, nonequivalent diagrams that formal charge has to choose between, and the odd-electron cases where the Lewis model runs out

Resonance HybridFractional Bond OrderOctet vs ChargeOdd Electron
Topic 2.6

Resonance and Formal Charge

Represent a molecule with a Lewis diagram that accounts for resonance between equivalent structures or that uses formal charge to select between nonequivalent structures.

Sometimes more than one valid Lewis structure can be drawn for the same arrangement of atoms. EK 2.6.A.1 says that when those structures are equivalent, resonance must be included as a refinement — and that this refinement is often needed to make qualitatively correct predictions about structure and properties.

Resonance does not mean the molecule flickers between forms. The real molecule is a single structure that is an average of all contributors, with the electrons delocalized over the whole set of positions. Consequences you can predict:

  • All bonds in the resonance set are identical in length and strength. All three N–O bonds in NO₃⁻ are the same, even though any single Lewis structure shows one double and two singles.
  • The bond order is fractional: NO₃⁻ has bond order 4/3, and its bonds are shorter than an N–O single bond but longer than an N=O double bond.
  • Delocalization stabilizes the species — the basis for why carboxylate ions are much more stable than alkoxide ions (see 8.6).

When candidate structures are nonequivalent, EK 2.6.A.2 says the octet rule and formal charge are the criteria for choosing the best model:

FC = (valence electrons) − (lone-pair electrons) − ½(bonding electrons)

The preferred structure has formal charges as close to zero as possible, and any negative formal charge placed on the most electronegative atom. Formal charges must sum to the overall charge of the species.

Finally, EK 2.6.A.3: as with any model, Lewis structures have limits. They fail particularly for species with an odd number of valence electrons — NO and NO₂ cannot give every atom an octet no matter how you draw them.

Key points

  • Resonance means delocalization, not oscillation. The real structure is one averaged structure.
  • Equal resonance contributors → all those bonds are equal in length and strength.
  • Best nonequivalent structure: formal charges nearest zero, negative FC on the most electronegative atom.
  • Odd-electron species (NO, NO₂) expose the limits of the Lewis model.

Equations

  • not on the sheetNot on the equation sheet; you are expected to know it.
    • valence electrons of the free atom
    • electrons in lone pairs on that atom
    • electrons in bonds to that atom

Common mistakes

  • The double-headed arrow (↔) is not an equilibrium arrow. The forms do not interconvert.
  • Formal charge is a bookkeeping device, not a real charge. Do not confuse it with oxidation number or with a partial charge δ.
  • Resonance requires the same atom positions. Moving an atom gives a different compound (an isomer), not a resonance form.
  • Formal charges must sum to the species charge. If they do not, the structure is wrong.

Worked example

Two Lewis structures can be drawn for OCN⁻: one with a C≡N triple bond and O–C single bond, and one with C=N and C=O double bonds. Use formal charge to decide which is the better model.

Structure A: O–C≡N with 3 lone pairs on O and 1 on N.
O: 6 − 6 − ½(2) = −1
C: 4 − 0 − ½(8) = 0
N: 5 − 2 − ½(6) = 0
Sum = −1 ✓ (matches the ion charge)

Structure B: O=C=N with 2 lone pairs on O and 2 on N.
O: 6 − 4 − ½(4) = 0
C: 4 − 0 − ½(8) = 0
N: 5 − 4 − ½(4) = −1
Sum = −1 ✓

Decision: both have formal charges of the smallest possible magnitude, so the tiebreaker is where the negative charge sits. Oxygen (EN 3.5) is more electronegative than nitrogen (EN 3.0), so structure A, with the −1 on oxygen, is the better model.

Full notes for topic 2.6 →