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General Chemistry I · Chemical Bonding & Molecular Geometry

Resonance

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On this page 7 sections
  1. In 30 seconds
  2. Why this matters
  3. The college version
  4. Eli explains
  5. Key takeaway
  6. Study tools
  7. Sources & references

In 30 seconds

Sometimes a single Lewis structure cannot accurately represent a molecule because the bonding electrons are not fixed in one location. Resonance occurs when two or more valid Lewis structures (differing only in electron placement, not atom positions) can be drawn. The true structure is a resonance hybrid — a single, real molecule in which the electrons are delocalized over several atoms. It does not flip back and forth between the individual forms.

Why this matters

Resonance explains why ozone's two oxygen–oxygen bonds are identical, why the carbonate and nitrate ions are symmetrical, and why benzene is a flat, especially stable ring (delocalized π electrons). It is central to organic chemistry and biochemistry — the stability of DNA bases, proteins, and many drugs depends on delocalized electrons. Getting resonance right also corrects a pervasive misconception: bonds in such species are neither single nor double but something in between.

The college version

Key Ideas

  • Resonance forms: two or more Lewis structures that differ only in the placement of electrons (never in the positions of atoms).
  • Resonance hybrid: the actual structure, an average of the resonance forms; it is more stable than any single form would predict.
  • Delocalization: the bonding electrons are spread over several atoms rather than confined to one bond.
  • Equivalent forms: all forms have the same energy and contribute equally (e.g., O₃, NO₃⁻, CO₃²⁻).
  • Non-equivalent forms: forms with different formal-charge arrangements contribute unequally; the form with charges closest to zero (and negative charge on the most electronegative atom) dominates.
  • Fractional bond order: when bonds are equivalent in the hybrid, the bond order is the total number of shared pairs divided by the number of bonds.
  • Resonance stabilizes a molecule: the delocalized electrons are lower in energy than they would be in any single structure.

Equations and Variables

  • Bond order (fractional) = (total shared electron pairs across the resonance forms) / (number of equivalent bonds in the hybrid).
    • Ozone O₃: 3 shared pairs over 2 bonds → bond order = 3/2 = 1.5.
    • Nitrate NO₃⁻: 4 shared pairs over 3 bonds → bond order = 4/3 ≈ 1.33.
  • Formal charge (FC = V − N − ½B) is used to rank non-equivalent resonance forms.

How It Works

  1. Draw the molecule's Lewis structure.
  2. If the double (or triple) bond can be placed on a different equivalent atom without moving any atoms, draw those additional forms.
  3. Recognize that no single form is "the" structure: the real molecule is the hybrid, with delocalized electrons.
  4. The hybrid's bond lengths are equal and intermediate between single and double (e.g., O₃ has two equal O–O bonds ≈ 128 pm, between a single ≈148 pm and double ≈121 pm).
  5. If the forms are non-equivalent, weight them by formal charge (favor charges near zero and negative charge on the more electronegative atom).

Worked Example

Draw the resonance forms of ozone (O₃) and state the bond order.

O₃ has 3 × 6 = 18 valence electrons. The two forms differ only in which terminal oxygen carries the double bond:

O=O–O and O–O=O

Each form has one double bond and one single bond. The hybrid delocalizes the double-bond character over both O–O bonds, so:

  • Bond order = (1 + 2)/2 = 1.5 for each O–O bond.
  • Both O–O bonds are equal in length (~128 pm), longer than a double bond but shorter than a single bond.
  • The central O carries a +1 formal charge and each terminal O carries a −1 formal charge in each form, averaged across the hybrid.

Second example — nitrate (NO₃⁻): three equivalent forms, each with one N=O double bond and two N–O single bonds. Bond order = 4/3 ≈ 1.33 for every N–O bond; all three are equal.

Common Confusions

  • "The molecule flips back and forth between the resonance forms." — Wrong: it is a single, static hybrid; the forms are just drawing conveniences.
  • "Resonance forms are different molecules in equilibrium." — Wrong: the ↔ arrow denotes resonance, not an equilibrium (which uses ⇌).
  • "Each O–O bond in ozone is sometimes single and sometimes double." — Wrong: both bonds are identical, always 1.5.
  • "You can move atoms when drawing resonance forms." — Wrong: only electrons move; the atom skeleton is fixed.
Eli, the EliExplains learning guide

Eli explains

The same idea, in plain words

Explain it like I’m 10

Imagine a three-lane road where a single car (the double bond) can travel in any lane, but there's only one car. The traffic report describes the car as "spread across all three lanes" — that's the resonance hybrid. The car is not teleporting back and forth between lanes; rather, it's genuinely smeared out. For ozone, the "extra" bond is smeared across both oxygen–oxygen connections, so each is halfway between single and double. (The analogy's limit: the "car" is really a pair of delocalized electrons, and the smearing is a quantum property, not literal motion.)

Key takeaways

  • Resonance forms differ only in electron placement, never in atom positions.
  • The real molecule is a resonance hybrid with delocalized electrons — it does NOT oscillate between forms.
  • Delocalization lowers energy → resonance stabilization.
  • O₃: two equivalent forms, bond order 1.5, equal O–O bonds (~128 pm).
  • NO₃⁻ and CO₃²⁻: three equivalent forms, bond order ≈ 1.33.
  • Equivalent forms contribute equally; non-equivalent forms are weighted by formal charge.
  • A double-headed arrow (↔) connects resonance forms (it does not mean equilibrium).
  • Resonance = two or more valid Lewis structures differing only in electron placement.
  • True structure = resonance hybrid (delocalized electrons), more stable than any single form.
  • NOT oscillating; NOT equilibrium.
  • Fractional bond order: O₃ = 1.5, NO₃⁻/CO₃²⁻ ≈ 1.33.
  • Equivalent forms contribute equally; non-equivalent forms weighted by formal charge.
  • Equal, intermediate bond lengths are the experimental fingerprint of resonance.

Keep learning

Ready to build on this? Continue to the next lesson.

Practice General Chemistry I

This lesson has no separate scored set. Practice draws from the subject’s question bank.

Study tools & related lessonsYou’ll learn to · Related

You’ll learn to

  • Explain why some molecules require more than one Lewis structure.
  • Define a resonance hybrid and contrast it with the incorrect "oscillating" picture.
  • Calculate fractional bond orders from resonance forms.
  • Distinguish equivalent from non-equivalent resonance structures.

Sources & references

  1. OpenStax, *Chemistry 2e*, "7.4 Formal Charges and Resonance."
  2. OpenStax, *Chemistry 2e*, "7.3 Lewis Symbols and Structures."
  3. OpenStax, *Chemistry 2e*, "8.3 Multiple Bonds."

This lesson was adapted from the open educational references above; their licenses and attributions are preserved. See Copyright & Licensing.

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