Resonance Structures: What They Are and How to Draw Them

Try to draw a Lewis structure for ozone, O₃, and you’ll find two equally good options: a double bond on the left and a single bond on the right, or the other way round. Yet experiments show that both O–O bonds in ozone are identical in length. Neither drawing is correct on its own. Resonance is how chemists handle this: the real molecule is a blend — a hybrid — of all the valid structures. This article explains what resonance really means, how to draw resonance structures correctly, and why it makes molecules more stable.

The problem: one drawing isn’t enough

Lewis structures assume each pair of electrons belongs either to one atom (a lone pair) or to one bond between two atoms. That works well for many molecules, such as water and methane. But in some molecules, electrons are delocalised — spread over three or more atoms. A single Lewis structure can’t show this.

Ozone is the classic example. Its central oxygen is bonded to two others. A valid Lewis structure is:

O=O–O⁻ (with formal charges +1 on the central O and −1 on the single-bonded O)

But you could equally draw ⁻O–O=O. Experiment shows:

• Both O–O bonds are 128 pm long.
• A typical O–O single bond is 148 pm; the O=O double bond in O₂ is 121 pm.

So each bond is somewhere between single and double. The electrons in the π bond are shared across all three oxygen atoms.

Resonance structures and the resonance hybrid

Resonance structures (or resonance forms, or contributing structures) are two or more valid Lewis structures for the same molecule that differ only in the positions of electrons — not atoms.

They are connected by a double-headed arrow (↔).

The real molecule is the resonance hybrid: a single structure with the electrons spread out, which is a weighted average of the contributors.

An important warning

Resonance does not mean the molecule flips back and forth between the structures. Ozone doesn’t switch between two forms millions of times a second. It has one real structure all the time — the hybrid. The individual resonance structures are just our imperfect way of describing it using Lewis drawings.

A helpful analogy: a mule is a hybrid of a horse and a donkey. It isn’t a horse some of the time and a donkey the rest — it’s always a mule. Similarly, a rhinoceros might be described to someone who’s never seen one as “a cross between a unicorn and a dragon”. The rhino is real; the unicorn and dragon aren’t. Resonance structures are like the unicorn and the dragon.

How to draw resonance structures

Rules

1. Atoms don’t move. Only electrons (lone pairs and π bonds) change position. If you move an atom, you’ve drawn a different molecule (an isomer), not a resonance structure.
2. Only π electrons and lone pairs move. Sigma (σ) bonds stay in place.
3. Every structure must be a valid Lewis structure: the same total number of electrons, and no period 2 atom with more than eight electrons.
4. The overall charge stays the same in every structure.
5. Show electron movement with curly arrows: each arrow starts at a lone pair or a π bond and points to where the electron pair goes.

Step-by-step method

1. Draw one valid Lewis structure, including lone pairs and formal charges (see formal charge).
2. Look for a π bond next to an atom with a lone pair, a π bond next to a positive charge, or alternating double and single bonds.
3. Use curly arrows to move a lone pair into a bond (making a new π bond) and a π bond onto an atom (making a new lone pair).
4. Recalculate the formal charges in the new structure.
5. Repeat until you’ve found all the structures.

Worked examples

The carbonate ion, CO₃²⁻

One Lewis structure has one C=O double bond and two C–O⁻ single bonds. There are three equivalent positions for the double bond, giving three resonance structures:

O=C(–O⁻)₂ ↔ ⁻O–C(=O)–O⁻ ↔ (⁻O)₂C=O (with the double bond in each position)

In the hybrid:

• All three C–O bonds are identical — about 129 pm, between single (143 pm) and double (about 120 pm).
• Each C–O bond has a bond order of 4 bonds ÷ 3 positions = 1⅓.
• The 2− charge is spread evenly: each oxygen carries −⅔.

The nitrate ion, NO₃⁻

The same pattern: three equivalent structures, each N–O bond with a bond order of 1⅓, and each oxygen carrying an average formal charge of −⅔. The nitrogen carries a formal charge of +1 in every structure, so the total is +1 + 3(−⅔) = −1.

The ethanoate (acetate) ion, CH₃COO⁻

Two resonance structures, with the double bond on one oxygen or the other. Both C–O bonds are equal (bond order 1½), and the negative charge is shared between the two oxygens. This delocalisation stabilises the ion, which is a big part of why carboxylic acids are much more acidic than alcohols: when ethanoic acid loses H⁺, the negative charge is spread over two oxygens, whereas in the ethoxide ion from ethanol it’s stuck on one.

Benzene, C₆H₆

Benzene is the most famous case. The Kekulé structure — a six-membered ring of alternating single and double bonds — can be drawn in two ways, with the double bonds swapped round. Both are resonance structures.

In reality:

• All six C–C bonds are identical, 139–140 pm long — between C–C (154 pm) and C=C (134 pm).
• The six π electrons are delocalised in a ring above and below the plane of the carbon atoms.
• Chemists often draw benzene as a hexagon with a circle inside to show this.

Benzene is also far less reactive than you’d expect for a molecule with three C=C bonds. It doesn’t decolourise bromine water under normal conditions, because an addition reaction would destroy the stable delocalised ring. It undergoes substitution instead.

Not all resonance structures are equal

When resonance structures aren’t equivalent, some contribute more to the hybrid than others. The major contributors are the structures that:

1. Have full octets on as many atoms as possible.
2. Have the fewest formal charges.
3. Put negative charges on more electronegative atoms (and positive charges on less electronegative ones).
4. Avoid placing like charges on adjacent atoms.

For example, in the thiocyanate ion (SCN⁻), the structure with the negative charge on nitrogen contributes a bit more than the one with it on sulfur, because nitrogen is more electronegative.

Resonance energy: why delocalisation stabilises

The resonance hybrid is always lower in energy — more stable — than any single contributing structure. The extra stability is called the resonance energy or delocalisation energy.

Benzene gives a measurable value. Adding hydrogen to cyclohexene (one C=C) releases about 120 kJ mol⁻¹. If benzene were really a ring with three separate double bonds, adding hydrogen to all three should release about 3 × 120 = 360 kJ mol⁻¹. The measured value for benzene is only about 208 kJ mol⁻¹. The difference, about 150 kJ mol⁻¹, is benzene’s delocalisation energy — the extra stability that comes from spreading its π electrons around the ring.

These data are a classic piece of evidence against the simple Kekulé structure. Bond length measurements by X-ray crystallography (Kathleen Lonsdale’s work on hexamethylbenzene in 1929 showed a flat ring with equal bond lengths) provide the other.

Why resonance matters

• Bond lengths and strengths fall between single and double values (see bond length and bond strength).
• Acidity: carboxylic acids and phenols are more acidic because their anions are stabilised by resonance.
• Reactivity: benzene prefers substitution to addition; amides are less basic than amines because the nitrogen lone pair is delocalised into the C=O group.
• Colour: long delocalised systems absorb visible light — the colours of dyes, carrots (β-carotene) and chlorophyll come from delocalised electrons (see chlorophyll).
• Conductivity: graphite conducts electricity because of delocalised electrons across its layers.

Common misconceptions

• “The molecule switches between resonance structures.” It doesn’t; the hybrid is a single, unchanging structure.
• “Resonance structures show different molecules.” They’re different drawings of the same molecule; atoms never move.
• “All resonance structures contribute equally.” Only equivalent ones do; others contribute in proportion to their stability.
• “Resonance adds energy.” It lowers energy — delocalisation stabilises.

Key takeaways

• Resonance describes molecules whose electrons are delocalised and can’t be shown by one Lewis structure.
• Resonance structures differ only in electron positions; they’re linked by ↔, and the real molecule is the hybrid.
• Move only lone pairs and π electrons with curly arrows; never move atoms.
• Delocalisation gives intermediate bond orders (1⅓ in carbonate, 1½ in ethanoate and benzene) and extra stability (benzene ≈ 150 kJ mol⁻¹).

To practise drawing the starting structures, try Lewis structures practice.