Source: Lecture notes / Organic Chemistry (The Ohio State University)
Tags: resonance, resonance structures, electron delocalisation, curved arrows, resonance hybrid, major resonance contributor, MRC, Lewis structures, pi bonds, lone pairs, organic chemistry
Difficulty: Introductory to Intermediate Prerequisites: Lewis structures, formal charge, electronegativity, atomic orbitals (sp, sp2, sp3 hybridisation), sigma and pi bonding.
Resonance is one of the first places in organic chemistry where the simple Lewis structure model stops being enough. A single Lewis structure often cannot capture how electrons are truly distributed in a molecule, so we draw multiple resonance structures and combine them into a more accurate picture called the resonance hybrid. This matters because electron distribution drives reactivity: where electrons sit determines where bonds form and break. If you are comfortable with Lewis structures and formal charge, you are ready for this material.
Resonance structures are different ways of drawing the same molecule, differing only in the placement of pi electrons and lone pairs while keeping the atoms fixed. The molecule does not flip between these forms; it exists as a single hybrid, a blend of all valid resonance structures. The structure that contributes most to the hybrid (the major resonance contributor) is the one that maximises octets, places charge on appropriate atoms, and minimises charge separation.
Resonance
The delocalisation of electron density across a molecule. In simple terms, this means electrons are spread out over multiple atoms rather than locked between just two.
Resonance structures
Two or more valid Lewis structures for the same molecule that differ only in the positions of lone pairs and pi bond electrons. Think of these as different "snapshots" of where the electrons could be, not different molecules.
Resonance hybrid
The single, real structure of the molecule, understood as a composite (weighted average) of all individual resonance structures. Dotted lines show delocalised bonds, and partial charges (δ+, δ−) show where electron density sits.
Major resonance contributor (MRC)
The resonance structure that best represents the actual electron distribution, sometimes called the "best resonance form." The molecule looks most like its MRC.
Curved arrows (electron-pushing arrows)
Arrows drawn on structures to show the flow of electrons from one position to another. Double-headed curved arrows move a pair of electrons (2e⁻); single-headed (fishhook) arrows move one electron (1e⁻). These are not reaction arrows.
Pi (π) bond
A bond formed by the sideways overlap of p orbitals, sitting above and below the plane of the bonded atoms. Pi electrons are the electrons that move in resonance.
Lone pair (LP)
A non-bonding pair of electrons on an atom. Lone pairs can participate in resonance when they occupy a p orbital adjacent to a pi system.
Formal charge
The charge assigned to an atom in a Lewis structure, based on the difference between its valence electrons and the electrons it "owns" (lone pair electrons plus half of bonding electrons).
Resonance is the delocalisation of electron density. This is sometimes called the "2nd Rule of O Chem": resonance is probably why something is happening in a given reaction or property.
Key point: resonance structures are not constitutional isomers. They represent the same molecule. The atoms stay in the same positions; only lone pairs and pi bond electrons move.
Same atomic connectivity. You cannot move atoms. If an atom is bonded to a different set of neighbours, you have drawn a constitutional isomer, not a resonance structure.
Same net charge. The total charge on the molecule must be identical across all resonance forms.
Valid Lewis structures. Each form must obey Lewis rules. Atoms can have fewer than an octet but never more (for second-row elements).
Only lone pair and pi bond electrons differ. Sigma bonds do not move. If you are moving a sigma bond, you are drawing a reaction, not resonance.
Curved arrows show the flow of electrons from a source (lone pair or pi bond) to a destination (forming a new pi bond or lone pair).
Double-headed curved arrow (↷): moves 2 electrons.
Single-headed curved arrow (fishhook, ↷ with half head): moves 1 electron.
These are distinct from reaction arrows:
A single forward arrow (→) means a forward reaction: A + B → C.
A double-headed equilibrium arrow (⇌) means a reversible reaction: A + B ⇌ C.
A double-headed resonance arrow (↔) means resonance: A ↔ B. This is the one used between resonance structures.
Do not confuse the resonance arrow (↔) with the equilibrium arrow (⇌). Resonance structures exist simultaneously as a hybrid; equilibrium implies two distinct species interconverting.
The molecule does not bounce between resonance forms. It exists as one thing: the resonance hybrid, which is the weighted average of all contributors.
In the hybrid:
Dotted or dashed lines replace some of the bonds to show delocalised electron density.
Partial charges (δ+ and δ−) replace full formal charges, reflecting the averaged distribution.
Example: formamide (HCONH₂) has two resonance structures. One places a double bond between C and O with a lone pair on N. The other places a double bond between C and N with a negative charge on O. The hybrid shows partial double-bond character in both the C–O and C–N bonds, with partial negative charge on O and partial positive charge on N.
When resonance structures are not equivalent, the one that contributes most to the hybrid is the MRC. Rank by these criteria, in order of importance:
1. Maximise octets
The structure in which more atoms (especially C, N, O) have a full octet is the better contributor. A structure that gives carbon an octet at the cost of putting formal charge on oxygen is usually favoured over one with an incomplete octet on carbon.
Example: for the cyanate ion (NCO⁻), the form :N≡C–O:⁻ is the MRC because all three atoms have octets, whereas :N=C=O:²⁻ leaves nitrogen without a full octet and stacks charge.
2. Place negative charge on more electronegative atoms
If both forms have full octets, the one with the negative formal charge on the more electronegative atom wins. Oxygen bearing a negative charge is more favourable than carbon bearing it.
Example: in an enolate, the form with negative charge on oxygen is the MRC over the form with negative charge on carbon.
3. Minimise separation of charges
A structure with charges on adjacent atoms (one positive, one negative) is higher in energy than one with no charge separation. The form with fewer formal charges, or charges closer to zero, is preferred.
Resonance is not a drawing exercise. It has three practical consequences:
Explains reactivity. Electron-rich sites (where electron density is delocalised to) attract electrophiles. Electron-poor sites attract nucleophiles. Resonance tells you where those sites are.
Shows how a molecule truly exists. A single Lewis structure is a simplification. The hybrid, informed by all resonance contributors, is closer to reality.
More resonance means more stability. Molecules with more valid resonance structures (more delocalisation) are generally lower in energy and more stable.
Resonance does not apply everywhere. Three structural cues tell you to look for it:
1. A charge or lone pair next to a pi bond (allylic position)
A lone pair or negative charge sitting on an atom directly adjacent to a C=C or other pi bond is the classic trigger. The electrons can delocalise into the pi system.
Example: an allyl anion, where the negative charge is next to a C=C double bond. The charge delocalises across three carbons.
2. A positive charge next to a double bond or lone pair
An empty p orbital (positive formal charge) next to a pi bond or an atom with a lone pair allows electron density to flow toward the positive centre.
Example: a carbocation adjacent to a C=C double bond, or adjacent to an atom bearing a lone pair (like O or N). The MRC often achieves an octet on the originally electron-deficient atom.
3. A pi bond between atoms of different electronegativity
When a pi bond connects two atoms that differ in electronegativity (e.g. C=O, C=N), one resonance form places extra electron density on the more electronegative atom. This is why carbonyl groups (C=O) have a δ+ carbon and a δ− oxygen.
When drawing resonance structures, do not violate rules unnecessarily. If a rule must be broken in every possible resonance form (for example, an expanded octet on a third-row element), that is acceptable. But if the violation is avoidable, that form is not a valid contributor.
The test is: does this resonance structure individually satisfy the rules of Lewis structures? If yes, it counts.
All resonance occurs in p atomic orbitals (the pi system).
Resonance is exclusively a phenomenon of electrons in p orbitals. Moving single (sigma) bonds would create constitutional isomers, not resonance structures. Only lone pairs and pi bonds participate, because those electrons occupy p orbitals that can overlap sideways.
1. The atom has no p orbital or pi system.
An sp3-hybridised atom with no adjacent pi bond cannot participate in resonance. All of its orbitals are used for sigma bonds and sp3 lone pairs (which sit in hybrid orbitals, not pure p orbitals).
Example: the nitrogen in a simple amine like NH₃ bonded to an sp3 carbon does not delocalise its lone pair, because there is no adjacent pi system for it to join.
2. The atom has both a lone pair and a pi bond but the lone pair is in the wrong orbital.
If an atom's p orbital is already occupied by a pi bond, its lone pair sits in an sp2 hybrid orbital instead, pointing in the plane of the molecule rather than perpendicular to it. That lone pair cannot overlap with the pi system.
Example: a pyridine-type nitrogen. The lone pair on N is in an sp2 orbital in the plane of the ring, while the p orbital perpendicular to the ring is already part of the aromatic pi system. So the lone pair does not participate in resonance with the ring.
"The molecule flips back and forth between resonance structures." It does not. The resonance hybrid is the one real structure; the individual forms are bookkeeping tools.
"Resonance structures are different molecules." They are not. They are different drawings of the same molecule. Atoms stay fixed; only electrons move.
"Any lone pair next to a pi bond will always participate in resonance." Not necessarily. The lone pair must be in a p orbital that can overlap with the pi system. If it is in an sp2 or sp3 hybrid orbital in the wrong orientation, it cannot delocalise.
"More resonance structures always means the molecule is more stable." Quality matters as much as quantity. A molecule with two strong (low-energy) resonance contributors can be more stabilised than one with five poor contributors.
⚠️ You will be asked to draw all valid resonance structures for a given molecule. Know the three cues (LP/charge next to pi bond, positive charge next to LP/pi bond, polarised pi bond).
⚠️ You will be asked to identify the MRC. Apply the three criteria in order: octets first, then electronegativity of charged atoms, then charge separation.
⚠️ Do not confuse the resonance arrow (↔) with the equilibrium arrow (⇌). Examiners test this directly.
⚠️ Drawing a resonance structure that moves sigma bonds is a common error that costs marks. Only lone pairs and pi bond electrons move.
⚠️ Understanding resonance is foundational. It will reappear in acid/base chemistry, carbonyl chemistry, aromatic chemistry, and reaction mechanisms throughout the course.
1. True or false: resonance structures are constitutional isomers of each other.
False. They are different electron arrangements of the same molecule with the same atomic connectivity.
2. Fill in the blank: in the resonance hybrid, bonds that are delocalised are shown as ______ lines, and charges are shown as ______ charges.
Dotted (or dashed); partial (δ+ / δ−).
3. True or false: the major resonance contributor is the form with the most charge separation.
False. The MRC has the least charge separation (all else being equal).
4. Fill in the blank: the three criteria for choosing the MRC, in order of priority, are (1) maximise ______, (2) place negative charge on the more ______ atom, (3) minimise ______ of charges.
(1) octets, (2) electronegative, (3) separation.
5. True or false: resonance can involve the movement of sigma bond electrons.
False. Only lone pair and pi bond electrons participate in resonance.
Q: What is the difference between a resonance arrow (↔) and an equilibrium arrow (⇌)?
A: A resonance arrow indicates that two structures are resonance contributors to the same hybrid, existing simultaneously. An equilibrium arrow indicates two distinct species that interconvert over time.
Q: A molecule has a lone pair on nitrogen adjacent to a C=C double bond. Can this lone pair participate in resonance? What condition must be met?
A: It can, provided the lone pair occupies a p orbital that is parallel to the p orbitals of the pi system. If the nitrogen's lone pair is in an sp2 hybrid orbital in the plane of the molecule (as in pyridine), it cannot overlap with the pi system and does not participate.
Q: You draw two resonance structures for an anion. In structure A, oxygen bears the negative charge. In structure B, carbon bears the negative charge. Both structures have full octets on all atoms. Which is the MRC and why?
A: Structure A is the MRC. When octets are equal, the negative charge should sit on the more electronegative atom. Oxygen is more electronegative than carbon.
Q: Why does increased resonance stabilisation generally lead to lower energy (greater stability)?
A: Delocalisation spreads electron density over a larger volume, lowering the energy of the system. Concentrating charge on a single atom is energetically costly; distributing it is favourable.
Q: An sp3-hybridised carbon is bonded to an oxygen with a lone pair. Is there resonance between the oxygen lone pair and that C–C single bond? Why or why not?
A: No. The carbon is sp3-hybridised and has no p orbital available for pi overlap. Resonance requires a pi system (p orbital participation). A lone pair adjacent to only sigma bonds cannot delocalise through resonance.
This material connects directly to acid-base chemistry: resonance stabilisation of a conjugate base makes the parent acid stronger. It also underpins carbonyl chemistry (the polarisation of C=O through resonance explains why nucleophiles attack the carbon) and aromatic chemistry (benzene's stability comes from full delocalisation of pi electrons around the ring). Understanding resonance here sets up every reaction mechanism you will encounter later in the course.
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