Imagine you're trying to draw a photograph of a friend who is laughing. A single still frame captures one expression, but it misses the movement, the energy, the in-between of the laugh. A single Lewis structure does the same thing to certain molecules — it freezes them into one arrangement of electrons, but the real molecule is more like a short video clip, with electrons moving smoothly between positions.
Take ozone, O3. If you try to draw a Lewis structure, you get a dilemma. You can put the double bond on the left:
O=O−O
Or on the right:
O−O=O
Both satisfy the octet rule. Both have the same atoms. But which one is correct? Neither, alone. The real ozone molecule has two identical O−O bonds — each is halfway between a single and a double bond. No single Lewis picture can show that.
The Solution: Resonance Structures
Resonance structures are a set of two or more Lewis structures that collectively describe the actual electronic structure of a molecule where a single Lewis structure is inadequate. They are connected by a double-headed arrow (↔) to show they are not different molecules, but different ways of drawing the same molecule.
Important
Resonance structures are not real, separate molecules that flip back and forth. They are imaginary "snapshots" that we average together to get the true structure. The real molecule is a resonance hybrid — a blend of all contributing structures.
The Rules (Precise Statement)
Same atomic positions. Only electrons (pi bonds and lone pairs) move; atoms never move.
Same total number of electrons. You are redistributing, not adding or removing.
Valid Lewis structures. Each resonance form must obey the octet rule (for second-period elements) and have correct formal charges.
Curved arrows show electron movement. An arrow from a lone pair or a pi bond points to where those electrons go next.
How to Draw Them: The Curved Arrow Method
Take the nitrate ion, NO3−. Start with one valid Lewis structure:
O∣∣O−N=O−
Now, push electrons:
Take the lone pair on the top oxygen (the one with the negative charge) and push it down to form a double bond with nitrogen.
Simultaneously, push the existing double bond on the right up to become a lone pair on that oxygen.
You get a second structure:
O=N−O∣O−−
Repeat the process from this new structure, and you get a third. All three are resonance structures of NO3−.
Tip
A quick way to spot resonance: look for a pi bond next to an atom with a lone pair (or a pi bond next to a positive charge). That's the classic "conjugated system" that allows electrons to delocalize.
The Hybrid: What the Molecule Actually Looks Like
The resonance hybrid is not an average of the bond lengths — it is the actual molecule. In NO3−, all three N−O bonds are identical, with a bond order of 131 (one and one-third). The negative charge is spread equally over all three oxygens, not stuck on one.
You represent the hybrid by drawing dashed lines for partial bonds and placing the charge in a circle (or using fractional charges) to show delocalization.
Linearity here follows the classic electron-count rule: 16-valence-electron triatomics with no lone pair on the central atom (BeCl₂, CS₂) are linear; NO₂ (17 electrons, odd electron on N) is bent, and NCO⁺ (14 valence electrons) does not adopt the linear 16-electron geometry. The answer is (i) and (iv).
Species by species
BeCl2 — beryllium contributes two bond pairs and keeps no lone pair; the two Be–Cl bonds spread to 180∘. Linear ✓
NCO+ — valence electrons: 5+4+6−1=14. The familiar linear species of this family (CO₂, NCO⁻, N₂O) all have 16 valence electrons; removing two electrons from cyanate changes the electronic structure so that the 16-electron linear picture no longer applies. Not grouped with the linear pair. ✗ …
Same / Similar Concept — real previous-year questions on the same or a closely similar concept, not this exact question.
CBSE 2026Set ANNUAL2 marks
Q.Draw resonating structures of CO3^2- (carbonate ion).
›Reveal solutionSolution
CO3^2- is a resonance hybrid of three equivalent Lewis structures, each with one C=O and two C-O(-) bonds in a different position.
Carbonate ion, CO3^2-, has a central carbon bonded to three oxygen atoms in a trigonal planar arrangement. A single Lewis structure would show carbon forming one C=O double bond to one oxygen and two C-O single bonds to the other two oxygens (each of those two oxygens carrying a formal -1 charge, giving overall charge -2). However, since all three oxygen atoms are chemically identical and equivalent, we can draw three such structures, differing only in which of the three oxygens carries the double bond:
Structure I: double bond to O(1), single bonds with negative charge to O(2) and O(3)
Structure II: double bond to O(2), single bonds with negative charge to O(1) and O(3)
Structure III: double bond to O(3), single bonds with negative charge to O(1) and O(2)
Ozone's true structure is a resonance hybrid of two equivalent Lewis structures, each with a double bond on one side and a single bond on the other, giving both O-O bonds equal, intermediate bond length.
Ozone (O3) is a bent, angular molecule with a central oxygen atom bonded to two terminal oxygen atoms. A single Lewis structure would require one O-O bond to be a double bond and the other a single bond, with a formal positive charge on the central O and a formal negative charge on the singly-bonded terminal O. But experimentally both O-O bonds in ozone are found to be identical in length (intermediate between a typical single and double O-O bond) -- this is explained by resonance.
Structure I: (terminal O)=O(central, +)-O(terminal, single bond, -) -- double bond on the left oxygen, single bond (with negative charge) on the right oxygen.
Structure II: (terminal O, -)-O(central, +)=(terminal O) -- single bond (with negative charge) on the left oxygen, double bond on the right oxygen.
When one Lewis structure can't capture a molecule's real bonding, chemists draw several valid alternatives and treat the true structure as a blend (hybrid) of all of them.
Definition. For certain molecules or ions, a single Lewis structure fails to explain all of the observed properties (such as equal bond lengths where a single structure would predict different single and double bonds). In such cases, the molecule is described using two or more Lewis structures, called canonical forms or resonance structures, which differ only in the position of electrons (not in the position of atomic nuclei). The actual structure of the molecule is considered to be a resonance hybrid — a weighted average of all these contributing structures — and does not correspond to any single one of them.
Example — Ozone (O3).
Ozone can be drawn with the terminal O=O double bond on either side, giving two equivalent canonical structures. Experimentally, both O-O bonds in ozone are found to have identical, intermediate bond length/bond order (between that of a single and a double bond) — consistent with the real molecule being a hybrid of both resonance structures, not switching between them.
Example — Carbonate ion (CO3^2-).
The carbonate ion can be drawn with the double bond to any one of the three oxygens, giving three equivalent canonical structures. The real ion has all three C-O bonds of equal, intermediate length (delocalised over all three oxygens).