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.
A hydrogen bond is a weak attractive force between an H atom covalently bonded to a highly electronegative atom (F, O, N) and another nearby electronegative atom, and it is much weaker than a normal covalent or ionic bond.
When hydrogen is covalently bonded to a small, highly electronegative atom such as fluorine, oxygen, or nitrogen, the shared electron pair is pulled strongly towards that atom, leaving the hydrogen with a significant partial positive charge (δ+). This partially positive hydrogen is then electrostatically attracted to a lone pair on an electronegative atom of a neighbouring molecule (or a different part of the same molecule), forming a hydrogen bond.
Hydrogen bonds are classified as:
Intermolecular hydrogen bonds: formed between two different molecules, e.g., between water molecules (O−H···O), between HF molecules (F−H···F), or between NH3 molecules (N−H···N).
Intramolecular hydrogen bonds: formed within the same molecule, e.g., in ortho-nitrophenol.
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