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S1.1 - Introduction to the particulate nature of matter S1.2 - The nuclear atom S1.3 - Electron configurations S1.4 - Counting particles by mass - The mole S1.5 - Ideal gases S2.1 - The ionic model S2.2 - The covalent model S2.3 - The metallic model S2.4 - From models to materials S3.1 - The periodic table - Classification of elements S3.2 - Functional groups - Classification of organic compounds R1.1 - Measuring enthalpy changes R1.2 - Energy cycles in reactions R1.3 - Energy from fuels R1.4 - Entropy and spontaneity AHL R2.1 - How much? The amount of chemical change R2.2 - How fast? The rate of chemical change R2.3 - How far? The extent of chemical change R3.1 - Proton transfer reactions R3.2 - Electron transfer reactions R3.3 - Electron sharing reactions R3.4 - Electron-pair sharing reactions

S2.2 - The covalent model

2.2.1 Covalent Bonds and Lewis Formulas 2.2.2 Bond Types 2.2.3 Co-coordination (Dative) Bonds 2.2.4 VSEPR Shapes of Molecules 2.2.5 Electronegativity and Bond Polarity 2.2.6 Polarity and Dipole Moments 2.2.7 Covalent Network Structures and Allotropes 2.2.8 Intermolecular Forces 2.2.9 Physical Properties of Covalent Substances 2.2.10 Chromatography and Intermolecular Forces 2.2.11 Resonance Structures (AHL) 2.2.12 Benzene and Resonance (AHL) 2.2.13 Expanded Octet and VSEPR (AHL) 2.2.14 Formal Charge (AHL) 2.2.15 Sigma and Pi Bonds (AHL) 2.2.16 Hybridization (AHL)

Resonance Structures and Delocalization HL Only

Specification Reference S2.2.11

Quick Notes:

  • Resonance occurs when a molecule or ion has more than one valid Lewis structure with different positions for double bonds.
  • The actual structure is a resonance hybrid – a blend of all possible resonance forms.
  • Electrons are delocalized across more than two atoms in the molecule.
  • We show resonance structures with double-headed arrows (↔) between forms.
  • Resonance structures have the same arrangement of atoms, only electron positions differ.
  • Common examples include O3, NO3−, benzene (C6H6).
  • Delocalization reduces bond length differences and stabilizes the molecule.

Full Notes:

What Are Resonance Structures?

Resonance structures are different valid Lewis structures for the same molecule or ion, where:

These structures represent the same molecule, not different forms. The actual molecule is a resonance hybrid, meaning the bonding is a mixture of the possible forms.

Delocalization of Electrons

In resonance structures, π (pi) electrons are delocalized (spread over multiple atoms). This creates partial bonding across the delocalized region, resulting in equal bond lengths and increased stability.

Delocalization helps explain unusual bond lengths, stability of polyatomic ions, and the special nature of aromatic compounds like benzene.

Common examples include ozone, carboante ions and benzene.

Example Ozone (O3)

IB Chemistry diagram of ozone showing resonance structures with alternating double bonds and resonance hybrid.

Two valid structures: one double bond between the left O and middle O, or between the middle O and right O. Actual structure: a resonance hybrid where both O–O bonds are intermediate in length and strength.

Example Carbonate ion (CO32−)

IB Chemistry diagram showing resonance structures of carbonate ion with delocalized bonding.

Three resonance structures, each with one C=O and two C–O− bonds. In the hybrid all C–O bonds are equivalent in length and strength.

Example Benzene (C6H6)

IB Chemistry resonance structures of benzene showing alternating double bonds and delocalized ring.

Six carbon atoms form a ring with alternating single and double bonds. The resonance hybrid has delocalized electrons forming a stable aromatic ring.

Drawing and Recognising Resonance

When drawing resonance structures:

Structure 1.3 – Linked Course Question

Why Are Oxygen and Ozone Dissociated by Different Wavelengths of Light?

Oxygen (O2) and ozone (O3) absorb different wavelengths of light due to differences in their bonding and electron delocalisation.

IB Chemistry diagram comparing the bonds in oxygen and ozone, showing delocalisation effects on bond strength and light absorption.

The oxygen-oxygen bond in O2 is a true double bond, making it shorter and stronger than the bonds in ozone.

In ozone (O3), bonding is delocalised because of resonance. Instead of one single and one double bond, ozone forms a resonance hybrid where both oxygen-oxygen bonds are identical, with bond order between a single and a double bond.

As a result, ozone has no pure O=O double bond, but two longer, weaker bonds than in O2. This makes ozone easier to break, so it absorbs lower-energy (longer wavelength) UV light than oxygen.

In short, the greater delocalisation in ozone reduces the bond strength, explaining why it absorbs different light compared to O2.

Summary