In simple terms
A friendly intro before the formal notes — no formulas yet.
Benzene's Stable Secret
Benzene has a unique flat ring structure with a cloud of shared electrons, making it unusually stable. This stability dictates its chemical reactions, favouring substitution over addition.
Imagine six friends holding hands in a circle to form a strong framework (the σ bonds). Now, instead of passing a single ball (a localised π bond) between just two people at a time, they are all juggling a shared cloud of six balls together, above and below their hands (the delocalised π system). It's much harder to break this cooperative group apart (high stability) than to just intercept one ball from a simple pair.
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Benzene C₆H₆ — planar hexagon, 120° bond angles.
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σ bonds: C–C and C–H framework; π cloud above/below ring.
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Electrophilic substitution preserves aromaticity.
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Resonance structures — average bond order 1.5.
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Full topic notes
Formal explanation with the rigour you need for the exam.
The Shape and Structure of Benzene
Experimental evidence, primarily from X-ray diffraction studies, reveals that benzene is a perfectly flat, or planar, molecule. The six carbon atoms are arranged in a regular hexagon. All the carbon-carbon bond lengths are identical at 0.139 nm, and all bond angles, both C-C-C and H-C-C, are 120°. This uniformity is a key feature that the earlier Kekulé model of alternating single and double bonds could not explain.
Formula: C₆H₆
Shape: Planar, regular hexagon
Bond angles: 120° throughout
C-C bond lengths: All are 0.139 nm (intermediate between single and double bonds)
Orbital Overlap: The σ and π Framework
To achieve a trigonal planar geometry with 120° bond angles, each carbon atom in the benzene ring is sp² hybridised. Each carbon atom forms three σ (sigma) bonds. Two of these are formed by the head-on overlap of its sp² hybrid orbitals with those of its two neighbouring carbon atoms. The third σ bond is formed by the head-on overlap of its remaining sp² hybrid orbital with the 1s orbital of a hydrogen atom. This network of σ bonds creates a strong, planar C₆H₆ framework.
After forming the σ framework, each carbon atom has one unhybridised 2p orbital remaining. These p-orbitals are oriented perpendicular (at 90°) to the plane of the ring. The six p-orbitals are close enough to overlap sideways with their neighbours on both sides. This continuous sideways overlap merges the six individual p-orbitals into two ring-shaped electron clouds, one above and one below the plane of the σ framework. The six electrons (one from each carbon's p-orbital) are no longer associated with a single carbon atom but are shared amongst all six. This is the delocalised π system.
When asked to describe the bonding in benzene, always discuss the formation of the σ framework first, then explain the formation of the delocalised π system from the remaining p-orbitals. Marks are often allocated for both parts separately.
Evidence for the Delocalised Model
The delocalised model of benzene is not just a theory; it is supported by robust experimental evidence. The three main lines of evidence are thermochemical data, bond length measurements, and chemical reactivity.
The third piece of evidence is benzene's characteristic reactivity. Alkenes, with their localised π bonds, readily undergo electrophilic addition reactions. For example, cyclohexene will decolourise bromine water at room temperature. Benzene, however, does not react with bromine water. It only reacts with bromine in the presence of a halogen carrier catalyst (e.g., FeBr₃) and undergoes electrophilic substitution, not addition. This is because an addition reaction would permanently break the delocalised system, losing the large delocalisation energy. A substitution reaction allows the stable aromatic ring to be reformed.
When comparing benzene's stability, always be specific. Don't just say 'benzene is stable'. Instead, say 'benzene is more stable than expected for a structure with three C=C bonds due to delocalisation energy'. This comparative and explanatory detail is what examiners look for.
Worked examples
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The enthalpy of hydrogenation of cyclohexene (C₆H₁₀), which has one C=C bond, is -120 kJ mol⁻¹. The experimental enthalpy of hydrogenation of benzene (C₆H₆) to cyclohexane (C₆H₁₂) is -208 kJ mol⁻¹. (a) Predict the enthalpy of hydrogenation for the hypothetical cyclohexa-1,3,5-triene. (b) Calculate the delocalisation energy of benzene. (c) Explain what this value indicates about the stability of benzene.
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(a) The hypothetical cyclohexa-1,3,5-triene has three localised C=C bonds. Predicted enthalpy change = 3 × (enthalpy change for one C=C bond) = 3 × (-120 kJ mol⁻¹) = -360 kJ mol⁻¹.
Explain why the C–O bond in phenol (C₆H₅OH) is shorter than the C–O bond in cyclohexanol (C₆H₁₁OH), with reference to orbital overlap.
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In phenol, the oxygen atom is bonded to an sp² hybridised carbon of the benzene ring. One of the lone pairs on the oxygen atom is in a p-orbital that is parallel to the p-orbitals of the ring.
How it all connects
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Glossary
Key terms for this topic — skim now; the Check step will test them.
- Benzene
Benzene is a planar, regular hexagonal molecule.
- Delocalisation energy
The additional stability of a molecule due to electron delocalisation. It is the difference in energy between the actual molecule and a hypothetical structure with localised bonds (e.g., benzene vs. cyclohexa-1,3,5-triene).
Quick check
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Revision flashcards
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Key takeaways
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Formula: C₆H₆
Shape: Planar, regular hexagon
Bond angles: 120° throughout
C-C bond lengths: All are 0.139 nm (intermediate between single and double bonds)
Practice — then mark it
The whole point: a real Cambridge question, marked mark-by-mark.
Describe the shape of delocalised benzene. Include the geometry of each carbon, the C-C-H bond angle and the type of bond(s) between the carbon atoms and between the carbon and hydrogen atoms.
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Checkpoint
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Discuss Shapes of aromatic organic molecules; σ and π bonds
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