Ch 3 · Shapes of aromatic organic molecules; σ and π bonds · AS & A Level Chemistry 9701

Topic 3 of 4 in An introduction to A Level organic chemistry — AS & A Level Chemistry 9701 — 48 lessons.

Introduction

AS Level met sp2 hybridisation in the planar C=C bond of ethene. Benzene extends the same idea around a six-membered ring, but with a crucial difference: instead of three isolated double bonds, its six π electrons are shared, or delocalised, over the whole ring. This single structural feature explains benzene's shape, its unusual carbon–carbon bond lengths, and its resistance to the addition reactions expected of a molecule with three C=C bonds.

Key ideas — sp2 hybridisation and the shape of benzene

Each carbon atom in benzene is sp2 hybridised: its 2s orbital and two of its three 2p orbitals mix to form three equivalent sp2 hybrid orbitals lying in a single plane at 120^° to one another. Two of these orbitals form sigma bonds to the neighbouring ring carbons, and the third forms a sigma bond to a hydrogen atom, giving benzene a flat, regular hexagonal, planar shape with every C-C-C and H-C-C bond angle equal to 120^°. The remaining, unhybridised 2p orbital on each carbon points perpendicular to the plane of the ring, one electron in each.

🧑‍🏫 Teacher's explanation

Imagine a group of six friends standing in a circle, each holding two hands with their neighbors to form a ring. If everyone stands perfectly flat on the ground, they form a hexagon. In benzene, the carbon atoms do exactly this using sp^2 hybrid orbitals. These orbitals act like the 'legs' that hold the molecule flat and steady in a single plane.

📖 Definitions

🔍 Worked example — Easy

Calculate the total number of σ (sigma) bonds in a benzene molecule (C_6H_6).

🔍 Worked example — Standard

Compare the bond length of a C-C single bond (approx 1.54 AA) and a C=C double bond (approx 1.34 AA). Predict the approximate bond length in benzene.

🔍 Worked example — Exam-style

Explain why benzene is more stable than the hypothetical molecule cyclohexatriene.

⚠️ Common errors

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🎯 Exam question

(a) Describe the bonding in benzene in terms of σ and π bonds. [3] (b) Explain why the C-C bond length in benzene is shorter than a typical C-C single bond but longer than a typical C=C double bond. [2] [5 marks]

Key ideas — The delocalised pi system

Because every carbon's unhybridised p orbital is parallel to its neighbours', they overlap sideways all the way around the ring rather than pairing up into three separate, localised π bonds. This continuous sideways overlap spreads the six π electrons into a single delocalised molecular orbital system, forming two symmetrical rings of electron density, one above and one below the plane of the carbon atoms.

🧑‍🏫 Teacher's explanation

Imagine a group of six people standing in a circle, each holding a single piece of string. In a 'normal' molecule like cyclohexene, people only hold strings with the person directly next to them, creating fixed pairs. In benzene, however, everyone holds hands with both neighbors simultaneously, and the string is passed around the whole circle continuously. This 'shared' connection is what we call delocalisation.

📖 Definitions

🔍 Worked example — Easy

Compare the bond lengths in benzene to those in ethene (C=C) and ethane (C-C). If C-C is 1.54 \text{ \AA} and C=C is 1.34 \text{ \AA}, what is the approximate bond length in benzene?

🔍 Worked example — Standard

Explain why benzene is more stable than the hypothetical molecule cyclohexatriene.

🔍 Worked example — Exam-style

A benzene ring has a total π electron count. Calculate the number of π electrons in one benzene ring.

⚠️ Common errors

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🎯 Exam question

Key ideas — Delocalisation lowers the energy of benzene

If benzene really contained three localised C=C double bonds alternating with three single bonds (a hypothetical structure called cyclohexatriene), its enthalpy change of hydrogenation could be predicted from that of cyclohexene, which has one C=C bond and a hydrogenation enthalpy of -120\ mathrmkJ\ mol^-1. Three such double bonds, each hydrogenated independently, would be expected to release 3 × (-120) = -360\ mathrmkJ\ mol^-1. The experimentally measured enthalpy of hydrogenation of benzene is only -208\ mathrmkJ\ mol^-1, considerably less exothermic. Since hydrogenation releases less energy than predicted, benzene must start at a lower energy level than the hypothetical cyclohexatriene; the difference, -208 - (-360) = +152\ mathrmkJ\ mol^-1, is the delocalisation (resonance) energy, the extra stability gained by spreading the π electrons around the whole ring rather than confining them to three separate double bonds.

🧑‍🏫 Teacher's explanation

Imagine you have three separate, small, fenced-in playgrounds. If you put children (electrons) in each playground, they are stuck in those small areas. Now, imagine you tear down all the fences and create one giant, circular park. The children can run around much more freely. Because they have more space to move, they are 'happier' and have lower energy. In chemistry, when electrons are spread out over a larger area, we call this delocalisation, and it makes the molecule much more stable.

📖 Definitions

🔍 Worked example — Easy

Calculate the predicted enthalpy change of hydrogenation for a hypothetical molecule containing three localised C=C double bonds, if the enthalpy change for one C=C bond is -120 kJ mol^-1.

🔍 Worked example — Standard

The experimental enthalpy change of hydrogenation for benzene is -208 kJ mol^-1. Calculate the resonance energy (delocalisation energy) released when benzene forms from the hypothetical cyclohexatriene.

🔍 Worked example — Exam-style

A substituted benzene derivative has an experimental hydrogenation enthalpy of -250 kJ mol^-1. If the predicted enthalpy for the corresponding cyclohexatriene is -380 kJ mol^-1, calculate the resonance energy.

⚠️ Common errors

🧠 Quick check 1

🧠 Quick check 2

🎯 Exam question

(a) Calculate the enthalpy change for the hydrogenation of a hypothetical cyclohexatriene. [1] (b) Calculate the resonance energy of benzene. [2] (c) Explain why the experimental value for benzene is less exothermic than the value calculated in part (a). [2] [5 marks]

Key ideas — Bond lengths in benzene

X-ray diffraction shows that all six carbon–carbon bonds in benzene are identical in length, at 0.139\ nm. This is intermediate between the length of a typical carbon–carbon single bond (0.154\ nm, as in ethane) and a carbon–carbon double bond (0.134\ nm, as in ethene). A structure with three alternating single and double bonds would show two distinct bond lengths around the ring; the fact that every bond is the same, intermediate length confirms that the π electron density is spread evenly over all six bonds, each having partial double-bond character, rather than being localised as three discrete double bonds.

🧑‍🏫 Teacher's explanation

Imagine a group of six friends standing in a circle, holding hands. In a normal molecule like cyclohexatriene, they would be holding hands very tightly in some spots (double bonds) and loosely in others (single bonds). This would create a 'wobbly' ring with different distances between people. However, in benzene, the friends are all holding hands with the exact same amount of tension. This is because the electrons aren't stuck between just two people; they are flowing in a continuous loop around the whole circle.

Ch 3 · Shapes of aromatic organic molecules; σ and π bonds · AS & A Level Chemistry 9701

All chapters in An introduction to A Level organic chemistry