Bonding (A-level only) — AQA A-Level Chemistry
Test yourself on Bonding (A-level only) with AQA A-Level practice questions.
A week of Premium from your first lesson · Then the free plan continues · No card, no charge
Bonding (A-level only) explained
Benzene, C₆H₆, is a flat regular hexagon with all bond angles 120°.
Read the full explanation
Each carbon uses three of its four outer electrons in σ bonds, two to neighbouring carbons and one to a hydrogen. The fourth electron stays in a p orbital that points perpendicular to the plane of the ring. Those six p orbitals overlap sideways all the way round, so the electrons are no longer tied to any one pair of atoms but are delocalised in a ring of electron density above and below the plane. Because every carbon-carbon bond has an identical share of that density, all six are the same length, 0.139 nm, between a single bond at 0.154 nm and a double bond at 0.134 nm, which is exactly what X-ray studies measure.
Delocalisation of p electrons makes benzene more stable than the theoretical molecule cyclohexa-1,3,5-triene.
Delocalisation lowers energy: spreading six π electrons over six carbon atoms puts them, on average, further from repulsion with each other and closer to more nuclei than three isolated double bonds would. Hydrogenation enthalpies measure the difference, because the product is the same in both cases. Adding hydrogen to one C=C in cyclohexene releases 120 kJ mol⁻¹, so the theoretical cyclohexa-1,3,5-triene, with three isolated double bonds, should release three times that, 360 kJ mol⁻¹. Benzene in fact releases only 208 kJ mol⁻¹. Less energy is given out because benzene started at a lower energy: it is about 152 kJ mol⁻¹ more stable than the model predicts, and that gap is the delocalisation energy.
Students should be able to: use thermochemical evidence from enthalpies of hydrogenation to account for this extra stability explain why substitution reactions occur in preference to addition reactions.
The thermochemical argument and the reactivity argument are the same argument. Hydrogenation data show benzene sits roughly 150 kJ mol⁻¹ below where three isolated double bonds would put it, so the delocalised ring is a stable arrangement that a reaction pays to destroy. An addition reaction uses up two of the ring electrons to form a new σ bond, which breaks the delocalisation and gives a product higher in energy than the starting material. A substitution replaces a hydrogen atom, so the ring is regenerated at the end of the mechanism and the stability is kept. That is why benzene does not decolourise bromine water, while an alkene does: the ring cannot polarise a bromine molecule on its own.
Your focus
- Describe the shape and bond angles of benzene and state the evidence that all six carbon-carbon bonds are identical.
- Explain how the p orbitals on the six carbon atoms give rise to a delocalised ring of electrons.
- Compare the carbon-carbon bond length in benzene with those in ethane and ethene, and account for the value.
Show all 9 objectives
- Calculate the delocalisation energy of benzene from the enthalpies of hydrogenation of benzene and cyclohexene.
- Explain why the theoretical value is three times the cyclohexene figure and why the same product must be used.
- Draw and label an energy level diagram showing benzene, cyclohexa-1,3,5-triene and cyclohexane.
- Decide from two enthalpies of hydrogenation which compound is the more stable, explaining why the less exothermic value belongs to the more stable compound.
- State what an addition reaction does to the delocalised ring and what a substitution restores, then use the delocalisation energy of about 150 kJ mol⁻¹ to say which product is the more stable.
- Predict the result of adding bromine water to benzene and to cyclohexene, and account for the difference.
Bonding (A-level only) exam tips
Marking Points
- one mark for the planar regular hexagon with bond angles of 120°
- one mark for each carbon contributing one electron from a p orbital perpendicular to the ring
- one mark for sideways overlap of those orbitals giving delocalised electrons above and below the plane
- one mark for all six carbon-carbon bonds being identical and intermediate in length between a single and a double bond
- one mark for stating that the π electrons are delocalised over all six carbon atoms
- one mark for this lowering the energy of benzene relative to the Kekulé model
- one mark for the predicted value calculated as three times the enthalpy of hydrogenation of cyclohexene
- one mark for the difference between the predicted and measured values identified as the extra stability
- one mark for the predicted enthalpy of hydrogenation calculated as three times the cyclohexene value
- one mark for comparing that prediction with the measured value and naming the difference as delocalisation energy
- one mark for stating that addition would break the delocalised ring and lose that stability
- one mark for substitution restoring the delocalised ring so the product remains low in energy
Examiner Tips
- 💡Use the circle-in-a-hexagon symbol whenever you are asked about bonding; save Kekulé structures for mechanisms, where you need a specific C-H bond.
- 💡Quote the bond length evidence if the question says 'explain, using evidence'; a number scores where a description may not.
- 💡Sketch the energy level diagram: theoretical triene at the top, benzene below it, cyclohexane at the bottom, then label the gap.
- 💡Watch the signs. Enthalpies of hydrogenation are negative, so 'less negative' means 'started lower' means 'more stable'.
- 💡Answer 'why substitution' in terms of the product's energy, not the reagent's strength.
- 💡The bromine water test is the quickest way to show a ring is delocalised rather than a set of alkene bonds; quote the observation as well as the conclusion.
Common Mistakes
- drawing alternating single and double bonds and then claiming all the bonds are the same length
- saying the delocalised electrons lie in the plane of the ring rather than above and below it
- stating that the delocalised electrons come from the C-H bonds
- calling benzene an alkene because it is shown with double bonds in the Kekulé structure
- arguing that benzene is less stable because 208 is a smaller number than 360
- multiplying the enthalpy of hydrogenation of benzene by three instead of that of cyclohexene
- comparing benzene with cyclohexane rather than with the theoretical triene
- calling the delocalisation energy the enthalpy of hydrogenation of benzene
- explaining the preference for substitution by saying only that benzene is unreactive
- claiming benzene cannot undergo addition at all, when it does with hydrogen over nickel and with chlorine in ultraviolet light
- using enthalpies of combustion in place of hydrogenation without saying the products are the same
- saying benzene decolourises bromine water because it contains double bonds