HookThe molecule chemists drew wrong for sixty years
In 1825 Michael Faraday distilled a clear, sweet-smelling liquid from the oily residue left in London's gas-lamp cylinders and called it 'bicarburet of hydrogen'. Its formula turned out to be \(\text{C}_6\text{H}_6\) — six carbons carrying only six hydrogens. A molecule that hydrogen-poor should be crammed with double bonds and at least as reactive as an alkene; it should strip the colour out of bromine water on contact. It does no such thing. Benzene sat in chemistry's in-tray for forty years, refusing to match any structure anyone drew for it, until August Kekulé proposed a hexagon of carbons with alternating single and double bonds — a structure he later claimed came to him in a daydream of a snake seizing its own tail.
Kekulé's hexagon was a triumph, and it is also wrong in a way that matters. Three separate lines of evidence — bond lengths from X-ray crystallography, an enthalpy of hydrogenation that comes out 152 kJ mol⁻¹ short, and benzene's flat refusal to do addition chemistry — force a better model: six electrons delocalised into rings of electron density above and below the carbon plane. That model earns its keep. It explains why benzene's characteristic reaction is electrophilic substitution — swap a hydrogen, keep the ring — and it sets up the two mechanisms AQA examines in detail: nitration, the gateway to dyes, drugs and explosives, and Friedel-Crafts acylation, the carbon-carbon bond-maker that modern ibuprofen manufacture still opens with.
ModelThree experiments, one verdict: the ring is delocalised
Take Kekulé's structure seriously and it makes testable predictions. Alternating bonds mean two different bond lengths around the ring — a normal C–C single bond is 0.154 nm and a C=C double bond is 0.134 nm — so the hexagon should be slightly lopsided. Three genuine double bonds should also hydrogenate like three alkenes and react with bromine like one. Every one of those predictions fails.
The geometric evidence came first. In 1929 the crystallographer Kathleen Lonsdale used X-ray diffraction to show that the benzene ring is a planar, perfectly regular hexagon: every carbon-carbon bond is identical, at 0.140 nm — squarely between the single-bond and double-bond lengths. There are no long and short bonds to alternate. The chemical evidence agrees: benzene does not decolourise bromine water at room temperature, a test any alkene passes instantly.
The delocalised model explains all of it. Each carbon uses three of its four outer-shell electrons for normal covalent bonds — one to a hydrogen, two to its neighbouring carbons — which builds the flat hexagonal frame with 120° bond angles. The fourth electron sits in a p orbital perpendicular to the plane, and the six p orbitals overlap side-on to form a continuous ring of electron density above and below the carbons. Those six electrons belong to no particular pair of atoms: they are delocalised, which is why every bond is the same length and why the ring is unusually stable.
Cyclohexene contains one C=C bond and its enthalpy of hydrogenation is \(-120\ \text{kJ mol}^{-1}\). Predict the enthalpy of hydrogenation of Kekulé's 'cyclohexa-1,3,5-triene', compare it with benzene's measured value of \(-208\ \text{kJ mol}^{-1}\), and interpret the difference.
Step 1 — the prediction. Three isolated double bonds should release three times the energy: \(3 \times (-120) = -360\ \text{kJ mol}^{-1}\).
Step 2 — the comparison. Real benzene releases only \(-208\ \text{kJ mol}^{-1}\), which is \(152\ \text{kJ mol}^{-1}\) less than the prediction.
Step 3 — the interpretation. Benzene starts 152 kJ mol⁻¹ lower in energy than the hypothetical triene: it is more stable than its own drawing. That gap is the delocalisation stability, and it prices every reaction benzene is offered. Addition would break the delocalised ring and forfeit the 152 kJ mol⁻¹; substitution restores the ring and keeps it. Quote all three numbers — prediction, measurement, difference — because each carries its own mark.
MechanismElectrophilic substitution — swap a hydrogen, keep the ring
The delocalised ring is a band of high electron density, so the reagents benzene attracts are electrophiles — electron-pair acceptors. But there is a catch. Because those six electrons are spread thinly across six carbons, the ring cannot polarise an approaching molecule the way an alkene's concentrated C=C can. Bromine that reacts with ethene on contact leaves benzene untouched. Benzene demands pre-made, aggressive electrophiles, and both of the mechanisms AQA sets begin with a catalyst manufacturing one: \(\text{NO}_2^+\) for nitration, \(\text{RCO}^+\) for acylation.
The mechanism itself is two moves. First, a curly arrow from the delocalised ring to the electrophile: a pair of ring electrons forms the new bond, and the intermediate is drawn with a horseshoe of partial delocalisation spread over five carbons, a positive charge inside it, and the attacked carbon holding both its hydrogen and the incoming group. Second, an arrow from that carbon's C–H bond back into the ring: the hydrogen leaves as \(\text{H}^+\), the electrons rejoin the π system, and full delocalisation — with its 152 kJ mol⁻¹ of stability — is restored.
That restoration is the whole logic of aromatic chemistry, and it is a favourite short-answer contrast: an alkene undergoes electrophilic addition because it has no delocalisation to protect; benzene undergoes electrophilic substitution because losing a hydrogen is cheap and losing the ring is not.
MechanismNitration — one small group, two big industries
Nitration replaces a ring hydrogen with \(-\text{NO}_2\). The reagents are concentrated nitric acid and concentrated sulfuric acid — the 'nitrating mixture' — warmed with benzene to about 50 °C. The electrophile is the nitronium ion, \(\text{NO}_2^+\), and generating it is the sulfuric acid's job: \(\text{HNO}_3 + \text{H}_2\text{SO}_4 \rightarrow \text{NO}_2^+ + \text{HSO}_4^- + \text{H}_2\text{O}\). Notice the role reversal — the stronger sulfuric acid forces nitric acid to accept a proton and then fall apart, so nitric acid spends this reaction acting as a base.
The \(\text{NO}_2^+\) ion then attacks the ring exactly as the general mechanism says: arrow from ring to electrophile, positively charged horseshoe intermediate, arrow from the C–H bond back into the ring. The expelled \(\text{H}^+\) recombines with \(\text{HSO}_4^-\) to regenerate \(\text{H}_2\text{SO}_4\) — which is what makes the sulfuric acid a genuine catalyst. Overall: \(\text{C}_6\text{H}_6 + \text{HNO}_3 \rightarrow \text{C}_6\text{H}_5\text{NO}_2 + \text{H}_2\text{O}\). Temperature is a control variable, not a detail: hold the mixture near 50 °C and you get mononitration; let it run hotter and a second \(\text{NO}_2\) group goes on.
Why anyone bothers is the exam's favourite context. Reduce nitrobenzene and you get phenylamine (section 3.3.11), the ancestor of synthetic dyes and a workhorse of pharmaceutical synthesis. Nitrate methylbenzene three times and you get 2,4,6-trinitromethylbenzene — TNT. One substitution reaction, feeding both the dye works and the demolition trade.
CaseFriedel-Crafts acylation — buying a carbon-carbon bond
To build anything ambitious on a benzene ring — a drug, a dye, a fragrance — you eventually have to attach carbon to carbon, and no everyday carbon electrophile is powerful enough. The fix, published by Charles Friedel and James Crafts in 1877, is to let anhydrous aluminium chloride manufacture the electrophile from an acyl chloride: \(\text{CH}_3\text{COCl} + \text{AlCl}_3 \rightarrow \text{CH}_3\text{CO}^+ + \text{AlCl}_4^-\). The \(\text{AlCl}_3\) accepts a chloride ion, leaving the acylium ion \(\text{CH}_3\text{CO}^+\) — a carbon electrophile aggressive enough for benzene's diffuse ring.
The acylium ion attacks the ring through the standard two-arrow mechanism, and the expelled proton closes the loop: \(\text{H}^+ + \text{AlCl}_4^- \rightarrow \text{AlCl}_3 + \text{HCl}\), regenerating the catalyst. Overall, with ethanoyl chloride: \(\text{C}_6\text{H}_6 + \text{CH}_3\text{COCl} \rightarrow \text{C}_6\text{H}_5\text{COCH}_3 + \text{HCl}\) — the product is the aromatic ketone phenylethanone. Conditions matter and are asked directly: heat under reflux, with anhydrous \(\text{AlCl}_3\) and scrupulously dry apparatus, because water destroys the catalyst. A useful bonus: the C=O group withdraws electron density from the ring, so the product is less reactive than benzene and the reaction stops cleanly at one substitution.
This is not museum chemistry. The 1960s Boots synthesis of ibuprofen took six steps with an atom economy around 40%; the BHC route that replaced it in 1992 takes three catalytic steps — the first of them a Friedel-Crafts acylation of isobutylbenzene — with an atom economy of roughly 77%, and it collected a US Presidential Green Chemistry Challenge Award in 1997. Every pack of ibuprofen on a supermarket shelf starts with the reaction you are drawing curly arrows for.
Benzene reacts with propanoyl chloride, \(\text{CH}_3\text{CH}_2\text{COCl}\), in the presence of anhydrous \(\text{AlCl}_3\). Write the equation for the formation of the electrophile, the overall equation, and name the organic product.
Step 1 — the electrophile: \(\text{CH}_3\text{CH}_2\text{COCl} + \text{AlCl}_3 \rightarrow \text{CH}_3\text{CH}_2\text{CO}^+ + \text{AlCl}_4^-\).
Step 2 — the substitution: \(\text{C}_6\text{H}_6 + \text{CH}_3\text{CH}_2\text{COCl} \rightarrow \text{C}_6\text{H}_5\text{COCH}_2\text{CH}_3 + \text{HCl}\).
Step 3 — the name: the acyl group arrives intact, carbonyl carbon first, so the product is the aromatic ketone 1-phenylpropan-1-one.
Step 4 — close the loop: \(\text{H}^+ + \text{AlCl}_4^- \rightarrow \text{AlCl}_3 + \text{HCl}\), so the aluminium chloride is regenerated — call it a catalyst and you collect that mark. The self-check examiners love: count the chlorine. The Cl that left the acyl chloride leaves the flask as HCl, and the \(\text{AlCl}_3\) finishes exactly as it started.
VocabularyKey terms the mark scheme pays for
TrapsMisconceptions that cost marks
ExamWhat examiners want
Evidence questions are number questions. When asked to justify the delocalised model, quote the data: predicted bond lengths 0.154 nm and 0.134 nm against the measured uniform 0.140 nm; predicted hydrogenation enthalpy \(-360\ \text{kJ mol}^{-1}\) against the measured \(-208\ \text{kJ mol}^{-1}\), a 152 kJ mol⁻¹ stabilisation; and no reaction with bromine water at room temperature. Candidates who reason verbally without figures cap themselves at half marks.
Mechanism drawing is where AQA is pickiest. The first curly arrow starts on the delocalised ring (the circle) and points to the electrophile. The intermediate's horseshoe must stop short of the carbon bonded to the electrophile, with the positive charge inside the horseshoe. The second arrow starts on the C–H bond — never on the H atom — and returns to the ring. Learn the electrophile-generation equations cold (\(\text{HNO}_3 + \text{H}_2\text{SO}_4 \rightarrow \text{NO}_2^+ + \text{HSO}_4^- + \text{H}_2\text{O}\); \(\text{RCOCl} + \text{AlCl}_3 \rightarrow \text{RCO}^+ + \text{AlCl}_4^-\)) and the catalyst-regeneration equations: each is routinely a standalone mark.
Conditions come with reasons attached. About 50 °C for nitration because higher temperatures give further substitution; anhydrous AlCl₃ and dry apparatus, heated under reflux, for acylation because water destroys the catalyst. 'State and explain' means the explanation carries the second mark.
For the classic extended response — why substitution rather than addition — argue in energy terms: addition breaks the delocalised system and costs 152 kJ mol⁻¹; substitution passes through a temporary loss but restores the ring in the product. And always name the mechanism in full: 'electrophilic substitution', both words, every time.