Hydrocarbons: Alkynes and Aromatic Hydrocarbons

Alkynes carry at least one carbon-carbon triple bond, general formula CnH2n-2, fewer hydrogens still than either an alkane or an alkene of the same carbon count, since a triple bond ties up even more of carbon's bonding capacity in carbon-carbon connections alone. Ethyne, better known as acetylene, is the first and simplest member, familiar from oxyacetylene welding torches, where it burns with oxygen at a temperature hot enough to cut through metal. Structurally, ethyne pushes the alkene story one step further still: each carbon is sp hybridised this time, not sp2, forming one sigma bond to the other carbon by head-on overlap and one sigma bond to a single hydrogen, in a straight line, H-C-C-H bond angles of exactly 180 degrees. That leaves two unhybridised p orbitals on each carbon, perpendicular to each other and to the sigma framework, and these overlap sideways in two separate directions at once, forming two distinct pi bonds rather than the alkene's one. The result is the strongest and shortest carbon-carbon bond of the three hydrocarbon families, 120 pm and roughly 823 kJ/mol, wrapped in a cylindrically symmetric electron cloud all the way around the internuclear axis, exactly the geometry that makes ethyne a genuinely linear molecule.

Alkyne names replace the alkane suffix with -yne, and position isomerism reappears exactly as it did with alkenes: but-1-yne and but-2-yne share the formula C4H6, differing only in where the triple bond sits. But alkynes carry one property alkenes and alkanes flatly do not: a hydrogen attached directly to a triply-bonded carbon is measurably acidic, reacting with sodium metal or with sodamide to release hydrogen gas and form a sodium acetylide. The reason traces straight back to hybridisation. An sp hybrid orbital has the highest possible s-character of the three, fifty percent, and s orbitals sit closer to the nucleus than p orbitals do; a carbon atom with that much s-character in its bonding orbitals holds its shared electrons unusually tightly, behaving, in effect, more electronegative than an sp2 or sp3 carbon would. That extra pull on the C-H bond's shared electrons makes the hydrogen genuinely easier to release as a proton, exactly the definition of acidity, and neither ethene's sp2 carbon nor ethane's sp3 carbon holds its own C-H bond nearly as tightly, which is exactly why this specific reaction, with sodium or sodamide, cleanly distinguishes a terminal alkyne from an alkene or an alkane using nothing more than a simple test tube.

Industrially, ethyne comes from calcium carbide: limestone heated to quicklime, quicklime heated with coke to calcium carbide, and calcium carbide reacted with water to release ethyne gas directly. In the laboratory, a vicinal dihalide, treated with alcoholic potassium hydroxide, can eliminate two separate molecules of hydrogen halide in succession, first to an alkenyl halide, then, with sodamide, all the way to the alkyne. Once formed, alkynes add across their triple bond in stages, much like alkenes but with room for a second addition: hydrogen, using platinum, palladium or nickel, adds twice in succession, alkyne to alkene to alkane; halogens and hydrogen halides add the same way, unsymmetrical alkynes again following Markovnikov's rule, this time the reaction pausing at a gem-dihalide, two halogens on the very same carbon, once one equivalent has added. Water itself does not normally touch a triple bond at all, but warmed with mercuric sulphate and dilute sulphuric acid, one molecule adds across it after all, landing not as a simple alcohol but rearranging immediately into a carbonyl compound instead, since the initial addition product is never stable enough to persist.

Ethyne polymerises in two genuinely different ways, and only one of them stays linear. Under the right conditions, ethyne molecules link end to end into polyacetylene, a long chain of alternating single and double bonds; kept thin as a film, this specific polymer conducts electricity, lighter and cheaper than a metal wire doing the same job. Passed through a red-hot iron tube instead, three ethyne molecules take a completely different path, closing into a ring rather than staying open: cyclic polymerisation, and the product is benzene itself, the single most direct laboratory route from an aliphatic starting material into the aromatic world explored for the rest of this chapter. It is a strange, almost poetic transition to notice: the same small, acidic, triple-bonded molecule that reacts so predictably on its own becomes, three copies at a time, the seed of an entirely different, far less predictable kind of chemistry.

Benzene, C6H6, was isolated by Michael Faraday in 1825, and its formula alone was already a puzzle: that few hydrogens for six carbons implied serious unsaturation, yet benzene stubbornly refused to behave like any other unsaturated compound anyone had studied. In 1865, August Kekule proposed a cyclic structure, six carbons in a ring, alternating single and double bonds, a hydrogen on each, and famously later described the idea arriving to him half-asleep, watching a snake seize its own tail in a dream. The structure looked reasonable, and it explained why benzene formed only one possible monosubstituted product, since all six carbons and hydrogens were now equivalent by symmetry. But it made a testable, and ultimately wrong, prediction: with fixed alternating bonds, two different 1,2-disubstituted products should exist, depending on whether the substituents landed on a double-bonded pair of carbons or a single-bonded one, and experimentally, only one ortho-disubstituted product was ever found. Kekule patched this by proposing the double bonds oscillate rapidly back and forth between two positions, but even that patch could not explain benzene's other headline anomaly: its unusual stability, and its strong preference for substitution reactions over the addition reactions every other unsaturated compound in this chapter has readily undergone.

The real resolution came from resonance theory, already established as a concept earlier in this thread. Benzene is not oscillating between Kekule's two structures, and it is not any single fixed-bond structure either; it is a genuine resonance hybrid of both, with six pi electrons delocalised evenly around all six carbons rather than paired up into three separate, localised double bonds. Orbital theory makes this concrete: all six carbons are sp2 hybridised, forming a flat hexagonal ring of sigma bonds, and each carbon's one remaining, unhybridised p orbital, all six of them parallel and perpendicular to the ring's plane, overlaps sideways with its neighbours on both sides at once, equally, not preferentially with just one. The resulting pi electron cloud forms two continuous, doughnut-shaped rings of electron density, one above and one below the plane of the carbons, and X-ray diffraction confirms exactly what this picture predicts: all six carbon-carbon bonds in benzene measure precisely the same length, 139 pm, sitting neatly between a pure single bond (154 pm) and a pure double bond (133 pm). No fixed-bond Kekule structure could ever produce that result; only genuine, complete delocalisation can, and that same delocalisation is exactly what makes benzene unusually stable and reluctant to give up its aromatic ring system through an ordinary addition reaction.

"Aromatic" originally just meant benzene-like, but the concept has since grown into a precise, checkable rule that applies to any ring system, whether or not it actually contains a benzene ring. A compound qualifies as aromatic when it satisfies three specific conditions together: it must be planar, its ring must have complete, uninterrupted delocalisation of pi electrons all the way around, and it must contain exactly (4n+2) pi electrons, where n is any whole number, zero, one, two and onward, a condition known as Huckel's rule. Benzene fits with n = 1, six pi electrons exactly matching (4 x 1) + 2. This single rule is precisely why some rings that look superficially similar to benzene fail to qualify as aromatic at all, missing planarity, missing full delocalisation, or simply having the wrong pi electron count, and why other rings that do not contain a benzene ring whatsoever can still genuinely earn the aromatic label, provided they meet all three conditions on their own terms.

Benzene is isolated commercially from coal tar, but three laboratory routes exist too: the cyclic polymerisation of ethyne already covered, decarboxylation of an aromatic acid's sodium salt with soda lime, and reduction of phenol vapours passed over heated zinc dust. Physically, aromatic hydrocarbons are non-polar, often pleasantly aromatic-smelling liquids or solids, immiscible with water but freely miscible with organic solvents, and they burn with a noticeably sooty flame, a giveaway of their high carbon-to-hydrogen ratio. Chemically, benzene's whole personality flows from the stability just established: because breaking its delocalised aromatic system costs real energy, benzene strongly prefers electrophilic substitution, trading one hydrogen for a new group while the aromatic ring itself survives intact, over addition, which would sacrifice that stability altogether. Addition does happen, but only under genuinely vigorous, forcing conditions: hydrogenation to cyclohexane needs a nickel catalyst at high temperature and pressure, and chlorination under ultraviolet light adds three whole chlorine molecules to produce benzene hexachloride, both reactions requiring far harsher conditions than an ordinary alkene ever demands for the exact same kind of addition.

Nitration, halogenation, sulphonation, and the two Friedel-Crafts reactions, alkylation and acylation, all look like different transformations, but underneath, every single one runs through the identical three-step mechanism. First, the electrophile itself has to be generated: concentrated nitric and sulphuric acid together generate the nitronium ion, NO2+; a Lewis acid like anhydrous aluminium chloride activates a halogen, an alkyl halide, or an acyl halide into Cl+, an alkyl cation, or an acylium ion respectively. Second, that electrophile attacks the ring's delocalised pi system, forming a carbocation intermediate called an arenium ion, or sigma complex, in which one ring carbon has temporarily become sp3 hybridised and the aromatic delocalisation is, for this one intermediate step only, genuinely broken. Third, and this is what restores everything: that same sp3 carbon loses a proton, snapped up by whatever conjugate base is floating nearby, and the ring's full aromatic delocalisation, and its stability, snap back into place. Different electrophiles and different Lewis acid activators produce nitrobenzene, halobenzenes, benzenesulphonic acid, alkylbenzenes and acylbenzenes respectively, five distinct products from one identical three-step machine.

Attach a second substituent to a benzene ring already carrying one, and the three theoretically possible positions, ortho, meta, para, are not chosen at random; the identity of the group already present decides almost everything. Electron-donating groups, -OH, -NH2, and alkyl groups among them, push electron density into the ring through resonance (or, for alkyl groups, through the inductive effect), and that donated density concentrates specifically at the ortho and para positions, exactly where the resonance structures place the extra electron pair, so incoming electrophiles are drawn there preferentially: these groups are ortho/para directors, and since they leave the ring more electron-rich than plain benzene, they also activate it, speeding up substitution overall. Halogens are the one genuine exception worth remembering: their strong inductive pull withdraws electron density and technically deactivates the ring, yet their lone pairs can still donate by resonance specifically into the ortho and para positions, so halogens end up simultaneously deactivating and ortho/para-directing at once, an apparent contradiction that dissolves once inductive and resonance effects are counted separately rather than lumped together. Electron-withdrawing groups, -NO2, -CN, -COOH and -SO3H among them, pull density out of the ring instead, leaving the meta position comparatively the most electron-rich of the three by process of elimination, and incoming electrophiles land there: these groups are meta directors, and since they leave the whole ring more electron-poor than plain benzene, they also deactivate it, slowing substitution down. And it is worth closing on a genuinely serious note: benzene and, more so, polynuclear aromatic hydrocarbons, several fused benzene rings joined together, forming as a by-product of incomplete combustion in tobacco smoke, coal and petroleum, are toxic, and a number of them are established carcinogens, damaging DNA once absorbed into the body, a sober, real-world consequence of the very same delocalised ring chemistry this chapter has spent so long admiring.