Haloalkanes and Haloarenes

A chlorine-containing antibiotic, chloramphenicol, treats typhoid. An iodine-containing hormone the body makes on its own, thyroxine, keeps metabolism running, its deficiency causing goitre. Chloroquine, a synthetic halogen compound, treats malaria; halothane, another, is used as a surgical anaesthetic. All of this real, consequential chemistry starts from the same small structural move: replacing a hydrogen atom on a hydrocarbon with a halogen atom instead. Do this to an aliphatic hydrocarbon and the result is a haloalkane, or alkyl halide; do it to an aromatic ring and the result is a haloarene, or aryl halide, and the difference between those two, a halogen sitting on an sp3 carbon versus an sp2 one, turns out to decide almost everything explored in this chapter, how reactive the compound is, which mechanism a reaction actually follows, and what happens to the molecule's own shape in space along the way. There is a serious cost side to this chemistry too, worth holding in view from the very first page: halogenated compounds resist breakdown by ordinary soil bacteria, and several of the most useful ones in this chapter turn out to persist in the environment for a genuinely long time after their job is done.

Alkyl halides sort further by exactly which carbon holds the halogen: primary, secondary or tertiary, depending on how many other carbons that specific carbon touches, precisely the same vocabulary already established for carbocations. Allylic halides place the halogen one carbon away from a double bond; benzylic halides place it one carbon away from an aromatic ring, both still on an sp3 carbon despite sitting right next to unsaturation. Vinylic halides and aryl halides, by contrast, put the halogen directly on an sp2 carbon, on a double bond itself or directly on the aromatic ring. Compounds carrying two halogens sort by position too: gem-dihalides carry both halogens on the very same carbon, vic-dihalides on two adjacent carbons instead. Every carbon-halogen bond is polarised, since every halogen is more electronegative than carbon, but the size of that polarisation, and the bond's length and strength, shift steadily down the halogen group: fluorine, smallest and most electronegative, forms the shortest bond, 139 pm, and iodine, largest, forms the longest, 214 pm, bond strength falling in step as the bond stretches.

Alcohols supply the most reliable route: their -OH group swaps directly for a halogen when treated with a concentrated halogen acid, a phosphorus halide, or thionyl chloride, and thionyl chloride is genuinely the preferred choice, since its only by-products, sulphur dioxide and hydrogen chloride, are both gases that simply escape, leaving the alkyl halide clean and easy to isolate. Alcohol reactivity toward a given acid follows the familiar order, tertiary fastest, primary slowest. Hydrocarbons supply two rougher routes: free radical halogenation of an alkane produces a genuinely messy mixture of isomers, hard to separate and rarely worth attempting when a cleaner method exists, while an alkene reacting with a hydrogen halide gives a single, predictable haloalkane, following Markovnikov's rule exactly as established earlier in this thread. Two further reactions solve a specific supply problem: iodides and fluorides are difficult to make directly and cleanly, so a chemist instead makes the chloride or bromide first, then swaps the halogen itself. The Finkelstein reaction converts a chloride or bromide to the corresponding iodide using sodium iodide in dry acetone, the reaction pulled forward as the by-product sodium chloride or bromide precipitates straight out of solution; the Swarts reaction achieves the equivalent fluorine swap using a metallic fluoride like silver fluoride or antimony trifluoride instead.

A halogen cannot simply swap into an aromatic ring the way it swaps onto an alcohol's carbon; the ring calls for a genuinely different kind of chemistry. Electrophilic substitution, already established for benzene itself, works directly here too: chlorine or bromine, activated by a Lewis acid catalyst like iron or iron(III) chloride, substitutes onto the ring, and since ortho and para isomers differ substantially in melting point, they separate from each other easily once formed. Iodination needs an oxidising agent alongside it, since the reaction is otherwise reversible, and fluorine is simply too reactive to control this way at all. Sandmeyer's reaction supplies a second, more roundabout route, starting from a primary aromatic amine rather than benzene itself: treated with sodium nitrite in cold aqueous acid, the amine converts into a diazonium salt, and that diazonium salt, mixed with cuprous chloride or cuprous bromide, swaps its diazonium group for the halogen directly; the same swap to iodine needs no copper catalyst at all, simply shaking the diazonium salt with potassium iodide.

Pure alkyl halides are colourless, though bromides and iodides slowly develop colour with light exposure, and many volatile ones carry a distinctly sweet smell. Their boiling points sit noticeably higher than a hydrocarbon of similar mass, since a polarised C-X bond adds real dipole-dipole attraction on top of the van der Waals forces a plain hydrocarbon relies on alone, and for the same alkyl group, boiling point climbs RF < RCl < RBr < RI, the larger, more polarisable halogen strengthening the van der Waals contribution further still, exactly the same branching effect established for alkanes reappearing here too, a more compact isomer boiling lower than an elongated one. Bromo, iodo and polychloro compounds are typically denser than water, growing heavier still as carbon count, halogen count and the halogen's own atomic mass all increase. Solubility runs low in water regardless: dissolving a haloalkane means breaking water's own strong hydrogen bonds first, and the new attractions formed with the haloalkane are never quite strong enough to repay that cost, while organic solvents, needing no hydrogen bonds broken in the first place, dissolve haloalkanes readily.

A haloalkane's carbon carries a genuine partial positive charge, exactly where the more electronegative halogen has pulled electron density away from it, and that makes it a real, specific target for a nucleophile, an electron-rich species with a pair of electrons to offer. Swap that halogen out for a nucleophile and the halide itself leaves as the departing leaving group, precisely what nucleophilic substitution means. What actually arrives in its place depends entirely on which nucleophile is chosen: hydroxide or water gives an alcohol, an alkoxide gives an ether, cyanide gives a nitrile, ammonia gives a primary amine, silver nitrite gives a nitroalkane, and even a Grignard reagent's own carbanion can arrive, extending the carbon chain into a new alkane entirely. Cyanide and nitrite are worth a special note: each can attack through either one of two different atoms, cyanide through carbon (favoured, since a C-C bond is more stable than a C-N bond, hence the alkyl cyanide product with KCN) or through nitrogen (favoured instead with the more covalent AgCN, giving an isocyanide), and nitrite similarly through oxygen or nitrogen. Species offering two distinct points of attack like this are called ambident nucleophiles, and which end actually attacks depends on exactly how ionic or covalent the nucleophile's own reagent happens to be.

Some nucleophilic substitutions happen in a single, smooth step: the incoming nucleophile approaches the carbon from directly behind the leaving halogen, exactly opposite it, and as the new bond forms, the old carbon-halogen bond breaks simultaneously, no separate intermediate ever forming along the way. This is the SN2 mechanism, bimolecular by name because its rate genuinely depends on the concentration of both the haloalkane and the nucleophile together, and the geometric consequence is dramatic: the three other groups on that carbon, previously pointing away from the halogen, sweep through the plane and flip to the opposite side entirely, exactly like an umbrella caught by a strong gust and turned inside out. This single-step approach makes SN2 acutely sensitive to crowding: methyl halides, with nothing but small hydrogens nearby, react fastest of all, while a nucleophile approaching a tertiary halide has to fight past three bulky alkyl groups blocking the direct line of attack, making tertiary halides the slowest of the three, an order, primary faster than secondary faster than tertiary, that runs in the exact opposite direction of carbocation stability.

Other substitutions take a genuinely different path: the carbon-halogen bond breaks entirely first, on its own, before any nucleophile is even involved, producing a carbocation and a departed halide ion as two separate species; only in a second, faster step does a nucleophile attack that already-formed carbocation. This is the SN1 mechanism, unimolecular, since its rate depends only on the haloalkane's own concentration, the slow, bond-breaking first step happening entirely on its own regardless of how much nucleophile happens to be nearby. Because that first step's product is a carbocation, its stability decides everything about how fast the whole reaction goes: tertiary carbocations form fastest and easiest, exactly as established earlier in this thread, giving SN1 the reactivity order tertiary faster than secondary faster than primary, the precise mirror image of SN2's own order. Allylic and benzylic halides join tertiary halides at the fast end too, for the very same reason: their carbocations gain extra stabilisation through resonance with an adjacent double bond or aromatic ring, spreading the positive charge out rather than leaving it stranded on one atom alone.

Plane-polarised light, passed through certain compounds, has its plane of polarisation rotated, some compounds turning it clockwise (dextrorotatory, marked +) and others anticlockwise (laevorotatory, marked -), and this property, optical activity, traces directly back to molecular shape. A carbon attached to four entirely different groups is called an asymmetric carbon, or stereocentre, and a molecule built around one cannot be superimposed onto its own mirror image, no matter how it is rotated in space, precisely the relationship between a left hand and a right hand: similar, but never truly identical when laid on top of each other. This non-superimposability is called chirality, and a chiral molecule is optically active; a molecule that can be superimposed on its own mirror image, like propan-2-ol, is achiral instead, and optically inactive. Two mirror-image stereoisomers of a chiral molecule are called enantiomers, and they share every physical property in common, melting point, boiling point, solubility, except one: if one enantiomer rotates polarised light clockwise, the other rotates it anticlockwise by the exact same amount. Mix equal parts of both enantiomers together and their rotations cancel perfectly, producing a racemic mixture with zero net optical rotation, marked (±) or dl before the compound's name.

Run an SN2 reaction on an optically active haloalkane and the product's configuration inverts completely, since the nucleophile only ever attacks from directly opposite the leaving halogen, systematically flipping the arrangement of the other three groups every single time, without exception; (-)-2-bromooctane reacting with hydroxide reliably gives (+)-octan-2-ol, the -OH group landing exactly where the bromide used to point away from. SN1 tells a genuinely different story: the carbocation intermediate is planar, sp2 hybridised, achiral in its own right, and a nucleophile can attack it from either face with equal ease, one face giving the same configuration as the original halide, the other giving the inverted configuration instead. Roughly equal amounts of both products usually result, and the outcome is racemisation, an optically active starting material producing an optically inactive, racemic product. This single fact turns stereochemistry into a genuine diagnostic tool: watching whether a reaction inverts, retains, or scrambles a chiral centre's configuration reveals, directly and unambiguously, which mechanism, SN2 or SN1, actually carried the reaction through.

Heat a haloalkane carrying a beta-hydrogen with alcoholic potassium hydroxide, and a genuinely different reaction can win out over substitution entirely: the halogen leaves from its own carbon while a hydrogen leaves from the adjacent (beta) carbon at the same time, forming a new double bond in place of both, beta-elimination. When more than one beta-hydrogen is available, more than one alkene becomes possible, and Zaitsev's rule, formulated in 1875, predicts which one dominates: the alkene with the greater number of alkyl groups on its doubly-bonded carbons, the more substituted, generally more stable product. Whether a haloalkane actually undergoes substitution or elimination at all is its own small contest, decided by the halide's own structure, and by the strength and bulk of whatever base or nucleophile it meets: a bulky base tends to grab a proton rather than fight through to the crowded carbon itself, favouring elimination, while a small, strong nucleophile favours substitution instead. Two further reactions put a haloalkane's reactive carbon to entirely different use: treated with magnesium metal in dry ether, it forms a Grignard reagent, RMgX, a species so reactive that even trace water destroys it instantly, releasing a hydrocarbon in the process, which is exactly why the reaction demands scrupulously dry conditions; treated instead with sodium metal in dry ether, two haloalkane molecules join directly into one, doubling the carbon count in the already-familiar Wurtz reaction.

A haloarene's carbon-halogen bond looks similar to a haloalkane's on paper, yet reacts almost not at all the same way, for three compounding reasons. First, resonance: the halogen's own lone pair delocalises into the aromatic ring, giving the C-X bond a genuine, if partial, double-bond character, and a bond with double-bond character is simply harder to break than an ordinary single bond. Second, hybridisation: the sp2 ring carbon, with more s-character than an sp3 carbon, holds its shared electrons more tightly, measurably shortening the bond, 169 pm in a haloarene against 177 pm in a comparable haloalkane, and a shorter bond resists breaking more than a longer one does. Third, mechanism itself fails: any phenyl cation formed by breaking that bond on its own would have nothing to stabilise it, no resonance available to a cation sitting directly on the ring, ruling out SN1 entirely, and an electron-rich nucleophile is, in any case, reluctant to approach an already electron-rich aromatic ring at all. Add an electron-withdrawing group like -NO2, though, and everything flips, but only from specific positions: -NO2 at ortho or para to the halogen dramatically speeds up nucleophilic substitution, since the intermediate carbanion this time is directly stabilised by resonance into the nitro group itself, while the very same -NO2 at meta has no such resonance path available and leaves reactivity essentially untouched, one nitro group, three positions, and only two of them actually matter.

Beyond nucleophilic substitution, haloarenes readily undergo the same electrophilic substitution reactions any benzene ring does, halogenation, nitration, sulphonation, Friedel-Crafts, and the halogen already present behaves exactly as established earlier in this thread: inductively electron-withdrawing, making the whole ring somewhat deactivated and slower to react overall, yet resonance-donating specifically at the ortho and para positions, directing whatever electrophile does arrive to land there anyway, deactivating and ortho/para-directing at once, no contradiction once the two effects are counted separately. Sodium metal in dry ether extends a haloarene's usefulness further still: mixed with an alkyl halide, it links the two into one alkylarene, the Wurtz-Fittig reaction; mixed with a second haloarene instead, it links two aryl groups directly into one biaryl, the Fittig reaction. Beyond single-halogen chemistry, several genuinely important compounds carry more than one halogen at once, and their real-world story is a study in trade-offs. Chloroform, once a surgical anaesthetic, slowly oxidises in light and air into poisonous phosgene gas, and is stored in dark, completely filled bottles specifically to prevent that. Carbon tetrachloride, once an everyday cleaning fluid and fire extinguisher, depletes the ozone layer once released into the atmosphere. Freons, prized for being astonishingly stable and unreactive, are stable for exactly the wrong reason: that same stability lets them drift unchanged all the way into the stratosphere, where they finally break down and disrupt the ozone layer. And DDT, a genuinely effective insecticide that once helped control malaria and typhus on a global scale, turned out to resist breakdown in the environment so completely that it accumulated steadily in animal fat over time, a discovery that led most countries, including the United States in 1973, to ban it. Every one of these compounds was adopted precisely because it was so chemically stable and unreactive, and it is that same stability, working exactly as intended, that later became each one's defining environmental problem.