Hydrocarbons: Alkanes and Alkenes

LPG in a kitchen cylinder, CNG in a bus, petrol and diesel at a pump, the polythene of a shopping bag: every one of these is, at its chemical core, nothing but hydrocarbons, compounds built from carbon and hydrogen alone, no other element involved. Depending on exactly what kind of carbon-carbon bonds hold a hydrocarbon together, it falls into one of three main families. Saturated hydrocarbons, alkanes, contain only carbon-carbon single bonds; if that carbon skeleton closes into a ring instead of staying open, the same single-bonded compound is called a cycloalkane instead. Unsaturated hydrocarbons contain at least one carbon-carbon double bond (alkenes) or triple bond (alkynes), and aromatic hydrocarbons form a further, genuinely special category built around benzene's distinctive ring chemistry, covered later in this chapter. This single structural fact, how many bonds sit between any two adjacent carbons, turns out to predict almost everything else about a hydrocarbon: how it is named, how many isomers it can have, and above all, how reactive it actually is.

Methane, CH4, is the simplest alkane there is; replace one of its hydrogens with a -CH3 group and the result is ethane, C2H6, and repeating that replacement generates the entire alkane family, general formula CnH2n+2. The first three alkanes, methane, ethane and propane, each have only one possible structure, but from butane onward, more than one arrangement becomes possible: C4H10 can be a straight four-carbon chain (n-butane) or a branched one (isobutane, formally 2-methylpropane), identical formula, genuinely different compounds, chain isomers. The count only grows from there: C5H12 has three chain isomers, C6H14 has five, C7H16 has nine, and C10H22 has a startling seventy-five. Every carbon atom in any of these structures earns a label based on exactly how many other carbon atoms it directly touches: a carbon attached to no other carbon, as in methane, or to just one, as in ethane, is primary; attached to two carbons, secondary; to three, tertiary; and attached to four other carbons at once, quaternary (or neo). This primary-secondary-tertiary-quaternary vocabulary is not just bookkeeping, it reappears constantly across the rest of organic chemistry, wherever a reaction's outcome depends on exactly how substituted a particular carbon is.

Petroleum and natural gas supply most of the alkanes the world actually uses, but a chemist who needs a specific alkane in the lab has several deliberate routes available. Unsaturated hydrocarbons supply the most direct one: pass hydrogen gas over an alkene or alkyne in the presence of finely divided platinum, palladium or nickel, and it adds straight across the double or triple bond, hydrogenation, converting it fully to the corresponding alkane. Alkyl halides offer two further routes: reduced with zinc and dilute hydrochloric acid, a halide simply swaps its halogen for a hydrogen; treated instead with sodium metal in dry ether, two alkyl halide molecules join together into one longer alkane, the Wurtz reaction, useful specifically for building alkanes with an even number of carbon atoms, since it joins two identical halide fragments. Carboxylic acids supply a third route: their sodium salts, heated with soda lime, lose carbon dioxide entirely, decarboxylation, and land on an alkane with exactly one carbon fewer than the original acid; electrolysing an aqueous solution of the same sodium salt instead, Kolbe's electrolytic method, joins two acid fragments at the anode into one alkane with double the carbon count, releasing carbon dioxide and hydrogen as by-products along the way.

Carbon and hydrogen sit close enough in electronegativity that C-C and C-H bonds are almost perfectly non-polar, and that near-total lack of polarity is the single fact behind almost every physical property alkanes share: they are essentially insoluble in water ("like dissolves like", and water is polar while alkanes are not), held together only by weak van der Waals forces, and follow an unmistakable size-driven pattern, C1 to C4 are gases, C5 to C17 are liquids, C18 and heavier are solids, boiling point climbing steadily as molecular mass grows and more surface area becomes available for those weak intermolecular forces to act on. Branching runs the pattern in reverse: among pentane's three isomers, the straight chain boils highest, and neopentane, the most compact and nearly spherical, boils lowest, since a sphere-shaped molecule simply offers less surface contact to its neighbours than an elongated one does. This same non-polarity explains alkanes' chemical personality too: with no polar bond to attract a nucleophile or an electrophile, alkanes resist acids, bases, and most other reagents almost entirely under ordinary conditions, a stubbornness that earned them their older name, paraffins, from the Latin for "little affinity".

"Barely reactive" is not "never reactive", and alkanes do undergo a specific, useful set of reactions once the right trigger is applied. Halogenation replaces one or more hydrogens with a halogen under heat or light, proceeding through a genuine three-stage free radical chain mechanism: initiation, light or heat homolytically splits a halogen molecule into two reactive radicals; propagation, one of those radicals rips a hydrogen off the alkane, forming a fresh alkyl radical that immediately attacks another halogen molecule, regenerating a halogen radical and keeping the whole chain running on its own; termination, two radicals eventually collide and combine, consuming reactants and quietly stopping the chain. Complete combustion, burning in plentiful oxygen, releases carbon dioxide, water and a very large amount of heat, exactly why alkanes serve as fuels; starved of enough oxygen, incomplete combustion instead leaves carbon black behind, collected commercially for ink and pigment. A regulated, catalysed supply of oxygen achieves something gentler still, controlled oxidation, converting methane specifically to methanol or methanal rather than burning it away entirely. Two further reactions rely on heat alone: isomerisation rearranges a straight-chain alkane into a branched one using aluminium chloride and hydrogen chloride as catalysts; pyrolysis, or cracking, breaks a larger alkane down into smaller alkanes and alkenes at high temperature, the exact industrial process behind turning heavier kerosene fractions into lighter, more valuable fuels.

A single bond, unlike a double bond, allows genuinely free rotation, and rotating one half of an ethane molecule relative to the other, around the shared C-C axis, generates an infinite number of momentary spatial arrangements, called conformations or conformers, each one interconvertible into the next by nothing more than that rotation. Two extreme cases matter most. In the eclipsed conformation, the hydrogens on each carbon line up directly behind one another, as close together as possible; in the staggered conformation, they sit as far apart as possible instead, each hydrogen tucked exactly between two hydrogens on the opposite carbon. These conformations are usually drawn one of two ways: a sawhorse projection, viewing the molecule from an angle with the C-C bond drawn as a diagonal line, or a Newman projection, viewing straight down the C-C bond itself, the front carbon a dot, the rear carbon a circle. Eclipsed and staggered are not equally comfortable: bringing hydrogens close together in the eclipsed form increases electron-cloud repulsion between them, a real, if modest, energetic cost called torsional strain, and the staggered conformation, with its hydrogens held maximally apart, carries the least torsional strain of all, making it the more stable, and therefore preferred, conformation. The energy difference between the two extremes is small, only about 12.5 kJ/mol, comfortably within reach of ordinary thermal collisions at room temperature, which is exactly why rotation around ethane's C-C bond is, for all practical purposes, free, and why no one has ever managed to isolate one single conformer of ethane in a bottle.

An alkene's carbon-carbon double bond is really two different bonds stacked together: one strong sigma bond from head-on overlap of sp2 orbitals, and one distinctly weaker pi bond from sideways overlap of the two carbons' leftover p orbitals. That pi bond is shorter than a single bond, yet its own bond enthalpy is noticeably lower than the sigma bond's, since sideways overlap is fundamentally less effective than head-on overlap, and this single fact, a weak, exposed pi bond sitting above and below the molecule's plane, decides almost everything about how alkenes behave. Loosely held, easily disturbed electrons make an ideal target for electrophilic reagents, species specifically hunting for accessible electron density, and it is exactly this vulnerability that makes alkenes noticeably more reactive than the alkanes explored so far, willing to convert their weaker double bond into two new, stronger single bonds whenever a suitable electrophile comes looking. Older chemists called alkenes olefins, oil-forming, since the very first member of the series, ethene, was observed reacting with chlorine to form a distinctly oily liquid, a name and observation that, without knowing why yet, was already pointing straight at that same reactive pi bond.

Alkene names replace the alkane suffix -ane with -ene, numbering the parent chain from whichever end reaches the double bond first: propene, but-1-ene, but-2-ene. Alkenes exhibit two distinct kinds of isomerism. Structural isomerism reappears here in its usual chain and position forms, C4H8 covering but-1-ene, but-2-ene and 2-methylprop-1-ene between them. Geometrical, or cis-trans, isomerism is genuinely new, though, and it depends entirely on the double bond's one defining restriction: unlike a single bond, a double bond cannot rotate freely, since rotating it would have to break the sideways pi overlap holding it together. Attach two different groups to each doubly-bonded carbon and two distinct arrangements become possible, not interconvertible by any amount of rotation: cis, with two identical groups on the same side of the rigid double bond, and trans, with them on opposite sides, best pictured as two pieces of cardboard joined by a nail through their middle, free to be positioned but never twisted past each other. Cis and trans isomers of the same compound are not identical substances; but-2-ene's cis form carries a measurable dipole moment (0.33 D) since its two methyl groups sit on the same side and their individual bond dipoles do not cancel, while the trans form's opposing methyl groups cancel almost perfectly, leaving it essentially non-polar, and this same asymmetry shows up in melting points too, the more symmetrical trans isomer typically packing into a solid more efficiently and melting at a higher temperature than its cis counterpart.

Alkynes supply one route: partial hydrogenation with hydrogen over Lindlar's catalyst, palladium deliberately poisoned to stop the reaction cleanly at the alkene stage rather than running through to a full alkane, delivers a cis alkene specifically, while reducing the same alkyne with sodium in liquid ammonia instead delivers the trans isomer. Alkyl halides supply a second route: heated with alcoholic potassium hydroxide, a halide eliminates one molecule of hydrogen halide entirely, dehydrohalogenation, losing a hydrogen from the carbon next to the one carrying the halogen and forming a double bond in its place. Vicinal dihalides, two halogens on adjacent carbons, supply a third: treated with zinc metal, they lose both halogens at once as zinc halide, dehalogenation, and gain a double bond exactly where the halogens used to be. And alcohols supply a fourth: heated with concentrated sulphuric acid, an alcohol loses an entire water molecule, acidic dehydration, and the two carbons that used to hold the -OH and its neighbouring hydrogen form a double bond between them instead. Four different starting materials, four genuinely different mechanisms, one identical destination: a carbon-carbon double bond, freshly created.

Add HBr across a symmetrical alkene like ethene and there is only one possible product; add it across an unsymmetrical alkene like propene and suddenly there are two, and only one of them dominates in practice. In 1869, Russian chemist Vladimir Markovnikov studied exactly this and formulated a rule that still carries his name: the negative part of the adding molecule attaches to the carbon that already carries fewer hydrogen atoms. The mechanism behind it is a direct consequence of carbocation stability, already established earlier in this thread. H+ attacks the double bond first, and it can form either of two possible carbocations; whichever one is more stable forms faster and dominates, and since alkyl groups stabilise a positive charge through inductive donation and hyperconjugation, the more substituted carbocation, the secondary one in propene's case, wins decisively over the less substituted primary alternative. Br- then simply attacks whichever carbocation actually formed, landing 2-bromopropane as the major product, exactly what Markovnikov's rule predicts, now with a mechanistic reason behind it rather than just an empirical pattern.

Add HBr to propene in the presence of a peroxide, though, and the product reverses entirely: 1-bromopropane forms instead, in direct defiance of the rule that seemed so reliable a moment ago. Discovered by Kharash and Mayo in 1933, this peroxide, or Kharash, effect is genuinely anti-Markovnikov, and it happens for a specific, narrow reason: peroxide changes the entire mechanism from an ionic, carbocation-based pathway to a free radical one instead. A bromine radical, generated first from the peroxide, attacks the double bond and forms whichever radical is more stable, again the secondary one, and hydrogen then adds to complete the reaction, landing the bromine on the terminal carbon this time, the opposite regiochemistry from the ordinary ionic pathway. This reversal is strikingly narrow in scope: it happens with HBr specifically, and not with HCl or HI, since the H-Cl bond is too strong for the bromine-radical-generating cycle to break efficiently, while the H-I bond is so weak that iodine radicals simply recombine into I2 rather than adding to the alkene at all. One single reagent, one single mechanism, and one single element of the halogen family for which it actually works.

Cold, dilute, aqueous potassium permanganate (Baeyer's reagent) converts an alkene gently into a vicinal glycol, two -OH groups on what used to be the double-bonded carbons, and the reagent's own purple colour fading away as this happens serves as a simple, reliable test for unsaturation. Push the same oxidation harder, with hot, acidic potassium permanganate or potassium dichromate, and the double bond does not just gain two -OH groups, it breaks completely, the two fragments oxidised further into ketones or carboxylic acids depending on how many hydrogens sat on each original alkene carbon. Ozonolysis pushes this cleavage idea to its most precise: ozone adds across the double bond to form an unstable ozonide, and treating that ozonide with zinc and water splits it cleanly into two carbonyl fragments, aldehydes or ketones, right at the exact position the double bond used to occupy. This makes ozonolysis a genuine detective tool rather than just another reaction: given the two carbonyl products of an unknown alkene's ozonolysis, a chemist can work backward and reconstruct exactly where that double bond originally sat, and even the full original structure, purely from the pieces it broke into.

Heat ethene under high pressure with the right catalyst, and instead of reacting as separate, individual molecules, thousands of them link end to end into one enormous chain, -CH2-CH2- repeating over and over: polythene, and the reaction that builds it, polymerisation, joining many small monomer molecules into one giant polymer molecule. Propene undergoes the exact same kind of reaction to form polypropene. Both materials are now genuinely everywhere, plastic bags, squeeze bottles, buckets, pipes, television cabinets, precisely because a long, uniform hydrocarbon chain turns out to be cheap to produce and remarkably versatile to mould. That same ubiquity carries a real cost, though: polythene and polypropene resist the natural breakdown processes that eventually consume most other organic waste, and their sheer volume of everyday use has made plastic accumulation a genuine and growing environmental concern, the exact same reactive pi bond that builds a useful material into thousands of copies of itself also being the reason none of those copies disappear again on their own.