What makes a molecule an alkane
From the earlier rungs you already know how to read a carbon skeleton off a zig-zag and that carbon's talent for chaining up with itself — its catenation — is what makes organic chemistry possible at all. Now we meet the plainest molecules that talent can build. A hydrocarbon is any compound made of only carbon and hydrogen, and an alkane is the simplest kind: a hydrocarbon in which every bond is a single bond. There are no double bonds, no triple bonds, no oxygen, no nitrogen — just carbons holding hands with each other and with hydrogens. Because every carbon is bonded to the maximum number of other atoms it can reach, alkanes are called saturated: they are carrying a full load and have no room for more.
Every carbon in an alkane is sp3 hybridized — the same picture you met when we built up hybridization. Each carbon points its four single bonds toward the corners of a tetrahedron, about 109.5 degrees apart, and each of those bonds is a strong, cylindrically symmetric sigma bond. That tetrahedral geometry is the whole reason a chain is drawn as a zig-zag rather than a straight rod, and — as the next guides in this rung will show — it is also why a single bond can quietly rotate, letting the molecule twist into different shapes without breaking anything.
A family that climbs by CH2
Line up the alkanes by size and something tidy appears. Methane is CH4, ethane is C2H6, propane is C3H8, butane is C4H10. Each step up adds exactly one carbon and two hydrogens — one CH2 unit — to the one before it. A series like this, where every member differs from its neighbour by a constant CH2, is a homologous series, and the alkanes are the textbook example. Because the pattern is perfectly regular, you can write a single general formula that fits every straight-chain or branched alkane at once: CnH2n+2. Plug in n = 5 and you instantly know pentane is C5H12; no drawing required.
open-chain alkanes CnH2n+2 (n = 1, 2, 3, ...) n=1 CH4 methane n=4 C4H10 butane n=2 C2H6 ethane n=5 C5H12 pentane n=3 C3H8 propane n=6 C6H14 hexane close the chain into ONE ring -> lose 2 H -> CnH2n (cycloalkane) cyclohexane = C6H12, not C6H14
There is a close cousin worth meeting now. Bend a carbon chain around until its two ends join, and you get a ring — a cycloalkane such as cyclohexane. Forming the ring costs you the two hydrogens that used to sit on the now-joined end carbons, so the formula drops from CnH2n+2 down to CnH2n. That single ring is exactly one degree of unsaturation, even though there is not a double bond anywhere in sight — a useful reminder that the degree of unsaturation counts rings and pi bonds together. The reason a six-membered ring like cyclohexane is so common and so comfortable (its famous chair shape) is a story this rung tells in its own dedicated guide.
Where they come from and how they behave
Alkanes are not laboratory curiosities — they are the bulk of the world's fuel. Petroleum (crude oil) and natural gas are vast underground mixtures of alkanes, the slow remains of ancient marine life cooked over millions of years. A refinery separates crude oil mostly by distillation, splitting it by boiling point into fractions: natural gas (methane through butane) at the light end, then gasoline, kerosene, diesel, and heavy waxes and asphalt at the bottom. The same homologous-series regularity that lets you predict a formula also lets a refinery predict where each molecule will boil.
Their physical properties climb just as predictably as their formulas. Methane, ethane, propane, and butane are gases; pentane up through about C17 are liquids; the long chains are waxy solids. The pattern follows directly from the polarity and intermolecular-forces ideas of the foundation rung. Alkanes are almost perfectly nonpolar — a C-H bond barely separates charge at all — so the only thing holding two alkane molecules together is the weak, fleeting London dispersion force. A longer chain has more surface to brush against its neighbours, so the intermolecular forces add up and the boiling point rises. Branching makes a molecule more compact and ball-like, shrinking that contact surface, which is why a branched isomer boils a little lower than its straight-chain twin. And because alkanes are nonpolar, they refuse to mix with water and float on top of it — the reason an oil spill spreads as a slick.
Why alkanes are so unreactive
Alkanes were once called paraffins, from the Latin for "little affinity" — they barely react with anything. At room temperature they ignore acids, bases, oxidants, and reductants that would tear other molecules apart. The reason is built into their bonds. A C-C or C-H bond is strong, with a high bond dissociation energy (around 410 kJ/mol for C-H), so it costs a lot of energy to break. Just as important, those bonds are nearly nonpolar: there is no electron-poor carbon for a nucleophile to attack and no electron-rich site for an electrophile, so the usual machinery of polar organic reactions simply has nothing to grab. An alkane is chemistry's blank canvas — stable, strong, and indifferent.
The two reactions they will do
Indifferent does not mean inert forever. Alkanes have exactly two reactions worth knowing, and both rely on brute force rather than the gentle electron-pushing of polar chemistry. The first is combustion: given enough heat to start, an alkane burns in oxygen to carbon dioxide and water, releasing a great deal of energy. This is the entire point of petroleum as fuel. Methane burning is simply CH4 + 2 O2 -> CO2 + 2 H2O. With too little oxygen the burn is incomplete and yields carbon monoxide or soot instead — which is exactly why a faulty heater is dangerous.
The second is radical halogenation. Shine ultraviolet light (or apply heat) on a mixture of an alkane and a halogen such as chlorine, and a hydrogen gets swapped for a halogen: CH4 + Cl2 -> CH3Cl + HCl. What makes this special is the kind of reaction it is. Because alkane bonds are nonpolar, they do not break unevenly into ions; instead light cleaves Cl-Cl straight down the middle into two free radicals, each carrying one lonely unpaired electron. That even, one-electron-each splitting is called homolysis, and the radical halogenation that follows is the headline reaction of this rung's later guides.
- Initiation: UV light splits a halogen molecule by homolysis — Cl2 absorbs a photon and breaks into two chlorine radicals, each with one unpaired electron.
- Propagation: a chlorine radical yanks a hydrogen off the alkane to form HCl plus a carbon radical; that carbon radical then grabs a chlorine from another Cl2, making the product and regenerating a chlorine radical.
- Because each propagation step hands off a fresh radical, one photon can drive many cycles — a self-sustaining chain. Two radicals meeting and pairing up ends a chain (termination).
Notice the honest limitation already lurking here. With chlorine and a real alkane you do not get one clean product — the radical can pull off a hydrogen from several positions, and once one halogen is on it can come back for more, giving mixtures of mono-, di-, and tri-substituted products. That messiness, and the question of *which* hydrogen comes off most easily (which ties straight back to bond strengths and radical stability), is the whole subject the later guides untangle. For now, the headline is enough: alkanes are the quiet baseline of organic chemistry, and these two reactions — combustion and radical halogenation — are the only doors that open without a more reactive group already present.