Waking up the boring molecule
Across this rung you have learned to name alkanes and to twist them into chair and Newman shapes, and along the way you may have noticed something almost insulting: nothing *happens* to them. An alkane is just carbon and hydrogen joined by strong, non-polar single bonds, with no lone pairs, no charge, and no obvious place to grab. Acids ignore it, bases ignore it, most nucleophiles slide right off. That inertness is exactly why we burn them for fuel and wrap our food in them — but it makes the alkane a frustrating place to learn your first reaction, because the molecule simply will not cooperate under ordinary conditions.
There is one door, though. Mix methane with chlorine gas in the dark and they sit politely forever. Add ultraviolet light or strong heat and they suddenly react, swapping a hydrogen for a chlorine to give chloromethane, CH3Cl, plus HCl. The same trick works with bromine. This is radical halogenation, and the reason it covers a haloalkane onto an otherwise unreactive chain is that light or heat opens a door no acid or base could — it breaks a bond straight down the middle. Understanding that single move is the whole point of this guide.
Breaking a bond down the middle
Almost every reaction you will meet later breaks bonds *unevenly*: both shared electrons leave with one atom, making a cation and an anion — that is heterolysis, the world of nucleophiles and electrophiles. Radical chemistry is the other way to break a bond. In homolysis, the two electrons of a bond split up like a divorce settlement — one electron goes to each fragment. The result is not an ion but a free radical: a neutral atom or group carrying a single, unpaired electron, lonely and desperate to pair up again. We draw this electron flow with single-barbed fishhook arrows, each showing one electron moving, in contrast to the full double-barbed arrow that moves a pair.
Why does the halogen break and not the alkane? Energy. The Cl–Cl bond is fairly weak, so a UV photon carries more than enough energy to snap it apart by homolysis into two chlorine atoms, each a radical with seven valence electrons aching for one more. The C–H and C–C bonds of the alkane are much stronger and survive the same photon untouched. The number that governs all of this is the bond dissociation energy — the energy needed to homolyse a bond into two radicals — and weak bonds break first. That is the entire reason the halogen is the one that opens the door.
The chain reaction, step by step
Here is the surprise that makes radical halogenation worth its own guide: one photon can drive thousands of molecules to react. That is because the reaction is a self-sustaining loop, a radical chain mechanism, built from three kinds of step. The first, initiation, is the only one that needs the light: it spends a photon to split Cl2 into two chlorine radicals. The second kind, propagation, is where the real work and the chain happen — and crucially, each propagation step consumes one radical but produces a new one, so the chain never stops itself. The third, termination, is when two radicals finally meet and pair up, quietly ending one strand of the chain.
- Initiation. A UV photon strikes Cl2 and homolyses it: Cl2 -> two Cl radicals. One photon spent; two reactive chlorine atoms born. Nothing useful is made yet — this only lights the fuse.
- Propagation, step one. A chlorine radical, hungry for one electron, plucks a hydrogen off the alkane: Cl radical + CH4 -> HCl + CH3 radical. The chlorine is satisfied as HCl, but a brand-new carbon radical is left behind — the radical is handed on, not destroyed.
- Propagation, step two. The carbon radical grabs a chlorine from a fresh Cl2: CH3 radical + Cl2 -> CH3Cl + Cl radical. The product CH3Cl is finally made — and a new chlorine radical pops out, ready to start step one all over again on the next alkane.
- Termination. Now and then two radicals collide instead of finding a fresh molecule — Cl radical + CH3 radical -> CH3Cl, or two CH3 radicals -> CH3CH3. Each such meeting removes two radicals and snaps that strand of the chain. Because radicals are rare, these collisions are uncommon, which is why one initiation can power a long run.
Read those two propagation steps together and the magic appears: step one eats a chlorine radical and makes a carbon radical; step two eats the carbon radical and makes a chlorine radical. The radical is never used up — it is passed around the loop like a baton, making product on every lap. A single initiation can therefore turn over many thousands of molecules before a chance termination finally ends that chain. That self-feeding turnover is the defining feature of any chain mechanism, and it is why a tiny spark of light can chew through a whole flask of alkane.
Why bromine is picky and chlorine is sloppy
A real alkane like propane has two kinds of hydrogen: six on the end carbons (primary) and two on the middle carbon (secondary). When the halogen radical pulls one off, which does it take? The answer divides chlorine from bromine and is the most useful thing in this guide. The carbon radical that forms is more stable when it sits on a more substituted carbon — the order is roughly tertiary > secondary > primary > methyl, the same ranking you will later see for carbocations, and for the same family of reasons (neighbouring C–H and C–C bonds help spread out the lonely electron). So all else being equal, both halogens *prefer* to make the more stable radical. The question is how strongly that preference shows up.
radical stability (more substituted = more stable): 3deg > 2deg > 1deg > methyl chlorination of propane (~room temp): ~43% on the 2deg H, ~57% on the 1deg H bromination of propane: ~97% on the 2deg H, ~3% on the 1deg H (there are six 1deg H and only two 2deg H, yet bromine still goes for the 2deg site)
The split comes from how *reactive* each halogen radical is. A chlorine radical is wildly reactive and a little indiscriminate: it abstracts a hydrogen so easily, in a step that releases energy, that it barely waits to find the best one — it grabs almost whatever it bumps into, primary or secondary, so chlorination gives a messy mixture. A bromine radical is far gentler and lazier; its hydrogen-abstraction step is uphill in energy and reluctant, so the bromine radical is choosy and waits for the easiest, most stabilising hydrogen — the one that yields the most stable radical. That is selectivity, and it follows a broad pattern: a less reactive reagent is usually a more selective one, because it can afford to be picky. We say bromination has high selectivity and radical stability is what it is selecting *for*.
Why does "slow and reluctant" mean "choosy"? Because a reluctant, uphill step has a late transition state — at the moment of reaction the new radical is already nearly fully formed, so the energy gap between making a stable radical and a shaky one is felt strongly, and the molecule steers hard toward the stable one. The fast, downhill chlorine step has an early transition state where the radical has barely begun to form, so that energy gap is hardly felt and chlorine cannot tell the hydrogens apart. This early-versus-late idea has a name you will meet again — the Hammond postulate.
Honest limits, and a door to the radical chapter
Be honest about what this reaction is and is not good for. Even bromination is not perfectly clean — "97% secondary" still means a few percent goes elsewhere, and once you have made one CH3Cl, that product still has hydrogens, so the halogen can come back for a second and a third substitution to give CH2Cl2, CHCl3 and CCl4. Chemists fight this by using a big excess of alkane, so a radical is far likelier to meet fresh starting material than a product. Radical halogenation is therefore a real but blunt tool: wonderful for making the rare symmetric or single-type-of-hydrogen substrate, frustrating for a precise lab synthesis. It earns its place here as your first reaction precisely because the alkane gives you no gentler option.
One more honest caution, because it trips up almost everyone. Selectivity is about *which* hydrogen, not *how much* product — and the two can disagree. Propane has six primary hydrogens and only two secondary ones, so even though each individual secondary H is much more reactive, chlorine's mild preference is outweighed by sheer numbers and you still get a roughly even split. The right way to think is per-hydrogen reactivity times how many of each there are. Bromine's preference is so steep that it wins anyway; chlorine's is so shallow that the head-count usually wins. Whenever you predict a major product, weigh both factors, never just the stability.
Keep this guide as a seed, not the full tree. A whole later chapter is devoted to radicals, and it will reward what you have just built: it sharpens radical stability into numbers, explains how an isotope swap (a heavier hydrogen, deuterium) reveals which step is rate-limiting through the kinetic isotope effect, and shows off cleverer, tamer radical reactions — selective allylic bromination with a reagent called NBS, radical additions that run the opposite way to the ionic ones, and the radical polymerization that builds the plastic in your water bottle. For now you have the core that everything else hangs on: homolysis makes radicals, a chain carries them, and reactivity buys selectivity. That is a real reaction, mechanism and all — your first.