HookThe tower that sorts a barrel of sludge without a single reaction
When a tanker of crude oil docks at Fawley on Southampton Water, what it delivers is close to worthless: a dark, treacly mixture of hundreds of different alkanes, useless as it stands for anything at all. Within hours it has become bottled camping gas, petrol, jet kerosene, diesel and road bitumen — and the first, largest separation involves no chemical reaction whatsoever. A fractionating column sixty metres tall exploits a single fact: bigger alkane molecules have stronger van der Waals forces and therefore higher boiling points. Heat the mixture and the fractions condense out at different heights of the tower.
Everything else in Section 3.3.2 is what you do to those fractions afterwards, and it is four linked ideas. Why alkane boiling points rise with chain length, and how fractional distillation turns that into a separation. How cracking breaks the surplus heavy fractions into the alkenes and petrol the market actually wants. What alkanes produce when they burn — including the pollutants combustion quietly makes. And the free-radical substitution mechanism that proves alkanes are not as inert as their reputation suggests. It is a topic heavy on equations and mechanisms, and Paper 2 rewards precision in both over any amount of vague prose.
ModelAlkanes and van der Waals — why size sets the boiling point
Alkanes are saturated hydrocarbons: only single \(\text{C}-\text{C}\) and \(\text{C}-\text{H}\) bonds, general formula \(\text{C}_n\text{H}_{2n+2}\). Carbon and hydrogen have almost the same electronegativity, so the molecules are essentially non-polar and the only forces between them are van der Waals forces (induced dipole–dipole attractions). This one fact drives their physical behaviour. A longer chain has more electrons and a larger surface over which molecules make contact, so the van der Waals attractions are stronger, more energy is needed to pull the molecules apart, and the boiling point rises steadily up the series.
Branching works the other way. A branched isomer is more spherical and cannot pack as closely against its neighbours, so the surface contact and the van der Waals forces both fall. Pentane boils at 36 °C, but its branched isomer 2,2-dimethylpropane boils at just 9.5 °C — same formula, weaker forces, lower boiling point. That single relationship — more contact means stronger van der Waals means higher boiling point — is the engine that makes fractional distillation possible.
DataFractional distillation — a column that boils by height
Crude oil is vaporised and fed into the base of a fractionating column that is hot at the bottom (around 350 °C) and progressively cooler towards the top. The vapours rise, and each fraction condenses at the height where the temperature matches its boiling range. Short-chain, low-boiling molecules with weak van der Waals forces — refinery gas and gasoline (petrol) — travel highest before condensing; long-chain, high-boiling molecules — fuel oil and bitumen — barely rise and tap off near the bottom. A fraction is a group of hydrocarbons with a similar range of chain lengths and boiling points, not a pure compound.
Critically, this is a physical separation: no covalent bonds are broken and no new substances are made — the process simply sorts molecules that were already there. The fractions, from top to bottom, are refinery gas (bottled LPG), gasoline (petrol), naphtha (a feedstock for the chemical industry), kerosene (jet fuel), gas oil (diesel), fuel oil, and the bitumen used for road surfaces and roofing.
CaseCracking — breaking big molecules the market cannot sell
Distillation leaves refineries with a mismatch: too much heavy, long-chain fraction and not enough petrol and reactive alkenes. Cracking fixes it by breaking long-chain alkanes into shorter alkanes and alkenes, and industry runs it two ways. Thermal cracking uses a high temperature (roughly 700–1000 °C) and high pressure (up to about 7000 kPa), proceeds by a free-radical mechanism, and yields a high proportion of small alkenes such as ethene and propene — the raw material of the polymer industry. Catalytic cracking uses a zeolite catalyst (an acidic aluminosilicate honeycomb) at a gentler temperature (around 450 °C) and only slight pressure, proceeds via carbocation intermediates, and yields branched alkanes, cycloalkanes and aromatic compounds prized for high-octane motor fuels.
The economic driver is the same in both cases: turn a low-value surplus into high-value products. The catalyst in catalytic cracking earns its keep by lowering the temperature needed, cutting the energy bill and speeding the process up.
Cracking conserves every atom, so the products must add back up to the starting alkane. Crack decane, \(\text{C}_{10}\text{H}_{22}\): one split gives octane plus ethene, \(\text{C}_{10}\text{H}_{22} \rightarrow \text{C}_8\text{H}_{18} + \text{C}_2\text{H}_4\). Check the atoms — carbon \(10 = 8 + 2\), hydrogen \(22 = 18 + 4\). A different split gives \(\text{C}_{10}\text{H}_{22} \rightarrow \text{C}_7\text{H}_{16} + \text{C}_3\text{H}_6\) (heptane plus propene), also balanced. There must be at least one alkene among the products: splitting a saturated \(\text{C}_n\text{H}_{2n+2}\) into two saturated pieces is impossible, because count the hydrogens and you always fall two short unless a \(\text{C}=\text{C}\) double bond forms to absorb the difference.
DataCombustion — the point of it all, and its pollutant bill
Complete combustion, with plenty of oxygen, converts an alkane to carbon dioxide and water and releases a great deal of energy — the reason alkanes are fuels at all. Incomplete combustion, with limited oxygen, produces toxic carbon monoxide (CO, which binds to haemoglobin and starves the body of oxygen) and/or carbon particulates (soot), alongside water; the sign is a yellow, smoky flame instead of a clean blue one. Burning fuels generates further pollutants: oxides of nitrogen (NO and \(\text{NO}_2\)) form when the fierce temperature inside an engine forces atmospheric nitrogen and oxygen to react — crucially, this nitrogen comes from the air, not the fuel — and sulfur dioxide forms from sulfur impurities in the fuel. Both oxides of nitrogen and sulfur dioxide cause acid rain.
Two clean-up technologies matter. Catalytic converters turn carbon monoxide, oxides of nitrogen and unburnt hydrocarbons into carbon dioxide, nitrogen and water, for example \(2\text{CO} + 2\text{NO} \rightarrow 2\text{CO}_2 + \text{N}_2\). Flue-gas desulfurisation removes sulfur dioxide from power-station exhaust using calcium oxide or calcium carbonate, \(\text{CaO} + \text{SO}_2 \rightarrow \text{CaSO}_3\). Neither removes the carbon dioxide itself, a greenhouse gas that combustion produces by design.
Balance the complete combustion of octane, \(\text{C}_8\text{H}_{18} + \text{O}_2 \rightarrow \text{CO}_2 + \text{H}_2\text{O}\). Balance carbon first (8 \(\text{CO}_2\)), then hydrogen (9 \(\text{H}_2\text{O}\)); the right side now holds \(16 + 9 = 25\) oxygen atoms, needing 12.5 \(\text{O}_2\), so double everything to clear the half: \(2\text{C}_8\text{H}_{18} + 25\text{O}_2 \rightarrow 16\text{CO}_2 + 18\text{H}_2\text{O}\). Now the desulfurisation follow-up: if a fuel burns to give 64 g of \(\text{SO}_2\) (\(M_r = 64.0\), so 1.0 mol), removing it with calcium oxide, \(\text{CaO} + \text{SO}_2 \rightarrow \text{CaSO}_3\), needs 1.0 mol of CaO \(= 56\) g. Every equation in this topic is marked on the balancing first and the chemistry second, so always check the atoms tally before you move on.
MechanismFree-radical substitution — three stages, and count the radicals
Alkanes are largely unreactive — strong, non-polar bonds give nucleophiles and electrophiles nothing to grip — but in ultraviolet light they react with halogens by free-radical substitution, in which a hydrogen atom is replaced by a halogen. The mechanism has three named stages. Initiation: UV light supplies the energy to break the halogen molecule by homolytic fission into two radicals, \(\text{Cl}_2 \rightarrow 2\,\text{Cl}\bullet\). Propagation: a two-step chain that regenerates a radical each time, so a handful of initiating events drives thousands of product molecules — \(\text{Cl}\bullet + \text{CH}_4 \rightarrow \bullet\text{CH}_3 + \text{HCl}\), then \(\bullet\text{CH}_3 + \text{Cl}_2 \rightarrow \text{CH}_3\text{Cl} + \text{Cl}\bullet\). Notice the radical count is conserved in propagation: one radical is consumed and one is made. Termination: two radicals combine and are removed from the system, for example \(\bullet\text{CH}_3 + \text{Cl}\bullet \rightarrow \text{CH}_3\text{Cl}\).
Its limitations are why it is a poor way to make a pure compound despite a tidy-looking overall equation. Further substitution replaces the remaining hydrogens too, giving a mixture of \(\text{CH}_3\text{Cl}\), \(\text{CH}_2\text{Cl}_2\), \(\text{CHCl}_3\) and \(\text{CCl}_4\); with longer chains you also get a mixture of positional isomers. The clean overall equation hides a messy product flask.
Chlorination of methane, overall \(\text{CH}_4 + \text{Cl}_2 \rightarrow \text{CH}_3\text{Cl} + \text{HCl}\). Initiation (UV light essential): \(\text{Cl}_2 \rightarrow 2\,\text{Cl}\bullet\), homolytic fission. Propagation step 1: \(\text{Cl}\bullet + \text{CH}_4 \rightarrow \bullet\text{CH}_3 + \text{HCl}\). Propagation step 2: \(\bullet\text{CH}_3 + \text{Cl}_2 \rightarrow \text{CH}_3\text{Cl} + \text{Cl}\bullet\). Add the two propagation steps and the radicals cancel exactly to give the overall equation — proof they are short-lived intermediates, not products. A termination step such as \(2\,\bullet\text{CH}_3 \rightarrow \text{C}_2\text{H}_6\) explains the trace of ethane found among the products, and that ethane is the experimental fingerprint that free radicals are real.
VocabularyKey terms the mark scheme pays for
TrapsMisconceptions that cost marks
ExamWhat examiners want
Equations are the currency of this topic and they are marked on balancing before chemistry, so always check the atoms tally and state the conditions: UV light for the initiation step, the temperature, pressure and catalyst for cracking. When a question says ‘outline the mechanism’ of chlorination, the overall equation earns nothing — you must show all three stages, label them, place the dot on each radical correctly, and use fishhook (single-headed) arrows if arrows are requested. A favourite one-mark question, ‘state the essential condition for the first step’, wants exactly two words: ultraviolet light.
Boiling-point explanations (AO1/AO2) must reach the cause, not stop at the symptom. ‘Bigger molecules boil higher’ is incomplete; the mark is for more electrons and greater surface contact, so stronger van der Waals forces needing more energy to overcome. For pollutants, examiners want the pollutant, its source and its effect — name the gas, say where it comes from, and give the consequence such as acid rain or toxicity.
For cracking, identify the type, give its conditions and its typical products, and be ready to state the economic reason — matching supply to demand for petrol and alkenes. This content sits on Paper 2 (organic and physical chemistry) and returns synoptically on Paper 3, where the free-radical mechanism reappears alongside the ozone chemistry of the next section.