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Alkanes are the simplest family of organic molecules — saturated hydrocarbons with the general formula CₙH₂ₙ₊₂, tetrahedral geometry at every carbon, and strong, non-polar C–C and C–H bonds that make them comparatively unreactive. Yet they dominate the global energy economy. Crude oil, an enormously complex mixture of alkanes, cycloalkanes and aromatic hydrocarbons, is separated by fractional distillation into fuel fractions, longer chains are cracked into shorter alkanes and alkenes, and the products are burned (complete and incomplete combustion) or chemically modified. The single A-Level reaction mechanism for alkanes themselves is free-radical substitution — methane reacting with chlorine under UV light. This lesson covers all of it: properties, fractional distillation, thermal and catalytic cracking, complete and incomplete combustion, environmental pollutants and their abatement, the full free-radical substitution mechanism with curly half-arrows, selectivity in higher alkanes, and the comparative chemistry of fluorine, chlorine, bromine and iodine.
Spec mapping (AQA 7405): This lesson maps to §3.3.2 (alkanes), specifically §3.3.2.1 (fractional distillation of crude oil), §3.3.2.2 (modification of alkanes by cracking), §3.3.2.3 (combustion of alkanes) and §3.3.2.4 (chlorination of alkanes). Cross-refer to L0 of this course (nomenclature and basic structure) for IUPAC naming of the chloroalkane products; to L4 (alkenes — electrophilic addition) for the contrasting reactivity of the C=C π bond; to L8 (mechanism master class) for the broader survey of A-Level mechanisms; and to §3.1.5 (kinetics) where radical chain reactions are revisited in the context of rate equations and chain propagation. Refer to the official AQA 7405 specification document for the exact wording of each section.
Assessment objectives: AO1 recall items include the general formula of alkanes, the products of complete and incomplete combustion, the definitions of homolytic fission, free radical, initiation, propagation and termination, and the role of UV light in chlorination. AO2 application items include writing balanced combustion equations for any specified alkane, writing the full free-radical substitution mechanism with curly half-arrows for each homolysis or radical-substitution step, identifying the products of incomplete combustion (CO and C), and predicting the position isomers formed when a higher alkane reacts with Cl₂ or Br₂. AO3 evaluation items include explaining and quantifying the environmental impact of CO₂, CO, NOₓ, SO₂ and particulates; predicting which mono-substituted product will be the major product in a higher alkane and justifying the answer using C–H bond enthalpies and radical stability; and comparing fluorination, chlorination, bromination and iodination quantitatively.
| Property | Trend | Explanation |
|---|---|---|
| Boiling point | Increases with chain length | More electrons → stronger London (induced-dipole) dispersion forces |
| Boiling point | Decreases with branching | More compact shape → smaller molecular surface area → weaker London forces |
| Solubility in water | Insoluble | Cannot form hydrogen bonds; energy cost of disrupting H₂O hydrogen-bond network is too high |
| Solubility in non-polar solvents | High | "Like dissolves like" — London forces between alkane and solvent compensate for those broken in pure liquid |
| Density | Less dense than water | All liquid alkanes float on water; this matters for oil-spill containment |
| Combustion enthalpy per CH₂ unit | ≈ −650 kJ mol⁻¹ | Adds approximately linearly with chain length — the basis of "calorific value" of fuels |
Key Definition: London dispersion forces (also called induced dipole–induced dipole forces, or temporary-dipole–induced-dipole forces) are weak, short-range intermolecular forces arising from instantaneous fluctuations in electron density that induce matching dipoles in neighbouring molecules. They are the only intermolecular forces present between non-polar molecules such as alkanes. Their strength scales with the number of electrons and the molecular surface area of contact.
Petroleum (crude oil) is a complex natural mixture of hundreds of hydrocarbons — predominantly straight-chain and branched alkanes, with cycloalkanes and aromatic compounds (the "BTEX" — benzene, toluene, ethylbenzene, xylenes) as minor components. Crude oil is not directly useful as a fuel; it must be separated into fractions of similar boiling-point range, each tuned to a specific application (transport fuel, industrial fuel, feedstock for the chemical industry).
Separation is achieved by fractional distillation in a tall steel column with a vertical temperature gradient — hottest (~400 °C) at the base, coolest (~25 °C) at the top. Vaporised crude oil enters near the bottom; each fraction condenses on a tray at the height where the column temperature falls just below the fraction's boiling-point range.
| Fraction | Approx. chain length | Boiling point range / °C | Principal uses |
|---|---|---|---|
| Refinery gases | C₁–C₄ | Below 25 | LPG (propane, butane); calor gas, camping stoves |
| Petrol (gasoline) | C₅–C₁₀ | 25–75 | Motor fuel for spark-ignition engines |
| Naphtha | C₈–C₁₂ | 75–190 | Petrochemical feedstock — cracking and reforming feed |
| Kerosene (paraffin) | C₁₀–C₁₆ | 190–250 | Jet fuel; domestic heating |
| Diesel (gas oil) | C₁₅–C₂₅ | 250–350 | Diesel engines; central heating oil |
| Fuel oil (heavy gas oil) | C₂₀–C₇₀ | 350–400 | Ship and power-station fuel |
| Bitumen (residue) | C₇₀+ | Above 400 (non-volatile) | Road surfacing; roofing |
Key Definition: Fractional distillation is the separation of a liquid mixture into components by exploiting differences in boiling point, achieved in a column with a vertical temperature gradient where each component condenses at the tray whose temperature matches its boiling point.
The economic problem with crude oil is that the supply profile (large amounts of long-chain heavy fractions) does not match the demand profile (large amounts of short-chain petrol and feedstock). Cracking solves this mismatch.
Cracking is the breaking of long-chain alkanes into shorter alkanes and alkenes. The general scheme is:
long-chain alkane → shorter alkane + alkene
For example: C₁₅H₃₂ → C₈H₁₈ + C₂H₄ + C₂H₄ + C₃H₆ (a typical product mix; specific products depend on conditions and where the chain breaks).
The shorter alkane is more valuable as a fuel (petrol-range), and the alkene is the feedstock for the polymer industry (poly(ethene), poly(propene)) and for further organic synthesis. There are two industrial variants of cracking:
| Feature | Thermal cracking | Catalytic cracking |
|---|---|---|
| Conditions | High temperature (700–1200 °C), high pressure (~70 atm), short contact time (~1 s) | Lower temperature (450 °C), slight pressure, zeolite catalyst |
| Mechanism | Homolytic C–C fission → free radical intermediates → β-scission | Heterolytic C–C fission → carbocation intermediates → rearrangement |
| Catalyst | None | Zeolite (aluminosilicate, e.g. ZSM-5 or faujasite); shape-selective microporous solid |
| Product distribution | High proportion of alkenes (especially ethene, propene) — feedstock-focused | High proportion of branched and aromatic petrol-range alkanes — fuel-focused, higher-octane |
| Industrial role | Steam crackers — petrochemical feedstock | Fluid catalytic crackers (FCC) — petrol refining |
Key Definition: A zeolite is a microporous aluminosilicate framework whose precisely-sized pores admit small molecules and exclude large ones. The acidic sites within the pores catalyse heterolytic bond fission via carbocation intermediates; the pore geometry imposes shape selectivity on the products.
The two mechanisms are mutually informative for A-Level: thermal cracking proceeds via the same kind of radical intermediates we will meet in chlorination of alkanes (this lesson); catalytic cracking proceeds via carbocations of the kind we will meet in electrophilic addition to alkenes (next lesson). Crude oil refining is a working demonstration of the two great mechanism families of organic chemistry.
Alkanes are the dominant transport and industrial fuels of the global energy economy precisely because their combustion is highly exothermic, their bond enthalpies are well-matched to safe storage and controlled ignition, and their fractions span the volatility range needed for spark ignition (petrol), compression ignition (diesel) and turbine combustion (jet fuel).
In excess oxygen, an alkane burns cleanly to carbon dioxide and water. The general equation is:
CₙH₂ₙ₊₂ + (3n + 1)/2 O₂ → nCO₂ + (n + 1)H₂O
| Alkane | Balanced equation |
|---|---|
| Methane | CH₄ + 2O₂ → CO₂ + 2H₂O |
| Ethane | 2C₂H₆ + 7O₂ → 4CO₂ + 6H₂O |
| Propane | C₃H₈ + 5O₂ → 3CO₂ + 4H₂O |
| Butane | 2C₄H₁₀ + 13O₂ → 8CO₂ + 10H₂O |
| Octane | 2C₈H₁₈ + 25O₂ → 16CO₂ + 18H₂O |
The enthalpy of complete combustion of octane is approximately −5470 kJ mol⁻¹; the standard enthalpy of combustion per mole of CH₂ unit is approximately −650 kJ mol⁻¹ regardless of chain length, which is why fuel calorific value scales nearly linearly with mass.
Exam Tip: Balance combustion equations by balancing C first (set the coefficient of CO₂), then H (set the coefficient of H₂O), then O last (set the coefficient of O₂). If you end up with a fractional coefficient on O₂, multiply the entire equation through by 2.
When oxygen is limited (poorly maintained boilers, blocked flues, fuel-rich engine settings), combustion is incomplete. Possible products are carbon monoxide (CO) and carbon (soot, C) alongside water:
In real engines and burners the exhaust is a mixture of all three carbon-containing products (CO₂, CO, C) in proportions that depend on the air-to-fuel ratio (the stoichiometric or lambda value).
Why incomplete combustion matters:
Combustion in real engines produces a wider range of pollutants than the CO₂/CO/C picture above suggests. The high temperatures and pressures inside a cylinder activate side reactions that have nothing to do with the fuel itself.
| Pollutant | Source | Environmental / health effect | Abatement |
|---|---|---|---|
| CO₂ | Complete combustion of any hydrocarbon | Greenhouse gas; main driver of anthropogenic climate change; atmospheric concentration risen from ~280 ppm (pre-industrial) to ~420 ppm (2024) | Reduce demand; carbon capture and storage (CCS); switch to renewables and nuclear |
| CO | Incomplete combustion (limited O₂) | Toxic — binds haemoglobin, prevents O₂ transport | Catalytic converter oxidises CO to CO₂ |
| Particulates (C, PM2.5) | Incomplete combustion, especially in diesel engines and biomass burning | Respiratory disease, cardiovascular disease, premature death | Diesel particulate filters (DPF); biomass cooker chimneys |
| NOₓ (NO + NO₂) | N₂ + O₂ → 2NO at the high temperatures inside engine cylinders (~2000 K); not from the fuel itself | Acid rain (HNO₃); photochemical smog; respiratory irritant; ozone formation | Three-way catalytic converter reduces NOₓ to N₂; selective catalytic reduction (SCR) with urea in diesel engines |
| SO₂ | Combustion of sulfur impurities in crude-oil-derived fuels | Acid rain (H₂SO₃ → H₂SO₄); respiratory irritant; corrodes stonework | Pre-combustion desulfurisation; flue-gas desulfurisation (FGD) with CaCO₃ in power stations |
| Unburned hydrocarbons (UHC) | Incomplete combustion, evaporative losses | Photochemical smog; some are carcinogenic (e.g. benzene, polycyclic aromatic hydrocarbons) | Catalytic converter; engine tuning; fuel-vapour recovery at petrol pumps |
The catalytic converter in petrol cars is a synoptic link to the transition-metal chemistry of the inorganic course: a honeycomb of ceramic substrate coated with platinum, palladium and rhodium nanoparticles catalyses three simultaneous reactions: 2CO + O₂ → 2CO₂; unburned hydrocarbons + O₂ → CO₂ + H₂O; and 2NO + 2CO → N₂ + 2CO₂. The "three-way" name reflects these three abatement reactions running in parallel.
Practical-skills box — Observing chlorination under UV: A capped boiling tube containing methane (or a methane-rich gas such as natural gas) and chlorine is held in front of a UV lamp behind a safety screen. The initially yellow-green Cl₂ colour fades over 30–60 seconds as Cl₂ is consumed in propagation. Damp universal indicator paper held at the tube mouth on opening turns red, confirming the production of HCl. Wafting the vapours gently towards the nose (do not inhale directly) gives the sweet smell of chloroform-like products (CHCl₃, CCl₄). Repeat under a normal incandescent lamp shielded from UV: no observable reaction, confirming the role of UV in initiation. Risk assessment requires fume cupboard operation, eye protection, and full Cl₂/HCl handling protocols.
The single A-Level reaction mechanism for alkanes themselves is free-radical substitution by a halogen under UV light. Methane reacts with chlorine to give chloromethane (and hydrogen chloride):
CH₄ + Cl₂ → CH₃Cl + HCl (overall, UV)
This is the only A-Level mechanism involving free-radical intermediates and single-headed (fish-hook) curly arrows. Master it now and you have a reusable mental template for the radical chain steps of polymerisation (next lesson) and the radical mechanism of stratospheric ozone destruction by CFCs (year 13).
| Detail | |
|---|---|
| Reagents | Alkane + Cl₂ (or Br₂) |
| Conditions | Ultraviolet light — provides the energy (~243 kJ mol⁻¹ photon energy) to break the Cl–Cl bond by homolytic fission |
| Type of mechanism | Free-radical substitution (chain reaction) |
| Type of bond fission | Homolytic — each bonded atom takes one electron from the shared pair |
| Curly arrows | Single-headed (fish-hook) — each arrow represents the movement of one electron |
UV light supplies the energy to break the Cl–Cl bond homolytically:
Cl–Cl → 2 Cl·
Two single-headed (fish-hook) curly arrows are drawn from the centre of the Cl–Cl bond, one pointing to each Cl atom; each arrow represents the movement of one electron. Each chlorine atom leaves with one electron and is now an electron-deficient species with an unpaired electron — a chlorine radical (Cl·).
Key Definition: Homolytic fission is the breaking of a covalent bond such that each bonded atom receives one electron from the shared pair, forming two free radicals. It is represented by two single-headed (fish-hook) curly arrows, in contrast to heterolytic fission, in which the bonded atoms separate as a cation and an anion (one taking both bonding electrons), represented by a single double-headed curly arrow.
Key Definition: A free radical is a species (atom, molecule or ion) with at least one unpaired electron in its valence shell. Radicals are usually highly reactive intermediates and are denoted with a single dot beside the chemical symbol, e.g. Cl·, ·CH₃.
Two propagation steps form a self-sustaining chain: the radical produced by step 1 is consumed in step 2 and replaced by step 2's product radical, and so on. Each "turn" of the chain consumes one Cl₂ and one CH₄ and produces one CH₃Cl and one HCl.
Propagation step 2a:
CH₄ + Cl· → ·CH₃ + HCl
The chlorine radical abstracts a hydrogen atom from methane. One fish-hook curly arrow is drawn from the C–H bond to the H atom (showing one electron of the bond leaving with H to make the new H–Cl bond), and a second fish-hook curly arrow from the unpaired electron on Cl· to the new H–Cl bond. The remaining electron of the original C–H bond stays on carbon, generating a methyl radical (·CH₃).
Propagation step 2b:
·CH₃ + Cl–Cl → CH₃Cl + Cl·
The methyl radical attacks a Cl₂ molecule. One fish-hook curly arrow is drawn from the unpaired electron on ·CH₃ to one Cl of Cl₂ (forming the new C–Cl bond), and a second fish-hook curly arrow from the centre of the Cl–Cl bond to the departing Cl atom (which leaves with one electron as a new Cl·). The Cl· produced feeds back into step 2a, sustaining the chain.
A single initiation event can in principle drive thousands of propagation cycles before a termination step ends the chain. The chain length (the average number of propagation cycles per initiation event) is typically 10³ to 10⁴ for chlorination of methane.
A chain ends whenever two radicals meet and combine, removing two reactive intermediates from the system and producing a closed-shell molecule:
Exam Tip: Examiners insist on (i) all three stages labelled (initiation, propagation, termination), (ii) correct single-headed curly arrows where curly arrows are asked for, (iii) explicit radical dots on every radical, and (iv) at least one propagation step showing the radical chain "feedback" — the regenerated Cl· that returns to the start. The detection of ethane (C₂H₆) in the product mixture is the classic experimental evidence that the mechanism is radical — only radical–radical combination can give a C–C bond formation product.
The general approach transfers directly. The overall equation is C₂H₆ + Cl₂ → C₂H₅Cl + HCl.
| Step | Equation |
|---|---|
| Initiation | Cl–Cl → 2 Cl· (UV) |
| Propagation 1 | C₂H₆ + Cl· → ·C₂H₅ + HCl |
| Propagation 2 | ·C₂H₅ + Cl–Cl → C₂H₅Cl + Cl· |
| Termination | 2 Cl· → Cl₂; Cl· + ·C₂H₅ → C₂H₅Cl; 2 ·C₂H₅ → C₄H₁₀ (butane) |
The same mechanism, with Br· instead of Cl·, accounts for bromination — though as we will see, bromination is markedly more selective than chlorination.
The first major limitation of free-radical substitution as a preparative method is further substitution. The product CH₃Cl still has three C–H bonds and can react again with Cl· in exactly the same way:
CH₃Cl + Cl₂ → CH₂Cl₂ + HCl CH₂Cl₂ + Cl₂ → CHCl₃ + HCl CHCl₃ + Cl₂ → CCl₄ + HCl
In practice, chlorination of methane gives a mixture of CH₃Cl, CH₂Cl₂, CHCl₃ and CCl₄ whose proportions depend on the Cl₂:CH₄ ratio. To favour the mono-substituted product, a large excess of methane is used (giving Cl· a much higher probability of meeting CH₄ than meeting CH₃Cl). The mixture is then separated by fractional distillation, but the yield of any one product is modest. This is why free-radical substitution is not normally used as a synthesis method for chloroalkanes in the laboratory — nucleophilic substitution of alcohols with HCl, or addition of HCl to alkenes, give much cleaner mono-substituted products.
The second selectivity problem is that higher alkanes have C–H bonds in different chemical environments, all of which can be attacked. For propane (CH₃–CH₂–CH₃) there are two types of H atoms:
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