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Spec Mapping — OCR H432 Module 4.1.2 — Hydrocarbons, covering complete and incomplete combustion of alkanes as fuels, the formation of the principal combustion pollutants (CO, NOₓ, SO₂, unburned hydrocarbons and particulates), the role of catalytic converters in modern vehicles, and the free-radical substitution reaction of methane and other alkanes with chlorine or bromine in the presence of ultraviolet light, including the three stages (initiation, propagation, termination), the use of half-arrows ("fish-hooks") to indicate homolytic bond fission, and the limitations of the reaction as a synthetic route (refer to the official OCR H432 specification document for exact wording).
Alkanes are saturated hydrocarbons whose only bonds are non-polar C–H and C–C σ bonds. Both bond types are strong and non-polar, so alkanes are essentially unreactive towards polar reagents (acids, alkalis, oxidising or reducing agents, nucleophiles). Yet two classes of reaction matter enormously: combustion in oxygen, which is the basis of essentially the entire fossil-fuel economy, and free-radical substitution with halogens under ultraviolet light — the textbook example of a radical chain reaction. Both reactions proceed through homolytic bond fission (each atom of the breaking bond keeps one electron, generating neutral radical intermediates) rather than the heterolytic processes that dominate the rest of organic chemistry. This lesson develops both topics in the depth OCR expects: the chemistry, the equations, the mechanisms, the environmental issues raised by combustion, and the kinetic and selectivity problems that limit radical substitution as a synthesis.
Key Mechanism: Free-radical substitution is a three-stage chain reaction. (1) Initiation: UV light homolyses a halogen molecule (e.g., Cl₂ → 2 Cl•). (2) Propagation: a radical reacts with a closed-shell molecule to give a different radical and a closed-shell product (two steps that together sum to the overall equation). (3) Termination: two radicals combine to form a closed-shell product, ending that particular chain.
The chemistry of an alkane is dominated by the kinds of bond it contains. The C–C bond has an average bond enthalpy of about 347 kJ mol⁻¹ and the C–H bond about 413 kJ mol⁻¹. Both are strong. Both are also virtually non-polar: the Pauling electronegativities of C (2.55) and H (2.20) differ by only 0.35, far below the ~1.7 threshold for significant polar/ionic character. With no permanent dipole on any C–H bond, alkanes present no δ+ atom to attract a nucleophile and no δ− atom to attract an electrophile. They have no lone pairs to donate, no π bond to act as a nucleophile, and no easily-broken polar bond to act as a leaving group. Polar reagents in solution simply ignore them. What alkanes can be persuaded to do is react under conditions that supply enough energy (heat or UV photons) to break a non-polar bond homolytically — splitting the bonding pair so each fragment keeps one electron. The two situations OCR cares about are combustion and halogenation under UV light.
Key Definition — Homolytic fission: breaking of a covalent bond in which each atom retains one of the bonding electrons, producing two neutral radicals. Drawn with half-arrows ("fish-hooks") on each atom. Compare with heterolytic fission, in which one atom takes both electrons and the products are oppositely-charged ions.
Combustion is the rapid reaction of a hydrocarbon with oxygen to give oxides of carbon and water, with the release of a large amount of heat. The reaction is highly exothermic, and that is why alkanes — from methane in natural gas to long-chain alkanes in jet fuel and ship diesel — supply the vast majority of human primary energy.
In complete combustion, the alkane burns in a plentiful supply of oxygen and is fully oxidised to CO₂ and H₂O:
CnH2n+2+23n+1O2→nCO2+(n+1)H2O
Worked examples:
The enthalpy of combustion becomes more exothermic with chain length, in line with the increasing number of C–H and C–C bonds broken and the larger number of new C=O and O–H bonds formed. The balancing trick is always to balance carbon first, then hydrogen, and finally oxygen — oxygen is the messy coefficient because it has to mop up after both.
If the oxygen supply is limited (poorly ventilated boiler, old car engine, blocked flue, faulty Bunsen burner) the alkane is only partially oxidised. The products may include carbon monoxide (CO) and carbon (soot) alongside water:
A real, dirty flame produces a mixture of CO₂, CO and C — a gas mixture and a yellow sooty deposit, rather than the clear blue flame of complete combustion.
Carbon monoxide is the silent killer of household gas appliances. It is colourless, odourless and tasteless, so victims cannot detect it without a CO sensor. Physiologically, CO binds to the iron(II) centre of haemoglobin about 240× more avidly than O₂ does, forming carboxyhaemoglobin (COHb). The binding is functionally irreversible on the time-scale of breathing, so even sub-percent atmospheric concentrations rapidly knock out a significant fraction of the blood's oxygen-carrying capacity. Symptoms progress from headache and dizziness through confusion and unconsciousness to death within hours at moderate exposure. Hundreds of accidental UK deaths a year are still attributed to CO from poorly-maintained heating appliances, despite mandatory carbon-monoxide alarms.
Sulfur dioxide (SO₂): crude oil contains sulfur impurities (typically up to 3 % by mass before refining). When the fuel burns, the sulfur burns with it:
S+O2→SO2
SO₂ dissolves in atmospheric moisture to form sulfurous acid and is oxidised by O₂ or NOₓ to sulfuric acid:
SO2+H2O→H2SO3;2SO2+O2→2SO3;SO3+H2O→H2SO4
The resulting acid rain damages limestone and marble buildings (CaCO₃ + H₂SO₄ → CaSO₄ + CO₂ + H₂O), acidifies lakes and rivers (mobilising aluminium ions toxic to fish), and damages forest foliage. Modern refineries now hydrodesulfurise crude oil to below 10 ppm sulfur before sale, dramatically reducing UK SO₂ emissions since the 1970s.
Nitrogen oxides (NOₓ): at the high temperatures inside an internal-combustion engine (~2000 °C in the cylinder), atmospheric N₂ — usually unreactive — reacts with O₂:
N2+O2→2NO;2NO+O2→2NO2
NOₓ contributes to acid rain (forming HNO₃), to photochemical smog (catalysing tropospheric ozone formation), and to respiratory illness. It is most acute in city-centre air, where diesel engines dominate.
Unburned hydrocarbons and particulates: small droplets of unreacted fuel and condensed soot (the "PM2.5" and "PM10" fractions of urban air pollution) deposit on lung tissue, contributing to asthma, COPD and lung cancer. The black smoke from old diesel engines is the visible end of this distribution.
Modern petrol cars contain a three-way catalytic converter in the exhaust system. A honeycomb of ceramic substrate is coated with platinum, palladium and rhodium nanoparticles, giving an enormous surface area in a small package. It catalyses three reactions simultaneously:
The net effect is to convert the three principal pollutants (CO, NOₓ, unburned hydrocarbons) into less harmful CO₂, N₂ and H₂O. Note that CO₂ is still a greenhouse gas and is not removed — the converter solves local-air-quality problems, not climate ones.
Alkanes react with chlorine (or bromine) gas in the presence of ultraviolet light to produce a haloalkane and a hydrogen halide. One hydrogen on the alkane is replaced by one halogen — it is a substitution reaction — and the mechanism is a radical chain reaction in three stages.
Overall reaction (methane + chlorine):
CH4+Cl2→CH3Cl+HCl
Conditions: UV light (to homolyse Cl–Cl). Without UV — no reaction at room temperature. In the dark, a sealed mixture of methane and chlorine is indefinitely stable.
graph LR
A["Initiation<br/>UV breaks Cl-Cl<br/>generates 2 Cl·"] --> B["Propagation Step 1<br/>Cl· + CH4 → HCl + ·CH3"]
B --> C["Propagation Step 2<br/>·CH3 + Cl2 → CH3Cl + Cl·"]
C --> B
C --> D["Termination<br/>two radicals combine<br/>chain ends"]
A photon of ultraviolet light (λ ≈ 250–360 nm; the C–H bond is too strong to be broken at these wavelengths but the Cl–Cl bond at 242 kJ mol⁻¹ is just weak enough) provides enough energy to homolytically break the Cl–Cl bond:
Cl2hν, UV2Cl⋅
This is the only step in which radicals are created from non-radical species. Draw the two electrons of the Cl–Cl bond as half-arrows, each terminating on one Cl atom.
In propagation, a radical reacts with a closed-shell molecule to give a different radical and a closed-shell product. The radical count is conserved: one in, one out. Two propagation steps must be written:
Propagation Step 1:
Cl⋅+CH4→HCl+⋅CH3
The chlorine radical abstracts a hydrogen atom from methane. The C–H bond breaks homolytically (one electron stays with C, one goes to form the new H–Cl bond). A methyl radical •CH₃ is produced.
Propagation Step 2:
⋅CH3+Cl2→CH3Cl+Cl⋅
The methyl radical abstracts a chlorine atom from Cl₂, giving chloromethane and regenerating a chlorine radical. The chlorine radical can then attack another methane molecule, and the cycle continues.
Crucially, the two propagation steps add up to the overall equation:
CH4+Cl2→CH3Cl+HCl
Because each Cl• can attack many CH₄ molecules before being destroyed, a single photon of UV can trigger thousands of substitutions. This amplification — a chain reaction — is what makes the reaction practically useful (and explains why a candle plus chlorine gas in a fume cupboard can give explosive flashes of HCl-rich smoke).
Termination steps consume radicals without producing new ones, ending the chain. Any two radicals can combine:
Termination explains why small quantities of ethane appear as a by-product even though the overall stoichiometry of the reaction does not require it. OCR may ask for three plausible termination steps; you should be able to write all three.
Despite a clean overall equation, free-radical substitution is rarely a useful pure-synthesis route. Three problems dominate:
Once CH₃Cl forms, it still contains three C–H bonds and reacts further:
The product is always a mixture of mono-, di-, tri- and tetra-substituted compounds, plus unreacted methane and HCl. Mitigation: running the reaction with a large excess of methane statistically favours mono-substitution (each Cl• is more likely to encounter CH₄ than CH₃Cl).
For larger alkanes, abstraction can occur from different carbons, giving isomeric products. Propane + Cl₂ gives both 1-chloropropane and 2-chloropropane. Tertiary C–H bonds are abstracted slightly faster than secondary, which are faster than primary, because the resulting alkyl radical is increasingly stabilised by hyperconjugation (3° > 2° > 1° > methyl). For chlorination the per-H reactivity ratio is roughly 1 : 4 : 5 (1° : 2° : 3°); for bromination it is much steeper (1 : 80 : 1700) because bromination is less exothermic and therefore more selective.
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