Alkanes — Preparation, Properties, Reactions

What Are Alkanes

Alkanes are saturated hydrocarbons in which every carbon atom is $sp^3$ hybridised and joined to its neighbours by single covalent bonds alone. Successive members differ by a $\ce{-CH2-}$ unit, and constructing the family member by member shows that the molecular formula always satisfies the relation $\ce{C_{n}H_{2n+2}}$. Because each carbon already carries the maximum number of hydrogen atoms permitted by its four bonds, alkanes are described as saturated.

The tetrahedral geometry around each carbon dictates a $\ce{C-C-C}$ bond angle close to $109.5^\circ$, so an extended alkane chain is not straight but zig-zag. The C–C and C–H bonds are nearly non-polar because the electronegativity difference between carbon and hydrogen is small, and this electronic character is the root of almost every property discussed below — from low solubility in water to chemical inertness. Alkanes are historically called paraffins, from the Latin parum affinis, meaning of little affinity.

Nomenclature and Isomerism

The first three alkanes — methane, ethane and propane — have only one possible structure each. From butane onwards, the same molecular formula can be arranged into more than one skeleton, giving rise to chain (skeletal) isomerism. Butane, $\ce{C4H10}$, exists as the straight-chain n-butane and the branched 2-methylpropane (isobutane); pentane, $\ce{C5H12}$, has three isomers: n-pentane, 2-methylbutane and 2,2-dimethylpropane (neopentane).

In IUPAC nomenclature the longest continuous carbon chain is selected as the parent and numbered so that substituents receive the lowest set of locants. Alkyl groups, obtained by removing one hydrogen from an alkane ($\ce{-CH3}$ methyl, $\ce{-C2H5}$ ethyl), are cited as prefixes with their position numbers. Hydrogen atoms themselves are classified by the carbon they sit on — a point that becomes decisive when we reach halogenation selectivity.

Preparation of Alkanes

Petroleum and natural gas are the principal natural sources of alkanes, but several laboratory routes allow targeted synthesis. NCERT §9.2.2 groups them by the starting material: unsaturated hydrocarbons, alkyl halides, and carboxylic acids. Each method has a characteristic constraint that examiners exploit, so the carbon-counting logic matters as much as the equation itself.

1. From unsaturated hydrocarbons (hydrogenation)

Dihydrogen adds across the multiple bonds of alkenes and alkynes over finely divided platinum, palladium or nickel — the Sabatier–Senderens reaction. The metal adsorbs $\ce{H2}$ on its surface and activates the H–H bond. Platinum and palladium operate at room temperature, while nickel needs higher temperature and pressure.

$$\ce{CH2=CH2 + H2 ->[Pt/Pd/Ni] CH3-CH3}$$ $$\ce{CH3-C#C-H + 2H2 ->[Pt/Pd/Ni] CH3-CH2-CH3}$$

2. From alkyl halides

Alkyl halides (except fluorides) on reduction with zinc and dilute hydrochloric acid give alkanes with the same carbon count:

$$\ce{CH3-Cl + H2 ->[Zn, H+] CH4 + HCl}$$

The more important route is the Wurtz reaction: two molecules of an alkyl halide treated with sodium metal in dry ether couple to give a higher alkane containing twice the carbon atoms of the halide. The product therefore always has an even number of carbons and a symmetrical skeleton.

$$\ce{2CH3Br + 2Na ->[dry\ ether] CH3-CH3 + 2NaBr}$$ $$\ce{2C2H5Br + 2Na ->[dry\ ether] CH3-CH2-CH2-CH3 + 2NaBr}$$

3. From carboxylic acids

Heating the sodium salt of a carboxylic acid with soda lime (sodium hydroxide and calcium oxide, typically in a 3:1 ratio) eliminates carbon dioxide and yields an alkane with one carbon atom fewer than the acid. This loss of $\ce{CO2}$ is called decarboxylation.

$$\ce{CH3COONa + NaOH ->[CaO][\Delta] CH4 + Na2CO3}$$

Kolbe's electrolytic method takes the aqueous sodium or potassium salt of a carboxylic acid and electrolyses it. At the anode the carboxylate loses an electron and then $\ce{CO2}$ to give an alkyl free radical; two such radicals combine, so the alkane formed has an even number of carbon atoms.

$$\ce{2CH3COONa + 2H2O ->[electrolysis] C2H6 + 2CO2 + H2 + 2NaOH}$$

Physical Properties

Alkanes are colourless, odourless and almost non-polar; they dissolve in non-polar solvents but are essentially insoluble in water, following the rule like dissolves like. Their molecules attract one another only through weak van der Waals forces, which is why the lightest members are gases. The first four members $\ce{C1}$ to $\ce{C4}$ are gases, $\ce{C5}$ to $\ce{C17}$ are liquids, and those with 18 or more carbons are solids at 298 K.

Boiling point rises steadily with molecular mass because the van der Waals forces grow with molecular size and surface area. Among isomers, branching lowers the boiling point: as a molecule becomes more branched it approaches a spherical shape, reducing the surface area available for contact and weakening the intermolecular attraction. Hence the three isomeric pentanes boil in the order $n\text{-pentane} > 2\text{-methylbutane} > 2{,}2\text{-dimethylpropane}$.

Chemical Properties

Alkanes are generally inert towards acids, bases, oxidising agents and reducing agents — the very inertness implied by the name paraffin. Under forcing conditions, however, they undergo a defined set of reactions. The most heavily examined of these is the substitution reaction, in which one or more hydrogen atoms are replaced by halogen, nitro or sulphonic acid groups.

Halogenation proceeds in diffused sunlight or ultraviolet light, or at high temperature ($573$–$773$ K). Chlorination of methane does not stop cleanly at one stage but proceeds through successive substitution to a mixture of products:

$$\ce{CH4 + Cl2 ->[h\nu] CH3Cl + HCl}$$ $$\ce{CH3Cl + Cl2 ->[h\nu] CH2Cl2 + HCl}$$ $$\ce{CH2Cl2 + Cl2 ->[h\nu] CHCl3 + HCl}$$ $$\ce{CHCl3 + Cl2 ->[h\nu] CCl4 + HCl}$$

The reactivity of the halogens themselves follows $\ce{F2 > Cl2 > Br2 > I2}$. Fluorination is too violent to control, while iodination is slow and reversible — it must be driven by an oxidising agent such as $\ce{HIO3}$ or $\ce{HNO3}$ to remove the $\ce{HI}$ produced. The ease with which different hydrogens are replaced follows the order $3^\circ > 2^\circ > 1^\circ$, which governs the major product when an alkane has more than one type of hydrogen.

Free-Radical Halogenation Mechanism

Halogenation proceeds by a free-radical chain mechanism in three phases: initiation, propagation and termination. The chain is set off because the weak $\ce{Cl-Cl}$ bond breaks more easily than the stronger $\ce{C-C}$ and $\ce{C-H}$ bonds when energised by light or heat.

Initiation — homolysis of chlorine produces two chlorine radicals: $$\ce{Cl2 ->[h\nu] 2Cl^.}$$ Propagation — a chlorine radical abstracts a hydrogen from methane to give a methyl radical; this radical then attacks a fresh chlorine molecule, forming the product and regenerating a chlorine radical, so the chain continues: $$\ce{CH4 + Cl^. -> CH3^. + HCl}$$ $$\ce{CH3^. + Cl2 -> CH3Cl + Cl^.}$$ Termination — the chain stops when two radicals combine, consuming radicals without making new ones. Combination of two methyl radicals explains why ethane appears as a by-product of methane chlorination: $$\ce{Cl^. + Cl^. -> Cl2}$$ $$\ce{CH3^. + CH3^. -> CH3-CH3}$$ $$\ce{CH3^. + Cl^. -> CH3Cl}$$

Combustion, Oxidation and Cracking

Beyond substitution, alkanes undergo combustion, controlled oxidation, isomerisation, aromatisation, reaction with steam and pyrolysis. These define their industrial value as fuels and feedstocks.

Combustion — in excess air or oxygen, alkanes are completely oxidised to $\ce{CO2}$ and water with the release of a large amount of heat, which is why they are used as fuels. The general equation is: $$\ce{C_{n}H_{2n+2} + ( \frac{3n+1}{2} )O2 -> nCO2 + (n+1)H2O}$$ With insufficient air, incomplete combustion gives carbon black, used in inks and pigments: $$\ce{CH4 + O2 -> C + 2H2O}$$

Controlled oxidation with a regulated supply of oxygen and the right catalyst halts at a useful intermediate such as methanol, methanal or acetic acid. Isomerisation over anhydrous aluminium chloride converts straight-chain into branched alkanes (raising the octane number of fuel), while aromatisation dehydrogenates and cyclises $\ce{C6+}$ alkanes into benzene and its homologues over oxides of vanadium, molybdenum or chromium on alumina. Finally, pyrolysis (cracking) breaks higher alkanes into smaller alkanes and alkenes by heat, the basis of converting kerosene into oil gas and lighter fuels.