Ch 1 · Alkanes · AS & A Level Chemistry 9701

Topic 1 of 2 in Hydrocarbons — AS & A Level Chemistry 9701 — 68 lessons.

Introduction

Alkanes are the simplest hydrocarbons and the least reactive class of organic compound, yet they are the starting point for almost the whole of organic chemistry, since cracking and free-radical substitution convert them into the alkenes and halogenoalkanes studied later in this module. Their unreactivity and the exceptions to it, under ultraviolet light and in combustion, are both directly examinable.

Key ideas — Making alkanes: hydrogenation and cracking

Alkanes can be produced by the hydrogenation of an alkene: an alkene is reacted with hydrogen gas over a nickel catalyst at about 150 °C, adding H_2 across the C=C bond, e.g. CH_2=CH_2 + H_2 xrightarrowNi CH_3CH_3. Alkanes are also produced industrially by cracking: heavier hydrocarbon fractions from crude oil distillation are broken down into smaller molecules, giving a mixture of shorter-chain alkanes and alkenes.

🧑‍🏫 Teacher's explanation

Think of an alkene like a flexible, slightly unstable spring. It has a double bond, which is a region of high electron density that is very reactive. Hydrogenation is like taking that spring and 'locking' it into a stable, single-bond structure by adding hydrogen atoms. It's a way to turn a reactive chemical into a stable, useful fuel.

📖 Definitions

🔍 Worked example — Easy

Calculate the mass of ethane produced when 10.0 g of ethene (C_2H_4) undergoes complete hydrogenation. (Ar: C=12, H=1)

🔍 Worked example — Standard

A hydrocarbon fraction is cracked to produce a mixture of ethene and propane. If the starting alkane is octane (C_8H_18), write a possible balanced equation for the cracking reaction.

🔍 Worked example — Exam-style

In the hydrogenation of propene (C_3H_6) to propane (C_3H_8), calculate the volume of H_2 gas required to react completely with 50.0 cm^3 of propene at RTP. (Molar volume at RTP = 24.0 dm^3 mol^-1)

⚠️ Common errors

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🎯 Exam question

Key ideas — Cracking heavier fractions

Fractional distillation of crude oil produces more of the heavier, higher-boiling fractions (e.g. long-chain alkanes used in fuel oil) than there is market demand for, and less of the lighter, more useful fractions (e.g. petrol) than is required. Cracking breaks C–C bonds in these long-chain alkanes, using either high temperature and a zeolite catalyst (catalytic cracking, favouring branched and aromatic products for motor fuels) or very high temperature and pressure alone (thermal/steam cracking, favouring alkenes such as ethene for the polymer industry). Cracking therefore matches supply to demand by converting an excess of heavy, low-value fractions into more useful, higher-demand shorter alkanes and reactive alkenes.

🧑‍🏫 Teacher's explanation

Think of a long, heavy chain of people holding hands. This chain is hard to move and doesn't do much. Cracking is like taking a giant pair of scissors and cutting that long chain into smaller, more energetic groups. Some groups might be small enough to run around easily (like petrol), while others might have 'open hands' ready to grab onto something else (like alkenes).

📖 Definitions

🔍 Worked example — Easy

A fraction of crude oil contains decane (C_10H_22). If this undergoes thermal cracking to produce ethene (C_2H_4) and an alkane, what is the formula of the alkane produced?

🔍 Worked example — Standard

Calculate the mass of ethene produced from the cracking of 100 g of octane (C_8H_18) if the reaction follows the equation C_8H_18 arrow C_2H_4 + C_6H_14.

🔍 Worked example — Exam-style

Explain why catalytic cracking is preferred for producing high-octane petrol compared to thermal cracking.

⚠️ Common errors

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🎯 Exam question

Key ideas — Combustion of alkanes

Alkanes burn readily in air, releasing energy as heat. Complete combustion, in excess oxygen, produces carbon dioxide and water only, e.g. CH_4 + 2O_2 arrow CO_2 + 2H_2O. Incomplete combustion, in a limited supply of oxygen, produces carbon monoxide, CO, or carbon (soot), together with water, e.g. 2CH_4 + 3O_2 arrow 2CO + 4H_2O.

🧑‍🏫 Teacher's explanation

Think of combustion as a high-energy dance between fuel and oxygen. When you have plenty of oxygen, every single carbon atom in the fuel can find two oxygen atoms to bond with, resulting in a 'clean' burn that produces only CO_2 and H_2 extO. This is like a well-organized party where every guest has a partner.

📖 Definitions

🔍 Worked example — Easy

Calculate the mass of CO_2 produced by the complete combustion of 32g of methane (CH_4).

🔍 Worked example — Standard

A sample of ethane (C_2H_6) undergoes incomplete combustion in a limited oxygen supply. If 1.5 moles of ethane reacts to produce 3.0 moles of CO, calculate the mass of CO produced.

🔍 Worked example — Exam-style

Propane (C_3H_8) is burned completely. If 22g of propane is used, what volume of CO_2 is produced at RTP? (Assume molar volume at RTP is 24 dm^3mol^-1)

⚠️ Common errors

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🎯 Exam question

Key ideas — Free-radical substitution mechanism

Alkanes react with chlorine or bromine only in the presence of ultraviolet light, undergoing free-radical substitution. Using methane and chlorine as an example:

🧑‍🏫 Teacher's explanation

Think of a free-radical substitution reaction like a high-stakes game of tag in a dark room. The 'radicals' are the players. A radical is a highly reactive molecule with an unpaired electron, making it extremely 'aggressive' because it wants to find a partner to become stable. When UV light hits a chlorine molecule, it provides the energy to snap the bond, creating two highly reactive 'taggers' (radicals).\n\nOnce the reaction starts, it becomes a chain reaction. One radical hits a stable methane molecule, steals an atom, and creates a new radical. This new radical then hits another chlorine molecule, creating our desired product (chloromethane) but also creating a new radical to keep the cycle going. This is why even a tiny bit of UV light can cause a massive reaction. The reaction only stops when two radicals happen to bump into each other and pair up, 'terminating' the chain.

📖 Definitions

🔍 Worked example — Easy

Identify the species produced during the initiation step when Br_2 is exposed to UV light.

🔍 Worked example — Standard

Write the two propagation steps for the reaction between ethane (C_2H_6) and chlorine (Cl_2) under UV light.

🔍 Worked example — Exam-style

In a reaction between methane and chlorine, a termination step occurs where two methyl radicals (CH_3^bullet) collide. Write the product formed.

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🎯 Exam question

Key ideas — Why alkanes are generally unreactive

Alkanes are largely unreactive towards most common reagents, including acids, bases and aqueous nucleophiles, for two reasons. First, the C-H and C-C bonds are strong (high bond enthalpy) and require a large input of energy to break. Second, since carbon and hydrogen have very similar electronegativities, the C-H bond is essentially non-polar, so alkane molecules present no region of significant partial charge to attract nucleophiles or electrophiles. This combination of strong, non-polar bonds explains why alkanes only react under forcing conditions such as ultraviolet light (free-radical substitution) or a naked flame (combustion).

🧑‍🏫 Teacher's explanation

Imagine you are trying to break into a high-security vault. To get inside, you need two things: a very strong drill to break the thick steel door, and a way to find the lock. Alkanes are like that vault. The C-C and C-H bonds are the thick steel; they are incredibly strong and require a massive amount of energy to snap.

📖 Definitions

🔍 Worked example — Easy

Calculate the total bond enthalpy for one mole of methane (CH_4) if the C-H bond enthalpy is 413 kJ mol^-1.

🔍 Worked example — Standard

Explain why alkanes do not react with aqueous sodium hydroxide (NaOH).

🔍 Worked example — Exam-style

A hydrocarbon molecule has a bond enthalpy for its C-C bonds of 348 kJ mol^-1 and C-H bonds of 413 kJ mol^-1. Calculate the energy required to break all bonds in one mole of ethane (C_2H_6).

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🎯 Exam question

Key ideas — Pollutants from combustion and their removal

Combustion of alkanes, particularly in vehicle engines, produces several environmentally damaging pollutants: carbon monoxide (from incomplete combustion, toxic as it binds to haemoglobin in place of oxygen), unburnt hydrocarbons (contribute to photochemical smog), oxides of nitrogen, NO_x (formed when the high temperature of combustion causes atmospheric N_2 and O_2 to react, contributing to acid rain and smog), and sulfur dioxide (from sulfur impurities in the fuel, contributing to acid rain). Modern vehicles use a catalytic converter, containing platinum, palladium and rhodium catalysts, to convert these pollutants into less harmful gases before they are released, e.g. 2CO + 2NO arrow 2CO_2 + N_2.

🧑‍🏫 Teacher's explanation

Think of an engine like a furnace. When you burn fuel perfectly, you get heat, water, and carbon dioxide. However, engines are rarely perfect. Sometimes there isn't enough oxygen to finish the job, leading to 'incomplete' combustion. This is like trying to burn wood in a sealed jar; you get smoke and soot instead of clean flames.

Ch 1 · Alkanes · AS & A Level Chemistry 9701

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