Alkanes

Welcome to Alkanes

Welcome to one of the most fundamental chapters in Organic Chemistry! Alkanes might sound like a fancy chemical term, but you interact with them every single day. The natural gas heating your home, the petrol fuelling cars, and the wax on a birthday candle are all made of alkanes.

In this chapter, we will explore the structure, physical properties, fractional distillation, cracking, combustion, and the step-by-step mechanism of free radical substitution. Don't worry if organic mechanisms seem intimidating at first—we will break down each concept step by step so you can master your CCEA AS examinations with confidence!

1. Structure and Bonding in Alkanes

Alkanes are the simplest family of organic molecules. They belong to a homologous series called saturated hydrocarbons.

• Hydrocarbon: A compound containing only carbon and hydrogen atoms.
• Saturated: A molecule containing only single covalent bonds between carbon atoms (no double or triple bonds).

General Formulas

• Aliphatic (straight or branched chain) alkanes: \(\text{C}_n\text{H}_{2n+2}\)
For example, if \(n = 3\), the alkane is propane: \(\text{C}_3\text{H}_{2(3)+2} = \text{C}_3\text{H}_8\).
• Cycloalkanes (ring structures): \(\text{C}_n\text{H}_{2n}\)
For example, cyclopropane is \(\text{C}_3\text{H}_6\). Note that cycloalkanes have two fewer hydrogen atoms because the ring closure uses two bonding positions.

Shape and Bond Angles

Every carbon atom in an alkane forms four single covalent bonds, known as \(\sigma\) (sigma) bonds. These bonds are formed by the direct head-on overlap of atomic orbitals.

Because there are four bonded electron pairs and zero lone pairs around each carbon atom, the electron pairs repel each other equally to minimize repulsion:

• Arrangement/Shape around each Carbon: Tetrahedral
• Bond Angle: Approximately \(109.5^\circ\)

Memory Tip: Picture a camera tripod with an extra leg pointing straight up into the air—that is the classic tetrahedral shape of \(109.5^\circ\)!

Key Takeaway: Alkanes have the general formula \(\text{C}_n\text{H}_{2n+2}\), contain only single \(\sigma\)-bonds, and have a tetrahedral geometry around every carbon atom with bond angles of \(109.5^\circ\).

2. Physical Properties of Alkanes

1. Polarity and Intermolecular Forces

Carbon and hydrogen have very similar electronegativities (\(\text{C} = 2.5\), \(\text{H} = 2.1\)). Consequently, \(\text{C}-\text{H}\) and \(\text{C}-\text{C}\) bonds are non-polar. This means alkane molecules are non-polar overall.

The only intermolecular forces holding alkane molecules together are weak van der Waals forces (also called London dispersion forces or induced dipole-dipole forces).

2. Trend in Boiling Points: Effect of Carbon Chain Length

As the carbon chain length increases:

1. The molecular size increases, meaning the molecule contains more electrons.
2. The surface contact area between neighbouring molecules increases.
3. This creates stronger van der Waals forces between molecules.
4. More thermal energy is required to overcome these intermolecular attractions, leading to a higher boiling point.

3. Trend in Boiling Points: Effect of Branching

If you compare straight-chain and branched-chain isomers with the same molecular formula (e.g., straight-chain pentane vs. branched 2,2-dimethylpropane):

• Straight-chain alkanes can pack closely together with large surface contact areas, giving stronger van der Waals forces and higher boiling points.
• Branched alkanes are more spherical and cannot pack as closely together. This reduces surface contact area, leading to weaker van der Waals forces and lower boiling points.

4. Solubility

• In Water: Alkanes are insoluble (immiscible). They cannot form hydrogen bonds with polar water molecules.
• In Non-polar Solvents: Alkanes dissolve well in non-polar organic solvents (like hexane or benzene) because the van der Waals forces between solute and solvent are similar in strength.

Key Takeaway: Boiling points increase with longer carbon chains (more electrons and greater surface contact) and decrease with increased branching (less surface contact).

3. Fractional Distillation of Crude Oil

Crude oil (petroleum) is a non-renewable fossil fuel consisting of a complex mixture of mostly unbranched and branched alkanes.

Because crude oil is not directly useful as a mixture, it is separated into simpler mixtures called fractions using fractional distillation.

How Fractional Distillation Works Step-by-Step

1. Vaporisation: Crude oil is heated in a furnace until most of it vaporises.
2. Entering the Column: The hot vapour is pumped into the bottom of a fractionating column.
3. Temperature Gradient: The column is hot at the bottom (around \(350\text{ }^\circ\text{C}\) to \(400\text{ }^\circ\text{C}\)) and cool at the top (around \(25\text{ }^\circ\text{C}\)).
4. Condensation: As vapours rise, they cool. When a fraction reaches a tray where the temperature is just below its boiling point, it condenses into a liquid and is tapped off.
5. Separation:
• Short-chain alkanes have low boiling points, so they rise to the top and exit as gases.
• Long-chain alkanes have high boiling points, so they condense near the hot bottom or remain as thick residue.

Main Fractions (from top to bottom of the column)

• Refinery gases (\(\text{C}_1-\text{C}_4\)): Bottled gas for domestic heating/cooking (methane, propane, butane).
• Petrol / Gasoline (\(\text{C}_5-\text{C}_{10}\)): Fuel for motor cars.
• Naphtha (\(\text{C}_8-\text{C}_{12}\)): Chemical feedstock for petrochemical industry.
• Kerosene / Paraffin (\(\text{C}_{11}-\text{C}_{15}\)): Jet engine fuel and domestic heating.
• Diesel oil / Gas oil (\(\text{C}_{15}-\text{C}_{19}\)): Fuel for diesel engines (trucks, trains, buses).
• Mineral oil / Lubricating oil (\(\text{C}_{20}-\text{C}_{30}\)): Lubricants for machinery.
• Fuel oil (\(\text{C}_{30}-\text{C}_{40}\)): Fuel for ships and industrial furnaces.
• Bitumen / Residue (\(>\text{C}_{40}\)): Surfacing roads and roofing.

Key Takeaway: Fractional distillation separates crude oil based on differences in boiling points. Shorter molecules condense near the cool top; longer molecules condense near the hot bottom.

4. Cracking of Alkanes

Why is Cracking Necessary?

Fractional distillation produces large amounts of long-chain fractions (like fuel oil and bitumen), but the market demand for these is low. Conversely, the market demand for short-chain fractions (like petrol) and alkenes (to make plastics) is extremely high.

Cracking solves this economic imbalance by breaking long-chain, less useful alkanes into shorter, more valuable alkanes and alkenes.

Comparison of Cracking Methods

1. Thermal Cracking:
• Conditions: High temperature (\(450\text{ }^\circ\text{C}\) to \(900\text{ }^\circ\text{C}\)) and high pressure (up to \(70\text{ atm}\)).
• Main Products: High percentage of alkenes (e.g., ethene, used to make poly(ethene)).

2. Catalytic Cracking:
• Conditions: Moderate temperature (around \(450\text{ }^\circ\text{C}\)), slight pressure, and a zeolite catalyst (an aluminosilicate mineral with a honeycomb structure for large surface area).
• Main Products: Branched alkanes, cycloalkanes, and aromatic hydrocarbons. These burn much more smoothly in car engines and boost fuel performance.

Example Cracking Equation

\(\text{C}_{10}\text{H}_{22} \rightarrow \text{C}_8\text{H}_{18} + \text{C}_2\text{H}_4\)
Notice that the total number of carbons and hydrogens is equal on both sides!

Key Takeaway: Cracking converts surplus long-chain alkanes into high-demand shorter alkanes and alkenes. Catalytic cracking uses a zeolite catalyst at \(450\text{ }^\circ\text{C}\).

5. Combustion of Alkanes and Environmental Chemistry

Alkanes are generally chemically unreactive because they have strong, non-polar \(\text{C}-\text{C}\) and \(\text{C}-\text{H}\) bonds. However, they readily undergo combustion (burning in oxygen) in highly exothermic reactions, making them excellent fuels.

1. Complete Combustion

Occurs when there is an excess of oxygen. The only products are carbon dioxide and water vapour.

\(\text{CH}_4 + 2\text{O}_2 \rightarrow \text{CO}_2 + 2\text{H}_2\text{O}\)
\(\text{C}_3\text{H}_8 + 5\text{O}_2 \rightarrow 3\text{CO}_2 + 4\text{H}_2\text{O}\)

2. Incomplete Combustion

Occurs when oxygen supply is limited. Products include water vapour along with carbon monoxide (\(\text{CO}\)) and/or solid carbon particulates (soot, \(\text{C}\)).

\(\text{CH}_4 + \frac{3}{2}\text{O}_2 \rightarrow \text{CO} + 2\text{H}_2\text{O}\)
\(\text{CH}_4 + \text{O}_2 \rightarrow \text{C} + 2\text{H}_2\text{O}\)

Environmental Pollutants from Internal Combustion Engines

• Carbon Monoxide (\(\text{CO}\)): A colourless, odourless, and highly toxic gas. It binds irreversibly to haemoglobin in red blood cells, preventing oxygen transport in the blood.
• Carbon Particulates / Soot (\(\text{C}\)): Cause global dimming and exacerbate human respiratory problems like asthma.
• Unburned Hydrocarbons (\(\text{C}_x\text{H}_y\)): Greenhouse gases that react with nitrogen oxides in sunlight to produce photochemical smog.
• Nitrogen Oxides (\(\text{NO}\) and \(\text{NO}_2\), collectively \(\text{NO}_x\)): The extremely high temperature and electric spark inside a car engine cause atmospheric nitrogen and oxygen to react:
\(\text{N}_2 + \text{O}_2 \rightarrow 2\text{NO}\)
\(\text{NO}\) then reacts further with oxygen to form \(\text{NO}_2\), which dissolves in water to form acid rain and triggers photochemical smog.
• Sulfur Dioxide (\(\text{SO}_2\)): Formed when sulfur impurities in fuels burn (\(\text{S} + \text{O}_2 \rightarrow \text{SO}_2\)). It reacts with water in the atmosphere to form sulfurous/sulfuric acid, causing acid rain.

Catalytic Converters

Modern cars are fitted with catalytic converters to remove pollutants from exhaust gases.

• Structure: A ceramic honeycomb structure coated with a thin layer of precious metals (platinum (\(\text{Pt}\)), palladium (\(\text{Pd}\)), and rhodium (\(\text{Rh}\))) to provide a maximum surface area.
• Key Reactions:
1. Oxidation of carbon monoxide and reduction of nitrogen monoxide:
\(2\text{CO} + 2\text{NO} \rightarrow 2\text{CO}_2 + \text{N}_2\)
2. Removal of unburned hydrocarbons:
\(\text{C}_8\text{H}_{18} + 25\text{NO} \rightarrow 8\text{CO}_2 + 9\text{H}_2\text{O} + 12.5\text{N}_2\)

Key Takeaway: Complete combustion yields \(\text{CO}_2\) and \(\text{H}_2\text{O}\); incomplete combustion produces toxic \(\text{CO}\) and soot. Catalytic converters turn harmful \(\text{CO}\), \(\text{NO}_x\), and hydrocarbons into safer \(\text{CO}_2\), \(\text{N}_2\), and \(\text{H}_2\text{O}\).

6. Halogenation of Alkanes: Free Radical Substitution

Alkanes react with halogens (such as chlorine and bromine) in a reaction called free radical substitution. This reaction requires ultraviolet (UV) light (or high temperatures) to get started.

Key Definitions

• Free Radical: A highly reactive species with an unpaired electron, represented by a dot (e.g., \(\text{Cl}^\bullet\), \({}^\bullet\text{CH}_3\)).
• Homolytic Fission: The breaking of a covalent bond where each bonded atom takes one of the shared electrons, forming two free radicals.
• Substitution: A reaction in which an atom or group of atoms is replaced by another atom or group of atoms.

The 3-Step Reaction Mechanism: Chlorination of Methane

Step 1: Initiation (Radicals are Formed)

UV light provides the energy to break the \(\text{Cl}-\text{Cl}\) single bond by photochemical homolytic fission.

\(\text{Cl}_2 \xrightarrow{\text{UV light}} 2\text{Cl}^\bullet\)

Note: The \(\text{C}-\text{H}\) bond in methane does not break in initiation because the \(\text{Cl}-\text{Cl}\) bond is weaker and absorbs UV radiation easily.

Step 2: Propagation (Chain Reaction)

Radicals are used up and newly regenerated in a continuous cycle.

• Propagation Step 1: The chlorine radical removes a hydrogen atom from methane, producing a methyl radical and hydrogen chloride.
\(\text{CH}_4 + \text{Cl}^\bullet \rightarrow {}^\bullet\text{CH}_3 + \text{HCl}\)

• Propagation Step 2: The reactive methyl radical attacks an unreacted chlorine molecule to produce chloromethane and regenerate a chlorine radical.
\({}^\bullet\text{CH}_3 + \text{Cl}_2 \rightarrow \text{CH}_3\text{Cl} + \text{Cl}^\bullet\)

Common Student Mistake: Never write \(\text{CH}_4 + \text{Cl}^\bullet \rightarrow \text{CH}_3\text{Cl} + \text{H}^\bullet\)! A free hydrogen radical (\(\text{H}^\bullet\)) is energetically unfavorable and is never formed in this mechanism.

Step 3: Termination (Radicals are Removed)

The chain reaction stops when any two free radicals collide and combine to form a stable covalent bond.

• Combination of two chlorine radicals:
\(\text{Cl}^\bullet + \text{Cl}^\bullet \rightarrow \text{Cl}_2\)
• Combination of a methyl radical and a chlorine radical:
\({}^\bullet\text{CH}_3 + \text{Cl}^\bullet \rightarrow \text{CH}_3\text{Cl}\)
• Combination of two methyl radicals:
\({}^\bullet\text{CH}_3 + {}^\bullet\text{CH}_3 \rightarrow \text{C}_2\text{H}_6\)

Did you know? The formation of trace amounts of ethane (\(\text{C}_2\text{H}_6\)) is the experimental proof that methyl radicals are formed during the reaction!

Limitations of Free Radical Substitution in Organic Synthesis

Free radical substitution is generally poor for preparing pure halogenoalkanes because it gives a complex mixture of products due to two main reasons:

1. Further Substitution (Poly-substitution):
Once chloromethane (\(\text{CH}_3\text{Cl}\)) forms, chlorine radicals can attack it further:
\(\text{CH}_3\text{Cl} \rightarrow \text{CH}_2\text{Cl}_2 \text{ (dichloromethane)} \rightarrow \text{CHCl}_3 \text{ (trichloromethane)} \rightarrow \text{CCl}_4 \text{ (tetrachloromethane)}\)
• How to minimise this: Use a large excess of methane so chlorine radicals are much more likely to collide with methane than with chloromethane.
• How to maximise tetrachloromethane: Use a large excess of chlorine.

2. Isomer Formation on Longer Chains:
With longer alkanes (such as propane or butane), substitution can occur at different carbon positions along the chain, producing a mixture of position isomers (e.g., 1-chloropropane and 2-chloropropane).

Key Takeaway: Free radical substitution proceeds in three stages: Initiation (homolytic fission by UV light), Propagation (chain reaction making products and regenerating radicals), and Termination (radicals combining). It produces mixtures due to poly-substitution and isomer formation.

Quick Review Summary Checklist

• General formula: \(\text{C}_n\text{H}_{2n+2}\) for alkanes; \(\text{C}_n\text{H}_{2n}\) for cycloalkanes.
• Bonding & Shape: \(\sigma\)-bonds, tetrahedral shape around carbon, \(109.5^\circ\) bond angles.
• Boiling points: Increase with chain length (more electrons, stronger van der Waals forces); decrease with branching (less surface contact area).
• Fractional Distillation: Physical separation based on boiling points.
• Catalytic Cracking: Converts long chains to short alkanes/alkenes using a zeolite catalyst at \(450\text{ }^\circ\text{C}\).
• Complete vs Incomplete Combustion: Complete gives \(\text{CO}_2 + \text{H}_2\text{O}\); incomplete gives toxic \(\text{CO}\) and soot \(\text{C}\).
• Free Radical Substitution: Requires UV light; occurs via Initiation \(\rightarrow\) Propagation \(\rightarrow\) Termination.