Organic Chemistry – Fuels (Cambridge IGCSE 0620 – Topic 11)
Learning Objectives
- Identify methane as the main constituent of natural gas.
- Recall and use the basic terminology for organic compounds (homologous series, functional groups, saturated/unsaturated).
- Write and interpret molecular, structural and displayed (line‑angle) formulae for the first three alkanes.
- Apply IUPAC/common naming rules to simple alkanes.
- Describe the three major fuel types required by the syllabus (coal, natural gas, petroleum) and the main fractions obtained from petroleum.
- State the general formulae for alkanes and alkenes, write balanced combustion equations and recognise a typical substitution reaction.
- Explain how cracking produces alkenes and how the bromine‑water test is used to identify unsaturation.
- Apply the concepts of states of matter, atomic structure, bonding and stoichiometry to fuel‑related calculations.
- Understand key safety information for handling gaseous fuels.
1. States of Matter (Syllabus 1.1)
- Solid – particles are closely packed in a regular arrangement; only vibrate in place.
- Liquid – particles are close but can move past one another; have a definite volume but no fixed shape.
- Gas – particles are far apart and move freely; both volume and shape are indefinite.
Diffusion: the spreading of particles from an area of higher concentration to lower concentration. Rate of diffusion increases with higher temperature and lower molecular mass (e.g., methane diffuses faster than propane).
Kinetic Particle Theory (KPT) – explains pressure, temperature and volume relationships:
- Pressure results from collisions of particles with the walls of a container.
- Temperature is a measure of the average kinetic energy of particles.
- PV = nRT (ideal‑gas equation) is useful for calculations involving gaseous fuels.
2. Atoms, Elements & Compounds (Syllabus 2.1–2.5)
2.1 Periodic Table & Group‑Number/Charge Relationship
- Groups 1, 2, 13–18 have predictable valence‑electron counts and typical ionic charges (e.g., Group 1 → +1, Group 17 → –1).
- Transition metals (Groups 3–12) form coloured compounds and variable oxidation states.
2.2 Isotopes
- Atoms of the same element with different numbers of neutrons.
- Relative atomic mass (Ar) is the weighted average of isotopic masses (e.g., carbon: ¹²C ≈ 98.9 %, ¹³C ≈ 1.1 %).
2.3 Ions
- Cations – loss of electrons (e.g., Na⁺, Ca²⁺).
- Anions – gain of electrons (e.g., Cl⁻, SO₄²⁻).
2.4 Bonding Types
- Ionic bonding – transfer of electrons; formation of a lattice of oppositely charged ions (e.g., NaCl). High melting points, soluble in water, conduct electricity when molten or aqueous.
- Covalent (molecular) bonding – sharing of electrons; discrete molecules (e.g., CH₄, H₂O). Low melting/boiling points, poor conductors.
- Metallic bonding – delocalised electrons in a lattice of metal cations; explains conductivity, malleability, and high melting points of metals.
2.5 Simple & Giant Covalent Structures
- Simple covalent molecules – discrete entities (e.g., CO₂, CH₄). Low boiling points.
- Giant covalent structures – extensive networks (e.g., diamond, graphite, SiO₂). Very high melting points, hardness, and (in graphite) electrical conductivity.
3. Stoichiometry (Syllabus 3.1–3.4)
3.1 Mole Concept & Avogadro’s Number
- 1 mol = 6.02 × 10²³ particles.
- Molar mass (M) = atomic/molecular mass in g mol⁻¹ (e.g., MCH₄ = 16.04 g mol⁻¹).
3.2 Calculations Using Mass, Moles and Molecules
General steps:
- Convert mass ↔ moles using M.
- Use the balanced chemical equation to relate mole ratios.
- Convert back to mass or volume (for gases use PV = nRT).
3.3 Example – Combustion of Methane
CH₄ + 2 O₂ → CO₂ + 2 H₂O
- From 16 g CH₄ (1 mol) we need 2 mol O₂ (≈64 g) and obtain 1 mol CO₂ (44 g) + 2 mol H₂O (36 g).
- At STP, 1 mol gas occupies 22.4 dm³, so 1 mol CH₄ gives 22.4 dm³ of CH₄ and 2 mol O₂ gives 44.8 dm³ of O₂.
3.4 Concentrations (Supplement)
- Mass concentration: g dm⁻³.
- Molar concentration (M): mol dm⁻³.
- Useful for preparing solutions of acids, bases or gases dissolved in water.
4. Key Terminology (Syllabus 11.1)
- Homologous series: Compounds that differ by the repeating unit –CH₂– and show regular trends in physical properties.
- Functional group: Specific atom group that determines characteristic reactions (e.g., –OH in alcohols, –COOH in carboxylic acids).
- Saturated hydrocarbon (alkane): Only single C–C bonds; general formula
CₙH₂ₙ₊₂.
- Unsaturated hydrocarbon (alkene): At least one C=C double bond; general formula
CₙH₂ₙ.
- Structural formula: Shows how atoms are bonded (e.g., H–C–H).
- Displayed (line‑angle) formula: Simplified drawing for larger molecules.
5. The First Three Alkanes (Naming & Formulae – Syllabus 11.2 & 11.4)
| Compound |
Molecular formula |
Structural formula |
Displayed (line‑angle) formula |
IUPAC / Common name |
| Methane |
CH4 |
H–C–H (four H atoms around C) |
⋮C⋮ |
Methane |
| Ethane |
C2H6 |
H₃C–CH₃ |
CH₃–CH₃ |
Ethane |
| Propane |
C3H8 |
CH₃–CH₂–CH₃ |
CH₃–CH₂–CH₃ |
Propane |
General alkane formula: CₙH₂ₙ₊₂ (n = number of carbon atoms).
Combustion of Alkanes
Complete combustion produces carbon dioxide and water vapour:
$$\text{C}_n\text{H}_{2n+2} + \left(\frac{3n+1}{2}\right)\text{O}_2 \;\longrightarrow\; n\text{CO}_2 + (n+1)\text{H}_2\text{O}$$
Example – methane:
$$\text{CH}_4 + 2\text{O}_2 \;\longrightarrow\; \text{CO}_2 + 2\text{H}_2\text{O}$$
Typical Substitution Reaction (Photochemical Chlorination)
$$\text{CH}_4 + \text{Cl}_2 \xrightarrow{\text{hv}} \text{CH}_3\text{Cl} + \text{HCl}$$
Illustrates a radical substitution of an alkane by chlorine.
6. Fuels Required by the Syllabus (Syllabus 11.3)
6.1 Overview of the Three Main Fuel Types
- Coal – solid fuel, mainly carbon with variable amounts of hydrogen, sulfur and nitrogen.
- Natural gas – gaseous fuel; the principal component is methane (CH₄).
- Petroleum (crude oil) – complex mixture of hydrocarbons separated by fractional distillation into useful fractions.
6.2 Natural Gas – Focus on Methane
- Colourless, odourless (odourants such as mercaptan are added for safety).
- Principal component: methane (CH₄).
- Key physical data
- Molecular weight: 16.04 g mol⁻¹
- Boiling point: –161.5 °C
- Density relative to air: 0.55 (lighter than air)
- Energy content: ≈55 MJ kg⁻¹
- Explosive range in air: 5 %–15 % (v/v)
- Combustion (see Section 5):
CH₄ + 2 O₂ → CO₂ + 2 H₂O
6.3 Petroleum – Fractional Distillation & Major Fractions
Crude oil is heated; vapour rises through a fractionating column. As temperature falls, components condense at characteristic boiling‑point ranges.
| Fraction (boiling‑point range) |
Typical carbon number (approx.) |
Common uses |
| Gases (C₁–C₄) |
1–4 |
Cooking gas, LPG, petrochemical feedstock |
| Naphtha (C₅–C₁₀) |
5–10 |
Ethylene, propylene production; solvent |
| Gasoline (petrol) (C₆–C₁₁) |
6–11 |
Motor‑vehicle fuel |
| Kerosene (C₁₂–C₁₅) |
12–15 |
Aviation fuel, domestic heating |
| Diesel (C₁₅–C₂₀) |
15–20 |
Road & marine engines |
| Fuel oil (C₂₀–C₃₀) |
20–30 |
Industrial boilers, power stations |
| Lubricating oil (C₃₀–C₄₀) |
30–40 |
Machinery lubrication |
| Bitumen (≥C₄₀) |
40+ |
Road surfacing, roofing |
Trend across the column: Boiling point, density and viscosity increase as carbon chain length increases; vapour pressure decreases.
7. Alkenes – Unsaturated Hydrocarbons (Syllabus 11.5)
- General formula:
CₙH₂ₙ (one C=C double bond).
- Typical example: ethene (ethylene), C₂H₄.
- Physical properties: lower boiling points than the corresponding alkanes; the π‑bond makes them chemically more reactive.
7.1 Cracking – Production of Alkenes
Thermal or catalytic cracking breaks a larger alkane into a smaller alkane and an alkene (plus hydrogen). Example:
$$\text{C}_{10}\text{H}_{22} \;\xrightarrow{\text{heat}} \; \text{C}_8\text{H}_{18} + \text{C}_2\text{H}_4$$
Cracking supplies the alkenes needed for the petrochemical industry and for gasoline‑blending.
7.2 Bromine‑Water Test for Unsaturation
Alkenes decolourise bromine water because Br₂ adds across the C=C double bond:
$$\text{C}_n\text{H}_{2n} + \text{Br}_2 \;\longrightarrow\; \text{C}_n\text{H}_{2n}\text{Br}_2$$
Observation: orange‑brown bromine solution becomes colourless → indicates an unsaturated hydrocarbon.
8. Comparison of Common Fuels
| Fuel |
Physical state |
Main component(s) |
Typical energy content (MJ kg⁻¹) |
| Coal |
Solid |
Carbon (C) with variable H, S, N |
24–30 |
| Natural gas |
Gas |
Methane (CH₄) |
≈55 |
| Petroleum – gasoline fraction |
Liquid |
Mixture of C₇–C₁₁ alkanes/alkenes |
≈44 |
| Petroleum – diesel fraction |
Liquid |
Mixture of C₁₀–C₂₀ alkanes |
≈45 |
9. Safety Note (AO2)
- Methane is highly flammable; explosive limits in air: 5 %–15 % (v/v).
- Domestic natural‑gas supplies are odourised (usually with a mercaptan) to enable leak detection.
- Always work in a well‑ventilated area, use approved gas detectors and keep open flames away from suspected leaks.
- Store cylinders upright, secure them, and never handle with bare hands if the valve is cold (risk of frostbite).
10. Suggested Diagrams (placeholders for teachers)
- Structural (tetrahedral) diagram of methane.
- Fractionating column showing temperature gradient and position of each petroleum fraction.
- Reaction scheme for the bromine‑water test (Br₂ + ethene → 1,2‑dibromoethane).
- Ideal‑gas PV = nRT diagram illustrating how volume changes with temperature for a fixed amount of natural gas.
- Dot‑and‑cross diagram of a C–C single bond (alkane) vs. a C=C double bond (alkene).
11. Quick Revision Questions (AO1–AO3)
- Write the molecular formula, structural formula and IUPAC name for the main constituent of natural gas.
- Balance the complete combustion equations for (a) methane and (b) propane.
- Explain why methane dominates natural‑gas composition (give at least three reasons, e.g., formation conditions, stability, low boiling point).
- State the general formulae for alkanes and alkenes and give one example of each with its structural formula.
- Describe the process of fractional distillation of crude oil and name three useful fractions obtained.
- Write the balanced equation for the thermal cracking of decane to produce ethene and a smaller alkane.
- Predict the result of adding bromine water to a sample of ethene and to a sample of propane. Explain the observations.
- Identify two safety precautions that should be taken when working with natural gas in the laboratory.
- Using the ideal‑gas equation, calculate the volume (at STP) of methane produced when 8.0 g of CH₄ is completely combusted.
- Calculate the mass of oxygen required to combust 2.0 g of propane completely.