Combustion Reactions and Fuel Burning

A combustion reaction is a rapid exothermic reaction between a fuel and an oxidizer, usually oxygen, that releases heat and light; complete combustion of a hydrocarbon gives carbon dioxide and water, while incomplete combustion with limited oxygen also produces carbon monoxide and soot. In brief: propane burns as C₃H₈ + 5O₂ → 3CO₂ + 4H₂O, releasing 2,220 kJ per mole, and burning 1.00 kg of octane (a main component of gasoline) produces about 3.08 kg of carbon dioxide.

Combustion supplies most of the world’s energy for transport, heating and electricity. Its products also affect the air and climate. This article explains how to write and balance combustion equations for different fuels, compares complete and incomplete combustion, lists the energy content of common fuels in tables, works through calculations of oxygen needed, carbon dioxide produced and energy released, and gives the history of the study of burning.

How do you write combustion equations?

In complete combustion, every element of the fuel forms its stable oxide: carbon becomes CO₂, hydrogen becomes H₂O, sulfur becomes SO₂ (and nitrogen in the fuel gives NO or N₂). To balance, follow the order carbon, hydrogen, oxygen.

For an alkane C_nH_2n+2 the general equation is C_nH_2n+2 + (3n + 1) ÷ 2 O₂ → n CO₂ + (n + 1) H₂O. For methane (n = 1): CH₄ + 2O₂ → CO₂ + 2H₂O. For octane (n = 8): C₈H₁₈ + 12.5O₂ → 8CO₂ + 9H₂O, or 2C₈H₁₈ + 25O₂ → 16CO₂ + 18H₂O. For an alcohol, such as ethanol: C₂H₅OH + 3O₂ → 2CO₂ + 3H₂O. The oxygen in the fuel is counted in the oxygen balance.

What is incomplete combustion?

When the oxygen supply is limited, carbon is not fully oxidized. The products can be carbon monoxide (CO), carbon (soot) and water.

• 2CH₄ + 3O₂ → 2CO + 4H₂O
• CH₄ + O₂ → C + 2H₂O

In practice a mixture of CO₂, CO and soot forms. Incomplete combustion releases less energy per mole of fuel, produces smoke and gives off toxic carbon monoxide. A yellow, sooty flame shows insufficient air, while a blue flame indicates good mixing.

How much energy do common fuels release?

Worked examples

1. Carbon dioxide from octane. How much CO₂ forms when 1.00 kg of octane burns completely? n(C₈H₁₈) = 1,000 ÷ 114.23 = 8.754 mol. From the equation, 8 mol of CO₂ forms per mole of octane: 8.754 × 8 = 70.03 mol. The mass is 70.03 × 44.01 = 3,082 g = 3.08 kg.
2. Oxygen needed. The same combustion needs 12.5 mol of O₂ per mole of octane: 8.754 × 12.5 = 109.4 mol, with a mass of 109.4 × 32.00 = 3,501 g = 3.50 kg.
3. Volume of air. At 25 °C and 1 atm, 109.4 mol of O₂ occupies 109.4 × 24.5 = 2,681 L. Air is about 21% oxygen by volume, so the volume of air is 2,681 ÷ 0.21 = 12,800 L, or 12.8 m³.
4. Energy released. 8.754 mol × 5,470 kJ/mol = 47,900 kJ = 47.9 MJ, in agreement with the table (47.9 MJ/kg).
5. Mass balance check. Reactants: 1,000 g octane + 3,501 g O₂ = 4,501 g. Products: 3,082 g CO₂ + water (8.754 × 9 × 18.015 = 1,419 g) = 4,501 g. ✓
6. Carbon dioxide per unit of energy. Methane: 44.01 g CO₂ per 890 kJ = 49.4 g/MJ. Octane: 8 × 44.01 = 352 g per 5,470 kJ = 64.4 g/MJ. Hydrogen produces no CO₂ at the point of use. Per megajoule, natural gas produces about 23% less carbon dioxide than gasoline: (64.4 − 49.4) ÷ 64.4 = 23%.
7. Balance an alcohol. Propan-1-ol, C₃H₇OH + O₂ → CO₂ + H₂O: C: 3CO₂; H: 8 atoms, so 4H₂O; O on the right: 6 + 4 = 10; the fuel supplies 1, so O₂ provides 9, which is 4.5O₂. Double: 2C₃H₇OH + 9O₂ → 6CO₂ + 8H₂O.

What are the products of burning fuels?

• Carbon dioxide: a greenhouse gas that is the main product of burning carbon-containing fuels.
• Water vapor: visible as condensation from exhausts.
• Carbon monoxide and soot (particulates): from incomplete combustion.
• Sulfur dioxide: from sulfur impurities in coal and oil, which causes acid rain.
• Nitrogen oxides: formed from nitrogen and oxygen of the air at high temperatures in engines (N₂ + O₂ → 2NO), which contribute to smog.
• Unburned hydrocarbons: when combustion is incomplete.

What does a fire need?

Combustion requires fuel, oxygen and heat to reach the ignition temperature. The minimum temperature at which a liquid forms enough vapor to ignite is its flash point, which is about −42 °C for gasoline, so its vapor can ignite at winter temperatures. Once started, the heat released keeps the reaction going. Fires are controlled by removing one of the three parts.

How did the understanding of burning develop?

In the 17th and 18th centuries, burning was explained by the phlogiston theory, which assumed that a combustible substance contains a fire-like component that is released during burning. In the 1770s, Antoine Lavoisier showed through careful weighing that burning is a combination with oxygen, and that substances gain mass when they burn in a closed vessel. Later, Marcellin Berthelot and others measured the heats of combustion, and the study of combustion became the basis of thermochemistry. The internal combustion engine of the 19th century made the combustion of liquid fuels a main source of energy for transport.

Which mistakes are common?

Mistake 1: Balancing oxygen first

Balance carbon and hydrogen first, and then oxygen.

Mistake 2: Forgetting oxygen in the fuel

Ethanol and glucose contain oxygen, which reduces the amount of O₂ needed.

Mistake 3: Confusing mass and moles for CO₂ output

More mass of CO₂ is produced than the mass of fuel burned, because the carbon combines with oxygen from the air: 1.00 kg of octane gives 3.08 kg of CO₂.

Mistake 4: Burning fuels indoors without ventilation

Incomplete combustion produces carbon monoxide, which is colorless and odorless and can be fatal.