IB DP Chemistry Reactivity 1.3: Energy from fuels

This guide teaches Reactivity 1.3, Energy from fuels, from the IB Diploma Programme Chemistry guide published in February 2023, with first exams in 2025. It covers every understanding in the sub-topic, Reactivity 1.3.1 to 1.3.5. The guide lists no additional higher level content for Reactivity 1.3, so all of it is for both SL and HL.

When you have read it, try the practice questions, or use the shorter revision notes for a final read. See where this unit sits on the IB DP chemistry hub, and tick it off on the printable checklist. Our free diagnostics can show which topics need work first.

What this unit covers

Guide reference	Skills expected	Level
Reactivity 1.3.1	Write balanced equations for complete combustion of metals, non-metals, hydrocarbons and alcohols	SL and HL
Reactivity 1.3.2	Write balanced equations for incomplete combustion of hydrocarbons and alcohols, giving carbon monoxide or carbon	SL and HL
Reactivity 1.3.3	Compare coal, crude oil and natural gas; work out how much CO₂ different fuels add; link CO₂ to the greenhouse effect	SL and HL
Reactivity 1.3.4	Explain how biofuels fix carbon through photosynthesis; tell renewable from non-renewable sources; weigh up biofuels	SL and HL
Reactivity 1.3.5	Write the electrode half-equations for hydrogen and methanol fuel cells	SL and HL

Reactivity 1.3.1 Complete combustion

Combustion is the reaction of a substance with oxygen, started by heating, that releases energy.

Metals form metal oxides:

• 2Mg(s) + O₂(g) → 2MgO(s)
• 4Li(s) + O₂(g) → 2Li₂O(s)

Non-metals form non-metal oxides:

• S(s) + O₂(g) → SO₂(g)
• P₄(s) + 5O₂(g) → P₄O₁₀(s)
• 2H₂(g) + O₂(g) → 2H₂O(l)

Hydrocarbons and alcohols burn completely in excess oxygen to give carbon dioxide and water only. In every combustion, oxygen is the oxidizing agent and the fuel is the reducing agent, which links this topic to Reactivity 3.2.

A fuel’s high activation energy is useful: it does not react with air until lit, so it can be stored safely. A spark starts the reaction and the heat released keeps it going.

A balancing rule for organic fuels

For a fuel with x carbon, y hydrogen and z oxygen atoms per molecule, burning completely:

moles of O₂ needed per mole of fuel = x + y/4 − z/2, giving x CO₂ and y/2 H₂O

This is just the balancing order: carbon, then hydrogen, then oxygen, subtracting any oxygen already in the fuel.

Worked example 1: butan-1-ol

Butan-1-ol, CH₃CH₂CH₂CH₂OH, has the molecular formula C₄H₁₀O.

1. Carbon: 4 C atoms give 4CO₂.
2. Hydrogen: 10 H atoms give 5H₂O.
3. Oxygen atoms on the right: (4 × 2) + 5 = 13.
4. The fuel brings 1 oxygen atom, so 12 must come from O₂, which is 6O₂.

C₄H₉OH(l) + 6O₂(g) → 4CO₂(g) + 5H₂O(l)

Worked example 2: hexane, where a half appears

Hexane, C₆H₁₄, gives 6CO₂ and 7H₂O. The right side has 12 + 7 = 19 oxygen atoms, so you need 9½ O₂. Either leave the fraction or double every coefficient:

2C₆H₁₄(l) + 19O₂(g) → 12CO₂(g) + 14H₂O(l)

Both forms are balanced; the fraction form is useful when you want the equation per mole of fuel.

Reactivity 1.3.2 Incomplete combustion

When the supply of oxygen is limited, organic fuels do not fully oxidise. The carbon ends up as carbon monoxide, CO, or carbon, C (soot). Water still forms from the hydrogen. Real flames give a mixture, but you write a separate equation for each carbon product.

Burning methane with the Bunsen burner air hole closed gives a yellow, smoky flame that leaves soot on a cold surface. With plenty of air, the flame is blue and clean.

Why it matters:

• Carbon monoxide is toxic. It attaches to haemoglobin in red blood cells more tightly than O₂, so the blood carries less oxygen. You cannot see or smell it.
• Soot particles irritate the lungs and blacken buildings and engines.
• Less energy is released per mole of fuel.

Worked example 3: nonane in limited oxygen

Nonane, C₉H₂₀, gives 10H₂O in every case.

Forming carbon monoxide: 9CO and 10H₂O contain 9 + 10 = 19 O atoms, so you need 9½ O₂.

2C₉H₂₀(l) + 19O₂(g) → 18CO(g) + 20H₂O(l)

Forming carbon: only the water contains oxygen, 10 O atoms, so you need 5O₂.

C₉H₂₀(l) + 5O₂(g) → 9C(s) + 10H₂O(l)

Complete combustion would need 14 O₂ per C₉H₂₀.

For an alcohol, remember that the fuel supplies one oxygen atom. Propan-1-ol, C₃H₇OH, forming CO: 3CO + 4H₂O has 7 O atoms, minus 1 in the fuel, so 6 O atoms or 3O₂:

C₃H₇OH(l) + 3O₂(g) → 3CO(g) + 4H₂O(l)

Reactivity 1.3.3 Fossil fuels and carbon dioxide

Fossil fuels formed from the remains of organisms buried and compressed over millions of years. The three main ones are coal (mostly carbon), crude oil (a mixture of hydrocarbons, separated into fractions such as petrol, kerosene and diesel) and natural gas (mostly methane).

Fuel	Advantages	Disadvantages
Coal	Large reserves; cheap; a solid that is easy to stockpile	Most CO₂ per unit of energy; often contains sulfur, giving SO₂; soot and ash; mining is dangerous
Crude oil	Liquid fractions are easy to transport and pump; high energy per unit mass; also a raw material for plastics and other chemicals	Finite; spills damage ecosystems; larger molecules in diesel and fuel oil tend to burn incompletely
Natural gas	Most energy per unit mass; least CO₂ per unit of energy; burns cleanly with little soot	Needs pipelines or pressurised storage; leaks are explosive; unburned methane is itself a greenhouse gas

Larger molecules burn less completely

Bigger hydrocarbons burn incompletely more easily. Each molecule needs more O₂ to burn fully: methane needs 2 O₂, octane 12½. They have a higher percentage of carbon by mass, and stronger London forces make them less volatile, so they mix less well with air. Candle wax and diesel are often sooty; methane is not.

Worked example 4: comparing fuels

Use these values for the enthalpy of combustion: methane −890 kJ mol⁻¹, octane −5470 kJ mol⁻¹, carbon (as a model for coal) −394 kJ mol⁻¹. Take Ar(C) = 12.01, Ar(H) = 1.01 and Ar(O) = 16.00.

Energy per gram = energy per mole ÷ molar mass.

• Methane: 890 ÷ 16.05 = 55.5 kJ g⁻¹
• Octane: 5470 ÷ 114.26 = 47.9 kJ g⁻¹
• Carbon: 394 ÷ 12.01 = 32.8 kJ g⁻¹

Mass of CO₂ per MJ (1000 kJ) released. Each mole of carbon in the fuel gives one mole of CO₂ (44.01 g).

• Methane: 1 mol CO₂ per 890 kJ, so 44.01 × 1000 ÷ 890 = 49.4 g per MJ
• Octane: 8 mol CO₂ per 5470 kJ, so 8 × 44.01 × 1000 ÷ 5470 = 64.4 g per MJ
• Carbon: 44.01 × 1000 ÷ 394 = 112 g per MJ

For the same energy, coal adds more than twice as much CO₂ as natural gas. Fuels with more hydrogen per carbon atom get more of their energy from forming water.

Carbon dioxide and the greenhouse effect

Sunlight reaching the ground, mainly visible and near-infrared radiation, is absorbed; the warmed ground then gives off longer-wavelength infrared radiation. Molecules such as CO₂, H₂O and CH₄ absorb some of this infrared, which makes their bonds vibrate. They then re-emit the energy in all directions, including back towards the surface. Without this, the planet’s surface would be much colder; the process is the greenhouse effect.

Burning fossil fuels releases carbon locked away for millions of years. NOAA reports that atmospheric CO₂ was about 280 ppm before the Industrial Revolution and averaged about 423 ppm in 2024. More CO₂ absorbs and re-emits more infrared, which enhances the greenhouse effect and raises global average temperatures.

Reactivity 1.3.4 Biofuels

A renewable energy source is replaced by natural processes at least as fast as we use it, such as sunlight, wind and biofuels. A non-renewable source is used far faster than it forms, so its supply is finite. Fossil fuels are non-renewable.

Biofuels are fuels made from living matter, such as ethanol from sugar cane or maize, biodiesel from plant oils, and biogas (mostly methane) from rotting waste. The carbon in them was fixed from the air by photosynthesis within the last few months or years:

6CO₂(g) + 6H₂O(l) → C₆H₁₂O₆(aq) + 6O₂(g)

Reactants: carbon dioxide and water. Products: glucose and oxygen. Light supplies the energy.

Worked example 5: carbon bookkeeping for bioethanol

Plants fix 6 mol CO₂ to make 1 mol glucose. Fermenting that glucose to 2 mol ethanol releases 2 mol CO₂ (C₆H₁₂O₆ → 2C₂H₅OH + 2CO₂). Burning 2 mol ethanol releases 2 × 2 = 4 mol CO₂. Total released: 2 + 4 = 6 mol, the same as was fixed.

On paper the cycle is carbon neutral. In reality it is not, because fossil fuels are burned to run farm machinery, make fertiliser, distil the ethanol and transport it.

Advantages of biofuels	Disadvantages of biofuels
Renewable: a new crop can be grown each year	Land used for fuel crops cannot grow food, which can raise food prices
Lower net CO₂ than fossil fuels, as the carbon was recently fixed	Forest may be cleared for crops, releasing stored carbon and harming biodiversity
Can be blended with petrol or diesel and used in existing engines	Need large amounts of water and fertiliser
Can make use of waste, such as biogas from food or farm waste	Ethanol releases less energy per gram than petrol

Reactivity 1.3.5 Fuel cells

A fuel cell turns the chemical energy of a fuel straight into electrical energy. Fuel and oxygen are fed in continuously but kept apart, so electrons travel through an external circuit. The fuel is oxidised at the anode, which is the negative electrode; oxygen is reduced at the cathode, the positive electrode. Unlike a primary (voltaic) cell, which contains a fixed amount of reactants and stops when they are used up, a fuel cell works for as long as fuel is supplied.

Hydrogen fuel cell

With an acidic electrolyte:

• Anode: H₂(g) → 2H⁺(aq) + 2e⁻
• Cathode: O₂(g) + 4H⁺(aq) + 4e⁻ → 2H₂O(l)

With an alkaline electrolyte:

• Anode: H₂(g) + 2OH⁻(aq) → 2H₂O(l) + 2e⁻
• Cathode: O₂(g) + 2H₂O(l) + 4e⁻ → 4OH⁻(aq)

Either way the overall reaction is 2H₂(g) + O₂(g) → 2H₂O(l). The only product is water.

Worked example 6: methanol fuel cell (acidic electrolyte)

Build the anode half-equation step by step. Methanol is oxidised to carbon dioxide.

1. Carbon is balanced: CH₃OH → CO₂
2. Add H₂O to balance oxygen: CH₃OH + H₂O → CO₂
3. Add H⁺ to balance hydrogen (4 + 2 = 6 H on the left): CH₃OH + H₂O → CO₂ + 6H⁺
4. Add electrons to balance charge (6+ on the right): CH₃OH(aq) + H₂O(l) → CO₂(g) + 6H⁺(aq) + 6e⁻

Cathode, scaled to accept 6 electrons: 3/2 O₂(g) + 6H⁺(aq) + 6e⁻ → 3H₂O(l)

Adding them and cancelling 6H⁺, 6e⁻ and one H₂O gives CH₃OH + 3/2 O₂ → CO₂ + 2H₂O, the same overall change as burning methanol. Liquid methanol is easier to store than hydrogen, but the cell releases CO₂.

You do not need the details of the membrane that separates the electrodes.

Common errors

• Forgetting the oxygen atom already in an alcohol when counting O₂.
• Including CO₂ in an equation meant to form carbon monoxide only.
• Saying CO₂ “reflects” heat or “traps” it, instead of absorbing and re-emitting infrared radiation.
• Calling biofuels carbon neutral without mentioning the energy used to grow and process them.
• Putting the electrons on the wrong side: the anode releases them and the cathode takes them in.

Official syllabus

Source: Chemistry guide, International Baccalaureate Diploma Programme, February 2023 (first assessment 2025). Taught here: Reactivity 1.3, Energy from fuels, understandings 1.3.1 to 1.3.5, common to SL and HL.