1The idea in one paragraph

A fuel is a substance that releases useful energy when it reacts, and nearly all the fuels we use release it by reacting with oxygen. When there is plenty of oxygen, the carbon in an organic fuel ends up as carbon dioxide and the hydrogen as water; when oxygen runs short, some carbon ends up as poisonous carbon monoxide or as soot. Every fuel is a trade-off: how much energy it gives per gram, how much carbon dioxide it adds to the air for that energy, how cleanly it burns, whether it will run out, and what it costs to produce. Coal, oil and gas carry carbon locked away for millions of years, and burning them is raising the CO₂ in the atmosphere and with it the greenhouse effect. Biofuels recycle carbon taken from the air recently. A fuel cell skips burning altogether and turns the fuel's chemical energy straight into electricity.

2Combustion: what burns, and how to write the equation

Combustion is a reaction with oxygen that releases heat, usually with a flame. The guide names three families of substances that burn when heated in oxygen.

Reactive metals form the metal oxide. Magnesium burns with a dazzling white light and leaves a white powder:

2Mg(s) + O₂(g) → 2MgO(s)

Calcium burns in the same way with a brick-red flame to give CaO, and iron wool burns in a shower of sparks to give iron oxides.

Non-metals form non-metal oxides, usually gases:

• C(s) + O₂(g) → CO₂(g)
• S(s) + O₂(g) → SO₂(g), the choking gas behind acid rain when sulfur-containing coal is burned
• 2H₂(g) + O₂(g) → 2H₂O(l), which produces no carbon dioxide at all

Organic compounds, such as hydrocarbons and alcohols, give carbon dioxide and water when the oxygen supply is plentiful. This is complete combustion.

Every combustion is also a redox reaction, which Reactivity 3.2 develops: the fuel is the reducing agent and is oxidized; oxygen is the oxidizing agent and is reduced. And every useful fuel has a high activation energy (Reactivity 2.2): petrol and methane do not burst into flame at room temperature, which is exactly why they can be stored safely until a spark starts them.

Balancing an organic combustion, in a fixed order. Carbon first, then hydrogen, then oxygen last, because oxygen appears in both products and is the only element that is free to adjust.

Worked example 1: octane, C₈H₁₈.

Both final lines are correct. Halves are allowed in combustion equations, and an enthalpy of combustion is always per mole of fuel, so the version with one C₈H₁₈ is the one that goes with ΔHc.

Worked example 2: propan-1-ol, C₃H₇OH. An alcohol already contains one oxygen atom, and that is the only difference: subtract it before you count what O₂ must supply.

The trap is the hydrogen count. Propan-1-ol has seven hydrogens in C₃H₇ and one more in OH: eight in all, so four waters. Write the formula out as C₃H₈O if it helps.

For any hydrocarbon CₓHᵧ the pattern is CₓHᵧ + (x + y/4)O₂ → xCO₂ + (y/2)H₂O. Knowing it is a check, not a substitute for the working.

3Incomplete combustion

When there is not enough oxygen for every carbon atom to reach CO₂, some carbon stops part-way. Incomplete combustion produces carbon monoxide, CO, or carbon itself as soot, alongside water. In these equations the hydrogen still ends up as water. Figure 1 shows the three outcomes for methane.

Write incomplete combustion equations with the same order as before, with the carbon going to CO or C instead of CO₂:

Real incomplete combustion usually gives a mixture of CO₂, CO and C. An exam question will tell you which product to aim for; give exactly that one.

What you see. Close the air hole of a Bunsen burner and the methane burns in a limited supply of oxygen. The flame turns yellow and luminous, because tiny particles of carbon glow in it, and it becomes smoky. Hold a cold evaporating basin in it and the underside is coated black with soot. With the air hole open, the flame is blue, almost invisible, and much hotter.

Why it matters for health. Carbon monoxide is colourless and has no smell, so it cannot be detected without an alarm. It binds to the haemoglobin in red blood cells far more strongly than oxygen does, so the blood carries less oxygen, and high levels are fatal. This is why faulty gas boilers kill. Soot is made of particulates: tiny solid particles that lodge in the lungs and are linked to breathing and heart disease. Incomplete combustion also wastes fuel, because a carbon atom that stops at CO or C has released less energy than one that reaches CO₂.

Why bigger molecules burn less cleanly. Figure 2 shows how much oxygen each alkane molecule needs for complete combustion.

A methane molecule needs 2 O₂ molecules; an octane molecule needs 12½. The more oxygen molecules one fuel molecule needs, the more likely it is that not enough of them meet it at the moment it burns. Larger hydrocarbons are also liquids or solids, which must evaporate and mix with air before they can burn, while methane is already a gas mixed with air. So larger hydrocarbons have a greater tendency to burn incompletely: coal and heavy oil fractions smoke; natural gas rarely does.

4Fossil fuels: coal, crude oil and natural gas

Fossil fuels formed from the remains of living things buried and changed by heat and pressure over millions of years. Coal formed mainly from land plants. Crude oil and natural gas formed mainly from tiny marine organisms such as plankton. They are non-renewable: they formed over millions of years and we are using them in centuries, so for human purposes the supply is finite.

The guide asks you to judge them on two particular points, energy per unit mass and the tendency to burn incompletely, and to evaluate the carbon dioxide each adds to the air.

Energy per unit mass, the specific energy, is the heat released per gram of fuel. You get it from the enthalpy of combustion and the molar mass:

specific energy (kJ g⁻¹) = |ΔHc| (kJ mol⁻¹) ÷ M (g mol⁻¹)

CO₂ per unit of energy is the fairer climate comparison, because what we want from a fuel is energy, not mass. Divide the mass of CO₂ one mole of fuel produces by the energy that mole releases.

Worked example 3: comparing three fuels. Take methane for natural gas, octane for petrol from crude oil, and pure carbon as a simple model of coal (real coal also contains hydrogen, water, sulfur and ash, so real values differ). ΔHc in kJ mol⁻¹: CH₄ −890, C₈H₁₈ −5470, C −393.5. M(CO₂) = 44.01 g mol⁻¹.

(1 MJ = 1000 kJ, so dividing by 0.890 MJ is the same as multiplying by 1000 ÷ 890.) Figure 3 puts these beside two alternatives.

Coal releases more than twice as much CO₂ as natural gas for the same heat. The reason is structure: methane carries four hydrogen atoms for every carbon, and burning hydrogen gives energy without CO₂; coal is almost all carbon. The more hydrogen a fuel has per carbon, the less CO₂ it releases per unit of energy. Hydrogen itself sits at the end of that line with none at all, at least where it is burned.

Putting it together:

The guide's nature of science line asks about the environmental, economic, ethical and social sides of burning fossil fuels. A good answer names at least one of each: climate change and air pollution (environmental); cheap, reliable energy that economies are built on, and jobs in mining regions (economic); the people most affected by climate change are often those who burned the least (ethical); energy security and the health of people living near coal plants (social).

5Carbon dioxide and the greenhouse effect

The Earth's surface is warmed by sunlight and cools by radiating energy back out. The greenhouse effect is the trapping of some of that outgoing energy by gases in the atmosphere. Figure 4 draws it.

Tell it in four steps, because each step is a marking point.

1. Radiation from the Sun, mostly visible and ultraviolet (shorter wavelength), passes through the atmosphere and is absorbed by the surface, which warms.
2. The warm surface emits energy as infrared (longer-wavelength) radiation.
3. Greenhouse gases such as CO₂, H₂O and CH₄ absorb some of that infrared: the energy makes their bonds vibrate.
4. They re-emit the energy in all directions, so part of it returns to the surface instead of escaping to space. The lower atmosphere and the surface are warmer than they would otherwise be.

The greenhouse effect itself is natural and keeps the planet habitable. The problem is its enhancement. Burning fossil fuels adds carbon dioxide that had been out of circulation for millions of years. Carbon dioxide in the atmosphere was about 280 parts per million before industrialisation; direct measurement shows it passed 420 parts per million in the 2020s. More CO₂ means more infrared absorbed and returned, and so a rise in average surface temperature: global warming, with its effects on climate, sea level and ecosystems.

Why does CO₂ absorb infrared when N₂ and O₂, which make up 99% of the air, do not? A molecule absorbs infrared only if the vibration changes its dipole. N₂ and O₂ are symmetrical diatomic molecules whose single stretch can never create a dipole; CO₂ has bending and asymmetric stretching vibrations that can. HL students meet this again with infrared spectroscopy in Structure 3.2. The global warming potential of a gas compares its warming effect with that of the same mass of CO₂ over a set time. It depends on how strongly the gas absorbs infrared and on how long it stays in the atmosphere, which is why methane, a stronger absorber, counts for much more than its small concentration suggests.

6Biofuels, renewable and non-renewable

A renewable energy source is replenished by natural processes as fast as, or faster than, we use it: sunlight, wind, and crops that regrow each season. A non-renewable source is used far faster than it forms, so it will run out: coal, oil, natural gas, and uranium.

A biofuel is a fuel made from recently living material. It is renewable because its carbon was fixed from the air by photosynthesis over a short period, a single growing season rather than millions of years.

Photosynthesis: 6CO₂(g) + 6H₂O(l) → C₆H₁₂O₆(aq) + 6O₂(g), driven by light energy absorbed by chlorophyll

Photosynthesis is endothermic: the energy stored in glucose came from sunlight. Burning a biofuel releases that stored energy. The main biofuels:

• Bioethanol, made by fermenting sugars from crops such as sugar cane or maize with yeast: C₆H₁₂O₆(aq) → 2C₂H₅OH(aq) + 2CO₂(g). It is blended into petrol.
• Biodiesel, made from vegetable oils or waste cooking oil.
• Biogas, mostly methane, from the anaerobic breakdown of manure, sewage or food waste.

Figure 5 shows why a biofuel's carbon counts differently.

The CO₂ a biofuel releases was taken from the air by the crop a few months or years earlier. In principle the loop is closed and the fuel is carbon neutral. In practice it is not quite: farm machinery, fertiliser production, transport and distillation mostly run on fossil fuels, so there is a net carbon cost, and it can be large if forest is cleared to grow the crop.

7Fuel cells

In a power station, a fuel burns, the heat boils water, steam turns a turbine and the turbine drives a generator. Energy is lost at every one of those conversions. A fuel cell converts the chemical energy of a fuel directly into electrical energy, with no burning. It does this by splitting the redox reaction of combustion into two halves at two electrodes, and making the electrons travel between them through a wire.

The anode is where oxidation happens: the fuel loses electrons. The cathode is where reduction happens: oxygen gains electrons. The electrons flow through the external circuit from anode to cathode, and that flow is the current. Ions move through the electrolyte between the electrodes to complete the circuit. Figure 6 shows a hydrogen fuel cell with an acidic electrolyte.

The overall reaction is the same as burning hydrogen. The difference is where the energy goes: into electricity, not into heat.

Deducing any fuel cell half-equation in acid follows four steps: balance the atoms that are not O or H; balance O by adding H₂O; balance H by adding H⁺; balance charge by adding electrons. Figure 7 shows the methanol cell, and the working below builds its anode equation.

Worked example 4: the methanol anode.

The cathode is the same as before, multiplied to take six electrons: 1½O₂ + 6H⁺ + 6e⁻ → 3H₂O. Adding the two and cancelling the 6H⁺, the 6e⁻ and one water gives CH₃OH + 1½O₂ → CO₂ + 2H₂O, which is just the combustion of methanol. Always check your half-equations by adding them: the result should be the combustion equation.

If a question uses an alkaline electrolyte, OH⁻ ions carry the charge instead. For hydrogen the half-equations become H₂ + 2OH⁻ → 2H₂O + 2e⁻ at the anode and O₂ + 2H₂O + 4e⁻ → 4OH⁻ at the cathode. The overall reaction is unchanged.

Why use a fuel cell? It turns a larger share of the fuel's energy into electricity than burning it in a power station does, because it skips the heat engine. A hydrogen cell's only product is water. It is quiet and has no moving parts. The difficulties are elsewhere: most hydrogen today is made from natural gas, which releases CO₂; hydrogen is hard to store because it is a gas of very low density; and the electrodes need expensive catalysts such as platinum. Methanol is a liquid and far easier to store, but its cell does release CO₂.

A fuel cell differs from a primary cell such as an ordinary battery (Reactivity 3.2) in one way that matters: a primary cell contains a fixed amount of reactants and is finished when they are used up, while a fuel cell is supplied with fresh fuel and oxygen from outside and keeps working as long as they flow.