Enthalpy of combustion

The standard enthalpy of combustion, ΔcH⦵, is the enthalpy change for burning one mole of a substance completely in oxygen, under standard conditions and with everything in its standard state. Combustion always releases energy, so ΔcH⦵ is always negative: for methane it is −890.8 kJ mol⁻¹, and for ethanol −1366.8 kJ mol⁻¹.

Enthalpy of combustion definition

AQA's A-level specification asks you to define the standard enthalpy of combustion. A complete definition has four parts:

• The enthalpy change (heat energy change at constant pressure)
• when one mole of a substance
• is completely burned in oxygen
• under standard conditions, with all reactants and products in their standard states.

AQA defines standard conditions as 100 kPa and a stated temperature, usually 298 K (25 °C). "Standard states" means the normal physical state at those conditions, so the water made is liquid, not steam.

The equation must show exactly one mole of the substance burned, even if that means fractions in front of oxygen:

CH₃OH(l) + 1½O₂(g) → CO₂(g) + 2H₂O(l) ΔcH⦵ = −726.1 kJ mol⁻¹

Enthalpy of combustion values

Values from OpenStax Chemistry, Table 5.2. Notice that the enthalpy of combustion of hydrogen is the same reaction as the enthalpy of formation of water, and the enthalpy of combustion of carbon is the enthalpy of formation of carbon dioxide.

The trend in a homologous series

Going from methanol to ethanol adds one CH₂ group, and ΔcH⦵ becomes 1366.8 − 726.1 = 640.7 kJ mol⁻¹ more negative. Each extra CH₂ means one more CO₂ and one more H₂O formed, so each step along a homologous series adds a similar amount.

How to measure the enthalpy of combustion

Measuring an enthalpy change is AQA A-level required practical 2. For a liquid fuel such as an alcohol, a simple calorimeter works:

1. Put a measured mass of water (often 100 to 200 g) in a copper can, clamped above a spirit burner, and record its starting temperature.
2. Weigh the spirit burner with its fuel and lid.
3. Light the burner and stir the water, until its temperature has risen by about 15 to 20 °C.
4. Put out the flame, record the highest temperature reached, and reweigh the burner straight away.

Then use q = mcΔT for the heat gained by the water, and divide by the moles of fuel burned. The enthalpy change formula guide covers q = mcΔT in detail.

Worked example: methanol

Burning 0.64 g of methanol (Mr = 32.0) heats 200 g of water from 19.0 °C to 34.5 °C. Take c = 4.18 J g⁻¹ K⁻¹ for water, and find ΔcH of methanol.

1. Heat gained by the water: q = mcΔT = 200 × 4.18 × 15.5 = 12 958 J = 12.958 kJ.
2. Moles of methanol: n = 0.64 ÷ 32.0 = 0.0200 mol.
3. ΔcH = −q ÷ n = −12.958 ÷ 0.0200 = −648 kJ mol⁻¹ (3 s.f.). The sign is negative because the reaction released the heat.

Why experimental values are too low

The methanol result, −648 kJ mol⁻¹, is only about 89% of the data value, −726.1 kJ mol⁻¹. Results from simple calorimeters always come out less negative than data book values, because:

• Heat is lost to the surrounding air and used to warm the can and clamp, so the water gains less than the fuel released. This is usually the biggest error.
• Incomplete combustion turns some of the carbon into soot or carbon monoxide, which releases less energy than making CO₂.
• Fuel evaporates from the wick, so the mass loss includes fuel that didn't burn.
• Non-standard conditions: the water produced is a gas at the flame, not liquid as in the standard definition.

Improvements include a draught shield, insulating the can, a lid on the calorimeter, keeping the flame close to the can, and putting the lid on the burner as soon as the flame is out. A bomb calorimeter, which burns the fuel in oxygen inside a sealed, insulated container, loses far less heat.

Using enthalpies of combustion with Hess's law

AQA asks you to calculate enthalpy changes for reactions from enthalpies of combustion. This is especially useful for enthalpies of formation that can't be measured directly. For example, carbon and hydrogen don't react cleanly to make methane, but each of them, and methane, can be burned.

For C(s) + 2H₂(g) → CH₄(g): ΔfH = [−393.5 + 2 × (−285.8)] − (−890.8) = −965.1 + 890.8 = −74.3 kJ mol⁻¹. Note the order: reactants minus products, the opposite way round to the formation enthalpy version. The enthalpy calculator does these calculations with working.

Common mistakes

• Giving a positive value. Combustion is exothermic, so ΔcH is always negative.
• Using the mass of fuel in q = mcΔT. m is the mass of the water being heated.
• Forgetting to convert J to kJ before dividing by moles to get kJ mol⁻¹.
• Balancing with 2 moles of fuel. The definition is for one mole, so use fractions of O₂ if needed.
• Leaving out "complete" or "standard". Both are part of the definition.

To practise energetics in a short timed paper, try today's A-level Chemistry: Year 1 paper.

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