Enthalpy of combustion from a calorimetry experiment

A-Level (AQA) · unreviewed

Fuel burned under a beaker of water; temperature rise and mass of fuel used are known.

Worked example

Burning 0.82 g of ethanol raised the temperature of 204 g of water by 20.7 K. Using c = 4.18 J g⁻¹ K⁻¹ and M(ethanol) = 46.0 g mol⁻¹, calculate an experimental value for ΔcH of ethanol.

1. Heat from specific heat capacity: solve for Heat absorbed by the water
$q = mc\Delta T$
= 17650 J
2. Moles from mass: solve for Moles of ethanol burned
$n = \dfrac{m}{M}$
= 0.01783 mol
3. Divide Heat absorbed by the water ÷ Moles of ethanol burned → Heat released per mole
= 990.2 kJ mol⁻¹
4. Multiply by a coefficient: -1 × Heat released per mole → ΔcH (negative: exothermic)
= −990.2 kJ mol⁻¹

Answer: ΔcH = −990 kJ mol⁻¹

Method

q = mcΔT uses the mass of water. Divide by the moles of ethanol burned (n = m/M) to get the heat released per mole, then give it a negative sign because combustion is exothermic. The value is much less negative than the data-book −1367 kJ mol⁻¹, mainly because of heat lost to the surroundings and incomplete combustion.

Common mistakes

Combustion is exothermic, so ΔcH is negative.
In q = mcΔT, m is the mass of the water being heated, not the mass of fuel burned.

Formulas used

Heat from specific heat capacity
$q = mc\Delta T$
Moles from mass
$n = \dfrac{m}{M}$

Practise this with fresh numbers

Enthalpy of combustion from a calorimetry experiment

Worked example

Method

Common mistakes

Formulas used

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