You can calculate an experimental value for the enthalpy change of combustion for a fuel very simply – this is a required practical so you need to be able to describe both the method and how you would use the data in your calculations.
The standard enthalpy of combustion, ΔcH⦵, is the enthalpy change when 1 mole of a substance reacts completely with oxygen under standard conditions, with reactants and products being in their standard states at 298K.
Fuels are often compared in terms of their energy density (kJ g-1 or kJ kg-1) which is easily calculated – all we need to do is work out how many moles of the substance there are in 1kg / 1000g and then multiply up the energy.
e.g. ΔcH⦵ (H2(g)) = -286 kJ mol-1 , Mr = 2 g mol-1
No. of mol of H2 in 1 kg = 1000g / 2 g mol-1 = 500 mol
Energy density of H2 (kJ kg-1) = -286 x 500 = 143000 kJ kg-1
We can measure the ΔcH⦵ for fuels in the lab simply (if inaccurately) using a spirit burner, a balance, a known volume of water and a thermometer.
e.g. calculate ΔcH⦵ for methanol from the following experimental data:
mass of spirit burner before lighting = 189.3g
mass of spirit burner after capping = 187.9g
initial temperature of water = 17.5 °C
final maximum temperature of water = 39.0 °C
- Heat energy is transferred to the water:
- So, if we burn 1.4g methanol we release 13480.5 J heat energy
- No. of mol of methanol combusted = 1.4g / 32 g mol-1 = 0.0438 mol
- And finally, we can scale up to find the the energy transferred per mol of methanol
The official value for ΔcH⦵ for methanol is – 726 kJ mol-1. Clearly this is not a very good experiment!
Sources of error:
- Heat loss to the surroundings which comprise the air and the beaker itself (consider using draught shields to reduce this)
- Incomplete combustion of the fuel (evidence for this is the vast amount of soot that collects on the beaker)
- Evaporation of fuel from the wick after the flame is extinguished
- Non-standard conditions compared with the definition
The ultimate solution is to use a bomb or flame calorimeter which is designed to minimise all of these errors.
Practice questions
Assume that the specific heat capacity of water is 4.18 J g-1 K-1 and the density of water is 1.0 g cm-3 for all questions.
- (a) Combustion of 1.20g hexane raises the temperature of 100g of water from 20.0°C to 37.5°C. Calculate a value for the enthalpy change of combustion of hexane in kJ mol-1.
(b) Suggest two reasons why your value is different to that in the data book.
- Plot a graph of ΔcH⦵ against molar mass for the alkanes in the table below. Use your graph to estimate the energy released when 1.60g of pentane is completely combusted in oxygen, showing all relevant working.
| Alkane | Methane | Ethane | Propane | Butane |
| ΔcH⦵ / kJ mol-1 | -890 | -1560 | -2219 | -2877 |
- A lighter filled with butane gas was used to heat a small beaker containing 50cm3 water. The lighter was weighed before and after the experiment and shown to have lost 0.45g in mass. The temperature of the water increased from 18°C to 69°C.
(a) Calculate the enthalpy change of combustion for butane using this data.
(b) Use the information in the table below to estimate the ‘data book’ value of ΔcH for butane.
| Alkane | ΔcH⦵ / kJ mol-1 |
| Methane | -890 |
| Ethane | -1560 |
| Propane | -2220 |
| Pentane | -3509 |
| Hexane | -4194 |
(c) Suggest two reasons, other than heat losses, why the value calculated from experimental data is less exothermic than the value estimated in part (b).
(d) Suggest a modification that could be made to this experiment to improve its accuracy.
(e) Suggest why it is not possible to measure the enthalpy of formation of butane directly by experiment.
Answers
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