4.Tare the balance and record the mass of the crucible, its lid and the contents.
5.Next place the crucible directly under the copper beaker, surround with heatproof mats (as to provide a draught shield) and again (to make sure no heat has transferred from the copper beaker to the water) take the water temperature.
6.Remove the lid from the crucible, ignite the contents, surround the open side with a heatproof mat and allow the fuel to burn away. Stir the water and record the highest temperature reached.
7.Finally, reweigh the crucible (including the lid) and record the final mass.
8.Repeat the procedure with the other three alcohols remembering to:
- use fresh water each time
- allow the copper beaker to cool
- stir before recording the temperature (temperature change of the water is required not the beaker)
- place the lid on the crucible as soon as the fuel is loaded and only removing when ready to ignite
For the candle the same procedure applies, however it is impracticable to have the candle burn until all the fuel is expended so record the mass of the candle (including any surrounding vessel), re-gauge the height of the copper beaker so it is 4cm from the top of the candle (the candle maybe, and is likely, to have different dimensions to the crucible) and ignite. As before surrounding the apparatus with heatproof mats to provide a draught shield. Once an appreciable temperature rise is seen extinguish the candle, stir the water and take the temperature. Finally, re-weigh the candle.
Results
The above results only show the temperature change and not the amount of energy transferred to the water. To do this the following formula need be applied:
J = Joules (unit of energy)
SHC = specific heat capacity
Energy transferred to water (J) = Mass of water(g) x SHC(J/g˚C) x Temp. change(˚C)
If the mass of water heated (as with our procedure) is 200g, we give the SHC of water as 4.2j/g˚C (this is a universally accepted approximation) and the change in temperature was 100˚C:
J = 200 x 4.2 x 100
.`. J = 84000
Or as kilojoules (1/1000 of a joule) 84KJ
The following table of results has been prepared using the above formulae to calculate the amount of energy transferred from the fuel to the water. The only item that has changed from the above example is the temperature change, for this I have substituted the results obtained from the test.
From this position however we cannot see which of the fuels is most efficient. To do that we need to standardise the results so all fuels are compared equally. At the moment the results are showing the candle as being far insuperior to the propanol however only 0.265g of wax was burnt as opposed to 3.732g of propanol, that’s over 14 times as much.
The following set of results show the amount of energy transferred per gram:
The results so far have shown the total energy released and also the energy released per gram of fuel burnt. Given the following costs per gram, we can also see which is the most cost effective:
Costs per gram
Candle Wax → 0.07p
Methanol → 0.44p
Ethanol → 0.33p
Propanol → 1.28p
Butanol → 1.30p
Conclusion
It is clear from the results that the candle transferred the most energy to the water per gram of fuel burnt. Being the least expensive of the fuels also meant that it is clearly the most economical. However the final temperature of the water when heated by the candle was cool when compared to the other fuels meaning, in a practical sense at least, although it is the most economical it is not the most convenient if wishing to heat things to high temperatures.
Discussion
All the fuels burnt share the same key characteristics in that they are all hydrocarbons; basically they are made up of hydrogen and carbon. They are all also stable at room temperatures (and above) and require the presence of oxygen in order for them to burn. Not only do they need oxygen but also an initial input of energy, in this case a match. This initial energy is called the activation energy. Once the initial energy has gone in to start the reaction there is a short period of endothermic activity. At this point the bonds between the hydrogen and carbon are being broken. The oxygen then reacts with the hydrogen to make water and with the carbon to produce carbon monoxide and then again to form carbon dioxide. All the while these bonds are being formed energy is being released and the reaction is said to be exothermic. In order for a reaction to be exothermic overall, the amount of endothermic energy (energy being taken/bonds being broken) cannot be greater than the amount of exothermic energy (energy being released/bonds being formed). The difference between the two is said to be the enthalpy change (ΔH) and of course ΔH can be either plus or minus depending on whether or not the reaction is overall endothermic or exothermic respectively. In terms of the fuels burnt there is clearly a release of energy, we saw the temperature of the water rise and we could see the release of energy in the form of light from the flame.
Although our results were consistent and reasonably reliable there were still a number of sources of error. Were the experiment to be completed again there are a number of things I would change in order to make it a fairer test. Firstly, whilst burning the fuels it was noticed energy was lost from around the sides of the apparatus, if we could feel this energy as heat it was not contributing to the transfer of heat to the water. The levels of loss could not be measured and could have been greater with some fuels tested than others, distorting the final results should one have experienced greater loss than another. This, of course, given the apparatus available was unavoidable. The only way around this I can see would be to have a sealed chamber between fuel and water, limiting substantially any loss of energy.
