I suggest the longer the hydro-carbon chain of the fuel being burnt the greater the energy released per gram of fuel burnt. However given the price of fuel per gram, typically, does not increase proportionately to the chain length, the fuel releasing the most energy may not necessarily be the most cost effective. Method Apparatus: Copper beaker/can Ceramic crucible (with lid) Frame and clamp Heat Proof Mats Thermometer Balance Tin foil Small lengths of string (to act as wicks) Measuring cylinder Distilled water Selection of FuelsFor this comparison -Methanol C H3 OH -Ethanol C2 H5 OH.
-Propanol C3 H7 OH -Butanol C4 H9 OH and -Candle Wax Procedure: 1. Start by setting measuring out 200ml of water, transfer to the copper beaker and take the water temperature. 2. Now clamp the beaker to the frame and adjust so that the bottom of the beaker is at a height exactly 4cm from the top of the crucible (when the crucible is directly under it). 3. Remove the crucible from its position, carefully pipette approximately 1g of methanol into it and replace the lid (this is to reduce any loss through evaporation). 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 Temperature Change Fuel Initial mass /g Final mass /g Difference /g Initial Temp. / i?? C Final temp. /
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/gi?? C) x Temp. change(i?? C) If the mass of water heated (as with our procedure) is 200g, we give the SHC of water as 4. 2j/gi?? C (this is a universally accepted approximation) and the change in temperature was 100i?? 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. Fuel Temp. change /i?? C Energy Transferred /J Energy transferred /KJ Candle Wax 11 9240 9. 24 Methanol 24 20160 20. 16 Ethanol 26 21840 21. 84 Propanol 57 47880 47. 88 Butanol 41 34440 34. 44 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: Fuel Fuel Burnt /g Energy transferred /KJ Energy transferred per gram (KJ/g) .
Butanol 2. 028 34. 44 16. 982 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 Fuel Energy Transferred (KJ/g) Cost per gram /p Energy Transferred (p/Kj) Pence per kilojoule.
Propanol 12. 83 1. 28 0. 100 10. 02 Butanol 16. 98 1. 3 0. 077 13. 06 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 (?H) and of course ? H 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.
Andrew McNally/ 11th February 2002 /Science Access..