Input variables:
- Types of alcohol burnt.
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Mass of water can be changed, as long as it is recorded.
- Temperature rise can be slightly different, as long as it is recorded. It would otherwise be hard to get exactly the same rise in temperature for each experiment.
Output variables:
- The energy (i.e. temperature rise) per mole of alcohol produced for the 3 different alcohols.
Sources of error:
- Ignored the rise in temperature and therefore the heat absorbed by the copper vessel/calorimeter. The rise in temperature of the copper vessel means that energy has gone towards heating the copper can instead of heating the water. Energy transfer is not 100% efficient, but this is not taken into account.
- The combustion of alcohol is not entirely efficient as carbon was deposited on the bottom of the copper can.
- Heat losses to the surroundings. Heat energy is lost to the surroundings, to the air instead of heating the water. This is known because heat can be felt surrounding the copper can.
· Type of beaker, copper because copper conducts heat efficiently, unlike glass.
· Temperature rise of about 30ºC because this is easy to measure, ensuring that the water temperature doesn’t rise so much that it evaporates. A temperature of about 50oC is easy to handle, so that the can doesn’t have to left for a long period of time before it can be touched.
· Surrounding temperature of around 23ºC to ensure that the rise in temperature is due only to the heat given off from the alcohol combustion.
· The distance from the wick to the base of the can is adjusted so that all of the flame impinges on the base. See diagram.
· Same set of scales, scales can be tuned slightly differently, be more or less accurate. Using the same scales ensures that the results will all be measured with equal accuracy.
· The spirit burner was weighed with the lid on so that the alcohol doesn’t evaporate. Alcohols are very volatile and would quickly evaporate into the air. Between the time that the alcohol is weighed and the time the experiment is repeated, a considerable amount of alcohol would have evaporated making the experiment unreliable.
Equipment:
- Copper can/calorimeter – good conductor, will not retain a lot of heat, but will conduct it to the water making the heat transfer more efficient, therefore making the experiment fairer.
- Thermometer to measure starting and finishing temperature of the water
- Stirrer- to stir the water ensuring that there are no hot spots in the water that would alter the reading on the thermometer.
- Draught preventer (large piece of cardboard) so that heat from the alcohol does not
- Small, flat piece of thin cardboard with a hole in the middle, to insulate heat and keep water from evaporating. The hole should be big enough so that a thermometer and a stirrer can fit through it.
- Bunsen burner and splinter (or matches), so that you have a close source of fire to begin you experiment.
- Stand and clamp, so that the height of the copper vessel can be adjusted to suit the size of the alcohol flame.
- Spirit (ethanol, propanol and butanol) burner.
- Scales
- Water
Method:
- The experiment was set up as shown above, so that the alcohol flame just impinges on the copper calorimeter. This allowed the maximum amount of heat to reach the calorimeter.
- The copper calorimeter was weighed and the result recorded.
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Water was poured into the calorimeter so it was about 2/3 full. The can was weighed with the water and the result recorded.
- The mass of the calorimeter + water was taken away from the mass of the calorimeter to find the mass of water. Result recorded.
- The flame from the alcohol burner must just impinge on the calorimeter (the height of the calorimeter was adjusted if needed).
- The alcohol burner was weighed and result recorded.
- The temperature of water was recorded at room temperature/ before experiment with the thermometer.
- The alcohol burner was placed under the calorimeter and set alight.
- The water was stirred (this evened out the water temperature in the can, so that there were no hot spots).
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When the water temperature reached about 50oC, or had risen by about 30oC, the flame was put out, and the lid was quickly replaced. Alcohol is very volatile, and would quickly evaporate without the lid.
- The temperature of water was recorded
- The rise in temperature was calculated
- The weight of the alcohol burner was recorded
- The decrease in mass was calculated.
- The calorimeter was washed out, making sure the carbon deposited on the bottom was washed off.
- The calorimeter was weighed again, making sure it is still the same weight. If the weight was more/ less, the new weight was recorded, taking it into account when working out the mass of water.
- Experiment was then repeated 5 times for each alcohol.
R e s u l t s
Using the relationship: heat = temperature rise x mass water x 4.18
No. of moles = mass/ RFM
ETHANOL:
Method of calculating results obtained:
1.) 30.0 x 54.2 x 4.18 = 6796.68 joules energy produced by 0.6g ethanol
no. of moles of ethanol = mass/RFM
= 0.6 / 46 = 0.013
As 0.013 moles of ethanol liberate 6796.68 joules energy
Therefore 1 mole of ethanol liberates 6796.68 / 0.013 = 522821.5 joules
or 522.8 kJmol-1
2.) 26.5 x 54.2 x 4.18 = 6003.734 joules energy produced by 0.45g ethanol
no. of moles of ethanol = mass/RFM
=0.45 / 46 = 0.00978
As 0.00978 moles of ethanol liberate 6003.734 joules energy
Therefore 1 mole of ethanol liberates 6003.734 / 0.00978 = 613878.7 joules
or 613.9 kJmol-1
simply from comparing this result to the other results, I can see that it is anomalous. However, I would need to draw a graph to compare how abnormal the result is. However, comparing the result with the literary values, this result may be the closest to the true values (least amount of heat lost to the atmosphere).
3.) 26.3 x 54.2 x 4.18 = 5958.4228 joules energy produced by 0.53g ethanol
no. of moles of ethanol = mass/RFM
= 0.53 / 46 = 0.0115
As 0.0115 moles of ethanol liberate 5958.4228 joules energy
Therefore 1 mole of ethanol liberates 5958.4228 / 0.0115 = 518123.7 joules
or 518.1 kJmol-1
4.) 26.0 x 54.2 x 4.18 = 5890.456 joules energy produced by 0.52g ethanol
no. of moles of ethanol = mass/RFM
= 0.52 / 46 = 0.0113
As 0.0113 moles of ethanol liberate 5890.456 joules energy
Therefore 1 mole of ethanol liberates 5890.456 / 0.0113 = 521279.3 joules
or 521.3 kJmol-1
5.) 27.0 x 54.2 x 4.18 = 6117.012 joules energy produced by 0.55g ethanol
no. of moles of ethanol = mass/RFM
= 0.55 / 46 = 0.01196
As 0.01196 moles of ethanol liberate 6117.012 joules energy
Therefore 1 mole of ethanol liberates 6117.012 / 0.01196 = 511455.85 joules
or 511.5 kJmol-1
Average values = Σ ΔHc / number of values
I will not include the result of the second experiment (613.9) into this calculation because I think that it is an anomalous result.
Average value for ethanol (excluding 2nd value) = 518.425
PROPANOL:
Average values = ΣΔHc / number of values
None of my results appear anomalous at the moment, so all will be included in calculation.
Average value for propanol = 844.811
BUTANOL:
Average values = ΣΔHc / number of values
Average value for butanol = 1337.9864
Analysis
Ethanol:
The experimental values obtained are reasonably constant i.e. average +- . Suggesting that the technique was followed accurately, apart from the 2nd experiment, where the higher value suggests that some part of the procedure had been varied unconsciously.
The average value obtained was less than half the literature value suggesting that the experiment design was not suitable for measuring accurate values.
Propanol:
The experimental values obtained were not very satisfactory, as they produced a large number of anomalous results. This suggests that the technique used was varied throughout the experiment.
The average value obtained was less than half of the literature value, suggesting that the experiment design was not suitable for measuring accurate values.
Butanol:
The experimental values obtained appeared quite constant, apart from 2 results that were slightly lower. This suggests that the technique used was varied during the 1st and last experiment.
The average value obtained was about half of the literature value, suggesting that the experiment design was not suitable for measuring accurate values.
As the values obtained lie on a straight line correlating with the prediction made, this experiment confirms the prediction made even though the degree of exothermicity determined was less than what was expected from the literature values.
It is not surprising that a simple experimental set up such as used here, would give erroneous values as no special precautions were taken to ensure that no heat or alcohol was lost.
Qualitatively, the data shows that all the alcohols burnt exothermically and that the larger the molecule, the more exothermic the reaction. However the discrepancy between the values found and the literature values means it would be unwise to pay too much attention to any quantitative conclusion that could be drawn from this experiment.
During a reaction (when substances react to form new a substance(s)), the bonds holding the substances together are broken, and then new bonds are formed. When bonds are broken energy is taken in (i.e. it takes energy to break the bonds), and when bonds are made, energy is given out. In this case, there is more energy given out than taken in.
Discrepancies are probably due to:
- inadequate prevention of heat losses- heat lost to surroundings.
- ignoring the effect of using a copper calorimeter (copper is heated up, meaning that all heat didn’t go towards heating the water.
- incomplete combustion (carbon deposited on base).
I found that as the molecule size of the alcohol became greater, the amount of carbon deposited on the base of the copper calorimeter also increased. This showed that the combustion of the alcohols was not efficient.
Evaluation
Comment on accuracy/ reliability:
The results obtained, if compared to the literature values, are by no means accurate. However, it is important to consider that with the equipment available, I could not have made the experiment completely efficient in heat transfer. However, the results obtained were all similar for each alcohol, excluding the small number of anomalous results. This means that the experiment was reliable, but the values were inaccurate. From the results obtained it is not possible to calculate the heat produced by each alcohol as the results are not accurate enough, but it is possible to compare the alcohols. By comparing the average results, it is possible to see their relationships with each other, and to determine a trend (i.e. that the larger the molecule of alcohol, the more heat is produced- butanol produces more heat than ethanol or propanol).
Anomalous results:
Anomalous results could be disregarded if the experiment is repeated several times, so that any anomalous result is made obvious as it is unexplainably different from all other results. Anomalous results are often due to experimental mistakes being made unconsciously. If the result is closer to the literature value, it does not necessarily mean that it is more accurate. With the equipment available, the best result that can be obtained is the average result.
Suitability of experiment:
The experiment is suitable to find relationships or rough trends between different types of alcohols. However, the experiment is by no means accurate enough to produce results that can give the heat produced by an alcohol. The equipment available did not ensure that the heat transfer was efficient. Carbon was deposited on the bottom of the copper calorimeter, showing that there was no complete combustion of the alcohols, therefore also suggesting that the heat transfer was not efficient.
Sufficiency of evidence:
The experiment was repeated five times for each alcohol. This gave only five results to plot on a graph. If there had been more anomalous results, I would not have had enough data to compare them with so as to determine the anomalous results.
Improvement:
Obvious improvements would be to make the combustion more efficient so that no carbon was deposited and no carbon monoxide was produced. This could be achieved by designing a burner rather than using a wick. As with a gas oven, the fuel needs to be in gaseous form for the most efficient combustion. One reason why the results obtained were inaccurate is that some of the heat liberated was used to vaporise the liquid prior to its combustion. A suitably designed burner could get around this – presumably using several fine holes rather than one big hole to get better mixing of the vapour with the oxygen. Another modification might be to use oxygen enriched air or even pure oxygen to ensure the complete combustion of the alcohol.
Perhaps a cylindrical reflective metal sheet could be placed around the burner and calorimeter to prevent the heat being lost to the surroundings e.g. a sheet of aluminium.
Further work to provide additional evidence or extend enquiry:
To further my enquiry I would investigate in the same way the combustion of methanol (CH3OH) and pentanol (C5H11OH). This would give me further evidence to conclude that: the larger the alcohol molecule, the more bonds will be broken and formed, and therefore the more heat will be produced. For this enquiry, I would predict that methanol would produce less heat than all of the alcohols, and that pentanol would produce the most heat.
In conclusion, the experiment was worthwhile as it showed the relationship between different sized alcohol molecules in a combustion reaction. However, the evidence is not accurate enough to show any other information, other than the fact that the larger the alcohol molecule the more heat is produced during combustion.
