-The set-up will look like this:
-Light the Burner using a splint.
-Stir the water with the thermometer whilst heating as to heat the water equally, providing better results.
-Heat until the temperature has risen 15oC and then extinguish flame.
-Wait a few seconds and record the maximum temperature reached. (This is because the water temperature will still rise for a while after the heat is removed.)
-Record the weight of the burner.
-Repeat this 3 times for each Alcohol used and take an average for use in calculating the Enthalpy Change of Combustion. When repeating this experiment it is important to keep everything the same apart from the alcohol used. The water however should be changed after every experiment, as it is important to use the same starting temperature each time, and changing the water is quicker than allowing it to cool.
-It is important to be cautious in this experiment as the alcohols are flammable, and there are naked flames being used. This means keeping stock bottles of alcohols away from flames, as well as any pieces of paper/workbooks you might be using. Long hair should be tied back and goggles should be worn at all times as some of the alcohols, (Methanol, Butan-1-ol) are harmful.
-The results will be more accurate if a lid on the burner is used, as it reduces evaporation of the alcohol.
I have chosen the named alcohols, as they are all straight chain alcohols progressively gaining a carbon atom from 1 – 4. This will allow for easy comparison of the alcohols.
I chose Propan-2-ol to compare structural isomers to other structural isomers.
Results Tables:
Analysing Evidence:
[C=12, H=1, O=16]
Methanol (CH3OH):
100 x 4.2 x 14.83 = 6.22 kj per 1.19g of fuel 6.22 = 5.22 kj/g
1.19
Molar Mass:
12 + (3x1) + 16 + 1 = 32
1 mol of Methanol = 32g 5.22 x 32 = 167.04 kj
Enthalpy Change of Combustion for Methanol = 167.04 kj
Ethanol (C2H5OH):
100 x 4.2 x 14.33 = 6.01 kj per 0.47g of fuel 6.01 = 12.78 kj/g
0.47
Molar Mass:
(2x12) + (5x1) + 16 + 1 = 46
1 mol of Ethanol = 46g 12.78 x 46 = 587.88 kj
Enthalpy Change of Combustion for Ethanol = 587.88 kj
Propan-1-ol (C3H7OH):
100 x 4.2 x 14.83 = 6.23kj per 0.41g of fuel 6.23 = 15.20 kj/g
0.41
Molar Mass:
(3x12) + (7x1) + 16 + 1 = 60
1 mol of Propan-1-ol = 60g 15.20 x 60 = 912.00 kj
Enthalpy Change of Combustion for Propan-1-ol = 912.00 kj
Butan-1-ol (C4H9OH):
100 x 4.2 x 15.16 = 6.37 kj per 0.24g of fuel 6.37 = 26.54 kj/g
0.24
Molar Mass:
(4x12) + (9x1) + 16 + 1 = 74
1 mol of Butan-1-ol = 74g 26.54 x 74 = 1978 kj
Enthalpy Change of Combustion for Butan-1-ol = 1987 kj
Propan-2-ol (C3H7OH):
100 x 4.2 x 14.66 = 6.16kj per 0.37g of fuel 6.16 = 16.65 kj/g
0.37
Molar Mass:
(3x12) + (7x1) + 16 + 1 = 60
1 mol of Propan-2-ol = 60g 16.65 x 60 = 999.00 kj
Enthalpy Change of Combustion for Propan-2-ol = 999.00 kj
This experiment shows that the more CH2 groups there is in the structure of these alcohols, the higher the enthalpy change of combustion. The Propan-2-ol had a higher enthalpy change than the Propan-1-ol. This is because there are more branched chains. These branched chains help link the molecules together, making it more difficult to break the bonds apart, using more energy.
Evaluating Evidence:
There will always be inaccuracies in the results due to a number of things. If we address these causes of inaccuracy, then we can perhaps improve them. The main source of error will more than likely be human error, e.g. lack of concentration, accidents, etc. This cannot be calculated, however many other sources of inaccuracy can in fact be calculated using this formula:
Percentage Error = Error x 100
Reading
-The Pipette has a percentage error of 0.06 cm3, the reading was 50 cm3. Using the formula I calculated the percentage error of the Pipette as 0.12 %.
-The Balance has a percentage error of 0.005 cm3, a selected reading was 0.41g. Using the formula I calculated the percentage error of the Balance as 4.17 %.
This means the Balance has quite a high amount of percentage error. There are other sources of error, such as the positioning of the calorimeter, use of a different batch of distilled water half way through the experiment, and doing the experiment over two days.
Overall, however, the experiments results are fairly reliable. They have been made so by doing repeats of the experiment and calculating an average. The more repeats done, the more accurate the results, but time allowed for 3 repeats per alcohol.
If I was to do this experiment again, I would more than likely use a longer time period to do it in, rather than spreading out the experiment over 2 days. I would also do more repeats per alcohol, possibly 5 repeats per alcohol. Another way to make the experiment better would be to do it in a small room. You would have to make sure the room however was well ventilated as many of the alcohols are dangerous to inhale. A better draft shield would be helpful, as well as putting the calorimeter closer to the flame. It would make the experiment a lot quicker if a 100 cm3 pipette was supplied, rather than the 50 cm3 pipette, which was supplied, as the majority of the time to do the experiment was taken up by filling the pipette to refill the water in the calorimeter.
