Dianna & Stephane
January 31st, 2012
Comparison of Heat Energy Released by Methanol and Ethanol Experiment Lab
Introduction:
Statement/Goal: To calculate whether methanol or ethanol would be the better fuel as
an alternate energy source in the future using ring stand and thermometer for
measurement.
Chemical Equations for Methanol and Ethanol combustion, respectively:
Variables:
Independent variable: Alcohol- methanol and ethanol.
Dependent variable: Heat energy released.
Controlled/Constant variables: Time.
Hypothesis:
If we burn the chemicals and then measure the heat produced by each of the alcohol,
ethanol would come out as the better fuel, based on researched information.
Background information:
Ethanol and methanol are considered alternate fuels that can be used on vehicles.
Both of them are alcohol based, with properties different from each other. They come
from different feedstocks and thus have somewhat different properties that affect the
engine operation. Methanol, also known as wood alcohol, is produced from natural
gas while Ethanol, often called grain alcohol, is principally made from fermented
corn. Ethanol works well in internal combustion engines, and is a high-octane fuel.
“Octane helps prevent engine knocking and is extremely important in engines
designed to operate at a higher compression ratio, so they generate more power.”
Methanol, on the other hand, can be used to make methyl tertiary-butyl ether
(MTBE), an oxygenate that is blended with gasoline to enhance octane and create
cleaner burning fuel, also lowering the risk of flammability when compared to
gasoline.
Procedure:
1. Put on apron and safety goggles.
2. Obtain two spirit bottles; one with ethanol and one with methanol.
3. Punch holes in a soda can and stick a stirring rod through to hold it on the ring
stand. Measure 100.0 mL of water in a 50mL graduated cylinder and pour it
into the soda can. Put the soda can on the ring stand.
4. Use a thermometer to measure the initial temperature of the water
5. Measure the mass of the spirit burner (with the wick but without the cap) and the
methanol.
6. Place the spirit burner on the stand about two centimeters below the soda can.
7. Light the wick and leave it burning for three minutes.
8. After three minutes, move the can and put out the fire. Measure the final
temperature of the water. Take note of the temperature change. Measure the
mass of the spirit burner without the wick and record the mass of the alcohol
burned (change in mass).
9. Clean the bottom of the soda can if blackened and replace the 100.0 mL of water.
10. Repeat steps 4-9 twice more with methanol and with ethanol three times.
Data:
Raw Data Collection in Ethanol and Methanol Comparison
Initial
mass (g)
±0.0005g
Final
Change in
mass (g) Mass (g)
±0.0005g ±0.5mL
Initial
temp
(˚C)
±0.5 ˚C
21.0
Final
temp
(˚C)
±0.5 ˚C
61.0
Ethanol
249.146
247.652
1.494
1
Ethanol
247.652
246.206
1.446
23.0
65.0
2
Ethanol
246.206
244.482
1.724
21.0
66.0
3
Methano 179.881
178.225
1.656
21.0
63.0
l1
Methano 178.225
176.411
1.814
21.0
60.0
l2
Methano 176.411
174.802
1.609
20.0
61.0
l3
*Absolute uncertainties for mass measurement: ±0.0005g
Calculation for Table#1:
Change in temperature = Final temperature - Initial temperature
Eg. 61-21 = 40˚C
Change in Mass = Initial mass - Final mass
Eg. 249.146 - 247.652 = 1.494 g
Change
in temp
(˚C)
±0.5 ˚C
40.0
42.0
45.0
42.0
39.0
41.0
Data Processing of Ethanol and Methanol in Comparison
Alcohol/Tr Chan
ial
ge in
Mass
(g)
Ethanol T1 1.494
Amount
of moles
(mol)
0.0324
Total
Heat
Releas
ed(kJ)
16.720
Heat
released
per mol
(kJ/mol)
515.7012
Percentage Heat
of error (%) (J/g)
62.19
Ethanol T2 1.446 0.0314
17.556 559.4609
58.98
Ethanol T3 1.724 0.0374
18.810 502.7638
63.14
Methanol
T1
Methanol
T2
Methanol
T3
1.656 0.0517
17.556 339.7764
53.13
1.814 0.0566
16.302 288.0259
60.27
1.609 0.050202
808
17.138 341.37532 52.91
63
11191.
43
12142.
08
10910.
67
10601.
45
8986.7
7
10651.
34
Calculation for Table #2:
Number of moles of alcohol burnt = Change in mass/molar mass (of
ethanol/methanol)
Eg. n = m/M
*For molar mass, refer to a periodic table.
Total heat released = cm∆T
*Divide result by 1000 to convert J to kJ.
Heat released per mol = total heat released / number of moles burnt
Eg. 16.720/0.0324 = 515.7012 kJ/mol
Percentage of Error = (Theoretical - Experimental) / Theoretical x 100
Heat (J/g) = Total heat released (in Joules) / Change in mass
Eg. 16720 / 1.494 = 11191.43 J/g
Percentage of Uncertainty in volume of water= 1mL/100mL x 100 = 1%
Conclusion:
The data we collected in the end supports our hypothesis in the beginning, with
statistics showing Ethanol is clearly the chemical that creates more heat than
Methanol when we burn it with the ring stands. The data processing of heat released
for Ethanol is greater than that of Methanol by roughly 100 to 300 joules per gram.
According to our experiment, the heat released per mol is also greater for Ethanol
with a difference of 200 to 300 kJ per mol. The table showed us not only the fact that
Ethanol is the chemical that generates more heat out of the two, but also that it
supports the background information we had researched in the beginning. Ethanol can
be concluded that it is a better fuel option than Methanol because it has “38% more
available energy than methanol” with 86,000 BTU's per gallon for ethanol, while
methanol only has 62,000 BTU's/gallon.
The percentage of error is quite high when compared to the percentage of uncertainty,
with an average result of 58.44% of error throughout the experiment. The percentage
of uncertainty, on the other hand, is 1% because we used a 50mL cylinder to measure
for 100mL of water.
Despite the errors and uncertainty in this experiment, methanol still has only half the
energy content per gallon of gasoline while Ethanol is two-thirds the intensity of
gasoline, which means Ethanol is the better fuel.
Evaluation:
Although our experiment was effective in determining whether methanol or ethanol
was the better fuel, there were certainly changes to be made in the procedure in order
to obtain more accurate results. For example, when reading the temperature at the end
of three minutes, we would sometimes wait for the mercury to stop rising before
recording the results, which meant that our experiment lacked standardization with
respect to timing. We should have been more careful to pay attention to the
temperature right at three minutes. Another area in which we lacked control over
variables was the containment of heat. We had our spirit bottle burning 2 centimeters
below the soda can, and no way of keeping the heat from escaping in the open air.
This was probably the reason why our values for the heat released per mol of alcohol
burnt were much lower than the theoretical values. To counteract this error, we could
have a cylinder of newspaper or styrofoam around (but not touching) the lower
portion of the soda can and enclosing the spirit bottle as well. While it would not
solve the problem completely, it would definitely help trap more heat.