A + B + X + Y
The average bond enthalpy is the average energy needed to break one mole of covalent bonds. It is given the symbol E and is measured in Kjmol-1. These values are average because the exact value depends on the compound the bonds are in.
The use of the above enthalpy cycles will help to predict the results of the investigation.
The first alcohol of the group will be used as a worked example.
H1
CH3OH(l) + 1½O2 (g) CO2 (g) + 2H2O(l)
H2 H3
C(s) + 4H(g) + 4O(g)
H1 = H3 – H2
= (-393.5) + 2(-258.8) – (-239.1)
= -672 Kjmoldm-3
H2 = Enthalpy change of formation of methanol.
= (-239.1)
H3 = Enthalpy change of formation of carbon dioxide and water.
= (-393.5) + 4(-285.8)
The below values show the enthalpy of combustion calculated by the above method.
Another method used to calculate the enthalpy of combustion is to work out how much energy is required to break all the bonds and how much energy is given out when the bonds are formed. Below shows a worked example using ethanol this time.
The following figures are the average bond energies from the data book
C – C = 347
C – H = 413
O – H = 464
C – O = 358
O = O = 498.3
C = O = 805
These figures are used to calculate the energy required to break the bonds and the energy released when the bonds are formed.
C2H5OH + 6O2 2CO2 + 3H2O
1(C – C) 4(C = O)
5(C – H) 6(H – O)
1(O – H)
1(C – O)
3(O = O)
1(347) + 5(413) + 1(464) + 4(805) + 6(464)
1(358) + 3(498.3)
= 4278.9 = 6004
4278.9 – 6004 = -1725.1 Kjmoldm-3
The values below are the rest of the enthalpy of combustion values calculated in this way.
The values provided in a data book are slightly different, as shown below
From these data values it makes it much easier to predict the results of this investigation. It can be seen that increasing the number of carbons in the alcohol chain increases the enthalpy of combustion.
Apparatus
Diagram of apparatus
Method
The apparatus will be assembled as shown in the diagram above. 100cm3 of cold water will be accurately measured using the pipette filler and poured into the copper calorimeter. The temperature must be recorded. Setup the spirit burner with the required alcohol and weigh. The wick length must be 1.5cm long and this length must be the same for each investigation. This ensures that the flame is the same distance from the bottom of the calorimeter in each case. Once the burner is weighed and replaced the wick was ignited. When the temperature is 15degrees higher than the original temperature the flame was extinguished. The burner must now be reweighed and recorded. This method must be repeated three times for each fuel. The method must be exactly the same in each case. For each different fuel the same spirit burner is to be used, changing only the wicks. This ensures that there are no differences in the methods providing accurate results.
The following equations will be used to calculate the energy produced per mole
The first E = MC x T
Where M = the mass of water heated
C = the specific heat capacity of water which is 4.17Kj-1g-1
T = the temperature change.
This provides the energy transferred to the water provided by the fuel.
The next equation is used to work out the energy produced per mole for each fuel. This is so a fair comparison can be made between the values.
Moles = Mass
Mr
Moles = energy per mole.
Energy
Results table
Moles = mass of fuel used
Mr
Methanol CH3OH
1.02 = 0.0319
32
6.18 = 193.73kjmoldm-3
0.0319
Ethanol C2H5OH
Moles = 0.805 = 0.0175
46
6.18 = 353.14 kjmoldm-3
0.0175
Propanol C3H7OH
Moles = 0.64 = 0.0107
60
6.18 = 561.82kjmoldm-3
0.011
Butanol C4H9OH
Moles = 0.55 = 0.0074
74
6.18 = 853.14kjmoldm-3
0.0074
Pentanol C5H11OH
Moles = 0.59 = 0.0067
88
6.18 = 922.39 kjmoldm-3
0.0067
Analysis
The results obtain are as predicted, however the values are much lower. The table below shows the comparisons between my values and the textbook values.
The results obtain from this experiment, although low were precise. All the results were 30% (± 1%) of the actual results from the text book. From the graph it can be seen that the result for butanol was an anomalous result as it does not fit on the line of best fit as the rest of the results do. The overall trend of the results supports the prediction. The results from the text book are 600 kJmol-3 apart from each other, where as there is no trend in the results of this experiment. There is a 159 kJmol-3 increase from methanol to ethanol, a 208 kJmol-3 increase from ethanol to propanol, a 291 kJmol-3 increase from propanol to butanol, and a 69 kJmol-3 increase from butanol to pentanol. This shows that the average increase is 181 kJmol-3. This increase is again 30% of what is should be. The low results are due to several faults in the experiment. Although the investigation is well planned there are factors that can not be controlled with the equipment provided. The alcohol, when oxidized did not completely combust to produce just carbon dioxide and water. Instead black deposits formed on the underneath of the copper calorimeter, indicating carbon had been formed, and this is a product of incomplete combustion.
When the alcohol is oxidized, it is vapourised, breaking the Van Der Waals forces between the molecules. In a larger alcohol these Van Der Waal forces are easier to break and therefore mean less energy is required to break the bonds and therefore more energy is released when making the bonds involved in Carbon Dioxide and in water, which have very weak Van Der Waals forces as they are gases. When carbon is formed it is solid, which means stronger Van Der Waals forces have to be formed which requires energy, which means not as much energy will be released and transferred to the water in the copper calorimeter.
Percentage errors
To calculate the percentage errors, the accuracy of the equipment is required. In all cases this is indicated by the equipment itself.
Each time the piece of apparatus is used the margin of error increases. As the pipette used was a 25ml it had to be used 4 times to measure out 100ml of water this increased the inaccuracy of the pipette.
To now calculate the percentage error all the error margins are added together to give a total.
Evaluation
The investigation produced results that supported the prediction. As the size of the fuel increased so did the enthalpy change. The results followed the trend but the alcohol did not produce as much energy as expected. The reason for this is because there was a large amount of heat loss which could not be calculated.
- The wind shield although effective was not perfect. A large amount of energy was lost to the surrounding environment through convection currents. The gap at the bottom of the wind shield meant that air could be blown through, disturbing the flame.
- A gap was left at the bottom of the wind shield to allow a sufficient supply of fresh oxygen to the flame for complete combustion to occur, however incomplete combustion still took place.
- The copper calorimeter required a large amount of energy to heat it up before the energy was transferred to the water it contained. This is the main reason for such low results.
- The size of the flame could not be controlled exactly. Although the size of the wick was measured the actual height and width of the flame was not known. This meant that the flames were of varying size, which would have affected the total amount of energy transferred to the calorimeter.
- Energy would have been lost not only to the surrounding, but also to the foil surrounding the calorimeter and spirit burner.
- The accuracy of the equipment used was calculated and accounted for. Better equipment would have provided better results. The lack of equipment in the laboratory had a major effect on the results. Better equipment would have allowed much better results that were closer to the text book results to be obtained. However the required equipment for much more accurate results was not available.
However the results are precise in that they all, except butanol are on the line of best fit.
Overall the experiment was successful. The results obtained give a clear conclusion, that the larger the alcohol chain the higher the enthalpy change.
Improvements
The improvements that can be made to this particular experiment are:
- The specific heat capacity of copper and the mass of copper can by multiplied by the temperature rise. This will give the value for the energy transferred to the copper calorimeter. This would account for a lot of the energy lost from heating the copper calorimeter.
- A broader base copper calorimeter would prevent the flames from flaring out around the base. This would reduce the heat lost heating the surrounding area of the calorimeter.
To improve the experiment all together a bomb calorimeter can be used.
In a bomb calorimeter, a large voltage is required initiate a reaction. The calorimeter is surrounded by an excess of oxygen to ensure complete combustion. A high pressure of 25 atmospheres ensures that there is an excess of oxygen and no air or moisture.
Bibliography
Nuffield book of data values
http://www.acornuser.org/education/thesis/bombcalorimetry.html
Chemical ideas
Chemical storylines
Chemistry in Context – G C Hill J S Holman
A- level Chemistry E. N. Ramsden