Apparatus List
-
Copper combustion calorimeter to hold a fixed volume of water (75cm3)
- Thermometer, 10-50 ˚C in 0.1˚C intervals to measure the temperature rise of the water
- Heat proof mats to act as heat (drought) shields to prevent heat loss and exclude draft which may interfere with the experimental procedure since it may extinguish the flame.
- Bunsen burner and splint for lighting spirit burner
- Ruler to measure the distance between the wick and the base of the calorimeter
- Spirit burner to contain the alcohol
- Clamps to hold the apparatus in position
Diagram of Copper Combustion Calorimeter
5 cm
Before carrying out the experiment, I will consider three factors which may influence the accuracy and reliability of my results. These factors include:
- Distance from wick and metal copper combustion calorimeter
- Amount of water in the copper calorimeter
- Temperature rise of the water
The distance between the wick and copper metal calorimeter should be kept constant throughout the experiment. The distance should not be very large because that would prolong the experiment and result in less efficient burning of the alcohol. Likewise, the distance should not be very small as it would also result in less efficient energy transfer and may put the flame on the wick out if in contact with the base of the calorimeter.
The amount of water in the metal calorimeter should be of a reasonable volume to maintain the temperature of the water and to avoid a lot of evaporation. To ensure a fair test, the same volume of water will be used in each experiment. The volume of water should be such that, each experiment can be carried out quickly and does not waste time heating unnecessary amounts of water.
The temperature rise in which the water in the calorimeter should rise by is an important final factor, for the reason being that the temperature rise has to be one of a suitable rise and be kept the same throughout. I had to be aware of the scale of inaccuracy with particular temperature rises.
Below, a table showing the percentage inaccuracy for 4 different percentage rises can be seen.
Our group decided to use a 25ºC temperature rise. Reason being, that its inaccuracy percentage is low compared to the rest and a 50ºC rise would prolong the experiment even further as it would take a long time to achieve the required temperature.
The safety precautions involved in this experiment include wearing goggles to avoid any injury to the eye and a lab coat to protect clothing and skin. Since alcohols are toxic and flammable when spilt or handled incorrectly, can be a severe hazard.
In order to gain an understanding of this investigation, I was familiar with the method to determine enthalpy changes of alcohols, as we had practiced a trial coursework experiment in the school laboratory in previous lessons. I could therefore use these skills and knowledge to my advantage in preparing for this coursework. The purpose of the preliminary experiment is to allow me to develop a hypothesis for my actual investigation.
The table below shows a summary of the enthalpy changes and percentage accuracy for the results I obtained from the preliminary experiment using only three alcohols.
To find the percentage accuracy, the calculated enthalpy change is divided by the proper enthalpy change (here it is the Nuffield enthalpy change) and then multiplied by 100. I have chosen to show the percentage accuracy in 2 decimal places, as it is a suitable degree of accuracy just as the calculated enthalpy changes are also in 2 decimal places.
Procedure
- Set up the apparatus as shown in the diagram above
-
Place a known volume of water (75cm3) into the calorimeter container
- Stir and record the temperature of the water using the thermometer
- Gently secure the calorimeter in place using two clamp stands on either side (on the diagram, they are marked with an X). Fit the lid with the thermometer in place.
- Take a small spirit burner that is almost full of methanol and place it on the balance to weigh to ± 0.01g. Weigh each of the alcohols being used. Butan-1-ol is the only exception because it is not available in the school laboratory. Once weighed, place the spirit burner beneath the calorimeter ensuring there is a distance of 5 cm between the wick and the calorimeter container.
- Surround the spirit burner with heat proof mats for the heat and energy to be shielded and to ensure draughts do not put out the flame.
- Use a splint to light the wick.
- Slowly stir the water throughout the experiment using a stirrer or glass rod.
- Allow the spirit burner to heat up the water by about 25˚C in each case. The aim is to get the same temperature rise in each experiment so that the heat losses involved in each experiment are about the same.
- Immediately extinguish the flame and reweigh the spirit burner to find out how much fuel (alcohol) has been used. Record the final mass in grams.
- Repeat the above procedure three times with each alcohol to ensure fair test, accuracy and reliability in the results. An average can be taken of the three results.
- N.B/ In each of the repeats, the same equipment and method of carrying out should be used.
Obtaining Evidence
1) The calculation of the enthalpy change of combustion for methanol is carried out as shown below, based on the following assumptions:
Volume of water = 75cm3
Temperature rise = 25° C
Specific Heat Capacity of water = -4.2 J g-1 K-1
Mass of alcohol burned = 1.06g
Average loss in mass of methanol = 1.14 + 1.00 + 1.05 = 1.06g
3
This means that in three trials, the average mass of methanol burned was 1.06g.
Note: To change the units from Joules (J) to Kilojoules (kJ), divide the answer to the above equation by 1000.
Energy transferred (J) = mass of water (g) × specific heat capacity × temperature change
Q (J) = M/g × -4.2 J g-1 K-1 × ∆T °C
Q (J) = 75 × -4.2 × 25
= -7875 J
The number of moles of methanol burned is calculated by the following steps:
Mass of 1 mol of methanol, CH3OH = 12 + (4×1) + 16
= 32g
Number of moles of methanol burned = mass burned mass of 1 mol
=1.06g = 0.031 mol
32
Enthalpy Change of Combustion of Methanol = Energy transferred
ΔHC Number of moles
= -7875 J
0.031mol
= -254032.2581 J
= -254032.2581 / 1000
= - 254.0 kJmol-1
So in this experiment, I have calculated that the enthalpy change for the combustion of methanol is -254.0 kJmol-1. The theoretical value for the standard enthalpy change of combustion of methanol is -726 kJmol-1. Clearly, this value is much greater than that obtained in this experiment. This difference is due to the large heat losses that occur in this experiment. For example, heat losses to the surroundings are quite large even though heat shields are used. Also heat energy from the flame heats up the copper container even though it is a good conductor of heat. Other methods by which experimental errors might be reduced are discussed in the evaluation section.
2) Ethanol - C2H5OH is calculated in the same way as methanol
Average loss in mass of ethanol = 0.94 + 0.95 + 0.96 = 0.95 g
3
This means that in three trials, the average mass of ethanol burned was 0.95g.
Energy transferred (J) = mass of water (g) × specific heat capacity × temperature change
Q (J) = M/g × -4.2 J g-1 K-1 × ∆T °C
Q (J) = 75 × -4.2 × 25
= -7875 J
The number of moles of ethanol burned is calculated by the following steps:
Mass of 1 mol of ethanol, C2H5OH = (12×2) + (6×1) + 16
= 46g
Number of moles of ethanol burned = mass burned mass of 1 mol
=0.95g = 0.158 mol
46
Enthalpy Change of Combustion of Ethanol = Energy transferred
ΔHC Number of moles
= -7875 J
0.158 mol
= -381315.7895 J
= -381315.7895 / 1000
= - 381.3 kJmol-1
3) Propan-1-ol – C3H7OH
Average loss in mass of propan-1-ol = 0.85 + 0.82 + 0.94 = 0.87 g
3
This means that in three trials, the average mass of propan-1-ol burned was 0.87g.
Energy transferred (J) = mass of water (g) × specific heat capacity × temperature change
Q (J) = M/g × -4.2 J g-1 K-1 × ∆T °C
Q (J) = 75 × -4.2 × 25
= -7875 J
The number of moles of propan-1-ol burned is calculated by the following steps:
Mass of 1 mol of propan-1-ol, C3H7OH = (12×3) + (8×1) + 16
= 60 g
Number of moles of propan-1-ol burned = mass burned mass of 1 mol
=0.87g = 0.0145 mol
60
Enthalpy Change of Combustion of = Energy transferred
Propan-1-ol ΔHC Number of moles
= -7875 J
0.0145 mol
= -543103.4483 J
= -543103.4483 / 1000
= - 543.1 kJmol-1
4) Pentan-1-ol – C5H11OH
Average loss in mass of pentan-1-ol = 0.60 + 0.64 + 0.60 = 0.61 g
3
This means that in three trials, the average mass of pentan-1-ol burned was 0.61g.
Energy transferred (J) = mass of water (g) × specific heat capacity × temperature change
Q (J) = M/g × -4.2 J g-1 K-1 × ∆T °C
Q (J) = 75 × -4.2 × 25
= -7875 J
The number of moles of pentan-1-ol burned is calculated by the following steps:
Mass of 1 mol of pentan-1-ol, C5H11OH = (12×5) + (12×1) + 16
= 88 g
Number of moles of pentan-1-ol burned = mass burned mass of 1 mol
=0.61g = 0.0069 mol
88
Enthalpy Change of Combustion of = Energy transferred
Pentan-1-ol ΔHC Number of moles
= -7875 J
0.0069 mol
= -1136065.574J
= -1136065.574J / 1000
= - 1136.1 kJmol-1
Table showing a summary for the results obtained and their accuracy (%)
NB: The percentage accuracy was found using the following equation:
Analysis
Summary Table showing the trend between enthalpy changes of combustion of the different alcohols using different sources of information.
Using this summary table, a graph was plotted to show the trend more clearly and to compare against the theoretical values.
Graph: to show the trend between successive alcohols and their enthalpy changes of combustion. It can be seen that as the number of carbon atoms increases from methanol to pentan-1-ol, the enthalpy change of combustion increases.
By looking at this graph, there are some obvious anomalies in my experimental results. Firstly, although my results show an increasing trend which proves my hypothesis even though butan-1-ol is missing, the trend does not show a strong correlation as the other sources of data show. There is a great difference between my experimental calculated values and the theoretical values suggesting experimental errors and inaccuracies. For example, the enthalpy change for the combustion of methanol is -254.0 kJmol-1. The theoretical value according to the Spreadsheet data and the Nuffield Data Book for the standard enthalpy change of combustion of methanol is -726 kJmol-1. Clearly, this value is much greater than that obtained in this experiment. This difference is most probably due to the large heat losses that occurred in this experiment. For example, heat losses to the surroundings are quite large even though heat shields are used. Also heat energy from the flame heats up the copper container even though it is a good conductor of heat. Other methods by which experimental errors might be reduced are discussed in the evaluation section.
Nevertheless the results I obtained are sufficient enough to prove my hypothesis, in which I predicted that the enthalpy change between successive alcohols would increase as you go down the homologous series. This is because on burning, each alcohol molecule forms one more CO2 molecule and one more H2O molecule than the previous alcohol. More energy is released in the combustion products than is required to break the bonds so an alcohol with a greater number of carbon atoms will release more energy, as can be seen from the graph and equations below.
CH3OH + 1½ O2 → CO2 + 2H2O
C2H5OH + 3O2 → 2CO2 + 3H2O
C3H7OH + 4½ O2 → 3CO2 + 4H2O
Since the enthalpy change is negative because the reaction is exothermic, the values become more negative down the group.
Since my experimental values were quite unreliable, to test my second aim I used the Nuffield Book of Data values.
From the results quoted above it can be seen that the difference between the enthalpy change of combustion value of one alcohol and the next is nearly the same, suggesting that the –CH2― group contributes a specific amount to the total. This also suggests that this amount to the overall enthalpy change of combustion may be made by specific bonds.
When one extra –CH2― group burns, extra energy must be supplied to break more bonds, but even more energy is released by the formation of extra C=O bonds and O―H bonds. I have illustrated this concept below for the burning of the two simplest alcohols, with highlighting in green to show the extra bonds broken and extra bonds formed.
Diagrams to show bond making and bond breaking when two of the simplest alcohols undergo combustion:
1) Methanol CH3OH + 1½ O2 → CO2 + 2H2O
H
Ι
H―C―O―H + O=O → O=C=O + O
Ι ½ O=O H H
H
O
H H
2) Ethanol C2H5OH + 3O2 → 2CO2 + 3H2O
H H
Ι Ι
H―C―C―O―H + O=O → O=C=O O
Ι Ι ½ O=O O=C=O H H
H H O=O
½ O=O O
H H
O
H H
The addition of a –CH2― group has a definite effect on the overall enthalpy change between successive alcohols and is responsible for both the increase down the group and the difference between values as the C―H bond must require a certain amount of energy to break it or to be released when products such as carbon dioxide and water are formed.
Evaluation
Overall I think the experiment went well as I was able to prove my hypothesis and achieve my aims.
The graph I plotted with my experimental results (shown in pink) does not fit in with the main expected pattern (blue and orange lines) indicating that there were possibly many errors during the experiment. Although it is a linear trend it is not a particularly strong correlation (perhaps because butan-1-ol is missing).
I found the experiment straightforward and simple to follow. I had previously read up about calorimeters in other textbooks to inform myself of the procedure and having carried it out in my preliminary, I was confident with the use of the equipment. During the experiment I had to make certain that all the factors which were not being tested were kept constant throughout to ensure fair testing and accuracy.
Although I carried out the experiment to the best of my ability, heat losses in a combustion calorimeter are considerable and too great for accurate work as my experiment results show. To reduce this effect, next time I could calibrate the calorimeter before use.
The results can be said to be both reliable and unreliable. They were reliable in the sense that many variables were controlled and the fact that the combustion of each alcohol was repeated three times. From these results, I obtained an average which was used in the calculations.
However other factors such as any condensation of the alcohol back into the spirit burner and incomplete combustion were not taken into account and therefore not measured. This could have affected the whole experiment as the results may have been very inaccurate. This is evident for pentan-1-ol. Although there is minute dip on the line graph from ethanol to propan-1-ol for the calculated enthalpy, which is hardly evident, pentan-1-ol does not fall anywhere near the straight line which suggests that there may have been a few limitations to the method used.
There were a few limitations in the method. One important limitation was the positioning of the heat proof mats to shield the wick when alight. When an alcohol had to be replaced by another alcohol, the heat proof mats had to be removed, to allow access to the alcohol burner. They then had to be put back into place, however each time they were arranged into different positions. Another limitation was weighing the alcohol burners. When I weighed one of the alcohol burners after combustion had taken place, it weighed more than what it weighed before combustion. This was not expected, so I had to weigh it on another set of scales, which provided me with a different recording. Due to this, other readings may have been manipulated and not very reliable. The amount of time we had in which we had to investigate the combustion of alcohols was also a limitation. If we were allowed more time or if there was no allocated time in which we had to carry out the investigation, then perhaps more data could have been collected and then a more accurate average calculated.
There are a number of modifications to the method which could be used to provide more accurate and reliable results.
- A narrow flame should be used. This is because if the flame is broader at the base, much of the alcohol is not completely burned and a deposit of carbon may be seen on the sides (which was evident during the experiment). This also prevents efficient energy transfer.
- When the hot spirit burner is extinguished, there is considerable evaporation. To reduce this effect, a metal cap could be placed over the wick to extinguish the flame and to trap any alcohol vapour so there is no inaccuracy in the mass of the alcohol when it is re-weighed at the end.
- Use a fume cupboard instead of heat proof mats as heat shields. I think that the fume cupboard would greatly reduce the effect of draught and heat loss.
- Although the copper combustion calorimeter used in this experiment is insulated, some heat will be lost from it. A bomb calorimeter could be used as an alternative. The apparatus is specially designed to avoid heat losses by completely surrounding the ‘bomb’ with water.
Diagram: A Bomb Calorimeter
- The experiment could be caried out in an oxygen rich area to prevent incomplete combustion.
For the investigation to be furthered, some factors could be altered to see their affect on the enthalpy changes of the alcohols. These altered factors could be the same as to those used in this investigation: the distance from the wick and metal calorimeter, the amount of water in the metal calorimeter and lastly the temperature rise in which the water in the calorimeter should rise by. All other factors not being measured should be carefully controlled while one factor is being studied.
One other way to further the investigation could be to look into the enthalpy changes for other homologous series such as alkanes and alkenes. Once results from the enthalpy changes for any of these other compounds have been established, they can be compared to those of the results from the combustion of alcohols and any similar trends can be noted down easily.
Bibliography
- Nuffield Advanced Science Book of Data
- Nuffield Advanced Chemistry Student’s Book
- Introduction to Chemistry by John Murray
- Heinemann Advanced Science Chemistry by Ann & Patrick Fullick
- Chemistry in Context by Graham Hill and John Holman
Computer Software
- Microsoft Word
- Microsoft Excel
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