Preliminary Experiment:
We all did a trial experiment to get a feel of the layout and processes involved in the experiment and also to establish how best to perform the experiment so we knew what we were doing when it came to the real thing. Each alcohol was tested twice to check for anomalies and to give more credence to our results – more sound evidence, more reliability.
Bond Energy Theory:
I’m basing my results on the theory of bond energy – which is defined as the amount of energy in kilojoules associated with the breaking or making of one mole of chemical bonds in a molecular mole or compound (in this case of the alcohols). The energy stored in the bonds is called the enthalpy and is given the symbol H. The ΔH shows the change in enthalpy and is called the heat of reaction. Or so my textbook “GCSE Chemistry – B. Earl & L. D. R. Wilford” says anyhow.
Bonds Made/ Bonds Broken:
In an endothermic reaction, the energy level of the products will be greater than that of the reactants, as the reactants have gained energy during the reaction. In an exothermic reaction such as this one, the reactants lose their energy to their surroundings and end up with less energy in their products than at the start.
Here is an energy level diagram of an exothermic reaction:
As you can see in the diagram above, the energy that the reactants possess at first is increased dramatically (in this experiment by the presence of heat in the form of a flame) so that the energy level rises high enough to penetrate the activation energy barrier and is lost to the surroundings in the form of heat.
Method:
- Obtain correct Apparatus in sufficient quantities. (I think I have listed them all more or less in the “Diagram/Apparatus” section.)
- Clean the copper can of any imperfections on the surface that could influence the results.
- Set up experiment (as shown in the Diagram/Apparatus section).
- Check that the thermometer is suspended in the can, and not touching any of the sides.
- Make sure the can height is constant for each experiment.
- Making sure that the lid of the alcohol burner is fixed on top of the burner (to stop the alcohol evaporating away), weigh the burner on a set of electronic scales – making sure the figures are accurate to 2 decimal places. Note down the starting mass of the alcohol burner.
- Then, place 50ml of water into the copper can that is suspended directly above the alcohol burner.
- Note down starting temperature.
- Light the wick of the alcohol burner and quickly set up an insulation barrier of heatproof mats to block out any draughts or unnecessary heat loss.
- When the temperature of the water has increased by 30ºC, hastily blow out the flame that’s burning merrily atop the alcohol burner (without spilling the alcohol burner as we managed to do once…) and immediately bung the lid onto the burner and reweigh the alcohol burner on the SAME set of electronic scales to the same degree of accuracy as before.
- When all has been completed to a reasonable level of satisfaction, breathe a sigh of relief and brace oneself for the repeat of the whole procedure for the same alcohol (do each alcohol twice to validate results) and move up the homologous series steadily using the procedure outlined above to complete the experiment.
Equations for the bond energy:
Obtaining Evidence:
Results:
Note that all enthalpies were worked out to 2 decimal places to truly preserve that fastidious level of accuracy that we have been told to implement for such important experiments. I have a faint idea that it may help to make my final results more accurate too, which I think would be the better reason out of the two I just gave for working out the enthalpies to 2 decimal places.
Here are the actual text book values for the enthalpies of these alcohols:
Analysis:
I found that our results varied much less in enthalpy than the data book values. I also found however, that there was a similar relative trend in my results because (though they could be viewed as anomalous) they had been obtained in a constant anomalous environment that resulted in the preservation of the relative trend throughout the homologous series. This was because the errors duplicated themselves every time we performed the experiment, since we did our experiments in the same lab, in the same place with similar starting temperatures and the same type of equipment to match. Therefore, though our results came out completely wonky in comparison to the text book values, they were in fact correct to a certain degree in showing the relative trend of the homologous series and their enthalpies.
The enthalpies of the alcohols that we found from our experiments compared with those found in the “Nuffield’s book of Data” were:
The formula for heat energy given to water:
[Heat energy given to water = mass of water (g) × 4.2 (Jg-1ºC-1) × temperature rise (ºC)]
and the formula for finding moles:
[Mol = Mass ÷ Molecular mass]
They were both crucial in finding the enthalpies of each of the alcohols.
Evaluation:
Errors:
We did happen to spill the alcohol burner (only once, luckily) as we tried to put the flame out on the alcohol burner and set fire to our workstation which did results in a bit of a lively fiasco, which was fun… But didn’t really help us in terms of the time lost clearing up the mess etc etc.
Also, since we had spread our experiments over several lessons, we hadn’t paid much attention to the distance between the can and the alcohol burner so it varied over each lesson. Luckily, we did have some idea of how far away the alcohol burner was from the can so we adjusted it accordingly otherwise there might have been some serious consequences on our results…
The lack of Butanol in the lab was a setback for us as we couldn’t obtain the results we needed for that alcohol but since we tested each alcohol twice, it didn’t really matter (that we didn’t get to test Butanol) as we could guess pretty much where Butanol’s enthalpy lay on our graph from the results we had obtained during our experiments.
The fact that there were no proper draught insulating materials available to us so that we had to make to with using heat proof mats as the draught shields in our experiments was a potential error. Also, there weren’t enough to go around so some experiments ended up having virtually no heat insulation and some experiments looking like crude igloos.
Possible Improvements:
We could all do with a bit more common sense here and there. Possibly our experiment may be executed quicker if we’d another go at it.
The distance from the can to the alcohol burner should be logged and kept as similar as possible through the experiments.
Some Butanol can be obtained for the next experiment as well because it may have helped…
Purchasing some proper draught insulating and heat insulating material wouldn’t go amiss either.