· Put the chosen alcohol under the beaker allowing the flame to just touch the beaker.
· Leave to heat up until the alcohol evaporates.
· Take the temperature when the alcohol has evaporated.
· Record the temperature of the water now.
The variables that must remain constant throughout the experiment are…
· Mass of the water 150ml
· Type of can, metal, aluminium.
· Surrounding temperature of around 21°C.
· The height of the can from the crucible.
· Same set of scales.
· I will stir the water with the thermometer throughout heating to keep the water at an even temperature.
The variable that must be changed is…
· The type of alcohol used.
The formulae of the alcohols that I will be using are…
· Methanol C H OH
· Ethanol C H OH
· Propanol C H OH
· Butanol C H OH
Prediction
I predict that the more bonds there are holding the carbon, oxygen and hydrogen atoms together, more energy will be required to break them apart. For example Ethanol has the formula C H OH. In this formula you have five C-H bonds, one C-C bond, one C-O bond and one O-H bond. To separate these types of bonds you require a certain amount of energy which I will show in a table.
TYPE OF BOND ENERGY REQUIRED TO BREAK THE BOND (kj)
C-H 410
C-O 360
O-H 510
O=O 496
C=O 740
C-C 350
To separate C-H bond you need to apply 410 joules of energy. There are five such bonds in ethanol so you multiply 410 by five to get 2050 joules. You do these calculations for all the other types of bonds that make up ethanol, add them all together and you get 3270 joules. All of the other alcohols can be broken up in this way. Below is a table showing the energy required to break up the bonds in each alcohol.
Type of alcohol Energy required to break the bonds in the alcohol (kj)
Methanol 2100
Ethanol 3270
Propanol 3740
Butanol 4560
As you can see a longer molecule takes more energy to break its bonds, in this case Butanol. Compared to a smaller molecule, methanol that requires less energy to do so. I can come to predict that the longer the molecular structure in the alcohol the more energy it will take to remove the bonds. So when I come to predicting results I can safely say that Butanol will evolve more energy than methanol simply because it has more bonds to break.
Results
Repeat readings
Average results
Table showing the temperature change compared to the amount of carbon atoms in each alcohol
Amount of energy produced from each alcohol
Energy = Mass of water x 4.2 x Rise in Temperature
Methanol
E = 150 x 4.2 x 5
= 3150 kj
Ethanol
E = 150 x 4.2 x 8
= 5040 kj
Propanol
E = 150 x 4.2 x 9
= 5670 J
Butanol
E = 150 x 4.2 x 10
= 6300 kj
Energy produced per gram = Energy evolved
Mass of alcohol
Methanol
3150 / 0.5
= 6300 kj x 32g
= 201.6 per mole/kj
Ethanol
5040 / 0.5
= 10 080 kj x 46g
= 463.680 per mole/kj
Propanol
5670 / 0.5
= 11 340 kj x 60g
= 680.4 per mole/kj
Butanol
6300 / 0.5
= 12 600 kj x 74g
= 932.4 per mole/kj
Energy produced per mole
Analysing and drawing conclusions
I think my results tables and graphs clearly show the pattern that I have found in this experiment. That is that heat combustion does increase when the amount of carbon atoms increases. I believe that my results do show a positive correlation and do show that the more carbon atoms there are the heat of combustion goes up. Another reason for these results is that the molecular length becomes longer in the bigger molecules increasing the surface area hence allowing more energy to be released. These results do support my initial prediction. After this I can conclude that my initial prediction was actually right but I didn’t allow for all of the experimental errors. I conclude that carbon atoms in alcohols do have an effect on the heat of combustion. As the amount of carbon atoms go up the heat of combustion does. This is because everytime you add another carbon atom you are also adding 15 onto the relative atomic mass that plays a big part in calculating the end results.
Evaluation
Sound and light energy could have been lost into the room. I could have placed heatproof mats around my experimental area, they could not have kept all of the heat in and much of this would have been taken away in the convection currents through the air. The tin that the water was being held in would have used up some of the heat energy to heat itself up. The alcohol containers had varying amounts of alcohol in them to start with along with varying sizes of wicks. This all contributed to the fact that the flame coming from the alcohol was varying in size so was sometimes not even touching the tin can. The room temperature would also have acted as a cooling agent. One of the less important factors could have been if there was a lacking of oxygen leading to incomplete combustion. Then the oxygen molecules would only form with one carbon molecule producing carbon monoxide but I doubt this actually happened.
This was a very difficult experiment to conduct in a classroom because there are lots of potential ways of losing heat because everything likes to gain heat energy. I think the thing that hindered our results the most was the fact that gusts of air and convection currents were taking the heat away from the experimental area and there was no way to stop this. Perhaps if I started the experiment below room temperature, so that the amount of gained energy, from room temperature, might equal the energy lost at temperatures higher than room temperature, then the experiment could produce better results. If there is a limited supply of oxygen then you get carbon monoxide (each carbon atom can only bond with one oxygen atom). This is when incomplete combustion has occurred. This is so because the carbon monoxide could react some more to make carbon dioxide. If the oxygen supply is very limited then you get some atoms of carbon released before they can bond with any oxygen atoms. This is what we call soot. Since heat is given out when bonds form, less energy is given out by incomplete combustion. So this is why it affects the outcome of the experiment. To overcome this problem, I would have to make sure a sufficient supply of oxygen was involved in the reaction.
