Comparing the enthalpy changes of combustion of different alcohols.

        H        OH        H

Butan-1-ol

H        H        H        H        

         |         |         |         |        

H   —        1C  —        2C  —        3C  —        4C  —        H

         |         |         |         |        

        OH        H        H        H        

Butan-2-ol

H        H        H        H        

         |         |         |         |        

H   —        1C  —        2C  —        3C  —        4C  —        H

         |         |         |         |        

        H        OH        H        H        

Pentan-1-ol

H        H        H        H        H

         |         |         |         |         |

H   —        1C  —        2C  —        3C  —        4C  —        5C  —        H

         |         |         |         |         |

        OH        H        H        H        H

Equipment required.

• Small copper can to act as a calorimeter
• Thermometer (to 1°C of accuracy)
• 100cm3 measuring cylinder
• Spirit burners containing –  Methanol

        Pentan-1-ol

        Propan-1-ol

        Propan-2-ol

        Butan-1-ol

        Butan-2-ol

• Digital balance (to two decimal points)

• Draught shielding x2
• Stirring rod
• Stand and clamp
• Ruler (to measure the height of the calorimeter above the spirit burner)
• Safety goggles and lab coat
• Box of matches
• Pen and paper to record the results
• Paper towels to help clean soot off the bottom of the can

Variables.

• Draught
• Amount of water
• Material of can
• Starting temperature of water
• Temperature rise
• Height of the calorimeter
• Length of wick
• Time burnt
• Amount of oxygen present

My dependant variables are measuring the water temperature before and after and the mass before and after the experiment of the alcohols. I will not be timing the experiment, as it will not affect my results by the time taken for the temperature rise of 15°C to be obtained, although it would be interesting to study.

My independent variable is the alcohol, as these are substances that I will be changing throughout the experiment. These will be methanol, butan-1-ol, butan-2-ol, propan-1-ol, propan-2-ol and pentan-1-ol.

How will my method give me accurate and reliable results?

The plan that I have used will provide accurate results as I am using equipment that has a fairly high degree of accuracy. Another measure of accuracy is how well I can read the equipment that I am using. For example it will be easy to accurately read the reading on the digital balance but it will be harder to accurately read the temperature on a thermometer or the amount of liquid in a measuring cylinder. I am keeping the variables that I do not wish to change as constant as I can; although some of these are hard to control e.g. amount of O2 present in the atmosphere.

The plan that I have used will provide reliable results as I plan to repeat the experiment three times. This will give a reliable portrayal of the amount of energy released when one mole of a substance is burnt completely in oxygen.

Fair test

In order to get reliable results I will need to keep the experiment fair. I will do this by doing the following:

• The amount of water will stay constant throughout the experiment.
• I will use the same measuring cylinder to measure the amount of water used
• The experiment will be carried out away from a window to prevent any extra heat entering.
• The same apparatus should be used for each experiment with each fuel.
• The room temperature should be kept fairly constant.
• The length of the wicks must be kept at similar lengths to stop the alcohols burning at different rates to each other and therefore having different amounts of heat loss.
• Stirring the water using a stirring rod and NOT a thermometer.

Some different aspects of fair testing cannot be done. One such example is measuring the amount of water to fill the copper can with. If you use a measuring cylinder there will always be some water droplets left on the side of the cylinder and this is why no liquids can accurately be measured.

Risk assessment.

When working with alcohols there are precautions that need to be taken.

Methanol

• Highly toxic when ingested, inhaled or absorbed by the skin
• If inhaled, it produces a narcotic effect
• Harmful on contact with the skin
• Highly Flammable

        

Pentan-1-ol

• Highly flammable
• Harmful- Causes lung damage if inhaled
• Extremely irritant

        

Propan-1-ol and propan-2-ol

• Highly flammable
• Irritant
• Harmful

        

Butan-1-ol and Butan-2-ol

• Irritant
• Harmful
• Flammable
• Risk of serious damage to the eyes. Irritates the skin and is harmful if swallowed

        

• Care must be taken with the thermometer.
• Take care not to touch the calorimeter immediately after the experiment, as it will be hot.
• Keep the wick length small or the flame will be very high.

Safety

In order to keep my experiment safe I will do the following;

• Wear safety goggles and protective clothing (lab coat).
• Keep all of the alcohols away from naked flames as they are all extremely flammable.
• Work in a safe environment – free from clutter.
• Carry all substances with care.
• Remove loose clothing (e.g. tie).
• Do not inhale or drink any of the fuels as they are harmful.
• Do not eat or drink during the experiment.
• If any substance is dropped on the floor contact the lab technician immediately.
• After the experiment is completed wash your hands thoroughly.
• Keep a window open to allow ventilation of the room as some of the fuels emit harmful toxins which can cause damage to your lungs.
• Follow normal laboratory rules.
• Do not remove the lid of calorimeter until it needs to be removed.
• Strike matches away from the body.
• Make sure that the calorimeter is straight, as when the water starts to heat up the temperature increases and there is a risk of the calorimeter falling off the stand.

Diagram

Method

• Measure out 200cm3 of water using the measuring cylinder – I will make sure that the measurement of the water is taken using the bottom of the meniscus.

• Put the 200cm3 of cold water in a copper calorimeter (check the calorimeter is dry, it could affect the results) put the thermometer into the cold water and leave for a few minutes. Record it’s initial temperature. Write this into the appropriate column of the table.
• Support the calorimeter making sure that the calorimeter is straight and not leaning to one side and place underneath a spirit burner containing the alcohol you are going to burn (the calorimeter should be exactly 16.5cms from base of spirit burner), as shown in the diagram.
• Arrange a draught exclusion around the stand– this is a variable. This will stop any sudden breeze or draught affecting my results.
• Make sure that the reading on the digital balance is 0.00g and then weigh the burner and the alcohol, with the lid on to avoid evaporation of the alcohol. Keep the lid on the calorimeter at all times until you are ready to start to burn the fuel. Record this initial weight in the results table.
• Place the burner under the calorimeter, remove the lid and light the wick using the matches, (hold the match in one hand and lift the lid with the other) as shown in the above diagram.
• Stir the water, using the stirring rod, all the time whilst the water is being heated. This is to help get more precise results, as the heat will be circulating and being evenly distributed throughout the mass of water and not just sitting on the top.
• I plan to allow the temperature of the water to rise by 15°C, so I will allow the waters temperature to rise by 13°C-14°C and then extinguish the flame using the lid of the calorimeter. Care will need to be taken, as the flame can be quite big. Any heat transfer from the copper can (calorimeter) to the water in the can will now not affect my results.
• Replace the lid of the spirit burner to avoid evaporation of the alcohol and weigh the burner to find the mass of the alcohol that has been burnt. Check that the digital scales show a reading of 0.00g.
• Record the final weight in the table.
• Empty the contents of the calorimeter into the sink and rinse with water to cool the can and wash off any soot that has accumulated on the bottom and sides of the can.
• Change the water in the copper can again using the measuring cylinder to pour the 200cm3 of water into the calorimeter, again measuring to the bottom of the meniscus
• Repeat the experiment three times.
• Repeat the same experiment with the other alcohols, using the method above, keeping all conditions the same.
• Tabulate all of the results.
• Work out the mass of alcohol burnt on each occasion

Sources used in making the plan.

• Salter’s teacher support: Coursework guidelines sheets OCR 2004, Second edition page 21
• Developing fuels worksheets – Salter’s advanced chemistry 2000. DF1.3 – Comparing the enthalpy changes of combustion of different alcohols.
• Salter’s advanced chemistry, Chemical ideas, second edition pages 304-307 chapter 13.2 alcohols and ethers
• Salter’s advanced chemistry – Chemical storylines, Second edition pages 38-40 DF7 Changing the fuel
• Hazcards
• Data sheets from Salters Advanced chemistry 2000, Table 20 bond lengths and bond enthalpies

Implementation and results

Change in temperature is ALWAYS 15°C.

Methanol

Pentan-1-ol

Butan-1-ol

Butan-2-ol

Propan-1-ol

Propan-2-ol

The mass of the alcohols both before, after and the change in mass is to 2 decimal places. This is because the digital balance weighs to this degree of accuracy. The measurement taken with the thermometer is to one whole number because this is the degree of accuracy that the thermometer has. 

Analysing evidence and drawing conclusions.

Calculations.

Now that I have my results, I will work out the enthalpy change of combustion (∆Hc) for each alcohol.

I will work out the ∆Hc for each replicate for each alcohol and then work out the average ∆Hc.

Here is an example of how I will work out the ∆Hc. I will use replicate one of Methanol.

1. It takes 4.2J of energy to heat 1g of water by 1°C. So the amount of energy taken in by the water is calculated using the equation

M x C x Ø

Mass of water x Change in temperature of water x Specific heat capacity of water

200g x 15°C x 4.2J = 12600J or 12.6KJ

200g represents the starting mass of the water in this experiment which was 200cm3.

15°C is the required raise in temperature of the water used in each experiment.

4.2Jg-1 K-1 is the specific heat capacity of water.

2. Working out the number of moles of methanol burnt using

     Mass___    

Molar mass

Molar mass of Methanol (CH3OH) = (12x1) + (4x1) + (1 x 16) = 32g

3.15g is the first reading for the replicate one of methanol for the change in mass of the fuel.

3.15g

 32g         = 0.0984375 moles

A mole is 6.02 x 1023 atoms of the chemical. It is also known as Avogadro’s constant.

3. If 0.0984375 moles of Methanol release 12600J of energy, 1 mole would release

  12.60kJ__

0.0984375         = 128kJmol-1

Worked out by Specific heat capacity of water in Kilojoules

    Number of moles of methanol burnt.

Therefore the ∆Hc of methanol is –128KJmol-1

(Enthalpy of combustion is ALWAYS negative because all combustion reactions are exothermic)

I have calculated the enthalpy change of combustion (∆Hc) using the same method for each of the replicates for each of the fuels.

The chart on the next page shows the enthalpy change of combustion for each alcohol and the average enthalpy change of combustion overall.

I have highlighted the anomalous results in red. I intend to redo these results in order to get a better set of results to compare. When the experiment for methanol was carried out attempt two and three were done on the same day, but attempt one was done the day before. This may have resulted in the anomalous result. Also this was the first experiment that I carried out so it could have been that the way I was carrying out this replicate was different to all the others (experimental technique). I cannot account for the anomalous results for pentan-1-ol. There is a chance that I mis-read the bottle and did another fuel e.g. picked up propan-1-ol instead of pentan-1-ol or used pentan-2-ol instead of

Pentan-1-ol.

Below you will find the revised results for the experiment.

Pentan-1-ol

Methanol

As you can see these results are a lot more reliable than the previous set of results. The value obtained for methanol is similar to the other values obtained for that particular fuel. The results obtained for pentan-1-ol are very close together and may show an improvement in my experimental technique as I have carried out the experiment.

From these results we can now conclude that there is a definite relationship between the number of carbon atoms in the chain and the average worked out enthalpy change of combustion. I have input these results into the charts and worked out the average.

I will now draw a graph to interpret my results including the isomers and also another graph comparing my results with the results from the data books. (This graph will be seen in the analysis).

From the results that I have obtained I can see that there is a general trend between the number of carbon atoms in the chain and the enthalpy of combustion. The enthalpy of combustion becomes more negative as the alcohol molecules become larger.

The way to explain this is because the existing bonds between the alcohol and the oxygen are broken. To do this we must provide the activation energy to start off the reaction. This requires a spark or a burning splint to light the flame to supply this activation energy. If the bond breaking value is bigger than the bond forming value then the reaction is exothermic. If the bond breaking value is smaller than the bond forming value then the reaction is endothermic. Once this has happened new bonds form to produce the new products, which for combustion is carbon dioxide and water. As you go through the alcohols you are adding one carbon and two hydrogen atoms each time to the formula of the alcohol.

If we take into account the existence of isomers of each alcohol, we can see that the enthalpy of combustion value of propan-2-ol is less than the enthalpy of combustion value of propan-1-ol. The reason for this is, as the alcohol molecule does not form a straight chain in propan-2-ol as it does in propan-1-ol, the molecules cannot be so close to each other. This is mainly due to the position of the OH (alcohol) group in the molecule which alters the position of the electromagnetic poles.

If we study butan-1-ol and butan-2-ol we can see that this is not the case according to the results. But it has been scientifically proven that this is what should happen; these anomalous results can be put down to the technique used in the experiment, i.e. the heat loss and the incomplete combustion.

The volatility of a compound is a measure of how capable the compound is of evaporating or vaporizing at low temperatures. The more volatile a compound is, the higher the enthalpy of combustion value. This is because the distance between the bonds increases.

So in summary the longer the hydrocarbon the greater the enthalpy change of combustion due to the low volatility.

If we compare the average values obtained by my experiment to the average values from a data book and work out the error this pattern can be further proven.

If we study the enthalpy change of combustion from the values in the data book we can indeed see that as the number of carbon atoms in the chain increases the enthalpy change of combustion becomes more negative.

This graph shows the relationship between the enthalpy change of combustion and the number of carbon atoms. In order to keep this particular graph simple I have only used propan-1-ol and butan-1-ol and have left out propan-2-ol and butan-2-ol (the isomers). If we study this graph we can see that there is a perfect straight-line relationship between the number of carbon atoms and the increasing enthalpy of combustion value. The line is not exactly straight because ethanol has been excluded.

The values from this graph are the values obtained from the data book.

The enthalpy change of combustion is measured in KJ mol-1.

On the graph you can see that I have drawn an estimation of the enthalpy change of combustion value for ethanol. I can estimate that the value for ethanol is approximately –1350 KJ mol-1. The actual value is -1367 KJ mol -1

This graph further proves the relationship between the number of carbon atoms and the enthalpy change of combustion.

I will now construct a graph comparing the results that I obtained through my experiment and the results obtained from a data book.

Again the enthalpy change of combustion is measured in KJ mol-1.

This graph shows that my results showed a similar pattern to the results in the data book. Due to the size and scale of the graph the pattern of the values that I obtained for the enthalpy of combustion cannot be seen as clearly as the pattern for the values from the data book.

Sources used in the analysis.

• Nuffield advanced science book of data, Chemistry physical science Physics 1975
• Nuffield advanced science, revised book of data1984

• Salter’s teacher support: Coursework guidelines sheets OCR 2004, Second edition page 25

Evaluating evidence and procedures.

Doing experiment and investigations are the best ways of obtaining the enthalpy of alcohols, but in this particular experiment it was too inefficient to gain any reliable results.

During the experiment I noticed a lot of heat loss. Although the draft excluder did limit the amount of heat lost to the surrounding area, the heat did manage to find other ways out. Therefore most of the heat did not go to the water. A lot of heat was also lost to the equipment. After heating the metal calorimeter I noticed that each time a layer of black soot was forming under the copper can and up the sides of it. This is a sign of incomplete combustion. I did not expect to get very close to the actual values for enthalpy of combustion, as there was too many ways of losing heat to the atmosphere.

Also some of the water may have evaporated so there was less to heat which caused the reading to increase.

In order to make sure that my results would be even more reliable I would need to conduct the experiment in an oxygen-rich atmosphere so that there would be excess oxygen for complete combustion to occur. I would also need to reduce the distance between the flame and the bottom of the calorimeter. Another way of making sure that the experiment is carried out fairly would be to have all of the lengths of the wicks the same. I could also cut the end of the wicks off after each experiment to make sure that there is the same amount of alcohol absorbed into the end of the wick. If the metal calorimeter had a bigger base then more heat would be absorbed into the water, which would in turn reduce the amount of heat lost to the atmosphere.

The most important aspect of the procedure, which resulted in the results only being approximately ⅓ to ¼ of the actual enthalpy of combustion value, was the amount of heat that was lost to the surrounding area. The next major cause for the results was the incomplete combustion of the fuels.

Incomplete combustion is the partial burning of fuels, which can be due to lack of oxygen or low temperatures.

Another discrepancy would be the equipment that was used to make many of the measurements.

Working out the percentage error of the equipment used.

 In order to work out the percentage error we use the following equation

Percentage error = Error x 100

    Actual reading

Below I have worked out the precision error for each piece of equipment used

1. Digital balance = ±0.005g
2. Thermometer = ± 0.5°
3. Measuring cylinder = ±0.5ml
4. Ruler = ±0.05cm

 The percentage error for each piece of equipment is,

1. Digital balance = (0.005x 100)/251.50 = 0.02%
2. Thermometer =  (0.5 x 100)/100 = 0.5%
3. Measuring cylinder = (0.5 x 100)/100 = 0.5%
4. Ruler = (0.05 x 100)/160.5 = 0.31%

From these calculations we can see that the thermometer and the measuring cylinder had the largest percentage error, closely followed by the ruler. The most significant piece of apparatus in causing any discrepancy of results would be these two. If you relate the percentage error/precision error to the error obtained by incomplete combustion or the heat lost to the surroundings, we can see that in comparison these instruments are insignificant when compared to the error that would have been caused by these factors.

In order to improve the instruments/equipment with the highest percentage/precision error I could use a digital thermometer which would be more sensitive to temperature changes. This would have had some effect on the accuracy of the experiment, although the only way to improve the experiment would be to improve the procedure and remove some of the errors in that. I could use a pipette (like the type used in titration’s) instead of a measuring cylinder to decrease the inaccuracy of this particular piece of equipment and make it more precise to measure liquids. I would also use a larger range of fuels e.g. hexanol, heptanol, octanol and decanol to give further evidence that there is a perfect relationship between the number of carbon atoms and the enthalpy change of combustion.

I would need to find a way of introducing an ‘oxygen feed’ into the experiment so that there will be enough oxygen for complete combustion and also clean off any soot that had accumulated on the copper calorimeter. The reason that the incomplete combustion would have such a large effect on the overall results would be that the products formed from incomplete combustion release less energy than the products that are formed from complete combustion.

The equations for the complete combustion of the alcohols used are,

Methanol

CH3OH(l) + 1½O2 O(g) → CO2(g)  + 2H2O(l)

Propanol

C5H11OH(l)  + 8O2(g)  → 5CO2(g)  + 6H2O(l)

Butanol

C4H9OH(l)  + 6½O2(g)  → 4CO2(g)  + 5H2O(l)

Pentan-1-ol

C5H11OH(l)  + 8O2(g)  → 5CO2(g)  + 6H2O(l).

Even if I could weigh out all of the soot and work out the number of moles of carbon, there would be no way of getting all of the soot off the can.

In the calculations we do not account for the energy lost as light energy. This would represent a significant amount of energy.

The best way would be to minimize the overall distance between the calorimeter and the flame and have a calorimeter with a larger surface area to absorb more heat from the flame.

In order to reproduce the experiment and obtain a higher degree of accuracy of results that are closer to the actual values in the data book, there are some subtle changes that will need to take place.

1. The calorimeter we used consisted of a basic copper can. In reality a calorimeter looks like this

I would use this piece of equipment. I could not use this in the experiment that I carried out because the school does not have access to this type of equipment.

2. The calorimeter would have a lid on it to stop any heat escaping that way.
3. As already said an oxygen feed to provide excess oxygen would be set up to ensure that complete combustion occurs.
4. The distance between flame and calorimeter would therefore be reduced.
5. Surface area has been increased.
6. Use distilled water instead of tap water.

Sources used in the evaluation.

• Chemical Ideas, second edition page 106 – Specific heat capacity
• Picture from Google. Site address:-http://www.personal.psu.edu/staff/d/c/dcw1/graphics/calorimeter.gif&imgrefurl=http://www.personal.psu.edu/staff/d/c/dcw1/graphic_art_examples.html&h=277&w=288&sz=15&tbnid=qg4P3zGdnlwJ:&tbnh=105&tbnw=109&start=1&prev=/images%3Fq%3Dcalorimeter%26hl%3Den%26lr%3D