In a primary alcohol, the carbon which carries the (-OH) group is only attached to one alkyl group. An alkyl group is a group such as CH3 methyl or CH3CH2 ethyl. The structural diagram of a primary alcohol is as follows.
Equipment
- A minimum of 5 spirit burners containing the various alcohols that I will chose from. They must be primary alcohols so there is only one variable that will change between them.
- 1 copper with at least 500 ml capacity for the water, used as a calorimeter.
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250cm3 of water, along with a cylinder of the same measurements. Used to pour water into the calorimeter.
- 1 electronic balance (with measurements to 2 D.P.), used to measure mass of water and fuels.
- 1 Thermometer to measure the temperature of the water. This will be clamped so not to touch the calorimeter as it would affect my results. This will help record the max temperature.
- 2 clamp stands. 1 will hold the calorimeter and the other will hold the thermometer.
- 4 heatproof mats. These will be used to keep my results accurate by working as a draught excluder and also to protect the work surface.
- Matches to light up the spirit burners.
- 1 stirring rod to help get the water to max temperature.
- A ruler to measure the distance between the wick and the calorimeter and also the height of the wick. I must make sure that these measurements are kept the same for all my alcohols and repeats.
Method
- When the experiment has started I will collect all the apparatus above and set up appropriately as I will show in a diagram.
- When collecting the calorimeter, place on the electronic balance of accuracy to 2 dp, and make sure it has been set to zero. I will then pour water using a measuring cylinder into the calorimeter until the mass of water reaches 250g.
- When the apparatus is set up according to the diagram, the testing can begin. The calorimeter consists of a copper can, I will use this instead of a bomb calorimeter which would ensure the alcohol is burned under standardised conditions.
- When the calorimeter is filled with the correct mass of water, I will record the starting temperature.
- I will choose the 5 alcohols that I intend to burn. Weigh each spirit burner each time I repeat the experiment and take note of each mass.
- Check that heat proof mats are in place to act as a draught excluder and also place on top of work surface so not to damage it. Place the spirit burner on heat proof mat and surround with 3 other heat proof mats as shown in the diagram.
- Measure the wick height of each of the spirit burners I will be using, making sure each is the same height. If these variables are not controlled then my results would be inaccurate as the size of the flame for each burner would be different. I will also make sure the distance from the flame to the calorimeter is the same for each spirit burner to make sure the heat travels at the same rate. I will use a ruler to do this.
- Using clamp stands, hold in place both the thermometer and the calorimeter. Making sure the thermometer is only touching the water and not the calorimeter like the diagram suggests. The calorimeter supported by the clamp stand should be the same distance from the flame as suggested in my previous step. To get more accurate measures and a faster procedure, the distance should be kept as low as possible to ensure as much heat as possible reaches the calorimeter.
- Light the wick using the matches. Make sure the wick is lighted very quickly to make sure no excess alcohol evaporates into the air while the cap is off.
- Stir the water using the rod while combustion is taking place. I will keep heating until the temperature change has reached about 15°C. I will try to control this variable by stopping the burning at 13°C or 14°C because the temperature will still be rising when the burner has been extinguished but only by 1 or 2 degrees. This will be the peak temperature. When the temperature change has been achieved I will close the cap which avoids any excess fuel evaporating into the atmosphere.
- Once the burning has finished I will weigh the spirit burner to determine the mass of alcohol that has been used up. I will do this for each experiment or repeat.
- When the burning has finished there will be some soot mixture still on the calorimeter. This is due to lack of complete combustion. I will clean this off using cold water before I use the same calorimeter and burning another alcohol.
- Repeat this for each of the 5 fuels that I have chosen. I will repeat the experiment 3 times for each fuel until I reach concordant results. When doing the experiment again, the apparatus and variables must stay the same. Only the alcohols will differ for which they are all primary alcohols. This means that the length of the carbon chain will be the only changing variable.
Diagram
This is how I will set up my equipment:
Why the plan which I devised is likely to provide precise and reliable results
During my experiment there will 2 types of error that are likely to occur but I will try to prevent. One of the types of error is called systematic error. This happens when something goes wrong while doing the experiment and are based on human error. An example of this maybe if when I measure the calorimeter with the water already filled, this would not take into account the mass of the calorimeter. The other type of error is random error. This is one that I cannot control and is sometimes inevitable. This happens when there is equal chance of something going wrong during the experiment. This type of error is countered by repeating the experiment.
Other points that contribute to prove precision and reliability:
- Using a draught excluder will direct the heat from the spirit burner onto the calorimeter and prevent any excess heat reaching the atmosphere. If available, I will use aluminum foil to attach to the heatproof mats. This will do a better job of deflecting heat onto the calorimeter.
- Measuring the height of the wick. Ensuring the same type of wick and same amount of wick is used for each spirit burner will provide me with more reliable results.
- Measuring the distance between the wick and the calorimeter and keeping this variable the same for each experiment and repeat. Keeping the distance from the flame to the calorimeter the same also. This will contribute to precision and reliability as the heat will reach the calorimeter at the same rate and at a faster rate with less heat going into the atmosphere or having to deflect of the draught excluder.
- Lighting the wick as quickly as possible once the cap is open. This prevents any excess fuel evaporating into the atmosphere. Closing the cap to stop the burning also prevents any excess heat that is not needed going into the atmosphere. This affects the mass of alcohol that will be measured after the burning has finished.
- Cleaning the excess soot that has been left behind on the calorimeter, this is caused by the lack of complete combustion.
- The temperature change must be 15°C for each experiment. To control this variable I must stop the burning when the thermometer read 13°C or 14°C. I will stir the water to reach the peak temperature after the burning has stopped so that it is 15°C or at max 16°C. Stirring the water is another factor that will contribute to the precision of the experiment.
- Repeating my experiment 3 times for each alcohol until I reach concordant results. This will keep my results for each alcohol the same and allows me to find a more accurate average at the end of the experiment.
- Conducting the experiment the same way for each repeat that I do for each alcohol. Overall I may have to do the experiment 15 or 16 times. It is important that I do the experiment in the same manner each time.
- The thermometer must be clamped and not be touching the calorimeter as this would affect the results. The temperature collected must reflect that of the alcohol being burned. This is why I stop the burning 2 degrees earlier to ensure my peak temperature is 15°C.
Health and Safety Precautions
The experiment I am conducting uses highly dangerous alcohols. It is important that I review the safety measures I take to ensure no harm is applied to me or my colleagues. First I will list the type of alcohols and their dangers:
Methan-1-ol-Highly Flammable, Toxic
Ethan-1-ol- Highly Flammable, Toxic
Propan-1-ol- Highly flammable and harmful
Butan-1-ol- Highly flammable and harmful
Pentan-1-ol- Highly flammable and harmful
These are the 5 alcohols that I will be choosing out of the six. When carrying out the experiment I will wear goggles to avoid any alcohol or hot water getting into my eyes. To avoid any alcohols catching fire, I would ensure they near no naked flames. I will also prevent the alcohol evaporating into the air because this will also cause risk to igniting.
Methanol and Ethanol are both toxic. It is important that neither is inhaled or swallowed. If any of these two fuels were to spill, I would inform the nearest member of staff rather than clean it up myself. The member of staff would instruct me what to do. Whilst doing the experiment, it would be necessary to carry it out in a fume cupboard so that no fumes get into the atmosphere.
For the alcohols that are harmful it is important that none of them get into contact with my or my colleagues’ skin. This provides even more reason to wear goggles and any other protective equipment such as a lab coat.
The water used may be spilt which may be hot because it is being heated. To ensure the water burns no one the temperature will be kept to a moderate degree. If any of the glass equipment smashes then I will inform the nearest staff and ask for instructions but I will be very careful with the equipment that I handle. This includes glass rods and beakers.
The clamp stands that I use will provide risk in case they are not stable as they are holding both the thermometer and the calorimeter. I must be extra careful incase the thermometer is filled with mercury as this is a poisonous substance. The calorimeter may also hit the spirit burner which would be extremely dangerous so I must be very sure that the clamps are stable.
Bibliography
-Teacher Support: Coursework Guidance AS/A Level GCE Chemistry (salters, blue sheet)
-Chemical Ideas text book, chapter 4, and 13.
-World Wide Web
-Teacher support.
How I worked out my results
During the experiment I measured the start mass and the end mass of the alcohol whilst burning 200ml of water then measured the temperature change of the water. From these measurements taken I calculated the heat absorbed by the water (Q) measured in Kilojoules.
To work out Q for methanol I used the equation:
Heat Absorbed (Q)=mass of liquid (m) x specific heat capacity of a liquid (c)
x change in temperature (Δθ)
The amount of liquid was 200ml of water. I had to transfer this into mass in grams. This came to 200g. The specific heat capacity of water is 4.2 joules per gram. The change in temperature varied for each alcohol. To show how I worked out the amount of heat absorbed for each alcohol I will use Methanol as an example. As I did 4 experiments for each alcohol I will use the average measurements in the sum. The sum I used was:
Heat Absorbed (Q) = 200g x 4.2 x 16.25°C
= 13650 J
= 13.65 KJ
I did this sum for all the alcohols depending on their change in temperature, for this case it was 16.25°C. The average heat absorbed by the water when heated my Methanol was 13.65KJ.
The heat absorbed by water in each experiment was important to find the enthalpy change of combustion for each alcohol. This is where the start and end mass of the alcohols burned are used. To find the enthalpy change of combustion, I had to find the amount of moles burned in each experiment. This is calculated by using the following formulae.
Moles = Mass/ Mr (Relative molecular mass)
To show how I found the amount of moles burned for each of my experiments I will use Methanol as an example. I will use the average of my measurements that I found to show the technique used. The sum I used was:
Moles = 1.26g/32
= 0.04 rounded to 2 d.p.
I did this sum for each of my alcohols using their specific amount of mass burned. The Mr for the repeats was the same, only the mass burned was different. The Mr for different alcohols was obviously different. The average amount of moles of methanol burned was 0.04.
Now that I have calculated the amount of moles of alcohol burned and the amount of heat absorbed by the water in each experiment it was very simple to calculate the enthalpy change of combustion. The sum used to calculate this measurement was easy. I simply divided the amount of heat absorbed by the water for each alcohol by the amount of alcohol burned to allow this heat transfer. I will use the average measurements for Methanol to illustrate how I calculated the enthalpy change of combustion for Methanol. The sum I used was:
∆ Enthalpy (KJ/mol) = Q (KJ) / Amount of Moles burned
= 13.65 KJ / 0.04 Moles
= 341.25 KJ/mol
This is different to the average enthalpy change for combustion of methanol shown in my results because in the results I simply did the mean of all the enthalpy changes for all my repeats. For this case I did the mean amount of Q divided by the mean amount of moles of methanol burned which of course gave me a different answer. However the technique was used to calculate the enthalpy change of combustion for all my experiments but not my average.
After working out the enthalpy change of combustion for all my experiments I compared them to the real values in the data sheets. Here is a graph to compare the real values to my calculated values:
From my results table I calculated the ratio of my enthalpy data to the text book enthalpy data and they were consistent throughout my results. The ratio was calculated by dividing my ∆ Enthalpy by the text book’s ∆ Enthalpy for combustion. The ratio for all my alcohols was very close to 0.5. The graph above compares the trend in results.
The reason why I did not get the same ∆ Enthalpy as the text book is because of the lack of complete combustion. The experiment I completed was not under standard conditions and much of the heat was lost even though the draught excluder was in place. Both pieces of data show linear relationships between number of carbons and ∆ Enthalpy. My experimental data shows a less steep line than the text book line. This means that from my experiment, as the carbon chain increases the ∆ Enthalpy was not as big as the real values suggest. I have to take into consideration what I did wrong in the evaluation to give these results.
Analysis of Results
From the graph, I can see that most of my results follow the same trend. The only alcohol that produced varying results in the repeats was Pentan-1ol. This was mainly due to the differing temperature changes and the number of bonds that needed to be broken. The line of best fit shows how linear the enthalpy change for combustion is.
From my results I can see that my prediction was correct. As the length of the carbon chain increases, the number of bonds that need to be broken increase and therefore more energy is given out.
As the length of the carbon chain increases the intermolecular forces between molecules become stronger meaning it will take more energy to break them. There are more Intra molecular bonds that need to be broken as well as formed. Intra molecular bonds are the forces within a molecule such as the bond that holds O-H together. Intermolecular bonds are forces between molecules.
A dipole is a molecule (or part of part of a molecule) with a positive end and a negative end. All alcohols have dipoles present in them because of their OH groups. The oxygen and the hydrogen unequally share electrons because of the substantially higher electro-negativity carried by the oxygen. This gives the hydroxide group a permanent dipolar charge which allows for hydrogen bonds to occur between other OH groups in different alcohols. These bonds are also known as intermolecular bonds. When this attraction occurs between hydroxide groups it is known as permanent dipole-permanent dipole attraction where two or more permanent dipoles attract one another.
However because each alcohol only has one OH group in their molecular structure it wouldn’t affect the energy transfer regardless of the number of carbons in the chain. The energy needed to break this OH bond is the same for every alcohol.
The other intermolecular force that occurs within the alcohol is the induced dipole-induced dipole force that operates between all molecules as a result of the temporary dipoles within the molecules. This happens with any molecule due to the electrons in constant motion, and at a particular instant the may not be evenly distributed over the two atoms. This means that one end of the molecule has a greater negative charge than the other end and becomes an intermolecular force. However the force is very weak.
The energy needed to break this bond is affected by the length of the carbon chain as the intermolecular force becomes stronger as the molecule becomes larger and contains more electrons. This induced dipole-induced dipole force in alcohols is the attraction between the alkyl part of one molecule and the alkyl part of another molecule.
In conclusion, the major factor affecting the enthalpy change was the number of carbons in the molecule and the energy needed to break these induced dipole-induced dipole forces rather than the hydrogen bonds between oxygen and hydrogen.
The bonds these charges can form are relatively weak but as the number of electrons available increases so does the frequency of the instantaneous dipole. However the alcohols I am using are all linear and all have the same chance to line up and form instantaneous dipoles.
I will now analyse the bond enthalpies of Methanol and Pentan-1-ol to show how much energy is given off when burning the alcohols. The amount of energy given off for Methanol is much less than that for Pentan-1-ol as my results showed. I will show how much energy is needed to break each bond. These bonds are intra-molecular forces and take much more energy to break them than inter-molecular forces. All bond enthalpies are ∆-ve as the reactions for enthalpy change for combustion is exothermic.
Methanol
CH3OH + 1.5O2 → C02 + 2H2O
Bonds broken: Bonds Formed:
3x (C-H) = 3x 413KJ=1239KJ/mol-1 2x (C=O) = 2x 805KJ=1610KJ/mol-1
1x (C-O) = 1x 358KJ=358KJ/mol-1 4x (O-H) = 4x 464KJ=1856KJ/mol-1
1x (O-H) = 1x 464KJ=464KJ/mol-1
1.5x (O=O) = 1.5x 498KJ=747KJ/mol-1
The enthalpy value is equal to (1239 + 358 + 464 + 747) – (1610+1856) = -658KJ/mol-1
All combustion reactions are exothermic meaning the energy is given out to the surroundings. This means that 658KJ of energy is given out to the surroundings per mole of methanol combusted.
Pentan-1-ol
C5H11OH + 7.5O2 → 5C02 + 6H2O
Bonds Broken: Bonds Formed:
11x (C-H) = 11x 413KJ=4543KJ/mol-1 10x (C=O)=10x 805KJ=8050KJ/mol-1 1x (C-O) = 1x 358KJ/mol-1 12x (O-H)=12x 464KJ=5568KJ/mol-1
1x (O-H) = 1 x 464KJ/mol-1
7.5x (O=O) = 7.5 x 498KJ= 3735KJ/mol-1
The enthalpy value for Pentan-1-ol is equal to (4543+358+464+3735)-(8050+5568) = -4518KJ/mol-1. This means that 4518KJ of energy is given out to the surroundings per mole of Pentan-1-ol combusted.
From these calculations I can see that the enthalpy change of combustion has increased significantly as the extra carbon atom is added making the molecule larger and needing more energy to break the attraction between intra-molecular forces.
Limitations of the practical procedures that I used
When looking through the equipment I used I worked out the precision errors. I found that the only reading that would prove to be inaccurate would be the thermometer. I took a number of thermometers and noted their room temperature to see if they were all the same. The range of results was 2°C between the highest and lowest value. This means that from my results I must calculate the enthalpy change with +/- 1°C for the temperature change. I can see if this changes the trend of results or affects the precision if I compare it to the text book values. This technique will also give me the highest and lowest possible energy transferred. The following is the lowest possible energy values:
The following are the highest possible energy values transferred:
From these tables I have only changed the temperature reading by +/-1°C to show the upper and lower energy values. This was done using the same technique for my original data.
In conclusion, even if these values were used and analysed, the trends would still be the same but slightly less consistant. When looking at the ratio which compares the result to the text book values we can see that the lower energy values produce less accurate readings and the higher energy values gave even closer readings than my actual data. However, the temperature increase did not give that much precision to my results as they are still very far from the text book values. The real reasons for the inaccuracy of my results are down to how poor the equipment I used was.
The limitation in the practical experiment was responsible for my results to be more or less half the real values consistently for all my alcohol data. The following graph compares the energy values for both the text book and my results:
The reason for the two lines being so far apart from each other at different gradients is due to the massive amounts of convection and radiation that was lost during the burning of my alcohols. It was also because of lack of complete combustion that caused the most significant causes to this inaccuracy. This was due to the lack of proper draught exclusion and insulation. To have achieved higher accuracy and precision I might have taped aluminium foil to the heat proof mats. This would have allowed more heat radiation to be deflected onto the water which was where I was trying to divert all the energy towards.
The container used to hold the water; the copper calorimeter was metal and would have radiated heat from its entire surface area. A better designed container or a dull outer surface would have minimized this.
However when comparing my results to the text book values, I could see that although they were not the same, they were consistently wrong by about 50%. The lack of abnormalities in my results showed that my experiment was carried out using a consistant method which allowed me to better analyse my results. The measures that I kept the same during the experiment were the height of wick and the distance the flame was to the calorimeter. This measurement may not be completely accurate due to possible human error which is why not all my results were dead on 50% but were in the vicinity of this value.
Other experimental techniques that I did for all my tests were stirring the water. This allowed the temperature to reach its highest peak using the particle collision theory. Once I stopped the burning, I carried on stirring the water to ensure that the temperature change was accurate in my results. However, as I explained the thermometers used were sometimes ineffective to 1°C.
Things I would have changed had I done the experiment again, would be to use a digital thermometer as this would have reduced the possible human errors that I might have done. Also I would have used aluminium foil to reduce loss of heat radiation.
Bibliography
-Teacher Support: Coursework Guidance AS/A Level GCE Chemistry (salters, blue sheet)
-Chemical Ideas text book, chapter 5
-Collins advanced Chemistry text book Chapter 13
-World Wide Web
-Teacher support.