Lab Report on the Molar Enthalpy Change of Combustion PDF

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An experiment to determine the energy change when one mole of a substance is completely burnt in air or oxygen...

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An experiment to determine the enthalpy change of combustion of alcohols (Methanol, Ethanol and Propan-2-ol) Introduction Alcohols are organic molecules that contain a hydroxyl group, –OH. Alcohols form a homologous series with the general formula, CnH2n+1OH. Alcohols are grouped into three categories based on their structure: primary, secondary and tertiary. In primary (1˚) alcohols, the carbon atom bonded to the –OH functional group is attached to only one alkyl group, such as ethanol. Methanol is an exception to this case as there are no alkyl groups attached to the carbon atom, only hydrogen atoms, yet it is still classed as a 1˚ alcohol. In secondary (2˚) alcohols, the carbon atom bonded to the –OH functional group is attached to two alkyl groups, such as propan-2-ol. In tertiary (3˚) alcohols, the carbon atom bonded to the –OH functional group is attached to three alkyl groups, such as 2-methylpropan-2-ol (Ryan and Norris, 2014). The more carbon atoms added to the homologous chain, the more complex the structure becomes, and their physical and chemical properties will also differ and follow a certain pattern. Structural formula:

Methanol

Ethanol

Propan-2-ol

Alcohols have a higher boiling point in comparison to other organic molecules. This is due to the strong intermolecular hydrogen bonding between the –OH functional group. The hydrogen bond forms between the lone pair of electrons on the oxygen atom of one alcohol molecule and the hydrogen atom of another alcohol molecule (Smith and Older, 2015). The δ+ hydrogen atom in the –OH functional group of the alcohol is strongly attracted to the highly electronegative oxygen atom. Van der Walls’ forces are the weakest form of intermolecular forces and exist in all molecules. They are essentially a result of electrostatic attraction between temporary and induced dipoles caused by random movement of electrons in a molecule. The longer the alkyl chain in an alcohol, the higher the boiling point as the number of carbon atoms increases and so the increase in van der Walls’ forces means more energy is required to break these intermolecular forces (Cann and Hughes, 2015). The combustion of alcohol is exothermic which produces a lot of energy. Alcohols burn in excess oxygen to produce carbon dioxide and water. CH3OH (l) + 1½O2 (g)  CO2 (g) + 2H2O (l) C2H5OH (l) + 3O2 (g)  2CO2 (g) + 3H2O (l) C3H7OH (l) + 4½O2 (g)  3CO2 (g) + 4H2O (l) Due to the complete combustion of alcohol, it only gives water and carbon dioxide as the products, they make good fuels that burn cleanly and quickly (Wooster and Eccles, 2008). Alcohols give out a lot of energy, are readily available and are easily stored so they are good as fuels.

Fundamentally, combustion is the oxidation of carbon compounds in air to produce heat as well as carbon dioxide and water. The standard molar enthalpy of combustion (∆H˚c), is the enthalpy change when one mole of a substance is burned in excess oxygen under standard conditions with the reactants and products in their standard states (Lister and Renshaw, 2009). Enthalpy changes are all measured under standard conditions in which the reactants and products are in their standard states at 100kPa and 298K (25˚c). It is measured in kilojoules per mole, kJ mol-1. Heat is a measure of the total energy of all the particles in a substance. The energy of every particle is included and so the more particles there are the more heat there is. (Lister and Renshaw, 2009). Enthalpy is the heat content stored in a chemical system and heat cannot be measured directly. Therefore, only enthalpy change (∆H) can be calculated when a chemical reaction takes place and there is an exchange of heat energy with the surrounding. Enthalpy change is calculated using this formula: ∆H = Hproducts - Hreactants Chemical reactions can be grouped into two categories: exothermic and endothermic. Exothermic reactions releases energy and so the temperature of the surrounding increases. The combustion of alcohols is an exothermic reaction. The reason why ∆H is always negative for exothermic reactions is because the enthalpy of the product is always lower than the reactant (Wooster and Eccles, 2008) which is shown in figure 1. Endothermic reactions absorb energy.

Figure 1: The potential energy (enthalpy) diagram for an exothermic reaction (BBC, 2011)

An exothermic reaction releases energy to the surroundings, therefore causing the temperature of the surrounding to increase. Therefore, the enthalpy change can be measured by the heat transferred to the surrounding. This can be done by using a calorimeter. The enthalpy change can be measured by burning a known mass of a substance and using heat released to raise the temperature of a known mass of water. Enthalpy change can be measured by using the following formula: Enthalpy change = mass of substance x specific heat capacity x temperature change Q = mc∆t The specific heat capacity of water is given as 4.18 J g-1 K-1. This means that it takes 4.18 joules to raise the temperature of one gram of water by one kelvin.

Hypothesis I predict that the longer the carbon chain in the alkyl group of the alcohol, the more negative the enthalpy of combustion. Propan-2-ol is a 2˚ alcohol and has the longest chain out of methanol and ethanol and will require the most energy to break the intermolecular forces.

Method The calorimeter was clamped upright to avoid any spillages and placed just above the flame from the wick of the spirit burner. 100cm3 of water was poured into the copper calorimeter and the initial temperature of the water was measured by a thermometer and recorded. The temperature of the water was used in order to calculate the enthalpy of combustion for the alcohols. It was imperative to avoid touching the bottom of the calorimeter with the thermometer as this would have given a false reading of the temperature of the water and so the enthalpy change measured would be wrong. The spirit burners were already prepared with half-filled alcohol (methanol, ethanol and propan-2-ol) and fitted with a screw top lid with a rope acting as the wick. The spirit burner containing the methanol was weighed with the cap on the burner and recorded. It was vital to ensure the cap was kept on the burner when it wasn’t lit to avoid evaporation of the alcohol. The wick of the methanol spirit burner was lit and placed underneath the copper calorimeter and the metal draught shield was positioned around the copper calorimeter to avoid any unwanted heat loss and kept the calorimeter insulated. The flame from the spirit burner must only be touching the bottom of the calorimeter. This was to avoid the calorimeter from burning and burning the rubber edges of the clamp. The water was stirred, ensuring the lid of the calorimeter stays put in order to keep it insulated. The water was stirred in order to keep the temperature of the water even in the calorimeter. Once the temperature increased to about 20˚c, the flame of the spirit burner was extinguished by placing the cap on it. This was done immediately to avoid evaporation of the alcohol. The final temperature of the water and final mass of the spirit burner was recorded in order to work out enthalpy change of methanol. This method was repeated with the ethanol and propan-2-ol spirit burners. It is critical to ensure that minimum heat is lost to the surroundings. Quantitative data can be recorded in this way: Table 1: data collected from experiment of the water temperature and mass of spirit burner CH3OH C2H5OH C3H7OH Water Before calorimetry 19.5 20.5 20.0 temperature After calorimetry 40.0 42.0 42.5 (˚c) 212.84 202.68 220.55 Mass of bottle Before calorimetry (g)

After calorimetry

211.21

201.50

219.82

Results Table 2: Temperature of water in calorimeter and mass of each alcohol Alcohols Temperature of water (˚c) Mass of alcohol burnt (g) CH3OH 20.5 1.63 C2H5OH 21.5 1.18 C3H7OH 22.5 0.73 The heat energy transferred can be calculated using the following equation: Q = mc∆T Q = enthalpy change (in J) m = mass of water (in g) c = specific heat capacity (J g-1 ˚c-1) ∆T = temperature change (˚c) For example, in methanol: Mass of water in calorimeter = 100g Specific heat capacity of water = 4.18 J g-1 ˚c-1 Temperature change of water = 20.5˚c Q = 100 x 4.18 x 20.5 Q = 8569 J Q = 8.569 kJ (3 d.p.) In order to determine the enthalpy of combustion of the alcohols, the number of moles of each alcohol must be calculated using the following equation: Number of moles = mass / Mr For example, in methanol: Mass of methanol = 1.63g Mr of methanol = 32.0 Number of moles = 1.63 / 32.0 Number of moles = 0.0509 mol The enthalpy change of combustion can then be calculating using this formula: ∆H˚c = Q / n For example, in methanol: Q = 8.569 kJ n = 0.0509 mol ∆H˚c = 8.569/0.0509 ∆H˚c = 168.3 kJ mol-1 (1 d.p.) Since enthalpy change of combustion is exothermic, ∆H = -168.2 kJ mol-1

Table 3: Enthalpy change of combustion of each alcohol Alcohols

Number of carbons

CH3OH C2H5OH C3H7OH

1 2 3

Enthalpy change of combustion (kJ mol-1) -168.2 -349.7 -773.0

Standard Enthalpy of Combustion (kJ mol-1)

Graph 1: Graph showing the relationship between the standard enthalpy of combustion and the number of carbons in an alcohol molecule

Conclusion/Discus From the results of propan-2-ol has the standard enthalpy shown in both grap indicates that there correlation betwee Nu carbon atoms and the standard enthalpy of combustion. This suggests that the more carbons there are, the more negative the standard enthalpy of combustion. The results support my hypothesis and shows that more energy is required to break the increase in van der Walls’ forces in the homologous series of alcohols.

Graph 2:

The trend of the experimental values is similar to the accepted values trend. However, there is a big gap between the two values and the experimental data is less negative than the accepted values as shown in graph 2. This could be due to excessive heat loss, jeopardising the accuracy of the results. During the experiment, there were many factors that caused heat to be lost from the calorimeter such as instrumental and human inaccuracies as well as unavoidable heat loss via conduction, convection and radiation. The negative linear correlation of the approved values concludes that the number of carbon atoms in an alcohol is directly proportionate to the standard enthalpy of combustion. Therefore, I can conclude that alcohols with a longer alkyl chain such as butan-1-ol would have a more negative enthalpy change of combustion. Despite the big difference in the experimental and approved values, the practical was completed well and the results given follow a similar pattern to the approved values. A bomb calorimeter would be able to produce more accurate values as it is the most precise way to measure the enthalpy of combustion. There are less sources of error as heat loss is prevented by the heater and insulating jacket. There is also a thermostat to maintain the constant temperature of the water in the outer tank as well as the water in the calorimeter. This in turn prevents heat loss as the surroundings of the calorimeter are always

Figure 2: Bomb Calorimeter (BBC, 2016)

at a constant temperature and will not fluctuate. Also, the top is fully covered so heat cannot escape (Lister and Renshaw, 2009).

Evaluation The results proved that propan-2-ol had the most negative enthalpy of combustion. Given both experimental and accepted values shown in graph 2, better experiment results could have been yielded due to the large gap between them. Heat loss from the calorimeter affected the results immensely and many of them were due to practical errors such as not placing the draught shield in a strategic place that would stop the movement of air from causing intense heat loss. Other instrumental inaccuracies included the movement of the calorimetry lid by the metal stirrer therefore releasing some of the heat. Also, the set up for each of the spirit burners differed from one another and so the distance of the tip of the wick and the bottom of the calorimeter would have affected the results. This would mean that the spirit burners with a longer wick would have used up less alcohol to increase the temperature of the water by 20˚c and therefore give unreliable results. Other factors were unavoidable, the lid had holes for the thermometer and the metal stirrer and so heat loss was prevalent. Also, the heat from the spirit burner did not all contribute to raise the temperature of the water, a lot of the heat would be lost in the surroundings in this manner. This was certainly the biggest factor for heat loss which impacted the experimental results. During the experiment, the metal stirrer would be used too vigorously which caused a lot of the heat to escape from the hole of the lid as well as moving the lid about causing further heat loss. This could have easily been improved by handling the equipment with more care as well as being mindful. Also, when cleaning out the calorimeter for the next alcohol spirit burner, the calorimeter would be adjusted differently to the clamp compared to the previous spirit burner and because of the different lengths of wicks, the distance between the bottom of the calorimeter and the tip of the wick would differ greatly between each alcohol tested. This could have been improved by ensuring that the wick length is all the same in the spirit burners and the distance between the wick and the calorimeter should be measured. Also, this experiment took place in a lab that may have fluctuated in temperature and so affecting the heat transferred from the alcohol to the calorimeter. In order to counter this, the practical could have been carried out under controlled temperature. When the temperature of the water rose to 20˚c, it was difficult to extinguish the flame from the spirit burner straight away, and so the short delay caused the loss of alcohol by evaporation which affected results. This could have been avoided by minimising the time to extinguish the flame by being more careful and focused. The 100cm3 of water measured may have not been as accurate as well due to rushing to finish the experiment. This could have been avoided by being more meticulous and measuring the water to 100g to make the measurement more accurate. The experiment could have been improved by determining the enthalpy change of combustion for other organic substances such as alkanes. This would have shown further that alcohols have a greater enthalpy of combustion due to their properties of having strong hydrogen bonding and the comparison between the alkanes and alcohols would highlight

the proportionality in the enthalpy change of combustion and number of carbons in the homologous chain.

Bibliography BBC (2011) Available at: http://www.bbc.co.uk/bitesize/higher/chemistry/calculations_1/potential_energy/revision/ 2/ (Accessed: 22 November 2016). BBC (2016) Available at: http://www.bbc.co.uk/education/guides/z8p72hv/revision/3 (Accessed: 23 November 2016). Cann, P. and Hughes, P.J.E. (2015) Cambridge International AS and A level Chemistry. London, UK: Hodder Education. Lister, T. and Renshaw, J. (2009) AQA Chemistry A2. Oxford, UK: Nelson Thornes. Ryan, L. and Norris, R. (2014) Cambridge international AS and A level Chemistry Coursebook. 2nd edn. Cambridge, UK: Cambridge University Press. Smith, M. and Older, J. (2015) OCR A Level Chemistry Student Book 1. London, UK: Hodder Education. Wooster, M. and Eccles, H. (2008) Revise AS chemistry for OCR A. Edited by Rob Ritchie. 2nd edn. Essex, UK: Heinemann....