Controlled Variables:
In this experiment, the controlled variables were: the volume of distilled water used i.e. 50 cm3 in both the reactions, the temperature and pressure at which the experiment was conducted i.e. room temperature and pressure in the school laboratory and the time interval after which the temperature of the system was measured and noted down i.e. 0.5 minute or 30 seconds.
MATERIALS AND METHOD
Apparatus Used:
- 1 polystyrene cup
- Glass rod
- Tripod stand
- Electronic mass balance
- 2 spatulas
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A 50 cm3 measuring cylinder
- Ceramic gauze
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A 100 cm3 beaker
- Bunsen burner
- Thermometer
- A stop watch
Chemicals Used:
- Distilled water
- Copper (II) Sulphate pentahydrate
Diagram;
Procedure:
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First I measured 50 cm3 of distilled water using a measuring cylinder and poured it into a polystyrene cup.
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Then I weighed 5.0 g of copper (II) sulphate pentahydrate using an electronic mass balance and a filter paper. I added the 5.0 g of copper (II) sulphate pentahydrate to the 50 cm3 of distilled water already present in the polystyrene cup which already had a thermometer inserted to measure the change in temperatures after every 30 seconds i.e. at intervals of 0.5 minute.
- Thereafter, I weighed 6.0 g of copper (II) sulphate pentahydrate using an electronic mass balance and a filter paper. I completely heated this in a beaker in order to obtain anhydrous copper (II) sulphate.
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Then again I weighed 3.2 g of anhydrous copper (II) sulphate using an electronic mass balance and a filter paper and added it to a polystyrene cup which already had 50 cm3 of distilled water and a thermometer to measure the temperature changes after every 0.5 minute.
Dependent Variables:
The temperature of the system which is the dependent variable in this investigation was measured or obtained by reading from the calibrated scale on the thermometer after a set interval of 0.5 minute or 30 seconds.
Controlled Variables:
The volume of distilled water used i.e. 50 cm3 was kept constant by using a measuring cylinder to measure the volume, the temperature and pressure at which the experiment was conducted was also kept constant by conducting the experiment from the beginning till the end in the same room and the room temperature was monitored and the time interval after which the temperatures of the system were measured and recorded i.e. 0.5 minute was controlled using a digital stop watch.
RESULTS
Data Collection and Processing
Data Collection:
Volume of water = 50 ± 0.5 cm3
Mass of Hydrated Copper (II) Sulphate = 4.99 g
Mass of Anhydrous Copper (II) Sulphate = 3.19 g
Table 1 below shows the changes in the temperature of the system as measured during the duration of the experiment:
Observations:
When the hydrated copper (II) sulphate was heated, its color changed from bright blue to white signifying the loss of water which indicates the formation of anhydrous copper (II) sulphate.
Upon heating the anhydrous copper (II) sulphate, the slight temperature rise caused to beaker to become hot.
Graph 1 showing the Change in Temperatures when hydrated copper (II) sulphate was added to water [Temperature against Time]:
Graph 2 showing the Change in Temperatures when anhydrous copper (II) sulphate was added to water [Temperature against Time]:
Data Processing and Presentation:
Molar Mass of CuSO4 = 159.5 g
Molar Mass of CuSO4 . 5 H2O = 249.5 g
No. of moles of CuSO4 = [Given mass of CuSO4] ÷ [Molar mass of CuSO4]
= [3.19 g] ÷ [159.5 g]
= 0.02 mole
No. of moles of CuSO4 . 5 H2O =
[Given mass of CuSO4 . 5 H2O] ÷ [Molar mass of CuSO4 . 5 H2O]
= [4.99 g] ÷ [249.5 g]
= 0.02 mole
Enthalpy Cycle for the Reaction:
From the enthalpy cycle included in the previous page, the enthalpy change that we are going to determine is ΔH. From the knowledge of Hess’s Law of constant heat summation, we can say:
ΔH + ΔHp = ΔHa
Rearranging the above expression, we can get:
ΔH = ΔHa – ΔHp
Therefore, in order to obtain ΔH, we will have to calculate the difference between ΔHa and ΔHp.
Calculating Change in Temperatures:
Change in Temperature = Highest Temperature – Lowest Temperature
From the data collected while conducting the experiment, the lowest temperature reached = 27 °C. However the estimated lowest temperature obtained from graph 1 attached earlier = 25 °C. Therefore this estimated temperature from the graph compensates for the amount of heat energy gained from the surroundings while conducting the experiment and hence, the lowest temperature reached is 25 °C and this value will be used hereafter in order to carry out the necessary and appropriate calculations.
Change in Temperature for Hydrated Copper (II) Sulphate = (29 ± 0.5) – (25 ± 0.5)
= (4.0 ± 1.0) °C
From the data collected while conducting the experiment, the highest temperature reached = 33 °C. However the estimated highest temperature obtained from graph 2 attached earlier = 35 °C. Therefore this estimated temperature from the graph compensates for the amount of heat energy lost to the surroundings while conducting the experiment and hence, the highest temperature reached is 35 °C and this value will be used hereafter in order to carry out the necessary and appropriate calculations.
Change in Temperature for Anhydrous Copper (II) Sulphate = (35 ± 0.5) – (27.5 ± 0.5)
= (7.5 ± 1.0) °C
Calculating the Change in Enthalpy:
Q = mcΔθ
Heat = mass of solution X specific heat capacity X change in temperature
(J) = (g) (J K-1 g-1) (K or °C)
Calculating ΔHa:
No. of moles of Anhydrous copper (II) sulphate = 0.020 moles
Mass of solution = 50 g
Specific Heat Capacity of water = 4.18 J g-1 K-1
Change in Temperature = (7.5 ± 1.0) °C
Change in Enthalpy (ΔHa) = mcΔθ
= 50 X 4.18 X (7.5 ± 1.0)
= 50 X 4.18 X (7.5 ± 13%)
= 1567.5 ± 13% ≈ 1568 ± 13%
= 1568 ± 204 Joules
Note: The above enthalpy change value obtained is for the total no. of moles of anhydrous copper (II) sulphate which were involved in the chemical reaction i.e. 0.020 moles and not for 1 mole of anhydrous copper (II) sulphate.
Therefore:
0.020 moles = (1568 ± 204 J)
1 mole = x
x = [(1568 ± 204 J) X 1] ÷ [0.020]
= [(1568 ± 13%) X 1] ÷ [0.020]
= (78400 ± 13%) = (78400 ± 10192) Joules per mole
= (– 78.4 ± 10.2 KJ mol-1)
Note: There is a negative sign added to the enthalpy change value above because the reaction was exothermic, which means there was a rise in temperature due to release of energy.
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ΔHa = (– 78.4 ± 10.2 KJ mol-1)
Calculating ΔHp:
No. of moles of Hydrated copper (II) sulphate = 0.020 moles
Mass of solution = 50 g
Specific Heat Capacity of water = 4.18 J g-1 K-1
Change in Temperature = (4.0 ± 1.0) °C
Change in Enthalpy (ΔHp) = mcΔθ
= 50 X 4.18 X (4.0 ± 1.0)
= 50 X 4.18 X (4.0 ± 25%)
= 836 ± 25% = 836 ± 209%
= 836 ± 209 Joules
Note: The above enthalpy change value obtained is for the total no. of moles of hydrated copper (II) sulphate which were involved in the chemical reaction i.e. 0.020 moles and not for 1 mole of hydrated copper (II) sulphate. Therefore:
0.020 moles = (836 ± 209 J)
1 mole = x
x = [(836 ± 209 J) X 1] ÷ [0.020]
= [(836 ± 25%) X 1] ÷ [0.020]
= (41800 ± 25%) = (52250 ± 10450) Joules per mole ≈ (52300 ± 10500) J mol-1
= (52.3 ± 10.5 KJ mol-1)
Note: Since the above reaction was endothermic, the change in enthalpy value is positive.
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ΔHp = (52.3 ± 10.5 KJ mol-1)
Calculating ΔH:
ΔH = ΔHa – ΔHp
Substituting the respective ΔHa and ΔHp values in the above expression, gives the following:
ΔH = (– 78.4 ± 10.2 KJ mol-1) – (52.3 ± 10.5 KJ mol-1)
= ((– 78.4) – (52.3)) ± (10.2 + 10.5) KJ mol-1
= (– 130.7 ± 20.7 KJ mol-1) ≈ (– 131 ± 21 KJ mol-1)
= (– 131 ± 16%) KJ / mole or KJ mol-1
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ΔH = (– 131 ± 21 KJ mol-1)
Conclusion:
After conducting the above mentioned experiment, observing the reactions and after carrying out all the appropriate and necessary calculations outlined in the previous pages, I hereby conclude that the molar enthalpy change for the chemical reaction between anhydrous copper (II) sulphate with water and hydrated copper (II) sulphate with water is (– 131 ± 21 KJ mol-1). This means that the change in enthalpy of formation of hydrated copper (II) sulphate is – 131 KJ per mole with an error of 16%.
The error involved in this experiment is pretty high and that is the reason why there are doubts and uncertainties evolving in our minds as to consider this method or investigation appropriate or not.
The original or literature value for the enthalpy change of formation of hydrated copper (II) sulphate obtained from the internet (source included in the bibliography at the end) was –134 KJ mol-1. Hence, the difference between the literature value and the calculated value = –3 KJ mol-1. Percentage difference between the two values = 2%. Therefore, from the above comparisons, I can hereby conclude that the method used in this investigation was appropriate as the difference obtained between the calculated and the literature value is as small as 2%, hence the calculated value obtained is almost equal to the literature value of the enthalpy change of formation of hydrated copper (II) sulphate.
After carrying out the experiment and the appropriate calculations I would also like to conclude that my hypothesis made earlier on page 2 is correct as there was a rise in temperature observed and noted when the two chemicals reacted with each other hence, proving that the chemical reaction between anhydrous copper (II) sulphate with water and hydrated copper (II) sulphate with water is an exothermic reaction. The heat energy gained due to rise in temperature was then lost to the surrounding as a gradual decrease in temperature was observed with time after the reaction had taken place and because the reaction was exothermic, the enthalpy change value obtained is negative.
Evaluation:
The method and materials used for this experiment was pretty good, but because an error has been experienced, some improvements could be made for a more accurate and correct result. The errors and improvements linked to this experiment include the following:
- Using the electronic balance to measure the mass of the hydrated and anhydrous copper (II) sulphate used was a good idea as it has no uncertainties or errors linked to it. Also, making sure that the volume readings were noted by reading the lower meniscus was a good practice, eliminating human errors and leading to more accurate and reliable readings and results.
- Replacing the analogue thermometer used in this experiment with a digital thermometer can help to reduce the relatively high uncertainty linked to the use of the analogue thermometer. Basically, it was this limitation i.e. use of the analogue thermometer which involved an error of 20% in the calculations carried out earlier. Hence, using a digital thermometer would avoid such errors and enable us to obtain more accurate and reliable reading and results.
- There could have been significant heat losses to the surroundings during the duration of this experiment. Hence, using a well covered and insulated calorimeter instead of an open beaker could avoid such heat losses to the surroundings and therefore not hindering with the readings and results obtained.
- If the calculations carried out and the results obtained are expressed in more significant figures it can lead to more accurate outcomes and results.
Bibliography:
The information included in the general background in the design section of this lab report on page 1 was obtained from the following sources:
The images included on page 3 of this lab report were obtained from the following web link:
The literature value of enthalpy change of formation of hydrated copper (II) sulphate used to compare with the calculated value was obtained from the following web link: