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CHEM 2001: Core Chemical Concepts and Techniques

Experiment 8: Bunsen Ice Calorimeter

Emily Dunn

20777392

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CONTENTS

Aim ......................................................................................................................................................... 3

Theory ..................................................................................................................................................... 3

Materials & Method ................................................................................................................................ 4

Conclusion .............................................................................................................................................. 7

References ............................................................................................................................................... 8

Appendix 1: Raw Data ............................................................................................................................ 9

Appendix 2: Calculations ...................................................................................................................... 11

Appendix 3: Error Calculations ............................................................................................................ 14

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Aim

The aim of this experiment is to determine the standard enthalpy of formation of magnesium

oxide. This was achieved by using a Bunsen ice calorimeter.

Theory

Magnesium oxide forms when magnesium is burned in oxygen. It is an exothermic reaction

and results in a large quantity of heat energy being released

This change in enthalpy is difficult to measure directly by experiment as the reaction occurs

quickly. This means the resulting heat energy is released quickly making it difficult to

quantify. An alternative way of determining the enthalpy of formation is by measuring the

enthalpies of additional reactions and by using Hess’ law.

Hess’ law states that the total enthalpy change during the complete course of a chemical

reaction is the same whether the reaction is made in one step or in several steps [1]. This

allows the enthalpy change for a reaction to be calculated even when it cannot be measured

directly. Hess’ law can be used to determine the overall energy required for a chemical

reaction when it can be divided into synthetic steps that are individually easier to

characterize. This means the enthalpy changes for chemical reactions are additive.

When applying Hess’ law it is important to establish a convention for the signs. Exothermic

reactions release energy and are assigned a negative value. Endothermic reactions absorb

energy and are assigned a positive value.

For the combustion of magnesium, the following set of reactions can be made:

Therefore, the enthalpy of formation of magnesium oxide can be determined by:

This experiment attempts to measure the enthalpy of reaction of both magnesium metal

(Mg(s)) and magnesium oxide (MgO(s)) in hydrochloric acid (HCl). These enthalpy values can

then be used to determine the enthalpy of formation of magnesium oxide using the above

equation.

This will be achieved using a Bunsen ice calorimeter. An ice calorimeter is a simple piece of

equipment used to measure enthalpy changes in fast exothermic reactions. A precision of 5%

is typically expected [2]. The test tube where the reaction takes place is immersed in excess

ice. As the reaction takes place it causes the ice to melt. Under these conditions, the volume

of ice melted can be used to calculate the standard enthalpy of formation of the reaction using

the equations below.

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The following equations were taken from the lab manual [2].

Where

The enthalpy of reaction ( ) can therefore be given by

Where

= enthalpy of fusion of ice = 334 J/g [2]

Lastly, the standard enthalpy of formation ( ) is given by

Materials & Method

Please refer to the Lab Manual for the full experiment procedure [2].

The following quantities were used:

Actual Weight of Mg(s) = 0.1349 ± 0.00005g

Actual weight of MgO(s) = 0.6711 ± 0.00005g

The following concentrations were used:

2M HCl Concentration = 2.005 ± 0.003mol/L

6M HCl Concentration = 6.000 ± 0.1mol/L

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Results

The raw data used in the figures can be seen in Appendix 1.

Figure 1: Graph of Calorimeter Pipette Volume over time for the reaction of Mg(s) in HCl

Figure 2: Graph of Calorimeter Pipette Volume over time for the reaction of MgO(s) in HCl

0

0.1

0.2

0.3

0.4

0.5

0.6

0 2 4 6 8 10 12 14 16 18 20

Vo

lum

e (

mL)

Time (min)

0

0.1

0.2

0.3

0.4

0.5

0.6

0.7

0 2 4 6 8 10 12 14 16 18 20

Vo

lum

e (

mL)

Time (min)

ΔV

ΔV

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Calculations

Detailed calculations can be seen in Appendix 2.

Section A: Reaction of Magnesium Metal in Hydrochloric Acid

Difference in Volume

Mass of Ice Melted

Enthalpy of reaction

Standard Enthalpy of Formation

Section B: Reaction of Magnesium Oxide in Hydrochloric Acid

Difference in Volume

Mass of Ice Melted

Enthalpy of reaction

Standard Enthalpy of Formation

Standard Enthalpy of Formation of MgO(s)

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Discussion

This experiment determines the values for the enthalpy of formation for Mg+2

(aq) and MgO(s).

The literature values for these enthalpies are -462.0 kJ/mol and -601.8 kJ/mol respectively

[3]. The literature values were compared with the experimental values as seen in Table 1.

Table 1: Comparison of Experimental Results to Literature Values (Appendix 3)

Experimental Value -213.9 ± 80.2 kJ/mol -418.6 ± 108.1 kJ/mol

Literature Value -462.0 ± 0.05 kJ/mol [3] -601.8 ± 0.05 kJ/mol [3]

Difference 248.1 ± 80.25 kJ/mol 183.2 ± 108.15 kJ/mol

Percentage Difference 53.7 ± 17.4% 30.4 ± 18.0%

The experimental value deviated from the literature value which is expected to an extent

because of the errors accounted for in the results. However both the literature results are

outside the error bounds of the experimental results.

As the result is dependent on the accuracy of the

result, this is likely

what led to the large amount of error in the value. However, it is unexpected that

the results would be so far off the literature value. This might be due to errors

not accounted for such as improper procedures in the experiment and/or an upgrade in

equipment being required.

Examples of these potential errors include; the scales slightly fluctuating when measurements

were being taken, the dependence on the acid strength being accurate in the calculations and

improper transfer of the reactants into the calorimeter. Additionally, if the magnesium metal

pieces were not cut into small enough pieces, this could slow the rate of reaction, leading to

the assumption that the system had reached steady state, when it had in fact not. Whether or

not the current procedure is acceptable could be confirmed through repeating the experiment

following the same procedure to see if the results continue to remain this far off the literature

values.

The results could also potentially suggest that the literature value is wrong. This is highly

unlikely due to the extensive research that goes into determining a literature value compared

to this experiment which was only performed once.

Another source of potential error is the timing of the experiment. It was difficult to determine

whether the calorimeter had reached a constant steady state which could have led to the

experiment being stopped too early. This would result in an enthalpy lower than what is

actually expected which is consistent with the produced results.

Conclusion

In conclusion, the experimental value obtained for was very different to the

literature value having 53.7% error. This suggests that something went wrong in the

experiment. The value for was closer having 30.4% error, however this is still a

large deviation from the expected value.

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Going forward, I would repeat the experiment multiple times with the same procedure. If the

results continued to remain far from the literature value this would suggest the problem is

with the procedure itself. However, if the replicated results were closer to the expected value

this would suggest that the error in the current results came from random errors and/or

incorrect procedures being followed.

References

1. Atkins, P. and J. de Paula, Atkins' Physical Chemistry2010: OUP Oxford.

2. UWA, Core Chemical Concepts and Techniques Laboratory Manual. CHEM2001,

ed. M. Baker. Vol. School of Chemisty and Biochemistry. 2015.

3. Sottery, T.W., Chemical Principles, Sixth Edition (Masterton, William L.; Slowinski,

Emil J.; Stanitski, Conrad L.). Journal of Chemical Education, 1985. 62(12): p. A325.

4. Ginnings, D.C., An Improved Ice Calorimeter-the Determination of its Calibration

Factor and the Density of Ice at ooe. 1947.

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Appendix 1: Raw Data

Table A1: Section A Raw Data

Time (min) Recorded Volume 1-(Recorded Volume)

0 0.88 0.12

1 0.85 0.15

2 0.83 0.17

3 0.815 0.185

4 0.8 0.2

5 0.79 0.21

6 0.78 0.22

7 0.77 0.23

8 0.76 0.24

9 0.75 0.25

10 0.74 0.26

11 0.7 0.3

12 0.665 0.335

13 0.61 0.39

14 0.55 0.45

14.5 0.52 0.48

15 0.48 0.52

15.5 0.46 0.54

16 0.45 0.55

16.5 0.44 0.56

17 0.43 0.57

17.5 0.43 0.57

18 0.43 0.57

18.5 0.43 0.57

19 0.43 0.57

20 0.43 0.57

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Table A2: Section B Raw Data

Time (min) Recorded Volume 1-(Recorded Volume)

0 0.93 0.07

1 0.915 0.085

2 0.9 0.1

3 0.88 0.12

4 0.865 0.135

5 0.85 0.15

6 0.835 0.165

7 0.82 0.18

8 0.805 0.195

8.5 0.74 0.26

9 0.68 0.32

9.5 0.62 0.38

10 0.575 0.425

10.5 0.55 0.45

11 0.52 0.48

11.5 0.505 0.495

12 0.49 0.51

12.5 0.48 0.52

13 0.47 0.53

13.5 0.46 0.54

14 0.45 0.55

15 0.435 0.565

16 0.425 0.575

17 0.415 0.585

18 0.41 0.59

19 0.405 0.595

20 0.4 0.6

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Appendix 2: Calculations

Section A: Enthalpy of the Reaction between Mg(s) and HCl

Difference in Volume

Note: The error in the volume readings was taken to be half the smallest graduation on the

pipette [2]

Mass of Ice Melted

Note: Errors in Density constants were taken from literature [4]

Enthalpy of reaction

Note: The error in the enthalpy constant is an assumption.

Number of Moles

Hydrochloric acid is the limiting reagent as the magnesium metal is added in excess.

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Therefore

Standard Enthalpy of Formation

Section B: Enthalpy of the Reaction between MgO(s) and HCl

Difference in Volume

Note: The error in the volume readings was taken to be half the smallest graduation on the

pipette.

Mass of Ice Melted

Note: Errors in Density constants were taken from literature [4]

Heat of Reaction

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Note: The error in the enthalpy constant is an assumption.

Number of Moles

Hydrochloric acid is the limiting reagent as magnesium oxide is added in excess.

Therefore

Standard Enthalpy of Formation

Standard Enthalpy of Formation of MgO

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Appendix 3: Error Calculations

Enthalpy of formation of Mg+2

Difference in values

Percentage Difference

Enthalpy of formation of MgO

Difference in values

Percentage Difference