Coffee Cup Calorimetry: Enthalpy of Neutralization
Coffee Cup Calorimetry: Enthalpy of Neutralization
Due: Thursday November 22, 2012 Due: November 14th, 2012 Introduction The purpose of this experiment is to determine the molar enthalpy of neutralization, and calculate the unknown concentration of the HCl. Changes in energy are always associated with chemical reactions. (Clancy et at., 2011). If energy is liberated in the form of heat the reaction is said to be exothermic. (Reece, 2011). If energy is absorbed, it is endothermic. (Reece, 2011).
Thermochemistry is the branch of chemistry which deals with the gain or loss of heat. (Clancy et at., 2011). A calorimeter is used to measure these changes in heat. (Clancy et at., 2011). It is an insulated container, in which a reaction can occur and the temperature can be measured. (Clancy et at., 2011). A simple coffee cup calorimeter is used in this experiment. The coffee cup is not a sealed system, and therefore the internal pressure is the same as the atmospheric pressure. (Clancy et at., 2011). It can be said that the pressure of the reactants initially, q (heat of reaction) is equal to the final pressure of the system △H (enthalpy change). (Clancy et at., 2011). All strong electrolytes completely dissociate in solutions. As long as the acid and base completely dissociate, the heat of neutralization will not be effected. Neutralization reactions are usually exothermic and have a negative △H for strong electrolytes. They can be modeled using the formula: H⁺ + OH⁻ --> HOH where q= -55.90 kJ per mol of H⁺. (Clancy et at., 2011). At a constant pressure, the heat of reaction is equal to the enthalpy change. Therefore △H is -55.90 kJ per mol of H⁺. (Clancy et at., 2011). For a weak electrolyte, the reaction is endothermic with a q of 25.3kJ/mol.
Experimental Procedure The experiment was outline in the CHEM 120L lab manual, Experiment #4. All steps were followed without deviations. Results Table #1 Average Temperature of the Acid and Base Experiment
Part A Trial #1
Temperature of Acid (°C)
Temperature of Base (°C)
Average Temperature of Acid and Base(°C)
21.8
21.5
21.7
Experiment
Temperature of Acid (°C)
Temperature of Base (°C)
Average Temperature of Acid and Base(°C)
Part A Trial #2
22.1
21.9
22.0
Part B Trial #1
22.2
22
22.1
Part B Trial #2
22.1
21.8
22.0
Part C Trial #1
22.1
22.4
21.8
Part C Trial #2
21.9
22.2
22.1
Part D Trial #1
22.2
21.9
22.1
Part D Trial #2
22.0
21.9
22.0
Table #2 Temperature Recorded Every 1 Second Temperature (°C) Time (s)
Part A Trial #1
Part A Trial #2
Part B Trial #1
Part B Trial #2
Part C Trial #1
Part C Trial #2
Part D Trial #1
Part D Trial #2
1
29.2
30.1
26.5
30.1
23.2
22.9
25.9
30.1
2
30.1
32
28.9
32.1
25.1
24.6
28.6
32.1
3
30.8
35.6
31.2
33.8
26.7
25.8
29.1
33.8
4
36.0
35.8
35.8
30.1
Table #3 Temperature Recorded Every 10 Seconds Temperature (°C) Time (s)
Part A Trial #1
Part A Trial #2
Part B Trial #1
Part B Trial #2
Part C Trial #1
Part C Trial #2
Part D Trial #1
Part D Trial #2
10
34.5
35.8
35.8
33.8
26.6
25.4
30.1
30.9
20
34.5
35.7
34.7
33.8
26.5
25.4
30
30.9
Temperature (°C) 30
34.5
35.7
34.7
33.8
26.5
25.4
30
30.9
40
34.2
35.7
34.7
33.7
26.3
25.2
30
30.8
50
34.1
35.6
34.7
33.5
26.3
25.2
29.9
30.6
60
34.1
35.6
34.7
33.4
26.3
25.1
29.9
30.6
Table #4 Temperature Recorded Every 30 Seconds Temperature (°C) Time (s)
Part A Trial #1
Part A Trial #2
Part B Trial #1
Part B Trial #2
Part C Trial #1
Part C Trial #2
Part D Trial #1
Part D Trial #2
30
34
35.6
34.7
33.3
26.3
25.2
29.9
30.6
60
33.9
35.4
34
33.3
26.3
25.1
29.8
30.4
90
33.8
35.2
33.8
33.3
26.2
25.1
29.8
30.2
120
33.1
35.1
33.7
33.2
26.2
25.1
29.8
30.2
150
33
35.1
33.5
33.2
26.1
25
29.8
30.2
180
32.9
35.1
33.5
33.2
26.1
25
29.7
30.1
210
32.9
35
33.5
33.2
26.1
25
29.7
30.1
240
32.8
35
33.4
33.1
26
25
29.7
30
270
32.8
34.9
33.4
33.1
26
25
29.7
30
300
32.6
34.9
33.3
33.1
25.8
24.9
29.7
28.9
Table 5: Results Of Each Trial Part
Trial
A
1
2.0
2
2.0
△T (°C)
Moles H2O
q (kJ)
q/mol
2.0
14.35
0.08
-5.40
-67.55
2.0
13.7
0.08
-5.16
-64.49
14.03
0.08
-5.28
-66.02
[Base] (M) [Acid] (M)
Average part A: B
1
2.0
2.0
14.10
0.08
-5.31
-66.38
2
2.0
2.0
17.65
0.08
-6.65
-83.13
15.88
0.08
-5.98
-74.75
Average part B: C
1
2.0
0.5
4.65
0.03
1.24
41.44
2
2.0
0.5
3.55
0.03
0.95
31.64
4.1
0.03
1.09
36.54
Average Part C:
Part
Trial
D
1
2
2
2
△T (°C)
Moles H2O
q (kJ)
1.5
8.05
0.06
-3.37
1.5
8.95
0.06
-3.74
8.5
0.06
-3.56
[Base] (M) [Acid] (M)
Average Part D:
q/mol
Discussion In this experiment, a device called a coffee cup calorimeter was used to measure changes in heat. For part A, B and C the heat was determined by multiplying the mass of the acid and base, and assuming they are equal to the density of water (1mL = 1g), then multiplying it by the specific heat of water and the change in temperature. The neutralization of a strong electrolyte has a generally accepted q value of -55.90kj/mol of water. (Clancy et at., 2011). Through this calculation it was found that -66.02kJ was released per mol of water for part A. This is higher than the accepted value of q. Similarly, in part B, like HCl, the HNO3 was a strong acid, and therefore a strong electrolyte with a q value of -55.90kj/mol of water. (Clancy et at., 2011). The calculated value for q, was -74.75kJ per mol of water. In part C, the acid was phenol, which is a weak electrolyte. The accepted q value is 25.3kJ/mol of water. It is positive since it is an endothermic reaction. (Clancy et at., 2011). The value of q in this experiment for part C was found to be 36kJ per mol of water. The deviations for the accepted value of q and what was obtained can be a result of the sources of error. One error is that some of the heat released which would increase the temperature goes to other parts of the thermometer, and some radiates into the room. Therefore the max value of the temperature is not as large as the actual temperature increase. In addition, inaccuracy of the temperature readings may have slightly altered the results. An assumption was made that the solution had the same specific heat capacity as water, and the heat absorbed by the coffee cup was ignored.
Conclusion The purpose of this experiment was to determine the molar enthalpy of a neutralization reaction with a strong and weak electrolyte. In addition, the concentration of the unknown HCl was determined. Assuming that the acid and base had the same density as water (1mL = 1g) the heat of reaction was determined using the formula q=mc△T. The final temperature was taken from the graph which was extrapolated, and the initial temperature was the average of the temperature of the acid and the base, before they were mixed. Once the heat of reaction was found, it was converted into kJ. The mols of water were calculated using the volume of the acid or base and the molarity in a limiting reagent problem. Finally the kJ per mol of water was determined by dividing these two numbers. For part D the number of joules released when water was formed for NaOH and HCl was known, and therefore it was used to calculate how many mols of water were formed. Due to the ratios, the mols of water were equal to the initial mols of HCl. The concentration of HCl was determined by dividing the mols by the volume of 0.04L. The concentration of the unknown HCl (#1) was calculated to be 1.5M. This experimental method was not very accurate, since it was quite difficult to make accurate readings of the temperature.
References Clancy, C., Farrow, K., Finkle, T., Francis, L., Heimbecker, B., Nixon-Ewing, B., Schroder, M., & Schroder, T. T. (2011). Chemistry 12. (pp. 396-301). Whitby, ON: McGraw-Hill Ryerson. Department of Chemistry 2012 Cell Biology Laboratory Manual Fall 2012. University of Waterloo, Waterloo. pp. 52-58. Reece, J. B. (2011). Campbell biology (9th ed., p. 69-80). Harlow: Pearson Education.
Due: Thursday November 22, 2012 Due: November 14th, 2012 Introduction The purpose of this experiment is to determine the molar enthalpy of neutralization, and calculate the unknown concentration of the HCl. Changes in energy are always associated with chemical reactions. (Clancy et at., 2011). If energy is liberated in the form of heat the reaction is said to be exothermic. (Reece, 2011). If energy is absorbed, it is endothermic. (Reece, 2011).
Thermochemistry is the branch of chemistry which deals with the gain or loss of heat. (Clancy et at., 2011). A calorimeter is used to measure these changes in heat. (Clancy et at., 2011). It is an insulated container, in which a reaction can occur and the temperature can be measured. (Clancy et at., 2011). A simple coffee cup calorimeter is used in this experiment. The coffee cup is not a sealed system, and therefore the internal pressure is the same as the atmospheric pressure. (Clancy et at., 2011). It can be said that the pressure of the reactants initially, q (heat of reaction) is equal to the final pressure of the system △H (enthalpy change). (Clancy et at., 2011). All strong electrolytes completely dissociate in solutions. As long as the acid and base completely dissociate, the heat of neutralization will not be effected. Neutralization reactions are usually exothermic and have a negative △H for strong electrolytes. They can be modeled using the formula: H⁺ + OH⁻ --> HOH where q= -55.90 kJ per mol of H⁺. (Clancy et at., 2011). At a constant pressure, the heat of reaction is equal to the enthalpy change. Therefore △H is -55.90 kJ per mol of H⁺. (Clancy et at., 2011). For a weak electrolyte, the reaction is endothermic with a q of 25.3kJ/mol.
Experimental Procedure The experiment was outline in the CHEM 120L lab manual, Experiment #4. All steps were followed without deviations. Results Table #1 Average Temperature of the Acid and Base Experiment
Part A Trial #1
Temperature of Acid (°C)
Temperature of Base (°C)
Average Temperature of Acid and Base(°C)
21.8
21.5
21.7
Experiment
Temperature of Acid (°C)
Temperature of Base (°C)
Average Temperature of Acid and Base(°C)
Part A Trial #2
22.1
21.9
22.0
Part B Trial #1
22.2
22
22.1
Part B Trial #2
22.1
21.8
22.0
Part C Trial #1
22.1
22.4
21.8
Part C Trial #2
21.9
22.2
22.1
Part D Trial #1
22.2
21.9
22.1
Part D Trial #2
22.0
21.9
22.0
Table #2 Temperature Recorded Every 1 Second Temperature (°C) Time (s)
Part A Trial #1
Part A Trial #2
Part B Trial #1
Part B Trial #2
Part C Trial #1
Part C Trial #2
Part D Trial #1
Part D Trial #2
1
29.2
30.1
26.5
30.1
23.2
22.9
25.9
30.1
2
30.1
32
28.9
32.1
25.1
24.6
28.6
32.1
3
30.8
35.6
31.2
33.8
26.7
25.8
29.1
33.8
4
36.0
35.8
35.8
30.1
Table #3 Temperature Recorded Every 10 Seconds Temperature (°C) Time (s)
Part A Trial #1
Part A Trial #2
Part B Trial #1
Part B Trial #2
Part C Trial #1
Part C Trial #2
Part D Trial #1
Part D Trial #2
10
34.5
35.8
35.8
33.8
26.6
25.4
30.1
30.9
20
34.5
35.7
34.7
33.8
26.5
25.4
30
30.9
Temperature (°C) 30
34.5
35.7
34.7
33.8
26.5
25.4
30
30.9
40
34.2
35.7
34.7
33.7
26.3
25.2
30
30.8
50
34.1
35.6
34.7
33.5
26.3
25.2
29.9
30.6
60
34.1
35.6
34.7
33.4
26.3
25.1
29.9
30.6
Table #4 Temperature Recorded Every 30 Seconds Temperature (°C) Time (s)
Part A Trial #1
Part A Trial #2
Part B Trial #1
Part B Trial #2
Part C Trial #1
Part C Trial #2
Part D Trial #1
Part D Trial #2
30
34
35.6
34.7
33.3
26.3
25.2
29.9
30.6
60
33.9
35.4
34
33.3
26.3
25.1
29.8
30.4
90
33.8
35.2
33.8
33.3
26.2
25.1
29.8
30.2
120
33.1
35.1
33.7
33.2
26.2
25.1
29.8
30.2
150
33
35.1
33.5
33.2
26.1
25
29.8
30.2
180
32.9
35.1
33.5
33.2
26.1
25
29.7
30.1
210
32.9
35
33.5
33.2
26.1
25
29.7
30.1
240
32.8
35
33.4
33.1
26
25
29.7
30
270
32.8
34.9
33.4
33.1
26
25
29.7
30
300
32.6
34.9
33.3
33.1
25.8
24.9
29.7
28.9
Table 5: Results Of Each Trial Part
Trial
A
1
2.0
2
2.0
△T (°C)
Moles H2O
q (kJ)
q/mol
2.0
14.35
0.08
-5.40
-67.55
2.0
13.7
0.08
-5.16
-64.49
14.03
0.08
-5.28
-66.02
[Base] (M) [Acid] (M)
Average part A: B
1
2.0
2.0
14.10
0.08
-5.31
-66.38
2
2.0
2.0
17.65
0.08
-6.65
-83.13
15.88
0.08
-5.98
-74.75
Average part B: C
1
2.0
0.5
4.65
0.03
1.24
41.44
2
2.0
0.5
3.55
0.03
0.95
31.64
4.1
0.03
1.09
36.54
Average Part C:
Part
Trial
D
1
2
2
2
△T (°C)
Moles H2O
q (kJ)
1.5
8.05
0.06
-3.37
1.5
8.95
0.06
-3.74
8.5
0.06
-3.56
[Base] (M) [Acid] (M)
Average Part D:
q/mol
Discussion In this experiment, a device called a coffee cup calorimeter was used to measure changes in heat. For part A, B and C the heat was determined by multiplying the mass of the acid and base, and assuming they are equal to the density of water (1mL = 1g), then multiplying it by the specific heat of water and the change in temperature. The neutralization of a strong electrolyte has a generally accepted q value of -55.90kj/mol of water. (Clancy et at., 2011). Through this calculation it was found that -66.02kJ was released per mol of water for part A. This is higher than the accepted value of q. Similarly, in part B, like HCl, the HNO3 was a strong acid, and therefore a strong electrolyte with a q value of -55.90kj/mol of water. (Clancy et at., 2011). The calculated value for q, was -74.75kJ per mol of water. In part C, the acid was phenol, which is a weak electrolyte. The accepted q value is 25.3kJ/mol of water. It is positive since it is an endothermic reaction. (Clancy et at., 2011). The value of q in this experiment for part C was found to be 36kJ per mol of water. The deviations for the accepted value of q and what was obtained can be a result of the sources of error. One error is that some of the heat released which would increase the temperature goes to other parts of the thermometer, and some radiates into the room. Therefore the max value of the temperature is not as large as the actual temperature increase. In addition, inaccuracy of the temperature readings may have slightly altered the results. An assumption was made that the solution had the same specific heat capacity as water, and the heat absorbed by the coffee cup was ignored.
Conclusion The purpose of this experiment was to determine the molar enthalpy of a neutralization reaction with a strong and weak electrolyte. In addition, the concentration of the unknown HCl was determined. Assuming that the acid and base had the same density as water (1mL = 1g) the heat of reaction was determined using the formula q=mc△T. The final temperature was taken from the graph which was extrapolated, and the initial temperature was the average of the temperature of the acid and the base, before they were mixed. Once the heat of reaction was found, it was converted into kJ. The mols of water were calculated using the volume of the acid or base and the molarity in a limiting reagent problem. Finally the kJ per mol of water was determined by dividing these two numbers. For part D the number of joules released when water was formed for NaOH and HCl was known, and therefore it was used to calculate how many mols of water were formed. Due to the ratios, the mols of water were equal to the initial mols of HCl. The concentration of HCl was determined by dividing the mols by the volume of 0.04L. The concentration of the unknown HCl (#1) was calculated to be 1.5M. This experimental method was not very accurate, since it was quite difficult to make accurate readings of the temperature.
References Clancy, C., Farrow, K., Finkle, T., Francis, L., Heimbecker, B., Nixon-Ewing, B., Schroder, M., & Schroder, T. T. (2011). Chemistry 12. (pp. 396-301). Whitby, ON: McGraw-Hill Ryerson. Department of Chemistry 2012 Cell Biology Laboratory Manual Fall 2012. University of Waterloo, Waterloo. pp. 52-58. Reece, J. B. (2011). Campbell biology (9th ed., p. 69-80). Harlow: Pearson Education.
