Thermodynamics Lab Report

Introduction In this lab, the purpose was to verify Hess’s Law. Four main topics were covered during this experiment including enthalpy of reaction, heat of formation, Hess’s Law, and calorimetry.
The enthalpy of reaction, ΔHrxn is the heat or enthalpy change for a chemical reaction. The energy change is equal to the amount of heat transferred at a constant pressure in the reaction. The change represents the difference in enthalpy of the products and the reactants and is independent of the steps in going from reactants from products. The heat of formation (ΔH°f), which is also known as standard enthalpy of formation, is defined as the change in enthalpy that accompanies the formation of one mole of a compound from its elements with all substances in their standard states. Hess’s Law was a concept that the lab concentrated on. Hess’s Law states that if a reaction can be carried out in a series of steps, the sum of the enthalpies for each step equals the enthalpy change for the overall reaction. For example, the three chemical equations used throughout the experiment were: Equation 1: NaOH(aq) + HCl(aq) NaCl(aq) + H2O(l) Equation 2: NH4Cl(aq) + NaOH(aq) NH3(aq) +NaCl(aq) + H2O(l) Equation 3: NH3(aq) + HCl(aq) NH4Cl(aq)
Resulting in Equation 1 plus Equation 2(reversed) will equal Equation 3. In this case, Hess’s Law gave the ΔH for Equation 3. Calorimetry is the science of measuring heat, and is based on observing the temperature change when a body absorbs or discharges energy as heat. A calorimeter is the device used experimentally to determine the heat associated with a chemical reaction.
Methods and Materials
A calorimeter made of two nested Styrofoam cups and a cover with a hole big enough to put a thermometer through was set-up on the lab table. 50.0 mL of distilled water was measured using a 50-mL graduated cylinder and was then transferred to the calorimeter. After the distilled water was put into the calorimeter, the cover was placed on top. The lab procedure called for the calorimeter to be placed on a magnetic stirrer and then to add a magnetic strring bar set to spin slowly. Because the magnetic stirrer and stirring bar was not available at the time this lab was conducted, a popsicle stick was used to stir the distilled water. After stirring the water for a couple of seconds, the thermometer was placed through the hole of the cover and the temperature was measured. The temperature was then recorded in the Part 1 Data Table (Table 1.1). Using the a hot plate, approximately 75mL of distilled water placed in a 250-mL beaker was heated to about 70 ˚C. 50.0 mL of the 70 ˚C distilled water is measured in a 50-ml graduated cylinder. The termperature of the 50.0 mL of hot distilled water is measured using the thermometer and then recorded in the Part 1 Data Table (Table 1.1). Immediately after the the temperature was measured and recorded, the hot water was poured into the room temperature water in the calorimeter. The cover was placed over the calorimeter, the thermometer was inserted into the hole in the cover and the popsicle stick was used to stir the water. As the water was being constantly stirred using the popsicle stick, the temperature was measured and recorded in the Part 1 Data Table (Table 1.2) every 20 seconds for a total of 3 minutes. After the 3 minutes, the calorimeter was emptied into the sink in the lab table. Once the contents of the calorimeter was emptied, it was dried thoroughly.
50.0 mL of a 2.0M HCl solution was measured in a 50-mL graduated cylinder and then transferred to the calorimeter. The temperature of the HCl solution is measured and then recorded in the Part 2 Data Table (Table 2.1). The 50-mL graduated cylinder is then rinsed with distilled water to avoid the mixing of different solutions. 50.0 mL of a 2.0 M NaOH solution is measured in a 50-mL graduated cylinder. The temperature of the NaOH solution is measured and recorded in the Part 2 Data Table (Table 2.1). Once again, the lab procedure called for the magnetic stirrer and stirring bar used but since it was not available at the time the lab was performed, a popsicle stick was used instead. The 50.0 mL of 2.0M NaOH solution was quickly added to the calorimeter (which already contained the HCl solution) and was covered. The thermomemter was inserted through the hole in the cover and a popsicle stick was used to stir the mixture. While being constantly stirred, the temperature of the solution was measured and recorded in the Part 2 Data Table (Table 2.2) every 20 seconds for a total of 3 minutes. The calorimeter, thermometer, popsicle stick, and graduated cylinder used for Reaction 1 was thoroughly rinsed and dried. This process was repeated using a 2.0M NH4Cl solution and a 2.0M NaOH solution. This procedure required to be performed in the fume hood. The temperatures of NH4Cl and NaOH were recorded in the Part 2 Data Table (Table 3.1). The temperatures of the mixed solutions were recorded every 20 seconds for 3 mintues in the Part 2 Data Table (Table 3.2). Again, The calorimeter, thermometer, popsicle stick, and graduated cylinder used for Reaction 2 was thoroughly rinsed and dried. The process used in Reaction 1 and Reaction 2 was repeated again for Reaction 3 using 2.0M NH3 solution and 2.0M HCl solution instead. The temperatures of NH3 and HCl were recorded in the Part 2 Data Table (Table 4.1). The temperatures of the mixed solutions were recorded every 20 seconds for 3 mintues in the Part 2 Data Table (Table 4.2) This procedure required to be performed in the fume hood. Once all information was recorded in the data tables, the solutions were discarded in containers at the front of the lab classroom as directed by the teacher.
Results
Part 1 Data Table – Determination of Heat Capacity of the Calorimeter
Initial temperature (˚C) [Table 1.1]
50.0 mL H2O (room temperature)
21.2 ˚C
50.0 mL H2O (heated)
59.6 ˚C

Mixing Data [Table 1.2]
Time (sec)
Temperature (˚C)
Time (sec)
Temperature (˚C)
20
41.7
120
41.1
40
41.6
140
41.0
60
41.4
160
41.0
80
41.2
180
40.9
100
41.2

Tmix = 41.8 ˚C Tavg = 40.4 ˚C qcal = -292.6 J Ccal = -14.2 J/ ˚C

????; speciq) - ---> data tables, the solutions were discarded in containers at the front of the lab classroom
Part 2 Data Table – Determination of Heats of Reaction
Reaction 1: HCl(aq) + NaOH(aq)  NaCl(aq) + H2O(l)
Initial temperature (˚C) [Table 2.1]
50.0 mL 2.0 M HCl
21.2 ˚C
50.0 mL 2.0 M NaOH
21.9 ˚C

Mixing Data [Table 2.2]
Time (sec)
Temperature (˚C)
Time (sec)
Temperature (˚C)
20
34.2
120
34.0
40
34.2
140
33.9
60
34.0
160
33.8
80
34.0
180
33.7
100
34.0

Tmix = 34.2 ˚C ΔH = kJ/mol qrxn = -5245.9 J qrxn = -[(grams of solution * specific heat of solution * (Tmix – Tinitial) + (Ccal * (Tmix – Tinitial))] qrxn = -[(103 * (4.18 J/g* ˚C) * (34.2 –21.6) + (-14.2 J/ ˚C * (34.2 – 21.6))] = -5245.9 J

Reaction 2: NH4Cl(aq) + NaOH(aq)  NH3(aq) + NaCl(aq) + H2O(l)
Initial temperature (˚C) [Table 3.1]
50.0 mL 2.0 M NH4Cl
21.4 ˚C
50.0 mL 2.0 M NaOH
21.7 ˚C

Mixing Data [Table 3.2]
Time (sec)
Temperature (˚C)
Time (sec)
Temperature (˚C)
20
22.6
120
22.6
40
22.6
140
22.6
60
22.6
160
22.6
80
22.6
180
22.6
100
22.6

Tmix = 22.6 ˚C ΔH = -4.2 kJ/mol qrxn = -416.3 J qrxn = -[(grams of solution * specific heat of solution * (Tmix – Tinitial) + (Ccal * (Tmix – Tinitial))] qrxn = -[(103 * (4.18 J/g* ˚C) * (22.6 –21.6) + (-14.2 J/ ˚C * (22.6 – 21.6))] = -416.3 J

Reaction 3: NH3(aq) + HCl(aq)  NH4Cl(aq)
Initial temperature (˚C) [Table 4.1]
50.0 mL 2.0 M NH3
21.2 ˚C
50.0 mL 2.0 M HCl
21.0˚C

Mixing Data [Table 4.2]
Time (sec)
Temperature (˚C)
Time (sec)
Temperature (˚C)
20
34.4
120
34.0
40
34.3
140
34.0
60
34.2
160
34.0
80
34.2
180
33.9
100
34.1

Tmix = 34.5 ˚C ΔH = -55.6 kJ/mol qrxn = -5578.9 J qrxn = -[(grams of solution * specific heat of solution * (Tmix – Tinitial) + (Ccal * (Tmix – Tinitial))] qrxn = -[(103 * (4.18 J/g* ˚C) * (34.5 –21.1) + (-14.2 J/ ˚C * (34.5 – 21.1))] = -5578.9 J

Table 1.1 consists of the initial temperatures of 50 mL of distilled water at both room temperature and after heating. Table 1.2 consists of the temperatures of the room temperature distilled water and the heated distilled water mixed together in the calorimeter. This table shows the temperature after every 20 seconds for a total of 3 minutes. Table 2.1 contains the temperatures of the HCl and NaOH solution. Table 2.2 consists of the temperatures of HCl and NaOH mixed together (Reaction 1) at 20 second intervals for 3 minutes. Table 3.1 contains the temperatures of NH4Cl and NaOH. Stored in Table 3.2 are the temperatures of NH4Cl and NaOH mixed together (Reaction 2) at 20 second intervals for 3 minutes. Table 4.1 consists of the temperatures of NH3 and HCl. Table 4.2 contains the temperatures of NH3 and HCl mixed together (Reaction 3) at 20 second intervals for 3 minutes. Tmix is the temperature of a stragiht line (in the figures) extended back to the maximum temperuture at time zero. Tavg is the average initial temperature ([cold temp. + hot temp.] / 2) of the distilled water. qcal is the heat gained by the calorimeter. Ccal is the heat capacity of the calorimeter. ΔH is the enthalpy change of each reaction. qrxn is the amount of heat evolved in each reaction. Figure 1.1 shows the trend (decreasing) of the temperature of the heated and room temp. water for 3 minutes. Figure 2.1 shows the trend (constant, decreasing, constant, decreasing) of the temperature of the HCl and NaOH (Reaction 1) mixed together for 3 minutes. Figure 2.2 shows the trend (constant) of the temperature of the NH4Cl and NaOH (Reaction 2) mixed together for 3 minutes. . Figure 2.3 shows the trend (decreasing) of the temperature of the NH3 and HCl (Reaction 3) mixed together for 3 minutes.
Discussion
The net ionic equations of the three chemical equations used throughout the experiment
Equation 1:
NaOH(aq) + HCl(aq) → NaCl(aq) + H2O(l)
Na+(aq) + OH-(aq) + H+(aq) + Cl-(aq) → Na+(aq) + Cl-(aq) + H2O(l)
OH-(aq) + H+(aq) → H2O(l)
Equation 2:
NH4Cl(aq) + NaOH(aq) → NH3(aq) +NaCl(aq) + H2O(l)
NH4+(aq) + Cl-(aq) + Na+(aq) + OH-(aq) → NH3(aq) + Na+(aq) + Cl-(aq) + H2O(l)
NH4+(aq) + OH-(aq) → NH3(aq) + H2O(l)
Equation 3:
NH3(aq) + HCl(aq) → NH4Cl(aq)
NH3(aq) + H+(aq) + Cl-(aq) → NH4+(aq) + Cl-(aq)
NH3(aq) + H+(aq) → NH4+(aq)
Equation 1 – Equation 2(reversed) = Equation 3
Equation 1 ΔH = -52.5. kJ/mol
Equation 2 ΔH = -4.2 kJ/mol
OH-(aq) + H+(aq) → H2O(l) ΔH = kJ/mol
NH3(aq) + H2O(l) → NH4+(aq) + OH-(aq) ΔH = +4.2kJ/mol
NH3(aq) + H+(aq) → NH4+(aq) ΔH = -48.3 kJ/mol

The purpose of this experiment was met because the group was able to verify Hess’s Law. From the calculations, ΔH3 was found to be -55.6 kJ/mol, but when Hess’s Law was used, ΔH3 was found to be -48.3 kJ/mol. The percent error calculated from these values ended up being 15.1%. Since the calculated answers were as close to the actual, the group was able to verify Hess’s Law.
The first error that occurred in the experiment was that the group had to assume the densities of the solutions. In the calculation selection of the lab handout, it stated to assume the density of the solutions is 1.03 g/mL. Instead of assuming this value, the group could have calculated it to have a more accurate answer. Like the error before, it would have a small effect on the data that was produced. The second error was that the initial temperature measurements for all solutions were not the same. For example, in the first reaction of Part 2, the initial temperature of hydrochloric acid was 24.5 °C and the initial temperature of sodium hydroxide was 23.9 °C. In order to find the initial temperature, the average of these two numbers had to be taken, which could have caused the calculations to be slightly higher or slightly lower than what they should have been. However, this error would have a minor effect on the data since there is not a large difference between the numbers. The third error that occurred in the experiment was that too much or too little water could have been added to the solution in the process of mixing the molar solutions. This could have increased or decreased the molarity of the solution trying to be made. This would affect the results but only in a minor way.