Calorimeter Essay, Research Paper
Introduction
A team was sent to the chemical manufacturing division of a small chemical company to help the technicians with experiments. Since the notes written by the technicians were inaccurate and unfinished, all of the experiments they had preformed needed to redone and documented correctly. The head of the company gave the new team the task of trying to figure out why some chemical reactions caused the reaction vessel to get cold and others caused the vessel to get hot. The group constructed “an apparatus to measure the quantity of thermal energy gained or lost during the chemical reactions” (Bellama, 193). This device was called a calorimeter. A series of different reactions were conducted using two different calorimeters. First, hot and cold water tests were preformed. Based on these results the scientists calculated the heat capacities of the calorimeter. The density and specific heat of pure water were used for these calculations. The other tests that were redone and recalculated were: salts in water, precipitation reactions, and acid base reactions. Then the question of whether the solution absorbed or gave off heat can be answered. Also, whether or not the concentration of an acid base reaction made a difference in the heat absorbed or lost can then be resolved. The goal is to determine if the reactions gave off heat or became cold. The factors that affect heat energy changes were identified (Cooper, 103).
Results
The results for the heat capacities of the calorimeters were determined using the hot and cold water tests. Data was gathered from this experiment and calculations were preformed that resulted in the figures shown in table 1.
Table 1 – Heat Capacity of Calorimeter 1 and 2
Calorimeter (s) Trail 1 Trial 2 Average
1 .03 kJ/℃ .03 kJ/℃ 0.03 kJ/℃
2 .053 kJ/℃ .054 kJ/℃ .054 kJ/℃
Salt in water tests were then done using the salts BaCl2 and NaCl in solid form. Then calculations using the data from the experiment were completed enabling us to determine the results found in Table 2.
Table 2: Caliometer 1 – Change in Heat for Salts in Water
∆H in Trial 1 ∆H in Trial 2 ∆H in Trial 3
BaCl2 0.9 kJ/mol 1.2 kJ/mol 0.9 kJ/mol
NaCl 0.6 kJ/mol 1.2 kJ/mol 0.9 kJ/mol
Table 3: Caliometer 2 – Change in Heat for Precipitation Reactions
∆H in Trial 1 ∆H in Trial 2 ∆H in Trial 3
NaCl & AgNO3 -4.4 kJ/mol -2.8 kJ/mol -2.9 kJ/mol
BaCl2 & Na2SO4 1.4 kJ/mol -1.3 kJ/mol -1.6 kJ/mol
Table 4: Calorimeter 2 – Change in Heat for Strong Acid & Base Reactions
Strong Acid & Base ∆H
(1M) HCl & (1M) NaOH -85.0 kJ/mol
(3M) HCl & (3M) NaOH -63.3 kJ/mol
(6M) HCl & (6M) NaOH -79.3 kJ/mol
Table 5: Calorimeter 1 – Change in Heat for Weak Acid & Base Reactions
Weak Acid & Base ∆H
(1M) CH3COOH & (1M) NH4OH -48.4 kJ/mol
(3M) CH3COOH & (3M) NH4OH -50.7 kJ/mol
(6M) CH3COOH & (6M) NH4OH -54.3 kJ/mol
Discussion
Construction of Calorimeters:
Two calorimeters were constructed using two Styrofoam cups, one placed inside the other lined with aluminum, and a square cardboard lid. An ideal calorimeter was a good insulator and would have a low heat capacity; the perfect heat capacity would be zero. To measure the heat capacity hot and cold water tests were done and recorded for the calculations for the heat capacities, which are needed for later tests. In order to perform the calculations in all of the experiments the team used the density and specific heat of water, which would cause error in our results. The results were shown in Table 1 and a sample calculation is shown below.
Sample Calculation for Calculating Heat Capacity of a Calorimeter:
Heat Of Hot H20 (lost) = Heat Of Cold H20 (gained) + the Calorimeter
Heat Of Hot H20 (lost) = SH of H20 x Mass of Water x ∆T
Heat Of Cold H20 (gained) = SH of H20 x Mass of Water x ∆T
Specific Heat of H20 (SH) = 4.184 J/g ℃
Amount of H20 Used = 50 grams
Ti Cold = 23.1 ℃Ti Hot = 81.4 ℃Tf = 50.3 ℃ ∆TCold = 50.3 – 23.1 = 27.2 ℃∆THot = 50.3 – 81.4 = -31.1 ℃
Heat of Cold H20 = 4.184 J/g ℃ x 50 g x 27.2 ℃ = 5690.24 J
Heat of Hot H20 (lost) = 4.184 J/g ℃ x 50 g x -31.1 ℃ = 6506.12 J
6506.12 J = 5690.24 J + Heat gained by Calorimeter
Heat gained by Calorimeter = 815.88 J
Heat Capacity = Heat gained by the Calorimeter / ∆TCold= 815.88 J / 27.2 ℃ = 30.0 J/℃ ( 0.03 kJ/℃ )
The two calorimeters used were consistent in their heat absorption which is proven by the results in Table 1. Since calorimeter 1 had the lowest heat capacity, it was the best of the two calorimeters
Salts in Water Tests:
The first experiment performed was the heats of reaction for salts in water. The results of the salts in water experiment were shown in Table 2. Sodium chloride and barium chloride were picked to conduct the tests for our salts, and preformed in calorimeter one. The initial temperatures of the salts, however, were not factored into our equations. Instead room temperature was used for the initial temperature of the salts; which caused a discrepancy. A sample calculation of ∆H for a salt in water is shown below. These calculations are, also, the same ones done for the following experiments.
Sample Calculation for ∆H for a Salt in Water (Calorimeter One)
Heat Change in the Reaction = Heat Change of Solution + Heat Change of Calorimeter
Heat Change of Solution = Total Mass of Solution * Specific Heat * ∆T
Heat absorbed by Calorimeter = Heat Capacity * ∆T
Total Mass of Solution = 50 g H2O + 5 g NaCl = 55 g
Specific Heat of H2O = 0.004184 kJ/g ℃
Ti = 17.1 ℃
Tf = 16.9 ℃
∆T = 16.9 ℃– 17.1 ℃ = -0.2 ℃
Heat Capacity of Calorimeter One = 0.03 kJ/℃.
-∆H = (total mass * sp ht * ∆T) + (ht cap * ∆T)
= (55 g * 0.004184 kJ/g ℃ * -0.2 ℃) + (0.03 kJ/℃ * -0.2 ℃)
= -0.052 kJ
= 0.052 kJ
5 g NaCl * 1 mol/58.4 g = 0.086 mol
0.052 kJ/0.086 mol = 0.607 kJ/mol
∆H = 0.607 kJ/mol
The results in Table 2 show that the heat changes are all positive which mean that energy was being added. A reaction that caused a decrease in temperature by removing heat from its surroundings is called endothermic.
Precipitation Reactions:
For the precipitation reactions, our team chose sodium chloride and silver nitrate for the first reaction and barium chloride and sodium sulfate for the second reaction. The reaction between sodium chloride and silver nitrate formed silver chloride, an insoluble salt. According to Umland and Bellama, “All chlorides are soluble [in water] except AgCl and Hg2Cl2” (117). The sodium nitrate would be in solution because all compounds formed of sodium are soluble in water. The reaction between barium chloride and sodium formed barium sulfate, an insoluble salt. According to Umland and Bellama, “All sulfates are soluble [in water] except PbSO4, Hg2SO4, SrSO4, and BaSO4” (117). The sodium chloride would be in solution because all compounds formed of sodium are soluble in water. The balanced equations of the two reactions can be seen below.
NaCl (aq) + AgNO3 (aq) ? AgCl (s) + NaNO3 (aq)
BaCl (aq) + Na2SO4 (aq) ? BaSO4 (s) + 2NaCl (aq)
The results for the precipitation reactions can be seen in Table 3. For our initial temperature, our team used the temperature of the solution already in the calorimeter before we added the second. Our team always placed either sodium chloride or barium chloride, depending on reaction, in the calorimeter before adding the second compound. Calorimeter Two was used to conduct the experiments.
According to Umland and Bellama, “Energy changes associated with exothermic changes [have] a negative sign because energy is lost by the system” (199). As seen in Table 3, all but one of the results gleaned from the experiment were negative. Therefore, the general trend of the results indicate that these two precipitation reactions released heat into their surroundings. The experiment itself further reinforced our interpretation of the results. The calorimeter became warmer as the reaction progressed. “Changes in which the system gives off thermal energy—that is, changes that heat their surroundings—are called exothermic” (Umland and Bellama 194). Thus, according to our experiments, reactions involving these two precipitation reactions we chose are exothermic processes. The reactions caused an increase in temperature by releasing heat from the system.
A discrepant event occurred in the first trial of barium chloride and sodium sulfate. The heat of reaction calculation was a positive 1.37 kJ/mol. None of the other calculations gave a positive number, and our team redid the calculation for the ∆H of the first trial of barium chloride and sodium sulfate again. The result was still positive. A possible cause for this discrepancy is that an error could have been made in the temperature reading during the actual experiment.
A general weakness of this experiment is that, again, our team assumed that the density and specific heat of the solutions were the same as pure water. Also, the initial temperature of the second solution was not factored into our equations. There should not have been much difference, as both solutions should have been at room temperature. However, due to these assumptions, our calculations of ∆H for the precipitation reactions are slightly inaccurate.
Acid and Base Reactions:
Our final experiments involved acid base reactions. The results of the acid base reactions can be found in Tables 4 and 5. For our experiments, our team chose to use acetic acid and ammonium hydroxide for our weak acid and base. For our strong acid and base, hydrochloric acid and sodium hydroxide were chosen. Due to time restraints, our lab instructor told the class to perform only a single run of the reaction at each concentration. The experiment for the weak acid and base were conducted in Calorimeter One, and the experiment for the strong acid and base was conducted in Calorimeter Two.
For the initial temperature, our team used the temperature of the solution already in the calorimeter before we added the second. Our team always placed the acid in the calorimeter first before adding the base.
As seen in Tables 4 and 5, all of the results gleaned from the weak and strong acid base experiments were negative. Therefore, the results indicate that these particular acid base reactions released heat into their surroundings. The experiment itself further reinforced our interpretation of the results as both calorimeters became warmer as the reactions progressed. Thus, according to our experiments, reactions involving these two acid base reactions are exothermic processes. The reactions caused an increase in temperature in the surroundings by releasing heat from the system.
As seen in Table 4, as the concentrations of the weak acid and base increased, so did the heat released by the reaction. One molar solutions gave off the least heat, and the six molar solutions gave off the most heat. Therefore, according to our results, the concentration did affect the amount of heat given off by the reaction in this experiment. The more concentrated the solutions of weak acid and base, the more heat is given off by the reaction. The same general trend is seen in Table 5 with the exception of the one molar solution of strong acid and base. The concentration does have an affect on the amount of heat released by the reaction. The one molar solution should have been the one that gave off the least heat, and the six molar solution should have given off the most heat.
A discrepancy in this experiment is the one molar solution of strong acid and base in Table 5. The ∆H was –85.0 kJ/mol, more than the three and six molar solutions. This number is wrong. The ∆H for the reaction of the one molar solution of strong acid and base should be below that of the three and six molar solutions, as concentration does affect the amount of heat released by the reaction. This errant temperature could be due to a misreading of the thermometer during the actual experiment. This discrepancy might have been solved by conducting more than a single trial of the reaction under each concentration. However, due to time restrictions, our lab instructor gave us permission to conduct only one trial at each concentration and warned us that our results might be a little strange.
A general weakness of the acid base reaction experiments is that, again, our team assumed that the density and specific heats of the solutions were the same as pure water. Also, only one run was conducted on each of the molarities. If more trials had been done, our team would have achieved better results. Additionally, the initial temperature of the base was not factored into our equations. There should not have been much difference, as both solutions should have been at room temperature. However, due to these assumptions, our calculations of ∆H for the acid base reactions are slightly inaccurate.
As shown in Tables 4 and 5, the weak acids and bases used in the experiment gave off less heat than the stronger acids and bases used.
A weakness of the entire study was the assumptions we made in our calculations
of ∆H. Our team assumed that the density and specific heats of our solutions were the same as pure water. The density of pure water is 1.00 g/mL; the specific heat of pure water is 4.182 J/g. According to the lab manual, “This is not strictly true, especially as your solutions get more concentrated” (Cooper 106). None of our experiments used pure, distilled water. For convenience, our team utilized tap water, which is water mixed with various minerals and impurities, to conduct all of the experiments. Therefore, our results are not as accurate as they could have been, but under the time restrictions we were placed under, they are good approximations.
Conclusion
After all of the experiments and calculations were concluded, the team was able to come to tentative assumptions. The data suggests that the salts (used in the salt and water reactions) were endothermic because they absorbed heat from the surroundings. The acid and base reactions, on the other hand, were exothermic because they released heat during the reaction. By comparing the strong acid and base reactions to the weak acid and base reactions, it can be determined that the strong acids and bases gave off more heat. It was also resolved that the concentrations of the acid and bases did effect the amount of heat that was released. It seemed that by increasing the concentration used in the reaction it caused the reaction to release more heat.
Experimental Procedures
Calorimeter Construction:
1. Two Styrofoam cups were obtained
2. A sheet of aluminum was wrapped around one of cups
3. The cup with the aluminum wrapping was placed inside the other cup
4. A square piece of cardboard was cut to function as the lid
5. A hole was poked through the cardboard to allow the insertion of the thermometer
6. The structure was labeled calorimeter one or two with a black marker; so the two different vessels would not be confused during the experiments
7. Then a second calorimeter was constructed using the same steps as seen above
Finding the Heat Capacity of Each Calorimeter:
1. 50 mL of water were measured using a graduated cylinder and put into a calorimeter
2. The temperature was recorded every 30 seconds using the thermometer. This was done for two to five minutes to make sure the temperature had stablized.
3. The temperature of the cold water was recorded as the initial temperature. The thermometer should not be removed.
4. 50 mL of water was again measured into a beaker using a graduated cylinder
5. The beaker was placed on a hot plate and heated for ten minutes
6. It was then removed and the intial temperature was recorded using a different thermometer
7. This procedure should be preformed twice on both calorimeters
Solution of Salts in Water:
1. 50 mL of water was measured in a graduated cylinder and put into a calorimeter.
2. The temperature was measured and recorded as the initial temperature of the water.
3. 5 grams of salt was measured
4. The room temperature was measured and recorded as the initial temperature of the salt
5. The salt was then poured into the calorimeter and the temperature was recorded immediately
6. The temperature was closely monitored and recorded every ten seconds until a rise in temperature was observed
7. This procedure was to be preformed three times with each of the chosen salts
Precipitation Reactions
Ten milliliters of sodium chloride solution were measured in a graduated cylinder and put into Calorimeter Two. The initial temperature of the sodium chloride solution was recorded. Ten milliliters of silver nitrate were measured into a graduated cylinder and put into the calorimeter. The temperature was immediately recorded. The mixture was monitored, and the temperature was recorded every ten seconds until it began to decrease. The mixture was properly disposed of. The above procedure was repeated two more times. All results were observed and recorded. Ten milliliters of barium chloride solution were measured in a graduated cylinder and put into the same calorimeter. The initial temperature of the barium chloride solution was recorded. Ten milliliters of sodium sulfate were measured into a graduated cylinder and added into the calorimeter. The temperature was immediately recorded. The mixture was monitored, and the temperature was recorded every ten seconds until it began to decrease. The mixture was properly disposed of. The above procedure was repeated two more times. All results were observed and recorded. Equations in the lab manual were used to calculate the ∆H for precipitation reaction.