### Calorimeter

Calorimeter 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.

We Will Write a Custom Essay Specifically
For You For Only \$13.90/page!

order now

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 sod …