Calorimetry Experiments

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1 Calorimetry Experiments Pre-Lab: Today s laboratory period will include a variety of activities designed to re-familiarize you with safe practices for chemistry laboratory, the space and equipment you will use this semester, and the expectations for the course. Access the facdata folder of Dr. Steel (wsteel ask if you don t know how to find these!) and watch the pre-lab video available for this experiment. Before coming to lab be sure you have answered these questions in your lab notebook (so that you make a copy of them) and then submit the white copy at the start of lab: 1. The specific heat capacity of water is How much heat is absorbed by a 34.0-g sample of water if its temperature increases by 7.25 C? 2. When 4.50 grams of ammonium nitrate were dissolved in water, the heat absorbed was kj. Determine the molar enthalpy of solution of ammonium nitrate this is simply the heat absorbed per mole of ammonium nitrate dissolved. 3. Use Hess s Law and the first two reactions listed below to determine the enthalpy of reaction for the last reaction. N 2 O 5 2 NO + O 2 H = kj NO + O 2 NO 2 H = kj 2 NO 2 + O 2 N 2 O 5 H =? kj Objective: This lab will allow you to have hands-on experience with heats of reaction. You will measure the heat absorbed or given off by means of a calorimeter and use your experimental values to predict the heat transferred for other reactions. Method: For reactions that occur at constant pressure, the heat transferred is a good estimate of the total change in the internal energy of the reaction system. The heat transferred in this instance is known as the enthalpy change. The heat gained in the course of a chemical process or reaction can be computed from the temperature change of the surroundings if the mass of the system and specific heat capacity are known: q mc T General Chemistry II Lab Manual 1

2 In the previous equation: q = heat gained by system (lost by surroundings) m = total mass of the reaction system C = specific heat capacity of the system T = temperature change (T final - T initial ) In this experiment, the reactions will be all carried out in an aluminum calorimeter to prevent heat loss to the surroundings. In many cases you will work with a liquid as part of the system, and thus the specific heat capacity of the system, can be approximated using the specific heat capacity of liquid. In our work today, the liquid in all our systems will be water, which has a specific heat capacity of When using the heat formula above remember that the mass is the total mass of the system (reactants and water combined). The sign on q is significant and is controlled by the sign of T. As defined above, a negative q indicates that heat is flowing from the system to the surroundings, as would be the case in an exothermic reaction. A positive q indicates that heat is flowing into the system from the surroundings, as would be the case in an endothermic reaction. Procedure: The calorimeter used in this experiment normally consists of two aluminum cylinders that are nested in one another with a layer of insulating air or styrofoam between them. For today s work you will replace the inner cup with a beaker this is done to avoid some side reactions that can occur between the aluminum and the acid we will use as part of our reactions. The lid of the calorimeter has two holes in it: one is for a thermometer the other is for a stirring rod in our case we will simply use a magnetic stir bar to accomplish stirring. Overall, these devices are obviously not perfectly insulated, but we will ignore the small amount of heat loss in our measurements and assume that the calorimeter is a perfect thermal insulator. The main goal of the work you do today will be to determine the enthalpy of reaction for the formation of two metallic oxides: magnesium oxide and zinc oxide. Their respective formation reactions are shown below: Mg(s) + O 2(g) MgO(s) Zn(s) + O 2(g) ZnO(s) Unfortunately there are a few obstacles to studying these reactions in the lab with our equipment. First, both reactions produce a large amount of heat perhaps enough to damage our simple calorimeter. Second, depending on the quantities we used, it would be quite 2 Calorimetry Experiments

3 challenging to measure the temperature of the system, even if we had a thermometer with a range large enough to observe it. So, since we cannot directly measure the heat of reaction for these formation reactions, we will make use of Hess s Law to indirectly calculate them. Hess s Law states that the heat of reaction of any reaction can be expressed as the sum of the heats of reaction of any set of reactions that can be summed to give the reaction of interest. In the case of the magnesium oxide formation reaction, we will make use of this group of reactions: Mg(s) + 2 HCl(aq) MgCl 2 (aq) + H 2 (g) MgO(s) + 2 HCl(aq) MgCl 2 (aq) + H 2 O(l) 2 H 2 (g) + O 2 (g) 2 H 2 O(l) The third reaction in this group has a well-known enthalpy of reaction of (You are being given this value due to the difficulty in determining it experimentally.) We will determine experimental values of the other two and combine the three values in the appropriate way so as to give the enthalpy of reaction of our formation reaction. Similarly, we can make use of a set of reaction involving zinc to determine the heat of formation of zinc oxide: Zn(s) + 2 HCl(aq) ZnCl 2 (aq) + H 2 (g) ZnO(s) + 2 HCl(aq) ZnCl 2 (aq) + H 2 O(l) 2 H 2 (g) + O 2 (g) 2 H 2 O(l) 1. With a partner, obtain a calorimeter, 150 ml beaker, digital thermometer, two graduated cylinders (50 or 100 ml), hot plate, and magnetic stir bar. 2. Remove the inner aluminum cup from the calorimeter and replace it with the 150 ml beaker. Place the stir bar into the 150 ml beaker and then place the calorimeter atop the hot plate. 3. Cut two strips of magnesium ribbon, each about cm long. Use some sandpaper to clean the strips of any oxide coating on it it should be shiny before you proceed with step three. 4. Record the mass of each strip of magnesium if the mass is greater than 0.5 grams trim off a small piece so that it is less than this. 5. Add 50.0 ml of 2-M HCl to your calorimeter. Place the calorimeter on the hot plate and stir the contents with the stir bar. AT NO POINT SHOULD THE HEAT BE TURNED ON ON YOUR HOTPLATE! General Chemistry II Lab Manual 3

4 6. Record the temperature of the hydrochloric acid in the calorimeter this value is the initial temperature of your system. 7. Add one of the strips of magnesium to the acid and stir with your thermometer. Observe the temperature of the system and record the highest temperature it reaches. This temperature is the final temperature of your system. 8. Empty the contents of the calorimeter into the waste container labeled Calorimetry Experiments Waste and rinse out the 150 ml beaker. Dry it out with paper towels. 9. Repeat steps 5 8 with your second strip of magnesium. 10. Using two weigh boats, weigh out two separate samples each about 0.5 grams of magnesium oxide and record their masses in your notebook. 11. Once your 150 ml beaker has again been cleaned and dried, add 50.0 ml of 2 M HCl to it. As you have done previously, set the calorimeter atop the hot plate and begin to stir the contents. 12. Record the temperature of the hydrochloric acid in the calorimeter this value is the initial temperature of your system. 13. Add one of the samples of magnesium oxide to the acid and stir continuously. Observe the temperature of the system and record the highest temperature it reaches. This temperature is the final temperature of your system. 14. Empty the contents of the calorimeter into the waste container and rinse out the 150 ml beaker. Dry it out with paper towels and return it to the calorimeter. 15. Repeat steps with your second sample of magnesium oxide. 16. Clean out the 150 ml beaker and reassemble the calorimeter with the stir bar. 17. Repeat steps 3 16 using zinc (mossy) in place of magnesium and zinc oxide in place of magnesium oxide. When doing so, change the masses to approximately 1.0 g of zinc and 1.2 g of zinc oxide. Dispose of your reaction materials after each trial, as you did above. 18. Return your supplies to the appropriate locations and clean-up your work area. 4 Calorimetry Experiments

5 Analysis: Before beginning any of the calculations, it is recommended that you open the electronic version of the lab report sheet that you will complete for this experiment. To do so, follow these steps: From the Start button (bottom left corner of screen), select My Computer Select the facdata2 on 'Files (files)' (N:) server Open the folder for wsteel Open the folder for CHM137 Templates Open the Calorimetry Report Sheet If you took General Chemistry I here at York, you will recall that there are several documents in this folder. These will be the starting points for several of your future lab reports. At this time you may want to copy all these files to your H: drive space or a jump drive in order to save time later this semester. The Calorimetry Report Sheet should be completed entirely electronically there is no reason for you to hand-write any answers on it. The series of calculations described below parallel the report sheet. Begin by recording the appropriate raw data in the report sheet volumes, masses, initial and final temperatures, and the specific heat capacity of water. 1. For each reaction determine the change in temperature. Always subtract the initial from the final temperature. 2. Determine the mass of each of your eight systems (the system is the combination of the liquid solution and the solid you added to it). We can estimate the density of each acid solution as 1.0 g/ml and use this density to determine the mass of liquid used. Then you can simply add the mass of the solid used in each reaction to arrive at the total system mass. 3. Use your system masses, changes in temperature, and the specific heat capacity of water (which is a good estimate for the heat capacity of the entire reaction system, too) to determine the heat produced in each reaction. 4. Determine the moles of solid used in each of your eight trials. 5. Take each value from your results in step 3 above and divide it by its respective mole quantity from step 4 this converts your heats (in joules) to molar enthalpies of reaction (J/mol). Convert you answer to kj/mol at this point. Calculate separate average enthalpies of reaction for the magnesium reactions, magnesium oxide reactions, zinc reactions, and zinc oxide reactions (four averages in all). Record the average enthalpies of reaction in the appropriate areas on the report sheet. General Chemistry II Lab Manual 5

6 6. To determine the enthalpy of formation of magnesium oxide you must combine the reactions you studied (and the water reaction) in such a way so that they combine to give you the formation reaction. Determine how the provided reactions must be combined to add to the target formation reaction. Rather than typing out the entire sequence, simply fill-in a coefficient for each reaction i.e., if you need the first reaction to occur in reverse, it would have a coefficient of According to Hess s Law, whatever you do to a reaction, you should also do to its enthalpy when determining the overall enthalpy. Keeping this in mind, determine the enthalpy of formation of magnesium oxide by combining the known enthalpies of reaction. 8. Use a reference (such as your text or the Handbook of Chemistry and Physics) to find a known value for the enthalpy of formation of magnesium oxide. Be sure to list your reference on your report sheet. 9. Calculate a percent error for the enthalpy of formation of magnesium oxide. 10. Repeat steps 6 9 for the zinc reactions so that you find the heat of formation of zinc oxide. Lab Report: For this lab you should complete your report sheet by answering the remaining discussion questions on the report sheet. Print out a copy of your report (or it to your instructor, if they are willing to accept it in that format). If you worked with a partner, you must each hand in separate reports; your partner s name must be included on your report sheet. Prior to coming to the next lab, be sure to complete the pre-lab exercises for the Alka-Seltzer Analysis experiment. As you hopefully did for this experiment, you should write out your solutions to these questions in your lab notebook, so that you have a carbon copy of your work in case you want to refer to it as you do your calculations or if it should get lost after you hand it in. 6 Calorimetry Experiments

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