# 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

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