Experiment 7: Enthalpy of Formation of Magnesium Oxide

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1 Experiment 7: Enthalpy of Formation of Magnesium Oxide Objective: In this experiment, a simple calorimeter will be constructed and calibrated, and Hess' law of constant heat summation will be used to determine the enthalpy of formation of magnesium oxide, MgO. When magnesium is burned in oxygen, a large quantity of heat is released which results in a large, negative H rxn value: (1) Mg(s) + 1/2 O 2 (g) MgO(s) ; H rxn = large negative number of joules, J The change in enthalpy of this reaction is quite difficult to measure directly by experiment. However, the heat of reaction can be obtained indirectly by measuring the enthalpies of other reactions and using Hess' law. Review of E and H. Virtually all chemical processes involve energy transfer of one type or another, and the branch of chemistry that explores this transfer is called thermochemistry. The energy change of a reaction, symbolized by the expression E rxn, is the sum of energy which is transferred as both heat, q, and work, w: E rxn = q + w For a reaction that occurs at constant pressure, the term enthalpy is used, which is the sum of the energy of all the atoms and molecules plus the work of displacing the environment, that is, the PV work: H = E + PV Thus, the enthalpy change for a chemical reaction at constant pressure is defined by the following equation: H rxn = E rxn + P V In the case of an enthalpy change which takes place at 1 atm of pressure, we use the notation H rxn. In accordance with Hess' Law, enthalpy changes for chemical reactions are additive. If two chemical equations are added together to form a third equation, the enthalpy change of the third equation, H rxn (3), is equal to the sum of H rxn (1) and H rxn (2). When applying Hess' law, it is important to establish a convention for the signs, positive or negative, of the enthalpy changes. An exothermic reaction is one which releases energy and is assigned a negative value. An endothermic reaction is one which absorbs

2 energy and is assigned a positive value. The additive property of H rxn values will be applied in this experiment in order to determine the enthalpy change associated with the burning of magnesium metal in air. For the combustion of magnesium, reaction (1) [previous page], one possible set of reactions is: (2) Mg(s) + 2 H + (aq) Mg 2+ (aq) + H 2 (g) H rxn(2) = negative number (3) MgO(s) + 2 H + (aq) Mg 2+ (aq) + H 2 O (l) H rxn(3) = negative number (4) 1/2 O 2 (g) + H 2 (g) H 2 O (l) H rxn(4) = negative number Reaction (1) is equal to reaction (2) minus reaction (3) plus reaction (4): H rxn (1) = H rxn (2) H rxn (3) + H rxn (4) The enthalpy of formation of liquid H 2 O has been measured and is given by: H rxn (4) = H f (H 2 O) = Joules/mole = kj/mol The enthalpies of reactions (2) and (3) are measurable quantities. Heat is released by the addition of excess strong acid to Mg, thus the quantity H rxn (2) will be a negative number. Similarly, the addition of excess strong acid MgO will release heat, and H rxn (3) will also be negative. Measuring the Heat Capacity of a Calorimeter A calorimeter is a reaction vessel that is well insulated in order to prevent heat transfer between the reaction and its surroundings. As heat is evolved by the reaction inside the calorimeter, the temperature of the contents increases, and it is this temperature change which is measured and related to the enthalpy of the reaction. Thus, all heat effects that are observed in the calorimeter should be the result only of the H value of the reaction or process occurring in the calorimeter. The relationship between the amount of heat absorbed or released by a system, q, and the resulting change in temperature, T = T final T initial, may be expressed in terms of the heat that is absorbed or released by the water and the heat that is absorbed or released by the calorimeter. The reactions you are observing today are both exothermic, so the system is losing heat and the water and calorimeter are gaining heat. q (gained by water) = s water x m water x T water and q (gained by calorimeter) = C cal x T cal At constant pressure H = q p. Therefore the H system = - (s water x m water x T water + C cal x T cal )

3 This is just like what you did last week. A calorimeter must be calibrated to determine its heat capacity, that is, the amount of heat required to raise the temperature of the calorimeter itself by one degree Celsius or Kelvin. The same C cal from last week will be used this week. To measure the ΔH o rxn values for reactions (2) and (3) a small amount of each solid will be added to 100.0mL of HCl. The heat released by the reaction goes into the calorimeter. The ΔH o rxn is then calculated by H o = - (s water x m water x T water + C cal x T cal ). The unit for this ΔH o will of course be in Joules (J). The ΔH o needs to be expressed as a per mole quantity however. Since the amount of solid will limit (limiting reagent) in all cases, the calculated ΔH o needs to be divided by the number of moles of the solid reacted. ΔH o rxn/mole = = - (s water x m water x T water + C cal x T cal ) /moles of reactant (Mg or MgO solid) Measuring ΔT Instead of a thermometer, a temperature probe interfaced to a computer will be used to acquire the temperature readings in this experiment. Because no calorimeter is free from heat leaks, it is possible that you will see some heat loss to the surroundings reflected in your data. Figure 1 below shows a temperature vs. time plot for a calorimetric experiment in which the reaction was exothermic. After the initial sharp rise in temperature when the reaction is begun (at time of mixing, t = 0 seconds), the temperature begins to decrease due to heat loss from the calorimeter. As the temperature equilibrium becomes complete, the loss becomes steady and the temperature decrease is linear. Extrapolation of this linear portion of the curve back to t = 0 seconds gives an accurate value for the maximum temperature attained, T f, if equilibrium were to take place instantaneously with no thermometer lag.

4 Procedure The Temperature will be monitored using the Logger Pro program. Open the program on the Mac and, if the Temperature probe is properly connected, you will see a screen labeled Temperature vs. time. The program is now ready for data acquisition. When you want to acquire data you click on the start button (green arrow). When you are finished collecting a data set, click stop (red square). For reaction (1) you should use ml of 1 M HCl and 0.07 g of Mg For reaction (2) you should use ml of 1 M HCl and.25 g of MgO Record the exact masses of each solid reagent. Run each experiment 3 times. Data You will need to construct your own data tables for this lab. You are going to perform the heat of reaction with Mg three times and the heat of reaction with MgO three times. Carefully consider all of the data you will need for all three trials when you prepare your data tables. Print a copy of your graphs of temperature vs time for each trial to include in your lab writeup. Analysis and Calculations: Use the data you collected to calculate the enthalpy change/mol (ΔH o rxn/mole) for Mg or MgO for each individual trial. ΔH o rxn/mole = - (s water x m water x T water + C cal x T cal ) /moles of reactant (Mg or MgO solid) Don t average the change in temperature (ΔT cal.) and don t average the mass. Do 6 separate calculations. Three for Mg and three for MgO. Show all calculations and clearly indicate your answers. 2. Determine an average value of ΔH o rxn/mole for Mg and MgO from your three trials for each. 3. Using the answer from question 2 along with the enthalpy of formation for water (given earlier in the lab) use Hess s law to calculate the enthalpy of formation for magnesium oxide (see equations 1-4 on pages 1 and 2). 4. Compare your value of the enthalpy of formation of magnesium oxide to the value in your text. Determine a value for percent error. Show all your calculations and clearly indicate your answer. Results See the lab report format for a description of the results table.

5 Discussion The discussion is a qualitative and quantitative description of the experiment. Describe what you did (qualitative) and what the result was (quantitative number). For each value where you determined a percent error indicate a possible source of error (indicate if it was random or systematic) and describe exactly how it affected your data. You should make sure the error is quantifiable and actually relates to your data. Post Lab Question 1. What is another way you could determine the enthalpy of formation of magnesium oxide besides the experimental method you used today?

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