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Calorimetery and Hess's

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Calorimetery and Hess's Calorimetery and Hess’s Law Overview: Calorimetry is the technique used to measure the heat required or evolved during a chemical reaction. Heat has units of joules, so one might expect to be using a joule meter to measure heat changes. Unfortunately, companies do not make joule meters. Instead, a thermometer is used in calorimetery to measure heat exchange between the reaction and the calorimeter. The heat capacity of the calorimeter, C(J/°C), is used to convert the measured temperature change to heat exchange. − qrxn (J)= C (J/°C) Δt (°C) The negative sign in the equation indicates that either 1) the heat is flowing out of the system and into the calorimeter or 2) heat is flowing from the calorimeter and into the chemical products. Is the thermometer part of the system or the surroundings? Experimental Procedure: Work in lab groups of two. Two experiments will be performed using coffee cup calorimeters to gain experience with the technique and the calculations of calorimetery. These experiments can be performed in any order. Your instructor will indicate which experiment your group should perform first. Measuring the Heat of Formation of MgO(s) Calorimetry will be investigated by you lab group by determining the enthalpy change for the formation of magnesium oxide. Mg(s) + ½ O2(g) → MgO(s) This is a spectacular reaction and is used in fireworks and emergency flares. There are hundreds of YouTube videos showing this combustion reaction. One short clip is: http://www.youtube.com/watch?v=tCH3ocXPJwQ If you watched one of these videos you are probably thinking, “This reaction is going to completely melt any coffee cup I have ever seen. It truly would. We will not be running this reaction in lab, but rather running similar reactions that can be combined, using Hess’s Law, to create the above combustion reaction. We will examine the reactions of Mg(s) and MgO(s) with HCl(aq). The two reactions are: Mg(s) + 2HCl(aq) → MgCl2(aq) + H2(g) MgO(s) + 2HCl(aq) → MgCl2(aq) + H2O(l) Which of these two reactions would be considered a redox reaction? Using the enthalpy of formation for water, H2(g) + 1/2O2(g) → H2O(l) ΔHf° = − 285.8 kJ/mol the three reactions above can be combined to yield the enthalpy of magnesium oxide formation. Experimental Procedure for Heat of Formation of MgO: A figure of the polystyrene calorimeter is shown in the figure below Obtain the equipment from your teaching assistant to construct the calorimeter in Figure 1 at your lab table. Gently tighten the clamp on the digital thermometer, otherwise the case will crack. Figure 1: Schematic of the coffee cup calorimeter. Mg(s) + 2HCl(aq) → MgCl2(aq) + H2(g) : In order to determine the energy released (or absorbed) by a chemical reaction, the heat capacity of the calorimeter must be known. In this lab we are going to assume that 100% of the heat released by the reaction is absorbed by the water in the calorimeter. This is strictly not true, a portion of the joules released by the reaction are absorbed by the coffee cup, thermometer and stir bar. This, however, is very small and in fact is balanced to some extent by the loss of heat to the room. With a better calorimeter a reaction of known enthalpy is run to determine the heat capacity of the calorimeter, to begin the experiment. • To prepare the reaction, add 100 mL of 1.0 M HCl to the calorimeter. • Start the stir bar stirring and be sure that the thermometer is submerged in at least 2 cm of the acid solution. Begin recording temperature in the calorimeter. • Measure and cut 10 cm of magnesium ribbon. • Lightly clean the oxide from the magnesium surface with a scouring pad. • Using weigh paper to protect the balance from the reactive magnesium, mass the clean magnesium at the balance. • When the temperature on the calorimeter is constant, record the initial temperature, ti. • Add the magnesium to the calorimeter with the stir bar stirring the solution and replace the lid. Observe the temperature change, recording the maximum temperature, tf. • After the maximum temperature has been recorded, clean the calorimeter by placing the used solution in the waste beaker in the hood. Use the magnetic stir bar wand to extract your stir bar. Rinse and dry the inside of the calorimeter. Be sure not to drop the stir bar in the waste beaker or the sink. Calculate the energy released in the reaction using the specific heat of H2O, sH2O = 4.18 J/(g °C). Assume that the heat capacity of the calorimeter is the same as the heat capacity of the 100 g of water it contains. The energy released is an extensive property of the reaction. Had you added more magnesium more joules would have been released. We will almost never report extensive properties in our laboratory work. For the reaction you have just performed, Mg(s) + 2HCl(aq) → MgCl2(aq) + H2(g) determine the intensive property ΔHrxn in units of kJ/molrxn. Calculate how many moles of the reaction were run in the experiment. MgO(s) + 2HCl(aq) → MgCl2(aq) + H2O(l): The enthalpy change of the second reaction is performed in a similar manner to the first. • Begin by obtaining 100 mL of 1 M HCl(aq) and filling the dry calorimeter. • Begin stirring the solution and making sure that the themomenter is in place. • Measure between 0.6 and 0.7 grams of MgO powder. Record the actual mass to three digits. Replace the cap quickly on the MgO reagent bottle. MgO powder will absorb water quickly. • When you have recorded a stable starting temperature, ti, in the calorimeter, add the magnesium to the calorimeter and replace the lid and thermometer. Observe the temperature change until the maximum value is reached. Record the maximum temperature as tf. Use the temperature change and the specific heat of water to determine the joules released in the reaction. Again, after determining the joules released in your reaction (an extensive property) calculate the enthalply of reaction as an intensive property (kJ/molrxn). Determining ΔHf for MgO(s): The formation reaction of MgO(s) is: Mg(s) + ½ O2(g) → MgO(s) Use the enthalpy change for the two reactions run in the calorimeter and the heat of formation for water, H2(g) + ½ O2(g) → H2O(l) ΔHf = − 285.8 kJ/mol to determine the ΔHf for MgO(s) using Hess’s law. The textbook refers to this as the “indirect method”. A good name considering that we dare not directly run the reaction in our calorimeter. If you have completed both calorimetery experiments, clean your work area and return your dried calorimeter and thermometer to your teaching assistant. Determining the Curie Temperature of a Metal: The second experiment evolves measuring the Curie point of iron. The Curie point is the temperature at which a magnetizable metal undergoes a phase change and becomes non magnetic. For iron this temperature is quite hot. Figure 2 below shows a schematic for suspending an iron bolt to a magnet above a calorimeter. Figure 2: Schematic diagram of the apparatus used for determining the Curie Point in iron. The suspended bolt is heated by a Bunsen burner until loses its ferromagnetism. When this temperature is reached, the iron bolt will drop from the magnet and into the calorimeter below. The temperature of the iron when it fell from the magnet is calculated from the specific heat of iron and the heat absorbed by the water in the calorimeter. qH2O = −qFe • Construct the apparatus shown in the figure above. • Charge the calorimeter with 200 mL of H2O. Check to see that the thermometer will be submerged 2 cm in the water when the lid is in place. Record the initial temperature of the water. • Determine the mass of the iron screw. • Attach the screw to the magnet and place the calorimeter directly below the magnet. Allow for three inches between the bottom of the screw and the calorimeter. • Light the Bunsen burner by turning on the gas and striking a match and placing the lit match over the burner. If you are unfamiliar with lighting a Bunsen burner, ask your teaching assistance for instructions. • Begin heating the iron bolt. Focus on heating from the side and take care not to heat the magnet. It will take a few minutes of heating to raise the temperature above the Curie point. When the iron rises above the Curie point, it will drop from the magnet. • When the iron drops into the calorimeter, place the lid and thermometer on top of the calrorimeter. Observe the temperature change and record the maximum temperature. • Determine the joules absorbed by the water using the temperature change, mass of the water and specific heat of water, sH2O = 4.18 J/(g °C). • Use the heat absorbed by the water and the specific heat of iron, sFe = 0.444 J/( g °C), to determine the temperature change of the iron. • Finally use Δt = tf − ti to determine the temperature of the iron when it dropped into the calorimeter. There are several problems with this being an accurate measurement of the Curie point of iron. Predominantly the heating is not uniform. The bottom of the iron screw is almost certainly hotter than the top. This should be viewed as an upper bound to the Curie point in iron. If you have completed both calorimetery experiments, clean your work area and return your dried calorimeter and thermometer to your teaching assistant. Turn in any notebook pages requested by your instructor as you leave the lab.
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