Calorimetry 1: The Enthalpy of Formation of MgO Overview Getting Started Techniques Procedure FAQ Full Lab Manual Introduction & Goals Chemistry & Background Key Questions Prelab Problems Safety Procedure In Your Write-up Experiments Index ChemLab Home		Techniques In this experiment, you will use a coffee cup calorimeter, consisting of a double styrofoam cup and a plastic top with a hole for a thermometer. (It's crude, but effective!) The use of a calorimeter is detailed in the ChemLab website, both in the Techniques section and in the Overview pages for this week's experiment. In addition to using the calorimeter properly, key techniques for obtaining accurate results are starting with a dry calorimeter, measuring solution volumes precisely, and determining T accurately. Careful experimenters deal with the first two items easily. The last is somewhat more difficult. The change in temperature is determined by measuring the initial temperature, T1, of the reactants, and the maximum temperature, T2, of the contents of the calorimeter during the exothermic reaction. The determination of a precise value for T2 is complicated by the fact that a small heat exchange occurs between the surroundings and the contents of the calorimeter, both during the reaction and after its completion. The rate of exchange depends on the insulating properties of the calorimeter and on the rate of stirring. A correction for this heat loss is made by an extrapolation of a temperature vs. time curve (see Figure 1 in your lab manual). By extrapolating the linear portion of the curve back to zero time (the time of mixing of the reactants in the calorimeter), T2 is obtained. The value of T2 obtained in this manner is the temperature of the product solution, had the reaction occurred instantaneously with no enthalpy loss to the surroundings. It is good practice to plot temperature vs. time in your notebook as you record the data, during the experiment. You may wish to use a computer program, like Least Squares on the Public file server, the Least Squares Applet on the ChemLab website, or the linear regression on your calculator, to determine the linear equation of your data and extrapolate to zero time more accurately. Instructions for using Least Squares are given at the beginning of this manual on p. 39. Procedure The data collection in this experiment should be done with a partner, but all subsequent computations are to be done independently. When taking data with your partner, be certain to trade jobs, so each partner does all parts of the experiment. Remember to reference your partner's lab notebook when copying original data into your own. The details of how to use the calorimeter are described in the Techniques section. Determining Ctotal HCl + NaOH Clean two 250 mL Erlenmeyer flasks with soap, then rinse and drain. Repeat this until they drain cleanly without leaving droplets adhering to the sides, using purified water for a final rinse. Clean and dry the calorimeter with a towel. Make sure it doesn't leak. From the solutions available in the laboratory, rinse one of the Erlenmeyer flasks with a few mL of 0.500 M HCl and then drain. Rinse the other flask with a few mL of 0.500 M NaOH and drain. Label the flasks to avoid confusion. Measure 100 mL of 0.500 M HCl into the first flask, as exactly as possible using a graduated cylinder. Measure the same amount of 0.500 M NaOH into the second flask. Use one graduated cylinder for HCl only and one for NaOH only; clearly label them. Pour all the HCl into the calorimeter. Remember to add a spin bar and replace the top plus the calorimeter thermometer. When the calorimeter and the solution reach thermal equilibrium, record the temperature vs. time for a few minutes, to the nearest 0.01 °C. Remove the thermometer and cap, touching off any clinging drops, and replace the cap to hold the recorded temperature. Rinse and dry the thermometer, and hold it in the NaOH solution. Adjust the temperature of this solution to match that of the acid within approximately 0.1 °C. This is most easily done by cooling the surface of the NaOH flask under tap water until its contents are a couple of tenths of a degree below the acid temperature, then warming it with one's hands while steadily and carefully stirring with the thermometer. To obtain an accurate temperature measurement, it is essential that the thermometer bulb be completely emersed in the solution. If reaching thermal equilibration takes a long time, you may wish to dry off the thermometer and recheck the acid temperature before the final mix. If the two temperatures differ a bit before mixing, take the initial mixed solution temperature to be the average of the acid and base temperatures. Finally, carefully pour the base into the acid, stir gently with the thermometer, and record the exact time of mixing. Refit the thermometer in the calorimeter, stir gently, and record the temperature every minute for approximately 10 minutes, using a magnifying glass and reading to ± 0.01 °C. Plot temperature vs. time in your notebook as you record the data. For a minute or two, the reading may be variable, but then it should settle down and perhaps exhibit a slow downward drift (maybe 0.02 °/5 min. or so). Plot a temperature vs. time curve and extrapolate this drift back to the time of mixing, to determine the true value of T2, the final temperature. You can do this most accurately using Least Squares or another linear regression. Determining for Mg + HCl Clean and dry the calorimeter, spin bar and thermometer. Using the graduated cylinder reserved for HCl, add 200 mL (measured as accurately as possible) of 0.500 M HCl to the calorimeter. Stir the solution with the thermometer. Replace the thermometer and cap (put the cap over the top, not the bulb, of the thermometer) and allow the calorimeter and the solution to come to thermal equilibrium. Record T vs. time for a few minutes, to the nearest 0.01 °C. Meanwhile, measure a 10-15 cm length of magnesium ribbon. This should be about 0.10-0.12 g, which is approximately 0.005 mol. Scour it lightly with a nylon abrasive pad to remove any surface oxide. Weigh the ribbon to the nearest 0.0001 g, fold it in half, and wrap it completely (around the ends too) with a spiral of copper wire that has been formed around a pencil. Insert the ribbon into the wire spiral and flatten the spiral around it. The copper will not react; its purpose is to prevent the magnesium from floating to the surface. Record the temperature of the solution in the calorimeter (to the nearest 0.01 °C), note the time, add the magnesium to the acid in the calorimeter, and replace the thermometer and cover. Stir gently, and record the temperature every minute for approximately 15-20 minutes until a definite linear decrease in temperature is noted. As before, graph your data in your notebook as you collect them. From your temperature vs. time plot, extrapolate the linear portion of the temperature change back to the time of mixing to determine T2. Determining for MgO + HCl Weigh a 0.6 to 0.7 g sample of MgO to the nearest 0.0001 g. Bottles of MgO must be kept tightly covered to avoid absorbtion of water. Replace the cap when you have weighed your sample. Do not weigh your sample until you are ready to add it to your calorimeter. It may absorb water, while waiting for you to get everything else ready. Clean and dry the calorimeter, spin bar, and thermometer. Using your graduated cylinder, add a fresh charge of 200 mL of 0.500 M HCl to the calorimeter. Replace the thermometer and cap, and allow the calorimeter and solution to come to thermal equilibrium. Record the temperature vs. time for a few minutes. Note the time, add the MgO to the acid in the calorimeter, and replace the thermometer and cover. Stir gently, and record the temperature every minute for approximately 10-15 minutes until a definite linear decrease in temperature is noted. From a plot of temperature vs. time, determine T2 as previously.
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