of Carbon Dioxide (CO 2

Transcription

1 CHAPTER 10 Phase Changes of Carbon Dioxide (CO 2 ) Objectives This experiment is an introduction to phase changes of pure substances and an introduction to some simple microtechniques for doing experiments in the laboratory. The objectives of this experiment are to: Observe liquid phase of carbon dioxide directly. Use a micromanometer to estimate triple point pressure. Use calorimeter techniques to measure heat of sublimation. Background Solid carbon dioxide is the familiar material dry ice, which vaporizes without forming any liquid under ordinary conditions. This property is a function of temperature and pressure; if the conditions under which we live were very different, dry ice would behave differently. These relationships are often expressed in a phase diagram, a kind of graph which shows the stable phases of a substance at different conditions of temperature and pressure. The phase diagram for carbon dioxide is shown in Figure 1. Liquid Pressure Solid Vapor Temperature Figure 1. Phase diagram for CO

2 Chemistry 141 Experimental Chemistry The pressure and temperature regions where solid, liquid, and gas are stable are labeled as such. One does have to be a little careful about what this graph means! At the surface of a solid or liquid, there are always a few molecules that have enough energy to escape into the gas phase, so solids and liquids always have a film or envelope of gas around them. But as you all know, if the temperature drops below freezing and stays there, all the water outside will gradually freeze into ice and will evaporate very slowly. Dry ice is different in that it does not form any visible liquid under ordinary conditions, and it evaporates rapidly (compared to water ice). How can we interpret this on the phase diagram? The diagram has two variables, temperature and pressure. Pressure is atmospheric pressure when you look at dry ice under normal conditions, so whatever is happening as the CO 2 evaporates must be happening along some horizontal line corresponding to that pressure. Temperature is changing after the dry ice evaporates; it has a temperature of about 80 C, and as soon as it vaporizes it begins to warm up to room temperature. The dry ice is not stable in contact with the room temperature air around it; heat (thermal energy) flows from the air to the dry ice, causing it to evaporate. But the fact that no liquid can be seen during this process tells us one other thing: that the horizontal pressure line which marks atmospheric pressure must be below the point on the phase diagram where the three lines meet, because it is only under these conditions of pressure that liquid does not exist. A possible position of the atmospheric pressure line is shown in Figure 2 the dotted line indicates atmospheric pressure. Liquid Pressure Solid Vapor Temperature Figure 2. Possible position of the atmospheric pressure line in the phase diagram for CO 2. If this reasoning about the relationship between temperature, pressure, and phases of a substance is correct, is there any way of observing liquid carbon dioxide? Yes, there should be, simply by raising the pressure on the system. This is not difficult to do because the necessary increase in pressure is small. This is the first experiment you will do today; detailed instructions and precautions are given in the experimental section. 174

3 Phase Changes of Carbon Dioxide (CO 2 ) Chapter 10 The second part of the experiment is measuring the heat of sublimation of dry ice. Sublimation is the process of evaporation from the solid directly to the gas phase. All phase changes of this kind require some energy, because the molecules in the substances are going from states of limited motion to states of less limited motion. In an earlier experiment, you measured the heat of fusion (melting) of ordinary ice. That was not a very difficult experiment to do, but calculating the heat of fusion was complex because you had to take into account several temperature changes over the course of the experiment. This time the experiment will be a little harder to do because of the temperatures and pressures involved, but the calculations will be similar. Safety Precautions Eye protection must be worn at all times in this experiment. Never handle the crushed dry ice with your bare hands use a spatula or scoop. The potential for minor explosions is high because you are containing carbon dioxide gas under pressure in the first part of the experiment, and the beryl pipettes are barely strong enough to withstand the pressure and may rupture. Materials Equipment Beryl pipettes and scoops, Styrofoam cups, scissors (supplied in lab) Micromanometer, IC temperature sensor, thermometer, tongs (check out from stockroom) Watch glasses (from your locker) Supplies Crushed dry ice Experimental Procedures Part I. Estimation of the Triple Point Pressure 1. Take a new beryl pipette and clip the end off the tube so that you can insert small pieces of dry ice. Use a beryl scoop or a spatula to transfer dry ice into the pipette until you have the bulb about 1/4 full 5 mm or so of chunks. The pipette will frost up as this is done, and there is really no way to stop it. 2. When you have enough dry ice in the bulb, take your tongs and use the part close to the hinge to clamp the tube of the beryl pipette tightly shut. Hold it near the benchtop, away from your face. You will have to hold the tongs tightly with both hands to prevent leaks your TA will demonstrate this to you. If you have the tube tightly closed it will take less than a minute for liquid to begin to form in the bulb. Continue to clamp it until the solid melts, but not after: the pressure rises very rapidly at that point, and usually the beryl pipette bursts. If you release the pressure on the liquid by relaxing your grip on the tongs, the CO 2 will solidify very suddenly to a white powder. So long as the beryl pipette is intact you can repeat this experiment several times with the same sample. 175

4 Chemistry 141 Experimental Chemistry 3. To make an estimate of the triple point pressure, you need some way of measuring the pressure. This will be done with a micromanometer: a small glass tube, sealed at one end and graduated, with a pigtail for handling and a drop of colored water as a marker. Take one of these devices by the pigtail and drop it down the tube of your Beryl pipette. Be sure that you still have a fair amount of dry ice in the bulb; you may wish to add a little more before inserting the micromanometer. 4. Use the pigtail to position the manometer so that you can close the tube with the tongs without breaking the manometer, and so that you can see the colored water drop and the scale. Read the position of the water drop before you clamp the tongs shut read the distance in cm from the closed end of the tube (you should be able to estimate to the nearest mm). Clamp the tube closed again with your tongs and watch the colored drop. If the tube is tightly closed you will see the drop move quickly toward the closed end of the manometer tube. Watch the sample again for the formation of liquid; when you have some visible liquid, and there is still some solid left, read the manometer again. You are making this reading when solid, liquid, and gas are all present in your sample, so it is called a triple point pressure. We will explain in the data analysis section how to calculate this pressure from these two readings and the barometric pressure. Be sure that you record the barometric pressure for the day before you leave the lab! Your TA will have it written on the blackboard. Data Analysis To calculate the pressure at the triple point you need to use Boyle s law in the form: 176 P 1 V 1 5 P 2 V 2 When you take your first reading the volume inside the manometer tube, V 1 is proportional to the length from the closed end, and you can use the reading of the length instead of the volume. Since V i 5 L i A (length cross-sectional area), P 1 L 1 A 5 P 2 L 2 A The cross-sectional area doesn t change, so the A s can be canceled. P 1, the pressure before you close the pipette tube, is just atmospheric pressure for that day. P 2 is the pressure at the triple point, which is what you want to calculate. P 2 5 P 1 (L 1 /L 2 ) Report your experimental triple point pressure both in mm Hg and in atmospheres. Use the Data Summary sheet at the end of this experiment to record your data, calculations, and results. Part II. Measuring the Heat of Sublimation of Carbon Dioxide 1. Program the MicroLAB to measure time and temperature, with a one second delay. Calibrate your temperature sensor. Record the high calibration temperature in your notes. After the temperature sensor is calibrated fill a nested pair of dry

5 Phase Changes of Carbon Dioxide (CO 2 ) Chapter 10 Styrofoam cups with fresh hot water, insert the sensor, and cover the cup with a watch glass. From here on everything needs to be done as quickly and efficiently as possible. A sample of dry ice must be weighed and transferred to the calorimeter with as little loss as possible. The program needs to be running as the dry ice is added, so that you can see the entire interval of the temperature change. 2. Weigh a dry Styrofoam cup and record that weight on the data sheet. Place about 1 1/2 tablespoons of crushed dry ice in the cup you have weighed, and cover with a watch glass to minimize condensation. Start the program and see that the temperature is plotting on-screen, and that the temperature is a little below your high calibration temperature. Take your cup of dry ice to the balance, weigh it without the watch glass, cover it again and return quickly to your workstation. Immediately add the dry ice to the warm water. Be sure that all of the dry ice transfers; some of it may stick to the bottom of the cup. Carbon dioxide gas will evolve vigorously for a few seconds, and the water will cool rapidly. Continue monitoring the temperature until it stabilizes. 3. Remove the thermistor from the calorimeter, carefully shaking any drops of water back into the calorimeter. You still have to weigh the calorimeter to determine how much water it contained during the experiment. (Why are you doing this after the experiment instead of before?) Weigh the calorimeter and its contents, empty it out, dry it gently with a paper towel, and weigh the two nested cups so you can determine the weight of water you had present. 4. Save your data to a disk file, then repeat the procedures to get an idea of how reproducible your measurements are. Use the Lab Works spreadsheet to determine the water temperatures before and after adding dry ice. Data Analysis As in the experiment where you measured the heat of fusion of water (Flames, Heat, and Calories), the heat loss from one part of the system must equal the heat gained by other parts. (You may want to review the Background and Data Analysis sections of that experiment.) Heat lost by water 5 Heat gained by dry ice There are several small heat losses that we are neglecting, just as we ignored the loss of vaporized CO 2 between the balance room and your workstation. You can ignore the heat that went to warming the CO 2 gas. Most of the gas escapes immediately, while it is still very cold, and does not take very much heat from the water. To calculate the heat lost by the water you will need the temperature change and the mass of the water: q 5 s 3 m 3 ΔT, where q is the heat loss in calories, s is the specific heat of water (1.000 cal/g- C), m is the mass of the water sample, and ΔT is the water temperature change. 177

6 Chemistry 141 Experimental Chemistry The water heat loss approximately equals the heat gained by the dry ice sample to change it from a solid to a gas. For your lab report, you need to calculate the heat of sublimation per gram of dry ice, called the specific heat of sublimation (in cal/gram of CO 2 and the molar heat of sublimation (in Joules/mole of CO 2 ). One calorie of heat is equal to Joules, and the molar mass (molecular weight) of carbon dioxide is 44 grams/mole. Can you write an equation relating the specific heat of sublimation and the molar heat of sublimation? Show your calculations and report your results on the Data Summary sheet. Pre-lab Questions Make sure you can answer these before you come to lab. 1. What information does a phase diagram give you? 2. In the heat of sublimation portion of this experiment, suppose you had 75.2 g of water that initially was 64 C. Upon the addition of 11.6 g of CO 2 the temperature reached 18 C. Calculate the heat lost by the water. 3. Use your answer from Question #2 to calculate the specific heat of sublimation in calories per gram of CO

7 Phase Changes of Carbon Dioxide (CO 2 ) Chapter 10 CHAPTER 10 Name ID No. Worksheet Instructor Partner s Name (if applicable Course/Section ate (of Lab Meeting) Data Summary Part I. Estimation of the Triple Point Pressure Barometric pressure: mm Hg Micromanometer reading before clamping beryl pipette (L 1 ): mm Micromanometer reading at the triple point (L 2 ): mm Calculated triple point pressure: mm Hg; atm Calculations: Part II. Heat of Sublimation of Carbon Dioxide irst Trial Second Trial Mass of Styrofoam cup and crushed dry ice g g Mass of Styrofoam cup g g Mass of dry ice g g Mass of nested Styrofoam cups and water g g Mass of nested Styrofoam cups g g Mass of water g g (Continued on back) 179

8 Chemistry 141 Experimental Chemistry Temperature of water before adding dry ice Final temperature of water Temperature change of water Calculations (report results at bottom of page) irst Trial Second Trial C C C C C C Heat loss: irst Trial Second Trial Specific heat of sublimation: irst Trial Second Trial Molar Heat of sublimation: irst Trial Second Trial irst Trial Second Trial Average Heat loss from water cal cal Specific heat of sublimation cal/g cal/g cal/g Molar heat of sublimation J/mole J/mole J/mole 180