Measuring Heat Changes in Chemical Reactions: Calorimetry and Heat Capacity, Lab Reports of Chemistry

Download Measuring Heat Changes in Chemical Reactions: Calorimetry and Heat Capacity and more Lab Reports Chemistry in PDF only on Docsity! EXPERIMENT #6: THERMOCHEMISTRY INTRODUCTION Overview Thermodynamics is the study of the energy changes that accompany chemical and physical transformations. In this laboratory, we will explore the first law of thermodynamics. A general statement of the first law that the energy of the universe is constant. How do chemists use the First Law? They use it to relate observed external changes in energy to equal-but-opposite internal changes. In particular, atoms and molecules – the ‘stuff’ of chemistry – release heat, light, sound and motion in the course of chemical reactions. These ‘external’ forms of energy can all be measured, and we then assert, in the spirit of the First Law, that the ‘internal energy’ of the atoms and molecules (the ‘system’ in thermodynamic language) must have changed by an equal and opposite amount. The concept of ‘internal energy’ or ‘chemical energy’ The idea that energy is conserved carries with it the thought that if we have a set of atoms and molecules in a known physical state, they will always have the same amount of internal energy no matter how they arrived in this state. Consider, for example, a glass of water sitting on a table. The water in that glass has a certain total internal energy that does not depend on how the molecules were produced. Some, perhaps, came from an ancient star in which H and O atoms reacted with each other, while others were created in the cells of living plants and animals as a by-product of very complicated reactions. According to the First Law, however, their history doesn’t matter: the total internal energy depends only on the current state (Pressure, Temperature, electric & magnetic fields, etc.) of the system. Calorimetry A calorimeter is a device that is used to measure the heat evolved or absorbed ( pq ) accompanying a chemical reaction or physical transformation (the calorie is an old unit of heat, equal to 4.184 J). In this experiment, the calorimeter consists of a stoppered thermos bottle partially filled with water and a temperature probe. The purpose of the thermos is to keep heat from being lost to the room, while the purpose of the water is to absorb the heat produced by the chemical or physical change under investigation. (Water is good for this since it has a high heat capacity, explained below.) In addition to the water, of course, some of the heat is absorbed by the calorimeter (thermos bottle and the temperature probe), and some (a negligible amount in this case) is ‘lost’ through the walls of the thermos bottle to the room. Using the First Law, we can write: ssurrondingpsystem q- q ∆H == where systemH∆ is the enthalpy change of the system (the heat transferred at constant pressure), pq is heat evolved or absorbed in the system, and gssurroundinq is the heat absorbed by the calorimeter (the water + thermos bottle + temperature probe), which is what we measure. But how, exactly, do we measure the heat absorbed by the calorimeter? When you think about measuring heat, you probably think of using a thermometer. But thermometers measure temperature, which is related to the intensity of the random motion associated with heat, not the amount of heat itself. (The analogy is the speedometer of your car, which tells you how fast you are going, but not how much kinetic energy you have.) What we need is the ‘connection’ between changes in temperature and the corresponding changes in ‘internal’ (thermal) energy. It seems reasonable that these two things should be proportional to each other*, so let us write TC q ∆= , where ‘C’ is called the heat capacity of the object whose temperature change we measure. Note that the constant ‘C’ connects ‘heat’ and ‘temperature’, which makes it conceptually as well as numerically important! The phrase ‘heat capacity’ suggests the ability of an object to ‘hold heat’, and that is exactly what it is. Every substance has its own particular set of internal motions (since every substance has a different set of atoms, joined in different ways); thus every substance has a different ‘heat capacity’. Indeed, different amounts of the same substance have different heat capacities, so when discussing the heat capacity of a pure substance it makes sense to specify the ‘heat capacity per gram’ or ‘per mole’. (Physicists and engineers usually think in terms of mass, while chemists think in terms of moles!) The heat capacity per gram is different for every different substance, but the heat capacity per mole of similar materials (all noble gases and most metals) turns out to be very similar. This is one piece of evidence supporting the idea that ‘heat’ is just energy associated with the internal motion of the atoms in the substance. You will check this idea in Part III of the experiment. The units of heat capacity are ‘energy per degree’, i.e. J/˚C, or J/˚K (Celsius and Kelvin have the same size degrees so it doesn’t matter which one you choose). The heat capacity tells us how much the temperature changes for a given amount of heat – for example, if an object has a heat capacity of 20 J/˚K, then adding 20 Joules will raise the temperature 1˚K. The heat capacity per gram is called ‘s’ (the ‘specific heat’), thus sm C ⋅= , where ‘m’ is the mass in grams. The heat capacity per mole is called c (read ‘c-bar’), hence cn C ⋅= , where ‘n’ is the number of moles of the substance. Values for s and c can be found in reference books, for example the CRC Handbook of Chemistry and Physics. ‘s’ can be converted to ‘ c ’ by multiplying by the molar mass; for example, ‘s’ for water is 4.184 J/g-˚C (something you should remember!), thus ‘ c ’ is 18.015g/mol∗4.184J/g.˚C = 75.38 J/mol.˚C. Purpose This part of the experiment has three main parts: * Note: it is possible to have heat absorbed without a temperature change if the object undergoes a 'phase transformation', for example, ice absorbs heat when it melts, but 0 T =∆ . This is a 'special case' which involves only changes in potential energy (separating molecules from each other), rather than kinetic as well as potential energy. 'Temperature' only measures the kinetic energy component of heat. PROCEDURE Check out a LabPro Calculator with a temperature probe from the stockroom. Read the appendix to learn how to operate the LabPro calculator and configure it to show a constant temperature reading on the screen. A thermos bottle with a #12 stopper will be supplied in the laboratory to serve as the calorimeter. The stopper will have a mercury thermometer inserted into a center hole to offer a check on the metal temperature probe that comes with the calculator. There is a side groove in the stopper into which the metal temperature probe is to be inserted to take temperature readings. A. Temperature Measurements 1. This section gives general directions on how to make temperature measurements that will be needed in Parts B., C., & D. 2. Refer to the Appendix for instructions on how to use the LabPro system and prepare the temperature probe for data collection. Once setup is complete, temperature readings will be continually displayed on the calculator (in the main screen). Alternately, a thermometer can be used to make temperature measurements. 3. initialT is equal to the average of the temperature measurements prior to mixing the solutions. After mixing, the temperature of the solution will rise to a maximum value and then will start to cool towards room temperature. finalT is equal to the maximum value. If the temperature of the solution drops after mixing, finalT is equal to the minimum value. B. Finding the Heat Capacity of the Calorimeter 1. Put exactly 200 ml of room temperature deionized water into the thermos bottle and put the #12 stopper with thermometer into the mouth of the thermos. Insert the metal temperature probe through the side grove as far as it will go to monitor the temperature of the liquid. Allow this system (now called the calorimeter) to stand until the temperature remains constant. Record this temperature (Tc = initialT ) to the nearest 0.2˚C. 2. Put exactly 50 ml of deionized water into a 250 ml beaker and cover with a cover glass. Heat the water to ~ 90˚C using a hotplate. Use the mercury thermometer in the bench drawer to monitor the progress of heating. 3. Measure and record the temperature ( hT ), of this hot water with the LabPro temperature probe and immediately transfer all of the hot water into the calorimeter containing the room temperature water. Stir the solution with either the thermometer in the #12 stopper being careful not to break the thermometer bulb, or stir it with the metal temperature probe to ensure good mixing. After 60 seconds, take a temperature reading to 0.2 degree. Record this value as fT ( fT = finalT ). 4. Calculate Ccal: The heat lost ( lostq ) by the hot water is gained by the calorimeter (qcal): C-joules/g 4.18 50g )T - (T q q fhcallost °××== The heat capacity of the calorimeter is therefore the heat gained by the calorimeter divided by the temperature change of the calorimeter: )T - /(Tq C cfcalcal = 5. Repeat the above procedure twice and determine an average Ccal. C. Heat of Fusion of Water 1. Fill the calorimeter with 200 ml of room temperature water and record the initial temperature, initialT (water). 2. Obtain ice cubes from the TA. The TA will say how many cubes are necessary. (Note: Crushed ice is not suitable because it contains appreciable surface water; therefore, do not used crushed ice for this part of the experiment). 3. Dry the ice cube off with a piece of paper towel and immediately add the dried ice cube to the calorimeter. {It is important to minimize the amount of thawed ice (water) added so that only the heat of fusion is measured and not also the cooling due to cold water on the surface of the ice.} Stir until all the ice melts, indicated by the temperature reaching a constant value, finalT (water+ice). If the temperature is decreasing, the ice is still melting, and you need to continue stiring the contents of the calorimeter. 4. Tare an empty Styrofoam cup. Pour the contents of the calorimeter into the cup and reweigh it. Weigh the calorimeter, mass (water + ice), to determine the mass of the ice cube(s). Repeat this procedure twice more. Prepare a table similar to that below in your notebook. 5. Measure the temperature of an ice bath (crushed ice in water) and record this value as the initial temperature of the ice cube, initialT (ice cube). Ccal (J/ oC) Mass of calorimeter contents (100mL water), g Mass of calorimeter contents+ mass of added ice Temperature of water before ice addition, oC Temperature of ice before addition to the calorimeter, oC (= T crushed ice in water) Temperature of the calorimeter after addition of ice, oC APPENDIX OPERATION OF THE LABPRO SYSTEM General Instructions The LabPro System consists of a LabPro unit, a TI-83 graphing calculator, and up to 4 analog & 2 digital probes. Each probe will collect a particular kind of data such as temperature, pressure, absorbance, pH, etc. The following instructions will guide you through the process of setting up the LabPro system for the probes that you will be using. 1. Make sure the TI-83 Graphing Calculator is attached to the cradle on top of the LabPro unit and that the link cables connecting the two are firmly inserted in the cable ports. 2. Plug the AC adapter into an outlet and attach it into the LabPro unit. The unit will turn on automatically and run a self test (accompanied by beeping noises). 3. Turn on the TI-83 calculator (refer to Appendix C for some general instructions for operating the TI-83 calculator). 4. Start the DATAMATE program on the TI-83 calculator. Press PRGM , then select the DATAMATE program from the menu. [Note: To select an item from TI-83 menus, use to highlight the menu item and press ENTER or simply press the number of the menu item on your calculator (in this case, you would press 1 )]. Press ENTER to run the program. 5. After the program begins, an introductory screen will appear and the calculator will try to communicate with the LabPro unit and check for auto-ID sensors. If this communication is successful, the main screen will appear; otherwise, there will be an error message and you will have to make sure the connections are tight before continuing. 6. Prepare the probes that you will be using for data collection by plugging them into one of the numbered channels on the LabPro unit. Note: You should always plug the probes in AFTER turning on the program. 7. Newer sensors that are attached to the LabPro unit (such as the gas pressure sensor, temperature probe, and colorimeter) will be automatically identified and the main screen will display the channels they are connected to and the current readings. 8. Older sensors (such as the pH probe) will have to be manually specified in the program, using the following instructions. Temperature Probe The temperature probe is composed of a thermistor whose resistance changes in response to temperature. The built-in calibration is automatically selected so no calibration steps are required. If you get a reading of "-999.9", remove the probe from the LabPro, wait 10 seconds, and then reinsert the probe while the program is running. Gas Pressure Sensor The gas pressure sensor is composed of a pressure-sensitive membrane. As the pressure changes, there is a change in the voltage across the membrane. While a manual calibration can be done using the pressure sensor, in practice it is difficult to do and the built-in calibration provides adequate results so no calibration steps are required. If you get a reading of "-999.9", remove the probe from the LabPro, wait 10 seconds, and then reinsert the probe while the program is running. pH Probe The pH Probe consists of a pH amplifier and a pH electrode. The pH amplifier should be plugged into one of the channels of the LabPro. Also, verify that the pH electrode is plugged into the amplifier. The pH probe is not auto-detected and will therefore need to be calibrated manually. To perform a manual calibration: 1. After you have selected the pH probe on the calculator (see the General Instructions above), select 2:CALIBRATE. A calibration menu will appear 2. Select 2:CALIBRATE NOW. 3. Rinse the tip of the electrode with distilled water and blot the excess water. The electrode is extremely delicate so use Kimwipes and be careful. 4. Immerse the electrode in standard pH = 4.0 buffer solution so that the liquid level is above the white line on the electrode. Stir briefly and gently with the electrode to remove bubbles from the electrode surface. 5. When the voltage reading stabilizes, press ENTER on the calculator and enter “4” when prompted for a value. 6. Repeat steps 3 through 5 with the standard pH = 10.0 buffer solution. You should enter “10” instead of “4” when prompted. 7. Select 1:OK to return to the setup menu and select 1:OK again to return to the main screen.