Heat of Combustion PCh 6-99

Transcription

1 UMR ChemLabs Heat of Combustion PCh 6-99 Gary L. Bertrand, Professor of Chemistry A calorimeter (calor L = heat + metron Gr = measure) should be literally a device to measure heat. In reality, most calorimeters are designed to minimize the transfer of heat between the system (inside) and the surroundings (outside). The confusion dates back to the caloric theory, in which heat was thought to be a property of materials rather than a transfer associated with the process by which a material undergoes a change of state. There are three basic types of calorimeters: adiabatic - in which the transfer of heat between system and surroundings is prevented, isoperibol - in which a small controlled heat exchange is permitted, and isothermal - in which either work or heat may be used to maintain constant temperature within the system. Essentially, these devices are measuring a heat equivalent for a change of state of the system. This experiment utilizes a commercial calorimeter of the isoperibol type, the Parr Oxygen Bomb Combustion Calorimeter (below). A sample (up to a gram or two) is contained under a pressure (~300 psig) of pure oxygen in a sturdy stainless steel bomb (A, about 1 liter) which is immersed in 2 liters of water (B) in a metal bucket (C). The apparatus contains a thermometer (D) to measure the temperature of the water and a stirrer (E) to insure uniform temperature throughout the system, and is insulated from the surroundings by two air spaces (F) and a non-conducting shell (G). Fig. 1 1

2 The unit is assembled with the temperature of the water slightly below room temperature, and a slight positive slope is observed in a graph of temperaturevs time. When the sample is ignited, the temperature within the bomb increases, then the temperatures of the water and the bucket increase until all are at the same temperature, usually slightly above room temperature. The temperature usually goes through a maximum, then a slightly negative slope is observed in the temperature vs time graph. In the adiabatic version of this calorimeter, hot and cold water are circulated through the outer jacket to match its temperature to that of the water in the bucket, thereby preventing the transfer of heat to or from the bucket. The temperature vs time graph is horizontal before ignition, and again after the temperature increase has been observed. The temperature-time lines (A,B) are extrapolated to a common point in time. B C A Fig 2 [These lines are parallel for the adiabatic apparatus, but not for the isoperibol apparatus] A vertical line is selected which will make the two shaded areas in the graph approximately equal. The temperature change for the experiment ( T) is calculated from the final and initial temperatures on this line. The process is considered to take place instantaneously along the vertical line. The shaded areas (one positive, one negative) are considered to be proportional to the heat leak. Making these areas equal is assumed to nullify the heat leak, making the observed temperature change the same as would be observed in an adiabatic process. The average temperature of the measurement is taken as the temperature at the crossover point (C) between these areas. 2

3 Thermodynamic Analysis: The system (bucket, water, bomb, sample, oxygen) is assumed to be at T initial and some pressure P initial. The sample is ignited and the temperature of the system increases to T final at pressure P final. The bomb and its contents undergo no change in volume, but there is a very slight expansion of the water in the bucket due to the change of temperature. There is zero heat for this (pseudo) adiabatic process. There are two negligible increments of work - that due to the expansion of the water against the atmosphere (which will be compensated later), and that due to the electrical current used to heat the fuse wire which ignites the sample. It is then possible to calculate the heat (Q) that must be transferred in order to cool the system (bucket, water, bomb, reaction products, unburned oxygen) back to T initial. The heat for this cooling is calculated as the heat capacity equivalent of the system (C eq ) multiplied by the change in temperature: Q = C eq (T initial - T final ) = - C eq (T final - T initial ) = -C eq T The heat capacity of the system might be determined in several ways, but the most common method is to measure the temperature change for the combustion of a standard material (usually benzoic acid). This net process may now be considered as: [Sample + Req Oxygen](T initial,p initial,v initial ) -> [Products](T initial,v initial,p )(I) with the remainder of the system (bucket, water, bomb, excess oxygen) undergoing no change. Since this change involves no work, U 1 = Q = -C eq T. The difference between P initial and P is usually small, and the derivative of energy with respect to pressure at constant temperature is also small for most materials (zero for an ideal gas), so the energy change for taking the products back to P initial, T initial, and some new volume is considered negligible (an accurate value can be calculated for extremely precise work). The energy change per mole of sample is [Sample + Req Oxygen](T initial,p initial,v initial ) -> [Products](T initial,p initial,v )(2) with a small correction for combustion of the fuse wire (q corr ). U m = -(C eq T - q corr )/n sample. The molar enthalpy change ( H m )for this reaction is [ H = U + (PV)], H m = U m + P initial (V - V initial )/n sample = -[(C eq T - q corr )/n sample ] + RT N gas (3) in which N gas is the difference between the numbers of moles of gaseous products and reactants per mole of sample reacted. This enthalpy change is for the reaction at the experimental temperature (the average temperature from Fig. 2) and pressure (about 20 atm). For very precise work, this value is adjusted to the standard state conditions (1 atm) with the properties of the materials involved. The correction is normally small, and for this work it is assumed that the observed enthalpy is the standard heat of reaction ( rxn H ). This value is combined with the standard heats of formation of the products (usually water and carbon dioxide) to obtain the standard heat of formation ( f H ) of the sample. 3

4 Assembly: Each group will calibrate their apparatus with a standard sample of benzoic acid and will measure the heat of combustion of one of the samples (lauric acid, myristic acid, palmitic acid, or stearic acid). 1. Make pellets of benzoic acid (0.9 to 1.0 grams) and your sample (0.9 to 1.1 grams), and record weights. 2. Measure about 10 cm of fuse wire and record the weight. 3. If the electrodes are not shiny, clean them up with emery cloth. Attach the wire to the electrodes for good electrical contact as in Fig. 3 below. Arrange the wire so that it is pushing down on the pellet in the center of the cup, with neither the wire or the pellet touching the sides of the cup. If the wire touches any metal between the two electrodes, it will short out and you ll have to take everything apart and start over (TEA-SO). If the pellet shifts in handling so that it doesn t make good contact with the wire, it won t ignite and you ll have to TEA-SO. TEA-SO is another way of saying that we re going to be here a lot longer than we really care to. Shake the electrode assembly to make sure everything will stay in place. Fig 34. Place 1 ml of distilled water in the body of the bomb and assemble. It is important that the vapor is saturated with water before and after the combustion so that any water produced in the reaction ends up as liquid rather than vapor. 5. Pressurize, purge, and re-pressurize the bomb to obtain an atmosphere of pure oxygen. The TA will show you how to do this. Open the valve on the filling hose SLOWLY and bring the pressure up to atm. Open the screw valve on the top of the bomb SLOWLY and bleed the pressure down to 4-5 atm. Close the screw valve and SLOWLY bring the pressure back up to atm. Open the relief valve to quickly close the system and relieve the external pressure. Detach the hose from the bomb. 6. Place the bomb on the raised circle in the bucket. 7. Measure 2 liters of distilled water into the bucket, avoiding air bubbles as you do. The exact quantity of water is not as important as the reproducibility of this quantity in all of your experiments. The temperature of the water should be C. Place the bucket on the footings in the shell. 8. Attach the leads to the bomb. 9. Place the lid on the shell. Rotate the stirring mechanism to make sure it is turning freely. 10.Put the belt on the stirring mechanism and turn on the stirrer. If it is not rubbing on the bucket or the bomb, then you re ready to start taking data. 4

5 Data Collection: 1.Allow the system to equilibrate for 5-10 minutes, then take temperature readings at 60-second intervals for 5 minutes. Make sure the temperature increments are constant. 2.Press the firing button and hold it down for several seconds. Try to observe the flash of the light on the box. 3.Start taking temperature readings every 15 seconds for three minutes. If the temperature does not rise abruptly during this period, you ve got a problem. If the light did not flash in step #2, you may have a correctable open circuit. If it did flash, you may have a correctable short-circuit. Check your connections, take 60-second readings again for 5 minutes and try to fire again. If you still don t get ignition -- TEA-SO! 4.Continue taking readings at 60-second intervals until you have at least six sequential readings with a constant temperature increment. 5.Disassemble the apparatus. Carefully open the screw valve and release the pressure before attempting to unscrew the locking ring. It is possible for particulate matter to be blown out of the valve, so cover the valve with a cloth while venting the bomb. 6.Inspect the bomb for evidence of incomplete combustion (unburned material, soot, etc.). Solid lumps are either iron or iron oxide. Press them with a spatula to crush any iron oxide, reclaim any unburned iron including the wire wrapped around the electrodes, and weigh it. 7.Wipe out the inside of the bomb with a paper towel, inspect to make sure you haven t left any lint or other material, and proceed with the next measurement. Data Processing: 1.Enter your data into a data processing program (Quattro, Excel, etc. Data will be shared as Excel 4.0 files in the RESULTS section of the Web Page. Name your file in the format: PCh2C3B, indicating Experiment #2, Section C, Group 3B. 2.Convert temperature readings to Centigrade if necessary, and produce a graph of each of your measurements similar to Fig 2. Identify the set of data points (Set A) that will define the linear portion of the graph before ignition, and the set (Set B) that will define the linear portion after ignition. Also identify the crossover point (C) in time (t*) and temperature (T*). 3.Obtain the equations describing data sets A and B with least-squares regression. Substitute t* into the equation for set A to calculate T initial, and into the equation for set B to calculate T final, then calculate the temperature change T with its uncertainty. 4.Calculate the heat capacity equivalent from the temperature change ( T BA ), the weight of the sample (W BA ), and the weight of iron burned (W Fe ) in the calibration run (Benzoic Acid). C eq (J/K) = (26,425 W BA + q corr )/ T BA ; q corr (J) = 5858 W Fe 5.Calculate H m for your sample with Eq. (3) from above. 6.Either obtain the H m values for the other samples from the other groups, or download their data files from the server and repeat steps 2-5 to obtain the values. 5

6 Report: As results of this experiment are reported, the data will be filed in the Chem242 folder on the server. Discuss the outcome of the experiment in terms of the results obtained by your group, the accumulated results of this class, and literature values. 1.Calculate rxn H and its uncertainty for the combustion: C X H Y O Z (s) + (X + Y/4 - Z/2) O 2 (g) -> X CO 2 (g) + Y/2 H 2 O(l) ; N gas = Z/2 - Y/4 2. Compare your result to a literature value of the heat of combustion of your compound. 3. Look up the standard enthalpies of formation of O 2 (g), CO 2 (g) and H 2 O(l), and use the values to determine the standard enthalpy of formation (and the uncertainty) of the compound(s) studied. References: 1 Atkins, Peter Physical Chemistry, 6th Ed., Freeman, NY, 1998, P Halpern, A. M. Experimental Physical Chemistry, 2nd Ed., Prentice-Hall, Upper Saddle River, NJ 1997, p Table: Heats of Combustion Handbook of Chemistry and Physics, xx Ed., CRC Press, Boca Raton, FL. 6