The Molecular Weight of Carbon Dioxide

Transcription

1 The Molecular Weight of Carbon Dioxide Objectives The objectives of this laboratory are as follows: To generate and collect a sample of carbon dioxide gas, then measure its pressure, volume, temperature and mass. To appy these measurements to the Ideal Gas Law in order to determine the molecular weight of this gas. To compare this experimentally determined molecular weight to the theoretical molecular weight, and to calculate the percent error between these values. Background The state of a gas is completely described by its pressure (P), volume (V), temperature (T) and mole quantity (n). The mathematical relationship between these properties is given by the Ideal Gas Equation: PV = nrt (1) where P has units of atmospheres, V is in Liters and T is in Kelvins. R, the Gas Constant, has a value of L atm/k mol. Recall now that the number of moles of gas can be obtained by dividing the mass of the gas (in grams) by its molecular weight (or molar mass, in g/mol): n = m (2) M W Substituting this into the Ideal Gas Equation gives m PV = RT (3) M W which may be re-arranged to yield the following expression for the molecular weight of a gas: mrt M W = (4) PV Thus, the molecular weight of a gas can be determined simply by measuring the temperature, pressure, volume and the mass of a sample of that gas. In this experiment, the molecular weight of carbon dioxide gas will be determined. The carbon dioxide gas is produced by the reaction of calcium carbonate with hydrochloric acid: CaCO 3 (s) + 2 HCl (aq) CO 2 (g) + CaCl 2 (aq) + H 2 O (l) The carbon dioxide gas will be collected in an Erlenmeyer flask that is covered with aluminum foil. Since carbon dioxide is denser than air, the simple foil cover is effective enough to inhibit the rate at which the carbon dioxide mixes with the outside air. Once collected the temperature, pressure, volume and the mass of the CO 2 will be measured. The Molecular Weight of Carbon Dioxide Page 1 of 5

2 The temperature of the gas will be measured by briefly placing a thermometer underneath the foil. Since the flask is not sealed, the pressure of the gas in the flask will be equal to the pressure in the laboratory. The laboratory pressure will be measured with a mercury barometer. The volume of the gas is equal to the volume of the flask. The volume of the flask will be measured by filling it to the brim with deionized water. From the measured mass of the water and the known density of water (see Table 1), the volume of the water, which equals the volume of the flask, will be determined. The mass of gas in the flask can be determined by weighing the gas filled flask and then subtracting the mass of the empty flask. In most experiments the mass of an empty container is actually the mass of the container plus the air it holds. This is generally not a problem since the mass of the air is relatively small compared to the mass of a liquid or solid sample and can therefore be neglected. However when weighing a gas, not accounting for the mass of the air in the container would be a serious systematic error. The mass of air in the flask will be estimated from the volume of the flask and the known density of dry air (see Table 2). It is reasonable to approximate the air as being dry since the mole fraction of water vapor in air is very small (even in air that is saturated with water vapor the mole fraction of water vapor is only 0.03). The mass of the empty flask will then be calculated by subtracting the calculated mass of dry air from the measured mass of the flask filled with air. After the flask has been filled with carbon dioxide, it will be weighed, and the mass of the gas will be determined by subtracting the calculated mass of the truly empty flask. The molecular weight of carbon dioxide can then be obtained by substituting these measurements into Equation (4). Once this experimental molecular weight (EV) is determined it will be compared to the theoretical molecular weight (TV) of carbon dioxide, and the percent error between these values will be calculated using the formula: EV TV Percent Error = 100 TV Table 1: Density of Liquid Water Temperature, C Density, g/ml (5) The Molecular Weight of Carbon Dioxide Page 2 of 5

3 Table 2: Density of Dry Air Temperature, C Density, g/l P = 750 torr P = 760 torr P = 770 torr *Tabulated Data is obtained from the CRC Handbook of Chemistry and Physics, 64th ed., Procedure Safety 6 M hydrochloric acid is capable of causing serious chemical burns and blindness and should be handled with care. If the acid comes in contact with your skin immediately rinse the affected area with water for several minutes. Materials and Equipment Chemicals: Solid CaCO 3 and 6 M HCl Equipment: Two 250 ml Erlenmeyer flasks, 50 ml beaker, thermometer, utility clamp, stand, analytical balance, triple beam balance, barometer, aluminum foil square (15 x 15 cm). From the Stockroom: A CaCl 2 drying tube, a two-holed rubber stopper with a thistle tube inserted into one hole (make sure the stopper snugly fits Erlenmeyer flask A) and a piece of bent glass tubing (connected to a piece of rubber tubing) inserted into the other hole, and a straight piece of glass tubing connected to a piece of rubber tubing. Instructions 1. Obtain one of the 250 ml Erlenmeyer flasks. This will be designated as Flask B. Make sure that the flask is both clean and dry. Record the mass of the flask and the foil to the nearest g using the analytical balance. Note that it is mass (flask + foil + air) that is being measured here. Also record the temperature of the air in the flask. 2. Place about 25 grams of calcium carbonate into the other 250 ml Erlenmeyer flask, designated as Flask A. Add about 10 ml of water, or enough to cover the chips completely. Insert the two-holed rubber stopper (with the thistle tube and the bent tube) into Flask A, making sure that the thistle tube is adjusted so that it is beneath the water but not touching the bottom of flask. The stopper must fit tightly into the flask. Also obtain about 25 ml of dilute (6 M) hydrochloric acid, HCl, in a small, labeled beaker. The Molecular Weight of Carbon Dioxide Page 3 of 5

4 3. Assemble the apparatus shown in the figure below. Insert the straight glass tube (the one that is attached to the flexible tubing but not attached to the one-holed rubber stopper) into Flask B by placing it between the foil and the flask and pressing the foil against it to hold it in place. Be careful to fold and shape the foil only as much as necessary since it is fragile and will easily tear. Attach the flexible end of the tubing to the drying tube. The small rubber stopper attached to the bent glass tubing should be inserted into the other end of the drying tube. 4. When all is ready, pour 5 10 ml of the hydrochloric acid into the top of the thistle tube and allow it to run through the tube into Flask A. The reaction should begin immediately as evidenced by gas evolution. Allow the reaction to continue for at least 20 minutes to displace all of the air in Flask B with carbon dioxide gas. During this time pay attention to what is happening in Flask A; if the gas evolution ceases, add additional HCl through the thistle tube. After 20 minutes remove the tube from Flask B (keep the foil in place) and immediately weigh the flask containing carbon dioxide on the analytical balance to the nearest g. 5. Reassemble the apparatus and allow gas evolution to continue and flow into Flask B for an additional 15 minutes. Again, weigh Flask B (with the foil). The two masses (before and after the 15 minutes) should agree closely (to within g). If the mass has increased by more than g, reassemble the apparatus and continue to collect carbon dioxide gas for an additional 5 minutes, then reweigh the flask. Continue this procedure until successive masses agree to within g (or until a decrease in mass is observed). Next, measure and record the temperature of the carbon dioxide in the flask. Use the lab barometer to read the atmospheric pressure. The Molecular Weight of Carbon Dioxide Page 4 of 5

5 6. Finally, in order to determine the volume of Flask B, fill this flask with deionized water to the brim, wipe any water from the outside of the flask and weigh it, with the foil, to the nearest 0.01 g using the triple beam balance. Next, measure and record the temperature of the water. The volume of water in the flask can now be calculated from the mass and density of water in the flask at the temperature recorded. Cleanup Neutralize the contents of Flask A and any unused HCl by adding solid sodium bicarbonate until the solution no longer fizzes upon addition of additional bicarbonate. This step is best performed in the sink as the flask tends to overflow. Decant off the liquid, leaving any solid marble chips behind. Rinse the chips with water and decant off the water. Pour the chips into the labeled recovery beaker. The Molecular Weight of Carbon Dioxide Page 5 of 5