# DETERMINING THE MOLAR MASS OF CARBON DIOXIDE

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1 DETERMINING THE MOLAR MASS OF CARBON DIOXIDE PURPOSE: The goal of the experiment is to determine the molar mass of carbon dioxide and compare the experimentally determined value to the theoretical value. INTRODUCTION: A re-arranged form of the ideal gas law can be used to calculate the molar mass of an ideal gas, Molar Mass = mrt PV where m is the mass of the gas (in grams), T is the absolute temperature (Kelvin), P is the pressure (in atmospheres), V is the volume of the gas (in liters), and R = L atm/mol K. In this experiment, carbon dioxide will be generated, using the reaction shown below, and collected at atmospheric pressure and room temperature. 2HCl(aq) + CaCO3(s) CaCl2(aq) + CO2(g) + H2O(l) By determining the exact mass and volume of the carbon dioxide that is generated, the molar mass can then be calculated. GAS COLLECTION EQUIPMENT: 125-mL or 250-mL Erlenmeyer flask for gas collection (must be absolutely dry do not wash) Rubber stopper with a hole in the center Glass tube that moves freely in the rubber stopper Cork that fits the hole in the rubber stopper Rubber tubing Drying tube CARBON DIOXIDE GENERATOR: 250-mL Erlenmeyer flask Thistle tube apparatus with right-angle connector ~30 ml 6 M HCl 30 g calcium carbonate (marble chips) PROCEDURE: For best results, do not handle the gas collection flask with bare hands and use the same analytical balance for all mass measurements.

2 A. Collection of Dry Air 1. Use rubber tubing to connect the glass tube to the drying tube, and to connect the drying tube to the compressed air valve. 2. Place the glass tube into the hole in the rubber stopper and place the rubber stopper in the gas collection flask. The glass tube should move loosely in the stopper to prevent a build-up of excess pressure and to allow the dry air to displace the moist air in the flask. 3. Slowly open the compressed air valve and allow the compressed dry air to fill the flask (~10 min). 4. Turn off the compressed air; remove the glass tube, place the cork in the stopper and mass the flask. 5. Reattach the compressed air and fill the flask for an additional 5 min; place the cork in the stopper and mass the flask. Repeat this step until two consecutive masses differ by less than g. 6. Record the laboratory temperature and atmospheric pressure using the digital barometer. Convert to appropriate units. 7. Calculate the density of dry air (in g/l) using the equation shown below. Density = M MP RT where MM is the molar mass of dry air (28.96 g/mol), P is atmospheric pressure (atm), T is the absolute temperature (Kelvin), and R = L atm/mol K. B. Collection of Dry Carbon Dioxide 1. Place the 30 g calcium carbonate in the 250-mL CO2-generator flask; add ~20 ml deionized water and attach the thistle tube apparatus. The end of the thistle tube should be under water. 2. Detach the rubber tubing from the compressed air and attach it to the right-angle tube in the carbon dioxide generator. Stabilize the setup with clamps. 3. Add ~10 ml HCl to the generator through the thistle tube to initiate the reaction. Check the setup for leaks. Note: The glass tube must be loose in the stopper of the collection flask in order to displace the dry air with the carbon dioxide! Everything else should be air-tight. 4. Collect carbon dioxide for ~30 minutes, periodically adding small portions of HCl to keep up the CO2 production. 5. Remove the glass tube; place the cork in the stopper and mass the flask. 6. Reattach the generator and collect carbon dioxide for ~10 minutes; place the cork in the stopper and mass the flask. Repeat this step until two consecutive masses differ by less than g. C. Collection of Deionized Water: to be completed only after gas collection is complete 1. Fill the gas collection flask to the brim with deionized water. 2. Replace the rubber stopper, followed by the cork. Let the excess water spill over the sides of the flask. 3. Dry the flask and mass it on a top-loading balance (not the analytical balance). 4. Calculate mass of water in the flask. 5. Record the density of water at the laboratory temperature.

3 Temperature ( C) Density (g/ml) Temperature ( C) Density (g/ml) Use the density and mass of water to calculate the volume of water in liters (which equals the volume of gas collected).

4 PRELABORATORY ASSIGNMENT NAME: DATE: 1. If 200. ml of an ideal gas exerts a pressure of 760 mmhg, what volume will the same gas occupy at 1450 mmhg, assuming constant temperature? 2. A sample of hydrogen has a volume of 245 ml at 21 C. What volume will the same gas occupy at 120 C, assuming constant pressure? 3. A sample of carbon dioxide at 35.0 C occupies a volume of ml at 20.0 psi. What volume will the gas occupy at STP? 4. A g sample of gas has a volume of 224 ml at 745 torr and 72.4 F. What is the molar mass of the gas? 5. What is the density of dry air at STP? 6. What volume of carbon dioxide is produced at STP when 30.0 g calcium carbonate is combined with 30.0 ml 6.0 M HCl?

5 RESULTS AND CALCULATIONS A. Collection of Dry Air Temperature (Kelvin): Atmospheric Pressure (atm): Mass of flask, stopper, cork, and dry air (g): Density of dry air (g/l): B. Collection of Carbon Dioxide Mass of flask, stopper, cork, and CO2 (g): C. Collection of Water Mass of flask, stopper, cork, and water (g): Mass of water (g): Density of water (g/ml): Volume of water (L): D. Molar Mass of Carbon Dioxide Mass of dry air (g): Mass of empty flask, stopper, and cork (g): (empty means no air) Mass of carbon dioxide (g): Experimental Molar Mass of Carbon Dioxide: Theoretical Molar Mass of Carbon Dioxide: (using the periodic table) Percent Error: Percent Error = experimental theoretical 100 theoretical

6 POST-LAB QUESTIONS 1. A 2.50-g sample of impure sodium acetate was heated in the presence of excess sodium hydroxide, producing methane gas and solid sodium carbonate. A volume of 415 ml of gas was collected over water at 27 C and 742 torr. a. What is the balanced chemical equation for the reaction? b. How many grams of methane were produced? The vapor pressure of water at 27 C is 27.8 torr. c. How many grams of sodium acetate are required to produce this quantity of methane? d. What is the percent pure sodium acetate in the impure sample? 2. Methane burns in the presence of O2 to produce carbon dioxide and water vapor. If 2.50 L of methane is burned in an open container at STP, what volume of O2 is required for complete combustion? What volumes of carbon dioxide and water vapor would be produced?

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