Molar Mass Of A Volatile Liquid Chm 111 Formal Lab Report

1039 words - 5 pages

Molar Mass of a Volatile Liquid
Introduction: Given the mass of vapor at conditions of known pressure volume, and temperature, one should be able to determine the molar mass of the vapor. Mathematically, The Ideal Gas Law relates the quantities of pressure (P in atm), volume (V in liters), temperature (T in Kelvin), and quantity of gas (n in moles). Using the expression PV = nRT, where R is the ideal gas constant (0.0821 L atm/mole K), one can determine the quantity of gas under given conditions of pressure, volume, and temperature.
In this experiment, an unknown volatile liquid is heated in a boiling water bath and is vaporized. The vapor forces air from the flask through a tiny pin-hole in the foil cover, until the pressure within the flask equals the barometric pressure of the lab. As with all gases the vapor occupies the entire volume of its container. The temperature of the vapor equals the temperature of the boiling water bath. After 8–10 minutes, the vapor is cooled. The mass of the condensed liquid is determined and the molar mass is calculated from the moles calculated using the ideal gas law.
Materials:
· Unknown liquid (#9)
· 3 inch square of aluminum foil
· Two 800 or 1000 mL beakers
· 3 or 4 boilings chips
· 250 mL Erlenmeyer flask
· Burette clamp
· Thermometer
· Hot plate
Experimental Procedure:
1. Weigh the aluminum foil square and a 250 mL Erlenmeyer flask on an analytical balance.
2. Pour about 1.5 mL of your unknown liquid into the flask. Carefully form the foil into a cap and arrange the apparatus as shown in the picture to the right. Be sure the thermometer does not touch the beaker.
3. Add 3 or 4 boiling chips to the water in the 800 mL or 1000 mL beaker and heat the water to the boiling point. Continue heating for 8 to 10 minutes after the water comes to a rolling boil. Measure the temperature of the boiling water bath. Convert the temperature to units of Kelvin. This is the temperature of the vapor.
4. Turn off the hot plate and wait until the water has stopped boiling (about half a minute) and then loosen up the clamp holding the flask in place. Remove the flask from the beaker of hot water. Immerse it in a 800 or 1000 mL beaker of cool water to a depth of about 2 inches. When the flask is at room temperature (is cool to the touch) remove it from the water and dry it with a paper towel to remove surface water.
5. Re-weigh the flask, aluminum foil cap and liquid. Your instructor will give you the atmospheric pressure reading from the barometer. Convert the atmospheric pressure to units of atm.
6. Repeat the procedure using another 1.5 mL of your liquid sample.
7. Determine the volume of the flask by carefully filling it with water and weighing it (with the aluminum cap) on the analytical balance. Do not determine the volume before performing the experiment, since the water vapor that might remain in the flask after performing...

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