Gas Laws: Determining the Molar Mass of Volatile Liquid

CHM 2045L Gas Laws: Determining the Molar Mass of Volatile Liquid OBJECTIVE: To use the ideal gas law to determine the molar mass of a volatile liquid. INTRODUCTION:

Charles, Boyle and Avogadro separately investigated the relationship between pressure, volume and temperature of an ideal gas. Combining their work and solving for the constant to tie their contributions together resulted in the ideal gas equation:

PV = nRT

where P = pressure (atm); V = volume (L); T = temperature (K); n = number of moles and R = ideal gas constant or 0.0821 L atm/K mole.

The number of moles (n) can be expressed in terms of the mass of the gas sample (m) and its molar mass (Mm)

n = m . MM

This expression can be substituted into the ideal gas equation to give

PV = m RT MM

Solving for the molar mass

MM = mRT PV

There is a useful means of determining the molar mass of a substance from measurements of mass, temperature, pressure and volume of a gas sample. This directly yields the molecular weight (MW), since the MW in amu is numerically equal to the molar mass (MM) expressed in grams/mole.

Many substances show ideal gas behavior at low pressures and at temperatures above the boiling point. In this experiment, a sample of a volatile liquid is vaporized in a flask of known volume. The air originally present in the flask is swept out by the gas produced from the liquid sample during the vaporization process. The vaporized sample is assumed to act as an ideal gas and the number of moles of vaporized sample in the flask is calculated from the ideal gas law using its pressure and its temperature, the volume of the flask and the ideal gas constant. The gas pressure is equal to the atmospheric (barometric) pressure because the flask is not sealed. Upon cooling the vaporized sample remaining in the flask condenses and air re-enters the flask. The mass of the volatile liquid that filled the flask in the vapor form is the difference between the mass of the flask containing the condensed sample and the mass of the empty flask.

Safety Precautions: Wear safety goggles at all times while in the laboratory.

PROCEDURE:  Read through the procedure as if you were doing the lab so you can answer the questions

1. Prepare a 125-mL Erlenmeyer flask by cleaning the flask and then drying it completely. The flask must be completely dry, since any water present will vaporize under the conditions of the experiment and will adversely affect the results.

2. Cut a square of thick (freezer) aluminum foil to serve as a cover for the flask. Trim the edges of the foil so that it neatly covers the mouth of the flask but does not extend far down the neck.

3. Prepare a beaker for use as a heating bath for the flask. The beaker must be large enough for most of the flask to be covered by boiling water when in the beaker. Add the required quantity of water to the beaker. Set up the beaker on a ring stand over a Bunsen burner, but do not begin to heat the water bath yet. Add a 2

nd ring around the beaker to stabilize it.

4. Weigh the dry, empty flask with its foil cover and record the value.

5. Add 3-4 mL of a volatile liquid to the dry Erlenmeyer flask. Cover the flask with the foil

cover, making sure that the foil cover is tightly crimped around the rim of the flask. Punch a single small hole in the foil cover with a needle or pin.

6. Heat the water in the beaker to boiling. When the water in the beaker begins to boil, adjust the flame of the burner so that the water remains boiling but does not splash from the beaker.

7. Immerse the flask containing the liquid in the boiling water so that most of the flask is covered

with the water of the heating bath. Clamp the neck of the flask to maintain the flask in the boiling water.

8. Watch the liquid carefully. The liquid will begin to evaporate rapidly, and its volume will decrease. The amount of liquid placed in the flask is much more than will be necessary to fill the flask with vapor at the boiling water temperature. Excess vapor will be observed escaping through the pinhole made in the foil cover of the flask.

9. When it appears that all the liquid has vaporized, and the flask is filled with vapor, continue to heat for 1-2 more minutes. Then remove the flask from the boiling water bath; use the clamp on the neck of the flask to protect your hands from the heat.

10. Set the flask on the lab bench, remove the clamp, and allow the flask to cool to room temperature. Liquid will reappear in the flask as the vapor in the flask cools. While the flask is cooling, measure and record the exact temperature of the boiling water in the beaker.

11. When the flask has cooled completely to room temperature, carefully dry the outside of the flask to remove any droplets of water. Then weigh the flask, foil cover, and condensed vapor to the nearest 0.0001 g.

12. Repeat the experiment with another 3-4 mL sample of the volatile liquid. 13. Fill the flask to the very rim with tap water, cover with the foil cover, and weigh the flask,

cover, and water to the nearest 0.0001 g. Determine the temperature of the tap water in the flask. (Assuming the density of water at 20oC is 0.99819g/mL, calculate the mass of the water in the flask since it is related to the volume of the water in the flask which represents the volume of the vapor.) Do not use the volume of the Erlenmeyer flask marked on the side - It’s not the max that the flask can hold.

Name: ______________________

Data Table 1. Data for trials 1 and 2 of Determining the Molar Mass of Volatile Liquid

Trial 1 Trial 2

Mass of the dry, empty flask with its foil cover in grams 202.8418g 204.0197g Temperature of the water bath in

o C: 18.0oC 18.0 oC

Temperature of the water bath in Kelvin: Volume of the Erlenmeyer Flask: in mL 269.1mL 268.4mL Volume of the Erlenmeyer Flask: in L Mass of the cooled condensed sample in the flask with its foil cover in grams

203.4792g 204.5935g

Mass of sample in grams

CALCULATE the following using the data provided in Data Table 1:

1. The molar mass of the volatile liquid from the 2 trials, and the average. Assume that the atmospheric pressure is 1.00atm because we are so close to sea level.

CONCLUSION DISCUSSION:

In a paragraph in your conclusion, comment on how well the objectives of the experiment were met. Assume that the accepted molar mass of the volatile liquid is 58.08g/mol. Discuss the accuracy of your results and suggest potential sources of experimental error.