Lab #3
Determination of the Molar Mass of Volatile Liquid
Background-Overview: The ideal gas law relates the four measureable properties of a gas: pressure, volume,
number of moles and temperature. In this experiment, the ideal gas law will be used to determine the molar
mass of gases a volatile liquid. A small sample of the volatile liquid is placed in a tared 250 ml Erlenmeyer flask.
The volatile liquid and flask are heated in boiling water to volatilize the liquid. The air and excess vapor escape,
leaving the flask filled only with the volatile liquid vapor at atmospheric pressure and at the temperature of
boiling water. The flask is then removed and cooled to condense the vapor. Once cooled, the flask is weighed.
By massing the same flask filled with water, the volume of the flask is calculated. The molar mass of the
volatile liquid is then determined using the mass of the condensed vapor, the volume of the flask, the
atmospheric pressure and the temperature of the boiling water.
Ideal Gas Law:
PV = nRT
P= Barometric pressure (in atm)
V=Volume (in L)
n= number of moles
R= 0.821 L atm/mol K
T= temperature (in Kelvins)
Molar Mass:
Molar mass = mass of condensed vapor/moles of condensed vapor
Purpose: The purpose of this experiment is to determine the molar mass of a volatile liquid(ethanol) based on
measurements of its vapor density. The ethanol is volatilized and condensed in a fixed volume. The condensed
vapor is massed and ethanol’s molar mass is calculated from the experimental data.
Safety:
-Ethanol is a flammable liquid and a fire risk. Ethanol is made poisonous by the addition of denaturant-it cannot
be made nonpoisonous.
-Wear chemical splash goggles.
-Exercise care when working with the hot water bath
-Wash hands thoroughly with soap and water before leaving the laboratory
Materials:
-Chemicals
Ethanol
-Equipment
250 ml Erlenmeyer flask
1000 ml beaker
Hot plate
Aluminum freezer foil
Needle or pin
Pre-Lab:
Carefully read the procedure. A determination of the molar mass of methanol yielded the following data:
Temperature of boiling water bath
99.5 °C
Barometric Pressure
738 mm Hg
Temperature of Room Temperature water bath
24.0 °C
Mass of Empty Flask
1.557 g
Mass of Flask and condensed methanol
1.571 g
Mass of Flask and Water
16.001 g
1. Determine the mass of the condensed methanol.
2. Determine the mass of water in the filled flask.
3. Use the relationship of mass and density to determine the volume of the flask.
4. Using the Ideal Gas Law, calculate the number of moles of methanol.
5. Using the number of moles of methanol and the mass of the condensed methanol, determine the molar mass
of the methanol.
6. Why is the vapor unlikely to behave as an ideal gas near the temperature at which the vapor would liquefy?
7. Explain the use of the mass spectrophotometer to determine molar mass.
Procedure:
1. Obtain a dry 250 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. An oven may be available for heating the flask to dryness, or your instructor may describe
another technique.
2. Cut a square of thick 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. Weigh the dry, empty flask with its foil cover and record in the data table.
4. Prepare a 1000 ml beaker for use as a heating bath for the flask. Add the required quantity of water to the
beaker so that most of the flask is covered by boiling water when in the beaker. Begin to heat the water to
boiling. When the water in the beaker begins to boil, adjust the temperature of the hotplate so that the water
remains boiling but does not splash from the beaker.
5. Add 3-4 ml of ethanol to the dry Erlenmeyer flask. Cover the flask with the aluminum 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. Immerse the flask containing the ethanol 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.
7. The ethanol 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.
8. When it appears that all of the ethanol has vaporized and the flask is filled with vapor, continue to heat for
1-2 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.
9. 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, as well as the barometric pressure in the laboratory.
10. 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.
11. Remove the foil cover and clean out the flask.
12. Determine the volume of the flask: Fill the flask to the very rim with tap water, cover with the foil cover
and weigh. Determine the temperature of the water in the flask. Record both the mass and temperature in the
data table.
13. Determine the barometric pressure and record in the data table.
13. Clean your lab station.
Data and Analysis:
1. Create a data table with the following information: Mass of flask and cover, Mass of flask, cover and
condensed liquid, Mass of flask, cover and water, temperature of the water in the flask, temperature of the
boiling water, barometric pressure.
2. Determine the mass of the condensed ethanol.
3. Determine the mass of the water in the flask.
4. Record the density of water at the temperature found in the flask.
5. Using the mass of the water in the flask and the density of the water, determine the volume of the flask.
Conclusion:
1. Using the ideal gas law, determine the number of moles of methanol present in the flask.
2. Using the moles of ethanol and the mass of the condensed ethanol, calculate the molar mass of ethanol.
3. Calculate your percent error.
Lab #3
Determination of the Molar Mass of Volatile Liquid
Table of Contents
Dates
Background-Overview:
Safety:
Materials:
Pre-Lab:
Carefully read the procedure. A determination of the molar mass of methanol yielded the following data:
Temperature of boiling water bath
99.5 °C
Barometric Pressure
738 mm Hg
Temperature of Room Temperature water bath
24.0 °C
Mass of Empty Flask
1.557 g
Mass of Flask and condensed methanol
1.571 g
Mass of Flask and Water
16.001 g
1. Determine the mass of the condensed methanol.
2. Determine the mass of water in the filled flask.
3. Use the relationship of mass and density to determine the volume of the flask.
4. Using the Ideal Gas Law, calculate the number of moles of methanol.
5. Using the number of moles of methanol and the mass of the condensed methanol, determine the molar mass
of the methanol.
6. Why is the vapor unlikely to behave as an ideal gas near the temperature at which the vapor would liquefy?
7. Explain the use of the mass spectrophotometer to determine molar mass.
Procedure:
Data and Analysis:
1. Create a data table with the following information: Mass of flask and cover, Mass of flask, cover and
condensed liquid, Mass of flask, cover and water, temperature of the water in the flask, temperature of the
boiling water, barometric pressure.
2. Determine the mass of the condensed ethanol.
3. Determine the mass of the water in the flask.
4. Record the density of water at the temperature found in the flask.
5. Using the mass of the water in the flask and the density of the water, determine the volume of the flask.
Conclusion:
1. Using the ideal gas law, determine the number of moles of methanol present in the flask.
2. Using the moles of ethanol and the mass of the condensed ethanol, calculate the molar mass of ethanol.
3. Calculate your percent error.