Molar Mass of Gas Lab

MOLAR MASS OF A GAS

According to Avogadro’s hypothesis, one mole of any gas at standard conditions (STP) of pressure (760 mm Hg) and temperature (0oC or 273 K) should occupy 22.4 liters of space. This allows you to calculate the molar mass of any gas if you measure the following characteristics for a sample of that gas: mass, volume, pressure, and temperature.

In this activity, you will use a simple butane lighter as your source of butane gas, which has the chemical formula C4H10. You will measure the above four parameters of a sample of this gas, convert its volume to STP, and calculate its molar mass.

Objectives

When you have completed this activity, you should be able to:

1.State Avogadro’s hypothesis and relate it to the molar mass of the gas.

2.Use the combined gas law to compute the volume a sample of the gas will occupy if the conditions are changed to STP.

3.Explain why the vapor pressure of water must be subtracted from the room pressure in order to determine the true pressure of any gas sample collected over water.

4.Apply the ratio and proportion method to Avogadro’s hypothesis in order to compute the molecular mass of any gas.

5.Compute the percent error of your findings.

Materials

butane lighter
sink for water
thermometer
100 mL graduated cylinder
balance for weighing the lighter
ruler

Procedure

1.Goggles should be worn when performing this investigation.

2.Fill the sink about ¾ full of water.

3.Lay a 100 mL graduated cylinder in the tray; allow it to fill completely with water, and then hold it inverted and upright without allowing any air bubbles into it.

4.Obtain a butane lighter. DO NOT LIGHT IT. Dry it as thoroughly as possible by using a paper towel and by blowing forcefully around the flint wheel and the release lever. Measure the mass of the lighter to the nearest 0.01 g. Record this mass in the Data Table.

5.Carefully hold the butane lighter under the water so that the released gas will go into the inverted graduated cylinder at the maximum rate until it contains about 50 mL of gas. When you read the volume of gas in the cylinder, make sure that the level of water inside the cylinder is the same as the level of water outside. This step ensures that the pressure inside the cylinder is the same as the atmospheric pressure outside the cylinder.

6.Remove the lighter and dry it thoroughly as you did before. Again, determine its mass to the nearest 0.01 g and record this mass in the Data Table.

7.Measure and record the temperature of the water in the sink. Use the Internet to find the atmospheric pressure in the room.

8.Dispose of the butane in the cylinder by releasing it into the well ventilated room.

DATA TABLE

Initial mass of butane lighter / g
Final mass of butane lighter / g
Volume of butane released / mL
Temperature of water / oC
Atmospheric pressure / mm Hg

Analysis

1.Since the gas in your graduated cylinder is a mixture of butane and water vapor, you must determine the partial pressure of the butane, Pbutane, alone. To do this, consult a reference and record the partial pressure of the water vapor, Pwater, at the temperature you recorded. Use the following formula to compute the partial pressure of the butane.

Pbutane = Atmosphere - Pwater

2.Use the following combined gas law formula and compute the volume that the butane sample will occupy at STP. (HINT: Convert both temperatures to Kelvin.)

Pbutane x Voriginal= Pstandard x Vfinal

TroomTstandard

3.Use the following ratio and proportion formula to determine the mass of butane needed to occupy a volume of 22.4 L at STP.

Grams of butane you used“X” grams of butane

mL of butane corrected to STP=22,400 mL

4.Compute the theoretical molar mass of butane based on its formula and the atomic masses on the periodic table.

5.Compare your experimental results from #3 to the theoretical value of #4, computing a percent error of your findings using this formula:

% error = measured value - accepted value x 100

accepted value