The ideal gas law, PV = nRT, gives an accurate description of the behavior of real gases at low pressures and at temperatures which are high relative to the boiling point. The ideal gas law is based on the assumption that the molecules experience no intermolecular forces and that the molecules occupy no volume. These assumptions are valid at low pressure and high temperature since under these conditions the molecular density is low. The molecules are too far apart to "feel" attractive forces exerted by other molecules. Furthermore, since the molecules are far apart, the volume occupied by the molecules is negligible compared to the total volume occupied by the gas.
In reality, intermolecular forces do exist and molecules do occupy space. The extent to which these factors cause a gas to deviate from the ideal gas law at a particular temperature and pressure depends upon its molecular structure.
The volume occupied by one mole of a gas is its molar volume,V*:
V* = R • T / P
For an ideal gas, this equation gives a value of 22.414 L/mol at Standard Temperature and Pressure (273.15 K and 1 atm).
The purpose of this experiment is to measure the molar volume of oxygen gas, and compare this measured value to the value predicted by the ideal gas law.
This can be accomplished by generating a known mass of oxygen gas, measuring its temperature, volume and pressure, and then using this data to calculate the molar volume at STP.
The oxygen is generated by the decomposition of potassium chlorate at high temperature according to the reaction:
2 KClO3(s) 2 KCl(s) + 3 O2(g)
The decomposition is carried out in the apparatus shown in Figure 1. A solid sample containing KClO3 and MnO2 is placed in the test tube and heated (MnO2 is a catalyst used to increase the rate of decomposition). The mass of oxygen produced is equal to the mass loss of the solid sample,
mass O2 = initial mass of KClO3 - final mass of KClO3
The oxygen gas produced in the reaction displaces water from the flask into the beaker. After the gas generation is complete, the pressure inside the flask is equalized to atmospheric pressure. At this point, the total pressure inside the flask is, according to Dalton's Law of Partial Pressures, equal to the sum of the partial pressures of oxygen and water vapor:
Patm = PO2 + PH2O
Where Patm is the atmospheric pressure obtained from a barometer, PO2 is the partial pressure of oxygen and PH2O is the vapor pressure of water, which can be obtained from the Handbook of Chemistry and Physics, (published by the Chemical Rubber Co).
The volume of the water displaced from the flask into the beaker is equal to the volume of oxygen produced. Assuming that Boyle's and Charles' laws apply, the volume of oxygen at STP is calculated as:
VSTP = VO2 • TSTP • PO2 / TO2 • PSTP
Where V is the measured volume of oxygen, T is its temperature, and TSTP, and PSTP are the temperature and pressure at STP. The molar volume at STP is then:
V* = VSTP / mols O2
Assemble the apparatus shown in Figure 1, except for the test tube. Fill a 500 mL Erlenmeyer flask nearly full with water. Use a rubber pipette bulb placed at the end of tube A to force water from the flask into the beaker, thus creating a siphon between the flask and the beaker (your instructor will show you the proper use of the bulb). Raise and lower the beaker to expel any air bubbles from tube B. Raise the beaker to fill the flask with approximately 400 mL of water. Note: water must not enter tube A! Close tube B with a pinch clamp. You now have a leak tight system from the Erlenmeyer flask to the 250 mL beaker.
Figure 1: The apparatus for the experiment.
CAUTION: KClO3 is a strong oxidizing agent and will react readily when heated with certain easily oxidizable substances such as grease and rubber. Clean and dry a test tube thoroughly before adding the KClO3 to it. Weigh the empty test tube, then add approximately 3 g of KClO3 to it and reweigh. All weighing should be to 0.001 g. To ensure no loss of KClO3 during transfer, weigh your test tube and add the KClO3 directly to it. Tap the test tube as you add KClO3 to ensure that it ends up in the bottom of the test tube. It is important that none of the KClO3 powder remains at the top of the test tube where it can contact the rubber stopper. After you have added the KClO3, place a piece of glass wool about half way down the test tube and reweigh. The glass wool will act as a filter to allow the oxygen to leave, but retain the KCl. The MnO2 catalyst has already been added to the KClO3 in a ratio of approximately 1:30.
Place the one hole stopper on tube A tightly into the test tube. Check the system for air leaks by removing the pinch clamp on tube B. If there are no leaks, water will not flow out into the beaker. If there is a leak, try wetting the two rubber stoppers and reassembling. If the leak persists, consult with your instructor.
Once your system is leak tight, replace the pinch clamp on tube B. Remove the tube from the beaker and dry it carefully. Weigh the empty 250 mL beaker, and then replace tube B into it.
Begin heating the KClO3 with a Bunsen burner and wait approximately 5 seconds before removing the clamp from tube B. Heat gently at first to obtain a moderate rate of oxygen evolution. If white vapors appear in the system, stop heating until they disappear. The end of tube B must remain immersed under water during the entire time after heating is begun. Air must not be allowed to enter the system. Continue heating until you have collected between 175 and 200 mL of water. NOTE: Do not exceed 200 mL or the combined weight of the 250 mL beaker and water will exceed the capacity of the balance.
When the heating is completed allow the apparatus to cool with tube B open, but with its end still immersed in water. Once the apparatus is at room temperature close tube B with a pinch clamp and remove it from the beaker. Carefully determine the weight of water collected. Use a thermometer to measure the temperature of the water in the beaker to 0.1°C. Based on this temperature, use these links to determine the vapor pressure of water in your oxygen sample and the density of water.
Weigh the test tube and its contents.
Obtain the atmospheric
the barometer (How to read an Eco-Celli barometer).
Waste Disposal: Remove the glass wool and permanganate plug from of your test tube (you may find that a little water and GENTLY tapping will help). The glass wool can be thrown in the trash. Place the permanganate plug into the marked waste container in the hood.
O.K., let's work with some sample data to see how all of this comes together to give us an experimentally determined value for the molar volume of a gas. REMEMBER, THIS IS AN EXAMPLE; YOUR NUMBERS WILL BE DIFFERENT! Assume that you started with a 4.000 g sample of pure KClO3 in your test tube. After heating the KClO3, you find that you have collected 180.000 g of water in the beaker. In addition, the temperature of the water in the flask is 22.3°C and the barometric pressure of the room is 761.3 Torr. From a Table of Water Density, the density of water at 22.3°C is 0.997701 g/mL. Therefore, the volume of oxygen generated is:
VO2 = 180.000 g / 0.997701 g/mL = 180.415 mL
After weighing the test tube, you find that the 4.000 g of pure KClO3 has been reduced to 3.770 g of mixed KClO3 and KCl. This 0.230 g difference represents the amount of oxygen gas generated during the experiment. The mols of oxygen are then calculated as:
mols O2 = 0.230 g • 1 mol/31.99 g = 0.00719 mols
Now you need to find out the volume that this gas would occupy at STP. Since you equalized the pressure inside the flask to the pressure in the room (atmospheric pressure), you can use the barometer to determine the pressure inside the flask. Unfortunately, the gas inside the flask is not just oxygen, it also contains water vapor. So the pressure you determined using the barometer is actually the sum of the partial pressures of the oxygen and the water vapor. To correct for the partial pressure due to water vapor, you need to use a Vapor Pressure Table for Water. Your experiment was run at 22.3°C which corresponds to 20.193 Torr of water vapor (you may need to interpolate if your exact temperature is not in the table). This means that the pressure of your gas that is due solely to the oxygen is 741.1 Torr (761.3 - 20.193). However, this was not collected at STP, so you must use the combined gas law to convert:
PSTP • VSTP / TSTP = PO2 • VO2 / TO2
VSTP = VO2 • TSTP • PO2 / TO2 • PSTP
VSTP = 0.180415 liters • (273.15 K • 741.1 Torr) / (295.45 K • 760 Torr)
VSTP = 0.1626 liters
Therefore, your calculated value for the molar volume of oxygen is:
V* = VSTP / mols O2
V* = 0.1626 liters / 0.00719 mols
V* = 22.6 liters/mol
The percent deviation from the accepted
value is calculated by subtracting the accepted value from your value
and then dividing by the accepted value:
(22.6 - 22.414) • 100 / 22.414 = 0.83%
For each trial calculate the following from your data. Show all calculations.
(Updated 7/13/13 by C.R. Snelling)