The Molar Volume of a Gas
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 will depend its molecular structure.
The volume occupied by one mole of a gas is its molar volume, . For an ideal gas
which gives a value of 22.414 L at STP ("Standard temperature and pressure" of 273.15 K and 1 atm).
One purpose of this experiment is to measure the molar volume of oxygen gas, and to compare the measured value to the value predicted by the ideal gas law. A second purpose is to determine the percentage KClO3 in an unknown sample.
The first purpose is accomplished by generating a known mass of oxygen gas, measuring its temperature, volume and pressure, and then using the 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:
|2KClO3(s) --> 2KCl(s) + 3O2(g)||(1)|
The decomposition is carried out in the apparatus shown in figure 1. A preweighed solid sample containing KClO3 mixed with KCl and MnO2 is placed in the test tube and heated ( The MnO2 is a catalyst for the decomposition of KClO3) . The mass of oxygen produced is equal to the mass loss of the solid sample:
mass O2 = initial mass of solid sample - final mass of solid sample
Figure 1: The apparatus for the experiment.
The oxygen gas produced in the reaction displaces water from flask B into beaker C. Once the pressure inside B is equalized with the atmospheric pressure, 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 flask B into beaker D is equal to the volume of oxygen produced. Assuming that Boyle's and Charles' laws apply, the volume of oxygen at STP is calculated from:
where V is the measured volume of oxygen, T is its temperature, and TSTP, and PSTP are the STP temperature and pressure.
The molar volume at STP is then:
The second purpose of the experiment is to determine the percentage KClO3 in an unknown sample.The percent KClO3 in the unknown is obtained from the amount of oxygen produced and the stoichiometry of reaction 1.
Stockroom: Things for each group to borrow and return on the same day.
Assemble the apparatus shown in Figure 1, except for the test tube. Fill flask B (to be referred to hereafter as "the flask") nearly full with water and add about 200 mL of water to beaker D (to be referred to hereafter as "the beaker") . Use a rubber pipet 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 . Raise and lower the beaker to expel any air bubbles from tube C. Raise the beaker to fill the flask to within about 2 cm of the short glass tube connected to tube A. Water must not enter tube A. Close tube C with a pinch clamp.
Obtain an unknown from the stockroom, and immediately enter its identification number into your notebook. 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 test tube E (to be referred to hereafter as "the test tube") thoroughly before adding the unknown sample to it. Weigh the empty test tube , and add to it 2 to 2.5 g of the unknown. All weighings should be to .01 g. Use a paper funnel to transfer the solid to the test tube to prevent it from coming into contact with the rubber stopper.
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 C. If there are no leaks, water will not flow out of the flask even if the water levels in the flask and the beaker are different. If there is a leak consult with your instructor.
Equalize the pressure in the flask with atmospheric pressure as follows: with the end of tube C immersed in water in the beaker , raise the beaker until the water levels in the beaker and the flask are equal. Air must not enter tube C. Close tube C with a pinch clamp and pour out the water in the beaker .
[Bunsen burner lighting.]
Remove the pinch clamp from tube C and begin heating the unknown mixture with a Bunsen burner. 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 C 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 no further oxygen is evolved. That is , until the water level in the beaker stops changing.
When the heating is completed allow the apparatus to cool with tube C open, but with its end still immersed in water. Once the apparatus is at room temperature equalize the pressure inside the flask with the atmospheric pressure, and close tube C with a pinch clamp.
Carefully measure the volume of water in the beaker with a 500 mL graduated cylinder. Measure the temperature of the oxygen gas by removing the stopper in the flask and inserting a thermometer into the gas.
Weigh the test tube and its contents. Obtain the atmospheric pressure from the barometer. Barometer instructions. The barometer reading must be corrected for the unequal expansion of the mercury and the brass scale. A table of correction factors will be supplied by the instructor.
Waste Disposal: Discard the used and excess unknown into the trash or wash it down the drain with lots of water.
Use the following table to guide your data collection:
Identification number of unknown:______
|Trial 1||Trial 2|
|Weight of test tube and contents before heating|
|Weight of test tube and contents after heating|
|Weight of empty test tube|
|Volume of oxygen collected|
|Barometer reading (uncorrected at___oC)|
|Temperature of oxygen|
|Vapor pressure of water at the temperature of the oxygen|
For each trial calculate the following from your data. Show all calculations.
Mass of oxygen
moles of oxygen
Absolute temperature of oxygen
Corrected barometric pressure
Partial pressure of oxygen in flask B
Volume of oxygen at STP
Molar volume of oxygen at STP
Moles of KClO3 decomposed
Mass of KClO3 in the unknown sample
Percent KClO3 in the unknown sample
Go on to calculate from the results of your two trials:
The average percentage of KClO3 in your unknown
The average molar volume of oxygen at STP
Report the average molar volume of oxygen at STP based on your experimental results. What is the percent deviation from the ideal gas value?
Report the average percent KClO3 in your unknown sample?
Report the average deviation of your two measurements of the percent KClO3 .