Experiment #9

Measurement of the Gas Constant and the Molar Volume of Oxygen

 

 

 

Purpose

This experiment will allow you to gain practical experience in the collecting and measuring the properties of gases.  From the data collected you will experimentally determine the values of the universal gas constant, R, and the molar volume of a gas, Vm, under standard conditions based upon your experimentally derived value of the gas constant.

 

 

Introduction

 

In this experiment, two values will be determined, the gas constant, R, and molar volume, Vm.  The calculations are based on the ideal gas law:

where P = pressure of the gas

V = volume of the gas

n = moles of the gas

T = temperature of the gas with units of Kelvin (K)

R = the gas constant

 

Moles are difficult to measure directly, however the number of moles of a substance is equal to the mass of the substance divided by its molecular mass so the equation above may be rearranged as below to replace moles by mass.

 

 

These equations can now be rearranged to solve for R and the molar volume Vm as shown below:

              

 

In this experiment you will measure the mass, volume, temperature, and pressure of a gas to experimentally determine values for R and Vm.  The literature values for these are listed below:

 

 

 

 

In this experiment, oxygen gas will be produced by the thermal decomposition of potassium chlorate in the presence of the catalyst manganese dioxide.  Catalysts are substances which increase the rate of a reaction but are not used up by the reaction.  Manganese dioxide, functions by lowering the temperature required to decompose the potassium chlorate.  The reaction for the decomposition of potassium chlorate is shown below:

 

 

 

 

 

Procedure

 

Safety Notes

¯  The KClO3/MnO2 mixture contains a strong oxidizing agent that may explode on excessive heating if the mixture becomes contaminated.  Be very careful not to get any glycerin in the mixture.

¯  Wear safety glasses at all times, especially when anyone in the laboratory is heating the KClO3/MnO2 mixture!

¯  All KClO3/MnO2 waste should be disposed of by dissolving the KClO3 in water and pouring it down the drain.  If solid KClO3 is put in the waste basket it may start a fire.

 

 

 

Text Box: Remember to use a milligram balance and record all masses to at least the nearest 0.001 g.Clean and dry a large (8-inch size) test tube.  Weigh the test tube.  Transfer 2.1-2.3g of the KClO3/MnO2 mixture into the test tube and reweigh it. 

 

Preliminary assembly

 

 

 

 

 

 

 

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Assemble the apparatus shown in figure 9.1.  A correctly assembled apparatus will be on display at the instructorÕs station as well.  Be sure that all stopper connections are tight so that no gas will be lost to leaks in the system.  Before attaching the test tube containing the weighed reaction mixture, use a pipet bulb to blow gently into the connecting tubing (from the end that is to be attached to the test tube) to start a siphon action between the one-liter flask and the 1000-mL beaker.  Siphon the water back and forth several times until you have removed all of the air bubbles.  Throughout the siphoning process, take care that no water enters the short piece of tubing in the two holed stopper that will ultimately be connected to the test tube.

 

Final assembly

 

1.     With the water level midway in the neck of the 1-liter flask, close the pinch clamp on the tube connecting the beaker and the flask.

2.     Attach the short piece of rubber tubing to the glass tube sticking out of the stopper of the test tube containing the KClO3/MnO2 mixture.  Make sure that all of the connections are air tight!

3.     This step comprises the equalization process.  In this step you are going to adjust the pressure inside of the 1-liter flask and the test tube (i.e., the system) so that it will be the same as the pressure on the outside, that is, the pressure in the room.  Then, the pressure inside of the assembly I known – it is simply the atmospheric pressure in the room which you will determine by reading the barometer.  The pressure of the system is to be equalized before the oxygen is produced and again at the end of the experiment just before the volume of the displaced water is measured.

Text Box: To equalize the system, open the pinch clamp and raise or lower the beaker till the level or height of the water in the beaker is the same as the level or height of the water in the flask. Close the clamp while maintaining the same water level in the flask and beaker. Again, the purpose of this process is to ensure that the pressure in the flask is the same as the pressure in the room, the atmospheric pressure.

4.     Take the glass delivery tube out of the beaker and empty the water out of the beaker.  Handle the delivery tube from the flask very carefully so as to not shake any water out of it.  Set the beaker down but do not dry it out with anything.  Return the beaker to its position under the delivery tube and reopen the pinch clamp.  From this point forward, DO NOT take the end of the delivery tube out of the water that will soon flow into the beaker.

5.     Only a few drops of water should flow into the beaker.  If water continues to flow into the beaker, there is a leak in the system and you must stop, fix the leak, and repeat the equalization process before you continue.  If you have a leak, make sure that the stopper is tightly seated in the flask.  Check that all of the rubber tubing-glass tubing connections are tight.  When you are sure they are all tight, repeat the procedure – flush the system and equalize the pressure in the system.

6.     At this point, BEFORE you start heating the system, you must have the system checked by the instructor.

7.     After the instructor gives you permission, make certain that the pinch clamp is open and start slowly heating the KClO3/MnO2 mixture in the tube.  After 1-2 minutes of gentle heating, turn up the heat level from the burner so that the mixture melts and starts to decompose.  Carefully watch the progress of the decomposition and regulate the heat level so that there is a slow but continuous evolution of oxygen.  Be sure to heat slowly or the oxygen will be generated faster than it can escape through the glass tube and it will shoot the stoppers out of the test tube or the flask meaning you must start over!  If the water level in the flask gets close to the bottom of the glass tubing that leads to the beaker, immediately stop heating the test tube.  If the water level lowers past the bottom of the glass tubing you must start over! As the reaction progresses, the mixture will begin to solidify and you will have to supply more heat.  When little gas seems to be evolving, heat the tube from all sides and from the end strongly to ensure that the last bit of KClO3 has been decomposed.  At this point, you should have 600 to 700 mL of water in the beaker.

The end of the tube that extends almost to the bottom of the flask must still be submerged under the water.  If the water level falls below the end of the tube, gas will escape and your results will be invalid.

8.     When no more oxygen is evolved, let the system cool to room temperature.  Make sure that the water delivery tube remains in the beaker of water and that the pinch clamp remains open. During the cooling, some water flows back into the flask to compensate for the contraction of the gas upon cooling.

9.     While you are waiting for the cooling, measure the barometric pressure in the room.  Your instructor will show you how to do this.

10.  When the system has cooled to room temperature, equalize the pressure and while the water levels are the same in the beaker and the flask, close the pinch clamp.

Be sure you have the pinch clamp securely in place before you do this next step! 

11.  Record the temperature of the gas in the flask by releasing the stopper slightly and inserting a thermometer into just the gas.  After the temperature comes to equilibrium, measure the temperature of the water remaining in the flask.

12.  When it is cool, weigh the test tube and KCl/MnO2, residue.   Measure the volume of water displaced into the beaker with a 1000-mL graduated cylinder.

Experiment #9

Measurement of the Gas Constant and the Molar Volume of Oxygen

 

Prelaboratory Exercise

 

1.     Write the balanced equation for the production of the oxygen gas in this experiment.

 

 

2.     How much of the KClO3 mixture are you to use in you experimental work? 

Between ______________ and ______________ grams

 

 

3.     In this experiment MnO2 is used as a catalyst. 

 

 

4.     What is the function of the MnO2 in this reaction?

 

 

5.     After you have set up your apparatus as shown in Figure 9.1 and you think that the system is air tight, what are you supposed to do next (i.e., before actually starting the experiment)?

 

 

 

6.     Before you start displacing the water with the oxygen has and again when you are all done, you must equalize the pressure inside of the system so that it is the same as that in the room.

a.     Why is this necessary?

 

 

b.     How is it done?

 

 

 

 

7.     A sample of oxygen gas is collected over water using the same procedure you will use in this experiment.  The system is equalized and the temperature of both the water and the gas is 22oC.  The atmospheric pressure is 745 mm Hg, what is the pressure of the O2 gas in the flask?

 

 

 

 

8.     Use the data below to answer the following:

 

mass of the empty test tube               18.767g

mass of test tube + sample                 20.989 g

mass of test tube + residue                 20.164g

mass of O2 liberated                            __________

volume of water displaced                  675 mL           

temperature of the O2                                     24oC

temperature of the water                    24oC

barometer reading                                742 torr

 

    Calculate the value of R in L atm/mol K from this data.

Table 9.1 Vapor Pressure of Water

 

T(oC)

Vapor Pressure of Water (mm Hg)

10

9.209

11

9.844

12

10.518

13

11.231

14

11.987

15

12.788

16

13.634

17

14.530

18

15.477

19

16.477

20

17.536

21

18.650

22

19.827

23

21.088

24

22.377

25

23.756

26

25.209

27

26.739

28

28.349

29

30.043

30

31.824

 

 

Calculation Hints

1.     From the  mass of the oxygen liberated, calculate the number of moles of O2 produced.  The pressure inside the system is equilibrated to atmospheric pressure.  As the oxygen is collected over water it is saturated with water vapor, so Patm=Poxygen+Pwater vapor.  The water vapor pressure at a variety of temperatures can be found in Table 9.1 above.  From the barometric pressure (i.e. the atmospheric pressure) and the water vapor pressure calculate the pressure of the O2.  From the volume or water displaced, determine the volume of O2 produced.

2.     Calculate the value of the gas constant, R, by substituting the known quantities from you measurements into the ideal gas equation and solving for R.

3.     Evaluate the molar volume, Vm, by using the value of R just calculated and the standard temperature and pressure (STP) conditions in the ideal gas equation.

4.     Compare both your value for R and for Vm to the accepted values.  This comparison should consist of percentage error calculated by

5.      

 

 

6.     Calculate the percent of KClO3 in the original mixture form the amount of oxygen gas produced.

 

 

Experiment #9

Measurement of the Gas Constant and the Molar Volume of Oxygen

Data and Calculation Sheet

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Data and Results Table

 

Mass of empty test tube (g)

 

Mass of test tube and KClO3/MnO2(g)

 

Mass of test tube assemble and residue (g)

 

Mass of O2, liberated (g)

 

Volume of H2O displaced (mL)

 

Temperature of O2 in flask (oC)

 

Temperature of H2O left in the flask (oC)

 

Barometric pressure (torr or mm of Hg)

 

Vapor pressure of water (mm Hg)

 

Calculated pressure of O2 gas (mm Hg)

 

 

 

 Experimental Value of R

 

Literature Value of R

62.4 L torr/mol K

Percentage error in R

 

Molar volume, Vm

 

Literature Value of Vm

22.4 L

Percentage error in Vm

 

Percentage of KClO3 in the original mixture

 

 

Sample Calculations

Post Lab Questions

 

1.     How would your result (i.e., the R value) be affected if you did not wait for the test tube to cool before measuring the volume of water?

 

 

 

 

 

 

 

2.     What is the volume of 1 mole of dry gas measured over water at a total pressure of 750 mm of Hg and 30¡C? (See Handbook for vapor pressure of water.)

 

 

 

 

 

 

 

 

3.     A mixture of gases at 760 mm pressure contains 65.0% nitrogen, 15.0% oxygen, and 20.0% carbon dioxide, by volume. What is the partial pressure of each gas in mm of Hg?

 

 

 

 

 

 

 

 

4.     What error would result from incomplete (say only 80%) decomposition of the KClO3 due to inadequate heating?

 

 

 

 

 

 

 

5.     Assume that 1.40 g of a KClO3 /MnO2 mixture (65% KClO3, by weight) was decomposed in the above experiment and all final measurements were made at 22¡C on a day when the atmospheric pressure was 716 torr. What volume of O2 would be evolved under the experimental condition? What is the STP volume of this O2?