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The Gas Laws: Finding the Molar Mass of Oxygen
PRE-LAB ASSIGNMENTS:
To be assigned by your lab instructor.
STUDENT LEARNING OUTCOMES:


Learn how to manipulate the Ideal Gas Law and Dalton’s Law of Partial Pressure.
Learn how to use the Ideal Gas Law to find the molar mass of a gas.
EXPERIMENTAL GOALS:
The goal of this experiment is to measure the volume of water displaced during the
decomposition of potassium chlorate to find the volume of oxygen gas generated, and to use that
to determine the number of moles of oxygen generated, and to find its molecular weight.
INTRODUCTION:
In this experiment oxygen gas is formed by the decomposition of potassium chlorate:
2KClO3(s) 
2KCl(s) + 3O2(g)


The potassium chlorate, with the catalyst manganese(IV) oxide, MnO2, is placed in a test tube,
and the weight is determined accurately. The test tube is then heated, and oxygen gas is evolved
as the potassium chlorate decomposes. (Potassium chlorate decomposes rather slowly when
heated. The MnO2 acts as a catalyst for the reaction, speeding the reaction up without itself
being consumed.) The entire weight loss of the test tube can be assumed to be due to the mass of
the oxygen gas which is formed. This oxygen gas is passed into a bottle containing water, and
displaces a volume of water equal to the volume of oxygen gas produced; the water is forced out
of the bottle through a hose, and is collected in a beaker, where its volume can be measured.
After measuring the temperature and pressure of the oxygen gas, the molar mass can be
calculated.
The apparatus is balanced against atmospheric pressure so that the pressure of the gas
which is collected is the same as that day’s atmospheric pressure. The gas which is collected is
actually a mixture of oxygen and water vapor, since gases bubbled through water pick up small
amounts of moisture. The total pressure of the gas is the sum of the pressure of the dry oxygen,
PO 2 , and the vapor pressure of water, Pw (from Dalton’s Law of Partial Pressures),
Ptotal  PO 2  Pw
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The pressure of the dry oxygen is determined by subtracting the vapor pressure of water from the
total pressure. Once the pressure of the dry oxygen gas is known, the ideal gas law is used to
determine the number of moles of oxygen:
PV = nRT
where P is the pressure of the gas, V is its volume, n is the number of moles of the gas, T is the
Kelvin temperature, and R is the universal gas constant (0.08206 L atm K-1 mol-1). The molar
mass of the oxygen can then be determined by dividing the mass of the sample by the number of
moles of the sample.
PROCEDURE:
1. Obtain the apparatus shown in Figure 1 (tubing assembly, 2.5 L bottle, 800 mL beaker, and
pinch clamp). Obtain from the stockroom a heavy-walled Pyrex test tube (ignition tube)
containing the KClO3/MnO2 mixture.
2. Fill the large (2.5 L) bottle with tap water. Insert the stopper/delivery tubes into the bottle.
Do not attach the test tube yet.
3. Fill the 800 mL beaker about one-third full of tap water. Open the pinch clamp and use a
squirt bottle labeled “MW of O2 Lab” to blow through the tubing at point A until water runs
into the beaker. Close the pinch clamp, keeping it toward the bottom of the rubber tubing.
4. Weigh the Pyrex test tube and KClO3/MnO2 mixture (without stopper). Spread the mixture
out along the lower portion of the test tube for a couple of inches.
Figure 1. Experimental setup.
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5. To make the pressure in the 2.5 L bottle equal to the room air pressure, attach the test tube to
the assembly. Keeping the end of the glass tube extension under water in the beaker, open
the pinch clamp and raise the beaker until the water level in the beaker is the same as the
water level in the 2.5 L bottle. Hold the beaker in this position for 10-15 seconds, then close
the pinch clamp and empty the beaker.
6. To test for leaks, open the pinch clamp. Only a very small amount of water, if any, should
dribble out of the glass extension tube; then it should stop. If water continues to flow, the
assembly has an air leak. Consult your instructor before proceeding.
7. If your apparatus has no air leaks (water stops running), you are ready to begin heating the
test tube. Leave the pinch clamp off and do not discard the small amount of water collected
in the beaker. Have your instructor check your apparatus before proceeding.
8. Heat the upper part of the KClO3/MnO2 mixture, slowly at first, then increase the heating
until a steady stream of water is flowing into the beaker. If necessary, move the heating zone
toward the lower end of the mixture. Continue heating until about 200-300 mL of water has
been collected in the beaker, then turn the Bunsen burner off, and let the entire assembly
stand undisturbed with the glass tube extension under water and pinch clamp open for ten
minutes. Do NOT attach the pinch clamp yet!
NOTE: Do not let the beaker overflow! If it starts to do so, catch any overflow in another
beaker. Also do not let the water in the 2.5 L bottle drop below the delivery tube. If it does
so, stop the experiment and start again.
9. Keeping the end of the glass tube extension under water, raise the beaker (or bottle) until the
levels of water are the same in both. Close the pinch clamp. Do NOT disconnect the
ignition tube until the other end of the apparatus is closed by the pinch clamp.
10. Measure the volume of water collected in a large graduated cylinder.
11. Remove the test tube (without the stopper) from the assembly and weigh it.
12. Record the barometric pressure of the room, which is written on the blackboard. Measure the
temperature of the water in the beaker.
13. Repeat the experiment, beginning at step 5 above. Since you were given enough
KClO3/MnO2 for two runs, use the final weight of the test tube and contents from the first run
as the initial weight for the second run.
14. After completing both runs, put your used sample test tubes in the beaker designated for that
purpose.
15. Dismantle the equipment and return it to its proper location.
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149
LAB REPORT
Finding the Molar Mass of Oxygen
Name ________________________________
Date _________
Partner ________________________________
Section _________
Report Grade ______
First
Determination
Second
Determination
1.
Weight of test tube + KClO3/MnO2
______________ ______________
2.
Weight of test tube after heating
______________ ______________
3.
Weight of oxygen evolved
______________ ______________
4.
Volume of water (= volume of oxygen evolved) (V)
______________ ______________
5.
Temperature of water (= temp. of oxygen) (T)
______________ ______________
6.
Barometric pressure (Ptotal)
______________ ______________
7.
Vapor pressure of H2O at T (from Appendix IIIB)
______________ ______________
8.
Partial pressure of oxygen (show calculations)
______________ ______________
9.
Partial pressure of oxygen in atm (PO2) (show ______________ ______________
calculations)
10. Moles of oxygen (use PV=nRT) (show calculations)
______________ ______________
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First
Determination
11. Molar Mass of oxygen (show calculations)
Second
Determination
______________ ______________
12. Average Molar Mass of oxygen
______________
13. Theoretical Molar Mass of oxygen (from periodic
table)
______________
14.
______________
% Error 
(your valu e - true value)
 100
true value
(show calculations)