Lab 2: Molar Mass of a Gas
Introduction
The ideal gas law is given by the following expression
PV = nRT
where
(1)
V = volume of the gas
P = pressure of the gas
n = moles of gas present
R = the ideal gas constant (0.08206 L atm mol-1 K-1 = 8.3145 J mol-1 K-1),
Note: J = Pa m³
T = temperature of the gas (in Kelvin)
The number of moles is equal to the mass of the gas, g, divided by the molar mass of the gas, M.
n
g
M
(2)
Substituting equation 2 into equation 1 and solving for M gives
M 
gRT
PV
(3)
If the values of g, T, P and V are all measured for a gas then the gas’s molar mass can be
determined by using equation 3.
The thermal decomposition of potassium chlorate (KClO3) is used to produce O2 gas.

2 KClO3(s)  2 KCl(s) +3 O2(g)
Manganese dioxide, MnO2, is used as a catalyst to speed up the reaction.
A weighed sample of potassium chlorate will be heated until all the oxygen is released. The mass
of oxygen, g O2 , generated will equal the difference between the mass of KClO3 before heating
and the mass of the residue, KCl, after heating.
g O2  g KClO3  g KCl
The temperature of the oxygen, TO2 , produced is assumed to be equal to the temperature of the
water.
The volume of oxygen, VO2 , produced will be equal to the volume of water displaced from a
water filled bottle. Since the volume of a gas depends strongly on its temperature, it is necessary
Document1
02/ 16
2-1
to allow the setup to cool to room temperature before the volume measurement is made. The
volume of displaced water is determined by measuring the mass of water displaced, g H 2O .
g H 2 O  gbeaker  H 2 O  gbeaker
Using the density, d H 2O , of water, the volume of water can be calculated. The density of water at
various temperatures can be found in the CRC Handbook.
VO2  VH 2O 
g H 2O
d H 2O
When the water levels in the flask and beaker are the same, then the pressure inside the flask is
equal to the barometric pressure of the room. The pressure in the flask is equal to the sum of the
pressure due to the oxygen and the vapor pressure of water, Proom  PO2  PH 2O . Rearranging
gives:
PO2  Proom  PH 2O
The vapor pressure of water at various temperatures can be found in the CRC Handbook.
Using these values for g O2 , TO2 , PO2 and VO2 along with the ideal gas constant R in equation 3,
the molar mass of oxygen is calculated.
Procedure
1.
Place the test tube in a small beaker and weigh them.
2.
Place approximately 1.0 grams of the potassium chlorate/manganese dioxide mixture into the
test tube.
3.
Accurately weigh the test tube containing the potassium chlorate mixture and beaker.
Record the weight.
4.
Nearly fill a 500 mL flask with water.
5.
Put about 200 mL of water into a 400 mL beaker.
Document1
02/ 16
2-2
6.
Setup the apparatus as shown in figure 1.
Tube A
Pinch Clamp
Tube B
7.
Figure 1
Place a rubber bulb onto tube A. With the pinch clamp open slowly squeeze on the bulb to
fill the tubes connecting the flask with the beaker. There should be no air bubbles in this
tube.
8.
Tightly fit the stopper with tube A into the test tube containing the potassium chlorate.
9.
Adjust the height of the beaker until the water in the flask is level with the water in the
beaker (figure 2). Close the pinch clamp. Empty the beaker. This sets the pressure inside
the flask equal to the room’s pressure.
Pinch Clamp
- - Level - KClO3
Figure 2
Document1
02/ 16
2-3
10. Place tube B back into the dry beaker and open the pinch clamp.
11. Gently heat the test tube with a Bunsen burner. Heat slowly first and more vigorously later.
Continue heating until no more water is going into the beaker.
12. Allow the test tube and setup to cool to room temperature. Keep tube B submerged into the
beaker of water
13. Adjust the height of the beaker until the water in the flask is level with the water in the
beaker. Close the pinch clamp.
14. Weigh the beaker and water. Record the weight.
15. Empty out the water and weigh the beaker. Record the weight.
16. Disconnect the apparatus and measure the temperature of the water in the flask. Record the
temperature.
17. Weigh the test tube and remaining solid in the same beaker as in step 3. Record the weight.
18. Record the atmospheric pressure.
19. Lookup and record the density and vapour pressure of water.
Document1
02/ 16
2-4
Lab 2: Molar Mass of a Gas
Date: _______________
Name: __________________________
Partner: _____________
Purpose:
Data and Results
Mass of test tube + beaker (g)
Mass of test tube + beaker
+ KClO3 before heating (g)
Mass of test tube + beaker
+ residue after heating (g)
Calculate the mass of oxygen generated (g)
Mass of beaker + water (g)
Mass of beaker (g)
Temperature of water (°C)
Density of water
Calculate the temperature of oxygen (K).
Document1
02/ 16
2-5
Calculate the volume of oxygen generated.
Barometric pressure
Vapour pressure of water
Calculate the pressure of oxygen.
Calculate the molar mass of oxygen (g/mol).
Calculate the % error in the molar mass.
% 𝐸𝑟𝑟𝑜𝑟 =
Document1
𝑀𝑒𝑎𝑠𝑢𝑟𝑒𝑑 𝑉𝑎𝑙𝑢𝑒−𝑇𝑟𝑢𝑒 𝑉𝑎𝑙𝑢𝑒
𝑇𝑟𝑢𝑒 𝑉𝑎𝑙𝑢𝑒
02/ 16
2-6
𝑥 100%
Questions:
1. What effect would a loss of O2 through a leak in the tubing have on the experimentally
determined molar mass? Explain.
2. Why is it necessary to allow the test tube to cool to room temperature before proceeding with
step 12? If this is not done, how would it affect the experimentally determined molar mass?
3. What effect would incomplete decomposition of the KClO3 have on the experimentally
determined molar mass? Explain.
Conclusions:
Document1
02/ 16
2-7