Chemistry Laboratory Manual
1
Experiment 11
1.1 Masses of equal volumes of gases
BACKGROUND
Avogadro’s Hypothesis states that ‘at the same temperature and pressure equal volumes of
different gases contain equal numbers of molecules’. If the mass of a known gas, say oxygen, in
a flask at a particular temperature and pressure is measured, and then the same flask is refilled
with another gas under the same conditions of temperature and pressure, and its mass
measured, the relative molecular mass of the second gas can be determined.
The purpose of this experiment is to determine the relative molecular mass of acetylene using
this method.
EQUIPMENT REQUIRED
Florence flask with rubber stopper (300ml)
Oxy-acetylene kit fitted with a delivery tube
Cotton wool
Thermometer (-10 - 100oC)
Barometer
Graduated cylinder (100ml)
Balance
PROCEDURE
#1
Weigh a stoppered florence flask on an accurate balance and record its mass.
#2
Remove the stopper and pass in oxygen from a cylinder for about 20 seconds. Wrap the
end of the delivery tube in a wad of cotton wool to prevent air diffusing back into the flask.
Carefully replace the rubber stopper and re-weigh the flask.
#3
Again pass oxygen into the flask for about 10 seconds, replace the stopper and re-weigh.
Repeat until the mass does not vary by more than one milligram.
SAFETY NOTE
 Acetylene is particularly flammable. Make sure there are no flames in the
vicinity.
#4
Fill the flask with acetylene in the same way and record its mass.
#5
Remove the stopper and fill the flask with water to the level of the bottom of the stopper.
Determine the volume of the flask by pouring the water into a graduated cylinder.
#6
Measure the room temperature and pressure.
PROCESSING OF RESULTS AND QUESTIONS
1. From the measured volume of the flask and the table below calculate the mass of air
initially in the flask. This will be used to calculate the mass of the empty flask and
stopper.
Mass of air in grams per litre
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Chemistry Laboratory Manual
at various temperatures and pressures
Pressure (kPa)
15
80.0
82.0
84.0
86.0
88.0
90.0
92.0
94.0
96.0
98.0
100.0
102.0
104.0
106.0
108.0
110.0
112.0
114.0
116.0
118.0
120.0
0.97
0.99
1.02
1.04
1.07
1.09
1.11
1.14
1.16
1.19
1.21
1.23
1.26
1.28
1.31
1.33
1.36
1.38
1.40
1.43
1.45
Temperature (oC)
20
25
0.95
0.97
1.00
1.02
1.05
1.07
1.09
1.12
1.14
1.16
1.19
1.21
1.24
1.26
1.28
1.31
1.33
1.36
1.38
1.40
1.43
0.94
0.96
0.98
1.01
1.03
1.05
1.08
1.10
1.12
1.15
1.17
1.19
1.22
1.24
1.26
1.29
1.31
1.33
1.36
1.38
1.40
30
0.92
0.94
0.97
0.99
1.01
1.03
1.06
1.08
1.10
1.13
1.15
1.17
1.20
1.22
1.24
1.26
1.29
1.31
1.33
1.36
1.38
2. Calculate the mass of the flask and stopper. Hence determine the masses of oxygen
and acetylene contained in the flask.
3. Determine the relative molecular mass of acetylene using the formula
mass of acetylene
mass of oxygen =
__relative molecular mass of acetylene__
relative molecular mass of oxygen (32.0)
Alternative method is to use the formula PV = nRT to determine the number of moles of
acetylene in the flask and then use
Number of moles =
__ mass of acetylene ________
Relative molecular mass of acetylene
to calculate the relative molecular mass of acetylene.
4. Acetylene has the formula C2H2. From this formula calculate its relative molecular mass.
How does your experimentally determined value compare?
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2
Experiment 12
2.1 Molar mass of butane
BACKGROUND
The purpose of this activity is to determine the molar mass of butane using gas from a butane
lighter. The experiment depends on the fact that the molar volume of an ideal gas is 22.4 L at
standard temperature and pressure (0oC and 101.3 kPa). By measuring the volume occupied by
a known mass of butane, and knowing that its molar volume is 22.4 L at S.T.P., it is possible to
calculate the molar mass of butane.
EQUIPMENT REQUIRED
Piece of glass tubing – approximately 60cm long with a diameter of approximately 25 mm,
stoppered at one end
Gas jar lid
Pneumatic trough (large)
Butane lighter
Retort stand, boss head and clamp
Thermometer (-10 – 110oC)
Graduated cylinder (250ml)
Marking pen
Barometer
PROCEDURE
#1
Almost fill the pneumatic trough with water.
#2
Fill the glass tubing with water and cover the open end with a gas jar lid making sure that
no bubbles of air remain. Invert the tubing and place the open end of the tube under the
surface of the water in the pneumatic trough. Remove the gas jar lid.
#3
Clamp the glass tubing in position.
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#4
Measure the temperature of the laboratory and obtain a reading of the atmospheric
pressure.
#5
Weigh the butane lighter on a balance to at least two decimal places.
#6
Hold the lighter upright, under water in the pneumatic trough and directly under the glass
tubing.
#7
Open the trigger of the lighter and allow the gas released to displace the water from the
glass tubing. Keep the trigger open until the water levels inside and outside the tube are
about the same. Remove the butane lighter
#8
Adjust the height of the glass tubing so that the water levels outside and inside the tubing
are the same. Mark this level on the outside of the glass tubing.
#9
Remove the glass tubing and determine the volume of gas collected by filling the tubing
to the mark with water and pouring this into a 250ml graduated cylinder.
#10
Very carefully dry the outside of the butane lighter with a paper towel, blow briskly over
the top of the lighter and dab carefully with the paper towel to ensure the lighter is
perfectly dry.
#11
Reweigh the lighter.
PROCESSING OF RESULTS AND QUESTIONS
1. Calculate the mass of butane released from the lighter.
2. Calculate the partial pressure of the butane by subtracting the vapour pressure of water
from the measured atmospheric pressure. Appendix 4 contains values of the vapour
pressure of water.
3. Calculate the volume the butane would occupy at S.T.P. and hence the number of moles
of butane released.
4. From the measured mass of butane calculate the molar mass of butane.
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