Molecular Mass of a Gas
Calculations
You will use the relationship:
PV = nRT
in your calculations. Since
Background
Since there is so much room between gas molecules at
normal pressures, the size or mass of the molecules does
not play a significant role in determining how many gas
molecules can fit into a volume. Avogadro’s hypothesis
states that equal volumes (T and P the same) of gases
contain the same number of molecules, no matter which
gases are involved. The gas law equation works because
of this property of gases. The following procedure could
be used for any gas. Lower molecular weight gases
would require more accurate balances to get decent
results.
g
M
(equation 2)
PV =
gRT
M (equation 3)
M=
gRT
PV n=
(M in this case stands for molar mass),
we can write:
€
Rearranging the equation, we have
Experiment
1. Weigh a 60 cm3 plastic syringe to the nearest 0.001
g, being careful that the syringe is not touching
anything but the balance pan (watch so it does
not touch the sides of the draft shield). Record the
weight on the data sheet.
2.At the fume hood, remove the cap from the syringe,
momentarily open, then close, the small black valve
on the regulator to fill the hose with gas from the
tank, and then attach the syringe to the hose on the
gas tank. Carefully open the small black valve on the
gas regulator, allowing the gas to push the piston up
a bit past the 60 cm3 mark. Turn the valve off, detach
the syringe from the hose, and push the piston in to
the 60.0 cm3 or less mark.
3. Replace the cap on the syringe and weigh. Record
the mass on the data sheet in the trial 1 slot.
4. Record the volume (to the nearest 0.1 cm3) on the
data sheet.
5.Use a thermometer to measure room temperature.
Record in the data table.
6. Back at the fume hood, push the gas out of the
syringe, then fill it a second time as before. Record
the mass under trial 2 on the data sheet.
7.Assume that the mass of the empty syringe remains
the same for trials 1 and 2. Record the volume for
trial 2.
8. Record room temperature for trial 2.
9.Assume that the temperature of the gas in the syringe
is the same as room temperature.
10. Read the atmospheric pressure on the barometer.
Record in the data table.
(equation 1)
€
(equation 4)
Rearrange the equation again so that grams are isolated:
€
g =
(equation 5)
Note: Check your algebra carefully for equation 5. You
will use the equation in your calculations.
Buoyancy Factor
We are all buoyed up by the air, just as we are buoyed
up by water. Of course, since the air is much less
dense than water, the amount of buoyancy is much
less. When weighing dense objects, buoyancy due to
air is usually not a factor. When weighing low density
objects, buoyancy must be taken into account. In order
to calculate the true weight of the gas in the syringe,
you will have to figure in a buoyancy factor. The filled
syringe is lighter by the weight of air displaced by the
filled syringe. To calculate that weight, use equation 5.
Air has a Mave = 29.0 g/mol. (Can you reason why this is
so?) V will be the volume of the gas in the second entry
of the calculation table. T is room temperature, and P is
barometric pressure.
Use R = 62.4 L·torr·mol-1·K-1 in your calculations.
Use the proper number of significant digits throughout
the calculations. Show the mathematical setup in the
calculation column.
1
Data Table
Trial 1
Trial 2
1. Mass of empty syringe and cap (to nearest 0.001 g)
g
2. Mass of filled syringe and cap (to nearest 0.001g)
g
g
3. Volume of gas (to nearest 0.1 cm3)
cm3
cm3
4. Room temperature (to nearest 1°C)
°C
°C
5. Barometric pressure (to nearest torr)
torr
Calculations
Calculation
Trial 1
Trial 2
Room temperature (Kelvin)
K
K
Volume of gas in liters
(Keep 3 significant digits)
L
L
Buoyancy factor
(use equation 5)
g
g
Corrected syringe weight
(#2 above + buoyancy factor)
g
g
Mass of gas
(corrected syringe weight - #1)
g
g
%
%
Molecular massexp (massexp) of gas
(use equation 4)
M exp − Mformula
x 100
Mformula
: substance is SF6 )
% error =
(Mformula
2
Name_________________________________________ Grade___________ Date ___________
questions
1. The density of a gas is usually given in g/L. Rearrange equation 3 on page 1 to isolate g/V, then calculate the
densities of the following gases at 25°C and 1.00 atm. V will be 1.00 L. Use R = 0.0821 L·atm/K·mol.
a) H2
b) N2
c) SF6
2. The reaction: NaHCO3(s) + HCl(aq) → CO2(g) + NaCl(aq) + H2O(l) produces carbon dioxide gas.
If the gas were produced at 20°C, where the vapor pressure of water is 18 torr, how many grams of NaHCO3 would
be required to produce 60.0 ml of wet CO2? Assume an atmospheric pressure of 763 torr, and an excess of HCl.
3. Show the Lewis structure for SF6. What is its VSEPR designation. Sketch the shape of the molecule. Is the
molecule polar? Give this same information for SF2 and SF4.
4. Sulfur hexafluoride is expensive. The tank used in today’s lab cost about $200, and originally contained about
5 pounds of the substance. How much did each syringe-full that you used today cost? (One pound = 454 grams)
1