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DETERMINATION OF THE MOLAR VOLUME OF A GAS
Purpose
To experimentally determine the molar volume of a sample of collected hydrogen gas at STP.
The volume of a sample of hydrogen gas is measured at room conditions. This measurement is
used in ideal gas law to calculate the molar volume of hydrogen gas at standard temperature and
pressure.
Background Information
In class, we have discussed general properties of gases and several laws that describe these
gases. In your background section, some topics to consider and write about include: gas
pressure, Kinetic Molecular theory, Boyle’s Law, Charles’s Law, and Molar Volume at STP.
In this experiment, you will produce hydrogen gas, H2 , by reacting magnesium with
hydrochloric acid according to the following equation:
Mg + 2 HCl  MgCl2 + H2
You will assume the hydrogen is an ideal gas, and you will measure its mass, volume,
temperature, and pressure. From these measured values, you will calculate the molar volume of
hydrogen and compare the result with the ideal value below.
For a specific amount of any ideal gas, the relation between the pressure of the gas, P, its
volume, V, its temperature, T, and its number of moles, n, is given by the ideal gas law:
PV=nRT
Here R is the proportionality constant, called the gas constant, and it has the same value for all
ideal gases under all conditions, namely, 0.0821 L·atm/mol·K. In all calculations in which this
constant is employed, pressure must be expressed in atmospheres, volume in liters, and
temperature in Kelvins.
o
Standard conditions are defined as exactly 1 atm pressure and 0 C (273 K). The molar volume
of a gas is the volume that 1.000 mol of it occupies under these conditions. It is the same for all
ideal gases:
V = nRT/P = 22.4 liters
Materials
Eudiometer (Buret), Hydrochloric acid (6 M), Magnesium ribbon, Copper wire, Thermometer
Hypothesis / Pre-lab Questions
1) What should the molar volume of any gas at STP be?
2) What type of reaction is this?
3) Find the volume of 80 g O2 at STP.
4) How many liters would 0.25 mol of any gas occupy at STP?
Procedure
1. Measure a piece of magnesium ribbon 18 mm long. Do not exceed 18 mm. This strip has
been pre-measured so that it will not produce more hydrogen than the collection tube will hold.
2. Weigh the ribbon to the nearest milligram (0.001 g) and record the mass on the report sheet. If
the balance is less precise, you may determine the mass by using the conversion factor of 0.1
g/cm of magnesium ribbon.
3. Produce and collect the hydrogen gas as follows: Claim a gas collection tube (already set up
in the lab). Fold up the magnesium ribbon into a small, tight bundle. Tie it with a piece of
copper wire that is 10 to 15 cm long.
4. Add about 10 mL of 6.0M HCl to the gas collection tube. CAUTION. Be sure to use
hydrochloric acid; others might react violently when the water is added. Then fill the tube
completely with tap water, until it is nearly overflowing.
5. Place the magnesium in the mouth of the tube so that it is about 3 cm below the surface of the
water. Fold the copper wire extension over the side of the tube. Insert a one- or two-hole stopper
into the opening so that the cage is held firmly in place.
6. Holding your finger over the stopper hole(s), invert the tube into a 400 mL beaker that is about
half-filled with water. Then clamp the tube in place as shown in the diagram, with its mouth
below the water's surface. (There is no need to rush this maneuver. The acid will take more than
a minute to diffuse down to the stopper, and by then it becomes dilute enough not to harm your
finger.
7. Observe the reaction. When no more hydrogen bubbles are visible, the reaction is complete.
Wait an additional 5 minutes so that the hydrogen gas comes to room temperature.
8. Cover the hole in the stopper with your finger and transfer the tube to a large cylinder or
battery jar filled with water. Lower or raise the tube until the liquid level on the inside of the
tube is the same as the outside. Record the volume of the hydrogen gas. (Be sure to read the
liquid level at eye level.)
9. Measure and record room temperature. Obtain the vapor pressure of water from the table.
Record.
10. To find the partial pressure exerted by the hydrogen, you must recognize that the atmospheric
pressure equals the partial pressure of hydrogen gas in the tube (when the water levels are equal),
plus the partial pressure of water vapor mixed with the hydrogen (according to the equation
below).
Patm = PH2 + PH2O
Temperature
(°C)
16
17
18
19
20
21
22
23
24
25
26
27
28
29
Vapor
Pressure
(mm Hg)
13.6
14.5
15.0
16.5
17.5
18.7
19.8
21.1
22.4
23.8
25.2
26.7
28.3
30.0
11. Find today’s barometric (atmospheric) pressure online and record in mm of Hg.
12. Calculate the absolute (kelvin) temperature.
13. Using the combined gas law equation (shown below), convert your measured volume of
o
hydrogen (from Step 8) to conditions of STP, 1 atm pressure and 0 C (273 K). Give the result in
units of liters (L).
PV/T (today’s conditions) = PV/T (STP)
14. From the mass of magnesium calculate the moles of magnesium reacted.
15. Using the balanced equation, calculate the moles of hydrogen produced.
Mg + 2 HCl  MgCl2 + H2
16. Calculate the molar volume by dividing the calculated volume at STP by the moles of
hydrogen produced.
17. Compare your experimental molar volume to the theoretical value (22.414 L) by calculating
the percent error in the molar volume. ( experimental value – theoretical value x 100 )
theoretical value
18. Enter the data after you answered the questions according to the corresponding step in the
procedure.
Data Table
1. Mass of magnesium strip (g)
2. Moles of Magnesium used (mols)
3. Temperature of the hydrogen (° C)
4. Temperature of the hydrogen (K)
5. Volume of the hydrogen collected (water levels must
be equal when measurement is made) (mL)
6. Barometric pressure (mm Hg)
7. Vapor pressure of the water (mm Hg)
8. Pressure of dry hydrogen gas (mm Hg)
9. Moles of hydrogen gas produced according to
stoichiometry
10. Volume of hydrogen gas at STP
11. Molar volume of hydrogen at STP in L
12. Percent error
Discussion Questions
1. For the following possible errors, give the direction they would cause your results to deviate
from the accepted value of 22.4 L/mol at STP. Tell whether the following errors would
increase, decrease, or have no effect on the experimental molar volume. Explain your answers.
a) the measured mass of the magnesium was too small
b) not all the Mg ribbon reacted
c) not all the MgO coating is removed from the Mg ribbon before beginning the experiment
d) the actual temperature of the hydrogen is less than room temperature.
e) the water level measured was lower than the actual level of water.
f) you read the volume when the water level in the tube was higher than outside the gas
collection tube in step 9
g) air bubbles were introduced when tube is turned over during step 8 (the transfer of the gas
collecting tube to the big jar)
2. Why did you have the percent error that you obtained?
3. All real gases deviate to some extent from the behavior of ideal gases. At standard conditions,
the density of O2 gas is 0.0014290 g/mL, that of H2 gas is 0.00008988 g/mL, and that of CO2 is
0.0019769 g/mL. Using these values and the molecular weights from your textbook (in g/mol,
not rounded off), calculate the molar volume of each of these, in mL/mol, to five significant
figures. (Hint: analyze the units for density, molecular weight, and molar volume.)
Correlate the values of these three gases with the molar volume of an ideal gas (22414 mL/mol).
a) Which gas deviates least?
b) Which gas deviates most?
c) Suggest a reason for this behavior.
4. What is the purpose in knowing the molar volume of a gas?
5. Why should the molar volume of all gases be the same?