Molar Volume of a Gas
PRE-LAB DISCUSSION
Avogadro’s hypothesis states that equal volumes of all gases contain equal numbers
of molecules under the same conditions of temperature and pressure. It follows from this
hypothesis that all gas samples containing the same number of molecules will occupy the
same volume under the same conditions of temperature and pressure. A special name is given
to the volume occupied by 1-mole samples of gases at STP. This volume is called the molar
volume. In this experiment, you will make an experimental determination of the molar
volume.
The basis of this experiment is the following reaction in which you will react a known
mass of magnesium with excess hydrochloric acid to produce the substances shown:
Mg(s) + 2HCl(aq)  MgCl2(aq) + H2 (g)
The hydrogen gas is the product that is of interest to you in this experiment. You will make
an experimental determination of the number of moles of hydrogen molecules produced and
the volume occupied by these molecules. The number of moles of hydrogen will be
determined indirectly. The balanced equations for this reaction shows that the molar ratio
of magnesium reacted to hydrogen gas produced is 1:1. Therefore, by determining the mass
of magnesium that reacts and the number of moles that this mass is equal to, you will also
be able to determine the number of moles of hydrogen gas produced. The volume of
hydrogen gas produced will be measured directly on the scale of a gas-measuring tube. The
gas laws of Boyle and Charles will be used to correct this volume, measured under laboratory
conditions, to the volume the sample of gas would occupy at STP. The collected data
(number of moles and volumes at STP) will be used to calculate that molar volume of the
hydrogen gas.
This experiment should aid in the understanding of the mole concept and the
concepts of molar volume of a gas.
PURPOSE
Determine the volume of 1 mole of hydrogen gas at STP using experimental data, known
mathematical relationships, and a balanced chemical equation.
EQUIPMENT & MATERIALS
Gas-measuring tube
Graduated cylinder, 10-mL
Graduated cylinder, 1000-mL
Metric ruler
Rubber stopper w/ one-hole stopper
Safety glasses
Thermometer
Copper wire – 15 cm (HCl) ~ 10 mL
Hydrochloric acid 3.0 M
Magnesium ribbon (Mg) ~ 5 to 7 cm
Water, distilled ~ 110 mL
PROCEDURE
1. Obtain a piece of magnesium ribbon from your teacher. Measure the length of the piece
of ribbon (to the nearest 0.1 cm). Record the length as (a) in your data table. Also
record the mass of 1 meter of magnesium ribbon. These data will be provided by the
teacher. Record as (b).
2. Obtain a piece of copper wire about 15 cm long. Wrap one end of the wire around the
piece of magnesium ribbon, leaving about 10 cm of thread free. Bend the piece of
magnesium so that it will fit easily into the gas-measuring tube.
3. Obtain about 10 mL of 3 M hydrochloric acid (HCl). CAUTION: Handle this acid with
care. Carefully pour the HCl into a gas-measuring tube.
4. Tilt the gas-measuring tube slightly. Using a beaker, slowly fill the gas-measuring tube
with water at room temperature. Try to avoid mixing the acid and water as much as
possible.
5. Lower the piece of magnesium ribbon (coiled and held in place by copper wire) 4 or 5 cm
into the gas-measuring tube. Drape the wire over the edge of the tube and insert the
one-hole rubber stopper into the tube as shown in Figure 1.
6. Add about 900 mL of water at room temperature to a 1-L graduated cylinder. Place your
finger over the hole in the rubber stopper and invert the gas-measuring tube. Lower the
stoppered end of the tube into the large graduated cylinder of water. Record your
visual observations as (g) in the data table.
7. Let the apparatus stand about five minutes after the magnesium has completely
reacted. Then, tap the sides of the gas-measuring tube to dislodge any gas bubbles that
may have become attached to the sides of the tube.
8. Move the tube up or down (to equalize pressure) until the water level in the tube is the
same as that in the graduated cylinder. On the scale of the gas-measuring tube, read
the volume of the gases in the tube. Record this volume as (c) in your data table.
9. In the data table, record the room temperature (d), and the barometric pressure (e).
10. If time permits and your teacher so indicates, repeat the experiment. If time does not
allow, then gather data from one or two other groups to average with your own.
Figure 1: Initial Setup
Figure 2: Final Setup
OBSERVATIONS AND DATA
Measurement
(a) length of Mg
ribbon
(b) mass of 1 meter
of Mg
(c) volume of H2 gas
in tube
(d) room
temperature
(e) barometric
pressure
(f) water vapor
pressure at room
temperature, PH2O
(from textbook
appendix)
Trial 1
Trial 2
Trial 3
cm
cm
g
g
cm
g
mL
mL
mL
°C
°C
°C
mm Hg
mm Hg
mm Hg
mm Hg
mm Hg
mm Hg
(g) visual
observations
CALCULATIONS
1. Find the mass of the Mg strip:
For example:
0.070 meter
1.00 meter
=
x grams
mass of 1-meter length
2. Calculate the number of moles of H2 formed (Hint: use stoichiometry. Convert from g
of Mg to moles of Mg to mol H2)
# moles H2 =
3. Find the pressure exerted by the H2 gas in the tube:
PH2 = Pbarometric - PH2O
4. Convert room temperature from °C to Kelvin:
K = °C + 273
5. Find the volume of the H2 gas at STP:
P1V1 = P2V2
T1
T2
OR
V2 = (V1 x P1 x T2) ,where:
(P2 x T1)
P1 = PH2 (from calculation 3)
V1 = experimental value of H2 (c)
T1 = room temperature (K)
T2 = 273 K
P2 = 760 mm Hg
V2 = volume of H2 at STP
6. Find the volume of 1 mole of H2 gas at STP:
Moles of H2 gas (from calculation 2)
Volume of H2 gas (V2)
=
1 mole
x mL
x = ______________ mL/mole = _____________L/mole
7. Calculate your percentage error for the volume of a mole of gas at STP.
DISCUSSION & CONCLUSION
In paragraph form, follow the instructions on the rubric.
Don’t skimp here, this is one of the most important parts of your report!