AP Chemistry
L5-1 Investigation of the Molar Volume Relationship in a Chemical Reaction
Purpose: To investigate the molar relationship involving mass, moles and volume.
Introduction: Both solids and gases often must be handled in the same experiment. The
amount of solid used or produced can be determined by measuring the mass of the
material on a balance but it is difficult to find the mass of a gas. For convenience the
chemist measures gas volume and calculates gas mass. It is, therefore, necessary for the
chemist to know the quantitative relationship between the molar mass and the molar
volume of a gas. Avogadro's hypothesis explains the relationship between the molar
volume, the molecular mass and the actual mass of a sample of a gas. Your
measurements will be taken at room temperature and pressure.
Objective: To investigate the chemical significance of Avogadro's hypothesis, the ideal
gas law, and Dalton's Law of partial pressures.
Equipment and materials:
•600 mL or 1000 mL beaker
•one hole rubber stopper (size 0)
•eudiometer tube
•utility clamp
•10 mL 3M HCl
•thermometer
•barometer
•ring stand
•3 cm Mg ribbon
•thin Cu wire
Things to remember:
Ideal gas law: PV = nRT where R = .08206 L•atm•mol -1• K-1
Dalton’s Law of Partial Pressure
Procedure:
1. On entering the lab, fill the beaker about one-half full of near room temperature tap
water and allow it to stand so that the temperature of the water may adjust to room
temperature.
2. Obtain a piece of Mg ribbon from Ms. Lowther. Measure the mass using the analytical
balance and record in your notebook.
3. Obtain enough Cu wire to make a basket to hold the Mg ribbon by wrapping around
the metal. Leave a portion of the wire loose to hang over the edge of the tube.
4. Prepare a ring stand with a utility clamp to support the tube.
5. Slowly pour about 10 mL of 3M HCl into the eudiometer tube.
6. Incline the tube (tip it) so air may escape and slowly fill the tube with tap water from
your beaker. Take care to mix the water and acid as little as possible.
7. With the tube completely full of water, insert the magnesium ribbon about 3 or 4 cm
into the tube. With the wire or thread against the side of the tube, insert a 1-hole stopper.
The stopper should force water and all air bubbles out of the tube and should hold the
thread or wire suspending the magnesium in place. Be sure there are no air bubbles
present! Tap the tube to dislodge any stubborn bubbles.
8. Cover the hole with your finger and invert the tube. Place the stoppered end into the
beaker full of room temperature water. Clamp the tube into place with the utility tube so
that the bottom of the tube is slightly above the bottom of the beaker. The reaction will
AP Chemistry
L5-1 Investigation of the Molar Volume Relationship in a Chemical Reaction
not start immediately. You have to wait for the acid to diffuse down through the water to
reach the Mg.
9. When the Mg has reacted completely and the evolution of gas has stopped, tap the
tube with your finger to dislodge any bubbles you see attached to the side of the tube.
10. Place your finger over the hole in the stopper and remove the tube from the beaker.
Lower the tube into the larger container of water used to equalize pressure and remove
your finger. Raise or lower the tube until the level of the water inside the tube is the same
as the level of water outside the tube. Read the scale on the tube as accurately as possible
(to the nearest 0.1 mL). This will give the volume of the gases (hydrogen and water
vapor) in the tube. Record the volume in your notebook.
11. Empty the contents of the tube and beaker and rinse both with tap water.
12. Record the temperature in degrees Celsius, the barometric pressure and the precise
mass of the Mg ribbon.
13. Clean up your area.
Calculations:
1. Calculate the moles of magnesium in the sample.
2. Calculate the partial pressure of hydrogen gas in the hydrogen gas-water vapor
mixture. (For the partial pressure of water vapor at different temperatures, consult a
reference.) Also, remember that the total pressure of a gas mixture is equal to the sum of
the partial pressures of each gas: PTotal = PH O+ PH . (Dalton's Law of Partial Pressures)
2
2
3. At this point, you know:
•the volume of the hydrogen gas (which you measured in the experiment)
•the temperature of the hydrogen gas (which you measured in the experiment)
•the pressure of the hydrogen gas (which you calculated in part 2).
You should now be able to use the ideal gas law to calculate the number of moles of
hydrogen gas present.
4. From your calculations, a fractional part of a mole of magnesium gave an
experimentally determined volume of H2 gas. Use this information to determine the
correct chemical equation. Start by writing down the reactants and products; this is the
skeleton chemical equation. Now compare the moles of magnesium to moles hydrogen
gas produced and this should give you an idea about the chemical equation.
?H+ + ?Mg = ?Mg+? + ?H2
5. Calculate how many liters the hydrogen would occupy if you had one mole of gas at
STP. Do this by comparing the ideal gas law for two situations: one you determined
experimentally and one for the ideal situation. Your only unknown should be the volume
of the ideal situation. Solve for V1 and substitute in your data for P,V, n, and T and the
values for P, T for STP and 1 mole for n. Calculate a % error by comparing to the
accepted value of 22.4 L/mol.
P1V1 n1RT1

P2V2 n 2 RT2

Situation 1 = ideal gas at STP
AP Chemistry
L5-1 Investigation of the Molar Volume Relationship in a Chemical Reaction
Situation 2 = your findings
P2V2
n2RT2