Gas Laws and the Stoichiometry of a Chemical Reaction
Introduction
Many chemical reactions involve the production of a gas from solid or liquid reactants. Examples include:
(1) CaCO3(s) à CaO(s) + CO2(g)
(2) Mg + 2HCl(aq) à MgCl2(aq) + H2(g)
(3) Zn + 2HNO3(aq) à Zn(NO3)2(aq) + H2(g)
Reactions (2) and (3) are similar in that their not ionic equations are identical:
metal + 2H+ à metal2+ + H2(g)
In this experiment, you will use pressure, volume, and temperature data to calculate the number of moles of H2 gas produced from a known mass of zinc or magnesium metal. Both reactions follow exactly the same pattern as the net ionic equation given above.
The Ideal Gas Law, PV = nRT, provides the necessary relationships to allow you to make your calculations.
Procedure
Production of Hydrogen Gas
You will do two trials, using Mg metal.
Step 1.
a. Obtain two pieces of Mg from the stockroom window. Weigh each piece and record the mass.
b. Fold the metal so that it can be enclosed in a small spiral of copper wire (Fig. 1) that has an extra length of wire attached to enable the case to be suspended in a gas collection tube. (Fig. 2) Do not place the metal sample in the gas collection tube at this time.
c. Set up a ring stand with a buret clamp to hold the gas collection tube in an upside down position (open end down) in your bowl filled 2/3 with water so that the open end of the tube is 1 inch below the surface of the water.
d. Disconnect the gas collection tube from the buret clamp. Add 10 ml, of 6 M HCl to the gas collection tube.
e. Tilt the tube slightly and slowly fill it with distilled water from a beaker. Try to minimize the mixing of the HCl and distilled water.
f. Holding the copper coil by the wire leader, lower the coil into the gas collection tube so that the metal in the cage is near the 50 mL mark on the tube. Hook the wire over the mouth of the tube and secure it in place by inserting a one or two holed rubber stopper into the mouth of the tube.
g. Make sure the tube is completely filled with distilled water, that no air bubbles exist in the tube, and that even the holes in the rubber stopper are filled with water.
h. Cover the hole in the rubber stopper with your finger and invert the tube. Place the tube back into position so that the rubber stopper is under water before removing your finger. This will prevent any air getting into the tube from the outside.
i. The metal will react with the acid and will eventually disappear leaving the copper coil untouched. Both reactions follow the same pattern: each metal atom reacts with two HCl’s to produce a +2 charged metal ion, two Cl- ions, and a molecule of H2. The tube should remain in position for five minutes after the reaction is complete. Any bubbles clinging to the sides of the tube can be dislodged by tapping the tube gently.
Step 2.
a. Record the volume of the gas to the nearest 0.1 mL.
b. Remove the gas collection tube from the water and pour the acid solution it contains down the sink. Rinse the tube with tap water.
Step 3.
a. Measure the atmospheric pressure using the barometer in the lab. Be sure to measure the pressure in units of mm Hg.
b. Measure the temperature of the water in the bowl, in ºC. Since the water and hydrogen gas have been in contact with each other, they should have identical temperatures.
Step 4.
The pressure of water for a specific temperature is listed in the CRC Handbook. Find and record the vapor pressure of water (mm Hg) for the temperature measured in step 3b. Hint: 1 KPa = 7.5006 mm Hg
Step 5.
The pressure in the gas collection tube was set to the atmospheric pressure (step 2a). The pressure inside the tube is the sum of the pressures of the water vapor and the hydrogen gas, according to Dalton’s Law of Partial Pressures. Calculate the pressure of the hydrogen gas in the collection tube.
Step 6.
Use your data and the ideal gas equation, PV = nRT, to find the number of moles of hydrogen gas in the tube. Make sure that your units are compatible for calculations!
Step 7.
Use the chemical equation and the number of grams of metal to find the number of grams of hydrogen that should have been produced (the theoretical yield.)
Step 8.
Convert the number of moles of hydrogen produced to the grams of hydrogen produced.
Step 9.
Calculate the percent yield for your experiment.
Report both sets of data and results in your formal report.
Date ______Name ______
Production of Hydrogen Gas
Recorded Measurements / Trial 1 / Trial 21. Mass of metal
2. Volume of H2 gas
3. Barometric pressure
4. Temperature of water
5. Vapor pressure of water
Calculations Show your mathematical set-ups for each trial
6. Pressure of hydrogen gas (Patm – PH2O):
Trial 1:
Trial 2:
7. Moles of hydrogen gas (based on n = PV / RT):
Trial 1:
Trial 2:
8. Actual yield of hydrogen gas in grams (based on moles from #7):
Trial 1:
Trial 2:
9. Theoretical yield of hydrogen gas in grams (based on mass of Mg used):
Trial 1:
Trial 2:
10. Percent yield of hydrogen gas:
Trial 1:
Trial 2:
Average:
11. % deviation of your trials ([High Trial – Low Trial] / Average) x 100:
Question
Stu Dent, a forgetful chemistry student, performed this experiment using iron metal and recorded the following data:
1. Mass of metal 0.089 g
2. Volume of hydrogen gas 40.5 mL
3. Barometric pressure 755.2 mm Hg
4. Temperature of water 24.2 ºC
When Dr. H performed the calculations to get the % yield of H2, he forgot to take into account the vapor pressure of water. Circle the most likely effect that Dr. H’s memory lapse had on the calculated % yield and explain?
a. The calculated % yield was too high.
b. The calculated % yield was too low.
c. The calculated % yield was over 100%.
d. The calculated % yield was not affected.
e. The effect on the calculated % yield cannot be determined unless more data is provided.
Explanation: