Lesson

Background

The Kinetic molecular theory of gases assumes that gas particles are in constant motion, have elastic collisions and no forces of attraction between particles. This theory allows us to treat gases under normal laboratory conditions as ideal gases. The Ideal gas law constant, R, relates the four variables that describe gases under normal laboratory conditions. The variables that describe gases include the pressure, volume, amount, and absolute temperature. These variables are related in the ideal gas law:

When insoluble gases are collected over water in a contained system, the total pressure of the gas collected is a sum of the partial pressure of the gas collected, and the vapor pressure of the water at that temperature. The pressure of the individual gas can be determined using Dalton’s Law of Partial Pressure:

Using both of these concepts along with reaction stoichiometry, the value of the ideal gas law constant can be determined experimentally.

Prelab Questions

1. Convert the following measurements:
   1. 735 mm Hg to atmospheres
   2. 23.5 mL to Liters
   3. 30.0˚C to Kelvin
2. What is the total pressure of a mixture of helium and hydrogen gas if the P_Helium is 225 mm Hg and the P_Hydrogen is 513 mmHg?
3. Write balanced single replacement reaction between aqueous hydrochloric acid and magnesium metal.
4. Determine the moles of hydrogen produced when 13.5g of magnesium reacts with excess hydrochloric acid.

Objective

• To learn how to collect a gas over water.
• To understand the meaning of a dry gas.
• To experimentally determine the value for the ideal gas law constant (R).

Materials

• 1 Piece of magnesium metal (2.0 – 3.0 cm)
• 20 mL of 3.0 M HCl
• 50 mL eudiometer
• Ring Stand with single buret clamp
• 1000 mL Graduated Cylinder
• Rubber stopper with hole(s)
• 10 cm String
• 600 mL Beaker
• Analytical balance
• Deionized water/distilled water

Safety

• Always wear safety goggles when handling chemicals in the lab.
• Wash your hands thoroughly before leaving the lab.
• Follow the teacher’s instructions for cleanup of materials and disposal of chemicals.
• When working with acids and bases, if any solution gets on your skin immediately rinse the area with water.
• Neutralize any acid spills with baking soda.
• Wear protective gloves when working with strong acids.

Procedure

1. Obtain a piece of magnesium ribbon between 2.0 and 3.0 cm long. Measure the mass of the magnesium using an analytical balance. Record in the data table.
2. Fold the piece of magnesium ribbon in half and tie one end of the string around it.
3. Tie the other end of the string through the hole in the rubber stopper.
4. Set up a ring stand and a utility clamp in position to hold a 50mL eudiometer. Place a 600 mL beaker 2/3 filled with tap water near the ring stand.
5. Tilt the eudiometer slightly and pour approximately 15-20 mL of 3.0 M HCl into it.
6. With the eudiometer still tilted, completely fill the eudiometer with deionized water. There should be no space left in the tube.
7. Place the magnesium ribbon into the eudiometer, add more deionized water if necessary and place seal the tube with the rubber stopper.
8. While wearing gloves, cover the hole(s) of the rubber stopper with your finger and invert the eudiometer into the 600mL beaker of water. Remove your finger after the stopper is below the surface of the water. Clamp the eudiometer in place so that it is not touching the bottom of the beaker. The denser acid will flow down the eudiometer and react with the magnesium.
9. Once the reaction stops (all of the magnesium should react), allow the eudiometer to sit 5 to 10 minutes so that it will return to room temperature.
10. Measure and record the temperature and pressure of the room in the data table.
11. Fill the large graduated cylinder with tap water.
12. While wearing gloves, transfer the eudiometer to the large graduated cylinder. Be sure to cover the hole(s) in the rubber stopper when transferring. Raise or lower the eudiometer until the level of water inside the eudiometer is even with the surface of water in the graduated cylinder. Record the volume of gas in the eudiometer. (This process ensures the pressure of the contained gas is equal to the pressure of the atmosphere).
13. Remove the eudiometer from the graduated cylinder and clean up.

Data

Calculations

1. Using the mass of magnesium measured, determine the number of moles of magnesium.
2. Convert the volume of gas collected from milliliters to liters.
3. Convert the atmospheric pressure of the room from inches Hg to mm Hg. Remember there are 2.54 cm in 1.00 inch.
4. Using Dalton’s Law of Partial Pressure and the vapor pressure of water, determine the partial pressure of dry hydrogen gas.
5. Convert the pressure of dry hydrogen gas from mm Hg to atmospheres.
6. Using the stoichiometry and the balanced chemical equation for the reaction, determine the moles of hydrogen gas produced in the reaction.
7. Using the Ideal gas law, calculate the value of ideal gas law constant, R.
8. Calculate the % error of the experimental value of the ideal gas law constant.

Analysis

1. Why should students wait 5 -10 minutes after the reaction is completed before measuring the volume of the gas?
2. What is the effect on the volume of gas measurement if bubbles were still in the eudiometer?
3. What effect on the ideal gas law constant would those bubbles have caused?
4. List two additional sources of error besides bubbles in the eudiometer.

Extension

1. The ideal gas law constant was calculated in this experiment, but real gases do not always behave ideally. What conditions cause the greatest deviations from the ideal gas law?
2. The Earth’s atmosphere is made up of many gases. What are the 4 most common gases of the atmosphere that a student breathes in with every breath?
3. Marcellus shale in Pennsylvania and West Virginia is a top source of what commercial gas mined in the United States? What is the name of the process used to mine the gas?