AP Chemistry Laboratory #6: Determination of the Molar Volume of a Gas Lab day: Thursday, October 12, 2017
Lab due: Friday, October 13, 2017 (end of hour)
Goal (list in your lab book): The goal of this lab is to experimentally determine the molar volume of hydrogen gas at STP. Introduction/Background (DON’T write in your lab book): Avogadro’s law states that equal volumes of gases contain an equal number of particles under the same conditions of temperature and pressure. It follows, therefore, that all gas samples containing the same number of particles will occupy the same volume if the temperature and pressure are kept constant. The volume occupied by one mole of a gas is called the molar volume. In this experiment, the molar volume of hydrogen gas at STP will be measured. The reaction of magnesium metal with hydrochloric acid provides a convenient means of generating hydrogen gas in the lab. If the reaction is carried out with excess hydrochloric acid, the volume of the hydrogen gas obtained will depend on the number of moles of magnesium as well as the pressure and temperature. The molar volume of hydrogen gas can be calculated if the volume occupied by the sample containing a known number of moles of hydrogen is measured. Since the volume will be measured under laboratory conditions of temperature and pressure, the measured volume will be corrected to STP before calculating the molar volume. Hydrogen gas will be collected by the displacement of water in an inverted gas measuring tube. The total pressure of the gas in the tube will be equal to the barometric pressure. However, the gas in the cylinder will not be pure hydrogen! The gas will also contain water vapor due to the evaporation of the water molecules over which the hydrogen is being collected. According to Dalton’s Law of Partial Pressures, the total pressure of the gases in the tube will be equal to the partial pressure of the hydrogen plus the partial pressure of the water vapor. The vapor pressure of water depends only on the temperature. Research questions (answer in your lab book in complete sentences; don’t write the questions): 1) Write and balance the reaction equation for hydrochloric acid reacting with magnesium. Include states of matter. 2) Write out the combined gas law (or rearrange the ideal gas law so that only number of moles and the gas constant are kept constant). 3) What is standard pressure (a) in atmospheres? (b) in mm Hg? 4) What is standard temperature (a) in Celsius? (b) in Kelvin? 5) State Dalton’s Law of Partial Pressures. 6) A reaction of 0.028 g of magnesium with excess hydrochloric acid generated 31.0 mL of gas. The gas was collected by water displacement in a 22 °C water bath. The barometric pressure in the lab that day was 746 mm Hg. (NOTE: Vapor pressure table on back of this lab handout or search for “Cabrillo College Vapor Pressure” online) (a) Use Dalton’s Law of Partial Pressures and the vapor pressure of water at 22 °C to calculate the partial pressure of hydrogen gas in the gas measuring tube. (b) Use the combined gas law or ideal gas law to calculate the “corrected” volume of hydrogen gas at STP. (c) What is the theoretical number of moles of hydrogen that can be produced from 0.028 g of magnesium and excess hydrochloric acid?
(d) Divide the corrected volume of hydrogen gas from question 6b by the theoretical number of moles of hydrogen gas in 6c to calculate the molar volume of hydrogen gas. Report your answer in L/mol. Materials (DON’T list in your lab book): 30 mL 2 M hydrochloric acid 1 barometer 1 balance 1 permanent marker 1 pair of scissors 2 600 mL beakers 15 cm copper wire 1 1000 mL graduated cylinder 1 one-hole stopper to fit gas measuring tube
4 cm magnesium ribbon 1 massing cup 1 50 mL beaker 1 test tube brush 1 ruler 1 gas measuring tube 1 digital thermometer 1 25 mL graduated cylinder 1 ringstand 1 test tube clamp
Hazards (list in your lab book): (include the safety contract and the hazards of the magnesium, hydrochloric acid, and hydrogen gas) Procedure (DON’T list in your lab book): 1. Prepare materials a. physically and chemically clean the gas measuring tube, 50 mL beaker, 25 mL graduated cylinder, and the scissors b. dry the 50 mL beaker, 25 mL cylinder, and the scissors c. label the 50 mL beaker and massing cup 2. Fill one 600 mL beaker with about 400 mL of tap water 3. Mass magnesium a. cut a 2 cm piece of magnesium ribbon b. mass the magnesium ribbon using a massing cup 4. Prepare the magnesium a. twist and fold one end of the copper wire around a pen to make a small “cage” into which the magnesium ribbon may be inserted (see Figure 2–NOTE: The cage must be able to fit into the gas measuring tube!) b. place the magnesium in the “cage” so it will stay c. insert the straight end of the copper wire into a one-hole rubber stopper so that the cage end of the copper is about 7 cm below the bottom of the stopper (see Figure 1) d. hook the end of the copper wire around the top of the stopper to hold the “cage” in place 5. Measure the hydrochloric acid a. obtain about 30 mL of 2 M hydrochloric acid in the 50 mL beaker b. measure out 15 mL of the HCl in the 25 mL graduated cylinder 6. Add the HCl slowly to the gas measuring tube 7. Carefully (so as not to mix the acid and water) fill the gas measuring tube completely (so it overflows!) with distilled water. (Dislodge any bubbles by tapping them!) 8. Insert the magnesium-copper wire-stopper assembly into the gas measuring tube
9. Place your finger over the hole of the rubber stopper 10. Invert the gas measuring tube 11. Lower the stoppered end of the tube into the 600 mL beaker containing the tap water. 12. Clamp the tube in place with a test tube clamp and ringstand (as in Figure 3). 13. Allow the reaction to take place completely (until all of the magnesium has reacted). 14. Fill a 1000 mL graduated cylinder with tap water using the second 600 mL beaker. 15. Cover the hole in the stopper with your finger 16. Transfer the gas measuring tube to the 1000 mL cylinder (see Figure 4). 17. Gently move the gas measuring tube up and down in the cylinder until the water level in the tube is the same as the water level in the cylinder (to equalize the pressure with the surrounding air) 18. Record the exact volume of hydrogen gas from the gas measuring tube. 19. Record the temperature of the water in the 1000 mL cylinder and the barometric pressure in the room. 20. Pour the contents of the gas measuring tube into the labeled waste container. 21. Repeat steps 3-20 for a second trial (but use the same cage!). 22. Clean up!
Post-lab calculations: 1) Calculate the theoretical number of moles of hydrogen gas in both trials using the reaction equation from your research. 2) Use a table listing the vapor pressure of water to calculate the partial pressure of hydrogen gas from both your trials. (NOTE: Vapor pressure table on back of this lab handout or search for “Cabrillo College Vapor Pressure” online.) 3) Use the combined gas law (or ideal gas law) to convert the measured volume of hydrogen to the “ideal” (or “dry”) volume the hydrogen gas would occupy at STP for both trials. 4) Divide the “ideal” volume of hydrogen at STP by the theoretical number of moles of hydrogen to calculate the molar volume of hydrogen at STP for both trials. 5) Average your trials. 6) Calculate your percent error. 7) Calculate the density of a mole (in grams per liter) of hydrogen using your average experimental molar volume and the molar mass of hydrogen gas. 8) What effect would leaking some air into the gas measuring tube while inverting have on the measured volume of hydrogen gas? Be specific! Explain your thinking! 9) What effect would using oxidized magnesium ribbon (which looks black or tarnished) have on the measured volume of the hydrogen gas? Be specific! Explain your thinking!
Lab handout based on the experiment “Determining the Molar Volume of a Gas” in Laboratory Experiments for Advanced Placement Chemistry (Second Edition) by S.A. Vonderbrink (Flinn Scientific, 2006)
