Chemistry Replacement of hydrogen by a metal Objectives In this experiment you will: 1. Collect the hydrogen gas formed by a known mass of magnesium reacting with hydrochloric acid. 2. Calculate the actual moles of the hydrogen gas from the volume of hydrogen collected. 3. Calculate the theoretical moles of hydrogen formed by a known mass of magnesium reacting with HCl acid. 4. Calculate the percentage error.
Equipment and materials magnesium ribbon, concentrated hydrochloric acid ring stand, test tube clamp, eudiometer, large glass cylinder, string, electronic balance, thermometer
Safety 1. Wear safety goggles and an apron. 2. Hydrochloric acid is corrosive and must be handled with care. If acid is spilled on your skin, wash it off immediately with water. If acid should get into your eyes, wash your eyes for 20 minutes at the eye wash fountain. Spilled acid should be wiped up with wet paper towels.
Procedure (Do twice) 1. Fill the large cylinder nearly full of tap water at room temperature.
2. Clean the surface of a magnesium ribbon and measure a mass that is 0.045 (3.0 cm) grams or slightly less if you are using a 50 mL eudiometer or that is 0.090 grams or slightly less (6.0 cm) if you are using a 100 mL eudiometer. Record the mass of the magnesium to the nearest 0.001 gram in the data table.
3. Roll the magnesium ribbon into a loose coil that will fit into the eudiometer and tie through a 2-hole stopper so the magnesium hangs about 5-6 cm. Do not place the magnesium into the eudiometer yet.
4. Pour 10-12 ml of concentrated HCl into the eudiometer. Gently add enough distilled water to fill the eudiometer completely; pouring slowly should keep the denser HCl at the bottom of the eudiometer.
5. Lower the magnesium coil just into the eudiometer opening. Put a stopper in the end of the eudiometer, leaving the free end of the thread outside of the eudiometer. You should not have any air bubbles in the eudiometer.
6. Covering the end of the stopper with a finger/thumb, quickly place the inverted eudiometer in the cylinder. The magnesium ribbon should be about an inch into the bottom of the eudiometer. Rest the eudiometer opening on the bottom of the beaker and hold the eudiometer vertically until the reaction is complete.
7. When the reaction has stopped, adjust the eudiometer until the liquid levels inside and outside are the same. If the gas has overflowed the eudiometer you used too much magnesium and will need to repeat the experiment.
8. Measure the volume of the hydrogen liberated and record in the data table.
9. Measure the temperature of the water in the beaker and record in the data table.
10. Obtain the barometric pressure reading in inches from the web.
Data Table Trial 1 Trial 2
1. Mass of Mg used ______g ______g
2. Volume of H2 collected ______mL ______mL
3. Temperature of H2 collected ______C ______C
4. Barometric pressure ______inches Hg ______inches Hg
Calculations (attach paper with your work, label trial 1 or 2, calculation # and show units)
1. Use the temperature of the beaker water as the temperature of H2 collected and convert to Kelvin.
2. Convert the atmospheric pressure to mm Hg
3. Write the partial pressure of water vapor at the H2 temperature from the partial pressure of water vapor table. Convert the water pressure to mm Hg.
4. Calculate the pressure of the dry hydrogen gas.
5. Calculate the percentage of H2 molecules in the collected gas using partial pressures.
6. Convert the volume of dry H2 gas collected into liters
7. Calculate the actual moles of dry H2 collected using the ideal gas law and the dry hydrogen pressure and volume.
8. Write a balanced equation for the reaction of magnesium with hydrochloric acid.
9. Calculate the theoretical number of moles of hydrogen formed from the starting mass of magnesium.
10. Calculate percentage error from the actual and theoretical moles of hydrogen gas
Questions Answer the following questions on your own paper
1. Identify an advantage of collecting gases over water instead of collecting them in a balloon.
2. Why subtract the water vapor pressure from the total pressure before calculating moles of hydrogen?
2. Explain why the ideal gas law constant must be used in calculating the actual moles?
3. Identify three possible sources of experimental error