Materials Included In Kit

Additional Materials Required

Prelab Preparation

Magnesium Ribbon “Density” Determination
To determine the g of Mg/cm:

1. Measure and cut 1.00 m of the magnesium ribbon.
2. Mass the 1.00-m segment on a balance precise to ±0.001 g.
3. Calculate the g/cm for the magnesium ribbon. Note: If the balance has a precision to only 0.01 g, use a 1.50-m segment of magnesium ribbon to obtain three significant figures for g/cm for the ribbon.

Safety Precautions

Hydrochloric acid is a corrosive liquid. Avoid contact with eyes and skin and clean up all spills immediately. Magnesium metal is a flammable solid. Keep away from flames and other sources of ignition. Wear chemical splash goggles and chemical-resistant gloves and apron. Wash hands thoroughly with soap and water before leaving the laboratory. Please consult current Safety Data Sheets for additional safety, handling and disposal information.

Disposal

Please consult your current Flinn Scientific Catalog/Reference Manual for general guidelines and specific procedures, and review all federal, state and local regulations that may apply, before proceeding. The water bath solutions remaining after the hydrogen gas has been collected will be slightly acidic. These should be neutralized with base (sodium bicarbonate is a good choice) and flushed down the drain with excess water according to Flinn Suggested Disposal Method #24b.

Lab Hints

• The laboratory work for this experiment can reasonably be completed within a normal 50- minute class period. The PreLab Questions may be assigned as homework in preparation for lab or may be used as part of a cooperative classroom discussion. Doing the prelab will help students understand the calculations needed to complete their lab report.
• The amount of hydrogen generated depends on the length of magnesium ribbon used, its linear density (mass in grams per centimeter) and the purity (freshness) of the metal. The linear density, in turn, depends on the thickness of the metal ribbon and will typically vary between 0.0075 g/cm and 0.014 g/cm. For best results, use fresh magnesium ribbon or buff the ribbon with sandpaper before weighing it.
• Calculate the length of magnesium ribbon that would produce between 42 mL and 46 mL of hydrogen. This is the length to be used by the students. The 3.3-cm length of ribbon listed in the procedure is from a ribbon with a linear density of 0.0118 g/cm.
• The concentration and volume of hydrochloric acid should be controlled variables. Optimum results were obtained using 15 mL of 2 M HCl, as recommended in the procedure. Using a smaller volume of more concentrated 3 M acid gave a more rapid reaction but the results were less reproducible.
• If a barometer is not available in the lab or classroom, students may consult an Internet site such as the national weather service site (http://weather.gov/) to obtain a current “sea-level” pressure reading for your area. Note that these are NOT actual barometric pressure readings. Meteorologists convert station pressure values to what they would be if they had been taken at sea level. The following equation can be used to recalculate the barometric pressure (in inches Hg) from the reported sea-level pressure (in inches Hg). Elevation must be in meters.

  barometric pressure = sea-level pressure – (elevation/312)

• The hydrogen gas may be collected in an ice-water bath at 0–5 °C rather than a room temperature water bath. This will reduce the size of the water-vapor contribution to the pressure and will also decrease the magnitude of the temperature correction in converting the measured volume to STP. If you live at high altitudes, or if the weather is stormy due to a low-pressure air mass, the experiment should definitely be carried out at 0–5 °C! Otherwise, the volume of hydrogen may exceed the 50-mL capacity of the eudiometer tube.
• If eudiometer tubes are not available, it is possible to use inverted 50-mL burets. The space between the 50-mL mark and the stopcock valve would have to be calibrated. Since the numbers invert when the buret inverts, students may have difficulties in taking the volume readings.

Teacher Tips

• Ask your students to come up with a list of interesting questions regarding the amount of gas in familiar, everyday objects. The questions (and answers) can be assigned as extra credit to reinforce gas law calculations or to prepare for tests. How much helium gas is needed to fill the Goodyear blimp? How much gas is present in an aerosol container pressurized at 2.2 atm? How many air molecules does it take to blow up a party balloon? What volume of butane gas would be generated if all the liquid butane in a barbecue lighter were converted to gas at STP?
• For an alternative microscale version of this experiment, use a 10-mL inverted graduated cylinder with a 1.0-cm sample of magnesium. See the Flinn ChemTopic™ Labs, Vol. 9, The Gas Laws for the complete procedure.

Further Extensions

AP^® Chemistry Standards
This lab fulfills the requirements for the College Board recommended AP Experiment #5: Determination of the Molar Volume of a Gas.

Answers to Prelab Questions

A reaction of 0.028 g of magnesium with excess hydrochloric acid generated 31.0 mL of hydrogen 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.

1. Use Dalton’s law and the vapor pressure of water at 22 °C (Table 1 in the Background section) to calculate the partial pressure of hydrogen gas in the gas collecting tube.

   Vapor pressure of water at 22 °C = 19.8 mm Hg Ptotal = 746 mm = PH2 + 19.8 mm Hg PH2 = 746 – 19.8 mm = 726 mm Hg

2. Use the combined gas law to calculate the “corrected” volume of hydrogen at STP. Hint: Watch your units for temperature and pressure!

   STP = 273 K (0 °C) and 1 atm

   {13899_PreLabAnswers_Equation_3}
3. What is the theoretical number of moles of hydrogen that can be produced from 0.028 g of Mg? Hint: Refer to Equation 1 for the balanced equation for the reaction.

   Theoretical number of moles hydrogen = number of moles of magnesium used

   {13899_PreLabAnswers_Equation_4}
4. Divide the corrected volume of hydrogen by the theoretical number of moles of hydrogen to calculate the molar volume (in L/mole) of hydrogen at STP.

   Molar volume = 0.0274 L/0.0012 moles = 23 L/mole

Sample Data

* The results of three trials are shown here to illustrate the reproducibility of the method. Only two trials are required in the Procedure. The “density” of the Mg ribbon was 0.0118 g/cm.

Results Table

Answers to Questions

1. Calculate the theoretical number of moles of hydrogen gas produced in Trials 1 and 2.

   Theoretical number of moles of H₂ = number of moles of Mg For Trial 1:

   {13899_Answers_Equation_5}
2. Use Table 1 in the Background section to find the vapor pressure of water at the temperature of the water bath in this experiment. Calculate the partial pressure of hydrogen gas produced in Trials 1 and 2.

   For Trial 1: Vapor pressure of water at 27 °C = 26.7 mm Hg
   P_total = 738 mm Hg = P_H2 + 26.7 mm Hg
   P_H2 = 738 mm Hg – 26.7 mm Hg = 711 mm Hg (0.936 atm)

3. Use the combined gas law to convert the measured volume of hydrogen to the volume the gas would occupy at STP for Trials 1 and 2. Hint: Remember the units!
   {13899_Answers_Equation_6}
4. Divide the volume of hydrogen gas at STP by the theoretical number of moles of hydrogen to calculate the molar volume of hydrogen for Trials 1 and 2.

   For Trial 1: Molar volume = 0.0389 L/0.0016 moles = 22 L/mole

5. What is the average value of the molar volume of hydrogen? Look up the literature value of the molar volume of a gas at STP in your textbook and calculate the percent error in your experimental determination of the molar volume of hydrogen.
   {13899_Answers_Equation_7}
6. One mole of hydrogen gas has a mass of 2.02 g. Use your value of the molar volume of hydrogen to calculate the mass of one liter of hydrogen gas at STP. This is the density of hydrogen in g/L. How does this experimental value of the density compare with the literature value? (Consult a chemistry handbook for the density of hydrogen gas at STP.)
   {13899_Answers_Equation_8}

   The literature value for the density of hydrogen gas is 0.0899 g/L at STP.

7. In setting up this experiment, a student noticed that a bubble of air leaked into the eudiometer tube when it was inverted in the water bath. What effect would this have on the measured volume of hydrogen gas? Would the calculated molar volume of hydrogen be too high or too low as a result of this error? Explain.

   The bubble of air that leaked in would cause the measured volume of hydrogen to be too high. Both the calculated volume of hydrogen at STP and the molar volume would be too high as a result of this error.

8. A student noticed that the magnesium ribbon appeared to be oxidized—the metal surface was black and dull rather than silver and shiny. What effect would this error have on the measured volume of hydrogen gas? Would the calculated molar volume of hydrogen be too high or too low as a result of this error? Explain.

   Oxidation of magnesium converts the metal to magnesium oxide (MgO). The magnesium oxide coating would still react with hydrochloric acid, but would not generate hydrogen gas. The combined reaction of Mg and MgO with HCl would generate less hydrogen than the reaction of Mg alone. Both the calculated volume of hydrogen at STP and the molar volume would be too low as a result of this error. Note: The black coating may be due to Mg₃N₂ or MgS.

Introduction

From blimps to airbags, gases are used to fill a wide variety of containers. How much of a particular gas must be produced to fill each of these containers? The amount of gas needed to fill any size container can be calculated if the molar volume of the gas is known.

Concepts

• Avogadro’s law
• Dalton’s law
• Ideal gas law
• Molar volume

Background

Avogadro’s law states that equal volumes of gases contain equal numbers of molecules under the same conditions of temperature and pressure. It follows, therefore, that all gas samples containing the same number of molecules 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 standard temperature and pressure (STP, equal to 273 K and 1 atm) will be measured.

The reaction of magnesium metal with hydrochloric acid (Equation 1) provides a convenient means of generating hydrogen in the lab.

If the reaction is carried out with excess hydrochloric acid, the volume of hydrogen gas obtained will depend on the number of moles of magnesium as well as the pressure and temperature. The molar volume of hydrogen can be calculated if the volume occupied by a 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 must be corrected to STP conditions before calculating the molar volume.

The relationship among the four gas variables—pressure (P), volume (V), temperature (T) and the number of moles (n)—is expressed in the ideal gas law (Equation 2), where R is a constant called the universal gas constant. The ideal gas law reduces to Equation 3, the combined gas law, if the number of moles of gas is constant. The combined gas law can be used to calculate the volume (V₂) of a gas at STP (T₂ and P₂) from the volume (V₁) measured under any other set of laboratory conditions (T₁ and P₁). In using either the ideal gas law or the combined gas law, remember that temperature must be always be expressed in units of kelvins (K) on the absolute temperature scale. Hydrogen gas will be collected by the displacement of water in an inverted gas measuring tube (also call a eudiometer tube) using the apparatus shown in Figure 1. The total pressure of the gas in the tube will be equal to the barometric (air) 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, the total pressure of the gas will be equal to the partial pressure of hydrogen plus the partial pressure of water vapor (Equation 4). The vapor pressure of water depends only on the temperature (see Table 1).

Experiment Overview

The purpose of this experiment is to determine the volume of one mole of hydrogen gas at standard temperature and pressure (STP). Hydrogen will be generated by the reaction of a known mass of magnesium with excess hydrochloric acid in an inverted gas measuring tube filled with water. The volume of hydrogen collected by water displacement will be measured and corrected for differences in temperature and pressure in order to calculate the molar volume of hydrogen at STP.

Materials

Prelab Questions

A reaction of 0.028 g of magnesium with excess hydrochloric acid generated 31.0 mL of hydrogen 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.

1. Use Dalton’s law and the vapor pressure of water at 22 °C (Table 1) to calculate the partial pressure of hydrogen gas in the gas collecting tube.
2. Use the combined gas law to calculate the “corrected” volume of hydrogen at STP. Hint: Watch your units for temperature and pressure!
3. What is the theoretical number of moles of hydrogen that can be produced from 0.028 g of Mg? Hint: Refer to Equation 1 for the balanced equation for the reaction.
4. Divide the corrected volume of hydrogen by the theoretical number of moles of hydrogen to calculate the molar volume (in L/mol) of hydrogen at STP.

Safety Precautions

Hydrochloric acid is a corrosive liquid. Avoid contact with eyes and skin and clean up all spills immediately. Magnesium metal is a flammable solid. Keep away from flames and other sources of ignition. Wear chemical splash goggles and chemical-resistant gloves and apron. Wash hands thoroughly with soap and water before leaving the laboratory.

Procedure

1. Fill a 400-mL beaker about 3⁄4-full with water.
2. Obtain or cut a 3.3-cm piece of magnesium ribbon. Measure and record in the data table the exact length of the magnesium ribbon to the nearest 0.1 cm.
3. Your teacher will provide a conversion factor in g/cm to calculate the mass of magnesium used in this experiment. Multiply the length of magnesium ribbon by this conversion factor to calculate the mass of the 3.3-cm piece of magnesium obtained in step 2. Record the mass of magnesium in the data table.
4. Obtain a piece of copper wire about 15-cm long. Twist and fold one end of the copper wire around a pencil to make a small “cage” into which the magnesium ribbon may be inserted (see Figure 2).
   {13899_Procedure_Figure_2}
5. Firmly place the 3.3-cm piece of magnesium into the copper-wire cage. For larger pieces of magnesium, bend or fold the magnesium ribbon first.
6. Insert the straight end of the copper wire into a one-hole rubber stopper so that the cage end containing the magnesium is about 7-cm below the bottom of the stopper (see Figure 1 in the Background section). Hook the end of the copper wire around the top of the stopper to hold the cage in place.
7. Obtain about 15 mL of 2 M hydrochloric acid in a 25-mL graduated cylinder.
8. While holding the eudiometer tube in a tipped position, slowly add the 2 M hydrochloric acid to the tube.
9. While still holding the eudiometer tube in the tipped position, slowly and carefully fill the tube with distilled or deionized water. Work slowly to avoid mixing the acid and the water layers. Fill the tube all the way to the top so no air remains in the tube.
10. Insert the magnesium–copper wire–stopper assembly into the eudiometer tube. The magnesium piece should be above the 50-mL line on the eudiometer tube (see Figure 3).
    {13899_Procedure_Figure_3}
11. Place your finger over the hole of the rubber stopper, invert the eudiometer tube, and carefully lower the stoppered end of the tube into the 400-mL beaker containing water. The tube should contain no air bubbles. Clamp the tube in place (see Figure 3).
12. Record any evidence of a chemical reaction in the data table.
13. If the magnesium metal “escapes” its copper cage, gently shake the eudiometer tube up and down to work it back into the acidic solution. Note: Do not lift the tube completely out of the water in the beaker.
14. Allow the apparatus to stand for 5 minutes after the magnesium has completely reacted. Gently tap the sides of the eudiometer tube to dislodge any gas bubbles that may have become attached to the sides.
15. Fill a 500-mL graduated cylinder with tap water and place the cylinder in the sink.
16. Cover the hole in the stopper with your finger and transfer the eudiometer tube to the 500-mL graduated cylinder (see Figure 4).
    {13899_Procedure_Figure_4}
17. Gently move the eudiometer tube up and down in the cylinder until the water level inside the tube is the same as the water level in the graduated cylinder. This is done to equalize the pressure with the surrounding air (barometric pressure). Note: Make sure the stoppered end of the eudiometer tube remains submerged in the water.
18. When the water levels inside and outside the tube are the same, measure and record in the data table the exact volume of hydrogen gas in the eudiometer tube.
19. Measure and record in the data table the temperature of the water bath in the graduated cylinder. Using a barometer, measure and record the barometric pressure in the lab in the data table.
20. Remove the eudiometer tube from the graduated cylinder and discard the water as directed by your instructor.
21. Repeat the entire procedure to obtain a second set of data. Record this as Trial 2 in the data table. If time permits, perform a third trial as well.

Student Worksheet PDF

13899_Student1.pdf