Materials Included In Kit

Additional Materials Required

Safety Precautions

Hydrochloric acid is a corrosive liquid and may cause skin burns and eye damage. Avoid contact with eyes and skin and clean up all spills immediately. Keep acid neutralizer (soda ash) on hand to clean up acid spills. Magnesium metal is a flammable solid. Keep away from flames and other sources of ignition. Wear chemical splash goggles, chemical-resistant gloves and a lab coat or chemical-resistant apron. Remind students to wash their 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 may be acidic. The water may be neutralized with base (sodium bicarbonate is a good choice) before rinsing 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 typical 2-hour lab period. The reaction of magnesium with hydrochloric acid on the microscale level recommended in this experiment takes about five minutes to proceed to completion. This allows students the opportunity to work through the procedure and conduct two or more runs in a typical lab period. The Prelaboratory Assignment may be used as part of a cooperative classroom discussion to review the calculations needed to complete the 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 0.015 g/cm and 0.020 g/cm. If the linear density of magnesium is greater than 0.020 g/cm, the volume of hydrogen generated in this reaction will exceed the 25-mL capacity of the graduated cylinder. Check to make sure that a 1-cm length of magnesium ribbon recommended in this experiment will not generate more than 24 mL of hydrogen gas. If necessary, cut back on the length of magnesium. For best results, use fresh magnesium ribbon or buff it with steel wool before use.
• The concentration and volume of hydrochloric acid should be controlled variables. Optimum results were obtained using 10 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.
• Not all graduated cylinders are created equal. The best results were obtained using glass, tall-form, 25-mL graduated cylinders marked with 0.2-mL minor increments. Economy-choice graduated cylinders are shorter and may be marked only in 0.5-mL minor increments—these are not recommended.
• If a milligram balance is not available, accurately measure out 10 cm of magnesium ribbon. Record the precise length to the nearest 0.1 cm, as well as the mass to two decimal places. Calculate the linear density (g/cm) and give this value to the students as a conversion factor to determine the mass of their 1-cm piece of magnesium ribbon.
• If a barometer is not available in the lab or classroom, students may consult an Internet site, such as the National Weather Service (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 a room temperature bath rather than in ice water. This will increase the size of the water-vapor contribution to the pressure and will also increase 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 will exceed the capacity of the graduated cylinder.
• This classic experiment can be transitioned to guided inquiry using the following approaches, which will improve student preparation and increase the level of student engagement and ownership of results.
  • Make it a challenge! Rather than having students use a prescribed amount of magnesium to generate an unknown volume of gas, challenge them to produce a specific volume of hydrogen gas. Students must work backwards using the gas laws to calculate how much magnesium metal to use.
  • Introduce the lab by demonstrating the general setup for generating and collecting a gas. Guide students to design the actual experimental procedure through a series of leading questions. What information (data) is needed to calculate the molar volume of a gas? What variables will influence the experimental data? Choose the independent and dependent variables for the experiment and describe the variables that should be kept constant during the experiment. Students could also discuss variables or other factors that will affect the accuracy of the results and how these may be controlled.
  • Extend the lab to incorporate consumer products, such as analyzing butane gas collected from a lighter.
• Ask 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 a 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?

Answers to Prelab Questions

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 at 22 °C. The barometric pressure in the lab was 746 mm Hg.

1. The vapor pressure of water at 22 °C is 19.8 mm Hg. Use Dalton’s law to calculate the partial pressure of hydrogen gas in the gas-collecting tube.

   P_total = P_H2 + P_H2O
   P_total = 746 mm = P_H2 + 19.8 mm Hg
   P_H2 = 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

   {14033_PreLabAnswers_Equation_5}
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 of hydrogen = number of moles of magnesium used

   {14033_PreLabAnswers_Equation_6}
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 mole = 23 L/mole

Sample Data

Laboratory Report

Results Table

*Rounded to two significant figures.

Answers to Questions

Enter the results of all calculations in the results table.

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

   {14033_Answers_Equation_7}
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 5 °C = 6.5 mm
   P_total = 738 mm Hg = P_H2 + 6.5 mm Hg
   P_H2 = 738 mm Hg – 6.5 mm Hg = 731.5 mm Hg (0.963 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!
   {14033_Answers_Equation_8}
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 = 0.0174 L/0.000782 mole = 22.3 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 and calculate the percent error in your experimental determination of the molar volume of hydrogen.
   {14033_Answers_Equation_9}
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.)
   {14033_Answers_Equation_10}
   The literature value for the density of hydrogen gas is 0.0899 g/L at STP; 2% error
7. In setting up this experiment, a student noticed that a bubble of air leaked into the graduated cylinder 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.

9. (Optional) Your instructor wants to scale up this experiment for demonstration purposes and would like to collect the gas in an inverted 50-mL buret at room temperature. Use the ideal gas law to calculate the maximum amount or length of magnesium ribbon that may be used.

   Assume that the maximum volume of hydrogen gas we can measure is 50 mL at a maximum “room” temperature of 30 °C (303 K) and a minimum atmospheric pressure of 0.95 atm. Substitute these values into the ideal gas equation to solve for n, the maximum number of moles of magnesium that can be used in the experiment. n = PV/RT

   {14033_Answers_Equation_11}

   Mass of magnesium = 0.0019 mole x 24.3 g/mole = 0.046 g (about 2.4 cm)
   Note: Assuming a higher than average temperature and a lower than normal atmospheric pressure gives us a maximum volume of hydrogen gas. This leaves some “wiggle room” for experimental error so that the actual volume will not exceed the capacity of the 50-mL buret.

Introduction

Avogadro, Boyle, Charles and Dalton—these scientists and their gas laws are well known. Together their work defined the relationships among four macroscopic gas properties: pressure, volume, temperature and the number of moles of gas. The gas laws have applications in physiology, meteorology, scuba-diving and even hot-air ballooning or airbag construction. How much gas must be generated to fill a hot air balloon or an airbag? The amount of gas needed to fill any container can be calculated if we know the molar volume of the gas. Answering this general question requires knowledge of all of the gas laws (ABCs + D)!

Concepts

• Avogadro’s law
• Dalton’s law
• Ideal gas law
• Molar volume
• Vapor pressure
• Water displacement

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 we will determine the molar volume of hydrogen gas at standard temperature and pressure (STP, equal to 273 K and 1 atm).

The reaction of magnesium metal with hydrochloric acid (Equation 1) provides a convenient means of generating small-scale quantities of 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 on the pressure and temperature. The molar volume of hydrogen can be calculated if we measure the volume occupied by a sample containing a known number of moles of hydrogen. 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 expressed in units of kelvins (K) on the absolute temperature scale. Hydrogen gas will be collected by the displacement of water in an inverted graduated cylinder using the apparatus shown in Figure 1. The total pressure of the gas in the cylinder 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 it 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 on 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 graduated cylinder 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

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 at 22 °C. The barometric pressure in the lab was 746 mm Hg.

1. The vapor pressure of water at 22 °C is 19.8 mm Hg. Use Dalton’s law 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/mole) of hydrogen at STP.

Safety Precautions

Hydrochloric acid is a corrosive liquid; it will cause skin burns and eye damage. 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, chemical-resistant gloves and a lab coat or chemical-resistant apron. Wash hands thoroughly with soap and water before leaving the laboratory.

Procedure

1. Fill a 600-mL beaker about ¾ full with ice water. The temperature should be between 0 and 5 °C.
2. Obtain or cut a 1-cm piece of magnesium ribbon. Measure and record the exact length of the magnesium ribbon to the nearest 0.1 cm. Do not exceed 1.2 cm.
3. Measure the mass of magnesium to the nearest milligram (0.001 g) using an electronic balance. Record the value in the data table.
4. Obtain a piece of copper wire about 12-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).
   {14033_Procedure_Figure_2}
5. Firmly place the 1-cm piece of magnesium into the copper-wire cage.
6. Insert the straight end of the copper wire into a one-hole rubber stopper so that the cage containing the magnesium is about 2-cm below the bottom of the stopper (see Figure 1 in the Background). Hook the end of the copper wire around the top of the stopper to hold the cage in place.
7. Obtain about 10 mL of 2 M hydrochloric acid in a 25-mL graduated cylinder.
8. While holding the graduated cylinder in a tilted position, slowly and carefully fill the graduated cylinder with distilled water from a wash bottle or a plastic pipet. Work slowly to avoid mixing the acid and water layers at this time. Fill the graduated cylinder all the way to the top so that no air remains in the cylinder.
9. Insert the magnesium–copper wire–stopper assembly into the graduated cylinder. The magnesium piece should be above the 25-mL line on the graduated cylinder (see Figure 1). Some water will spill over!
10. Place your finger over the hole of the rubber stopper, invert the graduated cylinder and carefully lower the stoppered end of the graduated cylinder into the 600-mL beaker containing ice water.
11. Record any evidence of a chemical reaction in the data table.
12. If the magnesium metal “escapes” its copper cage, gently shake the graduated cylinder up and down to work the metal back into the acidic solution.
13. Allow the apparatus to stand for 5 minutes after the magnesium has completely reacted. Gently tap the sides of the graduated cylinder to dislodge any gas bubbles that may have become attached to the sides.
14. Gently move the graduated cylinder up or down in the ice-water bath until the water level inside the graduated cylinder is the same as the water level in the beaker. This is done to equalize the pressure with the surrounding air (barometric pressure). Note: Be careful to make sure that the stoppered end of the graduated cylinder remains submerged in the water.
15. When the water levels inside and outside the cylinder are the same, measure and record the precise volume of hydrogen gas in the graduated cylinder.
16. Since the cylinder is upside down in the water bath, the volume reading obtained in step 15 must be corrected for the fact that the meniscus was read upside down as well. Subtract 0.2 mL from the volume recorded in step 15 and record the “corrected” volume of hydrogen gas in the data table.
17. Measure and record the temperature of the ice water in the beaker. Using a barometer, measure and record the barometric pressure in the lab.
18. Remove the graduated cylinder from the ice water and discard the contents as directed by your instructor.
19. Repeat the entire procedure to obtain a second set of data. Record this as Trial 2 in the data table.

Student Worksheet PDF

14033_Student1.pdf