Materials Included In Kit

Additional Materials Required

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 acidic. These should be neutralized by adding base (sodium bicarbonate is a good choice) and then 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 typical 50-minute class period. The reaction of magnesium with hydrochloric acid on the microscale level recommended in the procedure takes about five minutes to proceed to completion. This allows students the opportunity to conduct two or more trials in a typical lab period. The Prelab Questions may be assigned as homework in preparation for lab. Doing the prelab assignment will help students understand the calculations needed to complete their lab report.
• The amount of hydrogen gas 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 depends on the thickness of the metal ribbon and will typically vary between 0.0085 g/cm and 0.010 g/cm. If the linear density of magnesium is greater than 0.010‑g/cm, the volume of hydrogen generated in this reaction will exceed the 10-mL capacity of the graduated cylinder. Check to make sure that the 1-cm length of magnesium ribbon recommended in this experiment will not generate more than 10‑mL of hydrogen gas. If necessary, cut back on the length of magnesium. For best results, buff the magnesium with steel wool before weighing and cutting the metal for student use. The teacher must give the students the mass/length conversion factor (see step 3 in the Procedure)—pre-measure the mass and length of the 45-cm magnesium ribbon supplied with this kit.
• Optimum results are obtained using 5 mL of 2 M HCl, as recommended in the procedure. Using a smaller volume of more concentrated acid gives a more rapid reaction, but the results are less reproducible.
• Not all graduated cylinders are created equal. We have found that the best results are obtained using glass, tall-form, 10-mL graduated cylinders marked with 0.1-mL minor increments. Economy-choice graduated cylinders are shorter and are marked only in 0.2-mL minor increments—these are not recommended.
• 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 in 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 altitude, 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 10-mL capacity of the graduated cylinder.

Teacher Tips

• For an alternative lab activity dealing with molar volume and applications of the ideal gas law, see the demonstration “Molar Mass of Butane” in The Gas Laws, Volume 9 in the Flinn ChemTopic™ Labs series. This demonstration uses disposable butane lighters to generate a measured mass of butane gas. The molar mass of butane is calculated based on the theoretical molar volume of an ideal gas.
• See Construction of Gas Volume Cubes in the Demonstrations section of The Gas Laws for a follow-up assessment activity. Different groups of students are assigned different numbers of moles of gas and are asked to construct a cube to represent the volume of the given number of moles of gas at STP. The examples have been chosen so that all of the cubes will have dimensions to the nearest centimeter (e.g., 0.0446 moles of gas will have a volume of 1.00 L at STP, corresponding to a cube 10.0 cm in length on each side).
• 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?

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 in a water bath at 22 °C. 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.

   Vapor pressure of water at 22 °C = 19.8 mm Hg
   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

   {13961_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 hydrogen = Number of moles of magnesium used

   {13961_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 moles = 23 L/mole

Sample Data

*The conversion factor for the magnesium used in this experiment was 0.0094 g/cm.
†The results of three trials are shown here to illustrate the reproducibility of the method. Only two trials are required.

Results Table

Answers to Questions

See the results table in Sample Data for the results of the calculations (Questions 1–5).

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

   {13961_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 23 °C = 21.1 mm
   P_total = 764 mm Hg = P_H2 + 21 mm Hg
   P_H2 = 764 mm Hg – 21 mm Hg = 743 mm Hg

   {13961_Answers_Equation_8}
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!
   {13961_Answers_Equation_9}
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 = 0.00839 L/0.00039 moles = 22 L/mole

5. What is the average value of the molar volume of hydrogen gas? Look up the literature value of the molar volume of a gas in a textbook and calculate the percent error in the experimental determination of the molar volume of hydrogen gas.
   {13961_Answers_Equation_10}
6. 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 gas 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 gas 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.

7. 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 gas 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 gas 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 also be due to Mg₃N₂ or MgS.

8. (Optional) Your teacher wants to scale up this experiment for demonstration purposes and would like to collect the gas in an inverted 50-mL buret. Use the ideal gas law to calculate the maximum length of magnesium ribbon that your teacher should use.

   Assume that the maximum volume of hydrogen gas which 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

   {13961_Answers_Equation_11}

   Mass of magnesium = 0.0019 moles x 24.3 g/mole = 0.046 g Length of magnesium = 0.046 g/0.0094 g/cm = 4.9 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.

References

This activity was adapted from Flinn ChemTopic™ Labs, Volume 9, The Gas Laws; Cesa, I., Editor; Flinn Scientific: Batavia IL, 2003.

Introduction

Airbags have been required safety features on new cars since the 1980s and are credited with saving thousands of lives over that time. Airbags contain a compound that decomposes to release nitrogen gas upon impact from a collision. How much gas must be generated to fill an airbag? The amount of gas needed to fill any size container can be calculated if we know the molar volume of the gas.

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 will be determined 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 gas 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 gas. 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. In this activity, 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 is 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).

Table 1. Vapor Pressure of Water at Different Temperatures

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 in a water bath at 22 °C. 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 in the Background section.
4. Divide the corrected volume of hydrogen (Question 2) by the theoretical number of moles of hydrogen (Question 3) to calculate the molar volume (in L/mole) 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 0.9–1.0-cm piece of magnesium ribbon. Measure and record the exact length of the magnesium ribbon to the nearest 0.05 cm. Do not exceed 1.0 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 1-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 10-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).
   {13961_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 end containing the magnesium is about 1-cm below the bottom of the stopper (see Figure 1). Hook the end of the copper wire around the top of the stopper to hold the cage in place.
7. Obtain about 5 mL of 2 M hydrochloric acid in a tall-form, 10-mL graduated cylinder.
8. While holding the graduated cylinder at an angle, 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. Completely fill the graduated cylinder with distilled water 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 10-mL line on the graduated cylinder (see Figure 1 in the Background section). Excess water will spill over the top of the graduated cylinder.
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 400-mL beaker containing 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 it 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 attached to the sides.
14. Gently move the graduated cylinder up or down in the 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 exact 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 is being 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 water bath in the beaker.
18. Using a barometer, measure and record the barometric pressure in the lab.
19. Remove the graduated cylinder from the water bath and discard the water in the beaker as directed by your instructor.
20. Repeat the procedure to obtain a second set of data (the copper cage may be reused). Record this as Trial 2 in the data table.

Student Worksheet PDF

13961_Student1.pdf