Materials Included In Kit

Additional Materials Required

Prelab Preparation

Preparation and Use of Special Equipment

Plastic Syringe

1. Pull out the syringe plunger until it is at about the 50-mL mark (see Figure 3).
   {14001_Preparation_Figure_3}
2. Heat the tip of the nail in a burner and quickly force the nail through the plunger stem. (The nail easily melts through the plastic.)

Gas Delivery Apparatus

1. Assemble the gas delivery apparatus by first pushing the 1-quart plastic bag through the hole in the #10 stopper, leaving approximately 1 inch of the bag out of the opening (see Figure 4).
   {14001_Preparation_Figure_4}
2. Place the #2 one-hole stopper firmly into the freezer bag opening. The freezer bag should be held tightly between the walls of the two stoppers (see Figure 5).
   {14001_Preparation_Figure_5}
3. Carefully insert the tapered end of the medicine dropper through the hole of the #2 stopper. Place a drop or two of glycerin in the hole and slowly work the dropper back and forth until the tip is inside the bag (see Figure 6).
   {14001_Preparation_Figure_6}
4. Attach the short piece of latex tubing over the wide end of the medicine dropper.
5. Place the pinch clamp on the latex tubing. The apparatus is now complete.
6. Assemble a new gas delivery apparatus for each additional gas to be measured.

Filling with a Gas

1. Evacuate the bag: Remove the pinch clamp and attach the latex tubing to either a vacuum pump or aspirator. When the bag has been evacuated, replace the pinch clamp on the tubing.
2. Attach the end of the tubing to the stem on the valve of the gas source.
3. Remove the pinch clamp.
4. As slowly as possible, fill the bag assembly with the gas. The bag should be taut when filled, but not ready to burst. Note: This ensures that any leakage will not result in air diluting the gas sample in the bag.
5. Turn off the gas source and replace the pinch clamp.
6. Remove the tubing from the stem. The bag now contains a slightly pressurized gas sample.
7. Repeat steps 1 through 6 for each gas to be massed.

Pipet Holder

1. Take the 12-inch length of plastic tubing and cut ½" lengths to use as the pipet stem holders in Part 2.

Safety Precautions

Acetone, ethyl alcohol and isopropyl alcohol are all flammable and fire risks. Acetone and isopropyl alcohol are slightly toxic by ingestion and inhalation. Ethyl alcohol is made poisonous by the addition of denaturant—it cannot be made nonpoisonous. If ammonia or chlorine gases are used in Part 1, release these gases in an efficient working fume hood. Wear chemical splash goggles, chemical-resistant gloves and a chemical-resistant apron. Have students wash hands thoroughly with soap and water before leaving the laboratory. Please consult current Safety Data Sheets for additional safety 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. In Part 1, carefully release nontoxic gases into the atmosphere. Unknown or toxic gases should be released only under an efficiently operating fume hood. Ethyl alcohol may be disposed of according to Flinn Suggested Disposal Method #26b. Isopropyl alcohol and acetone may be disposed of according to Flinn Suggested Disposal Method #18a.

Lab Hints

• Because of the small quantities involved, an analytical balance gives the best results. Two significant figures should be obtained with a balance that weighs to ±0.001 g, three significant figures with a balance that weighs to ±0.0001 g.
• In Part 2, it is necessary to quickly remove the pipets from the hot water bath and immerse them in room temperature water. Be sure the outside of the pipets are dried completely before measuring their mass.
• There are various sources for gases to test. Propane and butane are available from hardware supply stores. Burner gas is a source for methane. Lecture bottles of gases, ranging from low molecular weight gases, hydrogen and helium, to the high molecular weight gas, sulfur hexafluoride, are available from Flinn Scientific.
• Step 2 in Part 1 requires a good amount of force to pull the syringe plunger out. Keep fingers away from the plunger shaft to avoid pinching fingers if the plunger should slip.
• An alternative method can be used to determine the true mass of each gas. Each gas is massed in the syringe to give its apparent mass. In cases where the gas has a low molar mass value, this apparent mass value will actually be negative (when the mass of the syringe plus air is subtracted from the mass of the syringe plus sample gas). The true mass is calculated by adding the mass of the same volume of air to this apparent mass. Use the density of air tables in the CRC Handbook of Chemistry and Physics to determine the mass of air.

Answers to Prelab Questions

1. The following data represents a determination of the “molar mass” of a sample of air by comparison with the mass of oxygen as the reference gas. Assuming the air is 79% nitrogen, 20% oxygen and 1% argon and that these gases act as ideal gases, calculate both the experimental and theoretical “molar mass” of air.

   Mass of evacuated syringe: 40.687 g

   {14001_PreLabAnswers_Table_3}
   (Hint: For determining the theoretical molar mass of air, assume the percentages represent the mole fraction of each gas in the solution of air.)
   1. Mass of gas/Mass of oxygen: (0.054/0.060) = 0.90
   2. Experimental molar mass: 0.90 x 32.00 g/mol = 29 g/mol
   3. Theoretical molar mass: (0.79)(28.01 g/mol) + (0.20)(32.00 g/mol) + (0.01)(39.95 g/mol) = 29 g/mol
2. A determination of the molar mass of methyl alcohol yielded the following data.

   Temperature of boiling water bath: 99.5 °C
   Barometric pressure: 738 mmHg
   Temperature of room temp. water bath: 24.0 °C
   Density of water at room temperature: 0.9973 g/mL

   {14001_PreLabAnswers_Table_4}
   Using the data, fill in the rest of table. Calculate the molar mass of methyl alcohol using Equation 3 and compare this value to the actual molar mass of methyl alcohol. The volume of the pipet is equal to the volume of water in the pipet. Use the relationship of mass and density to determine this volume. Once the volume of the pipet is determined, Equation 3 in the Background section can be used to calculate the molar mass of methyl alcohol.
   1. Mass of condensed methyl alcohol: Trial 1: 1.571 g – 1.557 g = 0.014 g
   2. Mass of water: Trial 1: 16.001 g – 1.557 g = 14.444 g
   3. Volume of water: Trial 1: 14.444 g / 0.9973 g/mL = 14.483 mL
   4. {14001_PreLabAnswers_Equation_7}
   5. Actual molar mass, CH₃OH: 32g/mol

Sample Data

Part 1. Molar Mass of Gas Samples
Mass of evacuated syringe: 40.726 g

*Data for sulfur hexafluoride, SF6

Part 2. Molar Mass of Volatile Liquids
Temperature of boiling water bath: 99.0 °C
Barometric pressure: 738 mmHg
Temperature of room temp. water bath: 24.0 °C
Density of water at room temperature: 0.9973 g/mL

Jumbo Pipets Volatile Liquids

Answers to Questions

Part 1. Molar Mass of Gas Samples

1. Why can the buoyancy force in this experiment be ignored?

   Buoyancy can be ignored in this experiment since the mass of the assembly with a vacuum inside has been determined. Whatever buoyancy force is present is identical for both the vacuum case and the gas case. The difference in mass between the syringe filled with gas and the evacuated syringe is the true mass of the gas in the syringe.

2. Determine the mass of each gas in the syringe. Enter these values in the Part 1 Data Table.

   Mass of gas = Mass of syringe assembly with gas – Mass of evacuated syringe (See sample Part 1 Data Table).

3. How should the number of molecules trapped in the syringe compare between the various gases? Explain.

   Since they are at the same pressure, temperature and volume, Avogadro’s hypothesis states they have equal numbers of molecules.

4. Determine the ratio of the mass of gas/mass of oxygen for each gas. Enter these values in the Part 1 Data Table.

   Determine the ratio by taking mass of gas (step 2) and dividing by mass of oxygen (two significant figures are shown; if a 0.0001-g balance is available, three significant figures would be allowed, See sample Part 1 Data Table).

5. How should the ratio of the mass of one molecule of gas/mass of one molecule of oxygen compare to the ratio of the mass of gas/mass of oxygen? Explain.

   The ratio should be the same since the factor of the number of molecules would cancel out.

6. Use the molar mass of oxygen as a reference to determine the molar masses for the other gases tested in Part 1. Enter these values in the Part 1 Data Table.

   Molar mass = (mass of gas/mass of oxygen) x 32.00

7. Determine the accepted molar masses for each gas used (including the air value calculated in Prelab Question 1).

   The accepted molecular weights from formulas are:
   Carbon dioxide = CO₂ = 12.0 + 2(16.0) = 44.0
   Burner Gas if pure CH₄ = 16.0
   Sulfur hexafluoride = SF₆ = 32.1 + 6(19.0) = 146.1
   Air is 79% N₂, 20% O₂ and 1% Ar, so molecular weight = 0.79(28) + 0.20(32) + 0.01(39.9) = 28.9

8. Determine the percent error in your molar masses.

   % error for carbon dioxide = (|45 – 44.0|/44.0) x 100 = 2.3% % error for burner gas = (|18 – 16.0|/16.0) x 100 = 12% (Assume burner gas is pure CH4 ) % error for sulfur hexafluoride = (|140 – 146.1|/146.1) x 100 = 4.2% % error for air = (|30 – 28.9|/28.9) x 100 = 3.8%; also this assumes dry air; if humidity is high, this air would contain water vapor (m.w. = 18) that would lower the experimental result.

9. How do the molar masses compare to the accepted values for each gas tested? Are there any patterns?

   The two values for gases whose molar mass is more than air came out less than true molecular weight. The one gas less then air came out higher than molar mass (however it is not likely pure). There is a systematic error due to a slight bit of air in the rubber tubing.

10. Which gases should have the greatest experimental uncertainty? Explain.

    A very light gas should have the greatest uncertainty since the mass of gas is a small value obtained by subtraction.

Part 2. Molar Mass of Volatile Liquids

1. Determine the mass of the condensed, volatile vapor for each pipet trial and each unknown in Part 2. Enter these values in the Part 2 Data Table.

   Mass of vapor = (mass of pipet and liquid) – mass of pipet (See sample Part 2 Data Tables).

2. Use the CRC Handbook of Chemistry and Physics to determine the density of water at the temperature of the room temperature water bath used in this experiment. Enter this density value in the Part 2 Data Table. Use this value and the mass of water in each filled pipet to calculate the volume of each pipet.

   From the CRC Handbook of Chemistry & Physics, the density of water at 24.0 °C is 0.9973 g/mL.
   Volume of pipet = mass of water / density of water

3. Determine the mass of the condensed volatile liquid for each run. Enter the values in the Part 2 Data Table.

   Mass of liquid = mass of pipet and condensed liquid – mass of pipet

4. Calculate the molar mass of the liquid used in each run and the average of the three runs for each volatile liquid.
   {14001_Answers_Equation_8}
   T = 99.0 + 273 = 372 K (See sample Part 2 Data Table).
5. Volatile liquids with lower boiling points often give better results then those with higher boiling points. Suggest a reason for this.

   A higher boiling point would indicate stronger attractive forces. This nonideal factor would increase the molar mass found. Also, a higher boiling point would mean more difficulty in having all liquid in vapor stage.

6. What effect would vapor condensation in the necks of the jumbo pipets have on the reported molar mass? How large an error might this introduce?

   Condensation would increase the apparent mass of vapor, hence causing an increase in the reported molar mass. It would be a tremendous error since the mass of the same volume of liquid is huge compared to the mass of that volume of gas.

7. Some liquids have enough attractions between molecules to form dimers. (Dimers are molecules formed from the combination of the identical molecules, A + A → A₂.) What effect would this have on the experimental molar mass?

   Dimers would increase the molar mass since some units that should be monomers will have twice the mass.

Introduction

The molar masses of compounds are used daily in the chemistry profession. The molar mass is defined as the mass, in grams, of 1 mole of any element or compound. How is molar mass determined and how is the molar mass of an unknown found? In this experiment, the molar masses of sample gases are determined directly and the molar masses of several volatile liquids will be calculated based on measurements of their vapor density.

Concepts

• Molar mass
• Ideal gas law
• Buoyancy

Background

The ideal gas law relates the four measurable properties of a gas (P, V, n, T). In this experiment, the ideal gas law will be used to determine the molar mass of gases and volatile liquids.

The number of moles (n) of any pure substance is equal to its mass divided by its molar mass. Substituting for n in Equation 1 and then rearranging produces the equation for the molar mass of a gas. In Part 1, the mass of several “unknown” gases (X) is measured and compared to the mass of the same volume of oxygen. If two identical volumes of different gases are at the same temperature (T) and pressure (P), then their mass ratio must be equal to their molar mass ratio (Equation 4). Rearranging: When measuring the mass of a gas, the effect of buoyancy must be taken into account. Air, like water, exerts a positive or upward buoyant force on all objects. This force is compensated for in balances when massing liquids and solids. When massing gases, however, this force is not compensated for and is real. The apparent mass of a gas is less than the true mass of the gas. In Part 1, the mass of each gas will be measured in a 60-mL gas syringe that has been evacuated of air and set to a fixed volume. The syringe is first massed with no gas in the fixed volume, then with a sample gas in the fixed volume. The effect of buoyancy is thus eliminated and the true mass of each gas, including air, can be determined directly.

In Part 2, the molar masses of several volatile liquids with boiling points well below the boiling point of water are determined. A small sample of the liquid is placed in a tared 15-mL plastic pipet and the pipet is then heated in boiling water to vaporize the liquid. The air and excess vapor escape, leaving the pipet filled only with the volatile liquid vapor at atmospheric pressure and at the temperature of boiling water. The pipet is then removed and cooled to condense the vapor.

Once cooled, the pipet is weighed. By massing the same pipet filled with deionized water, the volume of the pipet is calculated. The molar mass of the volatile liquid is then determined from Equation 3 using the mass of the condensed vapor, the volume of the pipet, the atmospheric pressure, and the temperature of the boiling water.

Experiment Overview

The purpose of this experiment is to determine the molar masses of various gases and volatile liquids. In Part 1, the gases are massed with a special gas syringe and their molar masses are determined by comparisons to data from oxygen measurements. In Part 2, liquids are volatilized and condensed in a fixed volume. The condensed vapor is massed and the liquid’s molar mass is calculated from the experimental data.

Materials

Prelab Questions

1. The following data represent a determination of the “molar mass” of a sample of air by comparison with the mass of oxygen as the reference gas. Assuming the air is 79% nitrogen, 20% oxygen and 1% argon, and that these gases act as ideal gases, calculate both the experimental and theoretical “molar mass” of air. See Equation 4 in the Background section.

   Mass of evacuated syringe: 40.687 g

   {14001_PreLab_Table_1}
   (Hint: For determining the theoretical molar mass of air, assume the percentages represent the mole fraction of each gas in the solution of air.)
2. A determination of the molar mass of methyl alcohol (CH₃OH) yielded the following data.

   Temperature of boiling water bath: 99.5 °C
   Barometric pressure: 738 mmHg
   Temperature of room temp. water bath: 24.0 °C
   Density of water at room temperature: 0.9973 g/mL

   {14001_PreLab_Table_2}

Using the data, fill in the rest of the table. Calculate the molar mass of methyl alcohol using Equation 3 and compare this value to the actual molar mass of methyl alcohol. The volume of the pipet is equal to the volume of water in the pipet. Use the relationship of mass and density to determine this volume. Once the volume of the pipet is determined, Equation 3 in the Background section can be used to calculate the molar mass of methyl alcohol.

Safety Precautions

Acetone, ethyl alcohol and isopropyl alcohol are all flammable liquids and fire risks. Acetone and isopropyl alcohol are slightly toxic by ingestion and inhalation. Ethyl alcohol is made poisonous by the addition of denaturant—it cannot be made nonpoisonous. If ammonia or chlorine gases are used in Part 1, release these gases in an efficient working fume hood. Wear chemical splash goggles, chemical-resistant gloves and a chemical-resistant apron. Exercise care when working with the hot water bath. Wash hands thoroughly with soap and water before leaving the laboratory.

Procedure

Part 1. Molar Mass of Gas Samples

1. Push the plunger of the 60-mL syringe to the bottom of the specially prepared syringe and attach the syringe tip cap to the tip of the syringe.
2. Evacuate the syringe barrel: Pull the plunger to the 60-mL mark and place the nail in the prepared hole in the plunger so that the syringe plunger, when let go, returns to about the 50-mL mark. Note: This step requires two people—one person pulls the plunger out past the 50-mL mark and the other person then inserts the nail in the prepared hole.
3. Find the mass of the complete evacuated syringe assembly from step 2 to the nearest 0.001 g. Record the mass in the Part 1 Data Table.
4. Remove the syringe tip cap to allow air to enter the syringe. Replace the syringe tip cap and measure the mass of the complete syringe assembly filled with air. Record the mass in the Part 1 Data Table.
5. Remove the nail and the syringe tip cap. Depress the plunger to expel the air from the syringe.
6. Go to the gas delivery bag of oxygen in the fume hood and attach the syringe to the latex tubing. (This can be done by angling the tip of the syringe at 45 degrees to the end of the tubing, then working the tubing over the tip of the syringe.)
7. Release the pinch clamp on the gas delivery bag and draw oxygen into the syringe until the plunger is slightly past the 50-mL mark.
8. Insert the nail into the hole in the plunger, then push the plunger forward so the nail rests on the syringe barrel.
9. Reattach the pinch clamp to the latex tubing.
10. Hold the plunger in while releasing the syringe from the latex tubing. Immediately attach the syringe tip cap.
11. Measure the mass of the complete syringe assembly filled with oxygen and record the mass in the Part 1 Data Table.
12. Remove the syringe tip cap and expel the oxygen (use a fume hood when releasing gases that are unknown or poisonous).
13. Repeat steps 5–12 with the other assigned gases.

Part 2. Molar Masses of Volatile Liquids

1. Place a 400-mL beaker on the hot plate and add about 300 mL of water to the beaker, along with several boiling stones. Turn on the hot plate to boil the water.
2. Obtain three 15-mL jumbo Beral-type pipets. With pliers, pull the thin stems of each so that a very fine “capillary” tip is formed where the stem has been pulled (see Figure 1).
   {14001_Procedure_Figure_1}
3. Cut the pipet as shown in Figure 1 so that the capillary tip is less than 1-cm long.
4. Label the pipets 1, 2 and 3 with a permanent marker.
5. Mass each pipet to the nearest 0.001 g and record this mass in the Part 2 Data Table.
6. Draw 2–3 mL of the ethyl alcohol from the labeled bottle in the hood into each of the previously prepared and labeled pipets.
7. Insert the tips of the pipets containing the ethyl alcohol into the short piece of plastic tubing, then secure the tubing with a test tube clamp (see Figure 2).
   {14001_Procedure_Figure_2}
8. Lower the pipets into the boiling water bath. Make sure the entire bulb of each pipet is below the water line.
9. Heat for at least five minutes.
10. Carefully remove the pipets from the water. Inspect each pipet. If any liquid remains in a pipet bulb, heat the entire assembly for another minute.
11. Cool the pipets by lowering the pipet assembly into a bath of room temperature water in a 400-mL beaker.
12. Record the temperature of the boiling water bath and the barometric pressure of the room in the Part 2 Data Table.
13. Dry the pipets with paper towels and mass each pipet, which now contains only the consensed vapor, to the nearest 0.001 g. Record these values in the Part 2 Data Table.
14. Fill a 250-mL beaker with room temperature deionized water.
15. Fill pipet 1 with the deionized water, then expel the water into the sink to flush the remaining ethyl alcohol from the pipet. Repeat this process several times.
16. To determine the volume of pipet 1: Fill the pipet completely with deionized water, dry the outside, and mass the pipet and water. Record the mass in the Part 2 Data Table.
17. Repeat step 16 for pipets 2 and 3.
18. Repeat steps 2 through 13 for acetone and isopropyl alcohol.

Student Worksheet PDF

14001_Student1.pdf