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Molar Mass Of A Volatile Liquid Advance Study Assignment

Description

A volatile liquid with a boiling temperature in the range of 70-90 °C is placed in a small, dry, weighed flask. The mouth of the flask is covered with aluminum foil. A small hole is pierced through the foil. The flask is heated in boiling water until the liquid boils, and for a few seconds thereafter. At that point, the entire space of the flask is filled with vapor at 100 °C. The flask is cooled, dried, and weighed. The volume of the flask is measured. From all these data, together with the barometric pressure, the molar mass of the volatile liquid is calculated.
Set
  • The term vapor is often used to describe the gaseous form of a substance which exists as a liquid or a solid under ordinary circumstances. Standard temperature is defined as 273 K; standard pressure is defined as 1 atmosphere, or 760 torr.
  • The molar volume of a gas at STP is 22.4 L.
  • Keep track of units, and use consistent units throughout.
  • The vapor was produced in a boiling water bath, so the temperature at which the vapor filled this flask was 373 K. Please do not confuse the measured temperature with the defined standard temperature which is for gases collected at the freezing point of water instead of the boiling point of water.

Procedure
  1. Set up a boiling water bath using a 400-mL beaker containing 250 mL of water.
  2. Tightly cover the mouth of a 125-mL Erlenmeyer flask with a small square of aluminum foil. Use a straight pin to make a small hole in the foil cap.
  3. Weigh the empty, capped flask.
  4. Remove the foil cap. Place a 2-mL sample of the liquid to be studied into the flask and replace the foil.
  5. Clamp the flask with a single buret clamp. Transfer the flask to the boiling water bath, immerse, and heat.
  6. Note the liquid refluxing inside the flask.
  7. Heat until liquid is no longer visible and no vapor condensate can be seen emerging from the pinhole. Continue heating 30 seconds beyond this time.
  8. Remove the flask; set it on a hot pad; remove the clamp; and wait for the flask to cool to room temperature.
  9. Dry the flask. Weigh the flask, cap, and condensed vapor.
  10. Dispose of the contents of the flask according to instructions. Fill the flask with water. Pour the water into a 250-mL graduated cylinder, measure the volume, and record.
  11. Measure and record the barometric pressure.

Handout

Name ___________________________ Class ________

Teacher__________________________

DoChem 078 Determining the Molar Mass of a Vapor

Mass of flask + cap =
Mass of flask + cap + condensate =
Volume of flask =
Barometric pressure =

Handout Makeup

Name ___________________________ Class ________

Teacher__________________________

DoChem 078 Determining the Molar Mass of a Vapor

Watch the movie.

  1. Identify the piece of equipment (balance, flask, barometer, graduated cylinder) that most limits the accuracy of this experiment. Justify your choice.
  2. Calculate the molar mass of the vapor using the sample data.
    • Mass of flask + cap = 84.15 g
    • Mass of flask + cap + condensate = 84.77 g
    • Volume of flask = 137 mL
    • Barometric pressure = 757 torr

Closure Questions:

The substance used here is 1,1,1-trichloroethane, CCl3CH3.

  1. Using the molar formula supplied by the teacher, calculate the actual molar mass of the unidentified substance.
  2. Compare the true molecular mass of the substance to the experimentally determined value. Find your percent error.
  3. Identify some likely sources of error in this experiment.
  4. A gas has a density of 1.25 g/L at STP. Find its molar mass.
  5. At STP, 10 liters of a gas has a mass of 13.4 g. Find the mass of 1 mole of this gas?

Teachers Guide

Purpose

To determine the molar mass of a substance from measurements of the density of its vapor.


Materials

(for 10 students working in pairs)

  • 5 125-mL Erlenmeyer flask
  • 5 400-mL beaker
  • 5 250-mL graduated cylinder
  • 5 support stand
  • 5 burner, iron ring, wire gauze, or
  • 5 hot plate
  • 5 hot pad
  • 5 straight pin
  • 5 balance
  • 5 single buret clamp
  • 30 mL of unidentified liquid (1,1,1-trichloroethane)
  • 1 roll of aluminum foil
  • 5 thermometers
  • 5 10-mL graduated cylinder

Lab Hints
  • A volatile solvent with little intermolecular attraction works best. Chlorinated hydrocarbons work well but present toxicity problems. 1,1,1-trichloroethane is a compromise. Be absolutely certain to restrict the amount of 1,1,1-trichloroethane put out in the laboratory.


Time

Teacher preparation: 25 minutes

Class Time: 45-50 minutes


Hazards

1,1,1-trichloroethane is toxic.

Hot objects cause burns.


Precautions
  • Control the amount of 1,1,1-trichloroethane put in the room to no more than 2.5 mL per 2 students. Provide adequate ventilation. Work in a hood if possible.
  • Use caution when handling hot objects.

Disposal

Save unused 1,1,1-trichloroethane for use in later years.


Presentation?

Presentation Question:

  • Identify the piece of equipment (balance, flask, barometer, graduated cylinder) that most limits the accuracy of this experiment. Justify your choice.
    • The weighing error is much larger than are any of the other errors imposed by the apparatus used.

Sample Data

(based on sample data for 1,1,1-trichloroethane):

Mass of flask + cap = 84.15 g
Mass of flask + cap + condensate = 84.77 g
Volume of flask = 137 mL
Barometric pressure = 757 torr

Calculations

(based on sample data for 1,1,1-trichloroethane)

Mass of the condensed vapor
= 84.77 g - 84.15 g = 0.62 g
Volume of the vapor at STP
Vstp = Vobs x (Pobs/Pstp) x (Tstp/Tobs)
Vstp= 137 mL x (757 torr /760 torr) x (273 K/373 K) = 101 mL
The molar mass of the gas sample:
= (0.62 g/ 0.101 L) x 22.4 L/mole = 1.4 x 102 g/mol
Percent error= ( | Experimental - Accepted value | / Accepted value) x 100%
= ( | 140 g/mol - 133.5 g/mol | / 133.5 g/mol) x 100% = 4.9%

Closure?

Closure Questions:

  1. Using the molar formula supplied by the teacher, calculate the actual molar mass of the unidentified substance.
  2. Compare the true molecular mass of the substance to the experimentally determined value. Find your percent error.
  3. Identify some likely sources of error in this experiment.
  4. A gas has a density of 1.25 g/L at STP. Find its molar mass.
  5. At STP, 10 liters of a gas has a mass of 13.4 g. Find the mass of 1 mole of this gas?

Answers to Closure Questions:

  1. The actual molecular mass of 1,1,1-trichloroethane is 133.5 g/mol.
  2. Answers depend upon the substance. For 1,1,1-trichloro-ethane, the true molar mass is 133.5 g/mol, and this experiment led to a value of 140 g/mol. This value is 4.9 % too high.
  3. Errors in mass determination are very important. An uncertainty of 0.01 g amounts to a 1.6 % change in the molar mass. When the liquid in the final flask has evaporated, the vapor pressure at room temperature has led to the displacement of some air. Since that air was in the flask at the outset, it should be weighed, too, so the resulting mass will be too low. Finally, the determination of the point, when no liquid is present, is difficult. Stopping too soon will result in large positive errors in the mass; stopping too late will result in small negative errors in the mass due to diffusion.
  4. molar mass = density (STP) x molar volume (STP)
    = 1.25 g / L x 22.4 L/mol
    = 28.0 g/mol
  5. molar mass = density (STP) x molar volume (STP)
    = (13.4 g/ 10 L) x 22.4 L/mol
    = 30.0 g/mol

Key Words
  • molar mass
  • density
  • vapor
  • volatile liquid
  • pressure
  • volume
  • intermolecular attraction
  • STP

This static was created at 12:31:54 PM on Monday, April 20, 2015

Experiment 11MOLAR MASS OF A VOLATILE LIQUID One of the important applications of the Ideal Gas Law is found in the experimental determination of the molarmasses of gases and vapors. In order to measure the molar mass of a gas or vapor we need simply to determine themass of a given sample of the gas under known conditions of temperature and pressure. If the gas obeys the IdealGas Law, PV = nRT (1) If the pressure P is in atmospheres, the volume V in liters, the temperature T in K, and the amount n in moles,then the gas constant R is equal to 0.0821 L ATM/(mole K). From the measured values of P, V, and T for a sample of gas we can use Equation I to find the number of molesof gas in the sample. The molar mass in grams, MM, is equal to the mass g of the gas sample divided by thenumber of mole n.n = PV Molar Mass = grams (2) RT No. of Moles This experiment involves measuring the molar mass of a volatile liquid by using Equation 2. A small amount ofthe liquid is introduced into a weighed flask. The flask is then placed in boiling water, where the liquid willvaporize completely, driving out the air and filling the flask with vapor at barometric pressure and the temperatureof the boiling water. If we cool the flask so that the vapor condenses, we can measure the mass of the vapor andcalculate a value for MM. WEAR YOUR SAFETY GLASSES WHILE PERFORMING THIS EXPERIMENTEXPERIMENTAL PROCEDURE Obtain a 600-mL Florence flask, a stopper, a cap (2 cm square of aluminum foil), and anabout 10-mLs of unknown liquid from the main workbench. If you should find a crack in theflask, report it to your instructor immediately so that it can be replaced. With the aluminumfoil cap loosely covering the neck of the Florence flask, weigh the empty, dry flask on theanalytical balance.Pour about 5-mL, about half your unknown liquid, into the flask. Assemble the apparatus asshown in Figure 1. Place the cap on the neck of the flask. Observe the liquid level in your flask.Add a two or three, at most, boiling chips to the water in the 600-niL beaker and heat the waterto the boiling point. Watch the liquid level in your flask; the level should gradually drop asvapor escapes through the cap. After all the liquid has disappeared and no more vapor comesout of the cap, continue to boil the water gently for approximately 5. Measure the temperatureof the boiling water. Shut off the burner and wait until the water has stopped boiling (about 1/2minute) and then loosen the clamp holding the flask in place. 11- 1


FIGURE .1Remove the flask from the beaker of water, using paper-toweling hold by the Florence flask bythe neck. To cool the Florence flask and the vapors, immerse the flask in a beaker of coolwater to a depth of about 5-cm. After holding the flask in the water for about 2 minutes. (Asthe flask cools the vapor inside condenses and air rushes into the flask.)Dry the flask with a towel to remove the surface waters and let it cool to room temperature.Reweigh the flask. Read the atmospheric pressure from the barometer.Repeat the procedure using the remaining 5-mL of your liquid sample.You may obtain the volume of the flask by measure its volume. Weigh the dry empty Florencewith the aluminum cap on a top-loading balance. Then fill the Florence flask with deionizedwater and again weigh on a top-loading balance.When you have completed the experiment, return the Florence flask to the main workbench. 11- 2


Data and Calculation:Molar Mass of a Volatile LiquidName _____________________ Instructor ___________ Date _______ Trial 1 Trial 2Unknown no.Mass of flask and stopper ________________g _______________gMask of flask, stopper, and ________________g _______________gcondensed vaporMass of flask, stopper, and water ________________g _______________gTemperature of boiling water bath _______________oC ______________ oCBarometric pressure ___________mm Hg ___________mm HgCalculations and Results ______________atm _____________atm ________________L _______________LPressure of vapor, P ________________K _______________KVolume of flask (volume of vapor), V ________________g _______________gTemperature of vapor, T K ____________moles ____________molesMass of vapor, gNumber of moles of vapor, nMolar mass of unknown, as found bv ________________g _______________gsubstitution into Equation 2Average Molar Mass of unknown ___________ grams 11- 3


Advance Study Assignment: Molar Mass of a Volatile LiquidName _____________________ Instructor ___________ Date _______ 1. A student weighs an empty flask with the aluminum cap and finds the mass to be 55.4415 g. She then adds about 5-mL of an unknown liquid and heats the flask in a boiling water bath at 100oC. After all the liquid is vaporized, she removes the flask from the bath, and lets it cool. After it is cool, she weighs the flask and the condensed vapor, obtaining a mass of 56.0397 g. She then drys the flask and fills it with deionized water weights the flask and water and obtains a total weight of 271.241 grams. The barometric pressure in the laboratory that day is 752.7 mm Hg and the temperature is 19.5oC. The density of water at 19.5oC is found in the handbook of Chemistry and Physics to be 0.996532 g/mL. a. What was the pressure of the vapor in the flask in atm? P = ______ atm b. What was the temperature of the vapor in K? T = ______ Kc. The volume of the flask in liters? V = _______ Ld What was the mass of vapor that was present in the flask? mass = _______ gramse. How many moles of vapor are present?f. What is the mass of one mole of vapor? (Eq. 2.) n = _______ moles 11- 4 MM = _________ g/mole


2. How would each of the following procedural errors affect the results to be expected in this experiment? Give your reasoning in each case. a. All of the liquid was not vaporized when the flask was removed from the water bath. b. The student continued to heat the apparatus for 30 minutes after all the liquid had evaporated. c. The flask was not dried before the final weighing with the condensed vapor inside. d. The flask was left open to the atmosphere while it was being cooled while the student took a 30-minute coffee break just before the final weighing. e. The flask was removed from the bath before the vapor had reached the temperature of the boiling water. All the liquid had vaporized. 11- 5