# Experiment_609_Determining the Molar Mass Using Ideal Gas Law_1_2_0

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 Student Name Laboratory Date:  Date Report Submitted: ___________________________ Student ID Experiment Number and Title Experiment 609:  Molar Mass Determination Using Ideal Gas Law

Experiment 609:  Molar Mass Determination Using Ideal Gas Law

Section 1:  Purpose and Summary

Determine the molar mass of an unknown volatile liquid by measuring the properties of its vapor phase after heating.  Using the properties of its vapor phase, calculate its molar mass using the ideal gas law.

In this experiment, students will determine the molar mass of an unknown volatile liquid by measuring the properties of its vapor phase after heating. The liquid is heated until it is completely vaporized. The vapor phase is trapped in a flask and the properties of the vapor are measured. The molar mass of the sample is then calculated using the ideal gas law:

PV = nRT

where P is the pressure in atmospheres, V is the volume in liters, n is the number of moles of the gas, R is the gas constant, 0.0821 $$\frac{\text {L× atm}}{\text {mol× K}}=$$ , and T is the Kelvin temperature.

Section 2:  Safety Precautions and Waste Disposal

Safety Precautions:

Use of eye protection is recommended for all experimental procedures.

Use tongs or mitts to handle hot flasks and beakers.

Waste Disposal:

Discard any excess unknown liquids in the organic waste container in the fume hood.

Section 3: Procedure

Liquid sample in this experiment may be any of the following:  acetone, isopropyl alcohol, ethyl acetate, or cyclopentane.  Additional liquids may also be added at the request of the instructor.

 1.  Obtain a liquid sample from your instructor.  Your instructor may choose to issue the liquid as an unknown sample. Record the name of the liquid sample here, or the identification number of the unknown:

Section 4. Calculations

 Trial 1 Trial 2 Convert your recorded barometric pressure (c) to atmospheres (atm). Assume that the pressure of the gas, P, is the same as the atmospheric pressure when the flask is full of the vapor of the unknown liquid.    Show your equation here: P = ____________ atm P = ____________ atm Convert the volume of the Erlenmeyer flask (e) to liters (l). Assume that the volume of the gas, V, is equal to the volume of the flask.    Show your equation here: V = ___________ l V = ____________ l Convert the temperature of water bath (b) to Kelvin. Assume that the temperature of the gas, T, is the same as the temperature of the water bath when the liquid has completely vaporized.     Show your equation here: T = ___________ K T = ____________ K Using the ideal gas law, PV = nRT, calculate the number of moles, n, of the gas.    Show your equation here: n = ___________ mol n = ____________ mol Calculate the mass of the sample: (d) – (a) ______________ g _______________ g Using the mass and number of moles of the sample, calculate the molar mass of the liquid unknown.    Show your equation here: ______________ g/mol ______________ g/mol Calculate the average molar mass of the liquid unknown. ______________ g/mol

Post Lab Questions:

1.  Obtain the identity of your unknown liquid sample from your instructor. From the chemical formula of the compound, calculate its molar mass. Compare this value with the average molar mass you obtained from this experiment. Calculate the percent error.

1.  There are some experimental errors or facts that could lead to a higher or lower calculated molar mass compared to the true value. In each case below, determine whether the experimental value would be higher or lower than the true value of the molar mass of the compound analyzed. Explain your answer briefly.

1. You took the flask out of the water bath before the entire sample had vaporized.

1. You took the flask out of the water bath before the entire sample had vaporized, and then put it back in the water bath to finish the vaporization process.

1. The temperature of the vapor inside the flask could be lower than the temperature of the water bath that you recorded.

1. You did not wipe off the traces of water outside the flask before weighing it.

1.  When you calculated the moles of the gas (i.e. vapor inside the flask), you used the ideal gas law. However, the vapor is probably not acting like an ‘ideal gas’ under the conditions of this experiment. Explain why this is so. How would this behavior affect the value of the calculated molar mass? Would it be higher or smaller than the true value?

1.  At the end of the experiment, you weighed your unknown sample in the flask after it has condensed back into a liquid.

1. Why can’t you just measure the mass of the liquid that you originally put into the flask at the beginning of the experiment for use in the calculations?

1. How do we take into account the presence of air in the flask?

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