Experiment 7

Determination of Universal Gas Constant, R, and Molar Mass of a Gas

You will experimentally determine the Universal gas constant, R, expressed in Liters, torr, moles and Kelvin.  Use the pressure in torr throughout (which you will determine by reading the barometer).  If you want to determine the value for R using atmospheres as the pressure unit, convert the above R determined above (L•torr/mol•K) into atmospheres (1 atm = 760 torr; L•atm/mol•K).

To determine experimentally the molar mass of a compound, you will follow the basic procedure used to determine R, but solve the Ideal Gas equation for molar mass.  I know, the Ideal Gas equation does not contain molar mass, but you can express it as a function of molar mass easily, since moles = grams/molar mass.

Part A: Determination of the Unversal Gas Constant, R (L•torr/mol•K)

Follow the procedure outlined in the published experimental protocols.  The only variation here, is a brief description of the procedure.

• Weigh carefully the dry Erlenmeyer flask containing a single boiling stone and a white rubber stopper attached (you must use a white stopper, since a black rubber stopper reacts with the organic chemicals and leeches non-volatile organics, some of which do not evaporate).  The piece of Al foil is not weighed before or after the experiment.  It is used only to keep the stopper and top part of the flask dry.
• Add 5-10 mL of methanol to the flask for evaporation.  (The amount of methanol you add is not important.  The methanol evaporates, leaving only enough methanol gas to fill the container at the appropriate temperature and pressure.  If you add a lot of methanol, it just takes longer for all of it to evaporate.  Neverthleless the same final result will be obtained.)
• After the liquid methanol as completed evaporated (it has boiled away), keep the flask in the boiling water bath for five (5) more minutes, to allow temperature equilibration.  (If you take the flask out of the boiling water bath immediately after the liquid has completely evaporate, all the the methanol gas will not be the temperature of the boiling water bath, but the temperature of the freshly boiled methanol, about 70^oC.)
• Cool the flask but be careful not to let any water go into the opening of the flask, otherwise water will contribute to the recovered mass.
• Remove the Al foil, and make certain the outside of the flask is dry.  Weigh the flask, containing the same boiling stone and stopper.  The difference in mass will be due to the condensed methanol, which was coverted to gas and now condensed.  This is the amount of methanol recoved in the experiment.

After two additional experiments, you will need to determine the volume of the flask.  Note where the stopper goes down into the mouth of the flask.  Fill water up to this level.  Then, using your 100-mL graduate cylinder, measure the total amount of water in the flask.  This will be the volume occupied by the gas during the experiment.

Now, you have the values you need for your experiment:

• Pressure (in torr) using the barometer
• Volume (in Liter) based on your volume measurement
• Temperature (boiling water bath temperature) expressed in Kelvin
• Moles (the number of grams of methanol divided by molar mass)

Before you leave today, please show your values for R (in torr) from each of your experiments, and the averaged value.  Determine a percent error (using the formula and information at the end of this protocol), compared to the real value of R, which is 62.36 L•torr/mol•K.

Experiment 1 value for R: ______________________________ L•torr/mol•K

Experiment 2 value for R: ______________________________ L•torr/mol•K

Average value for R: _______________________________L•torr/mol•K

Percent error for your average value for R: __________________ % (compared to 62.36 L•torr/mol•K)

Part B: Determination of Molar Mass of a Gas

The definition of Molar Mass (MM) is grams per mole of compound.  Using this definition, we can re-arrange the Ideal Gas Equation to calculate the Molar Mass of a chemical (using essentially the same procedure as followed for Part A of this experiment).  The equation for Molar Mass is re-arranged as shown below to solve for moles.  This definition of moles (grams/MM) will be inserted into the Ideal Gas Equation.

Using the new definition for moles above, the unit of moles is inserted into the Ideal Gas equation.  Molar mass can then be determined using the equation shown below:

Now, using this equation for calculating Molar Mass, you can determine the molar mass for one of the unknown chemicals.  

To do the experiment, perform the experimental protocol that you used in Part A to determine the amount of unknown chemical what filled the flask, and which condensed into liquid.  After you determine the mass of the condensed organic chemical for your experiment, use this mass in the Molar Mass formula above.  This should be the molar mass of your compound.  You do not need to do a percent error analysis because you will not know the exact Molar Mass of the chemical you are using.  However, the Molar Masses of the compound used in this experiment are shown below.  You will be graded on how close to the real value your experimental value is.

Useful information:

The chemicals used in this part of the experiment, and including the methanol from Part A, are shown below.  These chemical are listed by increasing Molar Mass, and do not reflect the unknown chemical order.

Chemical	Molecular Formula	Molar Mass (g/mol)	Boiling Point (oC)
Methanol	CH3OH (CH4O)	32.042	64.6
Acetone	CH3COCH3 (C3H6O)	58.08	56.5
2-Propanol	CH3CHOHCH3 (C3H8O)	60.09	82.5
2-butanone	CH3COCH2CH3 (C4H8O)	72.10	79.6
Hexane	CH3CH2CH2CH2CH2CH3 (C6H14)	86.17	69

Percent Error Analysis

To obtain %Error for a measurement, you need to know the Theoretical and Experimental values for your measurement, or set of calculations.   Use this formula to obtain the %Error for a measurement, which is always a positve (absolute) value.

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