Lab 1.3
m olar m ass of a gas
Introduction
You have often calculated the number of moles (n) or the number of atoms (N) in a sample using a
sample’s measured mass (m) and the com pound’s known molar mass (M). What would you do, however, if a
compound’s molar mass were not known? This would be the case, for instance, if a com pound’s formula were
not known. In such a situation, you could theoretically count one mole (or some other number) of molecules
of the substance and measure the mass of these molecules. But this would of course be impossible since you
could not actually count the individual molecules, nor could you count up to Avogadro’s number in your life.
In the case of gases, however, it is possible to (in effect) count molecules. Since the behavior of an ideal
gas is independent of the identity of the particular molecule or molecules of which the gas is composed, you
can calculate the number of molecules of a gas occupying a space by measuring the temperature, pressure,
and volume of the gas. If you additionally measure the mass of the gas occupying this space, you can
calculate the unknown molar mass of the gas. Measuring the mass of a gas sample may seem a difficult task.
All that is necessary, however, is to “trap” the gas by liquefying it so that it can be placed on the balance.
Pre-lab calculation
Prior to lab, calculate the molar mass of the substance for which these lab observations are made. Show your
work thoroughly. A clean, dry flask has a mass of 122.3155 g when sealed with foil. The foil is opened and
4.72 m L of liquid is poured in. Room temperature is 21.5°C and atmospheric pressure is 748.2 torr. The flask
is re-sealed and immersed up to its neck in 15.2°C tap water. The hot plate is turned on, and as the water
reaches 88.7°C the last liquid evaporates. The flask is immediately cooled in a 0.5°C ice bath, dried, and the
flask and liquid (still sealed with foil) have mass 123.4511 g. Finally the flask is found to hold 263 m L of water.
CAUTION!! Most of the compounds used in this lab are flammable! Therefore no flames
(i.e. no Bunsen burners) may be used. Additionally, some of the compounds are mildly toxic if
inhaled or absorbed through the skin. Therefore do not directly breathe the vapors from the
compounds, especially while heating them, and also avoid skin contact with them as liquids.
methanol
ethanol
acetone
methyl ethyl ketone
cyclohexane
Procedure
1. Tightly crimp a small square of aluminum foil over the top of a clean, dry Erlenmeyer flask.
2. Punch a very small hole in the foil using a pin. (The smaller the hole the better!)
3. Mass the flask (with the foil in place) on the analytical balance. (If the balance displays “- - -”, meaning
that its capacity has been exceeded, trade the flask for a lighter one.)
4. Measure 3-4 mL of the sample in a graduated cylinder. [This volum e need not be exact. Why?]
5. Open the foil, pour the liquid measured in step 4 into the flask, and replace the foil, being sure it is very
tightly sealed around the neck.
6. Attach a clamp to a ring stand and firmly clamp it onto the neck of the flask.
7. Place a large beaker on a hotplate and position the clamped flask as deeply in the beaker as possible.
8. Fill the beaker with as much tap water as possible. (You may re-use the heated water already in the
beaker for subsequent trials rather than starting with an empty beaker of cool water.)
9. Heat the water, monitoring the water temperature near the middle of the flask. (The water probably will
need to reach 70-100°C to accomplish what is described in the next step.)
10. Observe the liquid inside the flask; as soon as no more liquid remains, record the water temperature.
11. Immediately remove the flask from the water and quickly immerse it neck-deep in an ice bath to cool it.
12. Do not rem ove the foil. Dry the cooled flask’s exterior; be sure no water remains under the foil’s edges.
13. Mass the flask again on the analytical balance, with the foil in place and the liquid inside.
14. Open the flask (without damaging the foil, if possible); pour the liquid into the waste container; hold the
flask in the heated water to evaporate the remaining liquid, then dry the outside of the flask.
Repeat steps 1-14 at least twice more using different unknowns (A, B, and/or C). AFTER all those trials...
15. Fill the flask to the brim with tap water. Pour every bit of this water into a graduated cylinder and record
this as the interior volume of the flask. Repeat this measurement two more times. (Use the average of
these volumes in all subsequent calculations.)
16. Record the barometric pressure.
©Deerfield High School
AP Chem istry
1-3
Data prepare an appropriate table to contain data for three trials (ethanol and two unknowns)
Calculations and analysis of data
Complete the calculations necessary to determine the molar mass for each trial; your final lab report
should show a complete example (equation, substitution of data, solution with units) of these calculations for
the ethanol trial. Also complete these calculations for your other samples, although you need not show the
work for these. Finish these by the next class period; we will then pool the class data for each unknown so
you can compute a better average for each compound.
Your unknowns are among the following compounds:
acetone
cyclohexane
ethanol
ethyl acetate
methanol
methyl ethyl ketone
CH3COCH3
C6H12
CH3CH2OH
CH3COOC2H5
CH3OH
CH3COC2H5
Post-lab questions
1. Liquid ethanol has a density of 0.7849 g/mL. Suppose a student poured 3.5 mL of liquid ethanol into a
flask with interior volume 265 mL and performed this experiment at 95°C and an atmospheric pressure of
743 torr.
a. Calculate how many times the ethanol vapors could fill the flask as all of the liquid ethanol vaporizes.
b. On the basis of the above calculation, explain why it is reasonable to assume that there is essentially
no air left in the flask just before it is removed from the boiling water and plunged into the ice bath.
2.
In steps 10-11 of the procedure, you were instructed to stop heating the flask as soon as no more liquid
remained inside of it. But suppose you accidentally had continued heating the flask for 5 minutes after the
liquid was gone.
a. Calculate the rms speed of the nitrogen molecules in the room-temperature air outside the flask.
b. Calculate the rms speed of the vaporized ethanol molecules inside the heated flask.
c. Imagine that you are ~10 –9 m tall and are standing near the hole in the foil; considering your
calculations above, describe what you would expect to see (based on the above calculations) and the
consequences of this for your experimental results.
3.
The ethanol we used is denatured (i.e. it was made poisonous so that it cannot be drunk as liquor) by
adding isopropanol (CH 3CHOHCH 3) and/or benzene (C 6H 6).
a. Explain how this should affect the experimental value for the molar mass of ethanol.
b. Do the class data follow this prediction?
4.
Consider each of the following potential sources of error in this lab. For each error, (i) make a
reasonable estimate of its size based on the experimental conditions, (ii) calculate what percent error this
is, and (iii) specify whether the error would increase or decrease the apparent molar mass.
a. the uncertainty in reading the thermometer (assume a low reading)
b. the uncertainty in reading the graduated cylinder (assume a low reading)
c. water accumulation under the foil which was not removed before making the final mass measurement
1-4
AP Chem istry
©D eerfield High School