Introduction:

This lab was designed so that we, the students, could learn how to determine the molar volume of a gas effectively.

Method:

The first step that we took to accomplish our goal was to put on our safety goggles and choose a lab station to work at. We received one 400ml beaker, one polyethylene pipet, two test tubes with hole rubber stoppers, two small pieces of magnesium (Mg), one thermometer and a vial of hydrochloric acid (HCl). We took the 400ml beaker and filled it about 2/3 full of water (H20) that was 18 OC. Then we measured our pieces of Mg at 1.5 cm and determined that their mass was 1.36*10-2 g. We filled the pipet 2/3 full of HCl and poured it into one of the test tubes. Then, we covered the HCl with just enough H2O so that no H2O would be displaced when the stopper was inserted. After inserting the stopper, we placed the Mg strip into the hole, inverted the test tube and placed it in the 400ml beaker. HCl is heavier than H2O, so it floated from the tube, into the bottom of the beaker, reacting with the Mg along the way to produce hydrogen gas (H2). We then measured the volume of the H2, cleaned up our equipment and performed the experiment a second time.

Results:

The formula for reducing the volume of the H2 to STP in trial 1:

P1 V1 = P2 V2 766.0 mm Hg * 14.5ml = 760 mm Hg x

T1 T2 291 K 273 K

X = 13.71 ml

The equation for finding the molar volume in trial 1 with conversion to Liters:

Volume of H2 in STP 13.71 ml * 1L

Moles of Mg used 5.6 * 10-5 moles 1000 ml

Molar Volume in trial 1 = 24.48 L/mole

When the Mg interacted with the HCl, we observed a bubbling reaction.

Discussion:

The molar volume of the H2 in our experiment is very close to the theoretical molar volume, but I think that the deviation lies in the temperature of the H2O: in the first trial it is too high and in the second one too low. If I were to do the lab again I would change the temperature of the H2O and try to achieve the theoretical molar volume. In addition, this lab assumes that the ideal gas law is in place, so that skewed the results as well. The equations for determining the molar volume were extremely difficult for me and the first time I went through the lab I calculated a 93% and 78% error for trials 1 and 2 respectively. I realized that the 1.36 * 10-2 should be 1.36 * 10-3. After discovering this, I corrected it and came up with the much more reasonable percentage error.

Bibliography:

Chemistry In Microscale. New York: McGraw-Hill, 1991

Word Count: 477