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    Classic Experiment For Young Scientists
    By Enrico Uva | August 29th 2012 06:57 PM | Print | E-mail | Track Comments
    About Enrico

    I majored in chemistry, worked briefly in the food industry and at Fisheries and Oceans. I then obtained a degree in education. Since then I have...

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    The Space Age began when the Russians launched Sputnik in October of 1957. Thanks in part to the film October Sky and the book Rocket Boys, even today's youth realizes that early achievements from the race to space inspired many students to choose careers in science and technology. The inspirational sighting of Sputnik led a West Virginian coal miner's son to rocket-building, which in turn led to science fairs, industrial engineering and a career with NASA. In fact, the competitive juices stirred by seeing the technological feat of the "enemy" led to the creation of the agency that put men on the moon, and it made Americans revamp their science educational programs. They created Chem Study, PSSC Physics and BSCS biology.

    All 50 years of subsequent changes in senior high school science education on the continent have paled in comparison to the older quality programs. Rather than boring readers by venturing into the black hole of pedagogical debates, I will examine a classic chem study lab that reveals how a deceivingly simple measurement like gas volume can turn into a craft when we prioritize reproducibility and accuracy.

    A wide range of elemental metals will surrender their loosely held valence electrons to aqueous naked protons piggybacking on water molecules, in other words, to an acidic solution. Every pair of relinquished electrons will bond a pair of H+ ions, and they will escape from solution like nude bathers coming out of the bushes when they get their clothes back from a prankster. Now this hydrogen, H2, like any gas, will spread and assume the volume of any container, but how do we assure ourselves that it's not mixed with other gases like those of air? 

    If we have the acidic solution in a gas burette, the rising bubbles of hydrogen gas can push a solution out of the glass container. We will see later that this will not automatically guarantee purity, but before we cross that bridge, how do we set up the apparatus and trap the hydrogen so that we can measure its volume?

    We pour a small amount (10.0 ml) of a strong acid (6M HCl) into a burette and top off the rest with water. Next, with a ruler, we measure the length of the magnesium ribbon piece whose electrons will clothe the acidic protons. Why measure the magnesium in the first place?  Knowing how much reactant is associated with the production of hydrogen will help us check if our eventual measurement of H2 agrees with ideal gas theory. Why a ruler and not a balance? With a 50.00 ml burette and atmospheric pressure, it would be too easy to generate excess hydrogen and push all the water out. That would lead to the loss of gas, and we would not be able to record the volume. So we have to be stingy with the magnesium. But when we use only a small piece of magnesium to generate a conservative amount of hydrogen gas,  because of the type of balances typically found in the classroom, we would end up introducing a large error in the mass of the magnesium. For example, five centimeters of Mg ribbon has a mass of about 0.06 grams +/-0.01 gram, a 17% uncertainty, but measured with a ruler, the uncertainty associated with 5.00 cm is only 0.05 cm or 1%. We actually measure the mass of a full meter of magnesium 1.20 g  +/-0.01 , an uncertainly of only 0.8 %,  but use only 1/20 th and apply the factor to the full meter of Mg's mass.

    To keep the magnesium in place and making sure it all reacts and does not escape through the hole in the gas burette's rubber stopper, we place it in a little copper cage which itself is not capable of reacting with HCl. The cage is made out of copper wire that's wound around the rolled up magnesium, and one straight end of the wire is slipped through the hole. Making sure that there are no air bubbles (more water is added if some persist), we stopper the burette and invert it into a large cylinder, 3/4-filled with water.


    The inversion causes the dense acid to fall towards the bottom, where the caged magnesium awaits its arrival. One can follow the approximate movement of the acid because the concentrated acid and water do not have the same index of refraction, leading to a wavy appearance at the front of the invasion. Soon the bubbles of hydrogen appear, rise to the top of the glass and push down on the solution, forcing some of it out through the hole in the stopper.

    When the magnesium runs out, the hydrogen stops pushing the liquid out of the burette, which is graduated. But it's still too soon to read the volume of the gas produced. At constant pressure the volume of the gas will be proportional to absolute temperature, so if we don't let the gas cool to room temperature, we will end up with an inflated volume, just like you end up with inflated stock prices when investors get too optimistic. 

    The volume of the cooled gas still seems to fluctuate slightly if one bops the burette up and down the cylinder of water. The weight of water itself influences the total pressure (every centimeter of water-depth corresponds to an additional 0.1 kPa), so to make sure that we get a volume corresponding to the atmospheric pressure, we have to lift the burette until the level of remaining water inside the burette equals the outside water level of the cylinder.

    So now we have the volume of hydrogen, correct? Not quite. Any time you collect a gas over water, some water will evaporate and adulterate the collected gas. So the pressure of the gas, although now equal to the atmospheric pressure, is actually the sum of the hydrogen 's partial pressure and that of water. We usually provide students with a table of values of temperature and corresponding partial pressures for water, which unfortunately leads to the misconception that the values are derived from a simple gas law. It's actually calculated from the liquid's heat of vaporization, temperature and two constants: 


    Once this partial pressure of water vapor is obtained, it is subtracted from the atmospheric pressure, and the difference is hydrogen's pressure. The ratio of hydrogen's partial pressure to that of the total pressure multiplied by the burette-reading equals hydrogen's adjusted volume.

    Because at low pressures and relatively high temperatures there are virtually no attractions between hydrogen's molecules, the behavior of hydrogen's gas is ideal and its volume, V, is given by:

    V = nRT/P.

    After obtaining n, the number of hydrogen moles, from the moles of magnesium used, we can predict the volume of gas that should have been produced. Most students' measured volumes come within a couple of percentage points of the theoretical figure. If they receive the right guidance, they don't all become overwhelmed by all the nuances, as a pessimist might expect. Further discussion and lab exams also help consolidate all the concepts covered by this classic exercise.