In the Classroom
The Coupling of Related Demonstrations To Illustrate Principles in Chemical Kinetics and Equilibrium1 Richard A. Pacer Department of Chemistry, Indiana University–Purdue University, Fort Wayne Campus, Fort Wayne, IN 46805 While sample problems can be worked in lecture to illustrate how a rate law may be found from data obtained experimentally, student interest is heightened considerably if the set of data used is generated before the students’ eyes by means of a demonstration. The dependence of the rate of reaction of magnesium with HCl lends itself well to this purpose. If the length of a uniform magnesium ribbon and temperature are kept constant, a rate law Rate = k[H+]n can be found by varying HCl molarity and measuring the rate of formation of H2 gas. A novel feature of this demonstration is the use of a baby bottle to collect and measure the H2 gas. Davenport (1) notes that the baby bottle is the cheapest and most durable volumetric glassware on the market. The facts that it is graduated in cubic centimeters (as well as in ounces) and that the nipple serves as a pressure-sensitive two-way valve make it ideal for a number of chemistry experiments. It should be emphasized that what is being described here is an approximately 15-minute demonstration. More elegant means are available for determining n in the rate law for the Mg/HCl reaction. An example is the experiment described by Birk and Walters (2), based on careful pressure measurements. But theirs is a student experiment requiring some two hours of laboratory time, not a short lecture demonstration. Later in the semester, when students are introduced to ionic equilibria, equal lengths of magnesium ribbon can be placed in beakers containing equal volumes of equimolar HCl, CH3COOH, and H3BO3 . The differing rates can now be related to different concentrations of H+ provided by the three acids, due to different degrees of ionization and Ka values. This very simple demonstration shows up well on an overhead projector. Students can be reminded of the earlier rate law demonstration, and the two can be linked together.
a measure of the volume of H2 produced, permitting one to calculate the average rate of reaction. [CAUTION! Hydrogen pressure will force a stream of HCl out of the bottle if the opening is not covered! Be certain therefore that the bottle is not pointed at anyone as it is being removed. Safety goggles are absolutely essential. Although only small quantities of H2 are generated, the gas is explosive and calls for a well-ventilated room. If the distance between the instructor’s desk and the first row of students is small, use of a safety shield is highly desirable.] The experiment is repeated, using 0.60 M HCl and a fresh strip of magnesium ribbon. This time 60 seconds should be sufficient to give an adequate volume of H2. From the data, the reaction order with respect to HCl concentration may be calculated. (The 0.40 M and 0.60 M HCl solutions may be prepared by simple volumetric dilution from a common source, such as 6.0 M or 12 M HCl. Commercial grade HCl is adequate.) If done as a lecture demonstration (with calculations worked out on chalkboard), about 15 to 20 minutes of class time will be required.
Demonstration 2 The second demonstration is incredibly simple compared to the first. Three small beakers (such as 50-mL size) or Petri dishes are placed on an overhead projector. Into each is placed 30 mL of 1 M acid. The acids used are HCl, H3 BO3 , and CH3CO2H. All solutions should be at room temperature. (A 1.0 M solution of boric acid is fairly close to saturation, but should easily go into solution with mild heating and stirring. The experiment will also work well with slightly lower concentrations, such as 0.80 M acids.) A strip of magnesium (3.2 mm wide, commercial grade, cleaned with steel wool, if necessary) is cut into 1.0-cm lengths. A piece is dropped into each of the three acid solutions at essentially the same time, and results are noted on the overhead.
Procedure
Discussion
Demonstration I A 20-cm length of magnesium ribbon is cleaned with steel wool (if necessary), folded, and placed in the nipple portion of a baby bottle. It must be folded in an irregular manner (not wound), so that essentially all surface area is available for contact with acid. It should be fitted securely so that it will not drop out when the nipple is inverted. The bottle itself is filled to the top graduation mark (240 mL in the bottle I used) with 0.40 M HCl. The nipple is screwed onto the bottle, after which the bottle is inverted and the nipple placed below the surface of the water in a large pail or beaker. Begin timing as soon as the bottle is inverted. After 90 seconds, place your forefinger over the nipple opening and remove the bottle from the water bath. After pointing the nipple end of the bottle away from the audience, the forefinger may be removed. Measure by inspection the volume of HCl remaining in the (calibrated) baby bottle. The difference in volumes gives
In demonstration 1, one might anticipate an n value reasonably close to 2 (Birk and Walters [2], for example, reported a value of 2.06). Typical n values, however, range between 1.6 and 1.7, most likely because of diffusion rate limitations; hence, average rates are somewhat different from initial instantaneous rates. Nevertheless, useful data are generated before the students’ eyes, with which the instructor may use logarithms to evaluate n. A typical sample calculation is given below. Rate = k[HCl]n
R2 k(0.60 M)n = = (1.5)n R1 k(0.40 M)n 45 mL H 2 / min R2 = = 1.93 R1 23. 3 mL H 2 / min
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In the Classroom (1.5)n = 1.93 n log (1.5) = log (1.93)
n=
log (1.93) 0.286 = = 1.63 or 1.6 log (1.5) 0.176
From the data one may also calculate a rate constant. Using the rate law developed by Birk and Walters (2), Rate = k (surface area of Mg)a [H+]b where a = 1 and b = 2, one may use the data given above to calculate an average rate constant, kAVE, of 1.80 × 10{3 mL H2 s{1 mm{2 M{2. For advanced classes, one may wish to postulate a plausible reaction mechanism. The following might be offered for discussion: . 2H +(aq) + 2e{ → 2H (adsorbed on Mg surface) Mg(s) → Mg2+(aq) + 2e{ . 2H → H2(g) In Demonstration 2, students are not told (at least not initially) what the three acids are, but are asked to draw conclusions based on their observations. In a few minutes,
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the magnesium strip in 1 M HCl is completely consumed; the students readily conclude that that beaker must contain a strong acid. But the difference between the other two acids is both striking and puzzling. In 1 M acetic acid, H2 is evolved at a fairly significant rate, propelling the Mg strip about the beaker. But only an occasional bubble is seen forming on the strip in boric acid. This provides an excellent opportunity to involve students in a discussion of the meaning of Ka. Even though H 3BO3 and CH3CO2H are both weak, there is an enormous difference in their relative strengths. The Ka for acetic acid is 1.75 × 10{5, whereas that for boric acid is 5.81 × 10{10. Then, one can tie this demonstration to the earlier one, which showed the dependence of the rate of reaction of magnesium with acid on [H+]n, reinforcing the principles learned earlier. Note 1. Presented before the Division of Chemical Education at the ACS National Meeting in Denver, March 31, 1993 (Paper #333).
Literature Cited
Journal of Chemical Education • Vol. 74 No. 5 May 1997
1. Davenport, D. A. J. Chem. Educ. 1969, 46, 878–879. 2. Birk, J. P.; Walters, D. L. J. Chem. Educ. 1993, 70, 587–589.
