Every Year Begins a Millennium - American Chemical Society

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Award Address

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Every Year Begins a Millennium 2000 George C. Pimentel Award, sponsored by Union Carbide Corporation by Jerry A. Bell

In keeping with a millennial theme, year 2000 recipients of American Chemical Society awards were asked to reflect, in their award addresses, on the future of the area for which the award was given.1 Every year begins a millennium, a new thousand years. Reflecting on and learning from the experiences we have each year should influence how we teach in the future. This paper traces a few of the experiences that have enriched my understanding of teaching and taught me lessons about learning. None of these lessons will surprise you, but the examples may reinforce your own reflections and provide a more general message for the future. My first teaching experience was in 1954, my freshman year in college. As part of a social service program, I tutored a half dozen young men who were high school students in Boston’s North End. My tutees were at the North End Union to use the athletic facilities, but the director of the Union felt that the “price” for exercising their bodies ought to be some time devoted to exercising their minds as well. My tutees were average students; they were not failing their courses, but they had to struggle with math and science. For the first term, we went through problems from their texts and homework. Early in the second semester, we were working on solubility problems in chemistry and the students were having a lot of trouble. I couldn’t understand their difficulty; I knew they could do harder algebra than required for these problems. Finally, one of the students asked, “What does it look like?” I didn’t understand the question, so he and the others went on to elaborate with questions like, “Does the whole solution just turn solid?” I said that they must have seen this phenomenon in their laboratory and the response was, “What laboratory?” Since my own recent high school chemistry course had had two double periods each week for laboratory work, it never occurred to me that a chemistry course would not have a lab and I had never asked them.

Lesson: Find out what your students know.

The next week I rode the “T” to the tutorial session with a few test tubes, some straws, and a bottle of limewater in a cardboard box. My tutees and I blew through small samples of limewater and watched the precipitate form. Now the students knew “what it looked like” and could refocus their efforts on the algebra of the phenomena. The algebra no longer seemed so difficult and we could make rapid progress with the paper-and-pencil problems. Lesson: Engage students with phenomena before theory.

It’s not possible in an article, or even a live presentation, to recapture the excitement of that afternoon in the North End Union, but I would like you to experience some of the fascination of the phenomenon. Once your students have reasoned that it is probably the carbon dioxide in their breath that causes the precipitation reaction with limewater, Ca(OH)2 solution, you can ask them to hypothesize what would be observed if pure carbon dioxide is bubbled into limewater. Test the hypotheses by adding pure carbon dioxide to the limewater (1). A visually appealing way to do this (Fig. 1) is to add dry ice to the solution. The chemistry of the carbon dioxide–limewater system is shown in this series of reactions: CO2(g) → CO2(aq)

(1)

CO2(aq) + H2O → H2CO3(aq)

(2)

H2CO3(aq) + OH (aq) → HCO3 (aq) + H2O

(3)

HCO3 (aq) + OH (aq) →

(4)

⫺

⫺

⫺

–

CO32⫺(aq)

+ H 2O

CO3 (aq) + Ca (aq) → CaCO3(s) 2⫺

2+

(5)

Carbon dioxide dissolved in water reacts with water to form carbonic acid, which, in turn, reacts with hydroxide

Figure 1. Reaction of carbon dioxide with limewater. (a) A piece of dry ice is added to a cylinder of limewater; (b) bubbles of carbon dioxide rise through the solution; (c) the solution becomes cloudy; (d) as the carbon dioxide continues to bubble through the solution, the cloudiness disappears. Photos by Jerrold J. Jacobsen and Kristin Johnson

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Journal of Chemical Education • Vol. 77 No. 9 September 2000 • JChemEd.chem.wisc.edu


Award Address from the dissolved Ca(OH)2 to produce bicarbonate ion and then carbonate ion. The carbonate reacts with the calcium ion in solution to produce the observed white precipitate. The more carbon dioxide, the more carbonate and the more precipitate, so we predict that pure carbon dioxide will produce a more copious precipitate than we get by just blowing into the solution. Figure 1 shows that the precipitate does form. But, as more CO2 is added to the solution, the precipitate disappears. I first observed this phenomenon in the mid-1960s, when I was teaching at the University of California, Riverside. It is always a source of puzzlement for students and provides a wonderful teachable moment as we try to figure out what is going on in this chemical system. Adding just one reaction to those above rationalizes the disappearance of the precipitate: H2CO3(aq) + CO32⫺(aq) → 2HCO3⫺(aq)

(6)

When the hydroxide has all reacted, the carbonic acid reacts with the most basic substance left in the system, carbonate ion. As the carbonate ion concentration decreases, the precipitation reaction reverses and the calcium carbonate dissolves. This lecture experiment and the discussion that follows can lead to an exploration of the formation of limestone caves and the structures we find in them. Lesson: Welcome astonishing results that lead to further inquiry.

When I began teaching at Simmons College in 1967, undergraduate research was just getting underway at the College. A senior began work with me on a gas-phase photochemistry problem. Manipulations on vacuum systems are not always easy to visualize or understand, so I was working closely with her to try to be sure she would feel comfortable with the techniques. After a couple of weeks of work, we were preparing a sample and she said, “Please, let me do it myself.” I suddenly realized that I had, in fact, been doing the experiments and she had been almost a spectator. I had not been doing a very good job as a mentor and she taught me an important lesson. Lesson: Guide students to independence.

Later, we tried to express this lesson concretely in the design for the chemistry space in our new science center. A large area in the center of the department is occupied by an independent study laboratory in which each chemistry major has an individual carrel (Fig. 2) where she can carry out experiments alone or in small groups. Students can use their carrels for experiments, studying, discussion, or just hanging out. The objective is to create a “home” for students in the department that makes a statement about their ability to take a good deal of responsibility for their own learning. In the mid-1970s, I began teaching a sophomore-level biochemistry course for nutrition and medical technology majors. A great deal of biochemistry at this introductory level involves straightforward applications of the principles included in general chemistry, such as acid–base properties, equilibria, and kinetics. Many of the students’ problems with this course involved recalling the chemistry they had had two JChemEd.chem.wisc.edu • Vol. 77 No. 9 September 2000 • Journal of Chemical Education

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Award Address Photo by Jerry A. Bell

semesters previously. It seemed to me that a better integration of biochemistry and general chemistry, in greater depth than in the usual general organic–biochemistry course, could alleviate these problems and I have tried that in subsequent courses. Lesson: Integrate the chemistry subdisciplines.

Among the lecture experiments I have used as a part of this integration are the uncatalyzed and catalyzed reactions of aqueous solutions of carbon dioxide with hydroxide ion (2, 3). To a stirred, ice-cold solution of dissolved CO2, seltzer water, containing bromthymol blue indicator is added an amount of hydroxide ion equivalent to about one-half the CO2 (estimated as about 0.075 M at 1 atm pressure of CO2 and 0 °C). As shown in Figure 3, the initially yellow solution (acid form of bromthymol blue) turns blue (base form of bromthymol blue) and remains blue for many seconds before returning to the yellow acidic color.2 Neutralization of the added base is a surprisingly slow reaction. It is especially surprising if this experiment is preceded by one with acetic acid, for which the yellow-to-blue-to-yellow color change is complete within the second-or-two mixing time of the added hydroxide ion. Students recognize that they are faced with a puzzle. The process responsible for the slow disappearance of the base color is not likely to be slow reaction of hydroxide ion. Another explanation must be sought. Most dissolved CO2 is present as CO2(aq). The reaction of CO 2(aq) to form H2CO3(aq), reaction 2 above, is slow in both directions and this could have dire biochemical consequences. The reverse reaction is required to rid our blood of carbonic acid and bicarbonate ion as the blood passes through our lungs. If the reaction were as slow as in our experiment, we would asphyxiate. Blood contains an enzyme, carbonic anhydrase, that catalyzes reaction 2. Repeat the experiment with a drop of blood added to the solution3 and, as Figure 4 shows, the time for the blue color to disappear is much shorter. Even the very dilute enzyme has a striking effect on the reaction rate. These results are fascinating and astonishing. They make enzymatic catalysis real and help set the stage for understanding more quantitative treatments.

Figure 2. Simmons College Chemistry Department independent study laboratory with individual carrels for each chemistry major.

In 1992, I joined the American Association for the Advancement of Science (AAAS) to direct the science, mathematics, and technology education programs. Because of an intriguing challenge, I also became involved with the AAAS programs for scientists with disabilities. A luncheon for students and scientists with disabilities had been held at each AAAS annual meeting for several years. The objective was for students with disabilities to meet scientists with similar disabilities and see that the disability did not close the door to being a scientist. The challenge presented was to enliven these meetings by involving the participants in hands-on science activities and, in so doing, extend the message by having the students and scientists carry out the activities together. This effort, begun in 1993, has continued to the present and we have learned a lot about how to do such an activity. One of the staff members in the disability program at AAAS has cerebral palsy and she and I have worked together over the years to adapt activities, so they could be more easily done by people who are blind or have difficulty with manipulations (4). Often the adaptations are quite simple. Figure 5 shows the materials for an activity in which carbon dioxide gas is produced in a zip-seal plastic bag: two zip-seal bags, each containing a white powder (sodium bicarbonate and potassium hydrogen tartrate) and a plastic spoon; a third empty zip-seal bag; a 3-oz paper cup; and a capped bottle of water. To test the activity, a blindfolded volunteer adds a

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Photos by Jerrold J. Jacobsen and Kristin Johnson

Figure 3. CO2(aq) plus OH–(aq) reaction sequence. (a) Bromthymol blue indicator is added to a stirred, ice-cold solution of seltzer water; (b) a less-than-stoichiometric equivalent of hydroxide ion is added; (c) solution 10 seconds later; (d) solution 20 seconds later; (e) solution 25 seconds later.


Award Address

spoonful of each powder (distinguishable by touch in the bags) to the empty bag, fills the cup about half full of water, puts the cup in the bag without spilling it, seals the bag, and spills the water onto the solids. While mixing the contents of the bag, our volunteer reports observations made by touch, sound, or odor. The hydrogen tartrate ion is more acidic than hydrogen carbonate ion, so carbonic acid is formed in the sealed bag: Htart⫺(aq) + HCO3⫺(aq) → tart2⫺(aq) + H2CO3(aq) (7) The carbonic acid product can decompose to dissolved carbon dioxide, which can bubble out of the solution, the reverse of reactions 1 and 2. Thus, our volunteer reports hearing the fizzing and feeling the bag inflate as it fills with gas. Further, the volunteer reports that the reaction mixture gets quite cold, so we can also conclude that the overall reaction is endothermic. The usual activity to produce carbon dioxide in a bag involves mixing vinegar and baking soda. This adaptation uses water as the liquid, so spills are less of a problem. The reaction is also slower and much more endothermic, so the fizzing goes on longer and the temperature change is easier to detect. The adaptation makes it easier for an unsighted person to do and interpret the activity. Lesson: Adapt to meet the needs of students.

In the past decade or two, many of us have become much more aware that the tests we have been using may be inadequate to assess how much our students have learned about the concepts we are trying to teach. Tests often reward the


Figure 4. CO2(aq) plus OH–(aq) plus a drop of blood reaction sequence. (a) A small quantity of blood is added to a stirred, ice-cold solution of seltzer water; (b) bromthymol blue indicator is added to the stirred, ice-cold solution of seltzer water and blood; (c) a lessthan-stoichiometric equivalent of hydroxide ion is added to the solution; (d) solution 10 seconds later; (e) solution 15 seconds later.

facile use of algorithms, without probing more deeply into the concepts underlying the algorithms. In externally funded curriculum or professional development projects, it has become increasingly important to find ways to show funders that the projects lead to better student or participant understanding of science. Since I believe strongly in the efficacy of activity-based teaching and learning, a natural extension is to assess learning by finding out what the learners can actually do when faced with a new problem. Lesson: Find out what your students can do.

I have experimented with laboratory-based performance assessments and with take-home as well as in-class, small-scale experiments included as part of an examination. In the past several years, the participants in these experiments have been mainly teachers in professional development programs and workshops. One simple technique is to do a lecture experiment that is an extension of others you or your students have done and have the students write a short description of the chemistry that explains their observations. Figure 6 shows one such experiment that is based on the chemistry included in this article. Heaping spoonfuls of sodium hydrogen carbonate and calcium chloride are placed together with a few milliliters of water in a beaker. A gas is produced and the mixture gets warm. A white solid is still present in the beaker. What is the chemistry that is going on to produce these observations? The hydrogen carbonate ions produce a small amount of carbonate ion by disproportionation, the reverse of reaction 6. At their high concentration

JChemEd.chem.wisc.edu • Vol. 77 No. 9 September 2000 • Journal of Chemical Education

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Photos by Jerrold J. Jacobsen, Nancy S. Gettys, and Kristin Johnson

Figure 5. The reaction of cream of tartar and sodium bicarbonate in water can be done by a blind (or blindfolded) experimenter. (a) Materials required; (b) the experimenter has placed a teaspoon of each solid reactant and a cup of water into a zip-sealed bag; (c) the water is poured over the reactants inside the zip-sealed bag; (d) carbon dioxide inflates the bag. For a related activity see p 1104A.


Award Address

in this mixture, calcium ions react with carbonate ion to precipitate calcium carbonate, reaction 5. Loss of carbonate drives the disproportionation reaction to produce more carbonate and, in the process, produces more carbonic acid. The carbonic acid product can decompose to dissolved carbon dioxide, reaction 2, which can bubble out of the solution, reaction 1.4 Thus we have come full circle, from bubbling carbon dioxide into limewater to produce a calcium carbonate precipitate to a reaction in which the formation of calcium carbonate drives carbon dioxide out of solution. And we have come full circle in this article. The general message is simple: Use every experience you have to enrich your understanding of chemistry, of learning, and of effective teaching today, tomorrow, next week, and next year, remembering that every year begins a millennium. W

Supplemental Material

Supplemental material for this article is available in this issue of JCE Online. Several demonstration activities accompanied the presentation in San Francisco. Some of these, including the reaction of carbon dioxide with limewater, are re-created in this issue of JCE Online. Figures 1, 3, 4, 5, and 6 show single frames from the online version. Notes 1. This article is based on the award address for the year 2000 George C. Pimentel Award in Chemical Education, sponsored by the Union Carbide Corporation. The address was presented at the American Chemical Society Meeting in San Francisco, CA, on 28 March 2000. 2. The final color of the partially reacted solution of dissolved CO2(aq) is a yellowish green, because the solution is a buffer of carbonic acid and bicarbonate ion. The pH of

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this solution is in the blue-to-yellow transition range for bromthymol blue indicator. 3. The demonstrator can prick his or her finger with one of the devices used by diabetics to draw blood for glucose analysis, squeeze a drop of blood onto the skin, and swish the finger in a fresh solution of dissolved CO2(aq). Alternatively, but not as memorably, a solution of carbonic anhydrase can be used. 4. A reviewer pointed out that the reaction of sodium hydrogen carbonate and calcium chloride is the reverse of the net reaction that can be written for the Solvay process: CaCO3 + NaCl + CO2 + H2O → 2NaHCO3 + CaCl2 (The sodium hydrogen carbonate initially produced is heated to produce sodium carbonate plus water and carbon dioxide.) The spontaneous direction of this reaction is the one we demonstrated. Discussion of the Solvay process and inquiry into how it is driven “uphill” could be an interesting extension of this demonstration. Literature Cited 1. Scott, E. S.; Shakhashiri, B. Z.; Dirreen, G. E.; Juergens, F. H. In Chemical Demonstrations, Vol. 1; Shakhashiri, B. Z., Ed.; University of Wisconsin Press: Madison, WI, 1983; pp 329–337. 2. Bell, J. A. Am. J. Pharm. Educ. 1991, 55, 383. 3. Bell, J. A. In Chemical Demonstrations, Vol. 2; Shakhashiri, B. Z., Ed.; University of Wisconsin Press: Madison, WI, 1985; pp 122–126, and references therein. Also, see Vol. 3; 1989; pp 188–191 for an experiment relating biochemistry to acid– base chemistry. 4. Bell, J. A.; Summers, L. CHED Newslett. 1995, Spring, 24–28.

Jerry A. Bell is a member of the staff of the Education and International Activities Division, American Chemical Society, Washington, DC 20036; [email protected].



Figure 6. NaHCO3 plus CaCl2 plus water reaction sequence. (a) Solid NaHCO3 is added to water in a beaker containing a temperature probe; (b) solid CaCl2 is added to the mixture in the beaker; (c) mixture and temperature a few seconds later; (d) mixture and temperature 18 seconds later.

Every Year Begins a Millennium - American Chemical Society

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