Chemical Education Today
Letters Why Is Bismuth Subchloride Soluble in Acid The explanation for the demonstration “Why Is Bismuth Subchloride Soluble in Acid” by Damon Diemente (J. Chem. Educ. 1997, 74, 398) is missing a key reaction. Diemente attributes the dissolving of BiOCl in acid entirely to the role of the hydronium ion and omits the essential role of chloride ion in forming the BiCl4᎑ ion through the reaction BiO+ + 2H3O+ + 4Cl᎑
BiCl4᎑ + 3H2O
The complexation of Bi3+ with Cl᎑ is very favorable with a value of β4 about 106 (1) and cannot be ignored as Diemente suggests. That Cl᎑ ion is essential to the dissolution of BiOCl can be shown by carrying out the demonstrations using nitric acid in place of hydrochloric acid. The cloudy solution of 0.1 g of BiOCl in 10 mL of water becomes clear after addition of 1.6 mL of 6 M HCl but is still cloudy after addition of 4.5 mL of 6 M HNO3. If one now adds 0.8 mL of 6 M NaCl to the solution to which nitric acid was added, all solid dissolves. That both H3O+ and Cl᎑ ions must be involved can be seen by the fact that addition of 4.5 mL of 6 M NaCl alone also does not produce a clear solution, but addition of 0.5 mL of 6 M HNO3 after the NaCl causes all solid to dissolve. Similar results are obtained starting with BiCl3. Since BiONO3 is only slightly soluble, the same tests were performed starting with BiONO3 instead of BiOCl. The
results were essentially the same, confirming the importance of both the H3O+ and Cl᎑ ions. The behavior BiOCl is similar to that of SbOCl, which we have used for many years to help students understand coupled equilibria (2). That increasing the concentration of Cl᎑ should lead to BiOCl dissolving seems counter intuitive when only the dissolution equation BiOCl(s) BiO+ + Cl᎑ is considered. However, with a little assistance, students quickly recognize the greater sensitivity of the reaction forming BiCl4᎑ to the concentration of Cl᎑ because 4 ions are involved (as well as the dependence of H3O+). This leads directly to an appreciation of the importance of the exponents for each species in the equilibrium constant expression. Literature Cited 1. Martell, A. E.; Smith, R. M. Critical Stability Constants; Plenum: New York, 1976; Vol. 4, p 111. 2. Craig, N. C.; Carlton, T. S.; Ackermann, M. N.; Schoonmaker, R. C.; Renfrow, W. B. Composition, Reaction, and Equilibrium: Experiments in Chemistry; Addision-Wesley: Reading, MA, 1970. Martin N. Ackermann Department of Chemistry Oberlin College Oberlin, OH 44074
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Improving Introductory Chemistry In four fine articles in the July 1996 issue of this Journal (J. Chem. Educ. 1996, 73, 617–636), Gillespie, Spencer, and Moog, as a contribution of the Task Force on General Chemistry, point out how improvements can be made in the presentation of some of the procedures of the introductory chemistry course. These articles are also important in that they provide a particularly clear illustration of the way mainstream introductory chemistry has adopted the values or attitudes of modern society and modern culture. This societal or cultural emphasis, as Neil Postman has argued in several important books (see, for example, Technopoly: the Surrender of Culture to Technology; Alfred A. Knopf: New York, 1992), is on shared information and experiences, a sort of collective knowledge, rather than on individual understanding. In introductory chemistry courses, we enable our students to manage the information and procedures that we identify with introductory or general chemistry. We try to get students into chemistry. There is a kind of top-down learning in that students are presented with expert systems with which they can successfully manage the material of these courses. Modern courses contrast with the traditions of chemistry in which the focus is on getting chemistry into students. In this tradition, students are helped to build their own understanding of this corner of the physical world. Students must be able to fit the new material onto their own mental framework and then build their own understanding. What is learned must make sense to each student in a kind of bottom-up learning. Gillespie, Spencer, and Moog provide us with some clear, well-founded information-management procedures. Data from electron impact and photoelectron spectroscopy lead to information on shells, energy levels, and electron configurations. The electron domain concept provides a “simpler, quicker, and less mysterious route” to bonding and structure. A table of AVEE (average valence electron energies) and an equation for the partial electron charge on an atom can be used to characterize several electrical properties of atoms in molecules. A variety of reaction thermodynamic properties are obtained from tables of “atomization” properties by just a single type of calculation. In mainstream introductory chemistry courses, students must learn successful information-management strategies. Changes or improvements that help them do so are consistent with the values and goals of the course. The well thought out routes and strategies presented by Gillespie, Spencer, and Moog fit these requirements and can be incorporated into these courses. By using the guidance provided by these articles, teachers will be more successful in getting students to manage these parts of introductory chemistry courses. The presentation of the material of mainstream introductory courses can be improved, as the articles by Gillespie, Spencer, and Moog so well illustrate. But what do existing courses do for students and what will “new and improved” courses do? Not Much! Students can learn the material, but it makes little or no sense to them (Barrow, G. M. J. Chem. Educ. 1991, 68, 449). As a result, what is learned is of so
little intellectual value and so little practical value that improving mainstream introductory chemistry courses in this way is an exercise in elevato ad absurdum. These courses do not, as students quickly discover, consist of extensions of their own experiences. Students are not constructing their own mental models that, in the traditions of science, would serve as a basis for extending their own understanding. Once set adrift in the world of information management, students cannot make sense of the information they learn to process. Students gain no added respect for reason and for the individual—the heart of the tradition of modern science. Each student is deprived of the opportunity to experience that wonderful growth of confidence that comes when familiar bits of natural phenomena fit into a general scheme that they have had a hand in constructing or accepting. Many other courses that have little intellectual value earn their place in the curriculum by the usefulness of the skills that are taught. These mainstream chemistry courses do not. What is learned fits so precariously in each student’s understanding that it cannot be brought to bear in subsequent chemistry-type courses. In the job market, for example, having recipes for solving limiting-regent problems, displaying ground-state atomic electron configurations, and so forth, is not much of an advantage. Much time and effort are wasted in trooping all those high school, college, and university students through introductory or general chemistry courses. But the real tragedy of the mainstream courses, even with much improved presentations, is that so many promising students are turned away from science and maybe even away from serious attention to academic studies because of the senseless, even if fine-tuned, rigamarole of introductory chemistry courses. Some students would appreciate introductory chemistry courses that helped them build their own understanding and were stimulating and useful. Some would profit from courses based on information technology. Some would respond to courses that focused on industrial, environmental, and social issues. But we cannot get to the first of such alternatives and probably not to any other significant alternatives by “improving” existing courses. Gordon M. Barrow 2095 Thompson Rd. Gabriola, B.C., V0R 1X0, Canada
The authors reply: While we very much appreciate Barrow’s positive comments on our articles, we feel that our articles do more than just enable students to manage the information in the course. Our aim in these articles is to help students obtain a better understanding of the four topics we discuss. For example, we show that electron configurations are not just derived in some mysterious way from magic numbers (quantum numbers) for which no justification is (or can be) given at this level, but that they can be obtained from experimental data following the scientific method. We believe this approach improves a student’s understanding of electron configurations and removes some of the mystery that is associated with them in the student’s mind. Similarly, the use of enthalpies and en-
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Letters tropies of atomization enables a student to have a real and rather easily understood picture in his/her mind when doing calculations involving thermodynamic properties. We feel that we are helping students to build their own understanding rather than just learning to “manage” the information we provide. In contrast to Barrow’s comments that “students are not constructing their own mental models”, we developed or adapted the approaches outlined in these four articles with just such a goal in mind. At Franklin & Marshall College, for example, for the past three years we have been experimenting with a group learning approach in a few sections of general chemistry based on having students process information. In order to do this—that is, have students use data to build models rather than using data to support models that are thrust at them—we needed to find new ways to introduce the basics of chemistry. Students in our classes construct their own knowledge, based firmly on the learning cycle and constructivist ideas of teaching. We have consistently argued that memorization and the application of algorithms should not constitute a significant part of the course. Our attempt was to, yes, improve the existing course but to do much more. We wanted to provide for student understanding and the involvement of students in their own educational process. Nevertheless, we agree wholeheartedly with Barrow that improving existing courses is not enough and, indeed, is not the ultimate objective. As we have advocated for several years now, a drastic revision of the general chemistry course is urgently required (1, 2) to make the course more relevant to the student’s experience and needs and to help students understand chemistry rather than forcing them to manage— which we suspect most often means memorizing—the information we provide. An understanding of chemistry will be useful to them years later in whatever branch of science, medicine, or technology they ultimately work, whereas the information that they learned to manage in the general chemistry course will have been long forgotten. Although reforming the general chemistry course has been the subject of exhaustive discussion for at least the past ten years, except in a few isolated institutions little change has in fact occurred. In face of this considerable resistance to change, some of the reasons for which have been recently discussed (3), we feel any improvements we can make are better than none. And we anticipate that some of them at least will be useful in the revised course whatever form it takes. Barrow and ourselves have the same ultimate objective, and we hope he will contribute to the ongoing discussion more details of the type of course he is advocating. Literature Cited 1. Gillespie, R. J., J. Chem, Educ. 1991, 68, 192–194. 2. Lloyd, B. W., Spencer, J. N. J. Chem. Educ. 1994, 71, 206–209. 3. Gillespie, R. J., CHED Newsletter, Fall 1996. Ronald J. Gillespie Department of Chemistry, McMaster University Hamilton, ON L8S 4M1, Canada J. N. Spencer and R. S. Moog Department of Chemistry, Franklin and Marshall College Lancaster, PA 17604-3003
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Demystifying Introductory Chemistry The series of articles “Demystifying Introductory Chemistry” in this Journal (1) bring up what has bothered me for thirty years of teaching chemistry at all undergraduate levels. This is the confusing of properties of atoms and elements that has been perpetuated throughout these years in every textbook, journal article, and handbook that I have used or read. I have continually corrected this confusion for my students. To eliminate this confusion is very important for beginning students and also for the integrity of the principles of chemistry. This confusion is using “element” instead of “isolated atom in the gas-phase” in terminology involving ionization energies, electron affinities, average valence electron energies, atomic electronic configurations, etc. For example, the first, second, and etc. ionization energies are for isolated atoms in the gas-phase of a given element. To say or write “ionization energy for sodium element” is conceptually and experimentally incorrect. As stated this means the energy to remove an electron from solid sodium metal (more correctly, surface of sodium metal) which is the “work function”. The sodium atoms in sodium element are, of course, bonded together so the ionized electrons are coming from molecular orbitals and not atomic orbitals. When one is using ionization energies in the context of atomic structure, one is using, both conceptually and experimentally, “ionization energy of an isolated sodium atom in the gas-phase” and one should state this correctly. Throughout textbook chapters on atomic structure, “element” is used instead of “isolated atom in the gas-phase”. Then when the student reaches chapters on electrons and bonding in actual elements, confusion increases tremendously. If we are going to “demystify introductory chemistry”, we had better get our terminology correct. Also, since we can now “see” (chemists have mentally “visualized” for decades) and “move around” atoms, it is time to start at the “atom” in textbooks and build up to molecules, elements, and compounds in a simple to complex fashion instead of using older procedures such as macro elements and etc. down to “theoretical” atoms. Literature Cited 1. Gillespie, R. J.; Spencer, J. N.; Moog, R. S. J. Chem. Educ. 1996, 73, 617–636. Douglas Rustad Department of Chemistry Sonoma State University Rohnert Park, CA 94928
The authors reply: Rustad is right. Chemistry is difficult enough for beginning students without our making it unnecessarily more difficult. All of us, particularly textbook writers, should be much more careful to avoid the kind of confusing statement that he draws attention to. We know what we mean but we do not ask ourselves often enough if the beginning student will know what we mean. Both in the article he refers to and in my own textbook Atoms, Molecules, and Reactions: An Intro-
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duction to Chemistry (Prentice Hall 1994) there are tables headed “Ionization Energies of the Elements”. The text talks about isolated atoms but it would remove any possible confusion if in the heading, and elsewhere when necessary, the term element was qualified, as he suggests, by writing “Element (isolated atom in the gas phase)”. As I have said before, the general chemistry text is badly in need or reform (1). An important aspect of this reform would be to make a serious attempt to eliminate all the confusing statements such as Rustad refers to. The term element can be confusing in several other ways too. It can mean the pure bulk substance, one or more isolated molecules of the element, one or more isolated atoms of the element, or the element in a combined form as when we talk about the sodium in sodium chloride. An effort to remove these confusions from the textbook needs to be a cooperative effort because individually we often do not realize that terms that we use, and statements that we make, although perfectly clear to us, may not be clear to the student. I also agree with Rustad that now we can “see” atoms, and because it is essential for students to think in terms of atoms, the teaching of chemistry should begin with atoms, as I do in my own text. We can first define an element as a substance all of whose atoms are of the same kind and then later explain that this means that they have the same atomic number, and then build up to combinations of atoms or molecules.
Demonstration Explosion Last week I did a demonstration that produced a serious explosion. After putting methanol in a big glass carboy and rotating the carboy to build up some methanol vapor, I lit the mouth of the carboy. What normally happens is a "jet engine" effect out of the mouth of the carboy. In my case, the carboy exploded. Two polycarbonate blast shields were shattered and glass was blown as far as 15 feet away. I was not seriously cut and bruised, but had I not been using the two blast shields, I would have been severely injured. At this time, I am not sure what caused the explosion. I have done this demonstration around one hundred times with no problem using the exact same amount of methanol and technique. I think it is important to get the word out that this demonstration may be more dangerous than previously thought. I would also welcome any hypotheses concerning what caused the carboy to explode. Charles "Skip" Lee, Chemistry Instructor Rock Falls High School Rock Falls, Illinois 61071
[email protected] Literature Cited 1. Gillespie, R. J. J. Chem. Educ. 1997, 74, 484–485. R. J. Gillespie Department of Chemistry McMaster University Hamilton, ON L8S 4M1, Canada
We could not agree more with the comments of Douglas Rustad. Although our article was not written for students, his point is well-taken. Chemists have a vocabulary that is often “mystifying” to students. We know what we mean but the student encountering these terms or symbols for the first or even subsequent times is surely puzzled. We use the same symbol to convey the concept of an atom, a collection of atoms, or the element. There are many places in introductory chemistry where our terminology is not clear. For example, we refer to an impure compound, but a compound can't be impure. Some terminology is so ingrown it is unlikely that it can ever be eradicated. We speak of balanced equations, but an equation must be balanced or it isn't an equation. Instead of asking students to write a “balanced equation” we should ask them to write a chemical equation that describes the reaction of interest. We interchange mass and weight. We use terms familiar to chemists in ways that have particular meaning to chemists but rarely stop to consider that without years of experience in chemistry these terms could indeed be mystifying. J. N. Spencer and R. S. Moog Department of Chemistry Franklin and Marshall College Lancaster, PA 17604-3003
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