Laurence E. Strong
Earlham College Richmond, Indiana
Facts, Students, Ideas
M o r e than a half million1 students are introduced to chemistry each year in high school. If a Ph.D. degree represents a rough index of arrival a t the other end of the chemical academic road, about one per cent will complete the trip begun many years before in high school. Certainly the vast majority, however, will complete their study of chemistry by the end of one year's exposure. What is it that the chemistry instructor should hope to accomplish in that one year that will be worth a year of effort on his part and on the students' part? How is this t o he done both for the potential professional chemist and for the literate layman? How, in short, does the chemistry t ~ a c h e r escape being a baby sitter? My title is intended to suggest a strategy. Students can be introduced to modern chemistry as a study in the interrelation of facts and ideas. I t is to this dynamic interaction that modern chemistry owes its vigor and development. Indeed, the study of chemistry in such a fashion can give an important insight into the "tactics and strategy of science." The particular order of the words in the title can also suggest a sub-strategy which some instructors find important: that of keeping the student in thc middle and thus always susceptible t o being thrown off balance by an astute academic ploy. A most important strategy is that of getting the student t o find himself able to bring facts and ideas together. I n this the student operates as a sort of middleman. There are those who insist that experiment is the heart of chemistry while others will argue loudly for theory. Our experience is, however, that unreflective experimentation is unrewarding just as idle speculation is merely childish. To escape both these difficulties, classroom and laboratory facts and ideas should be interwoven, so that each contributes t o the student's developing understanding, for too often the laboratory serves only to illustrate what is already known to the student. I n ordinary jargon, laboratory work that is instructive should be a puzzle or a problem t o be solved bv the student t h r o u-~ hobservation and logic. The Chemical Bond Approach (CBA) Project2J is an Presented before the Division of Chemical Education, 140th Meeting of the ACS, Chicago, September, 1961. Based on work supported with grants from the NSF to Earlham College for the CBA Project. K. E., AND OROURN, E. J., "Offerings and Enroll'BROWN, ments in Science and Mathematics in Public High School," U.S. Dept. of Health, Education, and Welfare, Dept. of Education, Bulletin No. 5. ' LIVERMORE, A. H., AND STRONG, L. E., J. CHEM.EDUC.,37, 209 (1960). L. E., AND WILSON,M. X., J. CHEM.EDUC.,35, 56 STRONG, (1958).
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attempt to develop an introductory chemistry course which presents modern chemistry to beginning students. The presentation is intended to give students a preliminary understanding of what chemistry is about, rather than simply an encyclopedic collection of chemical reactions and laboratory techniques, or a mere overview of diverse conclusions held by chemists today. Such a course must he an organized one in which the pattern reflects the structure of the discipline itself. Since conceptual schemes play a major role in the organization of chemistry today, the organization of a course in chemistry is best based on conceptual schemes. This paper, then, will attempt to sketch how such an organization can be exploited and also some of the problems that it raises. Black Boxes and Mental Models
Beneath all our attempts a t scientific understanding is the problem in seeing in various phenomena something which can serve to tie one event to anot,her. William James has said that for a baby the universe is a "great, buzzing confusion." In what way do we grow out of this kind of babyhood? To illustrate this problem in epistemology CBA confronts the neophyte student of chcmistry with a "black box." The course argues that a black box is not only symbolic of an attitude toward science but, even more to the point, chemicals themselves are, for the chemist, black boxes. In its initial presentation the black box is an actual box with a loose object inside. The student studies it by shaking, rolling, twisting, and observing the various reactions produced within the box. From this study he attempts to match its behavior with deductions from a model he proposes for the contents of the box. The only essential ground rule is that the box cannot be opened for a direct view of the contents. Even if the artificial black box is eventually opened at the conclusion of the experiment, the point is made that a chemical system can never be opened for visual examination of the structures it contains. Chemicals are opened only in one's imagination. There are a number of imaginative mental models already available to chemists which provide ways of reasoning about the behavior of matter. Basically these mental models are devices for carrying out logical operations. Their value is judged not by appeal to their reality, but by the extent to which they provide logically satisfactory arguments and produce conclusions consistent with experimental data. The trouble with so much of the writing for modern introductory chemistry is that the models and the chemicals are treated as equals. CBA chemistry attempts, therefore, to present
several of the most powerful mental models presently available to chemists, and to demonstrate how they are effective for interrelating descriptive information. In fact, the title of the project refers to one such mental model, chemical bonds, with the implication that chemicals have structures. That substances can be thought about in geometrical terms is a very old proposition. It was discussed by the Greeks and probably even earlier by the Hindus. Modem science has found this view of great utility. A particularly useful feature is the idea that materials are made up of atoms held together by forces or bonds. Thus when we consider the difference in the hardness of diamond and graphite, there is no possibility of appealing to differences between the atoms since these must be assumed to be the same in the two different forms of carbon. The fact that both water and copper (11) sulfate are colorless yet their chemical mixture is blue can hardly be traced back directly to the atoms themselves. I n the logic of the atomic model, one must ascribe the colors of most aggregates of atoms to a feature developed as the result of structure or the bonds between atoms. But static models seem incapable of fostering a logic capable of dealing with such things as changes in state, and in particular the gaseous state. Closely interrelated with structure and order is the idea of disruption and unlimited expansion. For this concept, the kinetic molecular theory model is introduced and developed. Supplementary to the models of structure and of disruption is the model of energy. With these three logical developments available, the student can explore a variety of chemical systems. Let us look briefly a t each of these models to see how it may serve the purposes of an elementary course. Structure Models
Classroom trial thus far indicates that students are much intrigued by structural models and ths logic they suggest. A particularly simple way of introducing these ideas is being exploited in the CBA course. This is the so-called "charge cloud model" due largely to Dr. George Kimball.4 I n brief, this assumes that electrons behave as if they are electrically charged clouds. For simplicity, the cloud is assumed to be equivalent to a sphere with a uniform distribution of charge. Its only parameter is, then, its radius. Wave mechanics suggests the cloud will have a kind of kinetic energy which varies inversely with the square of its radius. Each cloud will have a potential energy determined by its position in relation to other electrons and atomic nuclei. Such an electrostatic potential energy will vary in inverse proportion to the cloud radius. Charge clouds overlap and the inherent potential energy relations involved are subject only to one important limitation, which is essentially a modification of the Pauli exclusion principle. By this limitation only two electrons may merge to occupy a single region of space. One way of expressing the deductions from such a model is by a two-term equation for the energy of a KIMBALL, G. E.,AND LOEBL, E. M., J. CHEM.EDUC.,36,233 (1959).
stable assembly of electrons and atomic nuclei.
Here R is the radius of a charge cloud while K , and K2 are calculated constants whose values are determined by fundamental constants descriptive of electrons and nuclear charges. Since a stable configuration is the one of minimum energy, we need to see how each term is to be made small. Our usual convention in electrostatic energy discussions, of course, leads to the smallest energy, quite often being the one represented by the largest negative number. In the first term is the wave mechanics or kinetic energy relationship for the charge cloud; it gets smaller as the volume of the electron gets larger. Electrostatic features of the model are in the second term which is made smaller by electron interaction with nuclei, while it becomes larger as the result of electrons interacting with electrons and nuclei interacting with nuclei. I n geometrical terms, then, an assembly of atoms will he most stable when the electrons can be as large as possible, consistent with keeping electrons away from each other but close to the nuclei, while keeping nuclei away from each other. Out of these geometrical relations we also draw the additional conclusions that a spherical two-electron cloud will be of particular advantage, while five two-electron clouds in a tetrahedral arrangement around one nucleus will be next most advantageous. Various other geometrical shapes will also turn up in appropriate situations. What all of this produces is a strategy by which we first view matter as inherently electrical, and then turn the electrical relations into a geometrical analogy. Out of this come directly the so-called "Rule of Two" and "Rule of Eight" which so usefully coordinate a great variety of compounds. One can also develop a strategy for understanding chemical reactions. Electrical structure models for atoms and molecules involved in a chemical reaction suggest that the change occurs in order to separate nuclei of high charge, to spread out electrons, and to surround nuclei with electrons more effectively. A fairly simple example of the application of these conclusions to chemical systems is provided by considering why H F transfers a proton to NHa to produce a set of charged ions out of uncharged molecules. I n NH2, a proton experiences less repulsion from the nitrogen nucleus than does a proton close to a fluorine nucleus in HF. For the same reason, neon atoms do not acquire protons simply because of the tremendous repulsion involved if a proton tries to get too close to the neon nucleus. Much of the chemistry of the isoelectronic series-methane, ammonia, water, hydrogen fluoride, and neon-can be interrelated by such a logical argument. After all, shouldn't studentsof the Br#nsted theory of bases a t least ask, "Why is not neon with four unshared pairs of electrons a powerful base?" Kinetic Molecular Theory
The neatness of structural theory is marred only by the realization that it won't work. Matter itself exists as solid, as liquid, and as gas. Moreover, one of these forms can usually be converted experimentally to Volume 39, Number 3, March 1962
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another. Structure models by themselves provide no means for such a contingency. A model must therefore be introduced to handle the logic of different physical states. Ready for this facet of experience is the model called kinetic molecular theory. This, by building into our structure system random and continual motion, provides for changes of state in response to changes in temperature. While the development of a student's ability to deal with gas law problems provides some limited experience with kinetic molecular theory, it is necessary to keep in mind that the theory plays a much more important role in our overall logic of the nature of matter. It is this logic to which the student needs to be exposed. Out of the logic comes the major conclusion that chemical reactions can also occur as the result of changes in randomness. Thus in assessing reaction possibilities, not only electrical energy decreases but also organizational changes must be considered. Until both have been properly evaluated, no conclusive answer can be given to reaction feasibility. This state of affairs can be summarized by the well known thermodynamic relation AF = AH
- TAS
where we now interpret the symbols so A F is a measure of reaction feasibility; AH is a measure of interaction energy changes among the atoms of the system and TAS is a measure of organizational energy change. Laboratory
Are the usual laboratory sessions of any value to the students? Presumably this can only be answered affirmatively in terms of the extent to which the laboratory work increases the student's grasp of chemistry. Ideally, the student should have a better understanding of chemistry when he leaves a laboratory session than he had when the session began. Does he? Consider an exercise as respectable as the preparation of copper(1) sulfide. Usually this is presented to the student as a demonstration of the Law of Definite Proportions. Now can we in good conscience claim that there is any logical connection between what the student observes and the generalization he studies? I t seems most unlikely. Of course, one can point out that the experiment is a fraud since the system does not lead to a stoichiometric compound anyway. But putting aside such a snide remark, there is still no real revelation in the laboratory that a compound is formed whose composition is essentially independent of changing circumstances. A proper experiment needs to confront the student with a reasonable variety of circumstances from which he can extract one constant feature, composition. Rather than bore you with a long discussion of a system with which this can be accomplished, I'll only refer you to Experiment 7 in the third edition of the CBA laboratory guide. But there is some good food for thought hidden in the reaction of copper and sulfur. A little calculation with handbook data or some moderately careful laboratory measurements reveal that the volume of copper(1) sulfide is less than the arithmetic sum of the volumes of the unreacted copper and sulfur from which the product is formed. This, in fact, seems to be the result in many exothermic reactions. In the light of what has 128
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been said about structure models, the student can be challenged to speculate on the reasonableness of such an experimental result. Unfortunately, the experiment with copper and sulfur is unsatisfactory for this purpose because the volume decrease is only about 3%. Other systems look much more promising as a basis for simple experimentation which can provoke a good deal of thought. The key to the logic is, of course, that since many reactions occur because the products compared to the reagents provide more compact packing of electrons around nuclei, then the products should be less voluminous thau the initial reagents. Energy
For chemists, the need for the concept of energy arises because of the common observation that most chemical changes are accompanied by temperature change. Energy is the most powerful imaginative model available for dealing with this relationship between temperature change and chemical change; and presumably energy is the'idea through which we deal logically with our repeated observations of the interconnectedness of change. I11 addition, energr provides a tremendonslg useful bookkeeping device to apply to reactions. Any attempt to fit our imagination to laboratory experience is best handled through some kind of bookkeeping arrangement. For chemists, -useful bookkeeping systems are provided by the concepts of mass, energy, and , charge, while geometry by contrast does not provide ,' any straightforward bookkeeping for sieehnd&ape. Stmudents should Lave some of the experiende of using the concept of energy i11 class and laboratory. There , are relatively simple experiments available.' More use .should be made of them in the laboratory. Again they should be used to set the student thinking and not ' simply to check up on hahdbook data.. In the CBA : laboratory program, students deal with the diffprenm between heat a,ad temperature (Exp. 9), calibration of a . calorimeter (Exp. 13), heat of vaporization (Exp. 14), thermal energy in relation to stoichiometry (Exp. 21), " and a heat of formation (Exp. 22). , '
Chemistry as a Liberal Art
What I have been describing is, in its essentials, an intellectual discipline. For this, chemistry primarily is. I t is one of mail's great attempts to convince himself that the universe is not a great buzzing confusion. Should not this be the burden of introductory chemistry? Perhaps this is even the true profession of the chemist rather than being a mere technician for the utilitarian features of society. I t is true that there are some deep troubles with such a broad view. Hustou Smith, Professor of Philosophy of M.I.T., wrote, "Frontier thinkers are no longer sure that reality is ordered and orderly. If it is, they are not sure that man's mind is capable of grasping its ~ r d e r . " ~ But my argument has been that our imaginative ideas of structure, disruption, and energy are logical devices, not necessarily the reality behind events. With these ideas, chemists do find it possible to int,errelate usefully the events that take place in test tube and beaker. Whether this points to an ultimate order Saanr, H., Saturday Evening Post, 234, 28 (1961).
or not is perhaps less important than the ability it gives to explore still further the physical universe. Of course, it may be that the ultimate order revealed by chemical study is that the behavior of chemicals in test tubes is essentially the same as the behavior of the chemicals that are our brains. There is then as much order, or lack of it, in the world external to us as there is in our heads. CBA chemistry is still under development, still full
of difficulties, still beinq revised. Our evidence strongly suggests, however, that such an approach is exciting to students and teachers in the classroom and the laboratory. And when all is said and done, the only real reason you and I pursue chemistry is because of its heady excitement, because it liberates our minds for exploration beyond any well-worn rut. To the extent that chemistry does indeed liberate the mind, it is aliberalart.
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