Modern Theory: A Tool in Teaching Elementary College Chemistry HARRY H. SISLER and CALVIN A. VANDERWERF University of Kansas, Lawrence, Kansas
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URING recent years students who have had trainmg UI . such sciences as mathematics and physics, and in their practical application in various branches of engineering, have shown a growing tendency to criticize adversely undergraduate courses in chemistry. The chief criticism which is heard is that, in these courses, the student is required to memorize a mass of more or less interesting and more or less related facts, but that insufficient attempt is made to bring out broad underlying principles in terms of which the student may evaluate and correlate the facts he has learned. The student of better than average intelligence is not to be satisfied by the study of a large body of almost purely descriptive data which is explained and correlated, if a t all, by empirical rules of thumb. A frank appraisal of most of the more recent chemistry texts, which presumably are in line with the subject matter taught in the majority of chemistry courses a t the present time, does not provide a sound basis for our rising in righteous indignation and labeling these criticisms as false and unjust. During the last century and earlier, i t was impossible for chemistry courses to be integrated in terms of broad fundamental principles, for such principles had not been developed or, a t best, were but imperfectly understood. This, however, is no longer true. Of course, our theories are and always will be imperfect, but recent developments in the fields .of physics and physical and theoretical chemistry have led to the discovery of a few broad, basic concepts in terms of which we believe it is possible to organize thoroughgoing, integrated, and internally consistent courses in inorganic and organic chemistry a t the elementary college level. Arguments for the use of modern principles in organizing subject matter of chemistry courses in our college curricula have been presented before1 and presumably have borne some fruit, for most chemistry teachers and textbook writers pay lip-service to modern theory and "modernize" their texts by putting in a chapter or two on atomic structure and the nature of the chemical bond. There is nothing wrong with this, of course. The fault is that the modernization stops there. For in many courses and texts we find that, far from providing a basis for the explanation and correlation of the descriptive material, these more recently developed fundamental concepts are rarely mentioned in the rest
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1 Wnnnam. "The need of modernizing the general J. C ~ MEnuc., . 12, 11-16 (1935).
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of the text or course. Instead, the old empirical explanations and rules of thumb are employed. The failure of most chemistry teachers to exploit to the fullest the possibilities of modem theoretical principles is illustrated by their insistence on retaining the old Mendeleeff form of the Periodic Chart even though the historical table is obsolete in the light of present day atomic theory. This point will be elaborated upon below. Teachers of chemistry whose presentations follow the lines of older classical texts will argue that the introduction of intangible concepts confuses and burdens the student unnecessarily and can be accomplished only by some slighting of the more practical aspects of the subject which they purport to emphasize. We should never lose sight of the fact that a theory, as Sir Joseph Thomson was wont to remark in his classes in the Cavendish Laboratory, is "a tool, not a creed," and certainly the theoretical tail shonld not be allowed to wag the pedagogical dog. But i t is the authors' contention that the use of modem concepts clearly stated and repeatedly applied greatly simplifies the study of practical chemistry, providing as it does, the continuous thread upon which otherwise isolated and incoherent facts may be strung to weave the fabric of a logical, unified, and integrated course. Such a treatment of the subject certainly is better calculated to develop the student's reasoning power and to encourage his scientific spirit of inquiry than is a mere memorization of facts. A student trained in such an approach carries with hi a broad picture into which new facts may be fitted, a ready tool by use of which additional knowledge may be interpreted, and a clear signpost pointing in the direction of further truth, rather than a loose bundle of unrelated facts which slip from his memory one by one. A clear-cut distinction must, of course, be drawn between fact and theory, but the student may soon be led to realize that shonld a present theory be discarded in favor of a better one, i t will probably have been the present concept which has fashioned the better implement and that the latter can best be understood and utilized by those who were thoroughly familiar with the former. As stated above, arguments for the use of modern principles in organizing subject matter in chemistry courses in our college curricula have been presented before. There is still, however, ample room for improvement of our chemistry courses along this line. It is the purpose of the authors to illustrate by a few
concrete examples how they attempt to apply modern theoretical principles in their teaching of elementary inorganic and organic chemistry. INORGANIC
The authors believe that a thorough, workable knowledge of the Periodic System and the relationship of this system to the electron configurations of the various atoms is an absolute necessity for any student who aspires to attain any thorough understanding of chemical science. This means that in considering the structures of the atoms the discussion must not be restricted to the elements in the first two series of the Periodic System, as is commonly the case, and that such subjects as the nature of the electronic configurations of the transition elements and the rare earths shall not be avoided but shall be discussed, and their significance to the building of the whole series of atoms explained. This has been done rather well on the graduate level by Rice2 and by Pauling8. It is up to the instructor to present the essentials of this matter in terms that elementary students can understand and apply. The authors have found that not only is such a presentation possible hut that i t pays dividends in increased student interest and understanding. Acceptance of this thesis immediately brings into focus the diiculties both from a scientific and a'pedagogical standpoint that follow from the use of the Mendeleeff style of Periodic Chart. Even if we use the Periodic Law as an empirical principle, making no effort to interpret it in terms of its theoretical basis, the Mendeleeff arrangement of the elements is not satisfactory. At least it is not unless we are willing to pass over such obvious incongruities as the placing of copper, silver, and gold in the same group as the alkali metals, or the placing of manganese in the same group as the halogens. Moreover, a glance a t a plot of some such atomic property as ionization potential against atomic number will show that the Mendele& table is not even in empirical accord with the periodicity of the properties of the atoms. Thus, not only from a theoretical but also from a practical point of view, a modern "long" form of the Periodic Table such as the Werner-Evans Table is much to he preferred over the Mendeleeff form. In this more recent form each of the "long series" is placed in a horizontal row and the transition elements fall in the middle part of the table. Moreover, such families of elements as that composed of copper, silver, and gold, and that containing chromium, molybdenum, tungsten, and uranium, are now in distinct columns in the chart rather than being placed as subgroups in the same column with another family of elements to which, in most cases, they bear only the most formal resemblance. Such a table fits itself admirably to the correlation of theory and descriptive material. 'RICE. "Electronic Structure and Chemical Binding," McGraw-Hill Boak Co., New York, 1940, pp. 84-103. VAULING. "Nature of the Chemical Bond." Cornell University Press. Ithaca. New York, 1940, pp. 24-30.
Another phase of chemical science which we belie7 to he a fundamental requirement for any intelligex consideration of the properties of matter is the stud of the various types of chemical bonds and the reasor for their formation. What is the nature of the force involved in the various types of combination whic atoms undergo? Of course, it goes without sayin that there is a limit to the depth to which this subje( may be explored a t the elementary level. This limi however, is seldom reached or even approached. Mor textbook writers content themselves by saying thz atoms may combine by the transfer of electrons to fon ions or by the sharing of pairs of electrons to fotr covalent compounds, and by giving a few examples c each. The authors believe that in addition the di! cussion should be carried to the point of inquiry int those characteristics of the combhing atoms whic determine whether the resulting compound will b ionic or covalent. Why, for example, is sodiur chloride an ionic compound, whereas aluminum chlc ride (anhydrous) is largely covalent? Or why d potassium, rubidium, and cesium ions form few, i any, complex ions, whereas beryllium, boron, an, aluminum form such complexes readily? The answers to these and other pertinent question which inevitably arise are to be found, to a certail extent a t least, in the consideration of the gaseou ionization potentials and the ionic radii of the elements Without going into any complicated discussion of th, manner in which these quantities are measured, we ma: present the ionization potential simply as a measur, of the energy required to remove an electron from a atom of the element under consideration, and the iouia radius as a measure of the size of the ion thus formed Admittedly such data do not tell the whole stoty, bu they do give qualitative and, in many cases, roughlj quantitative explanations of some of the combinin5 properties of the atoms. Once these terms have beer introduced and defined, the change in these quantitie: as we pass from element to element in the Periodic System is noted, explained, and correlated with othe) properties of the elements and their compounds a: they are studied. It is pointed out, for example, thai as one passes from left to right in a given series, thc ionization potentials of the elements show a genera increase, while the correspondimg ionic radii show : general decrease. These facts are explained in term! of the increasing nuclear charge on the atoms. Thus the experimental fact of the increasingly nonmetallic character of the elements and the increasing acidity ol their hydroxides is related directly to the change ir the structure of the atoms. Why, for example, i! sodium a strong reducing agent and chlorine a strong oxidizing agent? We no longer have to answer such a question with vague generalities. Why is sodium hydroxide a strongly basic hydroxide, whereas silicon hydroxide (silicic acid) is acidic? Simply because the sodium ion is a large ion which has only a small attraction for electrons and permits the breaking off of the hydroxyl ion, whereas silicon forms a small, highly
charged ion with high attraction for electrons, and does not permit hydroxyl ions to be released. In fact, the hydroxyl group is in this case so polarized that the release of protons is facilitated. The explanation of the more highly ionic character of sodium chloride than of aluminum chloride is exactly analogous. The lower ionic radius and the greater charge of the aluminum ion results in a much stronger attraction for a pair of electrons on the chloride ion than is the case in sodium chloride. It is further pointed out to the student that within a given family of the Periodic System, ionic radii increase and ionization potentials decrease with increasing atomic weight. (The exceptions to this rule in the case of the last member of some families due to the lanthanide contraction may or may not be explained, a t the discretion of the instructor.) The increasingly metallic character of the elements of a given family with increasing atomic weight, and the increasing basicity (or decreasing acidity) of their hydroxides and oxides, are readily explained by means of these facts. Furthermore, such points as the aforementioned low conductivity of beryllium chloride as compared with that of barium chloride are readily explained in terms of differences in ionic radii and, in some cases, differences in ionic charge. In this connection there is still another point which may be removed from the realm of pure memory work for the student and placed upon a reasonable theoretical basis. Instead of the student's memorizing that ferrous hydroxide is more basic than ferric hydroxide, that nitrous acid is a weaker acid than nitric, that hydrated stannic oxide is less basic than stannous hydroxide, the following general principle is stated and explained. Whenever an element forms oxides and hydroxides in more than one valence state, the oxide and hydroxide in the higher valence state are more acidic (or less basic) than the oxide and hydroxide in the lower valence state. This rule is readily explained to the student by showing that as the positive charge on the ion increases, its attraction for the pair of electrons on the hydroxyl ion will increase, thus lowering the basic nature of the compound. Moreover, this same tendency increases the polarization of the hydroxyl group and results in an increased tendency to release protons. In nitric acid there is one more pair of electrons shared with an oxygen atom than in nitrous acid. The fact that the extra oxygen shares a pair of electrons which are held exclusively by the nitrogen in nitrous acid results in an increase in the effective positive charge on the nitrogen atom. Hence nitric acid is a stronger acid than is nitrous acid. Obviously, it is necessary to make frequent use of electronic formulas in the development of these ideas. It may be said, however, that after a short time the students no longer regard these formulas as mysterious hieroglyphics, but welcome them as aids to a rational interpretation of the properties of the compounds. Space, of course, does not permit anything approaching a complete summary of the multitude of phe-
nomena which are explained and correlated by the use of the newer theoretical concepts. The variable valences of the transition elements as related to incomplete inner electron shells, the relationship of complex ion formation to ionic charge and ionic radius, the change in stability of the hydrides of the nitrogen family with increasing atomic weight, are only a few examples. It is not only the more academic aspects 09 the subject which may be correlated with theoretical principles, however. Such practical points as the methods for reduction of the ores of the metals, their reactivity (and hence their corrosion resistance), the forms in which they are found in nature, the dates of their discovery, their usefulness in industry, and many other interesting and important facts are readily correlated with the electronic structures and the ionization potentials of the atoms, and hence, with their position in the Periodic System. It cannot be overemphasized, however, that, in order that the full benefit of the use of these theoretical concepts may be obtained two things are necessary. First, these concepts must be introduced as early in the course as is feasible, and second, they must be repeatedly and consistently applied and expanded in the explanation of the descriptive material presented throughout the course. It is with regard to this second point, particularly, that most elementary inorganic texts fall short. ORGANIC
The authors believe that the electronic theory, particularly in its application to the problems of organic chemistry, has now been developed to the extent that its intelligent use not only provides a precise physical basis for this branch of the science, but also serves to systematize and to simplify the study of the subject for the beginner. Concepts of the electronic theory as extended to a treatment of organic chemistry have been reviewed thoroughly on an advanced level by Johnson4. It remains the problem of the individual teacher to present the salient generalizations of the theory in simple, concise terms understandable to students in the elementary course, and in such a manner that knowledge of the theory will actually expedite the study of practical organic chemistry. The authors feel that most of the present elementary organic texts which employ the electronic theory as a basis for study confuse the issue either by presenting the theoretical concepts so that they are much less intelligible than the facts themselves, or by employing the theory in the explanation of complex reactions, not regularly studied in the usual first year course, and neglecting to apply it in many simple, obvious cases. In order that they may be used throughout, the fundamental principles of the electronic theory should be presented to the student a t the outset of the course. Two good examples, preferably dealing with compounds .JJOHNSON, "Modern Electronic Concepts of Valence." GILMAN, "Orpnic Chemistry." 2nd ed., John Wiley and Sons. New York, 1943, Vol. 11.1821-1942.
already familiar to the student, should suffice to illustrate each principle. Certainly it seems illadvised to point out how the concepts may be applied in a given homologous series until the question is made real to the student when the study of this particular series itself is undertaken. The basic theoretical material can be covered adequately in three lectures. It is admitted that the student may not master all of the theoretical concepts completely when they are first presented, but as he learns to relate them repeatedly to the study of practical organic chemistry which follows, he will soon appreciate their significance and will find in them a broad basis upon which to organize and about which to correlate many of the multitudinous facts of organic chemistry. Only by consistent use of the theory in the interpretation of organic phenomena can satisfactory answers be offered for the host of questions which inevitably arise in the mind of the thinking student. Reviewing principles already familiar to the students, the authors begin the theoretical treatment with a discussion of the nature of the atom with particular emphasis upon electronic structure, followed by a summary of the nature of the ionic, covalent, and coordinate covalent bonds. The question of maximum covalency is covered, especially with reference to hydrogen, and conditions favoring the formation of the hydrogen bond are pointed out. Next the phenomenon of resonance is introduced and the enhanced stability and shortened interatomic distances of resonating systems are stressed. In further development of concepts previously studied in the inorganic field, rules used in deducing the normal electronic structure of organic compounds are explained and illustrated by means of one or two typical compounds. Pauling's Electronegativity Scale of the elements is interpreted in the light of its application to the question of bond energies and the existence of covalent bonds with partial ionic character. The significance of an electrical dipole is explained, and physical and chemical properties, such as melting point, boiling point, solubility, solvent power, association of molecules, conductivity, activity, etc., of the three broad classes of compounds, ionic, nonpolar covalent, and dipolar covalent, are reviewed. Finally, the activation of molecules arising through the displacement of electrons in the molecule giving rise to centers of high electron density or low electron density, and the relationship of such activation to chemical reactivity are considered. The two sidiple forms of electron displacements, the inductive and electromeric, are not further subdivided. Inductive effects are defined as electron displacements arising from an unequal sharing of the electron pair of a covalent bond. The following series6 listing various atoms and groups in the order of decreasing relative inductive electron-attraction effect is given: F > C1 > Br > I > OCHI > NHCOCH3 > CsH6 > CH= MAN. "Organic Chemistry." 2nd ed., John Wiley and Sons, New Ymk. 1943, Vol. 11, 1844 and 1894.
CH2 > H > CHa > Primary Groups > Secondary Groups > Tertiary Groups. Electrometric effects are designated as electron displacements associated with unshared electron pairs or with multiple covalent bonds. Obviously the two effects may oppose each other, and the generalization may be drawn that in cases where a hetero atom such as nitrogen, oxygen, or sulfur (in its lower covalent state) is joined to an unsaturated system by means of a single' link, the electromeric effect of the hetero atom overcomes its inductive effect in giving direct impetus to a certain course of reaction. Simple examples of each type of effect are cited and the meaning of "nucleophilic" (nucleus-seeking), "electrophilic" (electron-seeking), "carbanion," and "carbonium" ion is explained. Following the treatment of the theoretical concepts outlined above, each homologous series is discussed in the following sequence of topics: (1) type formula and homology; (2) nomenclature; (3) synthesis and preparation; (4) type electronic formula; (5) physical properties of the members of the series; (6) chemical properties and reactions; (7) important individual members of the series with emphasis on unusual reactions, uses, and commercial applications. Wherever possible, the consideration of the type electronic formula is made the basis for the study of the physical properties and chemical reactions of the members of the series, and every possibility of systematization, correlation, and generalization based upon electronic structure is exploited. A review of all of the possible applications of the electronic theory in the interpretation and integration of the f a d s of organic chemistry cannot be attempted in the brief space of this paper, but a few specific examples may be cited. Differences in the physical properties such as boiling points, melting points, and solubilities among various homologous series as, for example, the hydrocarbons, the alcohols, the ethers, the aldehydes, and the acids are readily explained and predicted on the basis of the dipolar character of the molecules and the possibilities of association through hydrogen bond formation. The same principles may be applied repeatedly, as in a comparison of the physical properties of isomeric primary, secondary, and tertiary amines. Consideration of the relative inductive series suggests that the central carbon atom of propane represents a center of high electron density compared to the two terminal atoms; hence, the central carbon atom should be more susceptible to attack by an electron-accepting reagent such as chlorine. This conclusion is coniirmed by the actual vapor phase chlorination of propane, in which it is found that the rate of replacement of the two hydrogen atoms attached to the central carbon atom is 3.25 times that for the remaining six hydrogen atomss. The relative inductive effect table likewise provides a logical basis for the prediction of the comparative reaction rates of the alcohols. Since a proton will
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a Hnss and
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escape more readily as the electron density of the OHgroup is diminished, reactions which involve the release of the proton from the alcohol, such as salt formation and esterification, should proceed most readily in the case of primary alcohols, and least readily for the tertiary alcohols. On the other hand, for reactions such as the replacement of the OH-group by X , and dehydration in the presence of an acid catalyst, in which the capacity of the oxygen atom to act as a donor is the rate-determining factor, the order is reversed. The intriguing question concerning the mode of addition of unsymmetrical reagents to unsymmetrical multiple bonds loses much of the mystery in which it has long been shrouded when viewed in the light of modern electronic theory. In the general case of the olefins, for example, the assumption that all of the alkyl groups have an effect of electron-release relative to H-C predicts that the odd pair of electrons of the double bond should be displaced toward the carbon bearing the larger number of hydrogen atoms. Hence, in the normal addition of reagents such as H-X or X-OH, the center of high electron density, -X or --OH, should be attracted to the carbon atom bearing the smaller number of hydrogen atoms. Consideration of the electronic structure of vinyl chloride, furthermore, reveals that the electromeric effect associated with the unshared pairs of electrons about the chlorine atom tends to increase the covalence between chlorine and the carbon atom, and should result in the formation of an active form showing the odd pair of electrons of the double bond displaced away from the carbon atom bearing the chlorine. As a result, the center of high electron density in the adding reagent should be attracted to the carbon atom bearing the chlorine. In the case of acrylic acid, on the contrary, the electromeric effect associated with the carbonyl oxygen favors the displacement of the odd electron pair toward the a-carbon, and in the normal addition of H-X, the halogen atom should add a t the p-carbon. All these predictions are confirmed in actual practice. Without demanding a study of the minute mechanisms proposed for the hydrolysis, alcoholysis, ammonolysis, and aminolysis of the acids, esters, amides, acyl halides, and acid anhydrides, the general concepts of the electronic theory provide a simple means for the correlation of the characteristic reactions of the organic acids and acid derivatives. In every case, the fundamental reactions of this group may be viewed as a sharing of an electron pair between the electrondonor atom in the molecule H-Y and the electrophylic 0
ing for the low electron density of the carbonyl carbon atom and by inhibiting the rupture of the C-Z bond. In general, the typical addition reactions of the carbony1 group likewise involve the attachment of a nucleophilic group to the carbon of the carbonyl. The concept of negative groups has played a prominent role in the development of theoretical organic chemistry. In electronic terms, the "negative" group of classical organic chemistry is simply one wflich has a tendency to withdraw electrons from the carbon atom to which i t is attached. A negative group will enhance the proton-escaping tendency and diminish the electron availability in any system into which i t is introduced. Thus, acetic acid is a stronger acid than methyl alcohol, chloracetic acid is a stronger acid than acetic acid, acetamide is a weaker base than methylamine, and the amino group in aniline may be "protected" by acetylation. Strong negative groups exert a marked effect in facilitating the ionization of hydrogen atoms bound to carbon, as noted in enolimtion and tautomerism phenomena and in such basecatalyzed or basic condensation reactions as the aldol and Claisen condensations, the alkylation of acetoacetic and malonic esters, the Perkin, Knoevenagel, and Michael reactions, the halogenation of aldehydes and ketones, and condensations involving cyanacetic ester, cyanacetamide, and aliphatic nitro compounds. Each of these synthetically important reactions can be explained in terms of the removal by the basic catalyst or condensing agent of a proton from a carbon atom attached to one or more strongly negative groups, followed by the addition of the resulting carbanion to the electrophilic system involved in the particular synthesis. In the aldol condensation, for example, the second step would be the usual carbonyl addition, in which the carbanion plays the role of the nucleophilic reagent. Consistent use of the electronic theory is admirably adapted to the teaching of organic chemistry, whether the aliphatic and aromatic compounds are studied together or separately. In the authors' courses, the latter method is followed, and i t is most gratifying to observe that to students schooled in the application of modern theory, the "benzene problem" is no problem a t all. In fact, such a student, given a clear picture of the electronic structure of aromatic nuclei, can actually predict aU of the distinguishing chap acteristics ordinarily associated with aromaticity. Diminished unsaturation and pronounced tendency to the formation and preservation of type follow from the increased stability of the benzene system due to resonance. The low reactivity toward oxidizing agents and the tendency to undergo substitution rather than 4 addition reactions may be explained on the same basis. carbonyl carbon atom of the acid derivative, R-CThe inert nature of the halogen atom linked directly Z, with the simultaneous splitting out of HZ. to an aromatic ring, and the acidic strength of the For a given electron-donor system, ease of reactivity phenols, as in the corresponding cases of vinyl halides among the various acid derivatives varies inversely and aliphatic enols, results from the electromeric effect with the electromeric electron release of Z, an effect which tends to diminish reactivity both by compensat(Continued on gage 496)
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MODERN THEORY: A TOOL IN TEACHING ELEMENTARY COLLEGE CHEMISTRY (Conlinued from page 483)
associated with the hetero atom which tends to increase its covalence with the a-carbon of the unsaturated system. Instead of resorting to involved empirical rules in order to classify ortho-para-, and meta-directing groups, the student places the whole question on a definite physical basis. Consideration of the stable resonance forms of any substituted benzene molecule in which the substituent, through the inductive or electromeric effect, tends to release electrons to the benzene ring, reveals that the effect is transferred mainly to the ortho- and para-carbon atoms, facilitating an increase in the electron density a t these positions. Hence, such a suhstituent will exert an ortho-paradirecting influence toward electron-seeking reagents, and, in general, substitution will be facilitated. The
opposite effect of a group which tends to withdraw electrons from the ring is likewise transmitted chiefly to the ortho- and para-carbon atoms, which, as a result, are centers of low electron density. Such a group is therefore meta-directing, with deactivation. It cannot be argued that a few broad concepts of the electronic theory afford a complete explanation of all the phenomena of organic chemistry, or even that all of the facts fit logically into the general picture. But the authors contend that wherever the theoretical treatment serves as a tool to simplify, correlate, and lend physical reality to the phenomena of organic chemistry, i t should he employed. The challenge to take full advantage of all that modem theory has to offer in improving the teaching of beginning organic chemistry rests with the individual instructor.
