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How We Teach Molecular Structure to Freshmen Michael O. Hurst Department of Chemistry, Georgia Southern University, Statesboro, GA 30460;
[email protected] In many ways the study of chemistry is the study of the molecule. Our understanding of the molecule has evolved enormously since the days of Dalton, or even since the days of Lewis. Although they are not perfect, we now have welldeveloped theories of molecular structure, and this is reflected in the modern chemistry curriculum. Much of the time spent in a college general chemistry course is time spent studying atomic and molecular structure. Unfortunately, as in many other topics in general chemistry (1) the new theories have been simply added on to the old ones in the textbooks. The following general chemistry textbooks were examined. Umland; General Chemistry; West (1993) Brady, Russell, and Holum; Chemistry: Matter and Its Changes, 3rd ed.; Wiley (2000) Chang; Chemistry, 7th ed.; McGraw-Hill (2002) Jones and Atkins; Chemistry: Molecules, Matter, and Change, 4th ed.; Freeman (2000) Petrucci and Harwood; General Chemistry: Principles and Modern Applications, 6th ed.; MacMillan (1993)
ionization energies and electron affinities (4 ). Since these were developed in the 1930s other definitions have also been developed (5). For the most part general chemistry texts have not kept up with this. Many do not even agree with themselves. Five of the 10 texts examined (Chang et al., Brady et al., Jones and Atkins, and Reger et al.) display a figure of the periodic table with the Pauling numbers in it for the students to use when they need to work with electronegativity numbers, but give a mathematical definition of electronegativity based on Mulliken. One (Zumdahl) used Pauling for both the numbers and the definition, three (Robinson et al., Petrucci and Harwood, and Kotz and Treichel) used Pauling for the numbers and did not give the mathematical definition, and only one (Umland) used a modern definition by AllredRochow (6 ). The teaching of electronegativity is thus based on the historical development of the idea and not on modern understanding. The use of more than one definition for the same concept is clearly confusing to students. Since electronegativity is central to an understanding of bonding and polarity, our students are starting off at a major disadvantage.
Robinson, Odom, and Holtzclaw; General Chemistry, 10th ed.; Houghton-Mifflin (1997)
The Nature of the Chemical Bond
Brown, LeMay, and Bursten; Chemistry: The Central Science, 7th ed.; Prentice-Hall (1997)
Bonding is central to the understanding of molecules. All of the texts studied, and indeed virtually all general chemistry courses, discuss chemical bonds in detail. Covalent, ionic, and metallic bonds are discussed as very different entities and used to distinguish between molecular, ionic, and metallic substances. Then polarity is introduced as a property of covalent bonds, and it is explained to students that the difference between the different bond types is not as obvious as they had been told (perhaps a month earlier in a typical semester course). So the student is now wondering what these different bonds really are. Different texts deal, or do not deal, with this in different ways. Three of the texts studied (Zumdahl, Petrucci and Harwood, and Brady et al.) discuss bonds in terms of their percent ionic character. Only one of those (Zumdahl) defines what this is. One text (Robinson et al.) simply defined bond character by the type of compound it is; another (Umland) based it on boiling points and conductivity of the compound in question. All of the texts stated that the ionic nature of the bond, however defined, is related to the difference in electronegativity between the atoms in the bond. Five (Chang, Brown et al., Brady et al., Jones and Atkins, and Reger et al.) gave numerical criteria for this, 2.0 usually being the dividing line between covalent and ionic (two used 1.7). One text (Jones and Atkins) stated in one section that ionic bond character is based on the electronegativity difference, and in the next section stated that it is based on the polarizability of the atoms in the bond. Modern chemistry no longer looks at bond type simply in terms of electronegativity differences. Sproul has shown that both the average electronegativity of the bonding atoms and the electronegativity difference must be considered when
Zumdahl; Chemistry, 4th ed.; Houghton-Mifflin (1997) Reger, Goode, and Mercer; Chemistry: Principles and Practice; Saunders (1993) Kotz and Treichel; Chemistry and Chemical Reactivity, 4th ed.; Saunders (1999)
All of them offer molecular structure in two chapters in which Lewis dot structures and VSEPR theory, valence bond theory, and molecular orbital theory are discussed. This means that students who came out of high school chemistry knowing that a molecule is a very small thing made of atoms and that it moves, and possibly knowing how to write a formula, are given three different explanations of how molecules are put together. Is it pedagogically sound or even necessary to give them three theories? De Vos and Pilot have argued that we actually give them six theories or concepts in acid–base chemistry, that we do it for historical rather than pedagogical reasons, and that we integrate them poorly and that this is why students have such trouble with acids and bases (1). The 10 texts listed above were examined in terms of how they teach molecular structure. How the texts discuss various aspects of molecular structure is examined below. Electronegativity Electronegativity has been central to an understanding of molecules and chemical bonding since it was first proposed by Pauling. Pauling defined it using bond energies (2, 3). Later, Mulliken developed an alternative definition using
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determining bond type (7 ). Using only electronegativity difference to predict bond type gives incorrect results, but is still done in all the texts examined. Multiple Theories of Molecular Structure Lewis dot structures combined with VSEPR theory is of course an oversimplification of molecular structure. This becomes apparent to students when the texts discuss the expanded octets of elements such as phosphorus and sulfur, and especially during the discussion of resonance. All of the texts in question handle the problem by introducing two new theories, valence bond theory and molecular orbital theory. Why this is done is questionable. Students are very easily confused by multiple theories for the same phenomenon. It would seem much more reasonable to simply explain that Lewis dot structures and VSEPR theory are a simplification and that resonance and the expanded octet concept are ways to explain molecular structure without getting into very complicated theories that are frequently not necessary at the freshmen level. After all, these same students are being told in their physics classes that Einstein showed Newton to be incomplete, but they will not proceed into a detailed discussion of relativity unless they take higher-level courses. They will not need hybridization in general chemistry except for crystal field theory, which is much later in the course and for which the instructor will probably need to reteach or review it anyway. Molecular orbital theory is not used again in general chemistry. As for future courses, most of organic is based on Lewis dot structures and VSEPR theory, and the organic texts reteach valence bond theory. This discussion relates to how we teach general chemistry overall. Ronald Breslow has said that there are two types of chemistry: (A) chemistry that does not refer to the special properties of particular molecules and (B) chemistry in which the properties of different substances depend on their particular composition and structure (8). He feels that we need less of A and more of B to interest more of the huge number of students who stop with our introductory courses. Reducing the number of theories we discuss that cover the same phenomenon would seem to help do this. This is even more important in countries besides the USA where organic is taught in the freshmen curriculum. Of course, it will not be as complete as if they had VB and MO theory as well, but most professors of higher-level courses would rather have students who have learned one theory well than students who
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have been exposed to several theories but know none of them well. This is introductory chemistry, and higher-level theories are for higher-level courses. Of course true textbooks could leave both VB and MO theory in but in separate chapters, and professors could add or omit them as they think best. However, many in the profession feel that the texts are already far too large and expensive. When asked what he thought the major change in teaching chemistry had been during his career, Harold C. Brown replied that it was the size of the textbooks (Brown, Harold, 1985, personal communication). They are much larger now than when he was a student, but he didn’t think the students were learning any more. This was in 1985; the texts are even larger today. There is also a contradiction in the way valence bond theory is currently taught in general chemistry texts. All of the texts examined use VSEPR to explain molecular geometry and teach it as being a consequence of Lewis structure theory. Then, usually in the next chapter, they discuss valence bond theory and use VSEPR to explain the geometry of the orbitals. They specifically instruct students to use VSEPR to determine the geometry of an orbital and then use that to determine the hybridization. How can a theory be declared incomplete or incorrect and then be used to make predictions for its replacement? To mix the theories this way is again quite confusing to students. Conclusion Much of general chemistry is taught the way it is for historical reasons. This results in overlapping concepts that easily confuse students. Bonding theory and related concepts are central to an understanding of general chemistry and need to be taught in a clear and uniform manner. Literature Cited 1. de Vos, W.; Pilot, A. J. Chem. Educ. 2001, 78, 494–499. 2. Pauling, L.; Yost, D. M. Proc. Natl. Acad. Sci. USA 1932, 18, 414. 3. Pauling, L. J. Chem. Educ. 1988, 65, 375. 4. Mulliken, R. S. J. Chem. Phys. 1934, 2, 782. 5. Spencer, J. N.; Moog, R. S.; Gillespie, R. J. J. Chem. Educ. 1996, 73, 627–631. 6. Moeller, T. Inorganic Chemistry: A Modern Introduction; Wiley: New York, 1982. 7. Sproul, G. J. Chem. Educ. 2001, 78, 387–390. 8. Breslow, R. Chem. Eng. News 2001, 79 (31), 5.
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