Letter pubs.acs.org/jchemeduc
Retire the Hybrid Atomic Orbital? Not So Fast Nivaldo J. Tro* Department of Chemistry, Westmont College, Santa Barbara, California 93108, United States ABSTRACT: This letter is a response to Alexander Grushow’s article, “Is It Time To Retire the Hybrid Atomic Orbital?” This letter suggests reasons for why valence bond theory and the associated hybrid atomic orbitals should continue to be taught in the chemistry curriculum. KEYWORDS: First-Year Undergraduate/General, Upper-Division Undergraduate, Physical Chemistry, Misconceptions/Discrepant Events, Covalent Bonding, Lewis Structures, MO Theory, Quantum Chemistry, VSEPR Theory, Valence Bond Theory
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n a commentary published in this Journal1 Alexander Grushow asks whether it might be time to retire the hybrid atomic orbital model of chemical bonding from our chemistry curriculum. He draws an analogy between hybrid orbitals and the phlogiston model of combustion and suggests that both models belong in the dustbin of outdated chemical theories. Grushow gives several reasons for why the hybrid orbital should be abolished, including that such orbitals do not exist and that they do not help our understanding of chemical bonding. In this short letter, I seek to clarify the bonding models currently taught in general chemistry courses and offer reasons for why we should continue to teach valence bond theory and the associated hybrid orbitals. Hybrid atomic orbitals are in themselves not a model for chemical bonding. Rather, they are a part of valence bond theory, which is in turn based on the quantum mechanical approximation technique known as perturbation theory. It is not clear from his commentary whether Grushow wishes to only eliminate hybridization from the chemical curriculum or whether he wishes to eliminate valence bond theory as a whole. However, because valence bond theory has limited value without hybridization, it seems that a call to eliminate one implies the elimination of both. In general chemistry, we have traditionally taught three models for chemical bonding: the Lewis model, valence bond theory, and molecular orbital (MO) theory. The progression is from simple to complex. The Lewis model represents electrons as dots and simply uses the octet rule to predict whether certain combinations of atoms are stable or not. Valence bond theory represents electrons as existing in atomic orbitals and applies perturbation theory to calculate how the energies of the electrons in those orbitals are affected when two or more atoms interact. Although hybridization of atomic orbitals is not necessary for valence bond theory in general (for example, unhybridized orbitals work for molecules such as H2 and H2S), it is necessary for molecules such as CH4 and H2O. Molecular orbital theory (which is based on the quantum mechanical approximation technique known as the variational method) calculates entirely new molecular orbitals that are properties of the molecule as whole. A lowering of the electronic energy in moving from the atomic orbitals of the isolated atoms to the molecular orbitals of the molecule results in bonding. This progression from simple to complex is, in my © 2012 American Chemical Society and Division of Chemical Education, Inc.
mind, a valuable exercise in chemical education, and removing the intermediate step eliminates an important part of this progression. Keeping valence bond theory as an intermediate step in understanding chemical bonding is important for several reasons. First of all, it helps to demonstrate the nature of models within scientific understanding. We teach our students many scientific models in general chemistry including the atomic theory, kinetic molecular theory, quantum mechanical theory, collision theory, and so on. These models, similar to all scientific models, are human constructs that help us to understand nature. Whether a particular model “exists” in the absolute sense of the word is a matter of scholarly debate. For example, kinetic molecular theory represents atoms as point particles. Should we throw out kinetic molecular theory because such “point particles” do not exist. No! To suggest that we throw out a model because it posits entities that “do not exist” betrays nearly all scientific models, because nearly all models contain elements that “do not exist”. So the question is not whether hybrid orbitals existI agree that they do notbut whether they are helpful in understanding chemical bonding. This question is similar to whether the octet rule or Lewis electron dots are helpful in understanding chemical bonding. No one suggests that electrons exist as dots arranged in pairs around an atom, but most chemists agree that depicting electrons this way in the Lewis model is helpful in understanding bonding. Chemical bonding is probably the only place in the curriculum where we teach multiple models for a single phenomenon. This exercise, if done correctly, teaches students a valuable lesson about the nature of scientific models. Valence bond theory should be kept, not only as a lesson in the nature of scientific models, but also because it helps us to better understand chemical bonding in ways that the other models do not. For starters, valence bond theory takes one step beyond the Lewis model and depicts electrons as existing in orbitals. The orbitals are localized atomic orbitals to be sure, but orbitals nonetheless, and this is a step in the right direction. Furthermore, depicting a chemical bond as the overlap between localized atomic orbitals is computationally sound. Valence Published: March 12, 2012 567
dx.doi.org/10.1021/ed2006289 | J. Chem. Educ. 2012, 89, 567−568
Journal of Chemical Education
Letter
bond calculations show that when two half-filled atomic orbitals overlap in space, the energies of the electrons in those orbitals decreases. Although these calculations are not explicitly carried out in a typical general chemistry course, they form the basis of the theory and can be easily understood. The idea that bonding lowers the energy of electrons is completely missing from the Lewis model, but is an integral part of valence bond theory. The idea of energy lowering is of course present in molecular orbital theory, but is more abstract and more difficult to grasp. Valence bond theory is also helpful because it can explain the rigidity of a double bond in ways that the Lewis model and valence shell electron pair repulsion (VSEPR) cannot. Although VSEPR can give us basic molecular shapes, it cannot explain why rotation about a carbon−carbon double bond is restricted. The overlap of half-filled, unhybridized p orbitals above and below the plane of a carbon−carbon double bond gives a nice mental picture of how a double bond differs from a single bond. Again, although MO theory can also give this picture, it does so only at a level that is seldom covered in general chemistry courses. I thank Professor Grushow for forcing us to more carefully examine the reasons that we teach valence bond theory and hybridization, even if it does not get the energies of the orbitals in methane correct. However, this deficiency is a relatively small price to pay for the degree of insight gained from valence bond theory. The Lewis model has more significant deficiencies, yet we continue to use it because of the insight it provides. The same can be said for many models that we continue to use. There may come a time when valence bond theory becomes obsolete, but I don’t think we are there yet.
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AUTHOR INFORMATION
Corresponding Author
*E-mail:
[email protected].
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ACKNOWLEDGMENTS Nivaldo Tro wishes to thank his colleague, Michael Everest, for helpful discussions on the topic of this response. REFERENCES
(1) Grushow, A. J. Chem. Educ. 2011, 88, 860−862.
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dx.doi.org/10.1021/ed2006289 | J. Chem. Educ. 2012, 89, 567−568
