Letter pubs.acs.org/jchemeduc
Retire Hybrid Atomic Orbitals? Roger L. DeKock* and John R. Strikwerda Department of Chemistry, Calvin College, Grand Rapids, Michigan 49546, United States ABSTRACT: We argue that localized bond orbitals can be used to interpret photoelectron spectra of molecules. KEYWORDS: First-Year Undergraduate/General, Second-Year Undergraduate, Upper-Division Undergraduate, Inorganic Chemistry, Organic Chemistry, Physical Chemistry, Misconceptions/Discrepant Events, Covalent Bonding, MO Theory, Valence Bond Theory
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form symmetric and antisymmetric stretching motions. These springs are best thought of as arising from localized bond orbitals. In summary, we agree with part of Grushow’s thesis and disagree with other parts. For example, we agree that the symbol sp3 as often used is simply an acronym for “tetrahedral”, and we needlessly complicate things for students by invoking a term that historically has a theoretical meaning. But we disagree on at least two points: (i) we do not feel that the opening two paragraphs that draw a comparison between phlogiston and localized orbitals is a valid comparison and (ii) we believe Grushow overreached by stating that such orbitals (i.e., localized) are inappropriate to interpret photoelectron spectra of molecules.
lexander Grushow wrote a commentary in a recent issue of this Journal entitled “Is it Time To Retire the Hybrid Atomic Orbital?”1 Grushow states “... photoelectron spectroscopic evidence indicates that hybrid atomic orbitals are inappropriate models for the description of electronic energies and electron density within a molecule.” Two sentences later he states “Despite significant experimental evidence and theoretical advances to indicate that hybrid atomic orbitals do not exist and do not appropriately describe molecular bonding, their description still permeates chemical education at many levels, and the model still finds its way into modern chemical literature.” Later in his article, Grushow states that he is not arguing against a localized bond orbital approach, but only against the concept of hybrid atomic orbitals. Yet, we believe that in the two sentences quoted above, the term “hybrid atomic” could be replaced by “localized”, and the same meaning would pertain with respect to the arguments that Grushow advances. We wish to comment upon the statement, made twice above, that hybrid atomic (i.e., localized) orbitals are not “appropriate”, in particular with respect to the ability to interpret photoelectron spectra of molecules. This question has been considered before in this Journal2 and in more than one monograph.3 In A Chemist’s Guide to Valence Bond Theory, Shaik and Hiberty include a section entitled “The Failure Associated with the Photoelectron Spectroscopy of CH4” (3a, pp 13−14). Two reasons are given why localized bond orbitals (LBOs) are appropriate to describe photoelectron spectra: First, as known since the 1930s, LBOs for methane or any molecule, can be obtained by a unitary transformation of the delocalized MOs. Thus, both MO and VB descriptions of methane can be cast in terms of LBOs. Second, if one starts from the LBO picture of methane the electron can come out of any one of the LBOs. A physically correct representation of the CH4+ cation would be a linear combination of the four forms that ascribe electron ejection to each of the four bonds.” The LBO description in chemistry is thus able to adequately account for the photoelectron spectrum of methane. Details of the relationship between the localized and delocalized molecular orbitals of methane have been published in this Journal.4 An LBO description lies at the heart of molecular mechanics in chemistry.5 In such a description of molecules, bond stretching is modeled by localized “springs”. The springs can couple together, for example, in a linear triatomic molecule, to © 2012 American Chemical Society and Division of Chemical Education, Inc.
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AUTHOR INFORMATION
Corresponding Author
*E-mail:
[email protected]. Notes
The authors declare no competing financial interest.
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REFERENCES
(1) Grushow, A. J. Chem. Educ. 2011, 88, 860−862. (2) Simons, J. J. Chem. Educ. 1992, 69, 522−528. (3) (a) Shaik, S. S.; Hiberty, P. C. A Chemist’s Guide to Valence Bond Theory; Wiley-Interscience: Hoboken, NJ, 2008. (b) Weinhold, F.; Landis, C. Valency and Bonding; Cambridge University Press: New York, 2005; page 125 briefly discusses the methane cation. (4) Bernett, W. A. J. Chem. Educ. 1969, 46, 746−749. (5) (a) Rappé, A. K.; Casewit, C. J. Molecular Mechanics across Chemistry; University Science Books: Sausalito, CA, 1997. (b) Boyd., D. B.; Lipkowitz, K. B. J. Chem. Educ. 1982, 59, 269−274.
Published: March 12, 2012 569
dx.doi.org/10.1021/ed200472t | J. Chem. Educ. 2012, 89, 569−569
