Lewis Structure Representation of Free Radicals Similar to ClO

1360 Journal of Chemical Education • Vol. 84 No. 8 August 2007 • www.JCE.DivCHED.org At the level of introductory chemistry, oxidation numbers and formal charges are tools useful in analyzing the electronic structure of molecules. Oxidation numbers are assigned by giving a negative charge to the more electronegative atom in a diatomic molecule, while formal charges are assigned by assuming that electrons in a bond are shared equally between atoms, regardless of relative electronegativities. Electronegativity may be defined as the ability of an atom in a bond to attract electrons toward itself and is therefore a relative scale. Various scales are in use including those of Pauling, Mulliken, Allred–Rochow, and Sanderson (1–4). The formal charge of an atom in a molecule can be calculated by taking the number of valence electrons originally surrounding the unbonded atom and subtracting the sum of the unshared electrons and half of the bonded electrons. When choosing the best possible Lewis structures, the conventional approach is to select species with lower absolute values of formal charges. These concepts, in combination with more advanced concepts such as conjugation and coordination are sufficient to interpret most observed chemical processes, and practicing chemists generally do not attempt more rigorous quantum chemical calculations unless high levels of accuracy are desired. The simple ideas above, therefore, provide a useful framework for understanding chemistry, as evidenced by their continued inclusion in introductory chemistry (5). However, despite their success, these concepts can fail even for simple systems. Students often find cases in which the rules are ambiguous or contradictory, and understanding these exceptions is important pedagogically. Some interesting examples of ambiguous cases are the ClO radical and its isoelectronic analogues. In the analysis of the Lewis structure of ClO, regardless of the scale chosen, students will recognize that oxygen is the more electronegative atom from an examination of periodic trends. Therefore, if one assumes a single bond and that the location of the unpaired electron is determined by electronegativity, it ought to be situated on the chlorine atom, since the more electronegative oxygen atom will attract enough electrons to form an octet (Figure 1A). However, in this structure, the formal charges are +1 and 1 for Cl and O, respectively, when calculated by the method described above. If the unpaired electron’s location is determined by formal charge, it ought to be on the oxygen atom (Figure 1B), since this structure leads to a formal charge of zero on both atoms. Another possibility is an increase in the bond order to two, which lowers the formal charges to zero. In that case, the chlorine atom has an odd number of electrons and its valence shell exceeds an octet (Figure 1C). One might also consider a three-electron bond, as Pauling (6) did (Figure 1D). In that case, the formal charge would be +0.5 for oxygen and 0.5 for chlorine, and the halogen would still exceed an octet. We propose to use density functional computational methods, in conjunction with a variety of other computational analysis methods, to accurately describe the electronic structures for these species and attempt to interpret the results in a Lewis structure framework. The structures shown in Figure 1 represent possible resonance structures of ClO and its analogues. The interpretation will be based on the comparison of the computed results to the Lewis structures in Figure 1, with the goal of identifying which resonance structures most strongly contribute to the behavior of the system. In addition to the pedagogical value of the clarification of Lewis structures, halogen oxide molecules are of interest since they promote stratospheric ozone depletion and tropospheric ozone production (7–10). A simple Lewis structure interpretation of these species may therefore be of interest to students of atmospheric chemistry.