Electronic Structures of Molecules XIV. Linear Triatomic Molecules, Especially Carbon Dioxide

Publisher Website
Google Scholar
Abstract
Electron configurations of the triatomic linear molecules CO2, BO2—, N2O, CS2, COS, NO2+, N3—, NCO—, NCS—, BeF2, HgCl2, ClCN, ClBO, etc. are given. The electron configurations probably of all of these are formally analogous, or formally identical in the case of isoelectronic molecules, such as CO2, N2O, BeF2, etc. The BeH2 type is also touched on. The relations of the electron configurations to ordinary valence conceptions are discussed, and the importance in practice of cases intermediate between those represented by simple valence formulas is emphasized. The usefulness of the present method in giving a simple description of intermediate cases of all sorts (between double-bonded A=B=C and single-bonded A–B–C, between homopolar A–B–C, etc., and heteropolar A—B++C—, etc., between A–B≡C and A+=B=C—, etc.) is brought out. The ionization potentials of CO2, CS2, N2O, HgCl2 are interpreted in relation to their electron configurations. The minimum potentials of CO2, of CS2, and of HgCl2, respectively, correspond to removal of a nearly nonbonding pπ electron of the O, S, or Cl atom. It is shown that the ultraviolet emission bands of CO2 probably represent transitions among four electron states of CO2+ which correspond to the four lowest ionization potentials of CO2; if this interpretation of the bands is correct, the CO2+ molecule has a linear form in these states. According to this interpretation, the four lowest potentials of CO2 are probably 13.72, 17.09, 18.00, and about 21.5 volts (or perhaps 14.63, 18.00, 18.91, and about 22.4 volts). These values are based on identification of the CO2 Rydberg absorption series limit at 18.00 volts with one of the potentials. The minimum potential above obtained (13.72 or perhaps 14.63 volts) differs considerably from the accepted value 14.3 or 14.4 volts based on electron impact. Radicals N3, NCO, NCS, and molecules NO2 and others are also touched on. A reason for the triangular form of NO2 is found in the fact that linear NO2+ (like CO2 or N2 or NO+) is composed of a stable system of closed shells, while the final electron added to make neutral linear NO2 would have to go into a probably fairly high energy antibonding orbital (as in NO). Evidently the actual NO2 avoids this by being triangular. Ultraviolet absorption spectra of CO2, CS2, and COS are also considered.