P444: Energetic order of atomic orbitals: Observed values versus textbook stories

Author: W. H. Eugen Schwarz, Theoretical Chemistry Center, Department of Chemistry, Tsinghua University, Beijing 100084, China (Presented by Michael W. Schmidt, Ames Laboratory, Iowa State University)

Co-Author: Jun Li, Tsinghua University, China

Date: 8/5/14

Time: 10:35 AM10:55 AM

Room: MAN 123

Related Symposium: S36

Any ‘atoms first in chemistry’ approach must begin with the shapes and energies of the atomic orbitals. The concept of atomic orbits and their energetic order in poly-electronic atoms was introduced by N. Bohr. In his widely published Noble lecture, he not only derived the spin-orbit split term energies Enlj of the core and valence shells of ‘all’ light and heavy elements from the diverse observed spectra, but also plotted the orbital orders correctly already since 1922. The common order is rather hydrogen-like. However, due to the strong shielding of the nuclear charge by the inner core electrons at the beginning of periods n = 4, 5, 6, 7 (i.e., the heavy s-block elements), the common hydrogenic orbital order (n-2)f < (n-1)d < ns becomes inverted to ns < (n-1)d < (n-2)f. Remarkably, the chemical community chose the latter exceptional orbital energy order as the standard order, and teaches it up to present times. The textbooks must then combine it with several further inconsistencies of logic and facts and ‘exceptions from the rules’ to achieve agreement with chemical deductions. We propose to teach the observed standard orbital order, which then does not require further mental maneuvers. For undergraduate students, one may just communicate this fact in general chemistry. At the graduate level, the exceptional orbital order in the first two groups of the periodic table can be theoretically explained by the nuclear shielding effect, which must be mentioned anyhow in connection with the s-p-d splitting.