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.

P445: Atoms first, but correctly: Consistent explanation of the periodic law

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: 11:10 AM11:30 AM

Room: MAN 123

Related Symposium: S36

The operational definition (Boyle 1661) and discovery of the mass of chemical elements (Lavoisier et al. 1787) called for a systematic ordering that was empirically achieved in the 1860s (Meyer, Mendeleyev, and many others). A logical explanation of the basic Periodic Law of chemistry is still missing today, despite the availability of sufficient theoretical concepts (since Bohr 1922, et al.). The periodicity of the elements is based on the periodicity of valency, i.e. of the number of holes in a strongly-bound atomic noble-gas core shell, or of weakly-bound valence electrons energetically high above the core. The particularly different groups of the halogens, noble gases, and alkali metals form the backbone of the periodic system. The origin is the particularly large orbital energy gap above the low-lying 1s2 and np6 noble-gas shells, and below the much more loosely-bound (n+1)s, nd and (n-1)f valence shells. The appearance and position of the large orbital energy gap is due to the shielding of the nuclear attraction of the more and less core-penetrating valence orbitals. We propose to teach the observed orbital energy gaps. One can then, in the 2010s, for the first time rationalize a) the structure and backbone of the periodic system, and b) why the p6 shell is special whereas the s2, d10 and f14 shells are not. For undergraduate students, one may just communicate this fact, while the theoretical explanation through nuclear shielding arguments can be taught to graduate students and beyond.