Lecture #7
Text sections: 12.13, 14, 15
  CURMUDGEON GENERAL'S WARNING. These "slides" represent highlights from lecture and are neither complete nor meant to replace lecture. It is advised not to use these as a reliable means to replace missed lecture material. Do so at risk to healthy academic performance in 09-105.
Lecture Outline The Periodic Table (continued)

Sizes

Isoelectronic species

Ionization energies
Continuing our discussion of effective nuclear charge, we move on to the sodium atom with Z=11. The electron configuration is 1s22s22p63s1 or [10Ne]3s1. The valence electron is the 3s electron, which can be promoted to the excited states illustrated. Note that because of the great penetrating probability of an electron in the s-orbital, the 4s becomes lower in energy than the 3d orbital. The yellow (590 nm) emission line of an excited Na is indicated.
A schematic view of the arrangement of orbitals for one-electron systems
A schematic view of the arrangement of orbitals for many-electron systems
A mnemonic device for reproducing the order of orbitals according to increasing energy for many-electron systems.
Here are configurations for four atoms (and one ion). Each atom is characterized now by the number of s and p valence electrons...7...accounting for the similar properties of these "halogen" elements in their chemistry.
Exceptions appear in the filling scheme owing to an "extra stability" associated with completely filled subshells and half-filled subshells.
Electron configurations for Group I elements, the alkali metals. Sizes are determined by "n" and Zeff for the last electron.
The trend in radii for the alkali elements
Generalization of the trend in radii for a Group (column) in the Periodic Table
The trend across a row of the Periodic Table, where "n" is fixed for the main groups, is determined by Zeff, which increases. (Main groups are the s-block and the p-block elements.)
The general trend across main group rows is a decrease in size from left to right in the Periodic Table
The trends for "Main Group" ("Representative" or "s- and p-block") Elements.
 Isoelectronic systems have the same electron arrangements. The example here can be generalized to any similar sequence.  
 All the rules we've seen about the electron configuration filling sequence apply to both atoms and ions of the main group elements (s-block and p-block) and to the transition metal atoms. For transition metal ions, the sequence changes!  
 To explain why the 4s-3d order shifts when ionizing an element, we resort to a familiar example, sodium, and look at its (excited) 4s and 3d excited one-electron orbital energies. Zeff refers to that for atomic Na in this illustration  
 The states are graphed at their proper energies, but we've indicated the effective nuclear charges drived from these energies. If an inner electron is removed completely from Na, we can estimate that the Zeff for the outer valence electron osup by approximately 1.  
 The other changes are similarly estimated and -- lo and behold -- the 3d is now more stable than the 4s orbital. When these are filled as in the transition metal ions, the 3d is occupied while the 4s is empty.

(Keep in mind that the values shown are just the effective nuclear charges, but the position of the energy level corresponds to E as determined by n and Zeff.)

  
For this course, 09-105, at CMU, we will adopt the general rule that notes a different configuration decision about transition metal ions as opposed to atoms and main group ions.