1. Quantum Theory & Structure of the Periodic Table
The orbitals that the electrons in an atom occupy can be illustrated in the form of an electron configuration.
This is simply a tool used to help us understand and predict chemical bonding.
Rules for filling up orbitals with electrons:
1. Aufbau Principle – fill lowest energy levels first before proceeding to next level.
2. Hund’s Rule – electrons must occupy single orbitals before pairing up with other electrons.

2. Filling orbitals with electrons
according to Aufbau follows
the pattern shown due to a
mixing of energy levels
shown below

3. Electron Configurations
Electron configuration for potassium:
19K = 1s22s22p63s23p64s1
Electron configuration for platinum:
78Pt = 1s22s22p63s23p64s23d104p65s24d105p66s24f145d8
Abbreviated configuration for platinum:
(choose nearest Noble gas in previous row and continue electron configuration)
78Pt = [54Xe] 6s24f145d8

4. Relationship between orbitals and structure of Periodic Table
Last electron for K is in 4s1 orbital
Last electron for Pt is in 5d8 orbital

5. Energy Level or Orbital Diagrams
These are representations that show the orbital orientations that the valence electrons occupy in an atom.
(use abbreviated electron configurations to show electrons in orbitals).
Eg. For uranium: abbrev. elec. conf. is 92U = [86Rn]7s25f4
orbital diagram is: [86Rn] __ __ __ __ __ __ __ __
7s f
fill orbitals following Hund’s Rule

6. Eg. For lead: abbrev. elec. conf. is
82Pb = [54Xe]6s24f145d106p2
orbital diagram is:
[54Xe] __ __ __ __ __ __ __ __ __ __ __ __ __ __ __ __
6s f 5d 6p
Orbital diagrams are useful in predicting magnetism in elements
paramagnetic – unpaired electrons present in orbitals that are attracted to a magnetic field
diamagnetic – no unpaired electrons present in orbitals and are not attracted to a magnetic field

7. Relating Quantum Numbers to Electron Configurations
The last electron(s) in the valence orbitals of an atom are responsible for chemical reactivity. The set of 4 numbers can identify the energy associated with these electrons. Eg. 19K ends in 4s1, quantum #’s are: (4,0,0,+1/2) 78Pt ends in 5d8, quantum #’s are: (5,2,0,-1/2) __ __ __ __ __ d

8. Eg. (6,1,1,-1/2) would represent the orbital electron: 6p6 __ __ __ Rn Eg. (5,2,2,+1/2) would represent the orbital electron: 5d5 __ __ __ __ __ Re

9. Exceptions to Electron Configurations
There is a certain stability that is seen in an atom when it can achieve a totally full (or exactly half full) orbital of electrons.
This can be seen in the configurations of elements in the d4 and d9 families.
Eg. 47Ag = [36Kr] 4s23d9 is really [36Kr] 4s13d10 since the ‘d’ orbital is completely filled and the ‘s’ orbital is now half filled with electrons.

10. Certain exceptions to the trends of ionization energy and electron affinity in the Periodic Table (learned about in grade 11) can be explained by the stability of these electron orbital arrangements.