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.
• Filling orbitals with electrons
according to Aufbau follows
the pattern shown due to a
mixing of energy levels
shown below
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
Relationship between orbitals and structure of Periodic
Table
• Last electron for K is in 4s1 orbital
• Last electron for Pt is in 5d8 orbital
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
5f
• fill orbitals following Hund’s Rule
Eg. For lead: abbrev. elec. conf. is
82Pb = [54Xe]6s24f145d106p2
orbital diagram is:
[54Xe] __ __ __ __ __ __ __ __
6s
4f
__ __ __ __ __ __ __ __
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
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)
__ __ __ __ __
-2
-1
0
5d
+1 +2
Eg.
(6,1,1,-1/2) would represent the orbital electron:
6p6 __ __ __
-1 0 +1
Rn
Eg.
(5,2,2,+1/2) would represent the orbital electron:
5d5 __ __ __ __ __
-2
Re
-1
0
+1 +2
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.
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.