
How are electrons
distributed in an atom?
Electron Configuration

 Shorthand method for representing electrons’
distribution in orbitals within subshells
 All orbitals have the same energy level—digenerate
 Orbitals– mathematical expressions of probability of
electron’s location
 Electrons occupy orbitals in a way that gives
LOWEST energy state
Orbital Diagrams

 Visual representation of electron configuration
 Represents electrons’ spins (, )
Example “spdf
notation”


2
3s
Example “expanded
spdf notation”


6
3p
Pauli Exclusion
Principle

 Only 2 electrons occupy each orbital
 Electron spins MUST be opposite/paired when 2
electrons occupy the same orbital
 +1/2, -1/2

Hund’s Rules

 Developed by Friedrich Hund
 Creates the most stable electron arrangement
 Based on electron spin
Hund’s Rules cont.

1) One electron MUST occupy each orbital BEFORE
electrons are paired in the same orbital.
2) Electrons added to subshell with the same spin
(+1/2, -1/2) so each orbital has one electron.
Using the periodic table
 Period numbers = principal quantum number of
valence shell electrons
 Subshells fill with electrons at different regions
within periodic table (s section, p section)

Ex. 1 Nitrogen

Ex. 2 Cr

Writing Electron
Configurations for Ions

 Anions
 Gaining electrons from valence shell
 Achieving full valence shell with gain
 SO, write appropriate noble gas configuration
 Cations
 Losing electrons from valence shell
 Main group metals---write noble gas configration
 P-block---cannot write noble gas configuration
Ex. S

-2
+3
Ga
Ex.

Electrons and Magnetic
Field

 Diamagnetism
 Exists when all electrons are paired
 Slight repelling from magnetic field
 Paramagnetism
 Exists when at least one electron is unpaired
 Unpaired electron attracted to magnetic field
Practice

1) Ca +2
5) Sb
2) P
6) Li +2
3) Mn
7) Br -1
4) S-2
Homework

 Practice Problems and textbook problems
pp. 331-333 #3-4, 21, 23, 25, 27, 29, 35, 41-42