ELECTRON
CONFIGURATION
LABEL THE SUBLEVELS (1S, 2S, …)
PRACTICE 1
Sulfur ( ______ electrons)
Electron configuration:
a. How many E.L. in an atom of S have e-?
b. How many e- are in the 2nd energy level?
PRACTICE 1 ANSWER
Sulfur ( __16____ electrons)
Electron configuration: 1s2 2s2 2p6 3s2 3p4
a. How many E.L. in an atom of S have e-?
Three
b. How many e- are in the 2nd energy level?
8 (2+6 - just add the superscripts)
PRACTICE 2
• Iron (________ electrons)
• Electron Configuration:
a. How many E.L. in an atom of Fe contain e-?
b. b. How many e- in 3rd energy level?
PRACTICE 2 ANSWER
• Iron (___26_____ electrons)
• Electron Configuration:
1s2 2s2 2p6 3s2 3p6 4s2 3d6
a. How many E.L. in an atom of Fe contain e-?
Four
b. How many e- in 3rd energy level?
14 (2+6+6)
NOBLE GAS NOTATION
(“SHORT HAND”)
• This is used to represent electron configurations of
noble gases using bracketed symbols.
SHORT HAND PRACTICE
• Use shortcut method to write e-configurations for Ar
and P.
• Ar: [Ne] 3s2 3p6 or [Ar]
• P:
[Ne] 3s2 3p3
GROUND STATE VS EXCITED STATE
Ground State – The most stable, lower energy arrangement of
the electrons in an atom.
(The electrons are in the correct order of orbitals.)
Excited State – When atom gains energy the electrons jump
into higher energy level.
Example: Identify the following atoms and describe their
state. (Ground or Excited)
a. 1s22s22p63s23p64s13d104p3
b. 1s22s22p63s23p64s23d104p5
c. 1s22s22p63s23p64s23d104p65s24d8
ELECTRON CONFIGURATION
A. Aufbau Principle – Electrons enter orbitals of lower energy first.
1s2, 2s2, 2p6, 3s2, 3p6, 4s2, 3d10, 4p6, 5s2, 4d10, 5p6, 6s2, 4f14, 5d10, 6p6, 7s2, 5f14, 6d10, 7p6
(or use PT)
B. Pauli Exclusion Principle – No more than __2______ electrons can
occupy the same orbital AND they must spin in opposite directions.
(Opposite spins help hold e- in an orbital by creating magnetic
attraction.)
C. Hund’s Rule – Orbitals of equal energy must EACH have _1____
electron with the same ____spin______ before any orbital is occupied by
a 2nd electron.
ELECTRON CONFIGURATION