Electron Configuration
Arrangements in the Quantum
Mechanical Model
Electrons are important because they
determine how elements will bond and react
with other elements.
Bohr versus the REAL
THING:
Quantum Mechanical Model of
the Atom (today’s model)
–Electrons DO occupy energy levels
but their path is not definite; it is
unknown!
–Heisenberg Uncertainty Principle It is impossible to know the exact
location and speed of an electron
Electron Configuration
Tells us the probable location of an
element’s electrons.
If YOU were an electron….
• Atom = school
• Energy Levels = hallways
• Sub levels = classrooms
• Orbitals = desks
• Electrons = students
ENERGY
Energy Levels –
give the general distance from the nucleus
Energy Levels –
give the general distance from the nucleus
Sublevels – the areas within the energy
levels where the electrons are found.
s
“sphere”
p
“propeller”
d
“double
propeller”
f
“flower”
(too complex)
Color YOUR Periodic Table!
1
1s
2
2s
2p
3
3s
3p
4
4s
3d
4p
5
5s
4d
5p
6
6s
5d
6p
7
7s
6d
7p
1s
6
4f
7
5f
Electron Configuration
2
1s
Energy Level –
Principal Quantum #
(possibilities are 1-7)
# of electrons
s: 1 or 2
p: 1-6
d: 1-10
f: 1-14
Sublevel
(s, p, d, or f)
If you add up all of the electrons, it should equal the atomic #
of the element
Orbital Diagram – visually diagrams
the location of the electrons.
• Only 2 electrons fit in each orbital.
• Example:
↑↓
↑↓
↑↓ ↑↓ ↑↓
↑↓
↑↓ ↑↓ ↑↓
↑↓
↑↓ ↑↓ ↑↓ ↑↓ ↑↓
1s
2s
2p
3s
3p
4s
3d
You must draw each orbital, even if it is empty.
Energy
Levels
Sublevels
# of
Orbitals
Maximum
# of electrons
1-7
s
2
6
d
1
3
5
2-7
3-6
p
4-5
f
7
10
14
↑
↑
4p
3 Main Rules of Orbital Diagrams
1) Aufbau Principle –
Electrons fill orbitals of the lowest energy
first.
The location of the
electron with the
highest energy will
help you identify
the element
Example:
1s2 2s2 2p6 3s2 3p6 4s2 3d6
Highest energy electron
Element = Iron
Aufbau Sequence – sublevels fill row by row,
from left to right
1
1s
2
2s
2p
3
3s
3p
4
4s
3d
4p
5
5s
4d
5p
6
6s
5d
6p
7
7s
6d
7p
1s
6
4f
7
5f
3 Main Rules of Orbital Diagrams
2) Hund’s Rule –
Each orbital must have 1 e- before any can
have 2 e↑↓ ↑
2p
↑
3 Main Rules of Orbital Diagrams
2) Hund’s Rule –
Each orbital must have 1 e- before any can
have 2 e3) Pauli Exclusion Principle –
Only 2 e- can be found in each orbital, and
they must be spinning opposite directions.
( )
Draw the orbital diagrams for the
following elements:
• Fe –
• O–
Now write the electron
configuration:
• Fe
1s2 2s2 2p6 3s2 3p6 4s2 3d6
• O
1s2 2s2 2p4
• Pb
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2
4d10 5p6 6s2 4f14 5d10 6p2
Lewis Dot Structures – show only
the valence electrons.
• Valence Electrons – outershell electrons
• You can count valence electrons by looking
at the group #.
1
8
1
1s
2
3
2
2s
2p
3
3s
3p
4
4s
3d
4p
5
5s
4d
5p
6
6s
5d
6p
7
7s
6d
7p
VALENCE electrons!
6
4f
7
5f
4
5
6
7 1s
Lewis Dot Structures – show only
the valence electrons.
• Valence Electrons – outershell electrons
• You can count valence electrons by looking
at the group #.
Remember Hund’s Rule when drawing Lewis
Dots BUT draw all “s” electrons (the first 2)
together on the right.
Lewis Dot Structures
Carbon:
C
4 valence e= 4 dots
Lewis Dot Structures
Carbon:
Chlorine:
C
Cl
7 valence e= 7 dots