Electron
Configurations
Pre-AP
Quantum Mechanics
Better than any previous model,
quantum mechanics does explain how
the atom behaves.
Quantum mechanics treats electrons
not as particles, but more as waves
(like light waves) which can gain or
lose energy.
But they can’t gain or lose just any
amount of energy. They gain or lose
a “quantum” of energy.
A quantum is just an amount of energy that the electron
needs to gain (or lose) to move to the next energy level.
In this case it is losing the energy and dropping a level.
Atomic Orbitals
Much like the Bohr model, the energy
levels in quantum mechanics describe
locations where you are likely to find
an electron.
Bohr
Remember that orbitals are
“geometric shapes” around the nucleus
where electrons are found.
Quantum mechanics calculates the
probabilities where you are “likely” to
find electrons.
Atomic Orbitals
Of course, you could find an electron anywhere if you looked
hard enough.
So scientists agreed to limit these calculations to locations
where there was at least a 90% chance of finding an electron.
Think of orbitals as sort of a "border” for spaces around the
nucleus inside which electrons are allowed. No more than 2
electrons can ever be in 1 orbital. The orbital just defines an
“area” where you can find an electron.
What is the chance of finding an electron in the nucleus? Yes, of
course, it’s zero. There aren’t any electrons in the nucleus.
Quantum Numbers/Address
1
1s
What element has an electron configuration of 1s1?
Energy Levels
Quantum mechanics has a
principal quantum number. It is
represented by a little n. It
represents the “energy level” similar
to Bohr’s model.
Red
Orange
Yellow
Green
Blue
Indigo
Violet
n=1
n=2
n=3
n=4
n=5
n=6
n=7
n=1 describes the first energy level
n=2 describes the second energy
level
Etc.
Each energy level represents a
period or row on the periodic table.
Sub-levels = Specific Atomic Orbitals
Each energy level has 1 or more “sub-levels” which
describe the specific “atomic orbitals” for that level.
n = 1 has 1 sub-level (the “s” orbital)
n = 2 has 2 sub-levels (“s” and “p”)
n = 3 has 3 sub-levels (“s”, “p” and “d”)
n = 4 has 4 sub-levels (“s”, “p”, “d” and “f”)
There are 4 types of atomic orbitals:
s, p, d and f
Each of these sub-levels represent the blocks on the
periodic table.
Periodic Table and Energy Levels
Orbitals
s
p
d
In the s block, electrons are going into s orbitals.
In the p block, the s orbitals are full. New electrons are going
into the p orbitals.
In the d block, the s and p orbitals are full. New electrons are
going into the d orbitals.
What about the f block?
Orbital Shapes
Levels and Sub-Levels
Energy
Level
Sublevels
Total Orbitals
Total
Electrons
Total Electrons
per Level
n=1
s
1 (1s orbital)
2
2
n=2
s
p
1 (2s orbital)
3 (2p orbitals)
2
6
8
n=3
s
1 (3s orbital)
2
18
Complete
the chart
your notes as we
p
3 (3p in
orbitals)
6 discuss this.
d level (n=1)
5 (3d orbitals)
The first
has an s orbital.10It has only 1.
There are no other orbitals in the first energy level.
s
1 (4s orbital)
2
32
We callp this orbital
1s orbital. 6
3 (4pthe
orbitals)
n=4
d
f
5 (4d orbitals)
7 (4f orbitals)
10
14
Electron Configurations
What do I mean by “electron
configuration?”
The electron configuration is the
specific way in which the atomic orbitals
are filled.
Think of it as being similar to your
address. The electron configuration tells
me where all the electrons “live.”
Rules for Electron Configurations
In order to write an electron
configuration, we need to know the
RULES.
3 rules govern electron
configurations.
Aufbau Principle
Pauli Exclusion Principle
Hund’s Rule
Using the orbital filling diagram at the
right will help you figure out HOW to
write them
 Start with the 1s orbital. Fill each orbital
completely and then go to the next one,
until all of the elements have been
accounted for.
Fill Lower Energy Orbitals FIRST
Each line represents
an orbital.
1 (s), 3 (p), 5 (d), 7 (f)
High Energy
The Aufbau Principle states
that electrons enter the lowest
energy orbitals first.
The lower the principal
quantum number (n) the lower
the energy.
Low Energy
Within an energy level, s
orbitals are the lowest energy,
followed by p, d and then f. F
orbitals are the highest energy
for that level.
No more than 2 Electrons in Any
Orbital…ever.
Wolfgang Pauli, yet
another German
Nobel Prize winner
The next rule is the Pauli Exclusion Principal.
The Pauli Exclusion Principle states that an atomic
orbital may have up to 2 electrons and then it is full.
The spins have to be paired.
We usually represent this with an up arrow and a
down arrow.
Since there is only 1 s orbital per energy level, only
2 electrons fill that orbital.
Quantum numbers describe an electrons position, and no 2
electrons can have the exact same quantum numbers. Because of
that, electrons must have opposite spins from each other in order
to “share” the same orbital.
Hund’s Rule
Hunds Rule states that when you
get to degenerate orbitals, you fill
them all half way first, and then you
start pairing up the electrons.
What are degenerate orbitals?
Degenerate means they have the
same energy.
So, the 3 p orbitals on each level
are degenerate, because they all
have the same energy.
Don’t pair up the 2p electrons
until all 3 orbitals are half full.
Similarly, the d and f orbitals are
degenerate too.
Start from the Beginning
NOW that we know the rules, we can try to
write some electron configurations.
Remember to use your orbital filling guide to
determine WHICH orbital comes next.
Lets write some electron configurations for the
first few elements, and let’s start with hydrogen.
Electron Configurations
Element
Configuration
Element
Configuration
H Z=1
1s1
He Z=2
1s2
Li Z=3
1s22s1
Be Z=4
1s22s2
B Z=5
1s22s22p1
C Z=6
1s22s22p2
N Z=7
1s22s22p3
O
1s22s22p4
F Z=9
1s22s22p5
Ne Z=10
1s22s22p6
(2p is now full)
Na Z=11
1s22s22p63s1
Cl Z=17
1s22s22p63s23p5
K
1s22s22p63s23p64s1
Sc Z=21
1s22s22p63s23p64s23d1
1s22s22p63s23p64s23d6
Br Z=35
1s22s22p63s23p64s23d104p5
Z=19
Fe Z=26
Z=8
Note that all the numbers in the electron configuration add up to the atomic
number for that element. Ex: for Ne (Z=10), 2+2+6 = 10
Looking Back
One last thing. Look at the previous slide and
look at just hydrogen, lithium, sodium and
potassium.
Notice their electron configurations. Do you
see any similarities?
Since H and Li and Na and K are all in Group
1A, they all have a similar ending. (s1)
Electron Configurations
Element
Configuration
H Z=1
1s1
Li Z=3
1s22s1
Na Z=11
1s22s22p63s1
K
1s22s22p63s23p64s1
Z=19
This similar configuration causes them to behave the
same chemically.
It’s for that reason they are in the same family or group
on the periodic table.
Each group will have the same ending configuration, in
this case something that ends in s1.