Transcript Atomic Orbitals and Electron Configurations

Atomic Orbitals and Electron
Configurations
Quantum Mechanics
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.
Energy Levels
Red
Orange
Yellow
Green
Blue
Indigo
Violet
n=1
n=2
n=3
n=4
n=5
n=6
n=7
Quantum mechanics has a
principal quantum number. It is
represented by a little n. It
represents the “energy level”
similar to Bohr’s model.
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 “sublevels” which describe the specific
“atomic orbitals” for that level.
Blue = s block
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.
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?
Each Orbital can fit 2 electrons…
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
in your notes as 6we discuss this.
p
3 (3p orbitals)
d
5 (3d orbitals)
• The first
level (n=1)
has an s orbital.10 It has only 1.
There are no other orbitals in the first energy level.
n=4
s
1 (4s orbital)
2
32
• We call
this
orbital
the
1s
orbital.
p
3 (4p orbitals)
6
d
f
5 (4d orbitals)
7 (4f orbitals)
10
14
Electron Configurations
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
acounted for.
Fill
Lower
Energy
Orbitals
FIRST
Each line represents
http://www.meta-synthesis.com/webbook/34_qn/qn3.jpg
an orbital.
1 (s), 3 (p), 5 (d), 7 (f)
High Energy
Low Energy
The Aufbau Principle states that
electrons enter the lowest
energy orbitals first.
The lower the principal quantum
number (n) the lower the 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.
The Pauli Exclusion Principle
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
Hund’s Rule states that
when you get to degenerate
orbitals, you fill them all half
way first, and then you start
pairing up the electrons.
Don’t pair up the 2p electrons
until all 3 orbitals are half full.
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
Z=8
1s22s22p4
F Z=9
1s22s22p5
Ne Z=10
1s22s22p6
(2p is now full)
Na Z=11
1s22s22p63s1
Cl Z=17
1s22s22p63s23p5
K Z=19
1s22s22p63s23p64s1
Sc Z=21
1s22s22p63s23p64s23d1
Fe Z=26
1s22s22p63s23p64s23d6
Br Z=35
1s22s22p63s23p64s23d104p5
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
Electron Configurations
Element
Configuration
H Z=1
1s1
Li Z=3
1s22s1
Na Z=11
1s22s22p63s1
K Z=19
1s22s22p63s23p64s1
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.