Transcript Chapter 4 - Educator Pages

Atomic Orbitals and
Electron Configurations
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
The energy levels in quantum mechanics describe
locations where you are likely to find an electron.
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
An electron can be found anywhere.
Scientists agreed to limit these calculations to locations
where there was at least a 90% chance of finding an
electron.
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?
Zero. There aren’t any electrons in the nucleus.
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
Blue = s block
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.
Orbitals
s
p
d
In the s block, electrons are in s orbitals.
In the p block, the s orbitals are full. New electrons fill the p
orbitals.
In the d block, the s and p orbitals are full. New electrons fill the d
orbitals
Orbitals
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
p
d
1 (3s orbital)
3 (3p orbitals)
5 (3d orbitals)
2
6
10
18
 Complete the chart in your notes as we discuss this.
n=4
s first level
1 (4s
orbital)
2 It has only 32
 The
(n=1)
has an s orbital.
1.
p
3
(4p
orbitals)
6
There are no other orbitals in the first energy level.

d
f
We
call
5 (4d orbitals)
(4f orbitals)
this 7orbital
the 1s
orbital.
10
14
Where are these Orbitals?
1s
2s
3s
2p
3p
4s
3d
4p
5s
4d
5p
6s
5d
6p
7s
6d
7p
4f
5f
Electron Configurations
The electron configuration is the specific way in which
the atomic orbitals are filled.
Think of it as being similar to an 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.
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.
What are degenerate orbitals? Degenerate
means they have the same energy.
The 3 p orbitals on each level are
degenerate, because they all have the same
energy.
Similarly, the d and f orbitals are degenerate.
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.