Special Topics for SOL 2
nd
2 Power Point
Atomic Orbitals and Electron Configurations (Chap 13)
Quantum Mechanics
http://www.meta-synthesis.com/webbook/30_timeline/310px-Bohr-atom-PAR.svg.png
 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
http://milesmathis.com/bohr2.jpg
 Much like the Bohr model, the energy
levels in quantum mechanics describe
locations where you are likely to find
an electron.
 Remember that orbitals are
“geometric shapes” around the
An old Bohr??
nucleus where electrons are found.
Mwwhaha!
 Quantum mechanics calculates the
probabilities where you are “likely”
to find electrons.
Atomic Orbitals
http://courses.chem.psu.edu/chem210/quantum/quantum.html
 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.
Energy Levels
http://www.chem4kids.com/files/art/elem_pertable2.gif
 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. It’s amazing how all this
stuff just “fits” together.
Sub-levels = Specific
Atomic Orbitals
 Each energy level has 1 or more
“sub-levels” 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
http://media-2.web.britannica.com/eb-media/54/3254-004-AEC1FB42.gif
http://upload.wikimedia.org/wikipedia/commons/thumb/e/e1/D_orbitals.svg/744px-D_orbitals.svg.png
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?
Objective C
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
Complete
p
d
1 (3s orbital)
the3chart
in your
(3p orbitals)
5 (3d orbitals)
p
3 (4p orbitals)
f
7 (4f orbitals)

notes
2
as6 we
10
18
discuss this.
 The first level (n=1) has an s orbital. It has only 1.
orbitals in the first
n = 4 There
s are no other
1 (4s orbital)
2 energy level.
32
 We dcall this orbital
the 1s orbital.
5 (4d orbitals)
6
10
14
Where are these Orbitals?
http://www.biosulf.org/1/images/periodictable.png
1s
2s
2p
3s
3p
4s
3d
4p
5s
4d
5p
6s
5d
6p
7s
6d
7p
4f
5f
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 Electon Configurations
https://teach.lanecc.edu/gaudias/scheme.gif
 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
an orbital.
1 (s), 3 (p), 5 (d), 7 (f)
High Energy
http://www.meta-synthesis.com/webbook/34_qn/qn3.jpg
 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.
http://www.fnal.gov/pub/inquiring/timeline/images/pauli.jpg
 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.
Wolfgang Pauli, yet
another German
Nobel Prize winner
 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
http://intro.chem.okstate.edu/AP/2004Norman/Chapter7/Lec111000.html
 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.
Objective D
 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
Z=8
1s22s22p4
F
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
Z=9
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
Objective D
 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 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.
The End