Quantum Mechanical Model of the Atom
How does the quantum mechanical model of the atom explain the properties of electrons? It
treats e- like a wave with quantized energy. These waves cannot be pinpointed and therefore
Quantum Mechanics places e- in orbitals and not orbits.
What is an Orbital?
A region in space where an e- of specific energy is 90% likely to be found.
Compare and Contrast the Bohr Model of the atom with the Quantum Mechanical Model of the
atom: The Bohr model uses orbits at specific energy levels, where the QM model uses orbitals
of specific mathematical shape to predict an orbital. However they both used energy levels to
predict where the electrons will be located.
Electron Configurations
Look at how electrons “fill up” atoms: From the nucleus out.
Electrons fill atoms by Level Sublevel Orbital
I. Levels
Levels: Principle Energy Levels : n
n
1
2
3
4
5
6
7
e- capacity of energy
level 2n2
2
8
18
32
50
72
98
# of sublevels
1
2
3
4
5
6
7
1
II. Sublevels
Sublevel Names: s p d f g h i
Sublevels are designated with the principle quantum # of the energy level, then the sublevel
letter.
n
1
2
3
4
5
6
7
Sublevel Names
Electron Capacity
1s
2s 2p
3s 3p 3d
4s 4p 4d 4f
5s 5p 5d 5f 5g
6s 6p 6d 6f 6g 6h
7s 7p 7d 7f 7g 7h 7i
2
8
18
32
50
72
98
Sublevels denote the different energies electrons within a level can have.
The type of sublevel the electron is in describes the shape of the probability distributions
described by the quantum mechanical model for that electron.
s sublevel:
p sublevel:
f sublevel:
d sublevel:
2
III. Orbital
Definition: A region in space where an e- of specific energy is 90% likely to be found. There
can be no more than 2 e- per orbital.
There are a given number of orbitals per sublevel:
Sublevel
s
p
d
f
g
h
i
Number of Orbitals
Orbital Designation
1
s
3
px py pz
5
dxy dyz dxz dx2-y2 dz2
7
Complicated math formulas
9
|
11
|
13
↓
See pictures in text and draw p orbitals here: (page 109)
See pictures in text and draw d orbitals here:
3
Fill up an atom with electrons filling Level Sublevel Orbital
Orbital
Level
Sublevel
1
2
2
2
2
3
3
3
3
3
3
3
3
3
4
4
4
4
5
5
5
5
6
6
6
7
1s
2s
2p
3s
3p
3d
4s
4p
4d
4f
5s
5p
5d
5f
6s
6p
6d
7s
1s
2s
2px
2py
2pz
3s
3px
3py
3pz
3dxy
3dyz
3dxz
3dx2-y2
3dz2
4s(1)
4p(3)
4d(5)
4f(7)
5s(1)
5p(3)
5d(5)
5f(7)
6s(1)
6p(3)
6d(5)
7s(1)
2 Electrons per
orbital
2
8
18
32
32
1s2
2s2
2px2
2py2
2pz2
3s2
3px2
3py2
3pz2
3d2xy
3d2yz
3d2xz
3d2x2-y2
3d2z2
4s2
4p6
4d10
4f14
5s2
5p6
5d10
5f14
6s2
18
6p6
6d10
2
7s2
2p6
3p6
3d10
4
Orbital Overlap
The order that electrons fill up energy levels is not 1, 2, 3, 4, 5, etc. There is some overlap in
sublevel energies.
The periodic table can be used as a tool to read the order of sublevel energies
n
ns
1
1s1
1s2
2
2s1
2s2
3
3s1
3s2
4
4s1
4s2
3d1
3d2
3d3
3d5
3d5
3d6
3d7
3d8
3d10
5
5s1
5s2
4d1
4d2
4d4
4d5
4d5
4d7
4d8
4d10
6
6s1
6s2
5d1
5d2
5d3
5d4
5d5
5d6
5d7
7
7s1
7s2
6d1
6d2
6d3
6d4
6d5
6d6
6d7
2
3
4
5
6
7
7
1s2
np
2p1
2p2
2p3
2p4
2p5
2p6
3p1
3p2
3p3
3p4
3p5
3p6
3d10
4p1
4p2
4p3
4p4
4p5
4p6
4d10
4d10
5p1
5p2
5p3
5p4
5p5
5p6
5d9
5d10
5d10
6p1
6p2
6p3
6p4
6p5
6p6
6d9
6d10
6d10
7p1
7p2
7p3
7p4
7p5
7p6
(n-1)d
n
(n-2)f
6
4f
4f
4f
4f
4f
4f
4f
4f9
4f10
4f11
4f12
4f13
4f14
4f14
7
5f0
5f2
5f3
5f4
5f6
5f7
5f7
5f9
5f10
5f11
5f12
5f13
5f14
5f14
Write electron configurations for:
H:
1s1
He:
1s2
Li:
1s22s1
Be:
1s22s2
B:
1s22s22p1
C:
1s22s22p2
N:
1s22s22p3
O:
1s22s22p4
F:
1s22s22p5
Ne:
1s22s22p6
5
Exceptions
Exceptions to the AUFBAU filling order: (filling up in order of increasing energy) Cr,
Mo, Cu, Ag, Au
These elements have electron configurations that are exceptions due to the stability of
the filled or half –filled d sublevel.
Cr:
1s22s22p63s23p64s13d5
Mo:
1s22s22p63s23p64s23d104p65s14d5
Cu:
1s22s22p63s23p64s13d10
Ag:
1s22s22p63s23p64s23d104p65s14d10
Au:
1s22s22p63s23p64s23d104p65s24d105p66s14f145d10
Configurations of Ions:
Anions: Gain e- so add p electrons to adjust for charge.
Cations: Lose electrons so remove electrons to adjust for charge.

For Representative elements, remove last electron to be added FIRST.

For Transition elements, remove s electrons FIRST, then d until the charge has been
accounted for.
Examples: S2- Ca2+ N3- Cu- Fe3+ Pb2+ Pb4+
Abbreviated Notation:
Back up to the previous noble gas and place that symbol in square brackets. Now add the rest
of the electron configuration.
Examples:
I
Au
Xe
[Kr]5s24d105p5
[Xe]6s14f145d10
[Kr]5s24d105p6
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Electron Spin
Electrons behave as though they are spinning. Spinning charges create: a magnetic field
1. If spinning clockwise: ↑
2. If spinning counterclockwise: ↓
When 2 electrons have the same spins, their spins are called: parallel ↑ ↑
Magnetic fields add constructively (reinforce eachother)
Paramagnetic: elements containing 1 or more unpaired e- . These elements will be drawn into
an external magnetic field.
Ferromagnetic: elements with extreme paramagnetism. Ex: Fe
When 2 electrons have opposite spins, their spins are called: paired ↑ ↓
Spins cancel: Magnetic fields add destructively
Diamagnetic: elements with all spins paired, magnetic field cancels.
These elements will NOT be drawn into an external magnetic field. They may be unaffected, or
they may actually be repelled by an external magnetic field.
To determine if elements are paramagnetic or diamagnetic, use the following:
 Hund’s Rule: Electrons occupy equal energy orbitals so that a maximum number of
unpaired electrons results.
Give each orbital in a sublevel one electron before you give any orbital in that sublevel a
second electron.
 The Pauli Exclusion Principle: Electrons cannot exist in the same orbital with the same
spin. Therefore if an orbital is full it will have one up spin and one down spin electron.
Draw orbital diagrams to determine if e- are paired or unpaired.
Use a dash to represent each orbital, and place orbitals within the same sublevel together. Fill
orbitals with electrons from lowest energy to highest applying Hund’s Rule and the Pauli
Exclusion Principle.
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Are the following paramagnetic or diamagnetic?
Cr:
Fe
Zn:
P:
Ar:
para
para
dia
para
dia
Excited States
So far, all configurations we have looked at are for elements and ions in the ground state, or
lowest possible energy state. This means all possible orbitals are filled from lowest level to
highest.
In an excited state, an electron absorbs energy and jumps up to any higher level, sublevel that
has empty space.
Example:
Ground State Configuration for Oxygen: 1s22s22p4
Some of the possible excited states of Oxygen: 1s22s02p6
1s22s12p5
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