Quantum Mechanical Model Notes (5.1-5.2)
Discoveries and facts that led to the
development of the Quantum
mechanical model
Dalton’s Atom Model
J.J. Thomson’s Plum Pudding Model
Rutherford’s Atomic Model
Problems with this model
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Light
Light is part of the electromagnetic
spectrum
electromagnetic radiation 
-Atoms are indivisible
-Atoms speckled with tiny negative charges called electrons
-Atoms contain a positively charged core called the nucleus that
is made up of protons. Electrons float around the nucleus.
-this model could not explain why electrons do not float
straight into the nucleus (opposite charges attract)
-electrons had no order
******************************************************************
- the spectrum is formed by electromagnetic radiation
- particles of energy that travel through space in waves
http://www.lcse.umn.edu/specs/labs/images/spectrum.gif
General Facts



Light is made of particles that travel in waves
Light is composed of tiny “packets” of energy called photons
Each photon carries a quantum of energy
Quantum 

Minimum amount of energy gained or lost by particles
(electrons)
Photoelectric Effect

It was observed that when light shines on metal, electrons are
emitted.
The photon of light hitting the metal must contain a certain
amount of energy in order for the electrons to be emitted.
It was concluded that electrons can only absorb a specific
amount of energy - too little and the electron will not moveelectrons must absorb or release a specific quantum of
energy.
Electrons at normal energy level
Electrons that have absorbed energy


Electrons in their ground state:
Electrons in their excited state:
What does all of this mean?


- The idea that electrons have a specific amount of energy helps
to explain why electrons do not float into the nucleus and stick
(due to opposite charges). The energy that electrons possess
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To summarize
The Bohr Model
Based on the above revelations:
Bohr tried to describe the location of
electrons with Energy Levels
Niels Bohr
helps them to overcome the attractive force of nucleus.
- The more energy each electron has, the farther it can be from
the nucleus.
- Low energy electrons float near the nucleus
- Higher energy electrons float farther away
Niels Bohr proposed that electrons exist around the nucleus at
discreet distances based on the energy each electron carries.
 The average “distance” (radius) of an electron from the
nucleus
o
There are 7 known energy levels and an 8th energy level
has been proposed.
o
The symbol for energy level is “n”
o
n = 1 means energy level 1.
o
View the example below (ignore all the series stuff – this
was the only reasonable simple image I could find online)
Example of energy levels:
http://www.frontiernet.net/~jlkeefer/balmer1.gif
Can electrons move between energy
levels?
According to Bohr, each energy level
can only contain a specific number of
electrons (due to electron-electron
repulsive forces)
 If an electron absorbs a specific amount of energy it will
become excited and jump up to the next available energy
level.
 When the electron gives off that extra energy (through heat
and light) it moves back down to its original energy level.
Energy
# of Electrons
levels (n=?)
N= 1
2
N=2
8
N= 3
18
N= 4
32
N=5
18 (could be 32-50 in
future if f and G sublevels
occur)
N= 6
8 (could be 18- 72 in
future if D, F, G, and H
sublevels occur)
N=7
2 (could be 8 - 98 in future
if filled)
You are responsible for knowing the highlighted
numbers
 As the distance from the nucleus increases, the space for
electrons to exist increases as well, that is why more electrons
can be added to each energy level
 Energy levels can also be called Principle Quantum Numbers
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The Quantum Mechanical Model:
Building on the information from the
Bohr model we begin to develop more
specific information about electrons
forming the Quantum mechanical
model
Sublevels
4 types of Sublevels (aka subshells)
Orbitals
Image of an S-sublevel
1 orbital (max 2 e-)
http://web.rollins.edu/~jsiry/atom-quantum.jpg
 A sublevel describes the space and shape within each energy
level in which an electron is most likely to be found.
 S, P, D, and F (some sources now include G and H sublevels,
but not yet in our standards)
o Each sublevel contains a specific number of
orbitals
o S can hold 2 e- has one orbital
o P can hold 6e- has three orbitals
o D can hold 10e- has five orbitals
o F can hold 14 e- has seven orbitals
 Spaces within a sublevel that can contain 2 electrons
o Helps one to understand how atoms bond
http://www.chem.umass.edu/genchem/whelan/111_Summer_2008_Index.htm
Image of a P-sublevel
3 orbitals (max 6 e-)
Image of a D-sublevel
5 orbitals (max 10 e-)
http://www.rmutphysics.com/charud/scibook/crystal-structure/porbital.gif
http://www.coldfusionenergyscience.com/theory
http://media-2.web.britannica.com/eb-media/47/7447-004-142D2365.gif
Images of an F-sublevel
7 orbitals (max 14 e-)
http://boomeria.org/chemtextbook/fig9-13.jpg
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Table from your book to summarize the
information
Principle Quantum
#
(Aka energy level)
Sublevels
#Orbitals
# of
Electrons per
Orbital
Total # of
Electrons
1
2
S
S
P
S
P
D
S
P
D
F
S
P
D
F
S
P
D
S
P
1
1
3
1
3
5
1
3
5
7
1
3
5
7
1
3
5
1
3
2
2
6
2
6
10
2
6
10
14
2
6
10
14
2
6
10
2
6
2
8
3
4
5
6
7

Apartment Analogy
Let’s try to use the apartment analogy
from your book to clarify the meanings
of parts of the electron cloud.
Analogy Key:
Apartment Building 
The Street 
The Floor of the Building 
The style of apartment 
The number of bedrooms 
The Electron Configuration:
4 Main Rules for Placing Electrons
around the Atom
1. Aufbau Principle:
2. Pauli Exclusion Principle:
3. Hund’s Law:
18
32
18
32
8
18
2
8
Red letters and numbers are not required for this class. I
have also not listed the G and H sublevels as those are still
being debated.
If an atom were like an apartment building and the street was
the nucleus then: The different floors of the apartment would
represent the energy levels (distances from the nucleus), the
different styles of apartments available on each floor would
represent the sublevels, and the number of bedrooms per
apartment would represent the orbitals. Electrons that have low
energy must live on the first floor (energy level). Electrons that
have more energy can climb the stairs and live on other floors.
The Atom
The nucleus
The Energy Level
The sublevel (spdf)
The orbitals
- fill in electrons of the lowest energy orbital first. Build your atom
from the bottom up.
- Only 2 Electrons may be placed in each orbital of a sublevel
(2 in the S, 2 in the Px, 2 in the Py, and 2 in the Pz etc.)
-Electrons in each orbital must be of opposite spin.
 Electrons of opposite spin are represented by up and down
arrows: ↑↓
-When filling a sublevel that contains more than one orbital,
place one electron in each orbital first and then if more electrons
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4. Heisenberg Uncertainty Principle:
Two Diagrams that show how electrons
are placed around the nucleus.
1. Orbital Diagram
Filling in the orbital diagram:
are to be added go back and fill in the remaining spaces.
 Using our analogy: each electron gets its own bedroom in
the apartment unless more electrons move in. As more
electrons are added, each room gets filled up.
↑↓ ↑↓
↑ ↑ ↑ add one more electron: ↑↓ ↑↓
↑↓ ↑ ↑
1S 2S
px py pz
1S 2S
px py pz
2P
2P
- Since electrons are in constant motion, one can never pinpoint
the exact location of an electron.
- We can only narrow down the possibilities of where an electron
may be found: Energy level  sublevel  orbital  ? (electron
is moving too much to be able to better define its location any
further than this)
- This diagram provide the following information about the
location of electrons in an atom:
 The energy levels (represent by #’s 1-7)
 The sublevel (spdf)
 The orbitals (represent by _ lines)
 The # of electrons (represent as up and down arrows ↑↓)
- According to the Aufbau principle, one must place electrons in
the lowest energy orbital first. This is not as easy as it sounds. The
problem is that not all of the orbitals of the same energy level get
filled at the same time; there is a fluctuating pattern that must be
followed – see image below
To determine the correct filling order:
One must follow each arrow as it goes
down. Whichever sublevel is touched
first by the arrow gets filled in first.
http://www.fordhamprep.org/gcurran/sho/sho/images/elecfill.gif
In other words, the order of filling sublevels is as follows:
1S 2S 2P 3S 3P 4S 3D 4P 5S 4D 5P 6S 4F 5D 6P 7S 5F 6D 7P
The orbital diagram for Carbon would
be:
2. Electron Configuration:
Carbon: 6 electrons
↑↓ ↑↓ ↑ ↑ _
1s 2s px py pz
2P
(notice that in the p-orbitals only one arrow is placed in each 
Hund’s Law)
-The electron configuration is a shorthand version of the orbital
diagram.
 The orbital diagram only shows the following:
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The energy level (represented by #s 1-7)
The sublevel (spdf)
The number of electrons (represented by a
superscript/exponent)
Example: Carbon’s diagram ↑↓ ↑↓ ↑ ↑ _
1s 2s px py pz
2P
Becomes this  1s2 2s2 2p2
o
o
o
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