Quantum
Model and
Electron
Configurations
Atomic Models:
Old
 Also
version = Bohr’s
known as the
planetary atomic
model
 Describes electron
paths as perfect orbits
with definite diameters
 Good for a visual
 New
version =
Quantum Theory
 Most accepted
 Diagrams electrons of
a atom based on
probability of location
at any one time
Bohr’s model:
 Nucleus
is in the center of an atom(like the sun) and
the electrons orbit the nucleus similar to the planets.
 Orbits are called shells




1st shell = 2 electrons
2nd shell = 8 electrons
3rd shell = 18 electrons
4th shell = 32 electrons
Last slide
QOD
What
is the
approximate
mass of an
electron?
0.000549
amu
VOCAB
 WHICH
IS A TRUE
STATEMENT?
 Compounds
can
be broken down
(decomposed) by
chemical means
 Compounds can
be decomposed
by physical means
Quantum
THEORY &
Mechanics
Study of how light interacts
with matter
Quantum Theory:
To
better the description of the atomic
structure, atoms were exposed to energy
(heat) which made the electrons go into
what is called the excited state (normal =
ground state).
When electrons returned to ground state
they emitted energy in the form of light.
Quantum Theory:
This method of study is called
spectroscopy (spectrum)
Visible light = part of the
electromagnetic spectrum
between 400-700 nm
Electromagnetic Spectrum
Quantum Theory:
 Electromagnetic
Spectrum
 From
crest to crest =
frequency which is
measured in hertz.
 This therefore can be
used to identify
elements (absorption
of energy and color
emitted is a fingerprint
of an element)
Kind of like wearing your team
colors.
Continuous spectrum of white light
 When you pass sunlight through a
prism, you get a continuous spectrum
of colors like a rainbow.
Line-Emission Spectrum
 However, when light from Hydrogen
& Helium gases were passed through
a prism, they found a dark
background with discrete lines.

WHY? This lead to the quantum theory.
H
He
A scientist, Bohr
suggested that electrons
must exist in Electron
Orbitals (shows the most
probable area to find an
electron of a certain
energy.)
Quantum Theory:
So whenever an excited
hydrogen atom falls to its
ground state or lower
energy level, it emits a
photon of light, which
means that energy levels
must be fixed.
 Video
Quantum Theory: Electron
Configuration
Electrons
(e-) of atoms
are the basis for every
chemical reaction.
In quantum theory,
electrons exist in
orbitals based on
probabilities and these
orbitals are arranged
within energy levels.
Notice… these orbitals
look different from Bohr’s.
This diagram is more
correct.
Quantum Theory: Electron
Configuration
Quantum Numbers
Quantum
numbers
specify the properties
of atomic orbitals
and the properties of
electrons in those
orbitals
We will define these
numbers & letters.
Example of Quantum
#:
2
3s
Quantum Theory: Electron
Configuration
Principle Quantum
Number (n)
Is
equal to the number of
the energy level (n).
The principle quantum #
corresponds to the energy
levels 1-7 which is the
period number (row) on
the periodic table.
Example of Quantum #:
2
3s
Blocks and Sublevels
P
E
R
I
1
3
O
4
D
6
S
1-7
d (n-1)
2
5
7
4
5
Quantum Theory: Electron Configuration
Maximum e- in Energy Levels
 The
maximum number of e- in
any one level is given by the equation 2n2
N=4, 32e
 Calculate
the maximum
number of electrons that can
occupy the 4th principal
quantum number (period 4).
 Solve: Use 2n2
2(4)2
32 electrons total
N=3, 18e
N=2,
8e
N=1, 2e
Quantum Theory: Electron Configuration
Sublevels and Orbitals
 An
energy level in made up of many
energy states called sublevels.
 The number of sublevels for each
energy level is equal to the value of
the principal quantum number.
EX: one sublevel in energy level one
(period 1)
two sublevels in level two (period 2)
three sublevels in level three (period 3)
*now lets find out what those sublevels
are called…
Example of Quantum #:
2
3s
Quantum Theory: Electron Configuration
Sublevels and Orbitals
 There
are 4 sublevels:
s
 Energy
p
d
f
levels and sublevels work together to form an
e- cloud.
 e- are repelled by one another and move as far apart
as possible.
 e- clouds take on characteristic shapes called
orbitals.
Sublevels and Orbitals (notice the
shapes)
Orbital Shapes
s orbitals are
spherical.
This diagram
represents an s
orbital.
d orbitals
contains 5
possible
orbital
shapes.
p orbitals
are
“dumbbell”
shaped.
This diagram
represents 1 of the
3 types of p
orbitals.
f orbitals
contain 7
possible
orbital
shapes.
Electrons & Orbitals
Orbitals
overlap and
change shape
as electrons
are added.
Each orbital
can only hold
2 electrons.
Example of Quantum #:
2
3s
Electrons and Orbitals
per orbit)
(count 2 electrons
Orbitals, and Electrons per Sublevel
Principal Quantum
Number (n)
Sublevel
# of
Orbitals
# of Electrons per
Orbital
1
s
1
2
2
s
p
s p
d
s p
d f
1
3
1 3
5
1 3
5 7
2
6
3
4
2 6
10
2 6
10 14
QOD
The
principal
quantum number
corresponds to the:
•Energy Levels
•Periods on the
periodic table
VOCAB
 Which
statement is
true: The
characteristic brightline spectrum (color)
of an element is
produced when its
electrons…
 Move to an excited
higher energy state
 Return to a lower
ground energy
state
Distribution of Electrons



1.
Atoms are electronically neutral. (for now)
There is an electron for every proton in the nucleus.
The larger the atom, the larger the electron cloud.
Pauli Exclusion Principle: only two e- can occupy the
same orbital due to
the opposite electronic spin .
Electron Filling Diagram
•Sublevels and
orbitals are filled as
indicated in the
diagram.
•Example:
1s2 2s2 2p6 3s2 3p6
4s2 …
energy
level
Notice… they don’t go in order !
Sublevel
orbital
# electrons
in the
orbital
Label your blank periodic
table.
Read it “like a book”
WRITE the Electron
Configuration
Now try:
1. C
2. Kr
3. Ca
4. Fe
5. Hg
QOD
What
is the total
number of electrons
that can be held in
the third principal
energy level?
2
2n
18
VOCAB
 Quantum Theory of
atomic structure states all
except:




Electrons orbit the nucleus
in perfect paths
Electrons form clouds
based on probability of
location
Electron clouds form
characteristic shapes due
to repelling of negative
charges
Electrons occupy the
lowest energy levels
before moving into higher
energy levels
Label your blank periodic
table.
Read it “like a book”
DRAW the Electron Configuration

Carbon has 6 e- (same as protons)

Start with lowest energy level and place
one electron in each orbital. Spins must
be in same direction within orbitals of the
same energy level.

If there are remaining e-, pair up singles
in same energy level before moving to
next highest energy level.
Electron Configuration
Carbon’s electron config. is:
1s2 2s2 2p2
1s
2s
2p
Superscripts total the number of
electrons
2+2+2=6
*Notice that you can write the electron
configuration based on the orbital diagram.
*When asked to draw or diagram, use arrow
configuration.
Last
*When asked to write, use 1s2,2s2… configuration.
slide
Electron Configuration – Noble
 Electron
Configuration
demonstrates a periodic trend, so
Gas
Configuration
you can write shorthand electron configuration using the
electron configuration of the noble gases in Group 18 of
the periodic table.
 Noble gases have stable configurations.
Noble
Gas Configuration
writing shorthand
e- config
 When
for an element, refer to the noble
gas in the energy level (period) justElectron
above the element.
Configuration
 Write the symbol of the noble gas in
Na = 1s22s22p63s1
brackets.
Al =
 Write out the remaining e-config
1s22s22p63s23p1
based on the energy filling
22s22p6
Ne
=
1s
diagram.
Shorthand
Electron
Configuration
Na = [Ne] 3s1
Noble Gas Configuration
EX: Na
Step 1: Na is in period 3 so refer to the noble gas in
period 2 which is Neon.
Step 2: Write Ne in brackets.
[Ne]
Step 3: Now write remaining electrons in standard
form. 3s1.
Step 4: Combine. [Ne]3s1
Noble Gas Configuration
EX: Br
Step 1: Br is in period 4 so refer to noble gas from
period 3 which is Argon.
Step 2: Write in brackets. [Ar]
Step 3: Write remaining electrons.
4s23d104p5
Step 4: Combine to form: [Ar] 4s23d104p5
*Check your work: Add the number of electrons from
the noble gas (18) to the subscripts of the remaining
e-config (17). 18+17=35 which is the electrons for Br.
Nobel Gas Configuration
Now try:
1. I
2. Kr
3. Na
4. Cu
Label your blank periodic
table.
Read it “like a book”
Electron Configuration with
Ions

When we write the electron configuration of a
positive ion, we remove one electron for
each positive charge:
Na
→
1s2 2s2 2p6 3s1 →

Na+
1s2 2s2 2p6
When we write the electron configuration of a
negative ion, we add one electron for each
negative charge:
O
1s2 2s2 2p4
→
→
O21s2 2s2 2p6
Electron Configuration
with Ions
Now try:
1. Ca+2
2. Fe-3
Label your blank periodic
table.
Read it “like a book”
QOD
What element
has completely
filled 3p
orbitals?
VOCAB
Which
of the
following is the
correct name for
Ca+1?
 Calcium
isotope
 Calcium
Argon (Ar)
1s2 2s2 2p6 3s2
3p6
 Calcium
ion
 Calcium with
extra electrons
Label your blank periodic
table.
Read it “like a book”
***
***
S - Block
1
1
D - Block
2
6
P - Block
1
1
2
3
4
5
2
2
3
1
4
3
5
4
6
7
2
3
4
5
6
7
8
9
3
10
4
5
1
1
6
5
F - Block
6
1
2
3
4
5
6
7
8
9
10 11 12 13 14
4
4
5
5
Last
slide