Chapter 3 Part II
How are the
electrons
arranged
around the
nucleus?
Normal or Rest Position
Wavelength
Crest
Amplitude
Wavelength
Amplitude
Trough
3
Parts of a Wave
Parts of a Wave
A. Amplitude: Height of the wave. The higher the
wave the greater the intensity.
B. Wavelength: (λ , “lambda”) in nanometers
(1 x 10-9 m) measured from crest to crest or from
trough to trough.
C. Frequency: (ν , “nu”) The number of times a wave
passes a fixed point. Measured in cycles/second (1/s)
1 cycle/second = Hertz (Hz)
ex) Radio FM 93.3 megahertz (MHz) is 93.3 x 106 cycles/sec.
Visible Light
Microwaves
Radio/TV
Radar
Ultraviolet
Infrared
Gamma Rays
X-Rays
Low
High
Long
Short
Low
High
Energy
Red
Orange
Yellow
Green
Blue
Indigo
Energy
Violet
The Electromagnetic Spectrum
Speed of Light
D. (c) 3.0 x 108 m/s or 186,000 miles/sec.
The relationship between wavelength and
frequency can be shown with the following
equations:
c
c
λ= or ν = this is an indirect relationship.
ν
λ
If λ  then ν .
Quantum Theory
A.
Planck’s Hypothesis:
(Max Planck 1900)
1. An object absorbs or emits light in
little packet or bundle called a
quantum (quanta –plural).
2. Energies are quantized.
(Think steps not a ramp.)
3. Equation relating energy (E)
to frequency (ν, nu)
(h= Planck’s Constant)
ee-
E = h
This is a direct relationship,
as ν  , E 
ee-
X
Light (Electromagnetic Radiation)
1. A quantum of light is called a photon.
2. Light travels through space in waves.
3. Light acts like a particle when it interacts with matter.
4. This shows the Dual Nature of Light.
3. The Atom
A. Atomic Emission Line Spectra: contains only
certain colors or wavelengths (  ) of light.
1. Every element has its own line spectrum (fingerprint).
Double Slit Experiment
This Experiment Proves the Dual Nature of
Light – Photons of Light travel through
spaces in waves, but act like particles
when they interact with matter.
Double Slit Experiment Video
Continuous Spectrum – White Light
Line Spectrum – Excited Elements
Line Emission Spectra of Selected Elements
Gas Discharge Tubes
• Electricity is added to the gas which causes the electrons to
jump to a higher or excited state. They immediately fall
back to the ground state and give off particular wavelengths
of light. We see a blending of wavelengths without the
spectroscopes.
White Light gives off a Continuous Spectrum
a blending of every possible wavelength
Spectroscope
• Uses a diffraction grating to diffract the light
into particular wavelengths of light.
A Line Spectrum results from excited elements - as
electrons of an element gain energy and rise to an
excited state they then fall back to their ground state in
the same pattern producing the same energy drop each
time which we see as individual wavelengths of light.
Atomic Spectra
Hydrogen
Helium
Lithium
Mercury
Although Bohr’s
atomic model
explained the line
spectra of hydrogen,
it failed for heavier
elements.
B. Bohr Model of Hydrogen: (1911)
1. Bohr said the energy of an electron was quantized
(only certain orbits that represented different amounts of
energy.)
2. Bohr labeled each energy level with a quantum number (n).
3. n=1 lowest level or ground state.
4. When electrons absorb energy they jump to a higher
(excited) state. n=2 n=3 n=4 n=5 n=6 n=7
5. Radiation (light) is emitted when an electron falls back from
a higher level to a lower level.
Electrons release energy as they fall
back to a lower energy level
Excited Atoms Emit Photons of Light:
Electrons absorb energy to rise to a higher or excited
state and emit energy in the form of a photon of light
as they fall back to their ground states.
Path of an excited electron as it “falls”
back to the Ground State
• When electrons gain
energy, they jump to a
higher energy level
(excited state).
• Electrons are not stable at
the excited state and will
immediately fall back to a
lower level or ground state.
• As they fall, they emit
electromagnetic radiation.
• Depending on how far
they fall determines the
type of radiation (light)
released.
Hydrogen Atom
electrons fall to n = 1 and
give off ultraviolet light.
electrons fall to n = 2 and
give off visible light.
electrons fall to n = 3 and
give off infrared light.
Infrared Light
Visible Light
Ultraviolet Light
Werner Heisenberg: (1927)
1. Heisenberg’s Uncertainty Principle: states the
position and momentum (speed, direction & mass) of
a moving object cannot simultaneously be
measured and known exactly.
2. You cannot predict future locations of particles.
3. He found a problem with the
Bohr Atom - no way to observe or
measure the orbit of an electron.
4. Quantum Mechanics
A. Quantum Mechanical Model of the atom
combines previous ideas and treats the
electron like a wave that has quantized energy.
Impossible to state the exact position or
momentum of an electron, but you can state a
probability of where the electron is located.
B. Electron Density
Where the density of
an electron cloud is
high there is a high
probability that is
where the electron is
located. If the electron
density is low then
there is a low
probability.
C. Atomic Orbitals - region around the
nucleus where an electron with a given
energy is likely to be found
(not the same as Bohr’s orbits!)
1. Orbitals have characteristic shapes, sizes, &
energies.
2. Orbitals do not describe how the electron moves.
3. The drawing of an orbital represents the
3-dimentional surface within which the electron
is found 90% of the time.
4. Sublevels can have 4 different shapes
s – orbital spherical
1s, 2s & 3s orbitals Superimposed on
one another
Electron-Cloud Models
p-orbital – dumbbell shaped
p-orbital - dumbbell shaped
d-orbital - double dumbbell or fan blades
s,p and d orbitals
z
z
z
x
x
x
y
y
x
y
s orbital
y
p orbitals
z
z
z
x
x
y
z
y
z
x
y
z
x
y
x
y
d orbitals
For a more complete representation and presentation of atomic orbitals go to
http://winter.group.shef.ac.uk/orbitron/
Models of d-orbitals
f-orbital –
more complex!
Quantum Numbers - Finding an
address for each electron:
1) “state” Principle Quantum Number (n) or the
energy level ranges from n=1 to n=7
2) “city”
Sublevel shape either s, p, d,or f.
3) “street” Orbital The orientation in space
ex) x,y,z axis
4) “house” Spin The cw or ccw motion of electrons.
1s, 2s,2p and 3s orbitals
superimposed on each other
• Model of s and p Orbitals Together
2. The number of Sublevels in an energy level
equals the Principle Quantum Number (n).
3. Orbital: There are always an odd number of orbitals.
s-sublevel has 1 orbital
p-sublevel has 3 orbitals
d-sublevel has 5 orbitals
f -sublevel has 7 orbitals
• Orbitals in higher principle levels get larger.
• A max of 2 electrons fit in each orbital.
Electron Spin
a. Two electrons in each orbital have opposite
spins. (clockwise and counterclockwise)
b. The opposite spin is written:  or ___
c. Pauli Exclusion Principle:
1. Each orbital can only hold 2 electrons.
2. The electrons must have opposite spins.
s-sublevel
p-sublevel
d-sublevel
f-sublevel
=
=
=
=
max 2 electrons
max 6 electrons
max 10 electrons
max 14 electrons
incorrect: ↑↑↑ incorrect: ↑↑ correct: ↑↓
d. Hund’s Rule:
• Electrons will spread
out among the
orbitals before they
pair up.
incorrect ↑↓ ↑
__
correct
↑
↑
↑
E. Electron Configurations:
1. Shows the distribution of electrons among the
orbitals. Describes where the electrons are found &
what energy they possess.
2. The Aufbau Principle: electrons are added up
one at a time to the lowest energy orbital.
Aufbau Diagram/ Diagonal Rule
Electron Configuration Examples:
Ex) electron configuration for Na:
1s2 2s2 2p6 3s1
Ex) orbital filling box diagram for Na:
x
y
z
      _
1s
2s
2p
3s
3. Electron Dot Diagrams:
Write the symbol for the element.
Place dots around the symbol to represent the
valence s & p electrons only.
Do NOT include d & f orbitals in diagram.
p orbital electrons
s orbital electrons
Electron Configuration
Aufbau Diagram Order
Aufbau Diagram
Visualizing The Electron~
Model of Bohr Atom - Electron Movement
Line Emission Spectra of Selected Elements
The Aufbau Principle
(Diagonal Rule)
1s2
2s2
2p6
3s2
3p6
3d10
4s2
4p6
4d10
4f14
5s2
5p6
5d10
5f14
6s2
6p6
6d10
6f14
7s2
7p6
7d10
7f14
Increasing energy
7s
6s
5s
7p
6p
5p
4p
4s
3p
3s
2p
2s
1s
6d
5d
4d
5f
4f
3d
Pauli Exclusion Principle: No more
than 2 e- are put in each orbital and
they must have opposite spin.
Hund’s Rule: electrons spread out
among equal energy orbitals in a
sublevel (like charges repel)
Aufbau Principle: Electrons fill
lowest energy levels first (n=1)
Electron Blocks on the
Periodic Table
Electron Configuration
Orbital Box Diagram
1s22s22p4
O
z
    
1s
2s
2p
x
35
17
y
x
16
8
Electron-dot Diagram
22s22p63s23p5
1s
Cl
y
z
x
y
z
        
1s
2s
2p
3s
3p
1s22s22p63s23p64s23d104p65s24d105p4
127
52
Te
x
y
z
x
y
z
x
y
z
x
y
z
                          
1s
2s
2p
3s
3p
4s
3d
4p
5s
4d
What does the Tellurium electron-dot resemble???
5p
The Copper Atom
Mark your Periodic Tables
1
2
13
14
15
16
17
18
Scanning Tunneling Microscope
In 1926, Erwin Schrodinger derived an
equation that described the energy and
position of the electrons in an atom

d

V 
8  m dx
h
2
2
2
2
 E
Equation for the
probability of a single
electron being found
along a single axis (x-axis)
Erwin
Erwin Schrodinger
Schrodinger