CHAPTER 5
The wave nature of light
Electromagnetic radiation: is a
form of energy that exhibits
wavelike behavior as it travels
through space.
Examples: Visible light
X – rays
Gamma rays
https://www.youtube.com/watch?v
=m4t7gTmBK3g
Electromagnetic spectrum
Transverse wave
Short
wavelength
Long
Wavelength
c = λν
All electromagentic waves
travel at the speed of light
C = 2.99 x 108 m/s
(3.0 x 108 m/s)
(sound only travels
at 340 m/s)
c = λν
All electromagnetic waves travel
at the speed of light
C = 2.99 x 108 m/s (3.0 x 108 m/s)
(6.7 x 108 miles per hour)
(Sound only travels at 340 m/s)
Wavelength is in meters
Frequency is in hz (1/sec)
c = λν
1. What is the frequency of green light,
which has a wavelength of 4.90 x 10 -9 m?
2. An X-ray has a wavelength of
1.15 x 10-10 m. What is its frequency?
3. Z100 broadcasts at a frequency of
100.3 MHz. What is the wavelength of the
broadcast? (HINT: UNITS)
c = λν
4. A compound emits blue light at a
wavelength of 4.50 x 10-5cm. What is
the frequency of the light?
5. What is the frequency of a wave if its
wavelength is 3.6 nm?
c = λν
6. Calculate the wavelength of an
electromagnetic wave that has a
frequency of 1.50 x 1013 Hz. (c = λν)
7. Is the wavelength longer or shorter
than red light?
E = hv
Frequency
Energy of
a photon
(J)
(Hz)
Plank’s Constant
6.626 x 10
-34
Energy and frequency have a
direct relationship
High energy = more violet light
J•s
E = hv
h = 6.626 x 10-34 J•s
An atom releases blue light with
a frequency of 6.8 x 10 12 Hz.
What is the energy of the photon
emitted?
c = λν
E = hv
1. Calculate the wavelength of an
electromagnetic wave that has a
frequency of 1.50 x 1013 Hz. (c = λν)
2. Calculate the energy of the same wave
Review: Relationships:
Direct or inverse?
Wavelength and frequency
Frequency and energy
Wavelength and energy
Emission spectrum
Incomplete spectra that are
characteristic of a substance
(like atom fingerprints)
Bohr’s model with the
hydrogen atom
Bohr Model
Electrons are in ‘shells’
Each electron ‘shell’ is
quantized
(the larger the radius, the
larger the energy)
Principal Energy levels (n) 1  7
Ground state vs. Excited State
Electrons gain energy and jump to
higher energy levels = Excited State
The excited state is
very unstable and the
electrons jump back
to their previous
energy levels
releasing a photon
(a bundle of energy :
LIGHT ENERGY!)
Filling order of the first
20 elements: 2, 8, 8, 2
Filling order:
For the first 20 elements: 2,8,8,2
For elements 21 – 57: 2,8,18,18,32
For elements 58-114 : 2,8,18,32,32
Valence e-
for representative elements
Organization of the periodic
table: Columns = groups
Rows = periods
A = representative
B = transition
Inner transition
Organization of the periodic table
Rrepresentative elements families: chemical
and physical properties are similar
1: Alkali metals –soft metals, very reactive
2: Alkaline metals – reactive metals
15: Pnictogens – Nitrogen column
16: Chalcogens – Oxygen column
17: Halogens – very reactive nonmetals
Fluorine: Very reactive and poisonous
Chlorine: Yellow – green gas
Bromine: Redish-brown liquid
18: Noble gases – very UNreactive gases
Quantum Mechanical
Model – ‘electron cloud’
• Based on the quantum theory, which
says matter has wave-like properties
• According to quantum theory, it’s
impossible to know the exact position
of an electron (Heisenberg’s
Uncertainty Principle)
https://www.youtube.com/watch?v=TQKE
LOE9eY4
Quantum Mechanical
model – ‘electron cloud’
(cont)
https://ww
The model uses complex
shapes called orbitals (volumes
of space in which there is likely
to be an electron)
w.youtube
.com/watc
h?v=cPDp
tc0wUYI
Principal energy level (n)
Sublevels
________
________
________
________
Sublevels
___S___
# orbitals _____
___P__ __D___ __F__
_____
_____
_____
Orbital = the most probable location for an electron (90%)
Shapes of s, p, and dOrbitals
s orbital
p orbitals
d orbitals
Each orbital can hold _____
electrons
# of electrons all together:
_____
_____
_____
_____
Ground State - Electron
configuration
“Address” for an atoms electrons: nLe
Aufbau principle: electrons occupy the
lowest energy orbit available in each
principal energy level first
Every principal energy level adds 1
new sublevel
Total number of electrons possible on
each sublevel = 2n2
Periodic Patterns
• Example - Hydrogen
1
2
3
4
5
6
7
1
1s
1st Period
1st column
of s-block
s-block
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
Energy levels
Overlap
4s3d4p
6s4f5d6p
Valence Electrons :
• The valence electrons are the electrons in the
outermost principal energy level. (Last energy
level on the bohr model)
• They are always ‘s’ or ‘s and p’ orbital electrons
• There can be no more than 8 valence electrons
Other electrons are called ‘kernel’ electrons
Carbon = 1s2 2s2 2p2
Kernel/core
Valence
Shorthand / Noble gas
notation
• Example -Calcium
Noble gas
that
precedes
element
1
2
3
4
5
6
7
2
[Ar]4s
4th Period
2nd column
s-block
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
Orbital diagrams
Uses boxes and
arrows to
represent electrons
Aufbau principle:
electrons occupy
the lowest energy
orbit available in
each principal
energy level first
Orbital diagrams
Uses boxes and arrows to represent
electrons
Pauli Exclusion principle: each electron in an
orbital must have a different spin (one up,
one down)
Hund’s rule: Electrons repel each other so
they must be distributed: single electrons
occupy each orbital before electron pairs
can be made