Electromagnetic Radiation
and Light
I. Models of the Atom
• Many different models:
– Dalton-billiard ball model (1803)
– Thompson – plum-pudding model (1897)
– Rutherford – Nuclear model of the atom (1911)
– Bohr – uses quantized energy of the atom (1913)
– Quantum Mechanical Model of the Atom (1926)
• Each new model contributed
to the model we use today.
• Even our current model,
does not give us an exact
model of how electrons
behave.
A. The Bohr Model
• Bohr used the simplest element, hydrogen, for
his model
• Proposed electron is found in specific circular
paths, or orbits around the nucleus
• Each electron orbit was
thought to have a fixed
energy level.
• Lowest level-ground state
• Any Higher Levelexcited state
The Bohr Model cont.
• One electron is capable of many different
excited states (e- jumping to higher level)
• Quantum: specific amount of energy an e- can
gain or lose when moving energy levels
• You can excite an e- with energy like electricity,
the sun, or magnets
Electron dropping from
higher level to lower-releases
energy
energy
B. Problems with the Bohr Model
• OOPS!- Model only works
with hydrogen
• Did not account for the
chemical behavior of atoms
• WRONG: Electrons do not
move around the nucleus in
circular orbits
• Still very helpful!!
II. How do Neon Signs work?
• They have
“excited” gases
in them.
Explanation
Step 1: an electron absorbs
energy and moves to a
higher energy level
Step 2: e- drops back down
to a lower energy level
Step 2
Step 1
•During drop it gives off
energy called a “photon”
•Sometimes this energy is
visible light (ROYGBIV)
• When a photon is emitted, energy is released. We can
calculate the energy released using the equation: E = h  ν
Application: Atomic Emission Spectrum
• Used to determine which elements are present in a
sample
• Used to determine which elements are present in a
star (because stars are gases)
• Each element has a unique spectrum
• Only certain colors are emitted because the energy
released relates to specific frequency
Spectroscope
• A spectroscope is needed to see the atomic
emission spectra, which acts similar to a
prism, separating different frequencies of light
Electromagnetic Spectrum
• Electromagnetic spectrum is the range of all
energies emitted from photons acting like waves.
Electromagnetic Spectrum with Visible
Light Spectrum
Lab:
Atomic
Emission
Spectra of
Several
Gases
Light
• Behaves like a particle
• Behaves like a wave
Characteristics of a Wave
• Wavelength  (lambda) – shortest distance between equivalent points on
a continuous wave [Unit = meters]
• Frequency  (nu) – the number of waves that pass a given point per
second [Unit = 1/second = s-1 = Hertz (Hz)]
• Crest – Highest point of a wave
• Trough – Lowest point of a wave
• Amplitude (a)– height from its origin to its crest (highest point) or trough
(lowest point) [Unit = meters]
Amplitude
Amplitude
(Wavelength)
(Wavelength)
Wavelength and Frequency
• Wavelength () and frequency () are related
• As wavelength goes up, frequency goes down
• As wavelength goes down, frequency goes up
• This relationship is inversely proportional
Wavelength and Frequency cont.
c = 
= c / 
=c/
c
Speed of light

wavelength

frequency
c = 
8
Speed of light (c) = 3 x 10 m/s
Question Time
• Calculate the wavelength () of yellow light if
its frequency () is 5.10 x 1014 Hz.
c


Question Time
• What is the frequency () of radiation with a
wavelength () of 5.00 x 10-8 m? What region
of the electromagnetic spectrum is this
radiation?
c


How Much Energy Does a Wave Have?
•
•
•
•
•
•
•
Energy of a wave can be calculated E
Energy
Use the formula E= h
E= Energy
h
Planck’s
constant frequency
 = frequency
h = Planck’s constant = 6.626 x 10-34 Joule . Sec
Joule is a unit for energy (J)
Energy and frequency are directly proportional,
as frequency increases, energy increases

Question Time
• Remember that energy of a photon given off by
an electron is
E =h
• How much energy does a wave have with a
frequency of 2.0 x 108 Hz? ( h = 6.626 x 10-34 J.s)
E = 1.3 x 10-34 Joule
Visible Light, Frequency, and Energy
• Red: longest wavelength (),
smallest frequency ()
• Red: frequency smallest (),
least amount of energy (E)
• Violet: smallest wavelength (),
largest frequency ()
• Violet: frequency largest (),
greatest amount of energy (E)
Flame Test
• The flame test is a way to determine the element
present in a sample
• When placed in a flame, each element gives off a
different color
• Operates same as neon signs; electrons excited
by heat and fall back down and give off
different colors.