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Transcript
Electromagnetic
Radiation & Light
• What are the atom models we know of?
2
Models of the Atom
A. There were many different models over
time
1. Dalton-billiard ball model (1803)
2. Thompson – plum-pudding model (1897)
3. Rutherford – Nuclear model of the atom
(1911)
4. Bohr – uses quantized energy of the atom
(1913)
5. Quantum Mechanical Model of the Atom
(1926)
• How do the atom models show the
importance of communication and sharing
information in the sciences?
4
B. Each new model contributed
to the model we use today.
Even our current Quantum
Mechanical model, does not
give us an exact model of
how electrons interact.
• Why would you use the simplest for of an
atom to make a model?
• What is the simplest atom?
• What is another name for a circular
pathway?
• What would be the difference if you were
in a grounded state vs an excited state?
• How could that relate to electron orbitals?
6
The Bohr Model of the Atom
A. Bohr used the simplest element, hydrogen,
for his model
B. He proposed an electron is found in
specific circular paths, or orbits around the
nucleus
C. Each electron orbit was thought to have
fixed energy levels
D. Lowest level-ground state
E. Highest level-excited state
• What could quantum mean?
• What would we need to do to excite an
electron?
8
F. One electron is capable of many different
excited states
G. An e- can gain or lose is a specific amount
or quantum of energy
H. You can excite an e- with energy like
electricity, the sun, or magnets
• What is released when an electron goes
to a lower energy level?
• What are photons?
10
Explanation of Atomic Spectra
•When an e- is excited it
moves to a higher energy
level
A.When the e- drops back
down to a lower energy
level, it gives off a quantum
of energy called a “photon”
•Only certain atomic
energies are possible and
emitted
•Photons have both wave
like and particle like
properties
step2
step1
E. When a photon is
emitted, we emit energy,
and use the equation E
=hν
Energy levels are like rungs of a
ladder. You cannot be in between a
rung
Energy levels in an atom’s electron
are unequally spaced. The higher
energy levels are closer together.
• What does spectrum mean? Give an
example
• What would the electromagnetic spectrum
mean?
13
Electromagnetic Spectrum
A. Electromagnetic spectrum is the range of all
energies emitted from photons acting like
waves.
Electromagnetic Spectrum with
Visible Light Spectrum
• What do these terms mean when related
to a wave?
Frequency
Wavelength
Crest
Amplitude
trough
16
Characteristics of a Wave
A. Wavelength  (lambda) – shortest distance between equivalent
points on a continuous wave [Unit = meters]
B. Frequency  (nu) – the number of waves that pass a given point
per second [Unit = 1/second = s-1 = Hertz (Hz)]
C. Crest – Highest point of a wave
D. Trough – Lowest point of a wave
E. Amplitude (a)– height from its origin to its crest (highest point) or
trough (lowest point) [Unit = meters]
(Wavelength)
Amplitude
(Wavelength)
• What is the relationship between
wavelength and frequency?
18
Wavelength and Frequency
A. Wavelength ( ) and frequency ( ) are related
B. As wavelength goes up, frequency goes down
C. As wavelength goes down, frequency goes up
D. This relationship is inversely proportional
c=
=c/
 =c/
c
Speed of light

wavelength

frequency
c=
8
Speed of light (c) = 3 x 10
m/s
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?
How do we calculate the energy of a wave?
What are the units for energy?
How are energy and frequency related?
What happens when an electron goes to a
lower energy level?
How much energy is given off by a wave with a
frequency of 2.0 x 108 Hz? ( h = Planck’s
constant = 6.626 x 10-34 Joule . Sec)
c


• Which colors of light have the most/least
energy?
• Is the high energy light going to have a
higher or lower frequency? Wavelength?
22
Visible Light, Frequency, and
Energy
A.Red
1. longest wavelength ( ),
2. smallest frequency ( )
3. least amount of energy (E)
B.Violet
1. smallest wavelength ( ),
2. largest frequency ( )
3. greatest amount of energy (E)
Flame Test
• The flame test is a way to determine the certain
elements that are present in a sample
• When placed in a flame, different elements give off
different colors
Why would they give off different
colors?
If no colors are emitted could there still be photons
released? How?
Why Do We See Different Colors in
the Flame Test?
A. The electrons are energized to an excited state
B. As electrons drop to lower levels, they give off photons
C. A photon is a particle of electromagnetic radiation with
no mass that carries a quantum of energy
D. If the photon’s frequency corresponds the visible light
range we can see different colors in the flame
E. If it is not in the visible light range, it may be giving off
other forms of electromagnetic radiation like radio,
microwaves, infrared, ultra violet, x-rays, or gamma
waves
• How could we use the flame
test/electromagnetic spectrum?
26
Atomic Emission Spectrum
A. When a current is passed through a vacuum
tube of gas at low pressure, a set of
frequencies of the electromagnetic waves are
emitted by atoms of the element
B. Used to determine which elements are present
in a sample
C. Used to determine which elements are present
in a star
D. Each element has a unique spectrum
E. Only certain colors are emitted meaning only
certain frequencies of light are emitted
Spectroscope
A. A spectroscope is needed to see the
atomic emission spectra, which acts
similar to a prism, separating different
frequencies of light
Problems with the Bohr Model
A. The model is fundamentally
incorrect and only works with
hydrogen.
B. Failed to explain the spectrum
of any other element
C. Did not account for the
chemical behavior of atoms
D. Electrons do not move around
the nucleus in circular orbits
E. Laid ground for later atomic
models.
Problems with the Bohr Model
cont.
Emission Spectrum for
Hydrogen
Absorption Spectrum for Hydrogen