Lecture 20
Light and Quantized Energy
Ozgur Unal
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Rutherford’s atomic model was not successful at explaining
the locations of electrons around the nucleus.
It also failed to explain the chemical behavior of elements in
the periodic table.
Example: Lithium, Sodium and Potassium have similar
chemical properties although they belong to different periods.
In order to understand atoms we need to consider the
quantum nature of particles, such as electrons, light etc.
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What do you think light is?
What descriptive words can you use for light?
Visible light is a type of electromagnetic radiation.
Electromagnetic (EM) radiation is a form of energy that
shows wavelike behavior as it travels.
Microwaves and X rays are examples of EM radiation.
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How can you describe waves?
The wavelength of a wave is the shortest distance between
equivalent points on a continuous wave.
The symbol for wavelength is λ.
It can be measured in meters, centimeters, nanometers etc.
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The frequency of a wave, denoted by ν,
is he number of waves that pass a given
point per second.
Hertz (Hz) is the SI unit of frequency.
1 Hz = 1 wave/second = 1/s = 1 s-1.
The amplitude of a wave is the wave’s height from the origin
to a crest, or from the origin to a trough.
The wavelength and
frequency does not affect
the amplitude of a wave.
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Compare the three waves below.
Which has higher amplitude,
frequency and wavelength?
All waves move, so they have speed.
All EM waves has a speed of 3 x 108 m/s in vacuum.
This speed is a universal value and denoted by c.
The speed can be expressed using frequency and wavelength
c=λ*ν
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White light passing through a prism separates into a
continuous spectrum as shown in the picture.
Each color has a unique wavelength and frequency.
Each color refracts at a different angle.
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The visible spectrum of light shown is only a small portion
of the complete electromagnetic spectrum.
The EM spectrum includes all forms of EM radiation, with
different frequencies and wavelengths.
Because all EM waves travel at the same speed, you can use
the formula c = λ * ν to calculate the wavelength or frequency
of any wave.
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Lecture 21
The Particle Nature of Light
Ozgur Unal
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Wave nature of light cannot explain some Physics
phenomena.
Objects glowing in different colors as they are heated up to
high temperatures
Photoelectric effect
When objects are heated, they emit glowing light.
As the temperature (average kinetic energy)
increases the color (frequency) of the
light changes from red to blue.
Wave model of light could not
explain this phenomenon.
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Wave model could not explain photoelectric effect either.
In the photoelectric effect,
electrons, called photoelectrons,
are emitted from a metal’s surface
when light of a certain frequency,
or higher than a certain frequency,
shines on the surface.
Wave model predicts that given enough time, low energy
(frequency) light would accumulate and supply enough
energy to eject photoelectrons from a metal.
In reality, a metal will not eject photoelectrons below a
specific frequency of incident light.
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At the beginning of 1900s, Max Planck was
able to explain the first phenomenon.
He proposed that matter can gain or lose energy
only in small, specific amounts called quanta.
A quantum is the minimum amount of
energy that can be gained or lost by an atom.
This means that, although the temperature seems to rise
continuously as you heat an object, in fact it rises in small,
discrete amounts.
Since this increase is so small, we perceive it as continuous.
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Planck also proposed that the energy of a quantum is related
to the frequency of the emitted (or absorbed) radiation by a
constant caleld Planck’s constant.
Equantum = h ν
h is Planck’s constant. h = 6.626 x 10-34 J*s
According to Planck’s theory, for a given frequency ν, matter
can emit or absorb energy only in whole-number multiples of
hν; that is 1hν, 2hν, 3hν and so on.
Analogy: This is similar to building a wall.
Only whole-numer multiples of bricks can be added to or
taken out from a wall.
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To explain the photoelectric effect, Albert Einstein proposed
in 1905 that light has a dual nature.
Light has wavelike and particlelike
properties.
Light can be thought of as a beam
of particles called photons.
A photon is a massless particle that
carries a quantum of energy.
Einstein extended Planck’s idea of quantized energy to light:
Energy of a photon = Ephoton = h ν
 Here, ν is the frequency of light.
The concept of photon explains photoelectric effect.
Photoelectric Animation
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Example: Sun not only emits visible light but also other
electromagnetic waves including radiowaves, microwaves, UV
etc. Calculate the energy of a photon from the UV portion of
the Sun’s light if it has a frequency of 9.231 x 1014 Hz.
Example: The blue color in some fireworks occurs when
copper (I) chloride is heated to approximately 1500 K and
emits blue light of wavelength 4.50 x 102 nm. How much
energy does one photon of this light carry?
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How is light produced in
neon signs?
When electricity passes through a tube filled with Ne atoms,
these atoms absorb energy and become excited.
These excited atoms emit radiation (light) to become stable.
This is how neon signs work.
If you pass the light emitted from Ne atoms through a prism,
you will see atomic emission spectra of Ne atoms.
Atomic emission spectrum of an element is the set of
frequencies of EM waves emitted by atoms of the element.
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Atomic emission spectrum of certain elements
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If atoms emits radiation of specific frequencies, that means
they absorb the same frequencies of light.
This gives us absorbtion spectrum.
Absorption spectrum of H
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By looking at the absorption
spectrum of the Sun, we can
find out the elements present
in the atmosphere of the Sun.
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