1. Light, Photon Energies, and Atomic Spectra

2. Electrons and Waves
Louis deBroglie proposed the dual nature of matter, which means that matter has matter-like properties and wave-like properties.
What are wave-like properties?
Wavelength
Frequency
Energy
These are other scientists who made contributions that allowed the current atomic model to be developed.

3. Wave Properties—Wavelength
Wavelength () is the distance from two consecutives higher or lower points of a wave (measured in meters “m” or “nm”)
wavelength

4. Wave Properties—Frequency
Frequency () is the number of times a wave completes a cycle in one second (cycles per second is “Hertz” or “Hz” or 1/ s or s-1)
Lower frequency
Higher frequency

5. Light is Electromagnetic Radiation
Electromagnetic energy is energy that has electric and magnetic fields
There are many types of Electromagnetic Radiation…visible is just one type!

6. Electromagnetic Spectrum
Longer Wavelength (l)
Smaller Frequencies (n)
Less Energy (E)
Shorter Wavelength (l)
Larger Frequencies (n)
Higher Energy (E)
Roy G. Biv
Longer Wavelength (l)
Smaller Frequencies (n)
Less Energy (E)
Shorter Wavelength (l)
Larger Frequencies (n)
Higher Energy (E)

7. Quick Check: Electromagnetic Spectrum
Which type of wave in the electromagnetic spectrum has the greatest energy?
Which type of wave has the longest wavelength?
List the waves of the visible spectrum in order of increasing energy.
Gamma Rays
Radio Waves
(Lowest  Highest)
Red, Orange, Yellow, Green, Blue, Violet

8. Relationship between Wavelength and Frequency
As the wavelength increases, the frequency of the wave decreases.
Important Note(s)
Wavelength must be in meters.
Frequency must be in Hertz.

9. Example #1 c =f l 3.00 x 108 m/s = (7.42 x 1014 s-1) l 7.42 x 1014 s-1
A purple light has a frequency of 7.42 x 1014 s-1.  What is its wavelength?
c =f l
3.00 x 108 m/s
= (7.42 x 1014 s-1) l
7.42 x 1014 s-1
7.42 x 1014 s-1
l = 4.04 x 10-7 m

10. Example #2 (3.18 x 10-7 m) 3.00 x 108 m/s = f 3.18 x 10-7 m
Certain elements emit light of a specific wavelength when they are burned. For example, silver emits light with a wavelength of 3.18 x 10-7 m. Determine the frequency of the wave emitted by silver.
c =f l
3.00 x 108 m/s
= f
(3.18 x 10-7 m)
3.18 x 10-7 m
3.18 x 10-7 m
f = 9.43 x 1014 s-1 or Hz

11. Example #3: Now You Try! 3.00 x 108 m/s= (5.09 x 1014 Hz) l
The yellow light given off by a sodium vapor lamp used for public lighting has a frequency of 5.09 x 1014 Hz. What is the wavelength of this radiation?
c =f l
3.00 x 108 m/s= (5.09 x 1014 Hz) l
l = 5.89 x 10-7 m

12. Relationship between Frequency and Energy
As frequency increases, the energy of the wave increases.
Important Note
Frequency must be in Hertz.

13. Example #4
What is the energy of a photon if it has a frequency of 6.82 x 1017 Hz? f
E = hf
E =
(6.63 x J•s)
(6.82 x 1017 Hz)
E =
4.52 x J

14. Example #5: Now You Try!
Determine the frequency of a wave that has a energy of 8.72 x J. E
E = hf
8.72 x J
= (6.63 x J•s) f
f = 1.32 x 1016 Hz

15. Putting the Two Together
Example #6
What is the energy of a photon of blue light that has a wavelength of 4.5 x 10-7m?
(6.63 x J•s)
(3.00 x 108 m/s)
E =
= 4.4 x J
4.5 x 10-7m

16. How color tells us about atoms Atomic Emissions
Atomic Spectrum
How color tells us about atoms
Atomic Emissions

17. A Closer Look at the Spectra and Bohr’s Model of the Atom
Energy in Atoms
The underlined phrase at the bottom of the page relates to an important animation. You will need the following login information.
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The animation shows energy being added to an electron. The energy comes in the form of a wave. When enough energy is added to an electron, it becomes excited or moves to an excited state which is farther from the nucleus. However, excited state electrons are unstable and tend to drop back down in energy levels (move closer to the nucleus). To do this, they must release energy, which often comes in the form of light. In this animation, it comes as different colors.
The two spectra on the slide show for one particular element what specific wavelengths or types of light that it will absorb and emit. Each element has their own unique spectrum, which allows us to use these spectra to identify elements.
They will then do the flame tests, which doesn’t demonstrate the complete spectrum for an element, but it does allow the students to see one characteristic line.
A Closer Look at the Spectra and Bohr’s Model of the Atom

18. What does this have to do with electron arrangement in atoms?
When all electrons are in the lowest possible energy levels, an atom is said to be in its GROUND STATE.
When an atom absorbs energy so that its electrons are “boosted” to higher energy levels, the atom is said to be in an EXCITED STATE.

19. Atomic Spectrum Each element gives off its own characteristic colors.
Can be used to identify the atom.
That is how we know what stars are made of.
Bright Line Emission Spectra
Energy Levels and Spectra movie

20. These are called line spectra
unique to each element.
These are emission spectra