1. Images:

2. Objectives… To use LED lights to… Understand how light is produced
Understand how LED lights work
Calculate wavelengths of light
Calculate frequencies of light
Measure voltage and convert it to energy
Calculate Planck’s Constant
Understand basics of Quantum Theory and how it relates to light

3. What do you know about light?
What is light?
Ask students what they already know about light then show the video “What the Heck is Light.” After the video go to the next slide and discuss that light is caused by “excited electrons.”
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Video: What the Heck is Light”
This song explains how light is formed, what the different types of electromagnetic radiation are and the significance of bright-line spectra. Lightning filmed June of Music and lyrics copyright 2007 by Mark Rosengarten. All rights reserved.

4. Different Forms of Light
Incandescent light bulb (and Halogens)
Florescent lamps
Neon lights
LED’s
All require “Excited ELECTRONS”
Review video and discuss that light is caused by “Excited Electrons.”
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YIPPEEEEEE!!!

5. The Bohr Model of Hydrogen
Neils Bohr came up with the idea of electrons in “fixed energy levels”
Discuss how Bohr determined that electrons were found in energy levels and the movement of electrons to an excited state and back to ground state emits photons of light. STRESS that the actual model of an atom is the Quantum model in the lower right hand corner and that the energy levels aren’t rings but areas around the nucleus where electrons are expected to be found.
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6. LED Lights LED- “Light-Emitting Diode”
Activity 1 “Get to Know Your LED Kit”
You have 10 minutes…
Have students read the directions to the student handout for Activity 1: “Get to Know your LED Kit” Give students about 5 minutes and discuss what they found.

7. Bohr Model Continued…
When electricity was passed through Hydrogen Gas, it glowed. The spectra of color looked like this:
Help students understand what Bohr saw by viewing different gas discharge tubes through the diffraction gradients. At a minimum show hydrogen and use this slide to show the students what they should see. Show as many tubes as available or as time allows.
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8. Every Element Has It’s Own Emission Spectra
Point out that each element has a specific emission spectra due to excited electrons emitting photons when they return to their ground state. An emissions spectra is a “finger print” of an element and is useful in determining the composition of mixtures. Compounds will also give off a specific emissions spectra.
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9. Light is Caused by Excited Electrons
Review concepts on this slide and next and point out the Noble Gases.
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10. What Bohr Concluded…
As the electrons absorb energy they become “excited” and move to a higher energy level
When the electrons move back to the original “ground state,” the energy absorbed is released as a photon of light
The color of light corresponds to a particular wavelength and frequency and therefore a specific, discrete amount of energy released.
Photon emitted
photon emitted
A photon is a “quanta” of energy because it occurs in specific, discrete amounts. Hence the term, “Quantum Model” of an atom.
Bohr’s conclusion.
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11. The Electromagnetic Spectrum
Review the entire EM Spectrum. Point out wavelength and frequency.
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12. Wave On a String
Use the “Wave on a String” applet to review parts of a wave. Set Amplitude and Frequency to 50, and Damping to 0. Slide the tension to “high.” Select “No End” and Oscillate. Show the students what happens to wavelength when the frequency is changed. Click on “Rulers” measure wavelength and amplitude. Show that changing amplitude doesn’t change the wavelength or frequency.

13. Activity 2 Measuring the Wavelengths of Light 20 minutes
Have the students follow the directions for Activity 2 on the student answer sheet. A picture of the set up is on the next slide.

14. Set Up

15. Activity 3 Finding the Voltage and Crunching the Numbers 15 minutes

16. What About Light from an LED?
Made mostly of semiconductor elements
The Group 14 elements are the Key!
Light is still caused by EXCITED ELECTRONS!
Review location of group 14 on Periodic Table.
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17. Si As Al Valence Electrons 3 4 5 • • • • • • • • • • •
Review location of group 14 on Periodic Table and point out the number of valence electrons.
Image: • Al •
• • Si As • • • • • • •

18. Silicon atom
Here is silicon covalently bonded with other silicon atoms. Note the 8 valence electrons in each silicon atom.
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19. Mobile Electrons and Holes?
Using group 13 elements because they have 1 less valence electron, “holes” are created where electrons can be. Since this creates a net positive charge, this is referred to as “P-type.”
Using group 15 elements because they have 1 more electron, “Free Electrons” are present. Since this creates a net negative charge, this material is referred to as “N-Type.”
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Image: P N

20. Band Gap Energy The Energy of “Excited Mobile Electrons.”
As the Excited Mobile Electrons “fall into ground state holes” that energy is released as photons of light!
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21. Periodic Properties Ga Ga Ga As N Ultraviolet Green Infrared
Since the bond energy is greater between Ga and more energy is released when the free electrons complete these bonds. This greater energy is due to the fact that nitrogen is a smaller and more eletronegative atom than P and As, and therefore holds electrons more tightly.
Image: Ga As
Infrared
Bond energy determines Band Gap Energy

22. “Tie it all together…”
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23. Energy Divided by Frequency
LED Data Table
Color
Actual Wavelength
(nm)
Voltage
(V)
Frequency
(1/s or Hz)
Energy
(J)
Energy Divided by Frequency
(J•s)
INFRARED
990
1.121 
3.o x 1014 
1.8 x 10-19  
6.0 x  
 RED 
 660
1.831 
 4.5 x 1014
2.9 x 10-19 
6.5 x 10-34 
ORANGE 
 620
 1.865
4.8 x 1014 
5.9 x 10-34 
YELLOW 
 590
1.898 
 5.1 x 1014
3.0 x 10-19 
GREEN 
 570
 2.036
 5.3 x 1014
3.3 x 10-19 
6.1 x 10-34 
BLUE 
430
3.348
 7.0 x 1014
5.4 x 10-19 
7.6 x 10-34 
Ask the students if they found similar data. Point out the last column. Ask if the values pretty similar. Ask what a good “ball-park” estimate would be. Ask if the number is big or small? Tell them they have discovered a fundamental constant in the world of chemistry, physics and Quantum Mechanics.

24. Max Planck The Father of Quantum Physics
Energy = Planck’s constant x wavelength
h = x J•s
Before 1900, light was considered to be a wave but explaining why light was emitted at different colors was not yet understood. Planck knew that the energy of light was related to the wavelength and used a mathematical trick in his equation by adding a constant, h. He assumed that he could use the constant just to set up the problem and then use a common mathematical technique to get rid of it. Much to his surprise, however, the results only made sense if he kept the constant around- if h had a very small but non-zero number. The result is that light can only occur in discrete “chunks” like particles. Every day waves can occur at any value but light is only emitted at discrete units of energy. Planck referred to these specific energy levels as “quanta” (the plural form of quantum from the Latin word for “how much.”) Therefore light at a given frequency might contain one quantum (one unit of energy, hf), two quanta, three quanta, and so on but never 2.5, 3.75, etc. The name of the quantum energy levels stuck, and came to be applied to the entire theory that grew out of Planck’s mathematical trick. Ironically, Planck didn’t like this idea that light came in discrete quanta but the constant he developed has been proven over and over again and helps explain the wave-particle duality of light. (Adapted from How to Teach Physics to Your Dog by Chad Orzel, Chapter 1.)
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Does this number look familiar?