1. Today is Friday (!), February 26th, 2016
In This Lesson:
Atomic Emissions
(Lesson 2 of 4)
Today is Friday (!), February 26th, 2016
Pre-Class: [choose one]
What is white light?
How are fireworks made to be different colors?
How are neon signs made to be different colors?

2. Today’s Agenda Atomic Emissions Flame Tests The Light Spectrum
Waves and Particles
In other words, we’re going to talk about how atoms give off light, then about light itself…and waves…
Where is this in my book?
P. 138 and following…

3. By the end of this lesson…
You should be able to explain what happens when energy is applied to an atom.
You should be able to describe the relationships between wavelength, frequency, and energy.

4. Emission Spectra Fireworks are exciting because of:
The noise they make.
The variety of color they display.
We’re going to focus on the color.

5. Identification
The multicolored lights created by fireworks occur because of the different elements that comprise the powder in fireworks.
Fireworks Filmed with a Drone video

6. Identification
Scientists have found that each element, when heated, gives off its own specific set of colors.
The element’s colors are its “fingerprints” and can be used to identify the element.

7. Element Colors Element Flame Color Sodium Yellow Potassium Violet
Rubidium
Pinkish-Red
Calcium
Orange-Red
Strontium
Red
Barium
Green
Copper
Blue-Green

8. Cesium
Blue

9. Calcium
Deep Orange

10. Sodium
Orange

11. Potassium
Violet

12. Copper
Jade Green

13. Flame Tests
Many elements give off characteristic light which can be used to help identify them.
Strontium
Sodium
Lithium
Potassium
Copper

14. Atomic Emission Spectrum of Barium using a Spectrometer
Electron Energy State
Electrons absorb energy from the flame.
When a certain amount is reached (a quantum), they jump to a higher energy level: the “excited state.”
Eventually, the electrons lose the energy in the form of light and fall back to the lowest, most stable energy level: the “ground state.”
Atomic Emission Spectrum of Barium using a Spectrometer

15. Lyman, Balmer, Paschen Series

16. Electromagnetic Spectrum

17. Sources of Energy
Where do electrons get energy to “jump” to the next higher energy level?
Collisions from other particles
Heat
Electricity
Light

18. Loss of Light?
As we learned, when electrons fall back to the ground state, they release energy in the form of light.
It’s complicated, but light can behave as a wave or a particle.
As a particle, a “unit” of light is called a photon.
Additionally, a quantum (plural: quanta) is the amount of energy needed to move an electron into an excited state.
A quantum of light is called a photon.

19. Particle-Wave Duality

20. Wave Statistics
Amplitude: The “height” of the wave from zero to crest (peak).
Wavelength: Distance between peaks in nanometers (nm) or meters (m).
Given by Greek letter λ (lambda).
Frequency: The number of cycles (wave peaks) that occur in a unit of time (per second or Hertz; Hz)
Given by Greek letter v (nu).

21. Wavelength, Frequency, and Energy
Long
Wavelength =
Low Frequency
Low Energy
Short
Wavelength =
High Frequency
High Energy

22. Takes less energy to do these hills…
How to remember?
How can you remember “high frequency = high energy?”
Imagine riding a bike over the wave peaks!
Takes less energy to do these hills…
…than to do these hills.
-Litz, 2014

23. Summary Electrons can move between energy levels.
Ground state: stable state; an electron is at the lowest energy level.
Excited state: unstable state; an electron is at a higher energy level.
Quantum: the amount of energy needed to move an electron from the ground to excited state.
Photon: a quantum of light.

24. Summary Wavelength and frequency are inversely related:
When wavelength increases, frequency decreases.
Frequency and energy are directly related:
When frequency increases, energy increases.
We only see a small part of all possible wavelengths/frequencies.
The visible spectrum.

25. Summary Variables: v (nu) – measure of frequency in hz (cycles/sec).
λ (lambda) – measure of wavelength in nm.
v (nu) – measure of frequency in hz (cycles/sec).

26. Summary: Emissions in Real Life
The reason most streetlights look a little “orange” is because they pass an electric current through sodium vapor.
Remember how sodium burns in orange color?
Compare LED light to Na vapor:

27. So now then… Let’s try some flame tests!
At each of your lab tables is one of seven different kinds of salt solutions.
This isn’t table salt.
In the salt solution is a wooden splint that has been soaking in it overnight.
You should take out your Bunsen burner (if it’s not already out) and light it.

28. Flame Tests I will turn off the lights.
At that point, each group will put ONE of the splints into the flame and record the color that is emitted.
The lights will come back on, and groups will rotate clockwise until all solutions have been tested.
There will be time for answering the questions that follow.

29. What NOT to do… Don’t let the splint burn.
Don’t place more than one splint into the flame.

30. Closure Which has higher energy, long or short wavelength?
Short wavelength (high frequency).
Exactly what is burning?
The various salt solutions (NOT the splint)
Did the electrons get closer to the nucleus or further away?
Further
What could we say happened to the electrons in terms of their Principal Quantum Number?
They briefly entered a higher energy shell (or principal quantum number) before falling back into their ground levels.