1. 5 Minutes to Finish Sheets – prepare a 15 to 30 sec blurb
Class Avg
% A’s
% B’s
% C’s
% F’s
2nd 76 27 9 36
3rd 85 42 43 5 10
4th 74 16 32
5th 73 25 20 45
6th 17
All
76.8
22.4
27.2
21.6
28.6

2. The Electromagnetic Spectrum
The electromagnetic spectrum is a series of waves that have different wavelengths.

3. Visible Spectrum
The visible spectrum is continuous (there are no breaks and the colors blend together).
White light is a combination of ALL colors of light. A prism breaks up white light into the separate colors so we can see them.
Each color has a definite frequency and wavelength.
The speed of these colors never changes – always speed of light (c)

4. Visible spectrum- Frequency & Wavelength
low energy colors
high energy colors
Red
Orange
Yellow
Green
Blue
Indigo
Violet

5. EM spectrum in real life!
UV Light
Television
FM Radio
Infrared Waves
Microwaves
& Radar
AM Radio
Visible Light
Telemetry
Millimeter waves
Shortwave Radio
X-rays and Gamma rays

6. Radio waves/TV/Short Wave are used in “classic” televisions and in radio.
Radar wave are what we know as the old “doppler” waves. These images are used from things such as weather or in the military.
Infrared Waves detect heat. These waves are used in night vision goggles.
UV waves are used in tanning beds, to kill bacteria, and irradiate food.
X-rays are used to see inside the body.
Waves can penetrate through the skin, but bounce off the more dense bone producing a reflection on the x-ray.
Gamma Rays are used in radiation therapy and gamma-surgery

7. Wave Model
Light consists of electromagnetic waves that travel at 3.00 x 108 m/s
That’s miles per hour!
That’s like 1200 lunar orbits around the earth a day!

8. Wavelength and Frequency
All EM waves move at the speed of light.
c = fλ
c = speed of light =
3.00 x 108 m/s
f = frequency (Hz)
λ = lambda = wavelength (m)
As wavelength increases,
frequency __________.

9. EM Waves
Frequency – number of cycles that pass a given point in a given amount of time.
Measured in Hertz (Hz)
1 Hz = 1 wave passes per second
400 Hz = 400 waves pass per second
½ Hz = ½ wave passes per second

10. Practice Problems c = f λ
If an EM wave has a wavelength of 630 nanometers, what is its frequency? (1 nm = 10-9 m)
If an EM wave has a frequency of 548 Hz, what is its wavelength?
c = 3.00 x 108 m/s
f = 548 Hz
λ = ?
c = 3.00 x 108 m/s
f = ?
λ = 6.30 x 10-7 m
c = f λ
3.00x108 m/s=(548 Hz)λ
c = fλ
3.00x108 m/s=f(6.30x10-7 m)
λ = m
f =4.76 x 1014 Hz

11. Practice problem λ = 6.85 x 10-6 m
Find the wavelength of light with a frequency of 4.38 x 107 MHz.
c = 3.00 x 108 m/s
f = x 107 MHz
λ = ?
c = fλ
3.00 x 108 m/s = (4.38 x 1013 Hz)(λ)
λ = 6.85 x 10-6 m

12. Particle Model The idea that light can act as a particle.
These particles are called photons, or quanta, and can move other matter.

13. Photoelectric Effect
The particle model was needed to explain why when you shine a high energy light on metal, electrons are ejected (moved) from the metal.
This powers solar-powered calculators.

14. Einstein Einstein proposed in 1905 that light can behave as both a
wave and a particle.
He defined a photon is a particle of electromagnetic radiation with no mass and carries a quantum of energy.

15. The energy contained in a photon (a quantum) depends on its frequency.
E = hf
E = energy (units are in Joules)
h = Planck’s constant = 6.63 x J.s
f = frequency (Hz)

16. Practice Problem E = hf
If a wave has a frequency of 230 Hz, what is the amount of energy of one photon (quantum) of this wave?

17. Practice Problem E = hf
If a wave has a frequency of 230 Hz, what is the amount of energy of one photon (quantum) of this wave?
E = hf
E = (6.626 x Js)(230 Hz)
E = 1.5 x J

18. Practice Problem E = hf, c = fλ
If a wave has a wavelength of 400. nm, what is the amount of energy of one photon (quantum) of this wave?

19. Practice Problem E = hf, c = fλ
If a wave has a wavelength of 400. nm, what is the amount of energy of one photon (quantum) of this wave?
400. nm = 4.00 x 10-7m
f = c/λ
E=hf
So, E = hc/λ = [(6.626 x 10-34Js)(3.00x108m/s)]/(4.00x10-7m)

20. Relating electrons to light
Remember that electrons occupy energy levels - the region surrounding the nucleus where an electron is likely to be found (think of rungs on a ladder, fixed levels with space in between)
When electrons are in their lowest energy level, they are said to be in their “ground state”
It is possible for electrons to jump from ground state to a higher energy level (called excited state) by absorbing energy.

21. Relating electrons to light
When electrons gain energy somehow, they can then occupy a HIGHER ENERGY LEVEL. They jump up to the next level and are said to be in their “excited state”.
When electrons lose energy they will fall back down to their GROUND STATE and release energy, and some of it is released as waves we can see – LIGHT!

22. Relating electrons to light
The electron configuration of sodium at ground state is 1s22s22p63s1.
The electron configuration of sodium at an excited state is 1s22s22p53s2.
What happened?

23. Practice Write the electron configuration of the following elements in their ground and excited states:
Element
Sodium
Fluorine
Oxygen
Ground state electron configuration
Ion electron configuration
Excited state electron configuration

24. Atomic Emission Spectrum
Continuous Spectrum (no breaks)

25. EM Spectroscopy
The use of a spectroscope to observe this EM radiation is called EM spectroscopy.
This EM radiation can be classified as
either an absorption spectra
or emission spectra.

26. Atomic Emission Spectrum

27. Absorption Spectra “Dark line spectra”
Electrons can jump to a higher energy level, but the atom must absorb energy in order to reach this excited state

28. Absorption spectra are different for each element
Hydrogen
Sodium

29. Emission Spectra “Bright line spectra”
When electrons leave the excited state and return to the ground state, they emit energy

30. Emission spectra are different for each element
Hydrogen
Iron

31. How do we use the absorption and emission spectra?
Each element gives off unique bright line and dark line spectra
Can be used to determine the composition of stars, comets, and planets by analyzing the received light
Can be used to determine the concentration of chemical compounds in a sample
Reports of its use in differentiating malignant tumors from benign.