1. I II III  Suggested Reading Pages 97 - 103  Section 4-1 Radiant Energy

2.  Light has characteristics of both waves and particles Dual Nature of Light

3.  Light – electromagnetic radiation  Amplitude  Wavelength  Frequency  Speed Properties of Light, Wave Description of Light

4.  Height of the wave  Measures brightness Amplitude

5.  Distance between crests  Visible light: 400nm (violet) – 750 nm (red) Wavelength

6.  All light moves through space at 3.00 x 10 8 meters per second. Speed

7.  The number of complete waves passing a fixed point in a given time.  Frequency =  = Frequency Speed of Light Wavelength C ___

8. Visible Light  Red – longest  Orange  Yellow  Green  Blue  Indigo  Violet - shortest

9. Electromagnetic Spectrum Heat Lamp

10.  Short Wavelength = High frequency = High energy  Long Wavelength = Low frequency = Low Energy Electromagnetic Spectrum

11. Photoelectric Effect  Electromagnetic Radiation strikes the surface of a metal, ejecting electrons. The flow of electrons creates an electric current.  Light consists of quanta of energy that behave like tiny particles.  Energy quanta = photons.

12. Photoelectric Effect Analogy Coins in … nothing happens until the correct amount is reached. Then a drink is ejected. Like energy in … nothing happens until the THRESHOLD FREQUENCY is reached. Then an electron is ejected.

13. Planck’s Theory  Energy exists in quanta.  Quantum: a small, specific amount of energy that can be gained or lost by an atom.  Plural = quanta

14. Planck’s Theory E = h E = quantum of energy in joules h = Planck’s constant (6.626 x 10 -34 j. s) = frequency of light in Hz or s -1

15. Put it all together:  Using Einstein’s Theory of Relativity formula: E = mc 2  And Planck’s formula: E photon = h  We are able to calculate the apparent mass of a photon.

16. Bohr Model of the H-Atom - 1913  Electrons can exist in one of only a certain number of allowed orbits  Electron’s energy is higher when it is in orbits that are farther from the nucleus.

17. Energy Levels  Electrons in the (ground state) absorb a quantum of energy, jumps to a higher level. (Excited State)  Jumps down to a lower level, releases a quantum of energy, which corresponds to a certain wavelength of light.

18. Energy Levels Each line represents a certain energy level jump.

19. Energy Levels Analogy  Energy Levels are like rungs on a ladder  You can’t stand in mid-air.  Electrons can’t exist in between levels.

20. Bohr’s Model  Only worked for Hydrogen.  Does not work for atoms with more than one electron.

21. Line Spectra – “Atomic Fingerprints”  Contain only certain colors or wavelengths.  Different from continuous spectrum.  Atomic emission spectrum.

22. What is Light – 2 min review