1. Electromagnetic Radiation
wavelength
Visible light
Ultraviolet radiation
Amplitude
Node

2. Figure 7.1

3. Wave motion: wave length and nodes

6. Long wavelength --> small frequency low energy
Short wavelength -->
high frequency
high energy
Long wavelength -->
small frequency
low energy

7. Rank the following in order of increasing frequency: microwaves radiowaves X-rays blue light red light UV light IR light

8. Long wavelength --> small frequency
Waves have a frequency
Use the Greek letter “nu”, , for frequency, and units are “cycles per sec”
All radiation:  •  = c
c = velocity of light = 3.00 x 108 m/sec
Long wavelength --> small frequency
Short wavelength --> high frequency

9. What is the wavelength of WONY?
What is the wavelength of cell phone radiation? Frequency = 850 MHz
What is the wavelength of a microwave oven? Frequency = 2.45 GHz

10. Quantization of Energy
Light acts as if it consists of particles called PHOTONS, with discrete energy.
Energy of radiation is proportional to frequency
E = h • 
h = Planck’s constant = x J•s

11. E = h • 
Relationships:

12. Rank the following in order of increasing photon energy: microwaves radiowaves X-rays blue light red light UV light IR light

13. E = h • 
What is the energy of a WONY photon?

14. Energy of Radiation
What is the energy of 1 mole of UV light with wavelength = 230 nm?

15. Energy of Radiation
What is the energy of 1 mole of IR light with wavelength = 1200 nm?

16. Where does light come from?
Excited solids emit a continuous spectrum of light
Excited gas-phase atoms emit only specific wavelengths of light (“lines”)

17. Light given off by solids

18. Light given off by Excited Hydrogen Gas

19. The Bohr Model of Hydrogen Atom
Light absorbed or emitted is from electrons moving between energy levels
Only certain energies are observed
Therefore, only certain energy levels exist
This is the Quanitization of energy levels

20. Line Emission Spectra of Excited Atoms
Excited atoms emit light of only certain wavelengths
The wavelengths of emitted light depend on the element.

21. Line Emission Spectra of Excited H Atoms
High E
Short 
High 
Low E
Long 
Low 

22. Line Spectra of Other Elements

23. Atomic Absorption and Emission

24. Origin of Line Spectra
Balmer series

25. For H, the energy levels correspond to:
Constant = 2.18 x J

26. Each line corresponds to a transition:
Example: n=3  n = 2

27. Name: ____________
_________

28. Quiz
Q1. Emission line with longest wavelength Q2. Absorption line with highest frequency Q3. Emission line with lowest frequency Q4. Transition that leads to forming H+

29. Matter Waves All matter acts as particles and as waves.
Macroscopic objects have tiny waves- not observed.
For electrons in atoms, wave properties are important.
deBroglie Equation:

30. Matter waves
Macroscopic object: 200 g rock travelling at 20 m/s has a wavelength:
Electron inside an atom, moving at 40% of the speed of light:

31. Can see matter waves in experiments

32. Heisenberg Uncertainty Principle
Can’t know both the exact location and energy of a particle
So, for electrons, we DO know the energy well, so we don’t know the location well

33. Schrodinger’s Model of H
Electrons act as standing waves
Certain wave functions are “allowed”
Wave behavior is described by wave functions: 
2 describes the probability of finding the electron in a certain spot
Also described as electron density

34. Example Wavefunction
Equation slightly simplified:

35. It’s all about orbitals
Each wavefunction describes a shape the electron can take, called an ORBITAL
Allowed orbitals are organized by shells and subshells
Shells define size and energy (n = 1, 2, 3, …)
Subshells define shape (s, p, d, f, …)
Number of orbitals is different for each subshell:
s = 1
p = 3
d = 5
f = 7

39. Quantum Numbers