1. Chapter 6: Electromagnetic Radiation
wavelength
Visible light
Ultraviolet radiation
Amplitude
Node

6. Figure 7.1

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

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

11. 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

12. 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

13. FCC Spectrum

14. 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

15. E = h • 
Relationships:

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

17. E = h • n
What is the energy of a WONY photon?

18. Energy of Radiation
What is the frequency of UV light with a wavelength of 230 nm?
What is the energy of 1 photon of UV light with wavelength = 230 nm?

19. Can this light break C-C bonds with an energy of 346 kJ/mol?
What is the energy of a mole of 230 nm photons?
Can this light break C-C bonds with an energy of 346 kJ/mol?

20. Does 1200 nm light have enough energy to break C-C bonds?

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

22. Light emitted by solids

23. Light emitted by hydrogen gas

24. 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

26. Emission spectra of gaseous atoms
Excited atoms emit light of only certain wavelengths
The wavelengths of emitted light depend on the element.

27. Line spectra of atoms

28. Energy Adsorption/Emission

29. For H, the energy levels correspond to:
Constant = 2.18 x J
Energy level diagram:

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

32. Explanation of line spectra
Balmer series

33. 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:

34. 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:

35. Can see matter waves in experiments

36. 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

37. 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

38. GPM

39. Example Wavefunction
Equation slightly simplified:

40. 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 orbital
p = 3 orbitals
d = 5 orbitals
f = 7 orbitals

42. “Allowed” Orbitals

43. NODES
Spherical Nodes

44. Quantum Numbers and Numbers of Orbitals