Understanding Light: Wavelengths, Photons, and Quantum Physics

1. wavelength Visible light Amplitude wavelength Node Ultraviolet radiation Electromagnetic Radiation

2. Figure 7.1

3. Wave motion: wave length and nodes

6. 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. 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 = 6.6262 x 10-34 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 10-18 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