Understanding Electromagnetic Radiation in Quantum Physics

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

3. Figure 7.1

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

7. 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

8. 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

9. 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

10. E = h •  Relationships:

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

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

13. 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?

14. 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?

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

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 emitted by solids

18. Light emitted by 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. Emission spectra of gaseous atoms • Excited atoms emit light of only certain wavelengths • The wavelengths of emitted light depend on the element.

21. Line spectra of atoms

22. Energy Adsorption/Emission

23. Constant = 2.18 x 10-18 J For H, the energy levels correspond to: Energy level diagram:

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

25. Explanation of line spectra Balmer series

26. 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:

27. 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:

28. Can see matter waves in experiments

29. 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

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

31. Example Wavefunction Equation slightly simplified:

32. 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

34. Shells, Subshells and Orbitals

35. Which subshell does not exist? • 5s • 2p • 2d • 4f • 15s

38. NODES Spherical Nodes

39. Quantum Numbers and Numbers of Orbitals