BUSINESS EXAM 2 THURSDAY NOVEMBER 4, 2010 MATERIAL COVERED: CHAPTERS 4, 5 & 6

1. BUSINESS • EXAM 2 THURSDAY NOVEMBER 4, 2010 • MATERIAL COVERED: CHAPTERS 4, 5 & 6 • TIME: 7:00PM-8:00PM • WHERE: (TO BE ANNOUNCED LATER) • WHAT TO BRING: CALCULATOR, ONE PAGE OWN NOTES • CONFLICT IN SCHEDULE? CONTACT ME TO MAKE SEPARATE TIME

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

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

6. The electromagnetic spectrum.

8. Which has the longest wavelength? • Infrared • Ultraviolet • X-rays • Radio waves

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

10. 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

11. In a vacuum, the speed of light, c, is 3.00 x 108 m/s. Therefore, The Wave Nature of Light • The product of the frequency, n (waves/sec) and the wavelength, l (m/wave) would give the speed of the wave in m/s. • So, given the frequency of light, its wavelength can be calculated, or vice versa.

12. The Wave Nature of Light • What is the wavelength of yellow light with a frequency of 5.09 x 1014 s-1? (Note: s-1, commonly referred to asHertz (Hz) is defined as “cycles or waves per second”.)

13. The Wave Nature of Light • What is the frequency of violet light with a wavelength of 408 nm?

14. 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

15. Metal cathode (-) vacuum window Anode (+) The Photoelectric Effect “Light” can cause ejection of e- from a metal surface. An anode (+) attracts e- Current is measured

16. The Photoelectric Effect Einstein proposed that “light”: • is quantized. • behaves like a stream of massless particles. • G. N. Lewis later named them photons. • Imagine photons (balls) hitting e- embedded in glue. • If the E of the ball: • is low, it can’t eject an e-. • exceeds the strength of the glue, an e- is released Higher intensity = more photons (balls). If E > threshold, more balls eject more e-.

17. 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

18. E = h •  Relationships:

19. Short wavelength light has: • High frequency and low energy • High frequency and high energy • Low frequency and low energy • Low frequency and high energy

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

21. 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?

22. What is the energy of a photon of 525 nm light? • 3.79 x 10-19 J • 4.83 x 10-22 J • 3.67 x 1020 J • 8.43 x 1023 J

23. Radio Wave Energy • What is the energy of a photon corresponding to radio waves of frequency 1.255 x 10 6 s-1?

24. 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?

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

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

27. Light emitted by solids

28. Light emitted by hydrogen gas

29. 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

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

31. Line spectra of atoms

32. Energy Adsorption/Emission

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

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

35. Explanation of line spectra Balmer series

36. Bohr Model of the Hydrogen Atom • Heated solid objects emit continuous spectra. • Excited atomic gases emit line spectra. • Each element has a unique pattern.

37. E = −RH n = 1, 2, 3. . .  1 n2 Bohr Model of the Hydrogen Atom Niels Bohr • Orbits the nucleus. • Different orbits are possible with different quantizedE values: Bohr (1913). The hydrogen e-: Rydberg constant 2.179 x 10-18 J If the e- has n = 1 (lowest, most negative E), the atom is in its ground state. If ionized (e- removed), n =  (E = 0).

38. infrared emission n ∞ 3 2 1 0 -¹⁄9RH -¼RH visible emission Energy -RH • 500 600 700 • wavelength (nm) ultraviolet emission ultraviolet absorption Bohr Model of the Hydrogen Atom absorption: ΔE > 0, n ↑ emission: ΔE < 0, n ↓ Bohr’s model exactly predicts the H-atom spectrum.

39. H-atom transitions: ΔE = −RH– 1 nf2 1 ni2 Bohr Model of the Hydrogen Atom Example Calculate the energy and wavelength (in nm) for an H-atom n = 4 → n = 2 transition.

40. Bohr Model of the Hydrogen Atom Calculate E and wavelength (nm) for an H-atom n = 4 → n = 2 transition.