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Photoelectric effect. EMS. crest. light. waves. amplitude. energy level. electrons. frequency. wavelength. origin. ROYGBIV. energy. http://youtu.be/xiMApWz5wsI. Electron.1. Wave nature of light Electromagnetic (EM) radiation- E emits wave like behavior  All waves have:

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  1. Photoelectric effect EMS crest light waves amplitude energy level • electrons frequency wavelength origin ROYGBIV energy

  2. http://youtu.be/xiMApWz5wsI

  3. Electron.1 • Wave nature of light • Electromagnetic (EM) radiation- E emits wave like behavior  • All waves have: • Wavelength- () distance from crest to crest • Frequency- () # waves past a given point per second (s-1) • Amplitude- height of wave from crest to origin • All light travels at speed of light, c = 3.00 x 108 m/s                    • c= 

  4. Electron -1 • EM spectrum • Radio  Microwave  Infrared Visible  UV  X-ray Gamma • Low EHigh E High            Low  Low           High  • Visible Light: Continuous spectrum • ROYGBIV • Low EHigh E High       Low  Low        High 

  5. Electromagnetic Spectrum

  6. ROYGBIV • Violet: 400- 420 nm • Blue: 420- 490 nm • Green: 490- 580 nm • Yellow: 580- 590 nm • Orange: 590- 650 nm • Red: 650-700 nm

  7. Electron -1 • What is the frequency of green light, which has a wavelength of 4.90 x 10-7 m? • An X-ray has a wavelength of 1.15 x 10-10 m. What is its frequency? • A popular radio station broadcasts with a frequency of 9.47 x 107 s-1. What is the wavelength of the broadcast? • Hz = waves/s c= 3.00 x 108 m/s

  8. Electron -1 • Particle Nature of light • Quantum- minimum amount of E gained or lost by an atom  • Quantitized E- E gained in packet (NOT continuous) • E=hn E= energy in J h= Planck’s Constant 6.6262 x 10-34Js = frequency (s-1) 

  9. Electron -1 • What is the energy of each of the following types of radiation? • 6.32 x 1020 s-1 • 9.50 x 1013 Hz • 1.05 x 1016 s-1 • What types of radiation are the above?

  10. Electron -1 • Photoelectric effect- • Photoelectrons are emitted from a metals surface when light of certain frequency shines on it • Ex. solar calculators, automatic doors • Each metal has its threshold for the photoelectric effect • If light is shined on metal that doesn’t have the correct frequency, no matter how long, e- will not be emitted 

  11. Lab Conclusion • Due tomorrow! Must be typed. • Explain what you did in the lab. • Explain what happened. • Explain the science behind what happened (photoelectric effect). • How you calculated wavelength, frequency, and energy & identified the unknown salts. • Reflection…what you liked, what you didn’t like (if any).

  12. Electron -1 • Atomic Emission Spectra • e- excited will jump to another E level, • As they fall they emit E (light) • n of the waves allow for a unique color (NOT a continuous color spectrum like a white light) • Atomic Emission

  13. Electron -2 • Bohr- e- only have “allowable E states” • Normally in ground state • e- around nucleus in orbits • Lower the E, closer the orbit to the nucleus • Quantum number (n)- lowest E state

  14. DeBroglie- all moving particles have • wave-like characteristics

  15. Electron -2 • Heisenburg Uncertainty Principle • Can’t know velocity & position of e- at same time 

  16. Schrodinger-quantum mechanical model of atom using wave properties of e- predicts e- will be found in orbitals, increase D of cloud = higher probability of e-

  17. Electron -2 • Principle quantum number (n)- tell relative sizes & shapes of orbitals • Higher n = bigger orbital = increased time of e- away from nucleus • Levels contain same number of sublevels as level number

  18. Electron -2 ** Each orbital can only hold 2 e-

  19. Electron -3 • Aufbau principle- each e- occupies lowest E orbital available • Electron Configuration- arrangement of e- in orbitals around atom • Lower E more stable than high E • Lowest E= ground state

  20. Example: • 1s2 • 1 is n • s is sublevel • 2 is number of e- in sublevel & is a superscript • All orbitals in each E sublevel are equal (the three orbitals are =E) • Fill s, p, d, f, for each level • Orbitals can overlap

  21. 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 7s 7p Electron Sequence Model Follow the yellow brick road

  22. 1s 1s 2s 2p 3s 3p 3d 4s 4p 4d 5s 5p 5d La 6s 6p 6d Ac 7s 4f 5f Electron Sequence by the Periodic Table

  23. 1 H 3 Li 4 Be 11 Na 12 Mg 19 K 20 Ca 37 Rb 38 Sr 55 Cs 56 Ba 87 Fr 88 Ra 58 Ce 59 Pr 60 Nd 61 Pm 62 Sm 63 Eu 64 Gd 65 Tb 66 Dy 67 Ho 68 Er 69 Tm 70 Yb 71 Lu 90 Th 91 Pa 92 U 93 Np 94 Pu 95 Am 96 Cm 97 Bk 98 Cf 99 Es 100 Fm 101 Md 102 No 103 Lr The Periodic Table 2 He 5 B 6 C 7 N 8 O 9 F 10 Ne 13 Al 14 Si 15 P 16 S 17 Cl 18 Ar 21 Sc 22 Ti 23 V 24 Cr 25 Mn 26 Fe 27 Co 28 Ni 29 Cu 30 Zn 31 Ga 32 Ge 33 As 34 Se 35 Br 36 Kr 39 Y 40 Zr 41 Nb 42 Mo 43 Tc 44 Ru 45 Rh 46 Pd 47 Ag 48 Cd 49 In 50 Sn 51 Sb 52 Te 53 I 54 Xe 57 La 72 Hf 73 Ta 74 W 75 Re 76 Os 77 Ir 78 Pt 79 Au 80 Hg 81 Tl 82 Pb 83 Bi 84 Po 85 At 86 Rn 118 Uuo 116 Uuh 89 Ac 104 Rf 105 Db 106 Sg 107 Bh 108 Hs 109 Mt 110 Uun 111 Uuu 112 Uub 114 Uuq s d p f

  24. Electron -3 • Electron Orbital Diagram: visually shows e- placement around the nucleus • Each orbital gets own box

  25.      Electron -3 • Pauli Exclusion Principle- only 2 e- can occupy an orbital.  Each w/ opposite spins show w/ arrow up and down NOT • Hund’s Rule- e- w/ same spin must occupy each E level in a sublevel before doubling up • Example: when filling the p sublevel with 4e-, each box gets 1 before doubling up one box NOT   

  26.      1s 2s 2p 3s 3p 4s 3d 4p          1s 2s 2p 3s 3p 4s 3d 4p        1s 2s 2p 3s 3p 4s 3d 4p                   1s 2s 2p 3s 3p 4s 3d 4p Electron Configurations • F – 1s22s22p5 • Cl – 1s22s22p63s23p5 • Al – 1s22s22p63s23p1 • Br - 1s22s22p63s23p64s23d104p5

  27. 1s 1s 1s 1s 2s 2s 2s 2s 2p 2p 2p 2p 3s 3s 3s 3s 3p 3p 3p 3p 4s 4s 4s 4s 3d 3d 3d 3d 4p 4p 4p 4p Electron Configurations • Sc • K • P • B

  28. Electron -3 • Other helpful hints- • #e= # p = atomic number if neutral atom • Add superscripts to get the #e, #p

  29. Electron -3 • Noble Gas Configuration • Go back to the last noble gas • Write symbol for noble gas in brackets • Write rest of configuration • Na Complete Configuration: • 1s22s22p63s1 • Na Noble gas Configuration: • [Ne] 3s1 • Exceptions to electron configuration: • e- want to be stable • Stable is a full or ½ full e- shell • Cr- [Ar] 4s23d4 [Ar] 4s13d5 • Cu- [Ar] 4s23d9[Ar] 4s13d10

  30. Electron -3 • Valence electrons- e- in outer most level • Put in noble gas configuration • Count e- in highest level • Ex: Na 1s22s22p63s1  has 1 valence e- • Cs [Xe] 6s1 has 1 valence e- • Cu [Ar] 4s13d10 has 1 valence e- • S [Ne] 3s23p4  has 6 valence e- • Lewis Dot Structures- shows valence e- around symbol  Li    N  Be   O  B    F  C    Ne

  31. Properties of the d and f-Block Elements • Magnetism – ability to be affected by magnet • Diamagnetism – all e- are paired, substance is unaffected or slightly repelled by magnetic field • Paramagnetic – unpaired electron in the valence orbital is attracted to magnetic field • Ferromagnetism – strong attraction of substance, ions can align in direction of field and form permanent magnet

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