waves amplitude frequency light crest ROYGBIV electrons EMS energy Photoelectric effect

 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 10 8 m/s  c=

 EM spectrum › Radio Microwave Infrared Visible UV X- ray Gamma  Low E High E High Low Low High › Visible Light: Continuous spectrum  ROYGBIV  Low E High E High Low Low High

 Violet : nm  Blue: nm  Green: nm  Yellow: nm  Orange: nm  Red: nm

 What is the frequency of green light, which has a wavelength of 4.90 x m?  An X-ray has a wavelength of 1.15 x m. What is its frequency?  A popular radio station broadcasts with a frequency of 9.47 x 10 7 s -1. What is the wavelength of the broadcast?  Hz = waves/sc= 3.00 x 10 8 m/s

 Particle Nature of light › Quantum- minimum amount of E gained or lost by an atom › Quantitized E- E gained in packet (NOT continuous)  E=h E= energy in J h= Planck’s Constant x Js = frequency (s -1 )

 What is the energy of each of the following types of radiation? › 6.32 x s -1 › 9.50 x Hz › 1.05 x s -1  What types of radiation are the above?

 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

 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).

 Atomic Emission Spectra › e - excited will jump to another E level, › As they fall they emit E (light) › of the waves allow for a unique color (NOT a continuous color spectrum like a white light) › Atomic EmissionAtomic Emission

 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

DeBroglie- all moving particles have wave-like characteristics

 Heisenburg Uncertainty Principle › Can’t know velocity & position of e - at same time

 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 -

 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 LevelSublevelsSublevel Called 11s 22s, p 33s, p, d 44s, p, d, f

** Each orbital can only hold 2 e - Sublevel# Orbitals# e - held s12 p36 d510 f714

 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

 1s 2  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

1s 2s 3s 4s 5s 6s 7p 6p 5p 4p 3p 2p 6d 5d 4d 3d 4f 5f 7s Follow the yellow brick road

1s1s La Ac 1s1s 5f5f 4f4f 2s2s 3s3s 4s4s 5s5s 6s6s 7s7s 2p2p 3p3p 4p4p 5p5p 6p6p 3d3d 4d4d 5d5d 6d6d

1H1H 3 Li 11 Na 19 K 37 Rb 55 Cs 87 Fr 4 Be 12 Mg 20 Ca 38 Sr 56 Ba 88 Ra 21 Sc 39 Y 57 La 89 Ac 22 Ti 40 Zr 72 Hf 104 Rf 23 V 41 Nb 73 Ta 105 Db 42 Mo 74 W 106 Sg 25 Mn 43 Tc 75 Re 107 Bh 26 Fe 44 Ru 76 Os 108 Hs 27 Co 45 Rh 77 Ir 109 Mt 28 Ni 46 Pd 78 Pt 110 Uun 111 Uuu 30 Zn 48 Cd 80 Hg 8O8O 16 S 34 Se 52 Te 84 Po 7N7N 15 P 33 As 51 Sb 83 Bi 6C6C 14 Si 32 Ge 50 Sn 82 Pb 5B5B 13 Al 31 Ga 49 In 81 Tl 9F9F 17 Cl 35 Br 53 I 85 At 2 He 10 Ne 18 Ar 36 Kr 54 Xe 86 Rn 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 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 24 Cr 29 Cu 47 Ag 79 Au 112 Uub 114 Uuq 116 Uuh 118 Uuo s d p f s1s1 s2s2 d1d1 d2d2 d3d3 d4d4 d5d5 d6d6 d7d7 d8d8 d9d9 d 10 p1p1 p2p2 p3p3 p4p4 p5p5 p6p6 f1f1 f2f2 f3f3 f4f4 f5f5 f6f6 f7f7 f8f8 f9f9 f 10 f 11 f 12 f 13 f

 Electron Orbital Diagram: visually shows e - placement around the nucleus › Each orbital gets own box Orbital# Orbitals# electrons held# boxes s121 p363 d5105 f7147

 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     

 F – 1s 2 2s 2 2p 5  Cl – 1s 2 2s 2 2p 6 3s 2 3p 5  Al – 1s 2 2s 2 2p 6 3s 2 3p 1  Br - 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 5   1s2s2p3s3p4s3d4p   1s2s2p3s3p4s3d4p   1s2s2p3s3p4s3d4p   1s2s2p3s3p4s3d4p

 Sc  K  P  B 1s2s2p3s3p4s3d4p 1s2s2p3s3p4s3d4p 1s2s2p3s3p4s3d4p 1s2s2p3s3p4s3d4p

 Other helpful hints- › #e= # p = atomic number if neutral atom › Add superscripts to get the #e, #p

 Noble Gas Configuration › Go back to the last noble gas › Write symbol for noble gas in brackets › Write rest of configuration  Na Complete Configuration:  1s 2 2s 2 2p 6 3s 1  Na Noble gas Configuration:  [Ne] 3s 1  Exceptions to electron configuration: › e - want to be stable › Stable is a full or ½ full e- shell › Cr- [Ar] 4s 2 3d 4  [Ar] 4s 1 3d 5 › Cu- [Ar] 4s 2 3d 9  [Ar] 4s 1 3d 10

 Valence electrons- e - in outer most level › Put in noble gas configuration › Count e - in highest level  Ex: Na 1s 2 2s 2 2p 6 3s 1 has 1 valence e -  Cs [Xe] 6s 1 has 1 valence e -  Cu [Ar] 4s 1 3d 10 has 1 valence e -  S [Ne] 3s 2 3p 4 has 6 valence e -  Lewis Dot Structures- shows valence e - around symbol Li N Be O B F C Ne

 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