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Unit 2B Notes: Electron Configuration Ch 6

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1 Unit 2B Notes: Electron Configuration Ch 6
Dunton Honors Chemistry

2 Newton- light consists of wave & particle properties
Waves- light is made of electromagnetic radiation      Wavelength- distance from trough to trough,  in m      Amplitude- distance from top to base line      Frequency- # waves past a point per unit time, v ·          Measured in Hz (hertz) = s-1 = 1/s        

3 Wave Anatomy

4 Speed of light- how fast all light goes, c
·  c= 3.00 X 108m/s ·  c = v Remember v is frequency in Hz or 1/s and  is wavelength in m (you may have to convert!) ** Notice that wavelength & frequency are inversely proportional Find the wavelength if the frequency is 5.10 X 1014 Hz.

5 Know this in order!! Electromagnetic spectrum:
Radio Micro Infrared Visible UV X-ray Gamma roygbiv High wavelength () Low  Low frequency (v) High v Low E High E

6 Electromagnetic Spectrum

7 6.2 Planck- energy is emitted and absorbed in packets (quanta)
Explains line spectrum E=hv h= Planck’s constant X Js v = frequency, Hz, 1/s, s-1 E proportional to v(frequency) Find the energy if the frequency of 5.00 X1015Hz. Find the wavelength of radiation if the energy of the wave is 2.66 X J. Use the chart on pg 276 to find the type of radiation.

8 Einstein- light in quanta= photons
E of photons is quantitized Photoelectric effect- metals eject e- called photoelectrons when light of a high E shines on them          Purple light has high v to have high E & can eject e-          Red light has low v so low E & can’t eject e-       If you want to increase E that will increase the speed of e-       If increase intensity, just increases # of e- not speed Ex. Solar calculators

9 7.2 Can find color or type of emission based on  & v
Emission spectrum- colored lines characteristic of elements Ø      Atoms absorb E then lose E give off light Ø      Can use this to ID elements Continuous spectrum- all colors (white light) Line Spectrum- specific colors, based on element Used to ID contents of stars & planets

10 7.4 Bohr’s Model: electrons orbit the nucleus; only orbits in certain energies are permitted Ground State- lowest E level Excited State- Higher than ground state The e- are raised to the next level, then release light when they return to ground state Must have enough E to raise to next level or won’t happen

11 6.4 DeBroglie- Quantum Mechanics- light behaves as wave & particles  = h/mv = wavelength h= Planck’s constant m=mass v= velocity Visible objects (baseball) have  too small to see, need very small object to detect  Heisenburg Uncertainty Principle- Can’t know the position & speed of electron at the same time

12 Energy levels- region around nucleus where e- likely to be found (electron density is high) 90%
Quantum- amount of E for e- to jump levels Continuous- ramp, no units Quantitized- fixed levels, fixed units

13 Flame Test Lab Sc3f Purpose: In this lab the student will observe light spectrum emitted from heated metal ions and calculate the frequency, wavelength and energy. Materials: salts, bunsen burner, Q-tips Safety: Wear goggles & apron. Wash hands & wipe down lab station after use.

14 Procedure: 1. Get two beakers of water. Label one “clean” and one “dirty”. 2. Dip the Q-tip into the “clean” water 3. Dip the Q-tip into the salt 4. Put the Q-tip into the flame 5. Record the color 6. Swish the Q-tip into the “dirty” water and leave it there. 7. Repeat with clean Q-tip for each sample. Observations: Create your own table with salt and color. There are 7 samples.

15 Conclusion: 1. Draw and label the parts of the atom.
2. Explain what happened on the atomic level in this lab. 3. How does the flame test provide support for quantitized energy levels? Explain. 4. List the metal ions present in your unknown solutions and provide reasons for your choices. 5. For each salt use the color to estimate the wavelength from the chart in your book. Use the wavelength to solve for the frequency and energy.

16 Schrodinger- estimates the probability of e- to be in certain area; uses wave and particle like properties to create quantum mechanics (a series of wave functions and mathematical equations)         A fuzzy cloud, more dense= more likely to find e- 90% of the time Orbitals-Wave functions with corresponding densities (shape and energy) **orbital is NOT the same as Bohr’s orbit

17 Quantum Numbers: 1. Principle Quantum Number- (n) 1-7
     Same as period, tells number of levels; the higher the number the further from the nucleus 2. Second Quantum Number aka the azimuthal quantum number -(l) determines shape Sublevels- same as # of n, shape of orbital          s- sphere l =0         p- dumbbell l=1        d- complex l=2       f- complex l=3    

18 Correlation of n & l Level (n) Sublevels (l) Sublevel Called 1 s 2
s, p 3 s, p, d 4 s, p, d, f

19 Quantum Numbers 3. Magnetic Quantum Number-ml describes orientation of orbital in space can be any number from –3 to 3 or zero 4. Electron spin quantum number- ms describe the direction of the electron spin producing a magnetic field; 2 opposite magnetic fields allow for close spacing; +1/2 or – 1/2

20 electron shells - Collection of orbitals with same n value
Subshell- set of orbitals with same n & l Orbital- Each orbital can hold 2 e- Formula for total number of orbitals =n2 Sublevel (l) Possible values of ml # Orbitals # Electrons s 1 2 p 1,0,-1 3 6 d 2,1,0,-1,-2 5 10 f 3,2,1,0,-1,-2,-3 7 14

21 6.7 Pauli Exclusion Principle- Because e- are negatively charged, and they repel each other, w/in an orbital one goes clockwise & one spins counter clockwise; no two e- can have the exact same set of quantum numbers Each electron has an “address”: Principal quantum number (n) Sublevel (l) Orbital (ml) Spin (ms)

22 Electron Configuration- way e- arranged around atom, each e- address
Using the PT: The principal quantum number for the outer electrons is the same as the period. There are blocks: s, p, d, f Noble gases have a full s & p level making them inert Alkali Metals- s1 Alkaline Earth Metals- s2 Transition Elements- outer s & inner d Inner Transition Elements- outer s & inner f

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

24 The Periodic Table 4 5 s1 s2 d1 d2 d3 d4 d5 d6 d7 d8 d9 d10 p1 p2 p3
Li 11 Na 19 K 37 Rb 55 Cs 87 Fr 4 Be 12 Mg 20 Ca 38 Sr 56 Ba 88 Ra 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 89 Ac 104 Rf 105 Db 106 Sg 107 Bh 108 Hs 109 Mt 110 Uun 111 Uuu 112 Uub 114 Uuq 116 Uuh 118 Uuo s f1 f2 f3 f4 f5 f6 f7 f8 f9 f10 f11 f12 f13 f14 d 4 5 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 p f

25 Electron Orbital Diagram: visually shows e- placement around the nucleus
Each orbital gets own box Orbital # Orbitals # electrons held # boxes s 1 2 p 3 6 d 5 10 f 7 14

26 Basic Rules to remember:
Aufbau Principle- e- enter the lowest energy level 1st Pauli Exclusion Principle- an orbital can only describe 2e-, Show each orbital w/its own box One is spinning clockwise & the other is counter clockwise, Show this with one arrow going up & one pointing down NOT   

27 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   

28 Aufbau Exceptions: Cr, Cu, Mo, Au, Ag
Want full or ½ full orbital and can shift e- to get it Ex: s2d4 s1d5 s2d9 s1d10

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

30 Electron Configuration
Cl Al Br

31 Electron Configuration
F 1s22s22p5 Cl 1s22s22p63s23p5 Al 1s22s22p63s23p1 Br 1s22s22p63s23p64s23d104p5

32 Orbital Diagrams   1s 2s 2p 3s 3p 4s 3d 4p   1s 2s 2p 3s 3p 4s

33 Electron Configurations
Sc K P B 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

34 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

35 8.1 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

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