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

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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 10 8 m/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 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 X10 15 Hz. 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 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 & planets7.2

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 11s 22s, p 33s, p, d 44s, p, d, f

19 Quantum Numbers 3. Magnetic Quantum Number-m l describes orientation of orbital in space can be any number from –3 to 3 or zero 4. Electron spin quantum number- m s 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 =n 2 Sublevel ( l ) Possible values of m l # Orbitals# Electrons s012 p1,0,-136 d2,1,0,-1,-2510 f3,2,1,0,-1,-2,-3714

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 (m l ) Spin (m s )

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- s 1 Alkaline Earth Metals- s 2 Transition Elements- outer s & inner d Inner Transition Elements- outer s & inner f

23 Electron Sequence by the Periodic Table 1s1s La Ac 1s1s 5f5f 4f4f 2s2s 3s3s 4s4s 5s5s 6s6s 7s7s 2p2p 3p3p 4p4p 5p5p 6p6p 3d3d 4d4d 5d5d 6d6d

24 The Periodic Table 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 d10d10 p1p1 p2p2 p3p3 p4p4 p5p5 p6p6 f1f1 f2f2 f3f3 f4f4 f5f5 f6f6 f7f7 f8f8 f9f9 f 10 f 11 f 12 f 13 f

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

26 Basic Rules to remember: Aufbau Principle- e- enter the lowest energy level 1 st 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: s 2 d 4  s 1 d 5 s 2 d 9  s 1 d 10

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

30 Electron Configuration F Cl Al Br

31 Electron Configuration 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

32 Orbital Diagrams   1s2s2p3s3p4s3d4p   1s2s2p3s3p4s3d4p   1s2s2p3s3p4s3d4p   1s2s2p3s3p4s3d4p

33 Electron Configurations Sc K P B 1s2s2p3s3p4s3d4p1s2s2p3s3p4s3d4p1s2s2p3s3p4s3d4p1s2s2p3s3p4s3d4p

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: –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

35 8.1 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

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