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Honors Chemistry Chapter 4 Notes

“Electron Configurations”. Honors Chemistry Chapter 4 Notes. Introduction…pg. 124. Using a different perspective gives us a different view…. Currently…. Coming Up…. Introduction…pg. 125. Read Aloud…. I. 4.1 Radiant Energy. A. What are the 4 characteristics of an electromagnetic wave?

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Honors Chemistry Chapter 4 Notes

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  1. “Electron Configurations” Honors Chemistry Chapter 4 Notes

  2. Introduction…pg. 124 Using a different perspective gives us a different view… Currently… Coming Up…

  3. Introduction…pg. 125 Read Aloud…

  4. I. 4.1 Radiant Energy A. What are the 4 characteristics of an electromagnetic wave? 1) 2) 3) 4) *pg. 126 B. Common Types X-rays, radio waves, gamma, visible, ultraviolet…

  5. C. Characteristics 1. Speed of light a. 3.00 x 108 m/s 2. Formula a. λ = c / ν λ = wavelength ν = frequency c = speed of light 3. Example: A radio wave has a frequency of 93.1 x 106 s-1. Determine the wavelength of this wave? λ = 3.22 m

  6. II. Electromagnetic Spectrum A. Based on wavelengths. B. Many components

  7. Assign: Do #’s 2 and 3 in your notes 1. Read 4.2 and 4.3 2. Calculate the wavelength of a radio wave that has a frequency of 96.5 e6 sec-1. 3. Would a wave that has a wavelength of 2.89 e-7 cm and a frequency of 1.04 e11 MHz exist? Explain your reasoning.

  8. III. Quantum Theory (4.2) Read pg. 130 Questions were raised…time for a new perspective of the atomic structure. A. Planck’s Theory (pg. 131) 1. quantum – fixed amount of energy that can be absorbed or emitted by a particle. 2. Car example (pg. 131)

  9. B. Photoelectric Effect (Einstein) 1. Photons – tiny particles with quanta that we see. 2. Quanta of energy caused electrons to jump and move…

  10. IV. Dual Nature of Light Summary: Light (photons) are dualistic in nature. They have a wave-like and particle-like property. This changes our perspective again. Where are electrons and why are photons produced sometimes and not others…??? Investigate with Flame Test Lab tomorrow…

  11. Honors Chemistry 4.3 “Another Look at the Atom”

  12. Results of Flame Test… Unknown…????

  13. I. Line Spectra A. Def – Color or wavelength that is seen. 1. Depends on material 2. Based on electrons movement

  14. II. Bohr Model A. Who was he??? 1. Danish physicist 2. Student of Rutherford B. Proposed – electron is found in specific paths or orbits 1. energy levels (n=1, n=2, n=3…) 2. electrons can move b/t each other (similar to rungs on ladder) a. based on energy absorbed or lost C. Only explained hydrogen atom… D. Uncertainty Principle – we don’t know where electrons are…we only know an area.

  15. Honors Chemistry 4.4 “A New Approach…”

  16. I. Quantum Mechanical Model A. What is it??? (pg. 141) 1. Based on probability (mathematically) a. Propeller blade 2. Restricts electrons to certain values or areas 3. No line paths

  17. II. Atomic Orbitals A. Def – a region of space where there is a high probability of finding an electron. B. Energy levels (n) 1. 1, 2, 3, 4…and so on 2. each energy level can have different #’s of orbitals with different types of shapes… C. The 4 Orbital Shapes…s, p, d, f 1. s = spherical 2. p = dumbbell 3. d = cloverleaf (4 out of 5) 4. f = complex *Every orbital can contain at most 2 electrons*

  18. D. When n = 1… 1. has one sublevel  1S 2. s always has 1 orbital E. When n = 2… 1. has two sublevels 2s and 2p 2. 4 orbitals total a. 2S = 1 b. 2P = 3 (Px, Py, Pz) (pg.144) -all perpendicular to each other F. When n = 3… 1. has three sublevels 3s, 3p and 3d 2. 3S = 1 3P = 3 3D = 5 orbitals 3. 9 orbitals total

  19. G. When n = 4… 1. 4S = 1 4P = 3 4D = 5 4F = 7 2. 16 orbitals total

  20. S P D F

  21. Honors Chemistry 4.5 “Electron Configurations and Orbital Notations”

  22. I. Electron Configurations • Aufbau Principle (pg. 148) 1. electrons enter the lowest energy levels first. 2. orbitals in a sublevel are of equal value. -ex: 2p = 3 orbitals (all have equal values) 3. energy levels overlap each other -ex: 3d requires more energy than 4s -Is a 4f orbital higher or lower than 5d? lower

  23. Pauli exclusion principle (pg. 148-149) 1. an atomic orbital may describe at most 2 electrons -orbitals hold 2 electrons 2. electrons must have opposite spins -clockwise or counterclockwise ↑ ↓ 3. Paired electrons ↑ 1s1 ↑↓ 1s2 hydrogen helium Electron configuration

  24. Hund’s Rule (pg. 149) 1. When electrons occupy orbitals of equal energy, one electron enters each orbital until all the orbitals contain one electron with parallel spins 2p ↑ ↑ ↑ ↓ Orbital notation

  25. 1s1 ↑ 1s2 ↑ ↓

  26. II. Abbreviated Configuration A. Rules 1. Write the most recent noble gas symbol 2. Place in brackets [Ar] 3. Complete electron configuration from the noble gas to the desired element. B. Ex: Silicon 1. Normal = 1s2 2s2 2p6 3s2 3p2 2. Abbreviated = [Ne] 3s2 3p2

  27. Ex: 1. What is the abbreviated electron configuration for the following elements? a. Lithium b. Chlorine c. Iodine d. Lead

  28. III. Exceptions in E.C. A. Chromium and Copper Groups 1. Normal Configuration -Cr 1s22s22p63s23p64s23d4 2. “½ Rule” -An element would rather have some orbitals half full than one partially empty

  29. normal 1s22s22p63s23p64s23d4 so… 1s22s22p63s23p64s13d5 Take 1 electron from 4s2 and carry it to the 3d orbital Both 4s and 3d are ½ full

  30. 3. Write the electron configuration for copper. 1s22s22p63s23p64s13d10 One orbital is ½ full and the other is full

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