1. Chapter 5: Electrons in Atoms 5.1: Models of the Atom

2. I. The Bohr Model A. Bohr model: “planetary model” of nucleus and electrons (e - ) B. E - in specific orbits around nucleus

3. II. Quanta A. E - orbits represent certain amount of energy for each electron B. E - can move energy levels if they are given or give off a specific amount of energy (Quantum)

4. III. Atomic Orbits A. Each orbit has a “Principle Quantum Number (n)”, n = 1, 2, 3, etc. B. Each Principle Level has sublevels labeled by letters (s, p, d, f) C. Each Principle Level has one more sublevel than the one before it

5. IV. E - Energy Energy of sublevels not clearly related to Principle Levels

6. V. Principle and Sublevel Order Short-Cut 7 s 7p 7d 7f 6s 6p 6d 6f 5s 5p 5d 5f 4s 4p 4d 4f 3s 3p 3d 2s 2p 1s Can also use Periodic Table to help

7. VI. Quantum Mechanical Model A. Bohr’s model only works for atoms having one e - B. Mathematical estimation of e - location (“Quantum Mechanical Model”), determines actual shape of e - orbits C. Orbital shape based on e - density diagram D. Shape based on 90% of e - positions

8. VII. Orbital Sublevel Shapes A. S-orbital: 1 spherical, holds 2e - B. P-Orbital: 3 figure- eight shapes, holds 6e -

9. C. D-orbital: 5 clover-leaf shapes, holds 10e - D. F-orbital: 7 orbitals, hold 14e -

10. VIII. Principle Orbitals A. Level 1: 1s (2e - ) = up to 2e - B. Level 2: 2s (2e - ) + 2p (6e - ) = up to 8e - C. Level 3: 3s (2e - ) + 3p (6e - ) + 3d (10e - ) = up to 18e - D. Level 4: 4s (2e - ) + 4p (6e - ) + 4d (10e - ) + 4f (14e - ) = up to 32e -

11. 5.2 Electron Arrangement

12. I. Electron Configuration Rules A. Aufbau principle: fill lowest energy level orbitals first, each orbital shown by box  B. Pauli exclusion principle: each orbital can only hold two e - s, each move in opposite direction shown by arrows C. Hund’s rule: electrons remain unpaired until each type of orbital is filled (ex. 4 electrons in P orbitals)

13. Boron Example

14. II. Configuration Short-cuts A. Use superscript numbers to indicate electrons in each sublevel B. Ex. Carbon is 1s 2 2s 2 2p 2 C. Use Noble gas immediately preceding element in place of electron configuration leading up to it D. Ex. Instead of Sulfur as [1s 2 2s 2 2p 6 ] 3s 2 3p 4 you can replace bracketed section as Neon E. Ex. [Ne] 3s 2 3p 4

15. 5.3 Quantum Mechanical Model

16. I. Light and Waves A. Electromagnetic radiation: waves of light

17. B. Crest: top part of a wave C. Trough: bottom part D. Amplitude: wave height from origin to crest E. Wavelength ( ): distance between crests Wavelength Crest Amplitude Trough

18. II. Wave Properties A. Frequency (  ): number of waves passing a point per second (Hertz = 1/second) B. Wave speed = wavelength ( ) x frequency (  ) C. Speed of any type of light is the same under the same conditions (3x10 8 m/s), called “C”

19. D. Long wavelength = less energy E. Short wavelength = more energy Demo: Color Diffraction F. Atomic Emission spectrum: atomic light through a prism gives off specific spectrum ArgonHelium Hydrogen Neon

20. III. Duality of Light C. Light acts as wave and particle that has discrete energies (only specific wavelengths) D. Light particles called “Photons” A. Experiment to See Nature of Light B. E - hit screen one at a time, but still interfere like waves

21. What the wave/particle nature of light might look like.

22. IV. Why Produce Light? A. If e - given energy, move to higher level B. When e - returns, emits excess energy as light

23. V. How Much is a Quantum? A. Energy = Planck’s constant (h) x frequency (  ) B. Planck’s constant (h) = 6.63x10 -34 Joules x Sec C. Energy measured in Joules (a.k.a. Kg m 2 /s 2 ) In honor of Niels Bohrs B-day.

24. VI. Hydrogen Spectrum A. “H” spectrum due to possible energies e - can emit B. For “H”: E = 2.178x10 -18 J(z 2 /n 2 ) E = 2.178x10 -18 J(z 2 /n final 2 - z 2 /n initial 2 ) C. “z” is nuclear charge, “n” is principle level D. Can also change Energy of photon equation to get wavelength E = hC/λ