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The Downfall of Classical Physics

The Downfall of Classical Physics. Niels Bohr – Quantum Mechanics. Niels Bohr applies Quantum Mechanics to Rutherford’s model and proposes that electrons are located around the nucleus in energy levels. . The Rise of Quantum Physics.

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The Downfall of Classical Physics

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  1. The Downfall of Classical Physics

  2. Niels Bohr – Quantum Mechanics Niels Bohr applies Quantum Mechanics to Rutherford’s model and proposes that electrons are located around the nucleus in energy levels.

  3. The Rise of Quantum Physics Spectral lines can be explained by the movement of electrons from one energy level to another.

  4. ENERGY LEVEL (n) • “Rings” of Bohr’s planetary model • 7 possible energy levels • Maximum number of electrons in each energy level = 2n2 • 1st level: • 2nd level: • 3rd level: 2 8 18

  5. ENERGY SUBLEVELS • 4 types of sublevels: • s • p • d • f • Every energy level begets a new sublevel • Energy level 1: 1 sublevel (s) • Energy level 2: 2 sublevels (s, p)

  6. ORBITALS • s sublevel = 1 s orbital • p sublevel = 3 p orbitals • d sublevel = 5 d orbitals • f sublevel = 7 f orbitals

  7. Summary (energy levels, sublevels)

  8. Summary (sublevels, orbitals, electrons)

  9. MAGNETIC SPIN • “Spin up” • “Spin down”

  10. Summary (energy levels, sublevels)

  11. Summary (sublevels, orbitals, electrons)

  12. RULES FOR ELECTRONS • An orbital can hold a maximum of TWO electrons • Pauli exclusion principle • Aufbauprinciple • Hund’s rule

  13. Pauli Exclusion Principle • no two electrons can have the same “address” • Energy level • Sublevel • Orbital • Spin

  14. AufbauPrinciple • Electrons will fill the orbitals with the lowest amount of energybefore filling in orbitals that have more energy • Order is: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p

  15. Hund’s Rule • Electrons will not share an orbital of the same energy if there is an empty orbital with that energy is available

  16. Orbital Diagrams Orbital Spin 1 s Energy level Sublevel

  17. Draw an orbital diagram for carbon: 1s 2s 2p

  18. How to draw orbital diagrams • Use periodic table to figure out where to end (energy level and sublevel of last electron) • Draw boxes (orbitals) for all energy levels and sublevels up to the ending point • 1 box for s, 3 boxes for p, 5 boxes for d, 7 boxes for f • Fill boxes with electrons (2 per orbital) with opposite spins • Follow Hund’s rule when in the p, d, or f sublevels

  19. Electron configurations • Based on orbital diagrams • Use numbers and superscripts along with s, p, d and f to show location of electrons in atoms • Coefficient = energy level • Letter = sublevel • Superscript = # of electrons in sublevel

  20. Example: Energy level 2 electrons 1s2 sublevel

  21. Write an electron configuration for nitrogen: 1s 2s 2p 1s22s22p3

  22. Practice • Draw the orbital diagram for oxygen • Write the electron configuration for oxygen 1s 2s 2p 1s22s22p4

  23. Shortcuts!! • Shortcut (shortened) Notation: • Use closest Noble Gas before the element • For bromine: • Closest Noble Gas before the element: • Argon: • Shortcut Notation: 1s22s22p63s23p6 4s23d104p5 1s22s22p63s23p6 [Ar] 4s23d104p5b

  24. What is the electron configuration of potassium? • 1s22s22p63s23p63d1 • 1s22s22p23s23p24s1 • 1s22s22p63s23p3 • 1s22s22p63s23p64s1

  25. Each period number in the periodic table corresponds to _______ A) an atomic mass B) an energy level C) an energy sublevel D) an atomic number

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