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Minooka -Electron Configurations Part 1
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Minooka -Electron Configurations Part 1

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  • 1. Orbitals and Electron Configurations Where are the electrons?
  • 2. The Rutherford Atom
  • 3. Problems with the Rutherford Atom
    • Electrons should be attracted to the nucleus and repel each other
    • Couldn’t answer the question
      • Why do the electrons stay in the electron cloud?
      • Why don’t the electrons collapse into the massively positive nucleus?
  • 4. The Bohr Atom
  • 5. Problems with the Bohr Atom
    • Fundamentally incorrect - only worked for the element Hydrogen
    • Couldn’t explain where the electrons were in atoms that had more than one electron
    • We don’t really know where an electron is at any one time, and we can’t predict it either
  • 6. Things Bohr got right
    • Energy Levels
    • Ground State
  • 7. What we saw in the flame test lab
  • 8. How we saw the light in the flame test lab
  • 9. How we explain the light in terms of energy levels of the electrons in the atom
  • 10. The difference between continuous and quantized energy levels
  • 11. Albert Einstein’s contribution
    • Developed quantum mechanics
    • Showed that Isaac Newton’s theories for motion do not give correct results when objects are traveling close to the speed of light
    • New equations in which the laws for motion are adjusted for the speed of light
  • 12. Max Planck’s contribution
    • German Physicist in the early 1900’s
    • Said that Light is made up of discrete bundles of energy called “quanta” (pleural of quantum)
    • Now known as “photon”
  • 13. Light has a Dual Nature
    • Behaves as both a particle (has properties of matter)
    • And a wave
      • Has wave properties such as
        • Wavelength
        • Frequency
        • Speed (velocity)
  • 14. Photons of red and blue light
  • 15. Light as both a wave and a packet of energy
  • 16. Schrodinger and De Broglie
    • Mid-1920’s
    • Louis Victor De Broglie from France
    • Erwin Schrodinger from Austria
    • Both young Physicists
    • Suggested that if light can act like a wave and a particle, then perhaps the same was true of the electron
  • 17. Wave-Mechanical Model
    • Also called the quantum-mechanical model
    • Electrons behave as both waves and particles (like light)
    • De Broglie and Schrodinger applied a mathematical analysis to their idea and found that it worked for all atoms, not just hydrogen
  • 18. Wave Mechanical Model
    • Electrons do not follow definite paths
    • Electrons are in a diffuse cloud of negative charge around the nucleus (like the Rutherford atom)
    • There are areas around the nucleus that correspond with certain energy levels (like the Bohr Model)
    • The areas around the nucleus where the electron probably is (energy levels) are called orbitals
  • 19. Firefly experiment
  • 20. Electron Probability
  • 21. The Hydrogen 1s Orbital
  • 22. Orbitals
    • Do not have distinct boundaries (like earth’s atmosphere)
    • Boundary is mapped at 90% electron probability (by convention)
    • Electrons can be found outside of this boundary
    • We can never map exactly where an electron is at any given moment
    • All elements have all of the orbitals
  • 23. The first four principle energy levels
  • 24. Sub-Levels
    • As the Energy Level number increases, the further away from the nucleus the electron is, and the higher the energy level
    • The further away from the nucleus the energy level is, the more space there is to divide up
    • Each Energy level is divided further into sub-levels
  • 25. How principle energy levels are divided into sub-levels (s,p,d,f)
  • 26. Second Principle Energy Level with sublevels corresponding to orbitals
  • 27. 1s and 2s orbitals (showing the relative size)
  • 28. The 2p orbitals (three of them)
  • 29. Diagram of Principle Energy Levels 1 and 2
  • 30. Relative size of the 1s, 2s, 3s orbitals
  • 31. The 3d orbitals
  • 32. Electron Filling
    • Aufbau Principle - electrons prefer the space closest to the nucleus
    • Therefore all of the electrons are arranged around the nucleus from lowest energy level to highest energy level
    • The most attractive orbital to any electron is the 1s orbital, then 2s, 2p, 3s, 3p, 4s, 3d, and so on
    • This corresponds to Bohr’s idea of the ground state
  • 33. Electron filling
    • Pauli Exclusion Principle - orbitals can hold a maximum of two electrons
    • Electrons repel each other and don’t want to share the same space (same negative charge)
    • Electrons will share the same space if they are spinning in opposite directions (like a magnet)