Electron Configurations

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Electron Configurations

  1. 1. Orbitals and Electron Configurations Where are the electrons? Tuesday, October 7th 2008
  2. 2. The Rutherford Atom
  3. 3. Problems with Rutherford’s Atom <ul><li>Electrons should be attracted to the nucleus and repel each other </li></ul><ul><li>Couldn’t answer the question </li></ul><ul><ul><li>Why do the electrons stay in the electron cloud? </li></ul></ul><ul><ul><li>Why don’t the electrons collapse into the massively positive nucleus? </li></ul></ul>
  4. 4. The Bohr Atom
  5. 5. Problems with the Bohr Atom <ul><li>Fundamentally incorrect - only worked for the element Hydrogen </li></ul><ul><li>Couldn’t explain where the electrons were in atoms that had more than one electron </li></ul><ul><li>We don’t really know where an electron is at any one time, and we can’t predict it either </li></ul>
  6. 6. Things Bohr got right <ul><li>Energy Levels </li></ul><ul><li>Ground State </li></ul>
  7. 7. What we saw in the flame test lab
  8. 8. How we saw the light in the flame test lab
  9. 9. How we explain the light in terms of energy levels of the electrons in the atom
  10. 10. Max Planck’s contribution <ul><li>German Physicist in the early 1900’s </li></ul><ul><li>Said that light is made up of discrete bundles of energy called “quanta” (plural of quantum) </li></ul><ul><li>Now known as “photon” </li></ul>
  11. 11. The difference between continuous and quantized energy levels
  12. 12. Wave-Mechanical Model <ul><li>Also called the quantum-mechanical model </li></ul><ul><li>Electrons behave like light and only have specific energy levels </li></ul><ul><li>This idea worked for all atoms, not just hydrogen (yay!) </li></ul>
  13. 13. Wave Mechanical Model <ul><li>Electrons do not follow definite paths </li></ul><ul><li>Electrons are in a diffuse cloud of negative charge around the nucleus (like the Rutherford atom) </li></ul><ul><li>There are areas around the nucleus that correspond with certain energy levels (like the Bohr Model) </li></ul><ul><li>The areas around the nucleus where the electron probably is (energy levels) are called orbitals </li></ul>
  14. 14. Orbitals <ul><li>Do not have distinct boundaries (like earth’s atmosphere) </li></ul><ul><li>Boundary is mapped at 90% electron probability (by convention) </li></ul><ul><li>Electrons can be found outside of this boundary </li></ul><ul><li>We can never map exactly where an electron is at any given moment </li></ul><ul><li>All elements have all of the orbitals, they just don’t use all of them all the time </li></ul>
  15. 15. The Hydrogen 1s Orbital
  16. 16. The first four principle energy levels
  17. 17. Sub-Levels <ul><li>As the Energy Level number increases, the further away from the nucleus the electron is, and the higher the energy level </li></ul><ul><li>Each Energy level is divided further into sub-levels </li></ul>
  18. 18. Principle energy levels are divided into sub-levels (s,p,d,f)
  19. 19. Second Principle Energy Level with sublevels corresponding to orbitals
  20. 20. 1s and 2s orbitals
  21. 21. The 2p orbitals (three of them)
  22. 22. Diagram of Principle Energy Levels 1 and 2
  23. 23. Relative size of the 1s, 2s, 3s orbitals
  24. 24. The 3d orbitals
  25. 25. Electron Filling <ul><li>Aufbau Principle - electrons prefer the space closest to the nucleus </li></ul><ul><li>Therefore all of the electrons are arranged around the nucleus from lowest energy level to highest energy level </li></ul><ul><li>The most attractive orbital to any electron is the 1s orbital, then 2s, 2p, 3s, 3p, 4s, 3d, and so on </li></ul><ul><li>This corresponds to Bohr’s idea of the ground state </li></ul>
  26. 26. Electron filling <ul><li>Pauli Exclusion Principle - orbitals can hold a maximum of two electrons </li></ul><ul><li>Electrons repel each other and don’t want to share the same space (because they’re both negative) </li></ul><ul><li>Electrons will share the same space if they are spinning in opposite directions (like a magnet) </li></ul>

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