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Electron configuration

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Electron configuration

  1. 1. • Principal Quantum Number, n • Azimuthal Quantum Number, l • Magnetic Quantum Number, ml
  2. 2. Exercise • What are the similarities and differences between the hydrogen atom 1s and 2s orbitals? • For n=4 what are the possible values of l and m? • give the values for n, l, m for a. 2p b. 5d c. 5f
  3. 3. Main Energy Level No. of Sublevel Identity of Sublevels No. of Orbitals (n2) Max. No. of Electrons (2n2) 1 1 1s 1 2 2 2 2s 2p 1 3 2 6 3 3 3s 3p 3d 1 3 5 2 6 10 4 4 4s 4p 4d 4f 1 3 5 7 2 6 10 14
  4. 4. Electron Spin Quantum Number, m • Two possible values are allowed + ½ and - ½ • Opposite directions
  5. 5. Can two electrons in an atom have the same set of quantum numbers?
  6. 6. Pauli Exclusion Principle • No two electrons in an atom can have the same set of four quantum numbers, n, l, ml, and ms. • An orbital can hold a maximum of two electrons and they must have opposite spins.
  7. 7. Electron Configuration • The way in which electrons are distributed among the various orbitals of an atom
  8. 8. Aufbau Principle • Electrons occupy orbitals of lower energy first.
  9. 9. The relation between orbital filling and the periodic table
  10. 10. The relation between orbital filling and the periodic table
  11. 11. Hund’s Rule • For degenerate orbitals, the lowest energy is attained when the number of electrons with the same spin is maximized.
  12. 12. Orbital Diagram
  13. 13. Anomalous Electron Configuration Example: Chromium Copper
  14. 14. • Consequence of the closeness of the 3d and 4s orbital energies. • Minor departures from the expected • Not of great chemical significance
  15. 15. Seat work Write the electron configuration and the electron orbital diagram for the following elements. 1. Cl 2. Os 3. Cs

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