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Electron Configurations

Electron Configurations. Section 4.3. Arrangement of Electrons. Electron configuration: the arrangement of electrons in an atom Electrons in atoms assume arrangements that have the lowest possible energies This makes them more stable

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Electron Configurations

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  1. Electron Configurations Section 4.3

  2. Arrangement of Electrons • Electron configuration: the arrangement of electrons in an atom • Electrons in atoms assume arrangements that have the lowest possible energies • This makes them more stable • Ground-state electron configuration: the lowest energy arrangement of the electrons for each element

  3. Rules for Electron Location • Aufbau principle: an electron occupies the lowest-energy orbital that can receive it • Pauli exclusion principle: no two electrons in the same atom can have the same set of four quantum numbers

  4. Last Rule • Hund’s rule: orbitals of equal energy are each occupied by one electron before any orbital is occupied by a second electron and all electrons in single occupied orbitals must have the same spin state

  5. Electron Configuration • Find the element on the periodic table • Start with the lowest energy level, fill the sublevels with electrons until you get to your element • The principle quantum number is written first, followed by the sublevel superscripted with the number of electrons in that sublevel

  6. Use the Periodic Table to Find the Sublevels for Your Element n - 1 n - 2

  7. Blocks of Sublevels 1 2 3 4 5 6 7 3d1

  8. Examples • Oxygen has 8 electrons • The first two electrons will go in the 1s orbital, 1s2 • The next two go in the 2s, 2s2 • The next 4 go in the 2p orbitals for 2p4 • Complete electron configuration: • 1s22s22p4 (make sure you have 8 e-)

  9. Another Way for Aufbau

  10. Exception • You would expect chromium’s electron configuration to be: • 1s22s22p63s23p64s23d4 • But it is more stable if the 3d sublevel is half full, so what really happens is: • 1s22s22p63s23p64s13d5 as Cr steals a 4s electron and puts it in the 3d sublevel

  11. Another Exception • You would expect copper’s electron configuration to be: • 1s22s22p63s23p64s23d9 • But it is more stable if the 3d sublevel is full, so what really happens is: • 1s22s22p63s23p64s13d10 as Cu steals a 4s electron and puts it in the 3d sublevel

  12. Orbital Notation (Diagrams) • Sometimes we want to see exactly which orbital is occupied and we want to specify the spin • Only two electrons can be in any one orbital, and they must have opposite spin • Every orbital in a sublevel gets one electron before any get two, and they all have the same spin

  13. Basic Directions • Write the electron configuration • Write lines based on the orbitals • One line for s, 3 lines for p, 5 lines for d • Draw arrows to represent the electrons, filling according to Hund’s rule and the Pauli exclusion principle

  14. Practice • For oxygen: • 1s22s22p4 • ___ ___ ___ ___ ___ • 1s 2s 2p • 8 electrons have been placed and Hund’s rule and the Pauli exclusion principle has been followed

  15. Shorthand or Noble Gas Notation • Noble gases: group 18 elements • Noble gas notation: AKA shorthand, when the preceding noble gas’s configuration is followed by the rest of the element’s configuration • Example: for vanadium write the preceding noble gas in brackets, [Ar] followed by 4s23d3 to be written as [Ar] 4s23d3

  16. Excited State vs. Ground State • An atom is in an excited state when the electron configuration shows electrons in higher than normal orbitals • Na would normally be: 1s22s22p63s1 • An excited state would be 1s22s22p53s2

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