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Electron Configurations

Electron Configurations. Review of an Atom . This isn’t an accurate depiction. . Electron clouds are 3D, not flat Electrons are spread out as much as possible, not usually round, and are moving rapidly These things are hard to see in a still picture. Orbitals.

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Electron Configurations

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  1. Electron Configurations

  2. Review of an Atom

  3. This isn’t an accurate depiction. • Electron clouds are 3D, not flat • Electrons are spread out as much as possible, not usually round, and are moving rapidly • These things are hard to see in a still picture

  4. Orbitals • Electrons spread out in orbitals • Orbitals have different SHAPES and ENERGY (distance from nucleus) • Quantum numbers describe orbitals • There are four quantum numbers

  5. Principal Quantum Number (n) • First of the four (n, #, #, #) • Describes the distance from the nucleus to the orbital (and therefore describes the ENERGY of the orbital) • Values: integers ≥ 1 • As the distance from the nucleus increases, n increases. • As the energy increases, n increases.

  6. Principal Quantum Number

  7. n relates to the Periodic Table! • Each period as a unique n value • Period 1: n=1 • Period 2: n=2 • Period 3: n=3 • ETC

  8. Lyman Series • The transition from n≥2 to n=1in a hydrogen atom • Result: ultraviolet emission lines of the hydrogen atom • Greater the difference in the principal quantum numbers, the higher the energy of the electromagnetic emission

  9. Lyman Series

  10. Balmer Series • The transition from n≥3 to n=2 in a hydrogen atom • Result: spectral line emissions of the hydrogen atom • As the n value increases, the wavelength emitted decreases (in nm)

  11. Balmer Series

  12. Paschen Series • The transition from n≥4 to n=3 in a hydrogen atom • Result: emission lines in the infrared band

  13. Paschen Series

  14. Angular Momentum (l) • Second of the four (n, l, #, #) • Shape of the sublevel • Range from 0 to n-1 (we will never deal with anything above l=3) • l=0 = s • l=1 = p • l=2 = d • l=3 = f

  15. s sublevel

  16. p sublevel

  17. d sublevel

  18. f orbitals

  19. Magnetic Number (ml) • Third of the four numbers (#, #, ml , #) • Denotes the orbital sublevel that is filled • s sublevel has ONE orbital (sphere has one orientation in space) • p sublevel has THREE orbitals (three orientations in space) • d sublevel has FIVE orbitals (five orientations in space) • f sublevel has SEVEN orbitals (seven orientations in space)

  20. Magnetic Number • s sublevel: one orbital • One orientation in space

  21. Magnetic Number • P sublevels: three orbitals • Three orientations in space

  22. Magnetic Number • d sublevels: five orbitals • Five orientations in space

  23. Magnetic Number • f sublevel: seven orbitals • Seven orientations in space

  24. Magnetic Number • Integers from -l to l • SO: • s: ml = 0 only since l= 0 • p: ml = -1,0,1 since l= 1 • d: ml = -2,-1,0,1,2 since l= 2 • f: ml = -3,-2,-1,0,1,2,3 since l= 3

  25. Spin (ms) • Fourth number (#, #, #, ms) • It is either -1/2 or ½ • Down or up

  26. What does all of this tell us? • There is a very specific order in which electrons fill orbitals. It is not random. There are some exceptions.

  27. Orbital Filling Order

  28. Three Important Principles • Aufbau (next) Principle • Pauli Exclusion Principle • Hund’s Rule

  29. Aufbau • Electrons fill the LOWEST energy sublevel before going to the next sublevel • 1s fills, then 2s fills, then 2p fills, then 3s fills, then 3p fills ….

  30. Pauli Exclusion Principle • Electrons pair according to OPPOSITE spins • ↑↓, not ↑↑ or ↓↓

  31. Hund’s Rule • Electrons spread out in equal energy sublevels before pairing electrons • (↑↑↑ and not ↑↓↑ _)

  32. Step by Step… • First level to fill is 1s level • Lowest energy sublevel • Holds two electrons • They are oppositely paired • A sublevel is represented by __ and holds 2 electrons

  33. Second sublevel is the 2s sublevel • It holds 2 electrons because of s • Electrons are oppositely paired

  34. So, we filled 1s, we filled 2s • Now comes 2p • Holds six electrons because p orbitals hold 6 electrons 1s22s22p6

  35. Finally.. • From 2p, 3s fills with 2 electrons • 3p fills with 6 electrons • 4s fills with 2 electrons • 3d fills with 10 electrons • 4p fills with 6 electrons • 5s fills with 2 electrons

  36. Example: Carbon • Neutral carbon: 6 electrons ↑↓↑↓↑↑ _ 1s 2s 2p • Six arrows for six electrons • 1s2 2s2 2p2

  37. Exceptions… • Some energy levels are SUPER close together • The levels are so close that electrons are able to move between these orbitals in order to minimize repulsion… • 4s and 3d orbitals are very close in energy • Exceptions exist for some period 4 d block elements • Cr is not 1s2 2s2 2p6 3s2 3p6 4s2 3d4 • Cr is 1s2 2s2 2p6 3s2 3p6 4s1 3d5 • It actually takes LESS energy to split the electrons between the 5 sublevels than it does to put them together in the 4s and 3d

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