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Electron Configurations & Quantum Numbers

Electron Configurations & Quantum Numbers

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Electron Configurations & Quantum Numbers

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  1. Electron Configurations & Quantum Numbers • Atomic structure defined by electron levels • Principal Quantum Number “N” • Small integer numbers allowed (e.g. 1, 2, 3 …) • N defines effective size (radius) of the electron orbit • The radius determines the energy (larger = more) • N defines shells, groups of electron s with same radius • Energy of the electron increases with larger values of “N”

  2. Quantum Numbers • Angular Momentum Quantum Number “L” • Defines the shape or path of the orbiting electron • Provides 3-D perspective • “S” is spherical • “P” is dumbbell shape • Also has small integer numbers, 0 to (N-1) allowed • NO negative values • Historical identification by letters as well as numbers • O  “s”, 1  “p”, 2  “d”, 3  “f”, etc. • Chemical structures are a direct result

  3. Quantum Numbers • Magnetic Quantum Number “m” • Also a set of small integers, negative values ARE allowed • Range of values “- L” to “+ L” (momentum numbers) • The + and – values relate to “spin” • Spin conceptually like current creating a magnetic field • Electron orbits usually have paired electrons opposite spin • Net magnetic moment is zero • Most materials are non-magnetic (spins cancelled) • Magnetic materials have unpaired spins • Fe, Co, Ni • Gives rise to permanent and reversible magnets

  4. Nomenclature

  5. Simplified chart of energy levels

  6. Another way to look at electronsorganized by increasing element number, not properties

  7. A few examples • Helium, element 2  1s2 • Boron, element 5  1s2 2s2 2p • Alternative description [He]2s2 2p • Neon, element 10  1s2 2s2 2p6 • Sodium, element 11  1s2 2s2 2p6 3s • Alternative description [Ne] 3s • Chlorine, element 17  1s2 2s2 2p6 3s23p5 • Alternate description [Ne] 3s23p5

  8. As electrons are added, the quantum numbers to build up orbitals and create new elements (see text table 3.2)

  9. Orbital Shapes: s, p, & ds=2e, p=6e, d=10e

  10. Writing Electron Configurations • Use the alphabetic abbreviation for shells (e.g. s, p, d, f) • List shells in numerical order 1, 2, 3 …. • List number of shell electrons in superscript • S shell ≤2, p shell ≤6, d shell ≤10, f shell ≤14 • Keep going until all electrons accounted for • Example • Sodium metal, Na0 Z=11 (protons = electrons) • 1s2 2s2 2p6 3s1 • Sodium ion, Na1+ Z=11 (11 protons + 10 electrons) • 1s2 2s2 2p6 • Alternative to utilize nearest (lower) inert gas shell • [Ne]3s1 for sodium metal, [Ne]1+ for sodium ion • A handy abbreviation for large atoms and ions • Electron shells of identical configurations are ISOELECTRONIC • Neon gas (1s2 2s2 2p6 ) &Sodium ion (1s2 2s2 2p6 )are Isoelectronic • A favorite exam question !

  11. Hund’s RuleElectrons add with same spin until all orbit positions have 1 electron, then pair up until shells are full.Gives rise to magnetic properties of transition elements

  12. Valence Electrons • Outer electrons undergo transfer • Outer electrons are most loosely held • Oxidation = loss of available electrons • Fe0 Fe++ + 2e- • Octet rule favors complete shell of 8 • Reduction of Chlorine = gain of electrons • Cl0 + e- Cl-

  13. Another way to look at electronsorganized by increasing element number, not properties

  14. Oxidation of Sodiumand resulting electronic configuration

  15. Electron Configurations by groupwhere n is principal quantum numbersum of superscripts is available electrons

  16. Another way to look at electronsorganized by increasing element number, not properties

  17. S orbitals are spherical

  18. Sketch of Neon, 1s22s22p6

  19. d-orbitals have complex shape10 electrons involved

  20. Electronic designation

  21. Hybrid Orbitals

  22. Quantum Numbers

  23. Electrons enter and fill orbitals according to four rules:

  24. Electron Addition Order

  25. Another sequence diagram …

  26. 2

  27. 3

  28. 4

  29. 5

  30. 6

  31. Another way to look at electronsorganized by increasing element number, not properties

  32. 7

  33. 8

  34. n=1, l=0, m=0

  35. Bohr-Sommerfeld

  36. Orbital shapes as defined by L = 0, 1, 2

  37. 9

  38. “3 d” orbitals

  39. 10

  40. n=3, l=2, m=1

  41. n=3, l=2, m=2

  42. n=4, l=2, m=2

  43. More representations …

  44. 12

  45. 14

  46. 15

  47. 16

  48. 17