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CH 8: Electron Configuration

CH 8: Electron Configuration

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CH 8: Electron Configuration

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  1. CH 8: Electron Configuration Renee Y. Becker Valencia Community College CHM 1045

  2. Electron Configuration of Atoms Rules of Aufbau Principle: • Lower n orbitals fill first. • Each orbital holds two electrons; each with different ms. • Half-fill degenerate orbitals before pairingelectrons. (p, d, & f)    NOT   __ 3px 3py 3pz

  3. Electron Configuration of Atoms

  4. Electron Configuration of Atoms Element Diagram Configuration Li (Z = 3)  1s2 2s1 1s 2s Be (Z = 4)  1s2 2s2 1s 2s B (Z = 5)  __ __ 1s2 2s2 2p1 1s 2s 2px 2py 2pz C (Z = 6)   __ 1s2 2s2 2p2 1s 2s 2px 2py 2pz

  5. Electron Configuration of Atoms Element Diagram Configuration O (Z = 8)  1s2 2s2 2p4 1s 2s 2px 2py 2pz Ne (Z = 10)  1s2 2s2 2p6 1s 2s 2px 2py 2pz S (Z = 16)  1s 2s 2px 2py 2pz 3s 3px 3py 3pz 1s2 2s2 2p6 3s2 3p6 or [Ne] 3s2 3p6 abbreviations using the noble gases valence vs. core electrons

  6. Electron Configuration of Atoms

  7. Electron Configuration of Atoms Tc (Z = 43) [Kr] 5s2 4d5 Technetium Ni (Z = 28) [Ar] 4s2 3d8

  8. Electron Configuration of Atoms

  9. Electron Configuration of Atoms

  10. Example 1: Electron Config. And NG Abb. • Sodium • Titanium • Argon

  11. Anomalous Electron Configurations • 19 of the predicted configurations from the periodic table are wrong • Largely due to unusual stability of both half-filled and fully filled subshells Cr (Z=24) expected configuration: 1s2 2s2 2p6 3s2 3p6 4s2 3d4   __ 4s 3d 3d 3d 3d 3d actual configuration: 1s2 2s2 2p6 3s2 3p6 4s1 3d5     4s 3d 3d 3d 3d 3d

  12. Atomic Radii

  13. Atomic Radii • ½ the distance between the nuclei of two identical atoms when they are bonded together.

  14. Example 2: Ionic Radii Which of the following in each pair has a larger atomic radius? • Carbon or Fluorine • Chlorine or Iodine • Sodium or Magnesium • O or O2- • Ca or Ca2+

  15. Example 3: Quantum Numbers and Electron Configuration What are the 4 quantum numbers for the following? Remember you are only interested in the last electron!! • C • Na+ • S • N3-

  16. Main Groups

  17. Ions and their Electron Configuration • Main-group metals donate electrons from the atom’s highest-energy occupied atomic orbital. • Na: 1s2 2s2 2p6 3s1 = [Ne] 3s1 • Na+: 1s2 2s2 2p6 = [Ne] • Mg: 1s2 2s2 2p6 3s2 = [Ne] 3s2 • Mg2+: 1s2 2s2 2p6 = [Ne] • Al: 1s2 2s2 2p6 3s2 3p1 = [Ne] 3s2 3p1 • Al3+ 1s2 2s2 2p6 = [Ne]

  18. Ions and their Electron Configuration • Main-group nonmetals accept electrons into their lowest-energy unoccupied atomic orbital. • N: 1s2 2s2 2p3= [He] 2s2 2p3 • N3–: 1s2 2s2 2p6 = [He] 2s2 2p6 = [Ne] • O: 1s2 2s2 2p4 = [He] 2s2 2p4 • O2–: 1s2 2s2 2p6 = [He] 2s2 2p6 = [Ne] • F: 1s2 2s2 2p5 = [He] 2s2 2p5 • F–: 1s2 2s2 2p6 = [He] 2s2 2p6 = [Ne]

  19. Example 4: Electron config. and NG Abb. • Cl- • F- • Ca2+ • Na+

  20. Ionic Radii or size • Atoms shrink when an electron is removed to form a cation • Dec. # of shells • Inc. Zeff : Less electrons, less shielding, outer electrons more attracted to nucleus, therefore smaller more compact

  21. Ionic Radii or size • Atoms expand when converted to anions • III A ns2 np1 __ __ __ • IV A ns2 np2 __ __ __ • V A ns2 np3 __ __ __ • VI A ns2 np4 __ __ __ • VII A ns2 np5 __ __ __ Adding one electron to each of these will not add another shell it will just fill an already occupied p subshell • Therefore the expansion is due to the decrease in Zeff and the increase in the electron-electron repulsions

  22. Ionization Energy, Ei • The amount of energy needed to remove the highest-energy electron from an isolated neutral atom in the gaseous state Increase Increase

  23. Ionization Energy, Ei • Some exceptions/irregularities to general trend • Ei Be > Ei B we would expect opposite • Be 4 e 1s2 2s2 • B 5 e 1s2 2s2 2p1 • 2s is closer to nucleus than 2p, Zeff for Be is stronger • 2s is held more tightly and is harder to remove

  24. Ionization Energy, Ei • Ei N > Ei O we would expect opposite • N 7e 1s2 2s2 2p3 __ __ __ • O 8e 1s2 2s2 2p4 __ __ __ • Only difference is that an electron is being removed from a half-filled orbital (N) and one from a filled orbital (O) • Electrons repel each other and tend to stay as far apart as possible, electrons that are forced together in a filled orbital are slightly higher in energy so it is easier to remove one • Therefore O < N

  25. Higher Ionization Energy, Ei1234… • Ionization is not limited to one electron M + Energy  M+ + e Ei1 M+ + Energy  M2+ + e Ei2 M2+ + Energy  M3+ + e Ei3 • Larger amts. Of energy are needed for each successive ionization, harder to remove an electron from a positively charger cation • The energy differences between successive steps vary from one element to another. Why? EC

  26. Higher Ionization Energy, Ei1234… • Easy to remove an electron from a partially filled valence shell • Difficult to remove an electron from a filled valence shell • Large amount of stability associated with filled s & p subshells • Na: 1s2 2s2 2p6 3s1 • Mg: 1s2 2s2 2p6 3s2 • Cl: 1s2 2s2 2p6 3s2 3p5

  27. Electron Affinity, Eea • Energy change that occurs when an electron is added to an isolated atom in the gaseous state. • The more neg. the Eea the greater the tendency of the atom to accept an electron • Group 7A (halogens) have the most neg. Eea, high Zeff and room in valence shell • Group 2A and 8A have near zero or slightly positive Eea

  28. Alkali Metals • Group 1A • Metallic • Soft • Good Conductors • Low MP • Lose 1 elec in redox, powerful reducing agent • Very reactive • Not found in elemental state in nature

  29. Alkaline Earth Metals • Group 2A • Harder, but still relatively soft • Silvery • High MP than group 1A • Less reactive than group 1A • Lose 2 e in redox, powerful reducing agent • Not found in elemental form in nature

  30. Group 3A • All but Boron • Silvery • Good conductor • Relatively soft • Less reactive than 1A & 2A

  31. Halogens • Group 7A • Non-metals • Diatomic molecules • Tend to gain e during redox

  32. Noble Gases • Group 8A • Colorless, odorless, unreactive gases • Ns2 np6 • Makes it difficult to add e or remove e

  33. Octet Rule • Group 1A tends to lose their ns1 valence shell electron to adopt a noble gas electron config. • Group 2A lose both ns2 “ “ • Group 3A lose all three ns2 np1 “ “ • Group 7A Gains one electron to attain NG • Group 8A inert, rarely lose or gain electrons