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CHEMISTRY 161 Chapter 8

CHEMISTRY 161 Chapter 8. ATOMIC ORBITALS. (n, l, m l , m s ). 3 s 3 p 3 d 2 s 2 p 1 s. E. H Atom Orbital Energies. energy level diagram H atom. energy depends only on principal quantum number. orbitals with same n but different l are degenerate. 4 d. 5 s. 4 p. 3 d. 4 s. 3 p.

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CHEMISTRY 161 Chapter 8

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  1. CHEMISTRY 161 Chapter 8

  2. ATOMIC ORBITALS (n, l, ml, ms)

  3. 3s 3p 3d 2s 2p 1s E H Atom Orbital Energies energy level diagram H atom energy depends only on principal quantum number orbitals with same n but different l are degenerate

  4. 4d 5s 4p 3d 4s 3p E 3s 2p 2s 1s MULTI-ELECTRON ATOM EXP I energy depends on n and ml orbitals with same n and different l are not degenerate EXAMPLES [Xe]

  5. Periodic Table of the Elements period ns2(n-1)dx g r o u p ns2np6 ns2 ns1 chemical reactivity - valence electrons

  6. PERIODIC TRENDS 1. ATOMIC RADIUS 2. IONIC RADIUS 3. IONIZATION ENERGIES 4. ELECTRON AFFINITIES

  7. ATOMIC RADIUS MAIN GROUPS ATOMIC RADIUS ATOMIC RADIUS EXP II

  8. 1s, 2s, 3s 1s 2s 3s

  9. 2px, 3px, 4px 4px 2px 3px

  10. ATOMIC RADIUS MAIN GROUPS ATOMIC RADIUS effective nuclear charge (shielding s vs p orbitals)

  11. IONIC RADII IONIC RADIUS IONIC RADIUS IONIC RADIUS

  12. cations are smaller than their atoms Na is 186 pm and Na+ is 95 pm one less electron electrons pulled in by nuclear charge anions are larger than their atoms F is 64 pm and F- is 133 pm same nuclear charge and repulsion among electrons increases radius O < O– < O2–

  13. EXAMPLES Which is bigger? Rb higher n, bigger orbitals Na or Rb K poorer screening for Ca K or Ca Ca or Ca2+ Ca bigger than cation Br or Br- Br smaller than anion

  14. QUESTION The species F-, Na+,Mg2+ have relative sizes in the order 1 F-< Na+<Mg2+ 2 F-> Na+>Mg2+ 3 Na+>Mg2+> F- 4 Na+=Mg2+= F- 5 Mg2+> Na+>F-

  15. QUESTION 1 F-< Na+<Mg2+ ALL 1s22s22p6 2 F-> Na+>Mg2+ ALL are isoelectronic 3 Na+>Mg2+> F- Na+ is 95 pm 4 Na+=Mg2+= F- Mg2+ is 66 pm F- is 133 pm 5 Mg2+> Na+>F-

  16. 3. IONIZATION ENERGIES energy required to remove an electron from a gas phase atom in its electronic ground state M(g) M+(g) + e- I1 > 0 first ionization energy (photon)

  17. I2 > 0 M+(g) M2+(g) + e- second ionization energy M2+(g) M3+(g) + e- I3 > 0 third ionization energy I1 > I2 > I3

  18. d shell insertion IONIZATION ENERGY IONIZATION ENERGY Why? first ionization energies decrease electrons closer to nucleus more tightly held

  19. IONIZATION ENERGY

  20. 1. closed shells are energetically most stable 2. half-filled shells are energetically very stable DERIVATION OF IONIZATION ENERGIES

  21. noble gases have the highest ionization energy

  22. 4. ELECTRON AFFINITIES the energy change associated with the addition of an electron to a gaseous atom X(g) + e– X–(g) electron affinity can be positive or negative

  23. general trend Why? ELECTRON AFFINITY ELECTRON AFFINITY

  24. 1. closed shells are energetically most stable 2. half-filled shells are energetically very stable DERIVATION OF ELECTRON AFFINITIES

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