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					Lecture 6 - Atomic Theory 3
     Review

        n1
     0  l  n-1
     -l  ml  l
 1 type of s orbital
3 types of p orbital
5 types of d orbital
7 types of f orbital
            Today...


Filling the orbitals with electrons

        Periodic Trends
There is a fourth quantum
         number

       spin, ms

   ms = +1/2 or -1/2
    (up or down)
Pauli’s Exclusion Principle:

 no two electrons in an atom
      have the same four
      quantum numbers
Filling Orbitals with Electrons

      “Aufbau” - filling up

 put two electrons in each type
of each orbital (one of each spin)
Filling Orbitals with Electrons

          Hydrogen:
 one electron in the 1s orbital

electronic configuration is 1s1
Filling Orbitals with Electrons

           Helium:
two electrons in the 1s orbital

electronic configuration is 1s2
     (first shell complete)
Filling Orbitals with Electrons

            Lithium:
 two electrons in the 1s orbital,
   then one in the 2s orbital

electronic configuration is 1s22s1
            or [He] 2s1
Filling Orbitals with Electrons

          Beryllium:

            1s22s2
              or
           [He]2s2
   Filling Orbitals with Electrons

Boron:      [He]2s22p1
Carbon:     [He]2s22p2
Nitrogen:   [He]2s22p3
Oxygen:     [He]2s22p4
Fluorine:   [He]2s22p5
Neon:       [He]2s22p6 (second shell full)
       Hund’s Rule

put electrons in a sub-shell
  spin-up until each orbital
        contains one

   then add spin-down
         electrons
     e.g. fluorine (1s22s22p5)

2p

2s
              degenerate p-orbitals
1s            (same energy level)
      e.g. oxygen   (1s22s22p4)



2p                  2p

2s                  2s

1s                  1s

     wrong               right
Orbital Filling Order

  1s
  2s   2p
  3s   3p   3d
  4s   4p   4d 4f
  5s   5p   5d 5f
  6s   6p   6d
s-block
                              p-block
          d-block




                    f-block
   Important Concept

Chemistry is determined by
  the outer shell electrons

    (valence electrons)
           Periodicity

   - regular changes of physical
properties across a period or down a
                group

  e.g. atomic radius, ionic radius,
ionization potential, electronegativity
Atomic Radius, e.g. I
     133 pm 133 pm
                     Periodicity - Atomic Radius
                                             Rb   Fr
                    300                 K
Atomic Radius, pm



                                   Na
                              Li
                    150
                                        Cl   Br   I
                                    F
                     0
                          0             Atomic Number   90
       Radius Down Group 6A

Atom    Outer Shell    Radius, pm
 O        2s22p4         66
 S        3s23p4         104
 Se       4s23d103p4     116
 Te       5s24d105p4     143
 Po       6s24f145d106p4 167
Radius Across the Third Period

Atom   Outer Shell   Radius, pm
Na       3s1           186
Mg       3s2           160
Al       3s23p1        143
Si       3s23p2        118
P        3s23p3        110
S        3s23p4        103
       Atomic Radius:

 Increases going down a group

Decreases going L to R across a
            period
   Making Ions

Metals lose electrons
 e.g. Na Na+ + e-

Non-metals gain them
 e.g. S + 2 e- S-2
         Na Na+ + e-

1s2 2s2 2p6 3s1 1s2 2s2 2p6 + e-

       [Ne] 3s1 [Ne] + e-
         S + 2 e- S-2

[Ne] 3s2 3p4 + 2e- [Ne] 3s2 3p6
                or
    [Ne] 3s2 3p4 + 2e- [Ar]
    Important Concept

 Atoms tend to attain electron
configurations of the noble gases
                 by:
            - ionizing
            - reacting
          “Octet Rule”
Isoelectronic Atoms


  Ne      1s2 2s2 2p6
  Na+     1s2 2s2 2p6
  Mg+2    1s2 2s2 2p6

  are isoelectronic
                Na, 186 pm Mg, 160 pm
Ionic Radii


Cations are
smaller than   Ne, 150 pm
neutral ions

                Na+, 95 pm Mg+2, 65 pm
       Anions are larger...

     e.g. Cl2 + 2 e-  2 Cl-

      99 pm            181 pm

(adding electrons to the same shell)
Removing electrons requires
         energy

 e.g. Na + energy Na+ + e-

   energy = I1 = 496 kJ/mol
    (1st ionization potential)
                   Ionization Potentials
    2500 He
                       Ne
                            Ar
I1, kJ mol-1



                                      Kr        Xe
                                                     Hg Rn
                                 Zn        Cd

                   Li Na K            Rb        Cs       Fr
               0
                   0             Atomic Number               100
     Adding electrons releases
             energy
      e.g. F + e- F- + energy

          EA = -328 kJ/mol
          (electron affinity)

EA increases going up and to the right
         on the periodic table

				
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