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ATOMIC ELECTRON CONFIGURATIONS AND PERIODICITY

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ATOMIC ELECTRON CONFIGURATIONS AND PERIODICITY Powered By Docstoc
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PERIODIC
 TRENDS
                               2


    PERIODICITY
Period Law-
    -physical and chemical
properties of elements are a
 periodic function of their
     atomic numbers
  General Periodic Trends
                                         3



       • Atomic and ionic size
          • Ionization energy
• Electron affinity, electronegativity
                           4
Effective Nuclear Charge
           Z*
Effective Nuclear Charge, Z*
                                          5




• Atom   Z* Experienced by Electrons in
         Valence Orbitals
•   Li   +1.28
•   Be   -------
•   B    +2.58         Increase in
•   C    +3.22         Z* across a
•   N    +3.85         period
•   O    +4.49
•   F    +5.13
       General Periodic Trends
                                             6




           Higher effective nuclear charge
           Electrons held more tightly




Larger orbitals.
Electrons held less
tightly.
                                           7
Lithium

                   Periodic Trend in the
                   Reactivity of Metals
Sodium




Potassium




            MOST
2.     Reactivity for Metals
 As you go down a group for metals
  the number of energy levels
  increase.
 Because of this, reactivity increases
  because the atom is more willing to
  give away its electron (react).
3.Nonmetalic Trends: Gain electrons
   Nonmetals on right side, form anions

   Going right elements are more
    nonmetallic (better gainers of electrons)

   Going UP elements become more
    nonmetallic (want to gain)
8. Reactivity nonmetals: Gain e
   The reason Across = fill the energy level

   Going UP a group, nonmetals have
    same valence but fewer total electrons

   Flourine is the most reactive nonmetal.
Atomic Radii                11




               Figure 8.9
                                            12

Atomic Size



• Size increases, down a group.
• Because electrons are added into
  additional energy levels, there is less
  attraction.
• Size decreases across a period.
• Because, increased effective nuclear
  charge.
                                           13


           Atomic Size
Size decreases across a period owing
  to increase in Z*. Each added electron
  feels a greater and greater + charge.




 Large                           Small
                                                                                        14

          Trends in Atomic Size
                            See Figures 8.9 & 8.10
Radius (pm)
250
                                           K

                            3rd period             1st transition
200
                                                   series
          2nd period         Na
           Li
150



                                                                              Kr
100
                                         Ar
                            Ne
 50


          He
  0
      0         5      10         15          20         25         30   35        40

                                       Atomic Number
                                         15

            Ion Sizes
                   +
                             Forming
Li,152 pm     Li + , 78 pm   a cation.
3e and 3p     2e and 3 p

• CATIONS are SMALLER than the
  atoms from which they come.
• The electron/proton attraction
  has gone UP and so size
  DECREASES.
                                          16

              Ion Sizes
                    -
                              Forming
F, 71 pm       F- , 133 pm    an anion.
9e and 9p      10 e and 9 p

 • ANIONS are LARGER than the atoms
   from which they come.
 • The electron/proton attraction has
   gone DOWN and so size INCREASES.
 • Trends in ion sizes are the same as
   atom sizes.
                               17

Trends in Ion Sizes




                 Figure 8.13
                                             18

         Ionization Energy

IE = energy required to remove an electron
  from an atom in the gas phase.




   Mg (g) + 738 kJ ---> Mg+ (g) + e-
                                             19

          Ionization Energy

IE = energy required to remove an electron
  from an atom in the gas phase.
   Mg (g) + 738 kJ ---> Mg+ (g) + e-




   Mg+ (g) + 1451 kJ ---> Mg2+ (g) +
                   e-
    Mg+ has 12 protons and only 11
 electrons. Therefore, IE for Mg+ > Mg.
                                                     20

              Ionization Energy
       1st IE: Mg (g) + 735 kJ ---> Mg+ (g) + e-
      2nd IE: Mg+ (g) + 1451 kJ ---> Mg2+ (g) + e-




3rd IE: Mg2+ (g) + 7733 kJ ---> Mg3+ (g) + e-
  Energy cost is very high to dip into a
  shell of lower n (core electrons).
  This is why ox. no. = Group no.
                                                                                                    21

           Trends in Ionization Energy
       1st Ionization energy (kJ/mol)
2500
            He
                             Ne
2000


                                                  Ar
1500
                                                                                               Kr

1000



 500



   0
       1    3    5   7   9    11   13   15   17    19      21   23   25   27   29   31   33   35
       H   Li                Na                        K
                                                                     Atomic Number
                                 22


   Trends in Ionization Energy

   • IE decreases down a
            group
• Because size increases.


• IE increases across a period
• Because effective nuclear
      charge increases
                                     23


     Electron Affinity
     A few elements GAIN
   electrons to form anions.
Electron affinity is the energy
    involved when an atom
 gains an electron to form an
             anion.
  X(g) + e- ---> X-(g)   E.A. = ∆E
                              24

Trends in Electron Affinity
                                        25



   Trends in Electron Affinity

• Affinity for electron   Atom EA
  increases across a      F   -328 kJ
  period (EA becomes      Cl -349 kJ
  more negative).
                          Br -325 kJ
                          I   -295 kJ
• Affinity decreases
  down a group (EA
  becomes less
  negative).
                                                26



       Electron Affinity of Oxygen


O atom [He]                 ∆E is
                               EXOthermic
                 + electron
                                because O
                              has an affinity
O- ion [He]          
                                 for an e-.
          EA = - 141 kJ
                                                    27



Electron Affinity of Nitrogen
N atom [He]                ∆E is zero for N-
                                      due to
                   + electron
                                    electron-
N- ion   [He]          
                                     electron
                                   repulsions.
              EA = 0 kJ
                                         28


      Electronegativity
     • So how is this different from
             electron affinity?
• Electron Affinity – is rating of how
  well an atom wants to gain an
  electron
• Electronegativity – is rating of how
  well an atom keeps the electron
  once it is bonded to another atom
                    29


Electronegativity
Electron Configurations
                          30


and the Periodic Trends
                        31


   “Your Best Friend”
• Periodic table

				
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