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					       Unit 1: Matter and
       Chemical Bonding
                   Trends in the Periodic Table

                                            Agenda
 Similarities in chemical and physical properties
Reactivity of alkali metals, alkaline earth metals,
                     non-metals, transition metals
                                             VIDEO
                                    Atomic Radius
                            Ionization Energy (IE)
                              Electron Affinity (EA)
                                Electron Negativity
Ionization Energy
   The ionization energy is the energy needed
    to remove an electron from an atom.

For an element X, it is written as:

         X + energy  X+ + e-
                         cation


The ionization energy is a measure of how
  tightly the electrons are held by the atom.
How does ionization energy
change down a group?
   The first ionization energy decreases as you
    move down a group.

   Why?
       The size of the atom increases.
       Electron is further from the nucleus.
How does ionization energy
change across a period?
   The first ionization energy increases as you
    move from left to right across a period.

   Why?
       Nuclear charge increases while shielding is
        constant.
       Attraction of the electron to the nucleus increases.
1st, 2nd, 3rd ionization energies
   Atoms with more than one electron have
    more than one ionization energy.

   energies correspond to the stepwise removal
    of electrons, one after another
Example:
        two valence electrons
        easy to remove first one, twice as hard to remove
         second one
        Remove one more major jump


            1st         2nd    3rd              4th
Be          899        1757 | 14,845           21,000
               Table 6.1, p. 173
Symbol First          Second       Third
  H    1312
  He   2731             5247
  Li   520              7297       11810
  Be   900              1757       14840
  B    800              2430       3569
  C    1086             2352       4619
  N    1402             2857       4577
  O    1314             3391       5301
  F    1681             3375       6045
  Ne   2080             3963       6276
Symbol First   Second        Third
  H     1312          Why did these values
                      increase so   much?
  He    2731   5247
  Li    520    7297          11810
  Be    900    1757          14840
  B     800    2430          3569
  C     1086   2352          4619
  N     1402   2857          4577
  O     1314   3391          5301
  F     1681   3375          6045
  Ne    2080   3963          6276
- The lower the ionization energy, the easier it is to remove the outer electron.
-The higher the ionization energy, the more difficult it is to remove the outer
electron.
Movie
   http://wps.prenhall.com/wps/media/objects/43
    9/449969/Media_Portfolio/Chapter_07/PeriTr
    ends-IonizationEnerg.MOV
Electron Affinity
   Under some circumstances, it is possible to get
    an atom to accept electrons.

   EA is the amount of energy released when one
    electron is added to an atom in the gaseous
    state.


             X + e- -----> X- + energy
                              anion
EA
   The same factors – for EA as for IE

   Across a period: (L to R)  INCREASE
        a valence shell that holds its electrons tightly will
        also tend to hold an additional electron tightly.

   DOWN a group  DECREASE
       A valence shell that loses electrons easily will
        have little attraction for additional electrons.
Electronegativity: the ability of
an atom in a bond to pull on
the electron. (Linus Pauling)
      Electronegativity
   When electrons are shared by two atoms a
    covalent bond is formed.

   When the atoms are the same they pull on
    the electrons equally. Example, H-H.

   When the atoms are different, the atoms pull
    on the electrons unevenly. Example, HCl
Reactivity
        Reactivity refers to how likely or vigorously an
         atom is to react with other substances.

        Factors:
    1)    how easily electrons can be removed
    2)    how badly they want to take other atom's electrons
          (electronegativity)

    b/c it is the transfer/interaction of electrons that is the basis of
         chemical reactions.
Metals - reactivity
   Period - decreases LR
    Group - increases  down a group

   Why?
       The farther to the left and down  the easier it is
        for electrons to be given or taken away
Non-metals Reactivity
   Period - increases LR
    Group - decreases down the group

   Why?
       The farther right and up  the higher the EN,
        resulting in a more vigorous exchange of electron.
Nuclear charge increases
Shielding increases
Atomic radius increases
Ionic size increases
Ionization energy decreases
Electronegativity decreases
                                           Summary
             Shielding is constant



             Nuclear charge increases
             Atomic Radius decreases

             Electronegativity increases
             Ionization energy increases

				
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