Periodic Table History & Organization - Fulton County Schools Home by zwftoTP

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									Ch. 6 - The Periodic Table &
        Periodic Law

              I. Development
               of the Modern
              Periodic Table
                 (p. 174 - 181)



                          I   II   III
A. Mendeleev

Dmitri Mendeleev (1869, Russian)
 Organized elements
  by increasing
  atomic mass
 Elements with
  similar properties
  were grouped
  together
 There were some
  discrepancies
A. Mendeleev

Dmitri Mendeleev (1869, Russian)
 Predicted properties of undiscovered
  elements
B. Moseley

Henry Moseley (1913, British)
  Organized elements by increasing
   atomic number
  Resolved discrepancies in Mendeleev’s
   arrangement
  This is the way the periodic table is
   arranged today!
    C. Modern Periodic Table

        Group (Family)
        Period
1
2
3
4
5
6
7
1. Groups/Families

Vertical columns of periodic table

Numbered 1 to 18 from left to right

Each group contains elements with similar
 chemical properties
2. Periods

Horizontal rows of periodic table

Periods are numbered top to bottom from
 1 to 7

Elements in same period have similarities
 in energy levels, but not properties
3. Blocks

     Main Group Elements
     Transition Metals
     Inner Transition
      Metals
    3. Blocks


1      Overall Configuration
2
3
4
5
6
7


         Lanthanides - part of period 6
         Actinides - part of period 7
Ch. 6 - The Periodic Table

  II. Classification of the
          Elements
       (pages 182-186)




                          I   II   III
    A. Metallic Character
         Metals
1        Nonmetals
2        Metalloids
3
4
5
6
7
1. Metals

Good conductors of heat and electricity
Found in Groups 1 & 2, middle of table in
 3-12 and some on right side of table
Have luster, are ductile and malleable
a. Alkali Metals

Group 1
1 Valence electron
Very reactive
Electron configuration
 ns1
Form 1+ ions
Cations
 Examples: Li, Na, K
b. Alkaline Earth Metals

Group 2
Reactive (not as reactive as alkali metals)
Electron Configuration
 ns2
Form 2+ ions
Cations
 Examples: Be, Mg, Ca, etc
c. Transition Metals

Groups 3 - 12
Reactive (not as reactive as Groups 1 or
 2), can be free elements
Electron Configuration
  ns2(n-1)dx where x is column in d-block
Form variable valence state ions
Cations
  Examples: Co, Fe, Pt, etc
2. Nonmetals

Not good conductors
Found on right side of periodic table –
 AND hydrogen
Usually brittle solids or gases
a. Halogens

Group 17 (7A)
Very reactive
Electron configuration
  ns2np5
Form 1- ions – 1 electron short
 of noble gas configuration
Anions
  Examples: F, Cl, Br, etc
b. Noble Gases

Group 18
Unreactive, inert, “noble”, stable
Electron configuration
 ns2np6 full energy level
Have a 0 charge, no ions
Examples: He, Ne, Ar, Kr, etc
3. Metalloids

Sometimes called semiconductors
Form the “stairstep” between metals and
 nonmetals
Have properties of both metals and
 nonmetals
Examples: B, Si, Sb, Te, As, Ge, Po, At
B. Chemical Reactivity

     Alkali Metals
     Alkaline Earth Metals
     Transition Metals
     Halogens
1    Noble Gases
2
3
4
5
6
7
     C. Valence Electrons

    Valence Electrons

      e- in the outermost energy level

    Group #A = # of valence e- (except He)
    1A                                              8A
1        2A                        3A 4A 5A 6A 7A
2
3
4
5
6
7
    C. Valence Electrons
     Valence electrons =
        electrons in outermost energy level
     You can use the Periodic Table to
       determine the number of valence electrons
     Each group has the same number of
    1A
       valence electrons                      8A
1     2A                        3A 4A 5A 6A 7A
2
3
4
5
D. Lewis Diagrams

Also called electron dot diagrams
Dots represent the valence e-
Ex: Sodium                Ex: Chlorine




                             Lewis Diagram
                             for Oxygen

								
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