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

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Ch. 6 - The Periodic Table & Periodic Law. I. Development of the Modern Periodic Table (p. 174 - 181). A. Mendeleev. Dmitri Mendeleev (1869, Russian) Organized elements by increasing atomic mass Elements with similar properties were grouped together There were some discrepancies.

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a mendeleev
A. Mendeleev
  • Dmitri Mendeleev (1869, Russian)
    • Organized elements by increasing atomic mass
    • Elements with similar properties were grouped together
    • There were some discrepancies
a mendeleev3
A. Mendeleev
  • Dmitri Mendeleev (1869, Russian)
    • Predicted properties of undiscovered elements
b moseley
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
C. Modern Periodic Table
  • Group (Family)
  • Period
1 groups families
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
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
3. Blocks
  • Main Group Elements
  • Transition Metals
  • Inner Transition Metals
3 blocks9

Lanthanides - part of period 6

Actinides - part of period 7

3. Blocks

Overall Configuration

a metallic character
A. Metallic Character
  • Metals
  • Nonmetals
  • Metalloids
1 metals
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
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
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
c. Transition Metals
  • Groups 3 - 12
  • Reactive (not as reactive as Groups 1 or 2), can be free elements
  • Electron Configuration
    • ns2(n-1)dxwhere x is column in d-block
  • Form variable valence state ions
  • Cations
    • Examples: Co, Fe, Pt, etc
2 nonmetals
2. Nonmetals
  • Not good conductors
  • Found on right side of periodic table – AND hydrogen
  • Usually brittle solids or gases
a halogens
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
b. Noble Gases
  • Group 18
  • Unreactive, inert, “noble”, stable
  • Electron configuration
    • ns2np6full energy level
  • Have a 0 charge, no ions
  • Examples: He, Ne, Ar, Kr, etc
3 metalloids
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
B. Chemical Reactivity
  • Alkali Metals
  • Alkaline Earth Metals
  • Transition Metals
  • Halogens
  • Noble Gases
c valence electrons

1A

8A

2A

3A 4A 5A 6A 7A

C. Valence Electrons
  • Valence Electrons
    • e- in the outermost energy level
  • Group #A = # of valence e- (except He)
c valence electrons22

1A

8A

2A

3A 4A 5A 6A 7A

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 valence electrons
d lewis diagrams
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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