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

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Ch 6 the periodic table periodic law l.jpg

Ch. 6 - The Periodic Table & Periodic Law

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


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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

  • Dmitri Mendeleev (1869, Russian)

    • Predicted properties of undiscovered elements


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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!


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C. Modern Periodic Table

  • Group (Family)

  • Period


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1. Groups/Families

  • Vertical columns of periodic table

  • Numbered 1 to 18 from left to right

  • Each group contains elements with similar chemical properties


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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


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3. Blocks

  • Main Group Elements

  • Transition Metals

  • Inner Transition Metals


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Lanthanides - part of period 6

Actinides - part of period 7

3. Blocks

Overall Configuration


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Ch. 6 - The Periodic Table

II. Classification of theElements(pages 182-186)


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A. Metallic Character

  • Metals

  • Nonmetals

  • Metalloids


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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


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a. Alkali Metals

  • Group 1

  • 1 Valence electron

  • Very reactive

  • Electron configuration

    • ns1

  • Form 1+ ions

  • Cations

    • Examples: Li, Na, K


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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


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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


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2. Nonmetals

  • Not good conductors

  • Found on right side of periodic table – AND hydrogen

  • Usually brittle solids or gases


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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


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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


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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


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B. Chemical Reactivity

  • Alkali Metals

  • Alkaline Earth Metals

  • Transition Metals

  • Halogens

  • Noble Gases


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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)


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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


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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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