Ch. 6 - The Periodic Table

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

• Ch. 6 - The Periodic Table Periodic Law

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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

• Deduced elements existed, but were undiscovered
  elements, their properties could be predicted

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B. Moseley

• Henry Moseley (1913, British)
• Organized elements by increasing atomic number
• Resolved discrepancies in Mendeleevs 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
• Each group contains elements with similar
  chemical physical properties (same amount of
  valence electrons in each column)
• 2 numbering systems exist
• Groups I through VIII with ea. followed by A
  or B
• A groups are Main Group Elements (sp electrons)
• B groups are Transition Elements (d electrons)
• Numbered 1 to 18 from left to right

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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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3. Blocks
Overall Configuration
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II. Classification of theElements(pages
182-186)

• Ch. 6 - The Periodic Table

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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
• Metallic properties increase as you go from left
  to right across a period

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a. Alkali Metals

• Group 1(IA)
• 1 Valence electron
• Very reactive, form metal oxides
  (ex Li2O)
• Electron configuration
• ns1
• Lowest melting points
• Form 1 ion Cations
• Examples Li, Na, K

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b. Alkaline Earth Metals

• Group 2 (IIA)
• 2 valence electrons
• Reactive (not as reactive as alkali metals) form
  metal oxides (ex MgO)
• Electron Configuration
• ns2
• Form 2 ions
• Cations
• Examples Be, Mg, Ca, etc

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c. Transition Metals

• Groups 3 12 (IB VIIIB)
• Reactive (not as reactive as Groups 1 or 2), can
  be free elements
• Highest melting points
• Electron Configuration
• ns2(n-1)dx where x is column in d-block
• Form variable valence state ions
• Always form Cations
• Examples Co, Fe, Pt, 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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2. Nonmetals

• Not good conductors
• Usually brittle solids or gases (1 liquid Br)
• Found on right side of periodic table AND
  hydrogen
• Hydrogen is its own group, reacts rapidly with
  oxygen other elements (has 1 valence electron)

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Nonmetal Groups/Families

• Boron Group IIIA typically 3 valence electrons,
  also mix of metalloids and metals
• Carbon Group IVA typically 4 valence electrons,
  also has metal and metalloids
• Nitrogen Group VA typically 5 valence electrons,
  also has metals metalloids
• Oxygen Group VIA typically 6 valence electrons,
  also contains metalloids

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a. Halogens

• Group 17 (VIIA)
• Very reactive
• Electron configuration
• ns2np5
• Form 1- ions 1 electron short of noble gas
  configuration
• Typically form salts (NaCl)
• Anions
• Examples F, Cl, Br, etc

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b. Noble Gases

• Group 18 (VIIIA)
• Unreactive, inert, noble, stable
• Electron configuration
• ns2np6 full energy level
• Have an octet or 8 valence e-
• Have a 0 charge, no ions
• Helium is stable with 1s2, a duet
• Examples He, Ne, Ar, Kr, etc

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B. Chemical Reactivity

• Metals
• Period - reactivity decreases as you go from left
  to right across a period. Group - reactivity
  increases as you go down a group
• Non-metals
• Period - reactivity increases as you go from the
  left to the right across a period. Group -
  reactivity decreases as you go down the group.

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C. Valence Electrons

• Valence Electrons
• e- in the outermost s p energy levels
• Stable octet filled s p orbitals (8e-) in one
  energy level

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C. Valence Electrons

• You can use the Periodic Table to determine the
  number of valence electrons
• Each group has the same number of valence
  electrons
• Group A of valence e- (except He)

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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