Unit 3

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I. History

• Unit 3 Modern Atomic Theory the Periodic
  Table
• The Periodic Table

2
Mendeleev

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

3
Dmitri Mendeléev
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Mendeleev

• Dmitri Mendeleev (1869, Russian)
• Predicted properties of undiscovered elements.

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Mendeleevs Periodic Table
Period 1 Group I II III IV V VI VII VIII
Period 1 H 1
2 Li 7 Be 9.4 B 11 C 12 N 14 O 16 F 19 F 19
3 Na 23 Mg 24 Al 27.3 Si 28 P 31 S 32 C 35.5
4 K 39 Ca 40 ? 44 Ti 48 V 51 Cr 52 Mn 55 Fe 56, Co 59, Ni 59
5 Cu 63 Zn 65 ? 68 ? 72 As 75 Se 78 Br 80
6 Rb 85 Sr 87 ? Yt 88 Zr 90 Nb 94 Mo 96 ? 100 Ru 104, Rh 104, Pd 106
7 Ag 108 Cd 112 In 113 Sn 118 Sb 122 Te 125 J 127
8 Cs 133 Ba 137 ?Di 138 ?Ce 140
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10 ?Er 178 ?La 180 Ta 182 W 184 Os 195, Ir 197, Pt 198
11 Au 199 Hg 200 Tl 204 Pb 207 Bi 208
12 Th 231 U 240

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Moseley

• Henry Moseley (1913, British)
• Organized elements by increasing atomic number.
• Resolved discrepancies in Mendeleevs arrangement.

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II. Organization of theElements

• Unit 3 Modern Atomic Theory the Periodic
  Table
• The Periodic Table

8
Periods

• Elements are arranged in seven horizontal rows,
  in order of increasing atomic number from left to
  right and from top to bottom.
• Rows are called periods and are numbered from 1
  to 7.
• Represent principal energy levels (n)

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Groups

• Elements with similar chemical properties form
  vertical columns, called groups (families).
• Groups 1A - 8A are the representative elements.
• s and p-blocks
• The B groups are the transition elements.

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Inner Transition Elements

• The two rows of 14 elements at the bottom of the
  periodic table are the inner transition elements.
• Lanthanide Series - the 4f row that includes 57
  (Lanthanum) - 71 Lu
• Actinide Series - the 5f row that includes 89 Ac
  (Actinium) - 102 No

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Metals, Nonmetals, Metalloids
1
Nonmetals
2
3
4
Metals
5
6
7
Metalloids
Zumdahl, Zumdahl, DeCoste, World of Chemistry
2002, page 349
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Metals

• To the left of the staircase (or below)
• Most are solids except for Hg which is a liquid
• Malleable, lustrous, ductile, good conductors of
  heat electricity

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Nonmetals

• To the right of the staircase (or above)
• Gases or brittle solids at room temperature
• Poor conductors of heat electricity (insulators)

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Metalloids

• Semi-metals
• These elements border the staircase and have
  properties of both metals and nonmetals.
• Include the following elements B, Si, Ge, As,
  Sb, Te, and At
• Dull, brittle, semi-conductors (used in computer
  chips)

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Group Names Properties

• Alkali Metals
• Alkaline Earth Metals
• Halogens
• Noble Gases

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Alkali Metals (Group 1A)

• very reactive
• good conductors
• end in s1
• need to lose 1e- to have noble gas configuration
• reactivity increases as you go down the group

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Alkaline Earth Metals (Group 2A)

• less reactive than alkali metals
• end in s2
• need to lose 2e- to have noble gas configuration

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Halogens (Group 7A)

• combine easily with alkali metals
• exist as diatomic molecules
• F2, Cl2, etc.
• electron configuration ends in p5
• need to gain 1e- to achieve noble gas
  configuration

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Noble Gases (Group 8A)

• full outer energy level
• electron configuration ends in p6
• do not form chemical compounds easily
• also called inert gases

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III. Periodic Trends

• Unit 3 Modern Atomic Theory the Periodic
  Table
• The Periodic Table

21
Periodic Law

• When elements are arranged in order of increasing
  atomic , elements with similar properties appear
  at regular intervals.

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Periodic Properties

• Atomic Radius
• Size of atom

• Ionization Energy
• Energy required to remove an e- from a neutral
  atom.

• Electronegativity
• Ability of an atom to attract electrons to itself
  in a chemical bond

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Atomic Radius

• Atomic Radius

• Increases to the LEFT and DOWN

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Atomic Radius

• Why larger going down?
• Higher energy levels have larger orbitals
• Shielding - core e- block the attraction between
  the nucleus and the valence e-
• Why smaller to the right?
• Increased nuclear charge without additional
  shielding pulls e- in tighter

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Ionic Radius

• Ionic Radius

• Cations ()
• lose e-
• smaller

• Anions ()
• gain e-
• larger

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Ionization Energy

• Ionization Energy

• Increases UP and to the RIGHT

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Ionization Energy

• Why opposite of atomic radius?
• In small atoms, e- are close to the nucleus where
  the attraction is stronger

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Electronegativity

• Electronegativity

• Increases UP and to the RIGHT

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Electronegativity

• Why larger going across?
• Stronger tendency to attract electrons and form
  negative ions
• Why smaller going down a group?
• Shielding - core e- block the attraction between
  the nucleus and the valence e-

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Examples

• Which atom has the larger radius?

• Be or Ba
• Ca or Br

Ba Ca
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Examples

• Which atom has the higher I.E.?

• N or Bi
• Ba or Ne

N Ne
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Examples

• Which atom has higher electronegativity?

• Li or C
• Cr or Br

C Br
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Examples

• Which particle has the larger radius?

• S or S2-
• Al or Al3

S2- Al
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The END of Unit 3!