Chemical%20Bonding

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• Chemical Bonding

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• Chemical Bond
• attractive force between atoms or ions that binds
  them together as a unit
• bonds form in order to
• decrease potential energy (PE)
• increase stability

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COMPOUND
more than 2 elements
2 elements
Ternary Compound
Binary Compound
NaNO3
NaCl
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ION
1 atom
2 or more atoms
Polyatomic Ion
Monatomic Ion
NO3-
Na
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Types of Bonds
COVALENT
IONIC
e- are transferred from metal to nonmetal
e- are shared between two nonmetals
Bond Formation
Type of Structure
true molecules
crystal lattice
Physical State
liquid or gas
solid
Melting Point
low
high
Solubility in Water
yes
usually not
yes (solution or liquid)
Electrical Conductivity
no
Other Properties
odorous
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Types of Bonds
METALLIC
e- are delocalized among metal atoms
Bond Formation
Type of Structure
electron sea
Physical State
solid
Melting Point
very high
Solubility in Water
no
yes (any form)
Electrical Conductivity
malleable, ductile, lustrous
Other Properties
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Ionic Bonds
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Ionic Bonding - Crystal Lattice
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Covalent Bonding - True Molecules
Diatomic Molecule
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Metallic Bonding - Electron Sea
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Bond Polarity

• Most bonds are a blend of ionic and covalent
  characteristics.
• Difference in electronegativity determines bond
  type.

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Bond Polarity

• Electronegativity
• Attraction an atom has for a shared pair of
  electrons.
• higher e-neg atom ? ?-
• lower e-neg atom? ?

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Bond Polarity

• Electronegativity Trend (p. 151)
• Increases up and to the right.

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Bond Polarity

• Nonpolar Covalent Bond
• e- are shared equally
• symmetrical e- density
• usually identical atoms

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• Polar Covalent Bond
• e- are shared unequally
• asymmetrical e- density
• results in partial charges (dipole)

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• Nonpolar
• Polar
• Ionic

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Bond Polarity

• Examples
• Cl2
• HCl
• NaCl

3.0-3.00.0 Nonpolar 3.0-2.10.9 Polar 3.0-0.92.1
Ionic
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• Chemical Bond
• attractive force between atoms or ions that binds
  them together as a unit
• bonds form in order to
• decrease potential energy (PE)
• increase stability

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Lewis Diagrams

• Molecular Structure

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Rule

• Remember
• Most atoms form bonds in order to have 8 valence
  electrons.

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A. Octet Rule

• Exceptions

• Hydrogen ? 2 valence e-

• Groups 1,2,3 get 2,4,6 valence e-

• Expanded octet ? more than 8 valence e- (e.g. S,
  P, Xe)

• Radicals ? odd of valence e-

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B. Drawing Lewis Diagrams

• Find total of valence e-.
• Arrange atoms - singular atom is usually in the
  middle.
• Form bonds between atoms (2 e-).
• Distribute remaining e- to give each atom an
  octet (recall exceptions).
• If there arent enough e- to go around, form
  double or triple bonds.

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B. Drawing Lewis Diagrams

• CF4

1 C 4e- 4e- 4 F 7e- 28e- 32e-
F F C F F
- 8e- 24e-
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B. Drawing Lewis Diagrams

• BeCl2

1 Be 2e- 2e- 2 Cl 7e- 14e- 16e-

• Cl Be Cl

- 4e- 12e-
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B. Drawing Lewis Diagrams

• CO2

1 C 4e- 4e- 2 O 6e- 12e- 16e-

• O C O

- 4e- 12e-
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C. Polyatomic Ions

• To find total of valence e-
• Add 1e- for each negative charge.
• Subtract 1e- for each positive charge.
• Place brackets around the ion and label the
  charge.

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C. Polyatomic Ions

• ClO4-

1 Cl 7e- 7e- 4 O 6e- 24e- 31e-
O O Cl O O
1e- 32e-
- 8e- 24e-
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C. Polyatomic Ions

• NH4

1 N 5e- 5e- 4 H 1e- 4e- 9e-
H H N H H
- 1e- 8e-
- 8e- 0e-
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D. Resonance Structures

• Molecules that cant be correctly represented by
  a single Lewis diagram.
• Actual structure is an average of all the
  possibilities.
• Show possible structures separated by a
  double-headed arrow.

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D. Resonance Structures

• SO3

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Molecular Geometry
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VSEPR Theory

• Valence Shell Electron Pair Repulsion Theory
• Electron pairs orient themselves in order to
  minimize repulsive forces.

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VSEPR Theory

• Types of e- Pairs
• Bonding pairs - form bonds
• Lone pairs - nonbonding e-

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VSEPR Theory

• Lone pairs reduce the bond angle between atoms.

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Determining Molecular Shape

• Draw the Lewis Diagram.
• Tally up e- pairs on central atom.
• double/triple bonds ONE pair
• Shape is determined by the of bonding pairs and
  lone pairs.

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Common Molecular Shapes

• 2 total
• 2 bond
• 0 lone

LINEAR 180
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Common Molecular Shapes

• 3 total
• 3 bond
• 0 lone

TRIGONAL PLANAR 120
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Common Molecular Shapes

• 3 total
• 2 bond
• 1 lone

BENT lt120
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Common Molecular Shapes

• 4 total
• 4 bond
• 0 lone

TETRAHEDRAL 109.5
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Common Molecular Shapes

• 4 total
• 3 bond
• 1 lone

TRIGONAL PYRAMIDAL 107
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Common Molecular Shapes

• 4 total
• 2 bond
• 2 lone

BENT 104.5
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Common Molecular Shapes

• 5 total
• 5 bond
• 0 lone

TRIGONAL BIPYRAMIDAL 120/90
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Common Molecular Shapes

• 6 total
• 6 bond
• 0 lone

OCTAHEDRAL 90
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Examples

• PF3

4 total 3 bond 1 lone
TRIGONAL PYRAMIDAL 107
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Examples

• CO2

2 total 2 bond 0 lone
LINEAR 180
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