Chemical Bonding

1
Chemical Bonding
2
Chemical Bonds

• Compound are formed from chemically bound atoms
  or ions
• Bonding only involves the valence electrons

3
Chemical Bonds

• Defn force holding two atoms together
• How are they formed?
• Atoms gain, lose, or share valence electrons
• Why does bonding occur?
• Stability achieve octet rule

4
Electron Dot Structure

• Shows valence electrons around atomic symbol

hydrogen
(group 1)
H



N


nitrogen
(group 5)




(group 7)
Cl

chlorine



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Types of Chemical Bonds

• 3 Types
• covalent bond
• ionic bond
• metallic bond

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Covalent Bond

• Defn two atoms share one pair of electrons

Electrons shared

A

B

A
B

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Covalent Bonds

• Where are these bonds found?
• - molecules (molecular compounds)
• - polyatomic ions

8
Ionic Bond

• Defn force holding cations and anions together

A

B

A
B-

Ionic bond
9
Ionic Bond

• Where are these bonds found?
• Ionic Compounds

10
Metallic Bonding

• Defn attraction of metallic cations

Occurs only in metals
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Covalent Bonding

• Whats going on?
• Molecule formed when 2 or more atoms bond
  covalently

Sharing of electrons
12
Two Types of Covalent Bonds

• i) nonpolar covalent equal sharing of e-
• ii) polar covalent UNequal sharing of e-

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Nonpolar vs. Polar
NONPOLAR
POLAR
14
Nonpolar vs. Polar
15
Nonpolar vs. Polar
16
Single Bond

• Defn one pair (2) of e- shared
• Lewis Structures represents how atoms in
  molecules are arranged
• atoms MUST obey octet rule (except hydrogen)

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Lewis Structures

• bonded electrons occur between bonded atoms

A
B
A
B
or

single bond
18
Lewis Structures

• Unshared or Lone Pairs electron pairs NOT
  involved in bonding

A
B
A
B





lone pairs
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Lewis Structures Examples

• H2O

(8 valence e- or 4 pairs)










O
H
H
H
O






H



O
H

H
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Lewis Structures Examples

• NHF2

(20 v.e. or 10 pairs)














N
F
F
F


















H
N
H





F











N
F
F






H
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Multiple Covalent Bonds

• Double Bond two pairs (4) e- shared

A
B
A
B


O2
(12 v.e. 6 pairs)














O









O
O
O
O
O



















O
O


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Multiple Covalent Bond

• Triple Bond three pairs (6) e- shared

A
B
A
B



N2
(10 v.e. 5 pairs)



















N


N


N
N
N
N














N
N



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Comparing single, double, and triple bonds

• Bond Strength
• Bond Length

Triple gt Double gt Single
Single gt Double gt Triple
The shorter the bond, the stronger it is
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Polyatomic Ions

• Defn CHARGED group of atoms covalently bonded
• - ex SO42-, NH41, NO31-

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Polyatomic Ions
SO42-
(32 v.e. 16 pairs)
2-


2-




O




O


















O
O
S
O
O
S


















O




O






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Polyatomic Ions
NH41
(8 v.e. 4 pairs)
1
1
H
H




H
H
N
H
H
N




H
H
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Ionic Bonding
giving/taking of valence electrons

• Whats going on?
• If I gave you a compound, how can you tell if it
  is ionic or not?
• combo of metal nonmetal

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Formation of Ionic Bonds

• NaCl

1-






Na
Cl
Na1





Cl







2s22p63s1
3s23p5
2s22p6
3s23p6
8 v.e.
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Formation of Ionic Bonds

• CaBr2

1-







Ca
Br




Br







Ca2





Br
1-







Br




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Using electronegativity to determine bond type

• Recall electronegativity how much an atom wants
  electrons
• Each atom is assigned a number between 0-4.0 to
  determine electronegativity strength

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Using electronegativity to determine bond type

• We know 3 types of bonds
• - nonpolar covalent
• - polar covalent
• - ionic
• To determine bond type, subtract
  electronegativity values and see scale

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Scale
Using electronegativity to determine bond type
polar covalent
nonpolar covalent
ionic
1.7
0.3
0
4.0
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Using electronegativity to determine bond type
H and Cl
3.0 2.1
0.9
polar covalent
C and S
2.5 2.5
0
nonpolar covalent
Na and F
4.0 0.9
3.1
ionic
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Metallic Bonding

• Defn bond formed from attraction between
  positive nuclei and delocalized electrons
• holds metals together
• Delocalized Electrons electrons detached from
  parent atom
• lost electron away from home

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Electron Sea Model

• Defn electrons move freely within other
  molecular orbitals

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Properties of Metals

• Electron sea model gives metals certain physical
  properties
• Shiny due to photoelectric effect
• Conduct electricity and heat electrons move
  easily from one place to another
• Malleable (pound into sheets)
• Ductile (put into wires)

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Why malleable and ductile?
atoms can also move from one place to another and
still remain in contact with and bonded to the
other atoms and electrons around them
Shape 1
Shape 2
shifted atoms
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Dipole Moment

• defn imbalance of electron density in a
  covalent bond
• Due to electronegativity of atoms

?- (partial negative) signifies more EN atom
? (partial positive) signifies less EN atom
shows direction of dipole moment
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Examples
?
?-
?-
?
H
O
Cl
C
H 2.2 C 2.6 N 3.0 Cl 3.2 O 3.4 F 4.0
?
?-
?-
?
C
F
N
H
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Intermolecular Forces

• Defn attractive forces between 2 molecules

44
Intermolecular Forces

• Dipole-Dipole attraction between oppositely
  charged polar molecules

?-
?-
?
?
?-
?
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Intermolecular Forces

• London Dispersion Forces very weak, very brief
  dipole moment created in nonpolar molecules

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Electrons evenly distributed
London force
Temporary dipole
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Intermolecular Forces

• Hydrogen Bonding strong bond between H and
  N,O, or F of another molecule
• - Water is prime example

?-
O
O
H
H
?
?
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?-
O
O
O
O
H
H
H
?
H
?
?-
O
hydrogen bond
O
H
H
?
?
O
O
O
O
H
H
H
H
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Strength Ranking
Hydrogen gt dipole-dipole gt London
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VSEPR

• Valence Shell Electron Pair Repulsion
• Defn determines the shape of molecule
• Electron pairs try to stay far away as possible

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lone pairs
atoms bonded to central atom
shape
4
0
tetrahedral
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Tetrahedral
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lone pairs
atoms bonded to central atom
shape
4
0
tetrahedral
trigonal pyramidal
1
3
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Trigonal Pyramidal
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lone pairs
atoms bonded to central atom
shape
4
0
tetrahedral
trigonal pyramidal
1
3
2
2
bent
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Bent
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lone pairs
atoms bonded to central atom
shape
4
0
tetrahedral
trigonal pyramidal
1
3
2
2
bent
trigonal planar
3
0
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Trigonal Planar
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lone pairs
atoms bonded to central atom
shape
4
0
tetrahedral
trigonal pyramidal
1
3
2
2
bent
trigonal planar
3
0
2
0
linear
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Linear