Chemical
Bonding
I
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
COMPOUND
2 elements
Binary
Compound
NaCl
more than 2
elements
Ternary
Compound
NaNO3
ION
1 atom
2 or more atoms
Monatomic
Ion
Polyatomic
Ion
+
Na
NO3
-
TYPES OF BONDS
IONIC
COVALENT
Bond
Formation
e- are transferred from
metal to nonmetal
e- are shared between
two nonmetals
Type of
Structure
crystal lattice
true molecules
Physical
State
solid
liquid or gas
Melting
Point
high
low
Solubility in
Water
yes
usually not
Electrical
Conductivity
yes
(solution or liquid)
no
Other
Properties
odorous
TYPES OF BONDS
METALLIC
Bond
Formation
e- are delocalized
among metal atoms
Type of
Structure
“electron sea”
Physical
State
solid
Melting
Point
very high
Solubility in
Water
no
Electrical
Conductivity
yes
(any form)
Other
Properties
malleable, ductile,
lustrous
IONIC BONDS
IONIC BONDING - CRYSTAL LATTICE
Covalent Bonding - True Molecules
Diatomic
Molecule
METALLIC BONDING - “ELECTRON SEA”
BOND POLARITY
 Most bonds are a blend of ionic
and covalent characteristics.
 Difference in electronegativity
determines bond type.
BOND POLARITY
Electronegativity
 Attraction an atom has for a shared pair of electrons.
 higher e-neg atom   lower e-neg atom +
BOND POLARITY
Electronegativity Trend (p. 151)
Increases up and to the right.
BOND POLARITY
Nonpolar Covalent Bond
e- are shared equally
symmetrical e- density
usually identical atoms
Polar Covalent Bond
 e- are shared unequally
 asymmetrical e- density
 results in partial charges (dipole)
+


Nonpolar
Polar
Ionic
BOND POLARITY
Examples:
Cl2
HCl
NaCl
3.0-3.0=0.0
Nonpolar
3.0-2.1=0.9
Polar
3.0-0.9=2.1
Ionic
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
LEWIS DIAGRAMS
Molecular Structure
I
RULE
 Remember…
 Most atoms form bonds in order to have 8 valence electrons.
A. OCTET RULE
 Exceptions:
F
F
 Hydrogen  2 valence e
F
B
F
 Groups F
1,2,3 get
2,4,6
valence
e
SO
F
H
O
H
N
 Expanded octet  more than 8
valence
e (e.g.
S, P, Xe)
Very
unstable!!
F F F
-
-
-
 Radicals  odd # of valence e-
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 aren’t enough e- to go around, form double
or triple bonds.
B. DRAWING LEWIS DIAGRAMS
CF4
1 C × 4e- = 4e-
4 F × 7e- = 28e-
F
32e
- 8e
24e-
F C F
F
B. DRAWING LEWIS DIAGRAMS
BeCl2
1 Be × 2e- = 2e-
2 Cl × 7e- = 14e16e
- 4e
12e-
Cl Be Cl
B. DRAWING LEWIS DIAGRAMS
CO2
1 C × 4e- = 4e-
2 O × 6e- = 12e16e
- 4e
12e-
O C O
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.
C. POLYATOMIC IONS
ClO4-
1 Cl × 7e- = 7e-
4 O × 6e- = 24e31e
+ 1e
32e-- 8e
24e-
O
O Cl O
O
C. POLYATOMIC IONS
NH4+
1 N × 5e- = 5e4 H × 1e- = 4e9e- 1e8e- 8e0e-
H
H N H
H
C. POLYATOMIC IONS
OH-
1 O × 6e- = 6e1 H × 1e- = 1e7e+ 1e8e- 8e0e-
O H
D. RESONANCE STRUCTURES
Molecules that can’t 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.
D. RESONANCE STRUCTURES
 SO3
O
O
O S O
O S O
O
O S O
MOLECULAR
GEOMETRY
I
VSEPR THEORY
Valence Shell Electron Pair Repulsion
Theory
Electron pairs orient themselves in order to
minimize repulsive forces.
VSEPR THEORY
Types of e- Pairs
Bonding pairs - form bonds
Lone pairs - nonbonding e-
Lone pairs repel
more strongly than
bonding pairs!!!
VSEPR THEORY
 Lone pairs reduce the bond angle between atoms.
Bond Angle
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.
Know the 8 common shapes
& their bond angles!
COMMON MOLECULAR SHAPES
2 total
2 bond
0 lone
LINEAR
BeH2
180°
COMMON MOLECULAR SHAPES
3 total
3 bond
0 lone
BF3
TRIGONAL PLANAR
120°
COMMON MOLECULAR SHAPES
3 total
2 bond
1 lone
SO2
BENT
<120°
COMMON MOLECULAR SHAPES
4 total
4 bond
0 lone
CH4
TETRAHEDRAL
109.5°
COMMON MOLECULAR SHAPES
4 total
3 bond
1 lone
NH3
TRIGONAL PYRAMIDAL
107°
COMMON MOLECULAR SHAPES
4 total
2 bond
2 lone
H2O
BENT
104.5°
COMMON MOLECULAR SHAPES
5 total
5 bond
0 lone
TRIGONAL
PCl5
BIPYRAMIDAL
120°/90°
COMMON MOLECULAR SHAPES
6 total
6 bond
0 lone
SF6
OCTAHEDRAL
90°
EXAMPLES
 PF3
F P F
4 total
3 bond
1 lone
F
TRIGONAL
PYRAMIDAL
107°
EXAMPLES
 CO2
2 total
2 bond
0 lone
O C O
LINEAR
180°