CH 9: Ionic and Covalent
Bonding
Renee Y. Becker
Valencia Community College
CHM 1045
1
Covalent Bonds
• Covalent bonds are formed by sharing at
least one pair of electrons.
• The attraction (nucleus/electrons) outweighs
the repulsions (electron/electron &
nucleus/nucleus)
2
Covalent Bonds
•Every
covalent bond has a characteristic
length that leads to maximum stability.
bond length
3
Strength of Covalent Bonds
•Energy
required to break a covalent bond in an
isolated gaseous molecule is called the bond
dissociation energy.
•Same
amount of energy released when the
bond forms
4
Example 1:
Which of the following is correct?
1. Energy is absorbed to form a bond
2. Energy is released when a bond is
formed
5
Polar Covalent Bonds
• Bond polarity is due to electronegativity
differences between atoms.
• Pauling Electronegativity: is expressed on a
scale where F = 4.0
6
Pauling Electronegativities
7
8
Electron-Dot Structures
• Using electron-dot (Lewis) structures, the
valence electrons in an element are
represented by dots.
.
• Lewis symbols
• Valence electrons are those electrons with the
highest principal quantum number (n).
9
10
Electron-Dot Structures
• The electron-dot structures provide a
simple, but useful, way of representing
chemical reactions.
• Ionic:
• Covalent:
11
Electron-Dot Structures
• Single Bonds:
C
H
H
H
H
• Double Bonds:
H
C
H
H
H
H
H
C
C
C
C
H
H
• Triple Bonds:
C
C
H
C
C
H
12
Drawing Lewis-Dot Structures
Rule 1: Count the total valence electrons.
Rule 2: Draw the structure using single bonds.
Rule 3: Distribute the remaining electron pairs
around the peripheral atoms.
Rule 4: Put remaining pairs on central atom.
Rule 5: Share lone pairs between bonded
atoms to create multiple bonds.
13
Drawing Lewis-Dot Structures
• NH2F Amino Fluoride: In this
molecule, nitrogen is the central atom.
• Rule 1: Number of electrons = 5 + (2 x 1) + 7
= 14 = 7 pairs
H N H
H N H
H N H
F
F
F
Rule 2
Rule 3
Rule 4
14
Drawing Lewis-Dot Structures
15
Drawing Lewis-Dot Structures
• Polyatomic molecules with central atoms
below the second row ten:
• In this compound there are 10 valence
electrons on bromine; this is called an
expanded octet. The extra pairs go into
unfilled d orbitals.
16
Example 2: Drawing Lewis-Dot Structures
•Draw electron-dot structures for:
C 3H 8
H 2 O2
CO2
N 2H 4
CH5N
C 2H 4
C 2H 2
Cl2CO
H3S+
HCO3–
17
Resonance Structures
• How is the double bond formed in O3?
Move lone pair from
this oxygen?
O
O
O
O
O
or
O
Or from this
oxygen?
O
O
O
• The correct answer is that both are correct,
but neither is correct by itself.
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Resonance Structures
• When multiple structures can be drawn, the actual
structure is an average of all possibilities.
• The average is called a resonance hybrid. A straight
double-headed arrow indicates resonance.
O
O
O
O
O
O
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Resonance Structures
• The nitrate ion, NO3–, has three
equivalent oxygen atoms, and its
electronic structure is a resonance
hybrid of three electron-dot structures.
Draw them.
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Formal Charge
• Formal Charge: Determines the best
resonance structure.
• We determine formal charge and estimate the
more accurate representation.



 # of bonding e  

Formal Charge = # of Valence e -  
 # of nonbonding e- 
 

 

2


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Example 3: Formal Charge
• Calculate the formal charge and
determine the most favorable of the
following electron dot structures:
SO2
NO3–
NCO–
N2O O3
CO32–
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Example 4:
What is the overall formal charge of the
following structure?
1.
2.
3.
4.
-2
-3
-1
0
O
O P O
O
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Example 5: Ionic Radii of Ions
• Compare ionic radii
– Fe & Fe3+
– Cl & Cl-
24