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)
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Covalent Bonds
•Every covalent bond has a characteristic
length that leads to maximum stability.
bond length3
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
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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
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Polar Covalent Bonds
• Bond polarity is due to electronegativity differences between atoms.
• Pauling Electronegativity: is expressed on a scale where F = 4.0
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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
.
Electron-Dot Structures
• The electron-dot structures provide a simple, but useful, way of representing chemical reactions.
• Ionic:
• Covalent:11
Electron-Dot StructuresElectron-Dot Structures
• Single Bonds:
• Double Bonds:
• Triple Bonds:
C
HHHH
C HH
H
H
C
H
HH
H
C C C
C HC C CH
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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.
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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
NH H
F
NH H
F
NH H
F
Rule 2 Rule 3 Rule 414
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:
C3H8 H2O2 CO2 N2H4
CH5N C2H4 C2H2 Cl2CO
H3S+ HCO3–
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Resonance Structures
• How is the double bond formed in O3?
• The correct answer is that both are correct, but neither is correct by itself.
O O O
O O O
O O O
or
Or from this oxygen?
Move lone pair from this oxygen?
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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.
Formal Charge= # of Valence e-
# of bonding e-
2
# of nonbonding e-
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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 CO3
2–
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Example 4:
What is the overall formal charge of the following structure?
1. -2
2. -3
3. -1
4. 0
P
O
O
OO
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