Chapter 1: Covalent Bonding and Shapes of Molecules
Aspartame (Nutrasweet®)
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I. Lewis Structures
A. Lewis symbols of elements
B. Ionic, covalent, and polar bonds
C. Lewis structures
D. Formal charge
E. Resonance structures
II. Molecular Shapes
A. VSEPR theory
B. Polarity of molecules
III. Valence Bond Model
A. Atomic and molecular orbitals
B. Hybrid atomic orbitals
IV. Functional Groups
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I. Lewis Structures
A. Lewis symbols of elements
Periodic Table
Review: periods; principle quantum numbers
s-block, p-block, d-block; groups
s and p atomic orbitals
rules for filling atomic orbitals
core and valence electrons
Electron Configuration
Lewis Symbol
H
C
N
O
F
Cl
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I. Lewis Structures
B. Ionic, covalent, and polar bonds
ionic bonds: transfer of electrons
Na + Cl
Na
+ Cl
covalent bonds: sharing of electrons
H + Cl
H Cl
electronegativity, : relative attraction for electrons in a bond
- increases going up and to the right in the periodic table
- Pauling electronegativity scale (arbitrary): Table 1.5
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I. Lewis Structures
B. Ionic, covalent, and polar bonds
H—H
 = 2.1 2.1
Cl—Cl
 = 3.0 3.0
D = 0  equal sharing of electrons
= nonpolar covalent bond
d+ d–
H—Cl
 = 2.1 3.0
D = 0.9  unequal sharing of electrons
= polar covalent bond
Na+Cl–
 = 0.9 3.0
D = 2.1  transfer of electrons
= ionic bond
generally:
when D < 1.9  covalent
> 1.9  ionic
nonmetal
+
nonmetal
metal
+
nonmetal
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I. Lewis Structures
C. Lewis structures
1. Count all the valence electrons; add one for each – charge
subtract one for each + charge
2. Draw single bonds between the atoms (the connectivity of the
atoms is determined experimentally and is usually given in the
problem).
3. Using the remaining valence electrons, place octets on all
atoms (exception H), in order of decreasing electronegativity.
4. If atoms do not have octets, use lone pair electrons on adjacent
atoms to form double or triple bonds to complete the octets.
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I. Lewis Structures
C. Lewis structures
CCl4
CH2O
C2H2
CH3OH
CH3CHCH2
HCN
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I. Lewis Structures
D. Formal charge
(Use the silly, complex formula in the textbook, or use this easier method:)
1. Divide the electrons in each bond equally between the two atoms sharing
them.
2. Count the number of electrons each atom now has and compare this
number to its normal valence.
• more electrons than normal valence  negative formal charge
• fewer electrons than normal valence  positive formal charge
H3 O+
CH3O–
CH3+
CO
N3–
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I. Lewis Structures
D. Formal charge
When two or more nonequivalent Lewis structures are possible, the
better (more stable) one is the one with:
1. fewer formal charges
2. more octets
3. a – charge on a more electronegative atom,
or a + charge on a more electropositive atom
COCl2
BF3
(CH3)2SO
In decreasing
order of
importance
HOCN
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I. Lewis Structures
E. Resonance structures
-two or more equivalent Lewis structures
-nuclei remain in fixed positions, but electrons arranged differently
delocalized
electrons
HCO2-
O
H C
O
H C
O
= H C
O
•neither of these accurately
describes the formate ion
•actual species is an average of
the two (resonance hybrid)
O½ O ½-
1.5 bond order
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I. Lewis Structures
E. Resonance structures
O
CH3 C
O
CH3 C
NH2
more stable
major contributor
NH2
dO
= CH3 C d+
NH2
less stable
minor contributor
Draw resonance structures for the following species. If the structures
are not equivalent, indicate which would be the major contributor.
CH3NO2
CH2CHO–
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II. Molecular Shapes
A. VSEPR theory
formula
Lewis structure
e– pair
geometry
molecular
geometry (angle)
C2H2
CH2O
CH4
HCN
NH3
H2O
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II. Molecular Shapes
B. Polarity of molecules
If the individual dipole moments in a molecule do not exactly cancel, then
the molecule will have a net dipole moment and be a polar molecule.
CCl4
CHCl3
CH2O
CO2
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III. Valence Bond Model
A. Atomic and molecular orbitals
H• + H•  H—H
Two electrons in s1s are lower energy than in the separate atoms
 covalent bond
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III. Valence Bond Model
B. Hybrid atomic orbitals
1. sp3 hybridization
H
CH4 facts:
C H
H
H
C
2s
2p
tetrahedral,
4 equivalent bonds
promote
electron
hybridize
2s
2p
H
4H
sp3 hybrid a.o.s:
C(sp3)
tetrahedral
C H
H
H
sp3
hybrid
a.o.s
s(sp3C + 1sH)
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III. Valence Bond Model
B. Hybrid atomic orbitals
1. sp3 hybridization
C
N
lone pairs
in sp3 a.o.s
H
H
H C C H
H
H
s(sp3C + sp3C)
O
H
H C N H
H
H
s(sp3C + sp3N)
H
H C O H
H
s(sp3C + sp3O)
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III. Valence Bond Model
B. Hybrid atomic orbitals
trigonal planar = sp2
2. sp2 hybridization
C2H4 facts: H
C
H
1s
H
1s
C
2p
sp
2
C
H
C
H
H
H
H
1s
1s
2p
2
sp
 bond
2p
H
all six atoms lie
in same plane
H
C C
H
H
overlap
p orbitals
s(sp2C + 1sH)
s(sp2C + sp2C)
H
H
C C
H
H
H
all atoms coplanar
for p orbital overlap
H
C
=
H
C
H
double bond =
1  bond +
1 s bond
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III. Valence Bond Model
B. Hybrid atomic orbitals
2. sp2 hybridization
H
H
H
C C
H
C O
H
H
C
H
H
C C
O
H
H
H
H
s(sp2C + sp2C) + 
C O
lone pairs
in sp2 a.o.s
s(sp2C + sp2O) + 
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III. Valence Bond Model
B. Hybrid atomic orbitals
3. sp hybridization
C2H2 facts: H
H
C C H
linear = sp
H
1s
C
sp
2p
C
sp
2p
2  bonds
2p
H
1s
C C
H
s(spC + 1sH)
s(spC + spC)
H
C C
H = H
C
C
H
triple bond =
2  bonds +
1 s bond
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III. Valence Bond Model
B. Hybrid atomic orbitals
3. sp hybridization
H C N
N
lone pair
in sp a.o.
H C N
s(spC + spN) + 2
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III. Valence Bond Model
B. Hybrid atomic orbitals
What is the hybridization of each indicated atom in the following compound?
CH3 OH
C
CH
HO
17-ethynylestradiol
(“The Pill”)
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IV. Functional Groups
Atoms or groups of atoms that behave similarly, regardless of the
structure to which they are attached.
CH3OH
CH3CH2OH
methanol
ethanol
OH
OH
b-phenethyl alcohol (lilacs)
geraniol
OH
retinol (vit A)
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IV. Functional Groups
Hydrocarbons (C & H only)
aliphatic
alkanes
C
aromatic
Heteroatomic compounds
alcohols R OH
ethers
C
alkenes
O
aldehydes
ketones
C C
alkynes
C C
cyclic compounds
R O R'
R C H
O
R C
R'
carboxylic acids
esters
amines
amides
O
R C OH
O
R C OR'
R NH2
O
R C NH2
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