Covalent Bonding
Attraction
Chapter 8: Covalent Bonding
Stable bond
Repulsion
•  Number of bonds = Number shared e- pairs.
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Covalent Bonding
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Single Covalent Bonds
Lewis structures:
dot = 1 e-. Line = 1 pair of e-
H−H
Potential Energy (KJ/mol)
H H
Single bond: one shared pair of e-.
Distance between nuclei (pm)
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Single Covalent Bonds
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Single Covalent Bonds
F
bonding
2F = 2(7) = 14 valence e- .
Share 2 e- to form octets.
H O H
H O H
H N H
H
bonding
lone pair
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F
O + 2H = 6 + 2(1) = 8 val. e-.
Two O-H bonds.
lone pair
H O H
4
5
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N + 3H = 5 + 3(1) = 8 val. e-.
3 N-H bonds.
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Single Covalent Bonds
Group
4A
# of
valence e-
# of e- shared
to form an octet
(8 - A group#)
4
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Guidelines for Writing Lewis Structures
1.  Count the valence e- in the molecule.
Example
2. Draw a skeleton structure. Join atoms with single
lines (pairs of e-).
3. Add e- pairs to form octets (except H). Start with
terminal atoms.
H
C in CH4
H–C–H
.. .. ..
..F – N – ..F
..F
H
3
N in NF3
6A
6
2
O in H2O
H – ..
O–H
7A
7
1
F in HF
H – ..
F
..
5
..
5A
..
..
..
..
..
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Guidelines for Writing Lewis Structures
Guidelines for Writing Lewis Structures
Phosphorus trifluoride, PF3
3. Build octets – start with terminal atoms.
1.
PF3 = 5 + 3 (7) = 26 valence eP (group 5A)
F
3 x F (group 7A)
2.  Skeleton (X is central in XYn ).
F
P
P
F
6 e- used in 3 bonds,
20 e- remain (10 pairs)
F
F
6 + 20 = 26 e-
F
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Guidelines for Writing Lewis Structures
3. Add e- pairs:
4 x O (grp 6A)
charge (-3) 3 e-
O P
O
2. Skeleton. P is central.
O P
O
32 e- used.
O
O
Add brackets and overall charge
to show this is an ion.
O
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8 e- used in 4 bonds,
24 e- remain (12 pairs)
O
= 5 + 4(6) + 3 = 32
P (grp 5A)
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Guidelines for Writing Lewis Structures
Phosphate ion, PO431. PO43-
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O
3-
O P O
O
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Single Covalent Bonds in Hydrocarbons
Single Covalent Bonds in Hydrocarbons
Larger alkanes:
Methane:
butane
H
H C H
H
2-methylpropane
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Single Covalent Bonds in Hydrocarbons
Single Covalent Bonds in Hydrocarbons
Cyclic alkanes:
Each alkane H-atom can be replaced by:
•  another atom.
•  a functional group.
1-chloro-2-methylpropane
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H
H
H
Cl C
H
C
C H
H
H C H
H
2-fluorobutane
H C C C C H
H H H H
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Multiple Covalent Bonds
Too few “dots” to complete all the octets?
Convert lone pairs to shared pairs.
Convert lone pairs to bond pairs.
3.  6 e- in bonds. Add the other
3 pairs to O (outer atom).
•  Each H shares 2 e•  C only “has” 6 e-.
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cyclopentanol
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Multiple Covalent Bonds
2. Skeleton.
OH
H F H H
H
H C O
H
H
H C O
H C O
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Each H shares 2 eC shares 8.
O shares 8.
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Multiple Covalent Bonds
Multiple Covalent Bonds in Hydrocarbons
O C O
2. Skeleton
propene C3H6
ethene C2H4
3.  4 e- in bonds. Add 3 pairs
to each O.
H
H
H C C H
H H
H C C C C H
H H
H H H
(propylene)
(ethylene)
O C O
H
H C C C H
H
4. Convert lone pairs to
bond pairs.
butene C4H8
H
(1 isomer shown)
(= is a line angle structure: see chapter 3 notes)
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Multiple Covalent Bonds in Hydrocarbons
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Double Bonds and Isomerism
Bond rotation:
propyne
H
H-C≡C-C-H
H
≡
ethyne (acetylene)
H-C≡C-H
≡
C3H4
C2H2
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Double Bonds and Isomerism
Cl
Cl
C=C
H
H
Cl
H
Cl
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Cl
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Double Bonds and Isomerism
Cl
Cl
H
Cl
Cl
H
C=C
C=C
H
H
Cl
Cl
Cl
Cl
C=C
H
No rotation around
a C=C bond.
Free C–C
rotation.
Cl
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Double Bonds and Isomerism
Bond Properties: Bond Length
Atom size and bond
type are important:
CH3
H 3C
H
H
H 3C
C=C
C=C
H
H
CH3
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Bond Properties: Bond Length
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Bond Properties: Bond Enthalpies
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Bond Properties: Bond Enthalpies
Bond Properties: Bond Polarity
Estimate ΔH for the following reaction from bond E.
Bonding pairs are not always equally shared.
CH4(g) + 2O2(g) → CO2(g) + 2H2O(g)
F
ΔH = {sum of bonds broken} – {sum of bonds formed}
F
= {4 DC-H + 2 DO-O} – {2 DC-O + 4 DH-O}
= {4(416) + 2(498)} – {2(803) + 4(467)}
= −814 kJ
(experimental value is −802 kJ)
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δ+
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H F
F has stronger e- attraction.
δ- Polar bond.
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Bond Properties: Electronegativity
Bond Properties: Electronegativity
•  Linus Pauling developed the first scale
•  Based on bond energies.
•  Pauling’s scale: F = 4.0 (arbitrary).
•  Unitless.
increasing
electronegativity
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Bond Properties: Electronegativity
Bond Properties: Electronegativity
Differences, ΔEN, determine bond polarity:
For the following bond pairs, indicate the δ+ and δatoms and choose the more polar bond:
Cl–F and Br–F; Si–Br and C–Br
ΔEN = 0.0
ΔEN = 0.7
ΔEN = 1.9
ΔEN = 3.0
increasing
electronegativity
covalent bond
polar covalent bonds
ionic bond
increasingly ionic
ΔEN (F and Br) > ΔEN (F and Cl)
BrF is more polar than ClF.
increasingly covalent
(Electronegativities: Br = 2.8, Cl = 3.0, F = 4.0
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Bond Properties: Electronegativity
Formal Charge
For the following bond pairs, indicate the δ+ and δ- atoms and choose the
more polar bond: Cl–F and Br–F; Si–Br and C–Br.
Used to study charge distribution in a molecule.
increasing
electronegativity
2.  Formal charge of each atom
Largest ΔEN between Si and Br
SiBr is more polar than SiC.
= (# of valence e-) – (e- “on” the atom).
Note: sum of formal charges = molecular charge
(Electronegativities: Si = 1.8, C = 2.5, Br = 2.8
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Formal Charge
Formal Charge
If there is choice between Lewis structures:
[ O C N ]–
Valence eLone pair e½ shared eFormal Charge
O
6
6
1
-1
•  Smaller formal charges are favored.
•  Negative formal charges should be on the most EN atoms
•  Like charges should not be on adjacent atoms
C
4
0
4
0
N
5
2
3
0
O N N
O N N
Formal charges: -1
+1
0
0
+1
-1
Check: Σ (formal charges) = ion charge = -1
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Formal Charge
Lewis Structures and Resonance
Which ClO2 structure is preferred?
Ozone has 2 equivalent structures:
-
Formal charges:
O Cl O
O Cl O
-1
-1
0
0
+1
O O
O
O
O
O
Both:
•  obey the octet rule
•  have the same number and types of bonds
•  have the same formal charges
-1
Experiments show that the OO bonds are identical.
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Lewis Structures and Resonance
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Lewis Structures and Resonance
Resonance in CO32-
O
O
O
O
O O
Experiment: All three CO bonds = 129 pm
•  Resonance structures differ only in e- pair
positions.
Typical bond lengths
C-O =143 pm; C=O = 122 pm.
•  Atom positions must not change.
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Exceptions to the Octet Rule
Fewer than Eight Valence Electrons
Often very reactive.
Be and B form e- deficient compounds:
H
Be
H
2 + 2(1) = 4 valence e-
H
F
F
F
B
F
3 + 3(7) = 24 valence e-
B
F
+
N
F
H
H
F
H
B
N
F
H
H
F
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Odd Number of Valence Electrons
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More Than Eight Valence Electrons
Some stable molecules have an odd number of
e -.
N O
NO2 5 + 2(6) = 17 valence e-
O N O
Very reactive. Most stable molecules have paired e-.
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More Than Eight Valence Electrons
ClF3
7 + 3(7) = 28 val. e-
Make octets on F.
e-
24 used, 4 remain.
[bonds (3 x 2); Lone pairs (3 x 6)]
•  Typically have strong, often pleasant, odors.
•  Contain a benzene or benzene-like ring
Cl
Its C-C bonds have equal lengths:
F
or
A solid ring
shows
resonance in
aromatics
Add 2 lone pairs to Cl – the 3rd period element.
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Aromatic Compounds
F
F
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Aromatic Compounds
Aromatic Compounds
Other aromatic examples:
Naming benzene rings:
ortho
C
O
=
CH3
CH3
CH3
CH3
H
benzaldehyde
para
meta
toluene
CH3
CH3
naphthalene
o-xylene
m.p./°C
b.p./°C
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m-xylene
-25.2
144.5
-47.8
139.1
CH3
p-xylene
13.2
138.4
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Aromatic Compounds
Systematic names:
CH3
6
2
CH3
3
5
CH3
CH3
1
CH3
4
1,2 –dimethylbenzene
1,3 dimethylbenzene
Cl
Cl
Cl
Cl
Cl
Cl
1,2,3–trichlorobenzene
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CH3
1,4 dimethylbenzene
Cl
Cl
1,2,4–trichlorobenzene
Cl
1,3,5–trichlorobenzene
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