Chapter 10 Chemical Bonding I: The Lewis Model
10.1 Bonding Models and AIDS Drugs (p. 393)
10.2 Types of Chemical Bonds (p. 394)
10.3 Representing Valence Electrons with Dots (p. 396)
10.4 Ionic Bonding: Lewis Structures and Lattice Energies (p. 397)
10.5 Covalent Bonding: Lewis Structures (p. 404)
10.6 Electronegativity and Bond Polarity (p. 406)
10.7 Lewis Structures of Molecular Compounds and Polyatomic Ions (p. 410)
10.8 Resonance and Formal Charge (p. 412)
10.9 Exceptions to the Octet Rule: Odd-Electron Species, Incomplete Octets, and Expanded Octets (p. 417)
10.10 Bond Energies and Bond Lengths (p. 422)
10.11 Bonding in Metals: The Electron Sea Model (p. 425)
Suggested Exercises: 39, 51, 53, 59, 61, 63, 65, 69, 71, 73, 75, 79, 81, 83
10.2 Types of Chemical Bonds (p. 394)
A chemical bond is the force that holds atoms together in a compound. It is a cohesive force that is
characterized by electrostatic attraction. The two types are ionic and covalent.
An ionic bond is a chemical bond formed by the electrostatic attraction between cations (positive ions)
and anions (negative ions).
A covalent bond is a chemical bond formed by the sharing of a pair of electrons between atoms such as in
HCl. Most often, covalent bonding is observed in compounds that are composed of nonmetals only.
10.3 Representing Valence Electrons with Dots (p. 396)
A Lewis symbol is a representation of an element in which the chemical symbol stands for the core
electrons and the valence electrons are represented by dots placed around the letter symbol of the
element. The number of valence electrons is the same as the group number for groups 1 and 2. For
groups 13 - 18, subtract 10 from the group number to determine the number of valence electrons. The
Lewis symbols for Li through Ne are as follows.
Li
Be
B
C
N
O
F
Ne
10.4 Ionic Bonding: Lewis Symbols and Lattice Energies (p. 397)
10.5 Covalent Bonding: Lewis Structures (p. 404)
Covalent bond: A chemical bond formed by the overlap of atomic orbitals. The electron pair between the
atoms is shared to some extent. A bond is considered to be covalent when the electronegativity difference
between the bonded atoms is 2.0 or less.
1. Single bond: A covalent bond of bond order = 1, consisting of two electrons in a sigma () bond. A
single bond does not have a pi bond ().
2. Double bond: A covalent bond of bond order = 2, consisting of four electrons; one pair in a sigma
bond and the other pair in a pi bond ().
3. Triple bond: A covalent bond of bond order = 3, consisting of six electrons; one pair in a sigma
bond () and the other pairs in pi bonds ().
4. Lone pair (non-bonding pair, unshared pair): A valence shell electron pair associated with one
atom, and not part of a covalent bond.
1
Examples
Nitrogen trichloride
Cl
N
Cl
Oxygen
Nitrogen
O O
N N
3-pentene-1-yne
H
H
Cl
Description
lone pairs: 10
single bonds: 3
lone pairs: 4 lone pairs:
2
single bonds: 1
single bonds: 0
H
C
C
H
H
C
C
C
H
lone pairs: 0
single bonds: 8
double bonds: 0
double bonds: 0
double bonds: 0
double bonds: 1
triple bonds: 0
triple bonds: 0
triple bonds: 1
triple bonds: 1
10.6 Electronegativity and Bond Polarity (p. 406)
Electronegativity (EN or ) is a measure of the ability of an atom in a molecule to draw bonding electrons
to itself. The general trend is that EN increases left to right along a period and bottom to top in a group.
The student should learn the values of H, B through F and the remainder of the halogens and S.
General Trend of Electronegativity ()
1
1A
1
H
2.1
3
Li
1.0
11
Na
1.0
19
K
0.9
37
Rb
0.9
55
Cs
0.8
87
Fr
0.8
18
8A
2
2A
4
Be
1.5
12
Mg
1.2
20
Ca
1.0
38
Sr
1.0
56
Ba
1.0
88
Ra
1.0
3
3B
21
Sc
1.3
39
Y
1.2
57
La
1.1
89
Ac
1.1
4
4B
22
Ti
1.4
40
Zr
1.3
72
Hf
1.3
5
5B
23
V
1.5
41
Nb
1.5
73
Ta
1.4
6
6B
24
Cr
1.6
42
Mo
1.6
74
W
1.5
7
7B
25
Mn
1.6
43
Tc
1.7
75
Re
1.7
8
8B
26
Fe
1.7
44
Ru
1.8
76
Os
1.9
9
8B
27
Co
1.7
45
Rh
1.8
77
Ir
1.9
10
8B
28
Ni
1.8
46
Pd
1.8
78
Pt
1.8
11
1B
29
Cu
1.8
47
Ag
1.6
79
Au
1.9
12
2B
30
Zn
1.6
48
Cd
1.6
80
Hg
1.7
13
3A
5
B
14
4A
6
C
15
5A
7
N
16
6A
8
O
17
7A
9
F
13
Al
1.5
31
Ga
1.7
49
In
1.6
81
Tl
1.6
14
Si
1.8
32
Ge
1.9
50
Sn
1.8
82
Pb
1.7
15
P
2.1
33
As
2.1
51
Sb
1.9
83
Bi
1.8
16
S
17
Cl
34
Se
2.4
52
Te
2.1
84
Po
1.9
35
Br
36
Kr
53
I
54
Xe
85
At
2.1
Classification of a bond is dependent upon the electronegativity difference between the bonded atoms.
Example
Electronegativity Difference (EN) Bond Type
0 – 0.4 (small)
Covalent
Cl2
0.4 – 2.0 (intermediate)
Polar Covalent HCl
> 2.0 (large)
Ionic
NaCl
2
Depicting Polar Covalent Bonds (p. 408)
There are two methods commonly used to depict an unequal sharing of electrons in a covalent bond.
 
One method uses a lower case Greek letter delta (δ).
H Cl
The δ+ signifies a partial positive charge.
The δ- signifies a partial negative charge.
The second method uses a cross-based arrow.
H Cl
The arrow points to the more electronegative element.
Based upon electronegativity, place δ+ and δ- signs over the appropriate atoms in the following bonds.
Br H
Br C
Br N
Br S
Br O
Br Br
10.7 Lewis Structures of Molecular Compounds and Polyatomic Ions (p. 410)
1. Determine the total number of valence electrons (bonding electrons), which is the element's group
number.
a. Add 1 for each unit of negative charge.
b. Subtract 1 for each unit of positive charge.
2. Draw a skeletal structure by determining the central atom and drawing single bonds to each of the
bonded atoms surrounding it.
a. The central atom is most often the least electronegative atom.
b. Hydrogen and fluorine are terminal bonded atoms.
c. In oxoacids (e.g., sulfuric acid – H2SO4) hydrogen is bonded to oxygen.
d. Molecules and ions are compact and symmetrical.
e. When Kr and Xe are present in a molecule, they are central atoms.
3. Complete the octet around each bonded atom. The octet rule is the tendency of atoms in molecules to
have eight electrons in their valence shell.
Octet Rule Exceptions
a.
Hydrogen requires only 2 electrons.
b.
Beryllium requires only 4 electrons.
c.
Boron requires only 6 electrons.
d.
Elements, starting with phosphorus and beyond, can exceed the octet rule.
4. Distribute the remaining electrons as lone pairs to the central atom.
5. If there are less than eight electrons on the central atom move one or two electron pairs from a
bonded atom to the central atom to form multiple bonds. Atoms often forming multiple bonds are C,
N, O, and S.
#1
SiCl4
#2
CH4
#3
NH3
#4
H2O
#5
#6
#7
#8
#9
#10
31+
PO4 NH3
H2SO4
O2
CO2
CO
Characteristic Bonding Patterns
Atom
Bonds
Lone Pairs
C
4
0
N
3
1
O
2
2
F
1
3
3
10.8 Resonance and Formal Charge (p. 412)
A resonance structure is one of several valid Lewis structures which have identical skeletal arrangements
but differing electron placement.
#1 Ozone #2 Carbonate ion
O3
CO32Formal charge (p. 414) is a hypothetical charge assigned to atoms in a Lewis structure.
Formal charge is a hypothetical charge assigned to atoms in a Lewis structure.
 It facilitates the determination of the best Lewis structure among competing Lewis structures.
 Assignment of a formal charge is based upon the assumptions that bonding electrons are equally
shared between bonded atoms and that the electrons of each lone pair belong completely to one
atom.
Formal charge = Group # - (Lone pair electrons + ½ of the bonding electrons)
Guidelines for Determining the Best Lewis Structure from Among Two or More Valid Lewis Structures
1.
Whenever you can write several Lewis formulas for a molecule, choose the one having the lowest
magnitudes of formal charge.
2.
When two proposed Lewis formulas for a molecule have the same magnitudes of formal charges,
choose the one having the negative formal charge on the more electronegative atom.
3.
Like charges should not be on adjacent atoms.
#1 Sulfuric acid
H2SO4
#3 Dinitrogen monoxide
N2O
#3 Cyanate ion
OCN1-
#4 Fulminate ions
CNO1-
Formal charge on the cyanate ion: OCN1Cyanate anion
N C O
0 0 -1
-1 0
0
-2 0 +1
N C O
I.
N C O
II.
N C O
III.
Formal charge on the fulminate ion: CNO1Fulminate anion
C N O
-1 +1 -1
-2 +1 0
-3 +1 +1
C N O
C N O
C N O
I.
II.
III.
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Odd-Electron Species (p. 418)
NO
NO2
ClO2
Nitrogen monoxide
Nitrogen dioxide
Chlorine dixide
Atoms and molecules with unpaired electrons are referred to as free radicals.
Incomplete Octets (p. 418)
H
Be
Hydrogen molecular
Beryllium dichloride
H2
BeCl2
Boron
Boron trihydride
BH3
Expanded Octets (p. 420)
Elements in Period 3 from phosphorus and beyond can exceed the octet rule. This is possible due to the
presence of the energetically accessible d orbitals in these elements.
#1
#2
#3
#4
H2SO4
AsCl5
XeF2
XeOF2
10.10 Bond Energies and Bond Lengths (p. 422)
Bond Enthalpies (bond energy) (p. 422)
Bond energy (D) is the average enthalpy change for the breaking of an A-B bond in a molecule in the gas
phase. Because it requires energy to break bonds (endothermic process), bond energies are always
positive numbers. When a bond is formed, the energy is equal to the negative of the bond energy (energy
is released, exothermic). See Table 10.3, p. 422.
 Bond formation – Exothermic
 Bond breaking – Endothermic
In general, the enthalpy of reaction (ΔH) is approximately equal to the sum of the bond energies (D) for
bonds broken minus the sum of the bond energies formed.
∆𝐻 = Σ𝐷𝐵𝑜𝑛𝑑𝑠 𝐵𝑟𝑜𝑘𝑒𝑛 − Σ𝐷𝐵𝑜𝑛𝑑𝑠 𝐹𝑜𝑟𝑚𝑒𝑑
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BOND ENERGY PROBLEM #1 Using bond energies, calculate the enthalpy of reaction for the complete
combustion of butane. {-5228 kJ}
2 C4H10(g) + 13 O2(g)  8 CO2(g) + 10 H2O(g) ΔH = ?
H H H H
2
H C
C
C
C H
+ 13
8
O O
O C O
+ 10 H
O H
H H H H
D (kJ/mol): C-H 414; C-C = 347; O=O 498; C=O* 799; H-O 464 [Table 10.3, p. 422]
∆𝐻 = Σ𝐷𝐵𝑜𝑛𝑑𝑠 𝐵𝑟𝑜𝑘𝑒𝑛 − Σ𝐷𝐵𝑜𝑛𝑑𝑠 𝐹𝑜𝑟𝑚𝑒𝑑
BOND ENERGY PROBLEM #2
Using bond energies, calculate the enthalpy of reaction for the complete combustion of benzene,
C6H6. {17,274 kJ}
H = ?
2 C6H6(g)
12 CO2(g) + 6 H2O(l)
+ 15 O2(g)
Benzene
H
H
H
C
C
C
C
C
H
C
H
H
D (kJ/mol): C=C 611; C-C 347; C-H 414; O=O 498; C=O 799; H-O 464 [Table 10.3; p. 422]
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