Chapter 3: Chemical Bonding
Compounds are formed from chemically bound atoms or ions
Bonding involves only the valence electrons
Ionic compounds – ionic radii and lattice energies
Molecular compounds
– covalent (polar or non-polar) bond
– bond order
– bond strength
– Lewis structures
Lewis Symbols – show the valence electrons as dots around the atomic symbol
1
The Octet Rule
Atoms tend to lose, gain, or share electrons until they have eight valence electrons
2
3 4
1

Bond orders and bond distances
Single bond distances
Bond distances of other bond types
5
Bond strength
Bond dissociation energy (the energy required to break a bond)
6
/mol
7 8
2

9
11
10
12
3

13
One more example: HNO
3
(HONO
2
)
O
O
O
N
O H
24e
O
O
N
O H or
O
N
O H
O
N
O O H
Exceptions to the octet rule
Elements in the 2 nd period can never have more than 8 electrons
Elements in the 3 rd or higher periods can have more than 8 electrons
Incomplete octet Expanded octet
N O O N O
Odd-electron molecules
15
Concept of Resonance
More than one possible Lewis structures. The actual structure of a molecule is taken as a blend of all the feasible Lewis structures do not confuse with structural isomers
14
16
4

What does writing resonance structures accomplish?
Electrons are often delocalized between two or more atoms.
Electrons in a single Lewis structure are assigned to a specific atom. Therefore, a single Lewis structure is insufficient to electron delocalization. Composite of resonance structures more accurately depicts electron distribution and describes a molecule.
Concept of formal charge
17
19
18
Take SCN- as an example illustrating the importance of the formal charge concept
( Don't confuse the formal charge with the oxidation state or the charge on an atom )
More examples to show the application
(formal charge and expanded octet
20
5

Octet
Lewis
Structure
SNF
3
F
N
S
F
F
Atom
Formal
Charge
S
N
2+
2-
SO
2
Cl
2 O
Cl S
Cl
O
S
O
2+
1-
SO
4
2-
O
O
S
O
O
S
O
2+
1-
SO
3
2-
O
O
S O
S
O
XeO
3
O
O
Xe O
Xe
O
1+
1-
3+
1-
IOF
5
O
F
F
I
F
F
F
I
O
1+
1-
Expanded
Octet Lewis
Structure
F
F
O
I
F
F
F
Atom
I
O
Formal
Charge
0
0
Expanded
to
F
N
S
F
F
Cl
O
S
Cl
O
O
O
S
O
O
S
N
S
O
S
O
0
0
0
0
0
0, 1-
12
12
12
O
O
S O
S
O
O
O
Xe O
Xe
O
0
0
0
0, 1-
10
14
14
21
Molecular geometry
The properties of a compound are very much determined by the size and shape of its molecules.
Expanded octet
An atom that has a d subshell in the valence electron shell can accommodate more than an octet of electrons
Hypervalence :
Hypervalent molecules : species having more than an octet around at least one atom
VSEPR model ( V alence S hell E lectron P air R epulsions)
The most stable arrangement of groups attached to a central atom is the one that has the maximum separation of electron pairs
(bonded or non-bonded)
22
23 24
6

Electron-pair geometry
Molecules in which the central atom has no lone pairs
25
Molecules in which the central atom has one or more lone pairs
CH
4
NH
3
OH
2
27
Using the examples to illustrate the VSEPR model
SF
4
, BrF
3
, XeF
2
, XeF
4
, OSF
4
26
28
7

Using the examples to illustrate the VSEPR model
F
F Br
F
F
Xe
F
SF
4
, BrF
3
, XeF
2
, XeF
4
, OSF
4
F F
S
F
F
F S
F
F
F
F
Br
F
F
F
F Xe
F Br
F
F
Xe
F
F
F
F
S
F
F
F
Xe
F
F F
F
Xe
F
F
F
SF
4 adopts a seesaw structure
29
Guidelines for applying the VSEPR model
30
31 32
8

Valence Bond Theory describes covalent bond in terms of the overlap of atomic orbitals
Valence bond model of H
2
33
Valence bond theory and molecular geometry
Hybridization -- Hybrid orbitals
Hybrid orbitals are mixtures of atomic orbitals with intermediate energy
The number of s, p and d orbitals in the mixture equals the number of hybrid orbitals
35
34
36
9

37
39
38
40
10

Examples with an sp 3 hybridization central atom
Hybridization of s, p and d orbitals
41
Pentagonal bipyramidal
Geometry (sp 3 d 3 )
Examples:
PCl
5
, SF
4
, OSF
4
Examples:
SF
6
, XeF
4
IF
7
42
Hybridization in ethylene, acetylene, etc.
Concept of
σ and
π bonds
Bonds that result from
σ overlap are called
σ bonds.
Bonds that result from
π overlap are called
π bonds.
43 44
11

Structure of
Ethylene
45
Bond polarity and dipole moment
46
47 48
12

Molecular Dipole Moments
Molecule must have polar bonds (necessary but not sufficient)
Need to know molecular shape because individual bond dipoles can cancel
49
Examples
51
50
13