CHAPTER 4 - CHEMICAL BONDING

sp hybridization

• sp hybridization is the mixing of 1 s orbital
and 1 p orbitals to form 2 equivalent sp
orbitals.

• The shape of the two hybrid orbitals is linear
with the angle of 180o

201

s px sp sp
202

• simplified drawing of sp orbitals:

sp sp Shown together
(large lobes only)

203

Example:

1) BeCl2
• Lewis structure :

• Valence orbital diagram;

Cl ground state : 3p
2p
3s

Be ground state :
2s

Be excited :

2s 2p

Be hybrid : 204

sp



p sp sp p

205

Example:

2) C2H2

• Lewis structure :

• Valence orbital diagrHam; C C H

C ground state : 2s 2p
C excited : 2s 2p

C hybrid : sp 2p

H ground state :

1s

206

Orbital Overlap: Example:



 C  
sp sp C
1s sp sp 1s

Molecular Shape : Linear

207

208

Example: 209
3) CO2
• Lewis structure :

• Valence orbital diagram;
Oground state :
C ground state :
C excited :
C hybrid :

• Orbital overlap:

3) CO2
 Lewis structure : O C O

 Valence orbital diagram;

C ground state : O ground state :

2s 2p O excited : 2s 2p
2p O hybrid : 2s 2p
C excited :
sp2 2p
2s

C hybrid : sp 2p

210

Orbital Overlap: O CO

 

s p2 s p2

 O

O sp2 sp C sp s p2 s p2
p
s p2

p pp

molecular shape: linear

211

sp3d hybridization
• sp3d hybridization is the mixing of 1 s orbital, 3

p orbitals and 1 d orbital to form 5 equivalent
sp3d orbitals
• The shape of the five hybrid orbitals is trigonal
bipyramidal with the angle of 120o and 90o

212

• simplified drawing of sp3d orbitals:

s p3d

s p3d

s p3d

s p3d

s p3d

213

Example:

1) PCl5 Cl
• Lewis structure : Cl P Cl

• Valence orbital diagram; Cl Cl

P ground state : 3d
3d
P excited : 3s 3p
3s 3p

P hybrid :

Cl ground state : sp3d
3s 3p
214

Orbital overlap: p
p
Cl

 p

s p3d  Cl

 P s p3d
Cl sp3d

s p3d sp3d 
Cl

p


Cl

p

Molecular Shape : Trigonal bipyramidal 215

Example:

2) ClF3
• Lewis structure :

• Valence orbital diagram;

Cl ground state :

Cl excited :

Cl hybrid :

F ground state : 216

Example: F
F Cl F
2) ClF3
• Lewis structure :

• Valence orbital diagram;

Cl ground state :

3s 3p 3d
3p 3d
Cl excited :

3s

Cl hybrid :

F ground state : sp3d 217
2s 2p

Orbital overlap: p
p
F


 s p3d s p3d s p3d
F s p3d
Cl

s p3d


F

p
Molecular Shape : T - shaped

218

sp3d2 hybridization

• sp3d2 hybridization is the mixing of 1 s
orbital, 3 p orbitals and 2 d orbitals to form 6
equivalent orbitals

• The shape of the six hybrid orbitals is
octahedral with the angle 90o

219

• Simplified drawing of sp3d2 orbitals:

sp3d2 sp3d2
sp3d2

sp3d2 sp3d2

sp3d2

220

Example: 221
1) SF6
• Lewis structure :

• Valence orbital diagram;
F ground state :
S ground state :
S excited :
S hybrid :

• Orbital overlap:

1) SF6 FF
 Lewis structure : F SF

 Valence orbital diagram; F F

S ground state : 3s 3p 3d
3d
S excited : 3s 3p

S hybrid : sp3d2

F ground state :

2s 2p 222

Orbital overlap: p
p
F
p
F   F p
p
s p3d2 s p3d2

s p3d2

Cl

 sp3d2 s p3d2 
F
F sp3d2


F

p

Molecular Shape : Octahedral

223

Example: 224
2) ICl5
• Lewis structure :

• Valence orbital diagram;
I ground state:
I excited :
I hybrid :
Clground state :

• Orbital overlap:

2) ICl5 Cl
 Lewis structure : Cl I Cl

Cl Cl

 Valence orbital diagram;

I ground state:

I excited : 5s 5p 5d
5s 5p 5d

I hybrid : sp3d2

Cl ground state :

3s 3p 225

Orbital overlap:

p

Cl

p Cl    Cl p

s p3d2

s p3d2 s p3d2

 sp3d2 I 

s p3d2

Cl s p3d2 Cl

p p

Molecular Shape : Square pyramidal

226

Exercise:

• For each of the following, draw the orbital
overlap to show the formation of covalent
bond
a) XeF2
b) O3
c) ICl4
d) OF2

227

a) XeF2 F Xe F
• Lewis structure :

• Valence orbital diagram;

Xe ground state : 5s

Xe excited : 5s 5p 5d
5p 5d

Xe hybrid : sp3d
F ground state : 2s 2p

228

Orbital overlap:

p

F


s p3d s p3d s p3d
s p3d
Xe

s p3d


F

p

Molecular Shape : Linear

229

b)O3 + -
• Lewis structure :
OO O
• Valence orbital diagram; 1 2

O+ ground state : 2p
2p
O+ excited : 2s
2s

O+ hybrid : sp2 2p

230

• Valence orbital diagram; + -

OO O
1 2

O1 ground state : 2p
2p
2s
2p
O1 excited :

2s

O1 hybrid :

sp2

O-2 ground state :

2s 2p

231

Orbital Overlap:



s p2  s p2

O2- sp2 sp2 O+

s p2 s p2 
2p
2p

O-

2p

Molecular Shape : V - shaped

232

C) ICl-4 -
 Lewis structure :
Cl Cl

 Valence orbital diagram; +

I- ground state: I Cl
Cl
I- excited : 5s 5p
I- hybrid : 5s 5p 5d

5d

sp3d2

Cl ground state :

3s 3p 233

Orbital overlap:

p Cl  s p3d2  Cl p

s p3d2 s p3d2 

 sp3d2 I-

32

sp d

Cl s p3d2 Cl

p p

Molecular Shape : Square planar

234

3) OF2 F
 Lewis structure : FO

 Valence orbital diagram;

O ground state :

2s 2p

O excited state : 2s 2p
O hybrid :

sp3

F ground state : 235

2s 2p

Answer Orbital Overlap:

Molecular Shape : V- shaped

s p3

 O

s p3 s p3

F2p -  sp3

F-

2p

236







4.4 Intermolecular forces

Learning Outcomes:

a) Describe intermolecular forces : van der Waals
forces, dipole-dipole interactions or permanent
dipole, London forces or dispersion forces,
hydrogen bonding.

b) Explain factors that influence van der Waals forces
and Hydrogen bond

c) Relate the effects of hydrogen bonding on the
physical properties:

i) boiling point

ii) solubility

iii) density of water compared to ice 241

Intermolecular Forces

 Intermolecular forces are attractive
forces between molecules.

• Intermolecular forces are responsible
for the physical properties such as boiling
point, melting point and solubilities of
molecular compounds.

242

Classification of intermolecular forces:

Intermolecular Forces

van der Waal Forces Hydrogen Bond

Dipole-dipole London forces
interactions

243

A. Van der Waal Forces

 Forces that act between covalent molecules.

 Two types of van der Waals forces:

(i) Dipole-dipole interactions
- The intermolecular attraction between
oppositely charged poles of nearby polar
molecules.

(ii) London forces
- The intermolecular attraction between
molecules as a result of instantaneous
polarisation of their electron clouds.

244

(i) Dipole-dipole interactions
(or permanent dipole forces)

 Exist between polar covalent molecules
 Polar molecules have permanent dipole due to

the uneven electron distributions

Example: - + -

+

H Cl H Cl

Chlorine is more Dipole-dipole forces; the partial
electronegative,
thus it has higher positive end of one molecule attracts
electron density
the partial negative end of the other

molecule 245

• The electrostatic forces between the permanent
+ pole of a molecule and permanent  pole of
adjacent molecules produce dipole-dipole
forces

• Molecules with higher polarity;
- stronger dipole-dipole forces

• Polar molecules experience dipole-dipole forces
as well as London dispersion forces

246

(ii) London Forces (or dispersion forces)

 Exist between non-polar molecules

 Result from the temporary (instantaneous)
polarization of molecules

 The temporary dipole molecules will be attracted
to each other and these attractions is known as
the London forces or dispersion forces

247

The formation of London forces

 Electrons move randomly. At any instant,
electron distributions in one molecule may be
unsymmetrical.

 The end having higher electron density is
partially negative (–) and the other is partially
positive (+).

 An instant dipole moment that exists in a
molecule induces the neighbouring molecule to be
polar.

 The attraction that exists between the

instantaneous dipole and induced dipole is
known as London forces.

248

Example:
London forces in Br2

Electrons in a molecule
Br Br move randomly about the

nucleus

At any instant, the - The temporary dipole
electron density might molecule induce the
be higher on one side Br + - neighbouring atom to
Br Br be partially polar
+

Temporary Br
dipole molecule

London forces

249

Factors that influence the strength of the Van
der Waals forces.

1. The molecular size

Molecules with larger molecular size have
stronger van der Waals forces as they tend to
have more electrons involved in the London
forces.

Example:

Molecular Boiling Point
Mass (oC)

N2 28 77.3 250
Cl2 71 239.1