Chemical Bonding O Level Slides

1. Chemical Bonding
~ BOND - Force
~ Chemical BOND - WHY - HOW
- WHAT
1 2 3
<INTRAmolecular
INTERmolecular
- WITHIN
- BETWEEN
H O
H
H O
H
INTRA
INTER
INTRAmolecular Forces
INTERmolecular Forces
STRONGER THAN

- Water BOILS at 100 oC
- Water DECOMPOSES at
Temp. > 2000oC
1
Chemical Bonding Lecture- 1
AS Level
Chemistry

2. TYPES OF BOND
CHEMICAL BONDS ~ INTRAmolecular
PHYSICAL BONDS ~ INTERmolecular
 IONIC (or electrovalent )
 COVALENT
 DATIVE COVALENT (or COORDINATE)
 METALLIC
 Ion – Dipole Interaction
 Hydrogen bonds
 Permanent dipole-dipole interactions
 Induced dipole-dipole interactions (London forces)
2
Chemical Bonding Lecture- 1
AS Level
Chemistry

3.  IONIC Bonding :
1.Dot and Cross Diagram
 Formation of NaCl
Na
.
. .
.
. .
.
.
.
.
Cl
x
x
x
x
x
x
x
x
x
x
x
x
x
x
x
x
x
.
1-
1+
2 , 8 2 , 8 , 8
~ An electrostatic attraction between
positive and negative ions OR
~ Transfer of e 1- from Metal to Non-metal
[ [
[ [2.Dot and Cross Formula
 Formation of Al2O3
Al .
]
[
3+
O
x
x
x
x
x
x]
[
2-
2 x X 3
.
 Mg S  Na3P  Mg3N2
Self-Assessment :
3
Chemical Bonding Lecture- 2
AS Level
Chemistry

4.  IONIC Compounds :
Properties
-- Hard Solids
-- High MP & BP
-- Electrolytes
-- Non-Volatile
-- NO MOLECULES
Strong Ionic Bond
Ionic Bond Strength
Factor :
~ CHARGE DENSITY =
𝐶ℎ𝑎𝑟𝑔𝑒
𝑅𝑎𝑑𝑖𝑢𝑠
~ High
~ Small
~ HIGH
~ MORE
- NaF NaCl
- Na2O MgO
1+ 1- 1+ 1-
1+ 2- 2+ 2-
X X
X X
~ Produces IONIC Solution ,
when dissolved in water
~ Can Show CONDUCTION
<~ STRONG
~ WEAK
DEFINITION :
-Ionize FULLY
-Ionize PARTIALLY
CONDUCTION
 Mobile ELECTRONS
 Mobile IONS
 VIBRATION OF ATOMS
~ Metals / Carbon
~ IONIC Compounds
~ POLAR COVALENT Compounds
~ Semi-Conductor ( Si , Ge )
- MOLTEN & AQUEOUS
Solution ~ CONDUCTOR
4
Chemical Bonding
Lecture- 3
AS Level
Chemistry

5. -- NO MOLECULES ~ Clusters of IONS
~ CUBIC Crystal LATTICE Structure
LATTICE Structure
- 3 D / 3 Dimensional
2
1
3
4
IONIC Crystal L.S.
SIMPLE Molecular L.S.
MACROmolecular L.S.
METALLIC L.S. Na / Cu
I2 / S8 / P4
SiO2 / SiC
NaCl / MgO
UNIT Cell
Binary IONIC Compound :
NaCl MgO Al2O3
Na1+
Cl1-
Cl1-
Cl1-
Cl1-
Cl1-
Cl1-
Na1+
Na1+
Na1+
Na1+
Na1+
Al3+ O2-
O2-
O2-
O2-
O2-
O2-
Al3+
Al3+
Al3+
Al3+
Al3+
CUBE~ +VE -VE
5
Chemical Bonding Lecture- 3
AS Level
Chemistry

6. Types Of IONS :
+VE IONS :
11Na = 2 , 8 , 1
12Mg = 2 , 8 , 2
13Al = 2 , 8 , 3
2 , 8
2 , 8
2 , 8
11Na1+ =
12Mg2+ =
13Al3+ = }~Isoelectronic
- Ne - NH3
-VE IONS :
15P = 2 , 8 , 5
16S = 2 , 8 , 6
17Cl = 2 , 8 , 7
2 , 8 , 8
2 , 8 , 8
2 , 8 , 8
15P3 - =
16S2 - =
17Cl1 - = }
~Isoelectronic
- Ar - H2S
~ Positive Ions are Smaller than the respective Atoms
~ Positive Ions are Smallest with More protons
~ Negative Ions are Larger than the respective Atoms
~ Negative Ions are Largest with More – ve charge 6
Chemical Bonding Lecture- 4
AS Level
Chemistry

7. Polarization
+ -
~ Polarization ( Covalent Character)
Factors:
 +ve ion ~ High Charge & Small Radius
 -ve ion ~ High Charge & Large Radius
}
More
Polarization
The Process of Distortion or Deformation of a –ve ion by
a +ve ion
Questions : 1 Na 1+ Mg 2+ Al 3+ 2
3
4
P 3- S 2- Cl 1-
Why AlCl3 is Covalent but AlF3 is Ionic in nature?
Why Al2O3 is Amphoteric in nature?
~ Na2O
~ SO2
~ Al2O3
- Basic
- Acidic
- Amphoteric
~ Ionic
~ Covalent
~ Ionic & Covalent
NaCl MgCl2 AlCl3
7
Chemical Bonding Lecture- 5
AS Level
Chemistry

8. COVALENT Bonding :
~ Electrostatic attraction between shared
electrons and Nuclei of two atoms
Questions :
1
2
3
Dot & Cross Diagram
Dot & Cross Formula
Lewis Structure /
Structural Formula
STEPS :
~ Central Atom
~ Group Number / Valence e 1-
Lesser in Number / More Bonds
SF2 H2O2
S F
F
o
o
o
o
o
o
x
x
x
x
x
x
x
x
x
x
x
x
x
x S F
F
o
o
o
o
o
o x
x x
x
x
x
x
x
x
x
x
x x
x
S
o
o
o
o
F
F
x
x
x
x
x
x
x
x
x x
x
x
~ H2 ; O2 ; N2 ; CH4 ; CO2 ; NH3
8
Chemical Bonding Lecture- 6
AS Level
Chemistry

9. 9
Chemical Bonding Lecture- 6
AS Level
Chemistry
COVALENT Bonding :
O
O
O
O
C O
O
O
O
O
O
O
CO2

10. COVALENT Compounds
MOLECULES
~ SIMPLE ~ MACROMOLECULES
- Soft Solids , Liquids , Gases
- LOW MP & BP
- Volatile
- Electrolytes ~ If POLAR
Non-Electrolytes ~ If Non-
POLAR
- Hard Solids
- High MP & BP
- Non-Volatile
- Non-Electrolytes
- Inert
RULES for Bond Formation
1
2
3
Duplet Rule
Octet Rule
Expansion of Octet
Rule
- H2SO4
S
o
o
o
o
o
o
x
x O
x x
xx
x
O
O
O
x
x x
x
x
x
x
x
x
x
x
x
x
x
x
x
x .
. H
H
10
Chemical Bonding Lecture- 7
AS Level
Chemistry

11. Dative COVALENT Bonding :
COVALENT
A A
B B
o x
A B
o
x
o
o
A B
o
o
~ Definition
~ Conditions
- One atom must have EMPTY Orbital
- Other atom must have LONE Pair
` Acceptor
` Donor / Ligand
11
Chemical Bonding Lecture- 8
AS Level
Chemistry

12. Dative COVALENT Bonding :
NH3 + H 1+ NH 4
1+
H2O + H 1+ H3O 1+
H3O 1+ + H 1+ H4O +2
N
o
o
o
o
o
H
H
H
H
x
x
x
1 +
N H
H
H
H 1 +
O
H
H
H
1 +
o
o
Questions :
1
2
3
AlCl3 & NH3
BeCl2 & NH3
AlCl3 & AlCl3
Al Cl
Cl
Cl
NH3
Be Cl
Cl
NH3
NH3
Al
Cl
Cl
Cl
Al
Cl
Cl
Cl
12
Chemical Bonding Lecture- 8
AS Level
Chemistry

13. COVALENT Bond Strength
~ Bond Energy / Enthalpy
Bond Length
- Number of Shared pairs
- Radius of atom
H-H ( g ) H ( g ) + H ( g )
O=O ( g ) O ( g ) + O ( g )
N=N ( g ) N ( g ) + N ( g )
Heat
Heat
Heat
Small
More
Short
High
Questions :
1
2
Which ONE of these is Thermally most Stable?
HF , HCl , HBr , HI
Explain the Trend of B.E. from Chlorine to Bromine
to Iodine?
3
C C
H
H
H
H
F
Cl
~ Maximum B.E.
~ Minimum B.E.
A
B
C
D
D
B
1 9
6
17
13
Chemical Bonding Lecture- 9
AS Level
Chemistry

14. Metallic Bonding
~ Definition
~ Structure of a Metal
~ Melting & Boiling point
~ Conductivity of Metals
1
2
3
4
 Strength of Metallic Bond
- Ionic Charge
- Ionic Radius
High
Small
Questions :
- Number of Mobile e MORE
- Radius of +ve ion SMALL
Explain the trend of :
~ M.P. & B.P. AND ~ Conductivity
- Across the Period Na Mg Al
- Down the Group Li Na K
14
Chemical Bonding Lecture-
10
AS Level
Chemistry

15.  MOLECULAR Orbital Theory: ( MOT )
Basic Principles :
Atomic Orbitals Overlap
1
2 BONDS Formed =Number of UNPAIRED Valence e1-
OR
6C = 1s2 2s2 2p2
2s 2p
16S = 1s2 2s2 2p6 3s2 3p4
3s 3p
Sulfur ( IV ) Oxide , SO2 Sulfur ( VI ) Oxide , SO3
HYBRIDIZATION :
 Application of MOT
 For ELEMENTS , When :
BONDS Formed =Number of UNPAIRED Valence e1-
15

16. HYBRIDIZATION :
- EXCITATION of e1-
- MIXING of Orbitals
from LOW energy to HIGH energy
to get SAME Energy &
SAME Shape of Orbitals
}
1
2
6C = 2s 2p 2s 2p
Ground State Excited State
16S = 3s 3p 3d
16S = 3s 3p 3d
16

17. sp
HYBRIDIZATION Types for CARBON :
Tetrahedral ; 109.5
Trigonal Plannar ; 120
Linear ; 180
17

18. HYBRIDIZATION Practice
- PCl5 - S2Cl2
- H2SO4 - Cr2O7
2-
-XeF4 ; ICl3 ; SF6 ; SnCl2 ; SO2; SO3
-H2O2
-N2H4
-C2H4
-C2H2
-H3PO4 ; H2CO3 ; HClO4 ; H2C2O4 ; BrO3
1- ; MnO4
1-
-O3
-CO
-P4
-N2O5
P
o
o
o
o
o
Cl
Cl
Cl
Cl
Cl
x
x
x
x
x x
x
x
x
x
x
x
x x
x
x
x
x
x
x
x
x
x
x
x
x
x
x
x
x
x
x
x
x
x
S S
o
o
o
o
o
o
o
o
o
o
o
o
Cl
Cl x
x
x
x
x
x
x
x
x
x
x
x
x
x
S
o
o
o
o
o
o
x
x O
x x
xx
x
O
O
O
x
x x
x
x
x
x
x
x
x
x
x
x
x
x
x
x .
. H
H
Cr
o
o
o
o
o
o
Cr o
o
o
o
o
o
O
O
O
O
O
O
O x
x
x
x
x
x
x
x x
xx
x
x
x
x
x
x
x
x
x
x x
x x
x
x
x
x
x
x
x
x
x
x
x
x
x
x
x
x
x
x
.
.
1 -
1 -
18

19. SHAPES of Molecules
~ Shape / Geometry - 3 D Arrangement of
atoms
VSEPR Theory
V
S
E
P
R
- Valence
- Shell
- Electron
- Pair
- Repulsion
}
Conclusions :
- Valence e 1- pairs of the Central Atom
- e 1- pairs ~ repel
} Order of repulsion
D
E
C
R
E
A
S
E
S
Lone pair – Lone pair
Lone pair – Shared pair
Shared pair – Shared pair
Conditions :
1
2
Polyatomic Molecules ONLY
Double / Triple Bond ~ One e 1- pair or region
19

20. 
 
  
20

21. Classification of Covalent Bond
1
2
3
Number of Shared pairs
Single
Double
Triple
H-H
O=O
N=N
Overlap Sigma / End-On
Pi / Side-wise
Charge Separation Non-polar
Polar
Sigma /
End-On
Pi / Side-wise
`1 Sigma 2Pi
Single C.B `1 Sigma Double `1 Sigma 1 Pi Triple 21

22. N2
CH4 C2H6
C2H4 C2H2
O2
22

23. Non-polar
Polar
Shared pair is attracted by both atoms
with Equal force
OR
OR
– Same Electronegativity
– Different Electronegativity
Shared pair is attracted by both atoms
with Unequal force
OR
Bond between Identical Atoms
Bond between Different Atoms
OR
When Vector Sum / Net Dipole is ZERO
When Vector Sum / Net Dipole is not ZERO
Except C-H
H H H F
-
+
O O
C
Al
Cl
Cl
Cl
C
H
Cl
Cl
Cl
Non-polar
Alkanes
Alkenes
Polar
Alcohols
Carboxylic acids
Polar Molecules are Reactive than Non-polar N N C O 23

24. Aqueous solution form
Intermolecular Forces
 Ion – Dipole Interaction
- NaCl in water
Na Cl
1+ 1-
O-H
H
O-H
O-H
O-H
O-H
O-H
O-H
O-H
H
H
H
H
H
H
H
` Hydration
Hydration Enthalpy - Exothermic
Ion – Dipole Interaction
~ CHARGE DENSITY =
𝐶ℎ𝑎𝑟𝑔𝑒
𝑅𝑎𝑑𝑖𝑢𝑠
~ High
~ Small
~ HIGH
~ MORE
Definition
Na
1+
(g)
Na
1+
(aq)
+ aqua
Al
3+
(g)
Al
3+
(aq)
+ aqua
-ve
-ve
24

25. Intermolecular Forces
 Hydrogen bonds
Conditions
`Hydrogen atom covalently bonded to a highly Electro-ve atom
`Lone pair on the Electro-ve atom
Definition
H F H F
+ - -
+
H O
H
H O
H
- -
+ +
+ +
H N
H
H
+
+
+ -
H N
H
H
+
+
+
- A. C2H5OH & H2O
B. HCHO & H2O
} `Electronegativity
`Average Hydrogen bonds formed
FON
25

26.  Induced dipole-dipole interactions (London forces)
Intermolecular Forces
 Permanent dipole-dipole interactions
Polar Molecules
H Cl
+ -
H Cl
+ -
CH3
CH3
C O
-
+
CH3
CH3
C O
-
+
Non-polar Molecules
PD-PD
H H H H
CH3
CH2
CH2
CH2
CH3
CH3
CH2
CH
CH3
CH3
CH3
CH2
C
CH3
CH3
`Number of Electrons
`Area of Contact
FACTORS :
F2 Cl2 Br2 I2
1 .
2 . C5H12
26

27. 27