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MOLECULAR SHAPE AND POLARITY
Subtopic 4.2
1
LEARNING OUTCOMES (covalent bonding)
1. Draw the Lewis structure of covalent molecules (octet rule such as
N𝐻3, 𝐶𝐶𝑙4, 𝐻2O, 𝐶𝑂2, 𝑁2𝑂4, and exception to the octet rule such
as 𝐵𝐹3, NO, 𝑁𝑂2, 𝑃𝐶𝑙5, 𝑆𝐹6)
2. Explain the concept of overlapping and hybridisation of the s and
p orbitals such as 𝐵𝑒𝐶𝑙2, 𝐵𝐹3, 𝐶𝐻4, 𝑁2, HCN, 𝑁𝐻3, 𝐻2O
molecules
3. Predict and explain the shapes of and bond angles in molecules
and ions using the principle of valence valence shell electron pair
repulsion, e.g. linear, trigonal planar, tetrahedral, trigonal
bipyramid, octahedral, v-shaped, seesaw and pyramidal
4. Explain the existence of polar and non-polar bonds (including C-
Cl, C-N, C-O, C-Mg) resulting in polar or/and non-polar molecules
LEARNING OUTCOMES (covalent bonding)
5. Relate bond lengths and bond strengths with respect to single,
double and triple bonds
6. Explain the inertness of nitrogen molecule in terms of its strong triple
bond and nonpolarity
7. Describe typical properties associated with ionic and covalent
bonding in terms of bond strength, melting point and electrical
conductivity
8. Explain the existence of covalent character in ionic compounds such
as 𝐴𝑙2𝑂3, 𝐴𝑙𝑙3, and Lil
9. Explain the existence of coordinate (dative covalent) bonding such as
𝐻3𝑂+, N𝐻4+
, 𝐴𝑙2𝐶𝑙6 , and [Fe (𝐶𝑁)6]³ˉ
MOLECULAR GEOMETRY
Molecular
Geometry
The 3D
arrangement of
atoms in a
molecule Affects
physical and
chemical
properties
Is predicted by
using Valence
Shell Electron
Pair Repulsion
(VSEPR) model
VSEPR THEORY
1. Valence Shell Electron Pair Repulsion theory
2. It states that each group of valence electrons
around a central atom is located as far away
as possible from the others in order to
minimize repulsion
3. To predict the molecular shape from the Lewis
structure * Kalau dekat, akan repulse each other
VSEPR
THEORY
Two type of electron pair around
the central atom:
bonding pairs and lone pairs
Electron pairs around the central
atom will repel one another and
arrange themselves as far apart
as possible from each other.
VSEPR THEORY
Why? To minimise the electron
pair-electron pair repulsion
around the central atom
ELECTRON GROUP
1. Electrons pair of lone electron that occupy localized
region around the central atom
2. The repulsion may occur either between :
Bonding pair (single bond or double bond or
triple) with bonding pair
Lone pair (odd electron) with bonding pair
Lone pair with lone pair
Bonding pair electron
Lone pair
electron
Note:
The electron pairs repulsion will determine the orientation of atoms in space
ELECTRON GROUP
Bonding pair-bonding pair repulsion
Lone pair-lone pair repulsion
Lone pair-bonding
pair repulsion
> >
Decrease of the repulsion force
FIVE BASIC MOLECULAR SHAPES
1. Linear (180°)
2. Trigonal planar (120°)
3. Tetrahedral (109.5°)
4. Trigonal bipyramid (90°, 120°)
5. Octahedral (90°)
ELECTRON GROUP ARRANGEMENT
1. Determined by number of electron groups around the
central atom
2. Can be one of the 5 basic shapes :
2 electron group/bonding pair : linear
3 electron group/bonding pair : trigonal planar
4 electron group/bonding pair : tetrahedral
5 electron group/bonding pair : trigonal bipyramidal
6 electron group/bonding pair : octahedral
1. Double bond and triple bonds can be treated
like single bond (approximation)
One electron group
2. Order of electron group repulsion
Lone pair – lone pair > lone pair – bonding pair
> bonding pair – bonding pair
SOME RULES WHEN USING VSEPR THEORY
SHAPE OF MOLECULE
1. Basic shapes are based on the repulsion between the
bonding pairs.
2. Tips to determine the molecular shape :
Step 1 : Calculate total number of valence electron
Step 2 : Draw Lewis structure of the molecule
Step 3 : Calculate the number of bonding pairs or
electron groups (Place bonding pairs as far as
possible to minimize repulsion)
Step 4 : Predict the shape using VSEPR theory
A. MOLECULAR SHAPE WITH 2 BONDING PAIR
Example: BeCl2
Lewis structure
shape
Linear
180° Be : 2e
2Cl : 2 X 7e = 14e
Total : 16 e
Cl .. : .. Cl Be
.. : ..
Electron group arrangement : linear
Molecular shape : linear (bond angle : 180°)
A𝑋2
B. MOLECULAR SHAPE WITH 3 BONDING PAIR
Example: BCl3 Lewis structure
B: 3e
3Cl : 21e
Total: 24e
B .. : ..
Cl
Cl
Cl .. :
..
..
: ..
Repulsive forces
between pairs are the same
120°
Trigonal planar
A𝑋3
Electron group arrangement : trigonal planar
Molecular shape : trigonal planar (bond angle : 120°)
C. MOLECULAR SHAPE WITH 4 BONDING PAIR
Example: CH4
Lewis structure
C
H
H H
H
Equal repulsion between bonding pairs – equal angle
109.5°
Tetrahedral
C: 4e
H : 4e
Total: 8e
Electron group arrangement : tetrahedral
Molecular shape : tetrahedral (bond angle :109.5°)
Pyramid is not tetrahedral
AX4
D. Molecules with 5 bonding pairs
Example: PCl5 Lewis structure
P
Cl
Cl
Cl
Cl
Cl
..
: ..
..
: ..
..
: ..
..
: ..
Shape:
120°
Trigonal bipyramidal
90°
Electron group arrangement : trigonal bipyramidal
Molecular shape : trigonal bipyramidal
bond angle :120°, 90°
E. Molecules with 6 bonding pairs Example: SF6
Lewis structure
S : 6e 6F : 42e Total : 48e
Octahedral
S
F
F
FF
F
F
90o
90o
Electron group arrangement : octahedral
Molecular shape : octahedral
bond angle : 90°
19
Class of molecules
Number of bonding
pairs
Number of lone pairs
Shape
AX2E
e.g. SnCl2
Sn: 4e
2Cl : 14e
Total : 18e
2 1
Bent / V-shaped
Bond angle : < 120o
Shape of molecules which the central atom has one or more
lone pairs (3 ELECTRON PAIRS)
Example: SnCl2 Lewis structure
Sn: 4e
2Cl : 14e
Total: 18e
AX2E
Electron group arrangement : trigonal planar
Molecular shape : bent or v-shaped (bond angle < 120°)
Sn
.. : ..
Cl Cl .. :
:
:
1. Calculate total number of valence
electron
2. Draw the Lewis structure
3. Calculate the number of bonding
pairs or electron groups (Place
bonding pairs as far as possible to
minimize repulsion)
4. Predict the shape using VSEPR theory
21
Class of molecules
Number of bonding
pairs
Number of lone pairs
Shape
AX3E
e.g. N𝐻3
3 1
Trigonal pyramidal Bond angle : < 109.5o
4 electron pairs in the valence shell of central
atom:
N : 5e
H : 3e
Total : 8e
22
Class of molecules
Number of bonding
pairs
Number of lone pairs
Shape
AX2E2
e.g. 𝐻2O
2 2
Bent / V-shaped Bond angle : < 109.5o
4 electron pairs in the valence shell of central atom:
O : 6e
H : 2e
Total : 8e
23
Class of molecules
Number of bonding
pairs
Number of lone pairs
Shape
AX4E
e.g. S𝐹4
4 1
Distorted tetrahedral (see-saw) Bond angle : < 90o, < 120o
5 electron pairs in the valence shell of central atom:
S : 6e
F : 28e
Total : 34e
24
Class of molecules
Number of bonding
pairs
Number of lone pairs
Shape
AX3E2
𝑒𝑔. 𝐵𝑟𝐹3
3 2
T-shaped Bond angle : < 90o
5 electron pairs in the valence shell of central
atom:
Br : 7e
F : 21e
Total : 28e
25
Class of molecules
Number of bonding
pairs
Number of lone pairs
Shape
AX2E3
eg. 𝐼3−
2 3
Linear Bond angle : 180o
5 electron pairs in the valence shell of central
atom:
I : 7e X 3
Total : 21e
26
Class of molecules
Number of bonding
pairs
Number of lone pairs
Shape
AX5E
e.g. I𝐹5
5 1
Square pyramidal Bond angle :90o and
180o
6 electron pairs in the valence shell of central
atom:
I : 7e
F : 35e
Total : 42e
27
Class of molecules
Number of bonding
pairs
Number of lone pairs
Shape
AX4E2
e.g. Xe𝐹4
4 2
Square planar Bond angle : 90o
6 electron pairs in the valence shell of central atom:
Xe : 8e
F : 28e
Total : 36e
28
KEEP IN MIND !
Elements period 3 or higher can form compounds that have both
octet and expanded octet valence shell
Example :
𝑃𝐶𝑙3 + 𝐶𝑙2 → 𝑃𝐶𝑙5
octet expanded octet
*Period 3 : Na, Mg, Al, Si, P, S, Cl, Ar
Never “expand” atoms of period 1 and 2 such as C and O.
They must be octet!
********************************************************
KEEP IN MIND !
Elements period 3 or higher can form compounds that have both
octet and expanded octet valence shell
Example :
𝑃𝐶𝑙3 + 𝐶𝑙2 → 𝑃𝐶𝑙5
octet expanded octet
*Period 3 : Na, Mg, Al, Si, P, S, Cl, Ar
Never “expand” atoms of period 1 and 2 such as C and O.
They must be octet!
Example :
********************************************************
Predict the geometry molecular shape for the
following molecules and ions by using VSEPR theory:
(a) PCl5 (b) IF
4+
(c) ClF3 (d) I
3-
(e) IOF5 (f) BrF
5
(g) XeF4 (h) O
3
EXERCISE 1
ClF3
I3
-
1. Valence electron
2. Lewis structure
3. Lone and bond pair
4. Predict the shape
ANSWER
32
Step 1 : Write the Lewis structure to see the number of electron groups (Bonding and lone pairs at the central atom)
Step 2 : Determine the electron group arrangement
Step 3 : Predict the ideal bond angle and any deviation caused by lone pairs
Step 4 : Draw and name the molecular shape by counting the bonding groups
GUIDELINES FOR APPLYING VSEPR THEORY
POLAR AND NON POLAR BONDS
1. Atoms with different electronegative from polar bonds (difference in EN)
2. Depicted as polar arrow :
3. Example :
C - Cl
BOND POLARITY
𝛿 − 𝛿 + C - Cl
# Polar bond
POLAR BOND
H F
electron rich
region electron poor
region
Polar bond
36
NON – POLAR BOND
F F
Non – polar bond
37
Example :
C – F
C – C (no difference of electronegativity)
BOND POLARITY
𝛿 + 𝛿 − C - F
# Polar bond
# Non polar bond
ELECTRONEGATIVITY
Electronegativity of an atom is the ability of an atom that is
covalently bonded to another atom to attract electrons to itself
An atom with a high electronegativity will attract electrons
towards itself and away from the atom with a lower
electronegativity
Electronegativity of an atom is inversely proportional to its size.
The smaller the atomic size, the stronger the attraction for the
bonding electrons and the higher electronegativity.
1. For covalent bond that consists of two identical atoms, the bonding electrons are shared equally between the two atoms and are attracted equally to both the atoms. This type of bond is called non-polar bond.
2. Diatomic molecules such as 𝐻2, 𝑂2, 𝐶𝑙2, 𝑁2, have non polar bonds because the atoms in the molecules have same electronegativity. These molecules are called non polar molecules.
NON POLAR BONDS
1. In covalent bond that contains two atoms that are not identical. The bonding electrons will be attracted more strongly by the more electronegative element and this result in unsymmetrical distribution of electrons.
2. Ex : 𝐻 − 𝐶𝑙 . The more electronegative chlorine atom attracts the bonding electron pair more strongly than hydrogen does.
3. HCl is a polar molecule and it contains polar bond.
4. The separation of charge in polar covalent bond like H-Cl is called polarisation.
5. When two electrical charge of opposite signs are separated by small distance, a dipole is established. Thus, HCl has a dipole because the H-Cl is a polar.
POLAR BOND
DIPOLE MOMENT (𝜇)
A quantitative measure of the polarity of a bond is its dipole moment ( µ ).
µ = Q r
Where : µ = dipole moment
Q = the product of the charge from
electronegativity
r = distance between the charges
Dipole moments are usually expressed in debye units(D)
1D (Debye) = 3.36 x 𝟏𝟎−𝟑𝟎Cm
RESULTANT (NET) DIPOLE MOMENT
1. Determined by molecular shape and bond polarity
Resultant dipole moment > 0 (polar molecule)
Resultant dipole moment = 0 (non-polar molecule)
2. Example :
∆𝐸𝑁 = 0 (charge in electronegativity)
𝜇 = 0
𝐼2 is a non polar molecule
I - I
RESULTANT (NET) DIPOLE MOMENT
3. Example :
∆𝐸𝑁 ≠ 0 (charge in electronegativity)
𝜇 ≠ 0
𝐼2 is a polar molecule
CI - I
RESULTANT (NET) DIPOLE MOMENT
4. Example :
𝐶𝑂2 shape =linear
The two bond dipoles cancel each other (molecule is symmetrical)
Resultant dipole moment, 𝜇 = 0
𝐶𝑂2 is a non polar molecule
𝐶𝑂2 𝑂 = 𝐶 = 𝑂
RESULTANT (NET) DIPOLE MOMENT
5. Example :
𝑂𝐶𝑆 shape =linear
The two bond dipoles do not cancel each other
Resultant dipole moment, 𝜇 ≠ 0
𝑂𝐶𝑆 is a polar molecule
OCS 𝑂 = 𝐶 = 𝑆
RESULTANT (NET) DIPOLE MOMENT
6. Example :
𝐵𝐹3 shape =trigonal planar
The two bond dipoles cancel each other
Resultant dipole moment, 𝜇 = 0
𝐵𝐹3 is a NON polar molecule
𝐵𝐹3
B
F
F
F
RESULTANT (NET) DIPOLE MOMENT
7. Example :
𝐵𝐹3𝐵𝑟 shape =trigonal planar
The three bond dipoles do not cancel each other
Resultant dipole moment, 𝜇 ≠ 0
𝐵𝐹3𝐵𝑟 is a polar molecule
𝐵𝐹2Br
B
Br
F
F
Carbon tetrachloride, CCl4
- molecular geometry : tetrahedral
- Chlorine is more electronegative than carbon,
- Dipole moment can cancel each other
- has no net dipole moment (µ = 0)
- therefore CCl4 is a nonpolar molecule.
- molecular geometry : tetrahedral
- Cl is more electronegative than C, C is more
electronegative than H
- Dipole moment cannot cancel each other
- has a net dipole moment (µ ≠ 0)
- therefore CH3Cl is a polar molecule.
Chloromethane, CH3Cl
Ammonia, NH3
- molecular geometry : tetrahedral
- N is more electronegative than H,
- Dipole moment cannot cancel each other
- has a net dipole moment (µ ≠ 0)
- therefore NH3 is a polar molecule.
BOND MOMENTS AND RESULTANT DIPOLE MOMENTS
H
1. Polar molecules possess polar bond.
2. A bond is polar when the two atoms that are participating
in the bond formation have different electronegativities. In
polar molecule, all the bonds collectively should produce a
polarity.
3. Though a molecule has polar bonds, it does not make the
molecule polar.
4. If the molecule is symmetric and all the bonds are similar,
then the molecule may become non polar.
5. Therefore, not all the molecules with polar bonds are polar.
DIFFERENCE BETWEEN POLAR BOND AND POLAR MOLECULES
The presence of polar bonds does not always lead to a
polar molecule
C−O is a polar bond
But, 𝐶𝑂2 is a non polar molecule
We have to CONSIDER BOTH (bond polarity and
molecular shape)
KEEP IN MIND!
𝑂 = 𝐶 = 𝑂
1. A molecule will be nonpolar if :
a) The bonds are non-polar
b) No lone pair in the central atom and all the surrounding atoms are the same
A SUMMARY ON HOW TO DETERMINE MOLECULAR POLARITY
CI – CI (non polar)
B
Br
F
F
B
F
F
F
POLAR NON POLAR
1. A molecule will be nonpolar if :
c) A molecule in which the central atom has lone pair electron will usually be polar with few exceptions
A SUMMARY ON HOW TO DETERMINE MOLECULAR POLARITY
POLAR
O
H H
N
H H H POLAR
1. A molecule will be nonpolar if :
c) A molecule in which the central atom has lone pair electron will usually be polar with few exceptions
A SUMMARY ON HOW TO DETERMINE MOLECULAR POLARITY
POLAR POLAR
Br
F
F
F
S
F
F
F
F
2. Exceptions : (NON POLAR MOLECULES)
A SUMMARY ON HOW TO DETERMINE MOLECULAR POLARITY
LINEAR
A
X
X
A
X X
X X
Xe
F F
F F
SQUARE PLANAR
60
SO2 ; HBr ; SO3 ; CH2Cl2 ; ClF3 ; CF4 ; H2O ; XeF4 ; NF3
Exercises :
Predict the polarity of the following molecules:
1. VSEPR theory : predict molecular shapes by assuming that electron groups tend to minimize their repulsions
2. But, it does not tell how those shapes (which is observed experimentally), can be explained from the interactions of atomic orbitals.
HYBRID ORBITAL OVERLAP AND HYBRIDIZATION
1. Covalent bonds are formed by sharing electrons from
overlapping atomic orbitals
2. Two types of bonds : σ bond and π bond
3. Example :
VALENCE BOND (VB) THEORY
1. Atoms in simple molecules or ions such as 𝐻2, HF, 𝑁2, etc. use pure s and/or p orbitals in forming covalent bonds.
2. Example : 𝐻2 (hydrogen molecules)
DIRECT OVERLAP OF s AND p ORBITAL
Example : HF (Hydrogen Flouride)
H = 1𝑠1
F = 1𝑠2 2𝑠2 2𝑝5
Example : 𝐹2 (Flourine molecules)
F = 1𝑠2 2𝑠2 2𝑝5
DIRECT OVERLAP OF s AND p ORBITAL
1. Mixing of two or more atomic orbitals to form a new
set of equivalent hybrid orbitals in the same energy
level
2. The spatial orientation of the new orbitals is cause
more stable bonds and are consistent with the
observed molecular shape.
3. 3 types of hybridization : sp, s𝑝2, s𝑝3 hybridization
HYBRIDIZATION
4.3.2 Formation Hybrid orbitals
• Overlapping of hybrid orbitals and the pure orbitals occur when different type of atoms are involved in the bonding.
• Hybridization of orbitals:
mixing of two or more atomic orbitals to form a new set of hybrid orbitals
• The purpose of hybridisation is to produce new orbitals which have equivalent energy
• Number of hybrid orbitals is equal to number of pure atomic orbitals used in the hybridization process.
Hybridization
• Hybrid orbitals have different shapes from
original atomic orbitals
• Types of hybridisation reflects the
shape/geometry of a molecule
• Only the central atoms will be involved in
hybridisation
HYBRIDIZATION
s orbital p orbital
sp orbital
TYPES OF HYBRID ORBITALS
Type Examples Electron group Electron group
arrangement
sp Be𝐶𝑙2 2 Linear
s𝑝2 B𝐹3 3 Trigonal planar
s𝑝3 C𝐻4 4 Tetrahedral
1. Draw Lewis structure
2. Predict electron group arrangement using VSEPR model
3. Deduce the hybridization of the central atom by matching
the arrangement of the electron groups with the hybrid
orbitals
4. Use partial the orbital diagram to explain the mixing of
atomic orbitals
DETERMINING HYBRID ORBITALS
Molecular formula
Lewis Structure
Molecular shape and electron group arrangement
Hybrid orbitals
1. Resulting from end to end overlap
2. Has highest electron density along the bond axis
3. Allow free rotation
4. All single bonds are 𝜎 bond
SIGMA (𝜎) BOND
73
+
bond
It formed when orbitals overlap from end to end
Example:
i. overlapping s orbitals
H H H H
bond
74
ii. Overlapping of s and p orbitals
H +
Px orbital
x x H
bond
75
iii. Overlapping of p orbitals
x x + x
bond
1. Resulting from side to side overlap
2. Has two regions of electron density
One above and one below the 𝜎 bond axis
3. One 𝜋 bond hold two electrons that move through both regions of the bond
4. 𝜋 bond restricts rotation
Pi (𝜋) Bond
bond It formed when two p-orbitals of the same orientation overlap
sideways
Double bond consists of one 𝜎 bond and one 𝜋 bond
Example :
y y
+
y y
bond
bond
Example :
𝐶𝑂2 has two 𝜋 bond and one 𝜎 bond
Triple bond always consists of one 𝝈 bond and two 𝝅 bond
𝑂 = 𝐶 = 𝑂 𝜋 𝜋
𝜎 𝜎
𝑁 ≡ 𝑁 𝜋
𝜎
𝜋
No of Lone Pairs +
No of Bonded Atoms Hybridization Examples
2
3
4
sp
sp2
sp3
BeCl2
BF3
CH4, NH3, H2O
How do I predict the hybridization of the central atom?
Count the number of lone pairs AND the number
of atoms bonded to the central atom
80
Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
Two equivalents sp hybrid orbitals that lie 180° apart
2 electron groups (from VSEPR theory)
Electron group arrangement : linear
Molecular shape : linear
FORMATION OF sp HYBRIDIZATION
sp sp
180°
Lewis structure
Valence electron configuration
Hybridisation
𝑩𝒆𝑪𝒍𝟐
FORMATION OF sp HYBRIDIZATION
2s 2p
O1 :
O2 :
2s 2p
Cl−𝐵𝑒 − 𝐶𝑙 𝜎 𝜎
Be ↑↓
↑ ↑
↑ ↑
sp orbitals
sp sp 2p 2p
Empty 2p orbitals
One s orbital + one p orbital → two equivalent sp orbitals
83
Three equivalents 𝑠𝑝2 hybrid orbitals that lie 120° apart
3 electron groups (from VSEPR theory)
Electron group arrangement : trigonal planar
Molecular shape : trigonal planar
FORMATION OF 𝑠𝑝2 HYBRIDIZATION
sp2
sp2
sp2
120°
Lewis structure
Valence electron configuration
Hybridisation
𝑩𝑭𝟑
FORMATION OF 𝑠𝑝2 HYBRIDIZATION
2s 2p
O1 :
O2 :
2s 2p
𝜎
𝜎
B ↑↓
↑ ↑
↑ ↑
𝑠𝑝2 orbitals
𝑠𝑝2 𝑠𝑝2 2p
Empty 2p orbitals
One s orbital + two p orbital → three equivalent 𝑠𝑝2 orbitals
B
Br
F
F
𝜎 ↑
↑
↑
𝑠𝑝2
TRY DRAW THEM 𝑠𝑝2 hybrid!!
Bond angle : 109.5°
4 electron groups (from VSEPR theory)
Electron group arrangement : tetrahedral
Molecular shape : tetrahedral
FORMATION OF 𝑠𝑝3 HYBRIDIZATION
sp3
sp3
sp3
sp3
109.5°
Lewis structure
Valence electron configuration
Hybridisation
𝑪𝑯𝟒
FORMATION OF 𝑠𝑝3 HYBRIDIZATION
2s 2p
O1 :
O2 :
2s 2p
𝜎
𝜎
B ↑↓
↑ ↑
↑ ↑
𝑠𝑝3 orbitals
𝑠𝑝3 𝑠𝑝3 𝑠𝑝3
Empty 2p orbitals
One s orbital + three p orbital → four equivalent 𝑠𝑝3 orbitals
𝜎
↑
↑
↑
𝑠𝑝3
↑
↑
↑
C H
H
H
H
𝜎
TRY DRAW THEM 𝑠𝑝3!!
Example : Methane, CH4
Ground state : C : 1s2 2s2 2p2
C H H H
H Lewis Structure
1s 2s 2p
Excitation: to have 4
unpaired electrons
Excited state : 1s 2s 2p
sp3 hybrid
shape: tetrahedral
sp3
sp3
sp3
sp3
H H
H
H
C
91
sp3-Hybridized C atom in CH4
sp3
sp3
sp3
sp3 1s
1s
1s
sp3 hybrid
• Mixing of s and three p orbitals
sp3
sp3
Other Example:
1. BF3
• Lewis structure :
• Valence Electron configuration :
F :
B ground state :
B hybrid :
• Molecular geometry
• Orbital overlap/Hybridisation:
Example: BF3
sp2
sp2 sp2
F : 1s22s22p5
Shape : trigonal planar
Pure p orbital
Other Example :
2. NH3
Lewis structure :
Valence orbital diagram ;
H :
N ground state :
N hybrid :
Orbital Overlap :
Molecular Geometry :
96
sp3
sp3
sp3
sp3
1s
1s
1s
97
Other Example:
3. H2O
• Lewis structure :
• Valence orbital diagram;
O ground state :
O hybrid :
• Molecular geometry
• Orbitals overlap:
HYBRIDISATION IN MOLECULES
CONTAINING DOUBLE AND
TRIPLE BONDS
Lewis structure
(ETHANE)
Valence electron configuration
Hybridisation
𝑪𝟐𝑯𝟒
FORMATION OF HYBRIDIZATION
2s 2p
O1 :
O2 :
2s 2p
C ↑↓
↑ ↑
↑ ↑
𝑠𝑝2 orbitals
𝑠𝑝2 𝑠𝑝2 2𝑝
UNHYBRIDIZED 2p orbitals
One s orbital + two p orbital → three equivalent 𝑠𝑝2 orbitals
↑
↑
↑
𝑠𝑝2
↑
↑
↑
𝜎
𝜋
𝐶 = 𝐶
H H
H H
HOW TO DRAW ?
101
bond
bonds
102
Lewis structure
(ACETYLENE)
Valence electron configuration
Hybridisation
𝑪𝟐𝑯𝟐
FORMATION OF HYBRIDIZATION
2s 2p
O1 :
O2 :
2s 2p
C ↑↓
↑ ↑
↑ ↑
𝑠𝑝 orbitals
𝑠𝑝 2𝑝 2𝑝
UNHYBRIDIZED 2p orbitals
One s orbital + one p orbital → two equivalent 𝑠𝑝 orbitals
↑
↑
↑
𝑠𝑝
↑
↑
↑
𝜎
𝜋
𝐶 ≡ 𝐶 H H
𝜋
HOW TO DRAW ?
105
Lewis structure
(BENZENE)
Valence electron configuration
Hybridisation
𝑪𝟔𝑯𝟔
FORMATION OF HYBRIDIZATION
2s 2p
O1 :
O2 :
2s 2p
C ↑↓
↑ ↑
↑ ↑
𝑠𝑝2 orbitals
𝑠𝑝2 𝑠𝑝2 2𝑝
UNHYBRIDIZED 2p orbitals
One s orbital + two p orbital → three equivalent 𝑠𝑝2 orbitals
↑
↑
↑
𝑠𝑝2
↑
↑
↑
BENZENE ???? (Look at the notes!)
108
ANSWER:
For each of the following, draw the orbital
overlap to show the formation of covalent bond
a) 𝐻2𝑂
b) N2
c) 𝐶𝐻3 Cl d) Al𝐶𝑙3
INERTNESS OF NITROGEN
MOLECULE 1. Nitrogen is a very electronegative element.
2. It is an inert (unreactive) element.
3. Inertness due to 2 factors :
a) Strong triple bond
b) Non polarity of 𝑁2
4. Bond energy 𝑁 ≡ 𝑁 is very high due to triple bond.
5. This strong bond must be broken first, then it can form with other
compounds.
6. A lot of energy needed to break
7. Only at high temperature, nitrogen can react with other elements
to form compounds.
8. Nitrogen molecules is non-polar. The absence of polarity on the
molecule explains why nitrogen is unreactive.
COVALENT CHARACTER IN IONIC
COMPOUNDS
1. Not all compounds are ionic, and not all compounds are covalent.
2. Polarisation of chemical bonds occur not only in covalent
compound but also in ionic compound.
3. Most ionic compounds have covalent character due to incomplete
transfer of electrons.
4. If a small cation with high electric charge approaches a large
anion, the cation will attract the electron cloud from the negative
ion.
5. Polarisation : The distortion of electron cloud of the anion by a
neighbouring cation.
6. Polarising power : The extent to which a cation (positive ion) can
polarise an anion (negative ion)
COVALENT CHARACTER IN IONIC
COMPOUNDS
7. Polarisation (distortion of electron cloud) of the anion produces
covalent character in the ionic bond because the valence
electron is partially shared between the cation and anion.
8. The greater the degree of polarisation of the anion, the greater
the amount of covalent character in the ionic bond.
9. The polarising power of a cation towards an anion is
proportional to the charge density.
𝑐ℎ𝑎𝑟𝑔𝑒 𝑑𝑒𝑛𝑠𝑖𝑡𝑦 = 𝑐ℎ𝑎𝑟𝑔𝑒
𝑖𝑜𝑛𝑖𝑐 𝑟𝑎𝑑𝑖𝑢𝑠 (𝑠𝑖𝑧𝑒)
10. Small and highly charged cations such as 𝐿𝑖+ and 𝐴𝑙3+ have
high charge density, so, high polarising power.
COVALENT CHARACTER IN IONIC
COMPOUNDS
7. Polarisation (distortion of electron cloud) of the anion produces
covalent character in the ionic bond because the valence
electron is partially shared between the cation and anion.
8. The greater the degree of polarisation of the anion, the greater
the amount of covalent character in the ionic bond.
9. The polarising power of a cation towards an anion is
proportional to the charge density.
𝑐ℎ𝑎𝑟𝑔𝑒 𝑑𝑒𝑛𝑠𝑖𝑡𝑦 = 𝑐ℎ𝑎𝑟𝑔𝑒
𝑖𝑜𝑛𝑖𝑐 𝑟𝑎𝑑𝑖𝑢𝑠 (𝑠𝑖𝑧𝑒)
10. Small and highly charged cations such as 𝐿𝑖+ and 𝐴𝑙3+ have
high charge density, so, high polarising power.
Question 1
Arrange the following chlorides in order of increasing covalent
character. Explain your answer.
NaCl, 𝑀𝑔𝐶𝑙2 , 𝐴𝑙𝐶𝑙3
Answer 1
NaCl < 𝑀𝑔𝐶𝑙2 < 𝐴𝑙𝐶𝑙3
The charge density of the cations increases in the order :
𝑁𝑎+< 𝑀𝑔2+ < 𝐴𝑙3+
Hence, polarising power of the cations towards the 𝐶𝑙_ ion increases
in the same order.
Polarising power of 𝐴𝑙3+ is high so, it is a covalent compound.
Question 2
Arrange two compounds of beryllium order of increasing covalent
character.
𝐵𝑒𝐹2 , 𝐵𝑒𝐼2
ANSWER 2
𝐵𝑒𝐼2 is expected to have a high degree of covalent character.
The size of 𝐼− ion is larger than 𝐹− ion, hence it is easier to be
polarised by the 𝐵𝑒2+ ion