Chemical bonding and molecular structure

Chemical Bonding & Molecular structures Page 1
UNIT 4
CHEMICAL BONDING AND MOLECULAR STRUCTURE
CHEMICAL BOND: The force which binds the two atoms together within a molecule is called chemical
bond.
Cause of formation:
 Attain the octet (Lewis-Kossel approach)
Atoms of different elements combine with each other in order to complete their respective octets
(i.e. 8 electrons in the outermost shell).
Note:
Contraction of octet: Central atom is electron deficient or does not complete its octet.
eg.
BeX2, BX3, AlX3
4 6 6
Expansion of Octet: Central atom has more than 8 electrons due to empty d orbital.
eg.
PCl5, SF6
10 12
Other exceptions:
a) Transition metal ions (Cr3+
, Fe2+
)
b) Pseudo inert gas configuration (Zn2+
)
c) Odd electronic species (NO, NO2)
d) Inter halogen compound (IF7, BrF3)
e) Compounds of Xenon (XeF2, XeF4)
 Minimize energy
 Gain stability
 Decrease reactivity
TYPES OF BONDS (Ranked in decreasing bond strengths)
1) Ionic Bond
2) Covalent bond
3) Coordinate bond
4) Hydrogen Bond
5) Van der wall bond

IONIC BOND
The bond formed by the complete transfer of one or more electron from the atom of metal to an
atom of Non-metal.
Remember: Ionic bond is a Non-directional bond.
Conditions for the formation of Ionic bond
 Metal must have low ionization energy.
 Non-metal must have low reactivity.
 Ions must have high electron affinity.
 Ions must have high lattice energy.
 Cation should be large with low electron negativity.
 Anion must be small with high electron negativity.
Some Important properties of ionic compound
i. Ionic compounds are hard.
ii. Ionic compound has high value of melting Point, boiling point & density.
iii. Conductor in fused, molten and aqueous state.
iv. Ionic compound show crystalline structure.
v. The energy released when the requisite number of gaseous positive or negative ion
combine to form 1 mole of ionic compound is called lattice enthalpy.
Add to your Knowledge:
In case of Univalent or Bivalent compounds, Lattice energy decrease as follows
e.g.
vi. Ionic Compounds are soluble in Polar solvents like water due to high dielectric constant
of these solvents.
Add to your knowledge:
Remember:
 Hydration enthalpy > lattice energy (Ionic compound is soluble)
 Hydration enthalpy < lattice energy (Ionic compound is Insoluble)
 Hydration enthalpy = lattice energy (Ionic compound is In equilibrium)

COVALENT BOND
The bond formed between two atoms by mutual sharing of electrons between them.
Remember:
 One atom can share max 3 electrons with other atom.
 When two atoms share one electron pair they are said to be joined by single covalent
bond, if two atoms share two pairs of electrons, the covalent bond between them is called
double bond & if shares three electron pairs then a triple bond is formed.
 The number of electrons contributed by each atom is known as co-valency.
Conditions for the formation of covalent bond:
 High ionization enthalpy of the combining elements.
 High nuclear charge & small atomic size of the combining elements.
 Nearly equal electron gain enthalpy & equal electro-negativities of combining elements.
Add to your Knowledge
The bond between two unlike atoms which differ in their electro-negativity is said to be polar
covalent bond.
Some important properties of covalent compound:
i. Solubility of covalent compounds follows the concept ‘like dissolve like’, i.e. non polar
solute dissolve in non-polar solvent & polar solute dissolve in polar solvent.
ii. Covalent compounds are non-conductors (exception- graphite).
iii. Covalent compounds are directional in covalent bond.
LEWIS STRUCTURE:
The structure which shows how valance electrons are distributed in a molecule is called lewis
structure.
Lewis dot structure of CO3
2- ion
Step-1: Total number of valance electrons of CO3
4+3x6=22 (6C- 2,4; 8O- 2,6)
Step-2: Total number of electrons to be distributed in CO3
2-
22+2=24 (for 2 unit negative charge)
Step-3: Draw the skeletal structure of CO3
O
O C O
Step-4: Put one shared pair of electrons between each C & O & complete the octet of oxygen
O
O C O
Step-5: Octet of C is not complete; hence, multiple bond is required between 1 C & 1O.
O
O C O
2-

FORMAL CHARGE:
FC = Formal Charge on an atom in a lewis structure
V = total number of valance electrons in the free atom
L = total number of non-bonding (lone pair) electrons
S = total number of bonding (shared) electrons
COORDINATE/DATIVE BOND:
The bond formed by donation of electron pair from one atom to another is called coordinate
bond.
Lewis Lewis
base acid
e.g.
Remember: There will be no change in hybridization of any of the two atoms during
formation of coordinate bond.
BOND PARAMETERS
BOND LENGTH:
Average distance between the centers of the nuclei of the two bonded atoms is called bond
length.
 For ionic compound, sum of ionic radius of cation & anion
 For covalent compound, sum of covalent radius
is considered as bond length.
Factors affecting bond lengths:
i.
ii. eg.
iii. eg. Sp3>sp2>sp
BOND ORDER:
In the Lewis representation of a molecule or ion, the number of bonds present between two
atoms is called bond order.
BOND ENERGY OR ENTHALPY:
The amount of energy required to break one mole of bonds of a particular type between two
atoms in a gaseous state.
Factors affecting bond energy:

BOND ANGLE:
The angle between the lines representing the directions of the bond, i.e. the orbitals containing
the bonding electrons is called the bond angle.
Factors affecting bond energy:
i.
ii. Bond angle is also affected by the electronic repulsion
DIPOLE MOMENT
It is the product of the magnitude of the charge & the distance between the centers of the positive
& negative charge.
Remember:
 It is used to measure the polarity in a molecule.
 It is represented by an arrow with its tail at the positive centre & head pointing towards a
negative centre.(electropositive to electronegative species or less electronegative to more
electronegative species).
Factor affecting dipole moment:
i. eg. HF>HCl>HBr
ii. eg.
iii. In case of polyatomic molecules the dipole moments depends on the spatial arrangement
on various bonds in the molecule. In such case, the dipole moment of a molecule is the
vector sum of the dipole moments of various bonds.
F
F Be F F B
Zero dipole moment F
Zero dipole moment
iv.
1)
2)
3)
N N
H F
H F
H F
The resultant dipole moment of NH3 is greater than that of NF3

-
+
Reason:
In case of NH3 the orbital dipole due to lone pair is in the same direction as the resultant dipole
moment of the N – H bonds, whereas in NF3 the orbital dipole is in the direction opposite to the
resultant dipole moment of the three N–F bonds.
RESONANCE STRUCTURE
Whenever a single Lewis structure cannot describe a molecule accurately, a number of structures
with similar energy, positions of nuclei, bonding and non-bonding pairs of electrons are taken as
the canonical structures of the hybrid which describes the molecule accurately. Such canonical
structures are called resonance structure.
It is observed due to delocalization of  electron.
Stability of Resonance structure:
 A Non-polar structure is always more stable than a polar structure.
 Greater the number of covalent bonds greater will be the stability.
CH3 C O > CH3 C O
 The resonance structure in which positive charge on electropositive atom & negative
charge on electronegative atom is more stable.
R R
C O > C O
R R
VALANCE SHELL ELECTRON PAIR REPULSION THEORY (VSEPR THEORY)
 The shape of a molecule depends upon the number of valence shell electron pairs
(bonded or non-bonded) around the central atom.
 Pairs of electrons in the valence shell repel one another since their electron clouds are
negatively charged.
 These pairs of electrons tend to occupy such positions in space that minimize repulsion
and thus maximise distance between them.
 The repulsive interaction of electron pairs decrease in the order:
lp-lp>lp-bp>bp-bp
Calculation of total number of electron pairs, bond pairs & lone pairs and predicting the
shapes of the molecules & Ions
i. Total no. of electron pairs
(No. of Valance electrons of central atom+ No. of mono-valent atom bonded to central atom)
ii. No. of bond pairs (shared Pairs) = No. of mono-valent atom bonded to central atom
iii. No. of Lone Pairs = Total no. of electron pairs - No. of bond pairs
+ +
+ - - +

No. of
bond Pair
2 3 2 4 3 2 5 4 3 2
No. of
lone Pair
0 0 1 0 1 2 0 1 2 3
Molecular
Formula
AB2 AB3 AB2L AB4 AB3L AB2L2 AB5 AB4L AB3L2 AB2L3
Geometry Linear
Triangular
Planar
Bent
(V-
Shape)
Tetrahe
dral
Trigonal
Pyramidal
Bent
Trigonal
Bipyrami
dal
See
saw
T-
shape
Linear
No. of
bond Pair
6 5 4
No. of
lone Pair
0 1 2
Molecular
Formula
AB6 AB5L AB4L2
Geometry Octahedral
Square
Pyramidal
Square
Planar
Bond angles of all geometric shapes:
Geometry Linear Trigonal Planar Tetrahedral Trigonal Bipyramidal Octahedral
Bond Angle 180 120 109.5 90 & 120 90
Remember:
 The shape should have been tetrahedral if there were bp in place of lone pair & due to the
repulsion between lp-bp (which is more than bp-bp repulsion) the angle between bond
pairs is reduced to 107 from 109.5.
N 107 O
H H H 104.5 H
H
 The shape should have been tetrahedral if there were all bp but two lp are present so the
shape is distorted tetrahedral or angular. The reason is lp-lp repulsion is more than lp-bp
repulsion which is more than bp-bp repulsion. Thus, the angle is reduced to 104.5° from
109.5°.
VALANCE BOND THEORY
Bond formation in Hydrogen molecule

 Consider two hydrogen atoms approaching each other
 When the two atoms are at large distance from each other, there is no interaction between
them.
 As these two atoms approach each other, new attractive and repulsive forces begin to
operate.
Attractive forces arises between
 the nucleus of one atom and its own electron
 nucleus of one atom and electron of other atom
Repulsive forces arise between
 Electrons of two atoms
 Nuclei of two atoms
Experimentally, it has been found that the magnitude of new attractive force is more than the
new repulsive force. As a Result two atoms approach each other.
 Ultimately a stage is reached where the net force of attraction balances the force of
repulsion and system acquires minimum energy.
 At this stage two hydrogen atoms are said to be bonded together to form a stable
molecule
eA eA
NA NB NA NB
eB
eB
Attractive Force Repulsive Force
Depending of type of overlapping, the covalent bonds are mainly of two types:
i. SIGMA (σ) BOND
When a bond is formed between two atoms by the overlap of their atomic orbitals along
the inter-nuclear axis (end to end or head on overlap), the bond formed is called sigma
(σ) bond.
i. s-s overlapping
+
1s orbitals Molecular orbital
of H-atom of H2 Molecule
ii. s-p overlapping
+
1s orbital 2pz orbital Molecular orbital

iii. p-p overlapping
+
pz orbital pz orbital Molecular orbital
ii. PI () BOND
Pi-bond is formed by lateral (sideways) overlapping of p-orbitals, i.e., by overlapping of
p-orbitals in the direction at right angles to the inter-nuclear axis.
+
p orbital p orbital Molecular orbital
S. No. Sigma (σ) bond Pi () bond
1
It is formed by the end to end overlap
of orbitals.
It is formed by sidewise overlap of orbitals.
2
The orbitals involved in the
overlapping are s-s, s-p or p-p.
These bonds are formed by the overlap of p-p
orbitals only.
3 It is a strong bond. It is a weak bond.
4
The electron cloud is symmetrical
about the line joining the two nuclei.
The electron cloud is not symmetrical.
5
It consists of one electron cloud,
which is symmetrical about the inter-
nuclear axis.
There are two electron clouds lying above or
below the plane of the atomic nuclei.
6
Free rotation about σ bonds is not
possible.
Rotation is restricted in case of -bonds.
HYBRIDIZATION
Intermixing or redistribution of energy among two or more half-filled, fully-filled, incompletely
filled or empty orbitals of comparable energy forms new orbital of comparable energies &
identical shapes. The new orbital thus formed are known a hybrid orbital.
FINDING HYBRIDIZATION
X = Number of Hybrid Orbitals
VE = No. of Valance electrons of the central atom
MA = No. of mono-valent atoms/ groups surrounding the central atom
c = Charge on cation if the given species is a polyatomic ion
a = Charge on anion if the given species is a polyatomic ion

Value of X Type of
Hybridization
Shape of the molecule Angle Example
2 sp Linear 180 BeCl2, C2H2
3 sp2
Triangular planer 120 BF3, C2H4
4 sp3
Tetrahedral 109.5 CH4, NH4
+
4 dsp2
Square Planer 90 [Ni(CN)4]2-
5 sp3
d Triangular bipyramidal 90 & 120 PCl5
6 sp3
d2
Octahederal 90 SF6
7 sp3
d3
Pentagonal bipramidal 90 & 72 IF7
Axial
Bond
Equitorial
bond
Triangular Planer Square Planner Triangular Bipyramidal
Octahedral Pentagonal
Bipyramidal
Remember: In PCl5, The three P-Cl bond lie in one plane & making an angle of 120 with each
other; these bonds are termed as equatorial bonds. The remaining two P-Cl bonds-one lying
above & the other lying below the equatorial plane make an angle of 90 with the plane. These
bonds are called axial bonds. As the axial bonds suffer more repulsion from equatorial bond,
therefore axial bonds are slightly longer & weaker than equatorial bonds.
MOLECULAR ORBITAL THEORY
(Based on LCAO- Linear combination of atomic orbitals Model)
i. Atomic orbitals undergo linear combination to form same number of molecular orbitals,
if they fulfill the following conditions
a. Atomic orbital must have comparable energies.
b. Atomic orbitals must overlap linearly for effective overlapping.
c. Atomic orbitals must have same symmetry along with the major molecular axis.

E
N
E
R
G
Y
E
N
E
R
G
Y
B.O. Decreases
ii. Molecular orbitals are formed due to constructive & destructive interference of atomic
orbitals.
iii. Constructive interaction of orbital lobes having same function  produces bonding MOs
like σ & , these are HoMOs (Highest occupied MOs)
iv. Destructive interaction of orbital lobes having different sign of  produces antibonding
MOs (LuMOs-Lowest unoccupied MOs) like σ*
& *
.
Facts related to HoMOs & LuMOs
i. Energy: LuMOs > HoMOs
ii. Wavelength: LuMOs < HoMOs
iii. Like atomic orbitals, MOs also follow Pauli Exclusion Principle, Hund’s Rule, Aufbau
Principle.
Remember:
 The MO obtained by addition of atomic orbitals is of lower energy than that of the atomic
orbitals itself & is called bonding orbital.
 The MO obtained by subtraction of atomic orbitals is of higher energy than that of the
atomic orbitals itself & is called anti-bonding orbital.
 Bond order (B.O.)
 Bond order of 1 is equivalent to single bond.
 Bond order of 2 is equivalent to double bond.
 Bond order of 3 is equivalent to triple bond.
 Bond order of 0 means no bond exist between the atoms of the molecule.
Total No. of
electrons
10 11 12 13 14 15 16 17 18
Bond order 1 1.5 2 2.5 3 2.5 2 1 1
H2 MOLECULE: anti-bonding MO
He2 MOLECULE:
1s 1s
σ1s
σ*1s
Atomic orbital
bonding MO
1s 1s
σ1s
σ*1s
He2 does not exist as B.O. is zero.
B.O. Decreases

2s
N2 will form triple bond.
1s
FOR THE MOLECULE LIKE B2, C2 & N2
N2 Molecule
N 1s2
2s2
2p3
Bond order
As all the electrons present are paired in MO
N2 is diamagnetic
Add to your knowledge:
There is no need to form all the MOs
1s, 2s & 2p. 2p alone is sufficient.
For the molecule other than B2, C2 & N2
O2 molecule
O 1s2
2s2
2p4
Bond Order
O2 will form double bond.
As the MOs are singly occupied by the electrons, O2 is paramagnetic though it contain even number of
electrons.
 Greater the bond order, the more stable the molecule or ion and shorter will be the bond
length & more will be the bond energy.
σ*2pz
2s
σ*2s
2p
σ2pz
σ*2pz
2p
π2px π2py
π*2py
π*2px
2p
2p
π*2py
π*2px
σ2pz
π2py
π2px
E
N
E
R
G
Y
E
N
E
R
G
Y
σ2s
1s 1s
σ*1s
σ1s
2s

HYDROGEN BOND
The bond between the hydrogen atom of one molecule & a highly electro-negative element of
same or another molecule is called as hydrogen bond.
Remember:
 This bond is weaker than covalent bond.
 This bond can only be formed by the smaller size electro-negative elements like F, O &
N.
 This type of bond is formed only in polar molecule in which H acquires partial positive
charge & the atom (which is covalently bonded with H atom) acquires partial negative
charge.
TYPES OF H-BONDS
1) Inter-molecular hydrogen bond
2) Intra-molecular hydrogen bond
Inter-Molecular Hydrogen bond
It is formed between two different molecules of the same or different compounds.
e.g. HF molecule, alcohol or water molecules, etc.
+δ - δ + δ - δ + δ - δ
H▬X H▬X H▬X
Covalent Bond Hydrogen Bond
Consequences:
 HF is a liquid having higher B.P.
 Alcohols are highly soluble in water & have high B.P.
Inter-molecular hydrogen bonding cause higher B.P., M.P., solubility, thermal stability,
viscosity, surface tension & occurrence of liquid state.
Intra-Molecular Hydrogen bond
It is formed when hydrogen atom is in between the two highly electronegative (F, O & N) atoms
present within the same molecule.
e.g. o-nitrophenol
Hydrogen bond
Intra-molecular hydrogen bonding cause lower B.P. & acidic nature of the molecule. But
increase the volatile nature.
O
N
O
O
H

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