Chemical Bonding - TestBag

Chemical Bonding
Chemical Bonding
Chemical bond.
A chemical bond is the force which holds the atoms of a molecule together. It may be of following type :
(a) Ionic (b) Covalent (c) Coordinate and (d) Metallic. In addition, attractive interactions between atoms
of different molecules give weak bonds of the following type: (i) Hydrogen bond and (ii) Vander Waal’s
interactions.
lonic or Electrovalent bond.
lonic or Electrovalent bond is formed by complete transfer of one or more electrons from the valence
shell of one atom to the valence shell of the other. The atom that loses electrons gets converted into a cation
while the other which gains electrons gets converted into an anion.
. .
Na.
+ . Cl : → Na [: Cl :]
. .
+
..
..
−
or
Na
+ Cl −
The strong electrostatic force of attraction between these oppositely charged ions is called lonic or
electrovalent bond and the number of electrons gained or lost by the atom is called its electrovalency.
Favourable conditions for forming a stable lonic bond are:
(1) Low ionization energy of the atom forming the cation.
(2) High electron affinity of the atom forming the anion.
(3) High lattice energy of the crystal formed.
Characteristics of ionic compounds are :
(1) Each ion is surrounded by a uniformly distributed electric field i.e., each ion is non- directional.
Therefore, ionic bond is also considered to be non directional. Hence ionic compounds do not show
stereoisomerism.
(2) These are highly soluble in solvent with high dielectric constant, such as water and other polar solvents, but
insoluble in non- polar solvents like benzene, ether, etc. (like dissolves like).
When an ionic compound dissolves in water, the ions are hydrated and the energy evolved in the process is
called hydration energy. Thus for an ionic solid to be soluble in water, the necessary condition is
∆H > H
Hydration energy
Lattice energy
Ionic compounds like CaCO 3 , Ca 3 ( PO 4 ) 2 , CaF2
, BaSO 4 , PbSO 4 , AgCl,
AgBr,
AgI etc., are insoluble in water
due to their high lattice energies.
(3) On account of strong electrostatic attraction between the ions in the crystal of an ionic compound, these
have high melting and boiling points. Order of melting and boiling points of certain compounds are examplified
below.
(a) NaF > NaCl > NaBr > NaI (b) MgO > CaO > BaO

Chemical Bonding
(4) Ionic compounds conduct electricity in solution as well as in molten form (fused state) because ions
become mobile. These free ions are able to move under the influence of an electric field and thus act as carriers of
current. In solid state, they are bad conductors as the ions are tightly held.
(5) In solution, ionic compounds show ionic reactions which are quite fast and instantaneous
Energetics of formation of ionic bond are represented by the Born-Haber cycle which is as followers :
∆H sub.
+ I.E.
Na (s)
+
1
Cl 2 ( g)
2
1
2
Na (g)
Cl(g)
∆H diss
−
− e + e ( −E.
A)
− U
(Lattice energy)
Na + (g)
Cl − (g)
Na
+
Cl
−
(s)
The heat of formation of an ionic solid is net resultant of the above changes.
1
∆ H f = ∆H
sub.
+ ∆H
diss + I.
E.
− E.
A.
− U
2
The lattice energy is one of the deciding factors of ionic bond formation. It depends upon :
(1) Size of the ions : Smaller the size of the ions greater is the attractive forces between the ions and
higher is the lattice energy.
(2) Charge on the ions : Greater the charge on ions greater will be attractive forces between the ions
and hence greater is the value of lattice energy
For example, lattice energy of bi-bivalent ions > bi-uni or uni-bivalent ions> uni-univalent ions. Order of
lattice energy of certain compounds are :
(a) LiX > NaX > KX > RbX > CsX (where X = F, Cl, Br or I); (size of alkali metal increases)
(b) MgO > CaO > SrO > BaO (size of alkaline earth metal increases)
(c) MgCO 3 > CaCO 3 > SrCO 3 > BaCO 3
(size of cation increases)
(d) Mg ( OH ) 2 > Ca(
OH ) 2 > Sr(
OH ) 2 > Ba(
OH ) 2
(size of cation increases)
As lattice energy increases, melting and boiling points of ionic compounds increases. This is due to strong
electrostatic force of attraction.
Fajan Rules.
In Ionic bond, some covalent character is introduced because of the tendency of the cation to polarise the
anion. In fact cation attracts the electron cloud of the anion and pulls electron density between the two nuclei.
+
–
+
–
CATION
ANION
According to Fajan rules, the magnitude of covalent character in the ionic bond depends upon the extent
of polarisation caused by cation. In general,
(1) Smaller the size of cation, large is its polarising power.
POLARISED
ELECTRON CLOUD
OF ANION

2
Chemical Bonding
(2) Among two cations of similar size, the polarising power of cation with noble gas configuration
6
10
( ns np nd ) is larger than cation with noble gas configuration ( ns ) e.g., Polarising power of
than
+
K
(3) Larger the anion, more will be its polarisability
Covalent bond.
2 np 6
+
Ag is more
Covalent bond is formed by mutual sharing of electrons so as to complete their octet or duplet in case of
H, Li and Be. Depending upon whether one, two and three electrons are shared by each atom, single, double
and triple bonds are respectively formed. The number of electrons contributed by each atom for sharing is
called its Covalency.
The formation of covalent bonds is best explained by a new approach according to which sharing of
orbitals is possible only with overlapping. According to this concept, the formation of covalent bond involves
the overlapping of half filled atomic orbitals of the atoms participating in bonding. The atomic orbitals
undergoing overlapping must have electrons with opposite spin.
When two half-filled atomic orbitals overlap along their internuclear axis, the bond formed is called the
sigma bond (or -bond) It may be formed by the overlap of two s-orbitals, two p-orbitals, one s- and one p-
orbital. Accordingly these bonds are dasignated as s-s, p-p and s-p sigma bonds. On the other hand, covalent
bond formed by sideways or lateral overlap of p-orbitals is called pi-bond (or –bond). It may be noted that
(1) All single bonds are –bonds.
(2) Multiple bonds contain one –bonds and the rest are –bonds.
(3) A –bond is never formed alone. First –bonds is formed and then the formation of the –bond takes
place.
(4) A sigma bond is always stronger than pi-bond because the extent of overlapping of atomic orbitals
along internuclear axis is greater than sideways overlapping.
Characteristics of covalent compounds are :
(1) Covalent compounds have low melting points and boiling points. As such they may be gases, liquids or
low melting solids.
(2) They are insoluble in water but soluble in organic (non-polar) solvents
(3) They do not conduct electricity in the solution or the molten state.
(4) Since covalent bonds are rigid and possess directional characteristics, therefore, they show
stereoisomerism.
(5) Their reactions are slow and molecular in nature and never proceed to completion.
Coordinate or Dative bond.
This type of bond formation occurs by one sided sharing of electrons, i.e., one atom donates a pair of
electrons while the other simply shares it so as to complete its octet. The atom that donates a pair of electrons
is called the donor while the other which accepts these electrons is called the acceptor. The coordinate bond
is usually represented by an arrow pointing from the donor towards the accept

Chemical Bonding
For example,
H
|
+
H — N : → H or
|
H
⎡ H ⎤
⎢ | ⎥
⎢H — N — H ⎥
⎢ | ⎥
⎢
⎥
⎣ H ⎦
+
Coordinate bond is also present in SO 2 , SO 3,
O3,
H 3O
, NO 3 or HNO 3
etc.
This coordinate bond has some polar character, it is also called dative or semi-polar bond.
In terms of VB theory, a coordinate bond is formed by overlap of a fully filled orbital containing a lone
pair of electrons with an empty orbital of another atom.
Examples of molecules in which all the three types of bonds, i.e., ionic, covalent and coordinate bonds are
CuSO . 5H
O,
NH Cl,
K [ Fe CN ] Cu( NH ) SO etc. besides these bonds, CuSO 4 . 5H
2O
also
present : ( )
contains a H-bond.
4 2 4 4
6
,
Characteristics of coordinate compounds are :
[
3
]
4
Since coordinate compounds are in fact covalent compound, therefore, their properties are almost similar
to those of covalent compounds. For example,
(1) Like covalent compound, they are insoluble in H 2 O
(2) They usually do not conduct electricity.
4
+
−
but are soluble in organic solvents.
(3) Their melting and boiling points are higher than those of covalent compounds but lower than those of
ionic compounds.
(4) Like covalent bonds, coordinate bonds are directional and hence these compounds also exhibit
stereoisomerism
Bond formula or Dash formula of molecules showing different type of bonding.
(1) Compounds having electrovalent bonds only (2) Compounds having covalent bonds only
Molecular formula / Dash formula
NaCl /
2
Na
+ Cl −
−
MgCl / Cl Mg
2
−
CaCl / Cl Ca
+
+ +
+ +
MgO / Mg + O
+
Na 2 S / Na S
2
−
CaH / H Ca
AlF / F
3
−
−−
+ +
3+
Al F
F
−
Cl
Cl
−−
Na
H
−
−
+
−
−
Molecular formula / Dash formula
HCl /
H − Cl
H 2 O / H − O − H
H 2 S / H − S − H HCN / H − C ≡ N
NH 3
/
4
C / H − C ≡ C − H
2 H 2
H
|
H − N
|
H
H
|
CH / H − C − H
|
H
H H
C / C C
| |
2 H 4
|
H
H
C 2 H 6
/
H H
|
H − C − C − H
|
H
PCl 3
/
=
3
|
|
|
H
Cl
Cl − P
|
|
Cl
H
|
PH / H − P
|
H

Chemical Bonding
(3) Compounds having electrovalent & covalent
bonds
Molecular formula / Dash formula
NaOH / Na
KCN / K
+
+
[ O − H]
[ C ≡ N]
−
−
2−
⎡ ⎤
++
CaCO 3
/
⎢ ⎥
Ca
⎢
O − C−
O
|| ⎥
⎢⎣
O ⎥⎦
(4) Compounds having covalent & coordinate bonds
CO / C = O N 2 O / N ≡ N → O
N 2O 3
/ O = N − N = O
↓
O
HClO / H − O − Cl → O
3
O
H 2O 2
/ H − O − O − H or H O → O
−
2
−
NO / O − N = O
H
HNO / H − O − N = O
2
↓
2
−
|
SO / O ← S = O
N 2O 4
/ O = N − N = O
↓ ↓
O O
N 2 O 5
/ O = N − O − N = O
↓ ↓
O O
HNO 3
/ H − O − N = O
↓
O
− −
NO 3 / O − N = O SO 3
/ O ← S = O
↓
↓
O
O
SOCl 2
/
HIO 3
/ H − O − I → O
O ← S − Cl
↓
|
O
Cl
Cl
SO 2Cl 2
/ O S → O
←
|
|
Cl
O
↑
HClO 4
/ H − O − Cl → O
↓
O
O
↑
H 2SO 4
/ H − O − S − O − H
↓
O
Molecular formula / Dash formula
2−
SO 3 /
−
H
−
−
2SO 3
/ O − S − O
O − S − O
↓
↓
O
O
O
↑
2−
SO 4 / H − O − S − O − H
↓
O
O
H 3 PO 4
/ ↑
H − O − P − O −
|
H
O
|
H
O O
H 4 P2
O7
/ ↑ ↑
H − O − P − O − P − O − H
| |
O O
| |
H H
H
H / H − O − P − O − H
3 PO 3
Al Cl /
2 6
( Anhydrous )
Cl
Cl
O 3
/
O
Al
O
|
↓
O
Cl
(5) Compounds having electrovalent, covalent &
coordinate bonds
Cl
O
Al
Molecular formula / Dash formula
4k
⎡ H ⎤
⎢
|
⎥
NH 4 Cl / ⎢H
− N → H ⎥
|
⎢
⎥
⎣ H ⎦
K 4 [ Fe(
CN ) 6 ] /
+
⎡
⎢
⎢
N ≡ C
⎢
⎢
⎢N
≡ C
⎢
⎣
C ≡ N
Fe
C ≡ N
+
Cl
Cl
Cl
−
⎤
C ≡ N
C ≡ N
⎥
⎥
⎥
⎥
⎥
⎥
⎦
4 −
Note : Molecules like
CuSO 4 . 5H2O
[ Cu(
NH 3)
4 ] SO 4 also have all three types
of bonding.
−

Chemical Bonding
Hybridization.
The concept of hybridization was introduced to explain the shapes of molecules. It involves the
intermixing of two or more atomic orbitals of slightly different energies but of the same atom so that a
redistribution of energy takes place between them resulting in the formation of an equal number of new
orbitals (called hybrid orbitals) having same energy, size and shape. It may be noted that:
(1) Both half – filled and completely filled orbitals can participate in hybridization
(2) Hybridisation never takes place in isolated atoms but occurs only at the time of bond formation.
(3) Hybrid orbitals form stronger bonds than pure atomic orbitals.
Types of hybridization
Type of
Hybridisatio
n
Character
sp s -character = 50%,
p -character = 50%.
2
sp s -character = 33.33%,
p -character = 66.67%
3
sp s-character = 25%,
p-character = 75%
(as s-character
decreases,
p-character increases,
bond angle decreases).
sp 3 d
s-character = 20%,
p-character = 60%,
d-character = 20%
Shape of the
molecule or ion
180 o
: A
:
Linear
. .
A
120 o
: :
Triangular planar
109 o 28′
A
: :
:
. .
. .
Tetrahedral
120 o
. .
90 o :
A
. .
. .
Trigonal bipyramidal
Example
BeF2 , BeH 2 , CO 2 , C 2 H 2 ,
HCN HgCl , CS , N ,
, 2 2 2O
Hg 2Cl
2 , [ Ag(
NH 3 ) 2 ]
BF 3, SO 2
, SO 3, C 2 H 4 ,
2−
−
O
CO 3 , NO 3 , H − C − H,
+
CH 3 , AlCl 3, Benzene,
graphite, C 2Cl 4, C 2H 2Cl 2,
C 2 H 5CHO
, BeCl 2.
CH 4, CCl 4
, SiF 4, H 2O,
NH 3, H 3O + ,
+
||
−
ClO 4 ,
−
+
2−
SO 4 ,
NH 4 ,[
BeF4
] , ClO 3 , NF3
,
XeO 3 , XeO 4 , CH 3 , 2 ,
−
−
− NH −
[ AlCl4 ] ,[ PH4]
, SnCl 4,
Diamond, Silica,
+
+
H 3 O ,
AsCl 3, Si(CH 3) 4, SiC,
[BF 4] – , ClO −
2 ,
NH +
4 , SF − −
2
, ClO 4 , NH 2
.
XeO 3F 2, XeF 2, ClF 3, SF 4,
PCl 5, PF 5, AsF 5, SbCl 5,
PCl
4
+ −
, PCl6
.

Chemical Bonding
sp 3 d 2 s-character = 16.66%,
p-character = 49.98%,
d-character = 33.33%
:
. .
A
90 o
:
SF 6 , MoF 6, XeF 4 , BrF 5,
XeOF 4, [BiCl 6] – , [PF 6] – ,
[Co(NH 3) 6] 3+ .
:
. .
:
Octahedral
sp 3 d 3 s-character = 14.28%,
p-character = 42.86%,
90 o A
: 90 o
. . . .
IF 7, XeF 6, [ZrF 7] 3– ,
[UF 7] 3– , [UO 2F 5] 3– etc.
d-character = 42.86%
72 o
:
. .
. .
Pentagonal bipyramidal
Method of predicting the hybrid state of the central atom in covalent molecules or polyatomic
ions
The hybrid state of the central atom in simple covalent molecule or polyatomic ion can be predicted by
using the generalised formula as described below:
Simple molecule Polyatomic Anion Polyatomic Cation
:
[ V G ]
1
X = + 2
[ V + G a]
1
X = + 2
X
= 2
1
[ V + G − c]
In the above formulae.
V = Number of monovalent atoms or groups attached to the central atom
G = Number of outer shell electrons in ground state of the central atom
a = Magnitude of charge on anion, c = Magnitude of charge on cation
Calculate the value of X and decide the hybrid state of central atom as follows:
X 2 3 4 5 6 7
Hybrid
state
sp 2
sp
3
sp
Examples : (1) Central atom is surrounded by monovalent atoms only
sp 3 d sp 3 d
2
BeF2 , BCl 3 , CCl 4 , NCl 3 , PCl 3 , PCl 5 , NH 3 , TeCl 4 , OF2
, H 2O,
SCl 2 , IF7
, ClF3
, SF4
, SF6
, XeF2
, XeF4
.
Shape and Structure of ClF3
Cl has seven electrons in Valence-shell and there are three monovalent atoms surrounding it.
1
H = [7 + 3 − 0 + 0] = 5
∴ Hybridisation = sp 3 d
2
(2) Central atom is surrounded by divalent atoms only – SO 2 , SO 3 , CO 2 , CS 2 , XeO 3
.
sp
3 d 3

Chemical Bonding
Hybridisation in SO 2
H =
1
[6
2
+ 0 − 0 + 0] = 3
∴ Hybridisation
2
= sp
(3) Central atom is surrounded by monovalent as well as divalent atoms – COCl 2 , XeO 2 F2
, POCl 3
.
Hybridisation in COCl 2
H =
1
[4
2
+ 2 − 0 + 0] = 3
(4) Hybridisation in anions –
∴ Hybridisation
2−
2−
2−
−
4 CO 3 , PO 4 , NO 2 ,
−
3
SO , NO .
2
= sp
Hybridisation in
H =
1
[6
2
+ 0 − 0 + 2] = 4
2−
SO 4 (Charge = 2)
∴ Hybridisation
3
= sp
Hybridisation in
−
NO 2
H
=
1
[5
2
+ 0 − 0 + 1] = 3
∴ Hybridisation
2
= sp
(5) Hybridisation in cations –
+
+
NH 4 , H 3O
.
Hybridisation in
+
NH 4 (Charge = 1)
H =
1
[5
2
+ 4 − 1 + 0] = 4
(6) Hybridisation in complex ions –
ICl
− −
4 I3
,
, ClF
∴ Hybridisation
−
2
3
= sp
Hybridisation in
−
ICl 4 (Charge = 1)
H
=
1
[7
2
+ 4 − 0 + 1] = 6
∴ Hybridisation
= sp
3 d 2
Hybridisation in I
− 3
1
H = [7 + 2 − 0 + 1] = 5
∴ Hybridisation = sp 3 d
2
(7) Hybridisation for complex ions like, [ Co(
NH 3 ) 6 ] , [ PtF6
] , [ Ni(
NH 3 ) 4 Cl 2 ] are given by counting
ligands in the co-ordination sphere.
Ligands 2 3 4 5 6
+ 2
2−
Hybridisation sp 2
sp
3
sp
or
2
dsp
sp 3 d
or
3
dsp
sp
3 d 2
or
d
2 sp 3

Valence shell electron pair repulsion (VSEPR) theory and shapes of molecules.
Chemical Bonding
This theory helps us in explaining the shapes of simple covalent molecules. According to this theory, the
electron pairs (bond pairs as well as lone pairs) present around the centra l atom repel each other and hence
move as far apart as possible so that there are no further repulsions between them. As a result, the molecule
has minimum energy and maximum stability. The direction of the electron pairs gives a definite geometry to
the molecule.
If the central atom in a molecule is surrounded by only one kind of atoms and bond pairs, the molecule is said
to possess regular geometry. If on the other hand, the central atom in a molecule is surrounded by both lone
pairs and bond pairs, the geometry of the molecule is distorted to some extent. Such molecules are said to
possess irregular geometry. This distortion is due to the reason that : Lone pair- Lone pair repulsion>Lone
pair-Bond pair repulsion > Bond pair-Bond pair repulsion (i.e., lp-lp > lp-bp > bp-bp). Depending upon the
number of bond pairs and lone pairs, the geometries of various types of molecules are summarized below:
Type
of
molecule
2
Total
no. of
electron
pairs
No.
of
bond
pairs
No.
of
lone
Pairs
Type of
hybridiza
t-ion
involved
Geometry of
molecule
Examples
AB 2 2 0 sp Linear BeF [ Ag ( ) ] +
2 , NH 3
2
AB 3 3 3 0
2
sp
Trigonal planar
BF
2− 2−
3 , AlCl 3 , NO 3 , CO 3
AB 2
L 3 2 1
AB 4 4 4 0
2
sp V-shaped SnCl 2 , PbCl 2
3
sp Tetraheral CH 4 , SiF4
, NH 4 , CCl 4
+
AB 3 L 4 3 1
3
sp
Trigonal
pyramidal
NH , PX 3 ( X = F,
Cl,
Br,
)
3 l
AB 2
L 2 4 2 2
3
sp
V-shaped
H
2 O OF2
, SCl 2 ,
, NH
−
2
AB 5 5 5 0
3
sp d
Trigonal
Bipyramidal
PF
5 , PCl 5 , SbCl 5 , Fe(
CO)
5
AB 4
L 5 4 1
AB 3 L 2 5 3 2
3
sp d See saw SF 5 , TeBr 4
3
sp d T-shaped ClF 3 , XeOF 2
AB 2 L 3 5 2 3 3
sp d
Linear
XeF
−
2 lCl2
,
, l
−
3
AB
3
6 6 6 0 sp
2
d Octahedral SF [ ] −
6 , SbF 6
AB 5 L 6 5 1
3
sp
2
d Square pyramidal lF5 , ClF5
,[ SbF 5
] , XeOF 4
2−
AB 4 L 2 6 4 2
3
sp
2
d
Square planar
SF
4 XeF4
,
, lCl
−
4
AB 7 7 7 0
3
sp
2
d
Pentagonal
bipyramidal
lF 7 , XeF 6

Chemical Bonding
Molecular Orbital (MO) Theory.
The main points of this theory are :
(1) When two atomic orbitals (AO’s) combi ne, they lose their identity and form new orbitals called
molecular orbitals (MO’s) ; one of which is bonding molecular orbital (BMO) while the other is
antibonding molecular orbital (ABMO).
Energy Level Diagrams.
There are two types of energy-level diagrams :
(1) For molecules upto N 2
i.e., B 2 C 2 and N 2
(Where the difference in energies between 2s and 2p-orbitals
is small and hence can interact) the order of filling of orbitals is
( 1s),
* (1 s)
(2s),
* (2s),
(2p
x ) = (2Py
), (2Pz
), * (2Px
) = * (2Py
) * (2Pz
).
(2) For molecules after N 2
i.e. O 2 , F 2
and NC 2
etc. (where the energy difference between 2s and 2p -
orbitals is large and hence cannot interact) the order of filling of orbitals is
( 1s),
* (1 s)
(2s),
* (2s),
(2p
z ) (2Px
) = (2Py
), * (2Px
) = * (2Py
) * (2Pz
).
Information obtained from MO Diagrams
1
(1) Bond order (B.O.)= ( N b − N a ) : Where N b
is equal to number of electrons in the bonding MO’s
2
and is number of electrons in the antibonding MO’s.
N a
Formation of molecules or molecular ions if bond order is greater than Zero, the molecule or ion exists
other wise not.
(2) Bond length : (i) A multiple bond (double or triple bond between two atoms) is always shorter than
the corresponding single bond.
(ii) Bond length decrease with the increase in s-character since a s- orbital is smaller than a p-orbital.
Thus
3
sp C − H = 1.093 Å
( as inalkanes )
2
sp C − H = 1.08 Å ,
( as inalkenes )
sp C − H = 1.057 Å
( as inalkynes )
(iii) Polar bond length is usually smaller than the theoretical non-polar bond length
(iv) For a given atom, the bond length increases with the size of the other atom bonded to it. For example,
HI > HBr > HCl > HF
(v) Since bond distance is the sum of the ionic or covalent (atomic) radii of the two concerned atoms, the
factors and trends observed in the ionic or atomic radii will apply on the bond distances.
(3) Bond energy or bond strength : It is defined as the amount of energy required to break one mole
of the bond (i.e., Avogadro’s number of bonds) and separate the bonded atoms in the gaseous state. It is also
known as the bond dissociation energy (D) of that particular bond
(i) Atomic size of bonded atoms : Since we know that bond distance ∝ Atomic size of the bonded
atoms. Hence smaller atoms form shorter bonds whose energy will be large. This is evident from the following
order of bond energies of the halogens,
Cl − Cl > Br − Br > I − I
Recall that the atomic size of the three halogens follows the order,
Cl < Br <
I

Chemical Bonding
(ii) Electronegativity of the bonded atoms (bond polarity) : Greater the electronegativity
difference, greater is the bond polarity and hence greater will be the bond strength, i.e., bond energy. This is
evident from the following order of bond energies of the different hydrogen halides,
H − F > H − Cl > H − Br > H − I
(iii) Extent of overlapping of atomic orbitals : A large extent of overlapping of the component
atomic orbitals imparts greater strength to the bond
(iv) Hybridisation : Hybrid orbitals form stronger bonds because they provide more extent of
overlapping than the pure atomic orbitals. Thus
than the p-p overlapping(i.e., bond).
3
3
sp − sp overlapping (i.e., bond) results in a stronger bond
(v) Percentage of s-character in a hybrid bond : The bond energy increases with the increase in the
percentage of s-character in a hybrid orbital. Thus bond energy increases in the following order.
% of s-character
3
sp < sp <
2
25 33.3
sp
50
(vi) Bond order : Since Bond energy ∝ Bond order, bond energy increases from a single bond to a triple
bond, i.e., C − C < C = C < C ≡ C
(vii) The bond energy decreases with the increases in number of lone pairs on the bonded
atom : For example, bond energies of the following single bonds having zero, one, two and three lone pairs of
electrons follow the following order,
C − C
81.6
>
⋅ ⋅
39
⋅ ⋅
N − N
⋅ ⋅
34.2
⋅ ⋅
> : O−
O : >
⋅ ⋅
⋅ ⋅
: F−
F :
⋅ ⋅ ⋅ ⋅
33.3 kcal mol
This is due to the presence of electrostatic repulsion between lone pairs of electrons on the two bonded
atoms.
(4) Magnetic properties : The MO diagram can predict whether the molecules or molecular ions are
paramagnetic if unpaired electron/s are present or diamagnetic if all the electrons are paired.
The MO diagrams, bond order and magnetic properties of some common molecules and their ions are
summarised below:
−1
Molecule/
ions
Molecular Orbital (MO)
Configuration
Bond order
Magnetic
character
H
2
2
( 1s)
+
H 2
−
H 2
+
He 2
2
( 1s)
1 Diamagnetic
1
( 1s)
2
2
1
* (1 s)
( 1s)
* (1 s)
1
2
1
2
1
2
Paramagnetic
Paramagnetic
Paramagnetic
2
Li KK ( 2s)
1 Diamagnetic
2 * 2
Be KK ( 2s)
( 2s)
0 (does not –
2

Chemical Bonding
exist)
B 2
C 2
N 2
+
N 2
KK ( 2s)
KK( 2s)
KK( 2s)
KK( 2s)
2
2
*
2
2
*
( 2s)
*
( 2s)
*
( 2s)
( 2s)
2
2
2
2
1
( 2Px
) (2Py
1
)
1 Paramagnetic
2
( 2Px
) (2Py
1
)
2 Diamagnetic
2
( 2Px
) (2Py
2
( 2Px
) (2Py
)
)
1
1
2
2P
)
3 Diamagnetic
( z
2
2P
)
2.5 Paramagnetic
( z
O 2
+
O 2
−
O 2
2−
O 2
F 2
KK ( 2s)
2
1
( 2Px
) (2Py
KK ( 2s)
2
*
( 2s)
)
1
*
( 2s)
1
( 2Px
) (2Py
KK ( 2s)
2
1
( 2Px
) (2Py
KK ( 2s)
1
( 2Px
) (2Py
2
KK ( 2s)
1
( 2Px
) (2Py
*
2
2P
)
( z
2
2
( 2Px
) (2Py
)
1
2
( 2s)
)
1
*
( 2s)
)
2
)
1
*
2
2P
)
( z
*
( 2Px
)
1
2P
)
( z
2
2
1
( 2Px
) (2Py
2
( z
2
2P
) a
2
* (2Px
) *(2Py
( 2s)
1
2
2P
)
( z
2
* (2Px
) *(2Py
2
)
)
1
1
)
)
2
2
2 Paramagnetic
2.5 Paramagnetic
1.5 Paramagnetic
1.0 Diamagnetic
1.0 Diamagnetic
Ne 2
KK ( 2s)
1
( 2Px
) *(2Py
2
*
( 2s)
)
2
*
( 2Pz
2
2P
)
( z
1
* (2Px
) (2Py
)
2
2
)
2
0
(does not exit)
Diamagnetic
CN
KK ( 2s)
2
*
( 2s)
2
1
1
( 2Px
) (2Py
) ( 2Pz
)
1
2.5 Paramagnetic
CN — 2
KK ( 2s)
( 2s)
*
2
1
( 2Px
) (2Py
)
1
2
2P
)
3.0 Diamagnetic
( z
NO
KK ( 2s)
2
*
( 2s)
2
1
( 2Px
) (2Py
* (2Px
)
1
)
1
2P
)
( z
2
2.5 Paramagnetic
+
NO
KK ( 2s)
2
*
( 2s)
2
1
( 2Px
) (2Py
)
1
2
2P
)
3.0 Diamagnetic
( z
CO
KK ( 2s)
2
* (2s)
2
1
( 2Px
) (2Py
)
1
2
2P
)
3.0 Diamagnetic
( z
2 *
Here KK represents non-bonding molecular orbitals of 1s-atomic orbitals, i.e., KK = (
1s)
(1 s)
2

Hydrogen bonds.
Chemical Bonding
When in a molecule, H atom is linked to a highly electronegative atom (F,O or N), this end becomes
negative while the H-end becomes positive. The negative end of one molecule is then attracted by the positive
end of the other. The bond thus formed is called hydrogen bond and is represented by a dotted line
+
−
+
−
+
−
..... H − F ..... H − F ........ H − F
Since hydrogen bond
8 to 42kJmol
−1
.
involves only a weak electrostatic attraction , therefore, its energy is only
The actual strength of any H-bond depends upon the electronegative difference between the
electronegative atom and the hydrogen atom. Thus, the strength of H-bond decreases in the order:
+
−
+
−
+
−
F − H ..... F > O − H ..... O > N − H ...... N
Hydrogen bond can be classified into the following two categories
(1) Intermolecular Hydrogen bond : This type of H-bond is formed between two molecules of the
same or different substance, i.e.,
H
| − +
H
| − +
H
| −
H — N ....... H — N ......... H — N …….
|
|
|
H
H
H
...... H
+ − + − + − + −
–
O ..... H –
O ....... H –
O ....... H –
Effects of hydrogen bonding : Intermolecular H-bonding increases the melting points, boiling points,
solubility, viscosity and surface tension while intermolecular H-bonding has opposite effects.
Metallic bond:
H R H H
(2) Intramolecular Hydrogen-bonding : This type of H-bonding occurs between H-atom and the
electronegative atom within the same molecule. For example
O
H
||
||
N +
O
o-Nitrophenol
C
O +
–
H ← Hydrogen bond
The metallic crystal consists of the assemblage of positive kernels occupying fixed positions and arranged
in some definite pattern immersed in the sea of mobile valence electrons. The attractive force between mobile
electrons and the metallic kernels is referred to as metallic bonds, eg., Gold, silver, copper, sodium etc.
Polarity in covalent bond & dipole moment .
A covalent bond, in which electrons are shared unequally and the bonded atoms acquire a partial
positive and negative charge, is called a polar covalent bond or a covalent bond between two
dissimilar atoms is a polar covalent bond.
O
O
.......
O –
H
Salicylaldehyde
← Hydrogen bond
Two kinds of notation are used to indicate a polar covalent bond e.g.
+ −
H − F
or
⎯⎯→
H − F

Chemical Bonding
Polar covalent bonds may be thought of as being intermediate between the non-polar bonds and
pure ionic bonds.
Bond polarity is described in terms of ionic character, which usually increases with increasing
difference in the electronegativity between bonded atoms.
H − F H − Cl H −Br
H − I
2.1 4.0 2.1 3.0 2.1 2.8 2.1 2.5
EN :
⎯⎯⎯⎯⎯⎯⎯⎯⎯⎯⎯⎯⎯⎯⎯⎯⎯
1.9 0.9 0.7 0.4 → EN=Electronegativity
Difference in EN Ionic character decreases as the difference inelectroneg ativity decreases
Pauling has estimated the approximate percentage of ionic character in various A-B covalent
bonds from the X − X ) values i.e., electronegativity difference of the two atoms forming the
covalent bond.
( A B
X A − X B
% ionic character Nature of A – B bond
0 0 Purely covalent
0.1 to 0.8 0.5 – 15 Covalent
0.9 to 1.6 19 – 47 Polar covalent
1.7 50 50% ionic and 50% covalent
1.8 to 3.2 55 – 93 Ionic
Hanny and smyth gave the following equation for calculating the percentage of ionic character in A–
B bond on the basis of the values of electronegativity of the atoms A and B.
2
% ionic character = [16( X − X ) + 3.5( X − X ) ]
A
B
A
B
This equation gives approximate calculation of % ionic character e.g., 50% ionic character
corresponds to x ~ x ) equal to 2.1
( A B
Dipole Moment : The percentage of polar character in a covalent bond is compared in terms of
dipole moment ().
It is defined as the product of magnitude of charge developed on any of the atom
and the distance between the atoms. In general, the molecules having = 0
molecules and molecules having > 0
(1) Dipole moment is a vector quantity.
are called polar molecules.
are called non polar
(2) Non polar diatomic molecules have a zero dipole moment while diatomic molecules formed
by atoms of different electronegativity possess dipole mements.
(3) Dipole moment of polyatomic molecules is taken as the resultant of all the bonds. For
example
(i) Molecules with regular geometry have a zero dipole moment, while
(ii) Molecules with irregular geometry (having b.p. and I.P) possess a definite dipole moment.
Dipole moment of some common substance

Chemical Bonding
Substance
Dipole moment
(D)
Substance
Dipole moment
(D)
HF
H 2 O
SO 2
NH 3
NF 3
1.91
1.84
1.60
1.46
0.24
CH 3 Cl
HCl
H 2 S
HBr
HI
1.86
1.03
1.10
0.78
0.38
The % ionic character can also be calculated as follows :
Experiment al value of dipole moment
% ionic character = × 100
Theoretica l value of dipole moment
Some Important Points.
(1) Formation of ionic bond was explained by Kossel, covalent bond by Lewis and coordinate covalent
bond by Sidgwick.
(2) Molecules having same number of atoms and electrons are known as isosters.
(3) Elements of group 1 and group 2 on combining with halogens, oxygen and sulphur generally, form
ionic bonds.
(4) Ionic compound is formed if Lattice energy + Electron affinity > Ionization energy.
(5) Lattice energies of bi-bivalent solids>bi-univalent solids > uni-univalent solids. For example, lattice
2+
2−
−1
2+
−
−1
+ −
−1
energy of Mg O (3932 kJ mol ) > Ca ( F ) (2581 KJmol ) > Li F (1034 KJmol ).
2
(6) For the same cation and different anions, the lattice energy of the ionic solids decreases as the size of
−1
the anion increases. For example lattice energy of LiF (1034 KJmol ) is much higher than that of
−1
Lil (740 KJmol ).
(7) Some electron deficient compound, i.e. in which the central atom has less than 8 electrons are BeCl 2 ,
BF 6, lF 3, AlCl 3.
IF 7.
(8) Some compounds in which central atom has expanded octet, i.e. more than 8 electrons are PCl 5, SF 6,
(9) To count the number of sigma and pi bonds in any molecule, write first its expanded structure e.g. for
H
vinyl cyanide ( CH 2 = CH − CN ), we can write C = C
| − C ≡ N hence number of -bonds = 6, -bonds
H
=3.
(10) Greater the overlapping, stronger is the -bonds formed. Thus the strength of -bonds is in the
order.
(11) Dipole moment values can be used to distinguish between cis- and trans -isomers. Usually cis –
isomers have higher dipole moments that the corresponding trans- isomers.
H

Chemical Bonding
Cl
H
C = C
Cl
H
Cl
H
C = C
(12) Dipole moments help to predict the geometry of the molecules. A molecule may contain polar
covalent bonds but its resultant dipole moment may still be zero, if it has a symmetrical structure
( BeF , CO , SO , BF CH CCl etc.).
2 2 2 3 ,
4
4
(13) Amongst isomeric dichlorobenzenes, the dipole moment decreases in the order: o > m > p
(14) If a molecule of the type MX 4 has zero dipole moment, the -bonds orbitals used by M( Z < 21)
must
3
cis-1,2-Dichloroethene (µ ≠
0)
be sp ( e.
g.
CH 4 , CCl 4 , SiF4
, SnCl 4 etc.).
(15) If a molecule of the type MX 3 has zero dipole moment, the -bonds orbitals used by M( Z < 21)
must
2
( 3 3
be sp e.
g.
BF , AlCl etc.)
.
(16) If a molecule of the type MX 2
has zero dipole moment, the -bonds orbitals used by M( Z < 21)
must be sp e.
g.
BeF ).
( 2
(17) If electronegativity difference between two combining atoms = 1.9, bond has 50%ionic and 50%
covalent character. If it is >1.9, ionic character >50% and bond is taken as ionic. If it is Bond pair-bond pair repulsions.
(23) Bond angles decreases in the order :
CH 4 > NH 3 > H 2O
( CH
4
has only bp-bp, NH 3 has lp-bp and H 2 O has lp-lp resulsions)
H
Cl
trans-1,2-Dichloroethene (µ = 0)
109.5 o 107 o 104.5 o
O > H S > H Se H Te (Electronegativity decreases in the order O > S > Se > Te)
H 2 2
2 > 2
104.5 o 92.1 o 91.0 o 90 o
(24) Bond order can assume any positive value including zero.
(25) The s-orbital of one atom cannot combine with the Px
or Py
orbital of the other atom because of
improper orientation.
(26)
2P x
or has only one nodal plane while 2Px
or 2Py
has two nodal planes.
2P y
*
*
(27) The stability (or bond dissociation energy) of N 2
1 1
(because their bond orders are 3, 2 , 2 and 2 respectively).
2 2
and its ions is in the order,
N
+ − 2−
2 > N 2 = N 2 > N 2

Chemical Bonding
(28) The bond lengths of O 2
and its ions are in the order,
O
2−
−
+
2 > O 2 > O 2 > O 2
(because their bond orders are 1.0, 1.5, 2.0 and 2.5 respectively).
(29) Out of
O
+
2 O 2 ,
, O
−
2
2−
2
the diamagnetic species is i.e., O .
−
H 2
(30) Though the bond order of
+
H 2
and
−
H 2 is the same i.e., 1/2
−
H 2
is slightly less stable than
has one electron in the antibonding orbital which results in repulsion and decreases the stability.
+
H 2
since
(31) Isoelectronic species have the same shape and same bond order, e.g. CO and
+
NO
or
2−
CO 3
and
−
NO 3 .
(32) As a result of resonance, the bond order change in many molecules or ions
Bond order =
Total number of bonds between two atoms in all the structures
Total number of resonating structure
2 + 1
For example, (a) In benzene B.O. = = 1. 5
2
(b) In
2−
CO 3
O
O –
|
C
O –
– O
O –
||
C
O –
– O
O –
|
C
O
2 + 1 + 1
B. O.
= = 1.33
3
(33) In PCl 5 , axial P–Cl bonds are longer than equational P–Cl bonds.
3 d 2
(34) The d- orbital taking part in
sp , the two d-orbitals involved are d 2 and d . 2 − 2
2
dsp hybridisation is d , that in d
2 − 2 sp 3 hybridisation, it is d 2
z
x
y
x
y
z
and in
(35) Intermolecular H-bonding decreases the volatility and increases the viscosity and surface tension of
substances.
(36) Strength of H- bond decreases in the order : H ...... F > H......
O > H.....
N.
(37) Cl has same electronegativity as N, yet there is no H-bonding in HCl because size of Cl is large.
(38) H-bonding formed is usually longer than the covalent bond present in the molecule (e.g. in
H 2O,
O − H bond=0.99 A but H-bond =1.77 A )
(39) Each H 2 O
through H-atoms.
molecule forms four H-bond; two through lone pairs of electrons on the O-atom and two
(40) The high specific heat of water relative to other liquids or solids suggests that large amount of energy
is required to break H-bonds in H 2 O .
(41) Intermolecular H-bonding in ice gives it an open cage like structure. As a result, ice has lower density
than liquid H 2 O .
(42) Water has maximum density at 4 o C(277 K) since upto 4 o C, the intermolecular H-bonds keep on
breaking there by decreasing the volume and increasing the density.

Chemical Bonding
(43) O H 2
because of intermolecular H-bonding has higher boiling point than the other hydrides of group
16 elements, i.e. H 2 S, H 2Se
and H 2Te.
the boiling points increase from H 2S
→ H 2Se
→ H 2Te,
increase in vander waal’s force of attraction.
due to
(44) HF because of intermolecular H-bonding has higher boiling point than other hydrogen halides of
group 17 elements, i.e. HCl, HBr and Hl. Here again the boiling points increase in the order :
HCl < HBr <
as the magnitude of vander Waal’s forces of attraction increase due to a corresponding increase in the size of
the halogen.
of
(45) Both HF and H 2 O
undergo intermolecular H-bonding but the boiling point of HF is lower than that
H 2 O . The reason being that the number of H-bonds formed by H 2 O
formed by HF.
HCl.
Hl
is double the number of H-bonds
(46) Due to intermolecular H-bonding, ethanol has higher boiling points that diethyl ether.
(47) Alcohols and ammonia are soluble in water because they form H-bonds with water.
−
2
(48) HF (or KHF ) exists but
2
−
HCl 2 (or KHCl 2
) does not because (there is H-Bond in HF but not in
(49) Due to intramolecular H-bonding in o-nitrophenol and intermolecular H-bonding in m-and p-
nitrophenols, o-nitrophenol is more volatile and hence can be separated from m- and p-isomers by steam
distillation.
(50) Resonating structures have same positions of atoms, same number of paired and unpaired electrons
and almost equal energy. They differ only in the arrangement of electrons.
(51) Resonance hybird has lower energy than any of the contributing structures and hence is more stable.
(52) Resonance energy is the difference of energy of the resonance hybrid and that of the most stable
contributing structure.
(53) Strength of bonds ; Ionic bond > Covalent bond > Metallic bond > H-bond.
***