CHEMICAL BONDING AND

MOLECULAR STRUCTURE

UNIT 4
CHEMICAL BOND
The attractive force
which holds various
constituents (atoms,
ions, etc.) together in
different chemical
species is called a
chemical bond.
 KÖSSEL-LEWIS APPROACH TO
CHEMICAL BONDING

The atoms of different elements combine

with each other in order to complete their
respective octets i.e. 8 electrons in the
outermost shell or duplet i.e. 2 electrons in
the outermost shell . The atoms of all other
elements are not stable and tend to gain
stability by acquiring an electronic
configuration of nearest noble gas.
 MODES OF CHEMICAL COMBINATION

1. By the complete transfer of electrons from

one atom to other---IONIC BOND
2. By the sharing of electrons between the
atoms
a. COVALENT BOND-shared electrons are
contributed equally by 2 combining atoms.
b. CO-ORDINATE BOND- shared electrons are
contributed by 1 atom but shared by both
atoms.
 OCTET RULE
Atoms can combine either by
transfer of valence electrons from
one atom to another
(gaining or losing) or by sharing
of valence electrons in order to
have an octet in their valence
shells. This is known as octet
rule.
 IONIC BOND

Bond formed by the

complete transfer of
valence electrons from a
metal atom to a
nonmetal atom.eg- NaCl,
MgO , CaF2
 Lewis Symbols:

In the formation of a molecule, only the outer

shell electrons take part in chemical
combination and they are known as valence
electrons. These valence electrons are
represented as dots around the
symbol of the atom.
The inner shell electrons are well protected
and are generally not involved in the
combination process.
 COVALENT BOND

The bond which is formed by

the mutual sharing of
electrons between two atoms
so as to complete their octet
or duplet is called covalent
bond.
 single covalent bond.

When two atoms share

one electron pair they
are said to be joined by
a single covalent bond.
 double bond.

If two atoms share two

pairs of electrons, the
covalent bond
between them is called a
double bond.
 triple bond

If two atoms share three

pairs of electrons, the
covalent bond
between them is called a
triple bond.
Draw and explain the
bonding in the following
molecules.
1.MgCl2
2.C2H4
3.F2
4.C2H2
5.CCl4
 Lewis Representation of Simple
Molecules (the Lewis Structures)

The Lewis dot structures

provide a picture of
bonding in molecules and
ions in terms of the shared
pairs of electrons and the
octet rule.
 Steps for writing Lewis structure

1. The total number of electrons required for

writing the structures are obtained by
adding the valence electrons of the
combining atoms.
2. For anions, each negative charge would
mean addition of one electron. For cations,
each positive charge would result in
subtraction of one electron from the
total number of valence electrons.
3. The total number of
electrons is distributed as
bonding shared pairs
between the atoms in
proportion to the total bonds.
4.In general the least electronegative
atom occupies the central position
in the molecule/ion. For example in
the NF3 and CO32–, nitrogen and
carbon are the central atoms
whereas fluorine and oxygen occupy
the terminal positions.
5.After accounting for the shared pairs of
electrons for single bonds, the remaining
electron pairs are either utilized for
multiple bonding or remain as the lone
pairs. The basic requirement being that
each bonded atom gets an octet of
electrons.
Formal Charge
 Let us consider the ozone molecule (O3).
The Lewis structure of O3 may be drawn as :
Hence we represent O3 along with the
formal charges as follows:
Formal charges do not indicate real
charge separation within the molecule.
Indicating the charges on the atoms in
the Lewis structure only helps in
keeping track of the valence electrons
in the molecule. Formal charges help in
the selection of the lowest energy
structure from a number of possible
Lewis structures for a given species.
Generally the lowest energy
structure is the one with the
smallest formal charges on the
atoms. The formal charge is a
factor based on a pure covalent
view of bonding in which electron
pairs are shared equally by
neighbouring atoms.
 IONIC OR ELECTROVALENT BOND

The formation of ionic compounds would

primarily depend upon:
• The ease of formation of the positive and
negative ions from the respective neutral
atoms;
• The arrangement of the positive and
negative ions in the solid, that is, the lattice
of the crystalline compound.
Obviously ionic bonds will be
formed more easily between
elements with comparatively low
ionization enthalpies and
elements with comparatively high
negative value of electron gain
enthalpy.
Ionic compounds in the crystalline state
consist of orderly three-dimensional
arrangements of cations and anions held
together by coulombic interaction
energies.
The crystal structure of sodium chloride,
NaCl is rock salt structure.
In ionic solids, the sum of the electron
gain enthalpy and the ionization
enthalpy may be positive but still the
crystal structure gets stabilized due to
the energy released in the formation of
the crystal lattice.
For example: the ionization enthalpy for Na+
(g) formation from Na(g) is 495.8 kJ mol–1 ;
while the electron gain enthalpy for the
change Cl(g) + e–→Cl– (g) is, – 348.7 kJ
mol–1 only. The sum of the two, 147.1 kJ
mol-1 is less than compensated for by the
enthalpy of lattice formation of NaCl(s)
(–788 kJ mol–1).
Therefore, the energy released in the
processes is more than the energy absorbed.
 Thus a qualitative measure of the
stability of an ionic compound is
provided by its enthalpy of lattice
formation and not simply by
achieving octet of electrons
around the ionic species
in gaseous state.
 The Lattice Enthalpy

The Lattice Enthalpy of an ionic

solid is defined as the energy
required to completely separate
one mole of a solid ionic
compound into gaseous
constituent ions.
For example, the lattice enthalpy of NaCl
is 788 kJ mol–1. This means that 788 kJ
of energy is required to separate one
mole of solid NaCl into one mole of Na+
(g) and one mole of Cl– (g) to an infinite
distance. It is not possible to calculate
lattice enthalpy directly from the
interaction of forces of attraction and
repulsion only.
BOND PARAMETERS
 4.3.1 Bond Length

Bond length is defined as the

equilibrium distance
between the nuclei of two
bonded atoms in a molecule.
 The covalent radius

The covalent radius is measured

approximately as the radius of an
atom’s
core which is in contact with the core of
an adjacent atom in a bonded situation.
The covalent radius is half of the distance
between two similar atoms joined by a
covalent bond in the same molecule
 The van der Waals radius

The van der Waals radius represents

the overall size of the atom which
includes its valence shell in a
nonbonded situation. Further, the van
der Waals radius is half of the distance
between two similar atoms in separate
molecules in a solid.
 Bond Angle

It is defined as the angle between the

orbitals containing bonding electron
pairs around the central atom in a
molecule/complex ion.
 Bond Enthalpy

It is defined as the amount of energy required

to break one mole of bonds of a particular
type between two atoms in a gaseous state.
The unit of bond enthalpy is kJ mol–1. For
example, the H – H bond enthalpy in hydrogen
molecule is 435.8 kJ mol–1.
H2(g) → H(g) + H(g); ΔaHV = 435.8 kJ mol–1
It is important that larger the bond
dissociation enthalpy, stronger will be the
bond in the molecule.
H2O(g) → H(g) + OH(g);
ΔaH1V = 502 kJ mol–1
OH(g) → H(g) + O(g);
ΔaH2V = 427 kJ mol–1
 mean or average bond enthalpy is
used. It is obtained by dividing total bond
dissociation enthalpy by the number of
bonds broken
Resonance Structures
 Resonance Structures

According to the concept of resonance,

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. Resonance is represented by a
double headed arrow.
Structure of carbonate ion in terms of resonance
Resonance stabilizes the molecule as
the energy of the resonance hybrid is
less than the energy of any single
canonical structure;
• The cannonical forms have no realexistence.
• The molecule does not exist for a certain fraction of time
in one cannonical form and for other fractions of time in
other cannonicalforms.
• There is no such equilibrium betweenthe cannonical
forms as we have between tautomeric forms (keto and
enol) in tautomerism.
• The molecule as such has a single structure which is the
resonance hybrid of the cannonical forms and
which cannot as such be depicted bya single Lewis
structure.
 dipole moment

It can be defined as the product of the

magnitude of the charge and the distance
between the centres of positive and
negative charge. It is usually designated by
a Greek letter ‘μ’. Mathematically, it is
expressed as follows :
Dipole moment (μ) = charge (Q) × distance
of separation (r)
.
Dipole moment is usually
expressed in Debye units (D).The
conversion factor is
1 D = 3.33564 × 10–30 C m
where C is coulomb and m is meter
Dipole moment is a vector quantity
and is depicted by a small arrow
with tail on the positive centre
and head pointing towards the
negative centre. For example the
dipole moment of HF may be
represented as :
The dipole moment in case of BeF2 is zero.
This is because the two equal bond dipoles
point in opposite directions and cancel the
effect of each other.
 In tetra-atomic molecule, for example
inBF3, the dipole moment is zero although
the B – F bonds are oriented at an angle of
120° to one another, the three bond
moments give a net sum of zero as the
resultant of any two
is equal and opposite to the third.
 dipole moment of NH3 is greater than that of NF3 .WHY?

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. The orbital dipole
because of lone pair decreases the effect of
the resultant N – F bond moments, which
results in the low dipole moment of NF3.
 Fajans rule

Just as all the covalent bonds have

some partial ionic character, the ionic
bonds also have partial covalent
character. The partial covalent
character of ionic bonds was discussed
by Fajans in terms of the following
rules:
• The smaller the size of the cation and the larger the
size of the anion, the greater the covalent character
of an ionic bond.
• The greater the charge on the cation, the greater
the covalent character of the ionic bond.
• For cations of the same size and charge, the one,
with electronic configuration (n-1)dn nso, typical of
transition metals, is more polarising than the one
with a noble gas configuration, ns2 np6, typical of
alkali and alkaline earth metal cations.
The polarising power of the cation, the
polarisability of the anion and the extent of
distortion (polarisation) of anion are the factors,
which determine the per cent covalent character of
the ionic bond.
 Molecular Geometry

Molecules of different substances have diverse shapes.

The shape of a molecule is very important for its physical

and chemical properties.

Lewis Concept is unable to explain the shape of

molecules.
The first simple theory that was put forward to explain
the shape of molecules is VSEPR THEORY.
 Valence Shell Electron Pair Repulsion
(VSEPR) Theory

HELPS TO PREDİCT THE SHAPE

OF A MOLECULE FROM THE
KNOWLEDGE OF THE GROUPS
OF ELECTRONS AROUND A
CENTRAL ATOM.
 The main postulates of VSEPR
theory are
as follows
1.The shape of a molecule depends upon
the number of valence shell electron pairs
(bonded or non bonded) around the central
atom.
2. Pairs of electrons in the valence shell repel
one another since their electron clouds are
negatively charged.
3.The repulsive interaction of electron
pairs decrease in the order:

Lone pair (lp) – Lone pair (lp) >

Lone pair (lp)– Bond pair (bp) >

Bond pair (bp) –Bond pair (bp)

4. These pairs of electrons tend to occupy
positions in space that minimise
repulsion and thus maximise distance
between them.
 For the prediction of geometrical shapes
with VSEPR theory,
it is convenient to divide molecules into two
categories as
(i) molecules in which the
central atom has no lone pair
(ii)molecules in which the
central atom has one or more
lone pairs.
 Molecules in which the central atom has no
lone pair
Compounds of type
AB2

Shape: Linear
 Compounds of the type AB3
trigonal planar

BF3, CO32−, NO3−, SO3

Compounds of the type AB4
Tetrahedral
ex. CH4, NH4+
Compounds of the type AB5
Trigonal bipyramidal
 Compounds of the type AB6

Molecular Geometry:Octahedral
 Two electron pairs in the valence orbital are
arranged linearly
Three electron pairs are organized in a trigonal
planar arrangement
Four electron pairs are organized in a
tetrahedral arrangement
Five electron pairs are arranged in a trigonal
bipyramid
Six electron pairs are organized in an
octahedral arrangement
 Molecules with central atom having one or more lone pairs

Compounds of the type AB2E

Molecular Geometry: Trigonal planar
Shape: Bent
 Molecules of the type AB3E
Arrangement of electron pair:Tetrahedral
Shape: Trigonal pyramidal
Four electron pairs are organized in a tetrahedral
arrangement
Shape: trigonal pyramidal
 Molecule of type AB2E2
Molecular geometry: Tetrahedral
Shape: Bent
 Molecule of type AB4E
Molecular geometry: Trigonal bi pyramidal
Shape: see saw
There are three molecular geometries:

•Tetrahedral, if all are bonding

pairs

• Trigonal pyramidal if one is a

nonbonding pair

Bent if there are two nonbonding

pairs
 Molecule of type AB3E2
Molecular geometry: Trigonal bi pyramidal
Shape:T shape
 Collabrative Activity

Find the geometry and shape of

PO43-

SO42-

PCl3
Assignment: 1.predict the shape of CO2 , HCN CH4, NH3 SO2
PCl5, SF6 and H2O by using VSEPR theory.