GOOD PPT

1. NAVODAYA VIDYALAYA
SAMITI NOIDA
E-CONTENT PREPARATION
FOR
CLASS 11TH
CHEMISTRY
CHAPTER-4: CHEMICAL BONDING AND
MOLECULAR STRUCTURE
PREPARED BY: C. RAJASEKHAR, PGT CHEMISTRY,
JNV CHIKKABALLAPURA, KARNATAKA

2. CHEMICAL BONDING AND
MOLECULAR STRUCTURE-VSEPR
THEORY
WE KNOW THAT LEWIS CONCEPT IS UNABLE TO
EXPLAIN THE SHAPES OF MOLECULES. THE
VALENCE SHELL ELECTRON PAIR REPULSION
(VSEPR) THEORY
VSEPR THEORY PROVIDES A SIMPLE PROCEDURE
TO PREDICT THE SHAPES OF COVALENT
MOLECULES.THIS THEORY IS BASED ON THE
REPULSIVE INTERACTIONS OF THE ELECTRON
PAIRS IN THE VALENCE SHELL OF THE ATOMS.

3. THE MAIN POSTULATES OF VSEPR
THEORY
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

4. 3. THESE PAIRS OF ELECTRONS OCCUPY SUCH
POSITIONS IN SPACE THAT MINIMISE REPULSION
AND THUS MAXIMISE DISTANCE BETWEEN THEM.
4. THE VALENCE SHELL IS TAKEN AS A SPHERE
WITH THE ELECTRON PAIRS LOCALISING ON THE
SPHERICAL SURFACE AT MAXIMUM DISTANCE
BETWEEN THEM.
5. A MULTIPLE BOND IS TREATED AS IF IT IS A
SINGLE ELECTRON PAIR AND THE TWO OR
THREE ELECTRON PAIRS OF A MULTIPLE BOND
ARE TREATED AS A SINGLE SUPER PAIR.
VSEPR MODEL IS APPLICABLE TO ANY ONE
STRUCTURE OF RESONANCE STRUCTURES

5. THE REPULSIVE INTERACTION OF
ELECTRON PAIRS IN THE ORDER
LONE PAIR(LP)- LONE
PAIR(LP) > LONE PAIR(LP)
–BOND PAIR(BP) >BOND
PAIR(BP)-BOND PAIR(BP)

6. The lone pairs localized on the
central atom ,each bonded pair is
shared between two atoms. As a
result , the lp electrons in a molecule
occupy more space compared to
bp of electrons This results in greater
repulsion between lone pairs of
electrons as compared to the lone
pair –bond pair repulsions. These
repulsion effects result in deviations
from idealized shapes and alterations
in bond angles in molecules.

7. As
(i)molecules in which the central atom has no lone
pair and
(ii) molecules in which central atom has one or more
lone pairs
VSEPR THEORY DIVIDES MOLECULES
IN TO TWO CATEGORIES

8. Below are the examples for linear, trigonal
planar, tetrahedral, trigonal bipyramidal and
octahedral
SUCH ARRANGEMENT CAN BE SEEN IN THE MOLECULES LIKE
BF3
,CH4
, AND PCL5
AS DEPICTED BY THEIR BALL AND STICK
MODELS

10. VALENCE BOND THEORY
► It is based on the knowledge of atomic orbitals, E.C of
elements, the overlap criteria of AOs and the principle of
variation and superposition . Let us consider the
formation of H2
molecule which is the simplest of all
molecules.
► Consider two H atoms A and B approaching each other
having nuclei NA
and NB
and electrons eA
and eB
.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
operate.

11. Attractive forces arise between :
► (i) NA
–eA
and NB
–eB
► (ii)NA
-eB
, NB
– eA
Repulsive forces arise between
► (i) electrons of two atoms eA
– eB
,
► (ii) nuclei of two atoms NA
– NB

12. Attractive forces tend to
bring the two atoms
close to each other
where as repulsive
forces tend to push them
apart

13. ORBITAL OVERLAP CONCEPT
► Partial merging of atomic orbitals is called overlapping of
atomic orbitals which results in the pairing of electrons.
► The extent of overlap decides the strength of a covalent
bond. Greater the overlap the stronger is the bond
formed between two atoms. Therefore the formation of a
covalent bond between two atoms results by pairing of
electrons present in the valence shell having opposite
spins

14. OVERLAPPING OF
ATOMIC ORBITALS
When two atoms
come close to each
other ,there is
overlapping of AOs.
The various
arrangements of s
and p orbitals
resulting in +ve, -ve,
and zero overlap are
depicted in fig

15. HYBRIDISATION
► Acc. to Pauling the AOs combine to form new
set of equivalent orbitals known as hybrid
orbitals.
► The hybridisation is defined as the process of
intermixing of slightly different energies so as
to redistribute their energies resulting in the
formation of new set of orbitals of equivalent
energies and shape.

16. SALIENT FEATURES OF
HYBRIDISATION
1. The no. of hybrid orbitals is equal to the no. of AOs that
are hybridised
2. The hybrid orbitals are always equivalent in energy and
shape
3. The hybrid orbitals are more effective in forming stable
bonds

17. IMPORTANT CONDITIONS FOR
HYBRIDISATION
(i) The orbitals present in the valence shell of the atom are
hybridised.
(ii) The orbitals undergoing hybridisation should have
almost equal energy.
(iii) Promotion of electron is not essential condition prior to
hybridisation.
(iv) Half filled or even filled orbitals of valence shell take
part in hybridisation.

18. TYPES OF HYBRIDISATION
SP hybridisation:
► It involves the mixing of one s and one p
orbital resulting in the formation of two
equivalent SP hybrid orbitals.
► Each SP hybrid orbital has 50% s- character
and 50% p- character and such a molecule
possesses linear geometry and is known as
diagonal hybridisation

19. EXAMPLE OF MOLECULE HAVING sp
HYBRIDISATION
BeCl2
: In Be one of the 2s electron is excited to vacant 2p
orbital to account for divalency and thus one 2s and one
2p orbitals hybridise to form sp hybrid orbitals and are
oriented in opposite direction forming an angle of 1800

21. sp2
Hybridisation
It involves the mixing of one s and two p orbitals resulting
three sp2
orbitals.
Ex: BCl3
in which one of the 2s electron is excited to
vacant 2p orbital resulting in three unpaired electrons
and these orient in a trigonal planar arrangement and
overlap with 2p orbitals of Cl atoms and the bond angle
between ClBCl is 1200

23. ► In case of H2
O molecule one 2s and three 2p
orbitals form 4 sp3
orbitals ,out of which two contain
one electron each and the other two contain a pair
of electrons.
► These four hybrid orbitals acquire a tetrahedral
geometry ,with two corners occupied by lone pairs
and the other two corners occupied by bond pairs
and the bond angle is reduced to 104.50
and the
molecule thus gets a V- shape or angular geometry

24. sp3
Hybridisation
► In CH4
mixing of one s orbital and three p orbitals
takes place resulting in four SP3
hybrid orbitals
having 25% s character and 75% p character and
these four hybrid orbitals are directed towards the
corners of a tetrahedron with a bond angle of
109.50.
Other examples are NH3
and H2
o and the
bond angles are 1070
and 104.50
respectively

26. STRUCTURE OF NH3
AND H2
O
► In NH3
the valence E.C of nitrogen in the
ground state is 2s2
2px
1
2py
1
2pz
1
having three
unpaired es in the sp3
hybrid orbitals and a
lone pair of electrons in the 4th
one. These
three hybrid orbitals overlap with 1s orbitals
of three H atoms to form 3 N-H sigma bonds.
► The molecule thus gets distorted due to more
repulsion between lp-bp than between two
bond pairs of electrons and the molecule gets
distorted tetrahedron with a bond angle of
1070

28. OTHER EXAMPLES OF sp3
, sp2
and sp HYBRIDISATION
► sp3
in C2
H6
molecule: In ethane both carbons assume
sp3
hybrid state.
► One orbital overlaps axially with other carbon to form sp3
–sp3
sigma bond while the other 3 hybrid orbitals form
sp3
–s sigma bonds with H atoms. In ethane C-C bond
length is 154 pm and each CH bond length is 109 pm

29. sp2
HYBRIDISATION IN C2
H4
► In ethene one of the sp2
orbitals of each carbon overlaps
axially with each other to form a C-C sigma bond while
the other two sp2
orbitals of each carbon forms sp2
–s
sigma bond with two H atoms. The unhybrid orbital of
each carbon overlaps side wise to form a pie bond which
has equal electron clouds above and below the plane .

31. sp HYBRIDISATION IN C2
H2
MOLECULE
► In this both carbons undergo SP hybridisation
having two unhybridised orbitals i.e. 2px
and
2py
► One sp orbital of each carbon overlaps with
each other to form c-c sigma bond and other
sp hybrid orbital overlaps axially with H
atoms forming sigma bonds .The two un
hybrid p orbitals present on each carbon
overlap side wise to form two pie bonds.

33. HYBRIDISATION OF ELEMENTS
INVOLVING d ORBITALS
Formation of PCl5
molecule(SP3
d hybridisation):
► The energy of 3d orbitals are comparable to those of
3s , 3p as well as 4s and 4p orbitals .therefore the
hybridisation involving either 3s ,3p and 3d or 3d, 4s
and 4p is possible. For ex: trigonal bipyramidal-sp3
d
► Square planar- dsp2
,square pyramidal-sp3
d2
and
octahedral - sp3
d2
or d2
sp3

34. FORMATION OF PCl5
On excitation one of 3s electrons is promoted
to vacant 3d orbital making five unpaired
electrons for bonding with five Cl atoms
forming a trigonal bi-pyramidal structure in
which three bonds called equatorial bonds lie
in a plane and the other two bonds called
axial bonds are longer than equatorial bonds
due to repulsive interaction from equatorial
bond pairs and are weaker .

36. FORMATION OF
SF6
On excitation six
orbitals with unpaired
electrons are
available for bonding
with six f atoms
forming SP3
d2
hybridisation with an
octahedral geometry

37. MOLECULAR ORBITAL THEORY
Salient features:
(i) The electrons in a molecule are present in various MOs
(ii) The AO s of comparable energy and proper symmetry
combine to form MOs
(iii) AO is monocentric where as MO is polycentric
(iv) The no. of MOs formed are equal to the no. of AOs combined
(v)BMO has lower energy and greater stability than AMO
(vi)The electron probability around a group of nuclei is given by
MO
(vii)The MOs are filled in accordance with Aufbau principle,
Hund’s rule and Pauli’s exclusion principle.

38. FORMATION OF MOLECULAR ORBITALS BY
LINEAR COMBINATION OF ATOMIC
ORBITALs
► The formation of molecular orbitals can be understood in
terms of the constructive or destructive interference of
the electron waves of the combining atoms.
► The sigma MO is formed by the addition of atomic
orbitals called bonding moleclular orbital and pie MO is
formed by the subtraction of atomic orbitals called
antibonding molecular orbital

40. ► Electron density in BMO is located between the nuclei of
bonded atoms and the repulsion between the nuclei is
very less and in case of ABMO electron density is
located away from the space between the nuclei.
► There is a nodal plane between the nuclei and the
repulsion between the nuclei is very high.
► Electrons placed in a BMO tend to hold the nuclei
together and stabilise a molecule.Therefore BMO always
possesses lower energy and ABMO higher energy

41. ► The electrons placed in ABMO destabilise the
molecule and the energy is raised above the
energy of the parent AOs and the energy of
BMO is lowered than the parent AOs.
► The total energy of two MOs remains the
same as that of original atomic orbitals

42. CONDITIONS FOR THE COMBINATION OF
ATOMIC ORBITALS
1. The combining AOs must have the same or nearly the
same energy means 1s can combine with 1s only
2. The combining AOs must have same symmetry about
the molecular axis. For ex 2pz
can combine with 2pz
only
but not with 2px
or 2py
3. The combining AOs must overlap to the maximum
extent. Greater the extent of overlap, the greater will be
the electron density between the nuclei of a molecular
orbital

43. TYPES OF MOLECULAR ORBITALS AS σ,
π,δ etc
► The sigma MOs are symmetrical around the bond
axis while pie MOs are not symmetrical.
► The linear combination of 1s orbitals of two atoms
produces two MOs namely σ1s and σ*
1s .Similarly a
linear combination of 2pz
orbitals of two atoms
produces σ2pz
and σ*
2pz
.
► Molecular orbitals formed from 2px
and 2pz
are not
symmetrical around the bond axis because of the
presence of positive lobes above and negative lobes
below the molecular plane Such MOs are π and π* .
The π* AMO has a node between the nuclei

44. ENERGY LEVEL DIAGRAMS FOR
MOLECULAR ORBITALS

45. The increasing order of energies of various molecular
orbitals for O2
and F2
is given below:
For molecules such as B2 , C2 , N2 , etc. the
increasing order of energies of various molecular
orbitals is

46. ELECTRONIC CONFIGURATION AND MOLECULAR
BEHAVIOUR
Stability of molecules: If Nb
is the no. of es in BMO and
Na
is the no. of es in ABMO then
(i) the molecule is stable if Nb
> Na
(ii) the molecule is unstable if Nb
< Na
Bond order: = ½( Nb
–Na
) . A positive bond order means
a stable molecule while a negative or zero bond order
means an unstable molecule.

47. Bond length : The bond length decreases as bond order
increases
Magnetic character : If all the MOs in a molecule are
paired, the molecule is diamagnetic or if one or more
MOs are singly occupied it is paramagnetic ( attracted by
magnetic field) Ex: O2
molecule

48. HYDROGEN BONDING
► It is defined as the attractive force which
binds H atom of one molecule with the
electronegative atom(F,N or O) of another
molecule.
Cause of formation of hydrogen bond:
Due to high EN of X attached to H the
electrons are displaced towards X (acquires a
partial –ve charge ) and H acquires a partial
+ve charge resulting in the formation of a
polar molecule having electrostatic force of
attraction as H-X---H-X—H_X----.The
magnitude of H bonding is maximum in solid
state and minimum in gaseous state and it
has strong influence on the structure and
properties of the compounds

49. TYPES OF HYDROGEN BONDS
1. Intermolecular hydrogen bond: It is formed
between two different molecules of the same or
different compounds. For ex H-bond in case of HF
molecule ,alcohol or water molecules etc.
2. Intramolecular hydrogen bond: It is formed
when H atom is in between the two highly
electronegative (F,O,N)atoms present within the
same molecule. For ex ,in ortho-nitrophenol the
hydrogen is in between the two oxygen atoms.