- Organic chemistry deals with carbon-based compounds. Carbon forms covalent bonds with other elements by hybridizing its atomic orbitals. - Carbon can form single, double or triple bonds depending on whether it undergoes sp3, sp2 or sp hybridization, respectively. This allows carbon to achieve the correct molecular geometry and bond energies. - The type of bonding (ionic, covalent or coordinate covalent) between elements depends on their difference in electronegativity. Covalent bonds can be nonpolar or polar depending on this difference.

1. Chemistry for Textile
5. Bonding in organic
compounds
L8: Primary & secondary bonding,
covalent, ionic and coordinate covalent
bonding

2. Organic Chemistry
 Organic chemistry includes those
compounds whose main structure is
based on carbon, while hydrogen is the
second most abundant element found
in organic compounds.
 The compounds containing only carbon
and hydrogen are called hydrocarbons.
 Frequently elements other than carbon
and hydrogen also appear in organic
compounds. These are usually oxygen,
nitrogen, sulfur and halogens
 Most of the bonds in an organic
compounds are of covalent nature.

3. Electronic configurations
1A 2A 3A 4A 5A 6A 7A 8A
1
H
1s1
2
He
1s2
3
Li
1s2
2s1
4
Be
1s2
2s2
5
B
1s2
2s22p1
6
C
1s2
2s22p2
7
N
1s2
2s22p3
8
O
1s2
2s22p4
9
F
1s2
2s22p5
10
Ne
1s2
2s22p6
11
Na
[Ne]
3s1
12
Mg
[Ne]
3s2
13
Al
[Ne]
3s23p1
14
Si
[Ne]
3s23p2
15
P
[Ne]
3s23p3
16
S
[Ne]
3s23p4
17
Cl
[Ne]
3s23p5
18
Ar
[Ne]
3s23p6

4. Electronic configuration of carbon
in atomic form

5. Atomic structure
 The hydrogen atom has only
one electron which is present in
the first or K shell. The
hydrogen atom requires one
more electron to complete it’s
duplet.
 Atomic number of carbon is 6,
so it has six electrons in all out
of which two reside in the first
shell. The other four electrons
occupy the second (the last)
shell. To become stable it
needs four more electrons to
complete it’s octet.

6. Atomic structure
 The shape of atom is spherical like
a ball with the electron cloud
occupying most of the space taken
by the atom. The nucleus of atom
is situated at the center of sphere
and occupies a very small space
as compared to the space
occupied by the electrons. The
electrons around the nucleus are
present in different shells.
 The hydrogen atom having a single
electron has only one shell around
the nucleus.
 The carbon with six electron has
two electronic shells. The inner
shell contain two electrons and the
outer shell contains four.

7. Why compounds are formed
 The atoms whose
outer shell is
incomplete (contains
less than 8 electrons)
are unstable and tend
to react with similar or
dissimilar atoms to
make compounds.

8. Why some compounds are ionic
while others are covalent?
 It depends on the
relative electron
attracting power of
the bonding atoms.

9. Electro-negativities of elements
 Since the size and the number of protons and electrons varies in atoms of
different elements, their power to attract electrons towards nucleus also
differs. This property of atoms is called electro-negativity and it determines
the nature of bond forming between two atoms. Electro-negativity of some
of the elements is given in table below.

10. Covalent bond
 When there is minor or no
difference of electro-negativity
between the bonding atoms,
the atoms only share their
electrons to form the bond.
The resulting bond is called a
covalent bond. The bonded
atoms cannot move away from
each other until the bond
between them breaks.
 When there is somewhat
greater difference of electro-
negativity between the bonding
atoms the bond is still covalent
but becomes polar.
H2.2
Cl3.16
H2.2
C2.55
OHO = 3.44
Cl3.16

11. Covalent bond

12. Polar covalent bonds
 A covalent bond becomes polar
when there is some difference of
elecro-negativity between the
bonding atoms.
 Hydrogen chloride is a covalent
compound but since there is
greater difference of electro-
negativity between these two
atoms, the electron cloud of the
molecular orbital is slightly shifted
towards the more electronegative
atom chlorine and hydrogen is
slightly deprived of electron cloud.
 Due to shifting of electron cloud
the chlorine atom acquire a partial
negative charge and hydrogen
acquires a partial positive charge.
Such bonds are called polar
bonds.

13. Effect of varying
Electronegativity on polarity
 The polarity in covalent bonds occurs due
to difference of electro-negativity in
bonding atoms.
 The strength of polarity varies with varying
difference of electro-negativity, i.e.
increasing difference will increase polarity
and decreasing will decrease polarity.

14. Non polar and polar bonds in
organic compounds
 The saturated hydrocarbons like
methane and ethane are non polar
organic compounds. Similarly
unsaturated hydrocarbons like ethene
(or ethylene) propene (or propylene)
are also non polar compounds.
 When carbon is bonded to nitrogen or
oxygen the bond becomes polar due
to greater difference of
electronegativity. Acetone and ethyl
amine are polar compounds.

15. Some characteristics of polar and
non-polar covalent compounds
Non-Polar compounds Polar compounds
They have lower melting and
boiling points.
They have higher melting
and boiling points.
Liquid compounds show
greater volatily.
Liquid compounds show
lesser volatility.
They tend to dissolve in non-
polar solvents.
They tend to dissolve in
polar solvents.
They show lower reactivity They show higher reactivity

16. Polar vs non-polar compounds
 Methane is a non-polar
compound which occurs
in gaseous form while
methanol is polar and
occurs in liquid form.
 Similarly non-polar
ethane is a gaseous
compound while ethanol
is polar and liquid.

17. Ionic bond
 When the difference of electro-
negativity becomes too high,
the electron from less
electronegative atom moves to
the more electronegative atom.
 Such a bond is called ionic
bond because such
compounds produce ions in
solution.
 The ions in solution can move
freely, hence the bonded
atoms move away from each
other.
H2.2
Cl3.16
OHO = 3.44
Na0.90

18. Ionic bond
 Hydrogen chloride (HCl) is
a covalent compound
where the difference of
electo-negativity is less
than one, while sodium
chloride (NaCl) is an ionic
compound where the
difference of
electronegativity is more
than 2.

19. Co-ordinate covalent or dative
bond
 It is similar to the covalent
bond in that electrons are
shared between the bonding
atoms but the difference is that
both the shared electrons are
donated by only one atom.
 Since both the shared
electrons are donated by one
atom, the donor atom becomes
electron deficient and hence
gains positive charge. The
acceptor atom become
electron efficient and gains
negative charge.
 This kind of bond has both
ionic and covalent character

20. Co-ordinate covalent or dative
bond
 Another interesting example of coordinate covalent or dative bond is
the ammonium ion which is formed when ammonia reacts with
hydrochloric acid to make ammonium chloride.
 Hydrogen atom leaves it’s electron with chlorine and makes
coordinate covalent bond with ammonia. The resulting species are
ammonium and chloride ions.

21. Atomic structure
 A simple structure of atom implies that all
the electrons in a single atomic orbital
have same energy.
 The simple structure of atom does not
however represent a true picture of atom.
It is because even in a single atomic
orbital all the electrons do not have the
same energy level and all atomic orbital
except the first one consist of more than
one sub levels or orbitals which are
designated as s,p,d and f orbitals.
 In case of carbon the second (last) shell
cotains two sub-shells namely s and p. S
is of lower energy and p is of higher
energy. Two of the four electrons in last
shell occupy s orbital and the other two
Px and py orbitals.
21

22. Atomic structure
 In atomic state the valance
shell of carbon has 2 electrons
in 2s orbital, one electron in
2px orbital and one electron in
2py orbital, while the 2pz
orbital is empty.
 In this situation there are only
two unpaired eletrons which
suggests the carbon to be
bivalent.
Hybridization 22

23. Atomic structure
 In fact in all carbon
compounds’ the
carbon is found in
tetravalent state.
 For e.g. in methane
carbon is in
tetravalent form.
Hybridization 23

24. Hybridization 24
Electronic configuration and
bonding in carbon
 Carbon can make four single
bonds with four other species
or one double and two single
bonds with three other species
or one triple bond and one
single with two other species.
 Looking at the electronic
configuration it seems difficult
for carbon to make such
bonds. For making one sigma
and one or more pi bonds p
orbitals of both atoms must be
parallel which is not possible
under these conditions.

25. Hybridization 25
Valence shell atomic orbitals of
carbon

26. Hybridization 26
Mutual orientation of 2s and 2p
orbitals

27. Hybridization 27
Bonding with carbon
 Considering the
electronic configuration of
carbon, we see that 2s
orbital is completely filled,
2px and 2py are partially
filled while the 2pz orbital
is empty. Under this
condition carbon can
make only two single
bonds with two other
species.

28. Hybridization 28
Bonding with carbon
 On the other hand if we
consider the promotion of
one of the 2s electrons to
the empty 2pz orbital,
four unpaired electrons in
the 2s and 2p orbitals will
be obtained and hence
carbon should now be
able to make four single
bonds with four other
species like hydrogen.

29. Hybridization 29
Bonding with carbon
 The problem which now arises
is that the low lying s orbital
will not be able to make an
effective bond with another
species due to hindrance
offered by p orbitals.
 Furthermore all the four bonds
will not be of equal bond
energy, however in actual
practice all four single bonds
which carbon makes with four
other similar species, like
hydrogen, are of same energy.

30. Hybridization 30
Hybridization
 Carbon can make four single bonds with four separate atoms, two
single bonds and one double bond with three separate atoms or one
single and a triple bond with two separate atoms.
 All sigma bonds posses the same bond energy.
 This cannot be justified by simply promoting one 2s electron to 2p
orbital.
 The solution to this problem was provided by suggesting the idea of
hybridization of 2s and 2p atomic orbitals to give the same number
of hybridized orbitals having the same bond energy and shape.
Hybridization is the mixing up of atomic orbitals of different
energy to give a new set of same number of hybrid orbitals
having same energy.

31. Hybridization 31
Shapes of hybridized and
unhybridized orbitals

32. Hybridization 32
sp3 hybridization
 When carbon makes four single bonds, one 2s and three
2p orbitals hybridize to produce four hybrid orbitals
called sp3 orbitals.
 These four hybrid orbitals have same energy value and
have same bond angles (109.5º), directed towards the
corners of a regular tetrahedron.

33. Hybridization 33
sp3 hybridized carbon

34. Hybridization 34
Methane molecule

35. Hybridization 35
sp2 hybridization
 When carbon makes two single bonds and one double
bond, one 2s and two 2p orbitals hybridize to produce
three hybrid orbitals called sp2 orbitals. The third 2p
orbital do not hybridize.
 The three hybridized orbitals have the same energy
value and bond angle (120º), while the unhybridized 2p
orbital have different energy value and helps in making
pi (Л) bond.

36. Hybridization 36
sp2 hybridization

37. Hybridization 37
sp hybridization
 When carbon makes one single bond and one triple bond, one 2s
and one 2p orbital hybridize to produce two hybrid orbitals called sp
orbitals. The two remaining 2p orbitals do not hybridize.
 The two hybridized orbitals have the same energy value and the
bond angle between them is 180º, while the unhybridized 2p orbitals
have different energy value and helps in making triple bond.

38. Hybridization 38
Sp hybridization C – C triple bond

39. Hybridization 39
Sp3 hybridization in Nitrogen,
Oxygen and Halogens
 Nitrogen has three unpaired electrons in its 2p orbitals and two
paired electrons in s orbital. While bonding it hybridizes to sp3
geometry making three single bonds with other species.
 Oxygen has two unpaired electrons in its 2p orbitals and four
electrons in two pairs occupying 2s and one of the 2p orbitals.
While bonding it hybridizes to sp3 geometry making two single
bonds with other species.
 Halogens have one unpaired electron in its 2p orbital and six
electrons in three pairs occupying 2s and two of the 2p orbitals.
While bonding it hybridizes to sp3 geometry making one single
bonds with another specie.

40. Hybridization 40
sp3 Hybridization in Nitrogen,
oxygen and halogens

41. Hybridization 41
sp2 hybridization in nitrogen and
oxygen
 Nitrogen has three unpaired electrons in its 2p orbitals and two paired electrons in 2s
orbital. While making double bond it hybridizes to sp2 geometry making two sigma
bond and one pi bond with other species.
 Oxygen has two unpaired electrons in its 2p orbitals and four electrons as two pairs
occupying 2s and one of the 2p orbitals. While making double bond it hybridizes to
sp2 geometry forming one sigma and one pi bond with the other specie.

42. Hybridization 42
sp hybridization in nitrogen
 Nitrogen has three unpaired electrons in its 2p
orbitals and two paired electrons in 2s orbital.
While making triple bond it hybridizes to sp
geometry making one sigma bond and two pi
bonds with the other specie.

43. Primary and secondary bonding
 A primary bond is a true bond
which is formed by sharing of
electrons between two atoms
or by transfer of an electron
from one atom to the other.
Primary bonds are covalent,
ionic and coordinate covalent
or dative bonds.
 A secondary bond on the other
hand is one where no true
sharing of electrons or transfer
does occur. Examples of
secondary bonding are
hydrogen bonding, dipole-
dipole interaction, Van der
Waal’s forces etc.