Nature of Bonding in Organic Molecules

1. NATURE OF BONDING
IN ORGANIC
MOLECULES
- SAHANA KAMATH
- I M.Sc. ANALYTICAL
- REG. NO. -199327

2. INTRODUCTION:
 ◘Organic chemistry deals with study of compounds of carbon.
◘Carbon possesses an unique character due to its position in periodic table.
◘ Carbon atom has six electrons surrounding its nucleus which are arranged in different orbitals of
varying energy.
◘ Total number of electrons that can accommodate in an orbital is two provided they are in opposite
spins.
◘ Electrons will never part as long as empty orbitals are available.
◘ In organic chemistry mainly s and p orbitals of carbon are utilized in covalent bond formation.
◘ s orbital is spherically symmetric.
◘ p orbital is dumb bell shaped.

3. HYBRIDIZATION
◘ According to the electronic configuration of carbon,it is found that the bonding capacity of carbon is four. Hence
has tetravalency. Ex :- CH4 , CCl
◘ This is possible only if 2s orbitals are involved in bond formation along with 2px and 2py
◘ The new electronic configuration is 1s2,2s1,2px1,2py1,2pz1.
◘ These pair with electrons of other atom to form covalent bonds.
‘Mixing of orbitals having nearly same energy and redistribution of the energies to form new orbitals
having identical shapes equivalent energies is known as hybridization.’
There are 3 types of hybridization:- sp3 hybridization, sp2 hybridization , sp hybridization.

4. sp3 hybridization:
◘ Four orbitals that are 1s and 3 p orbitals are hybridized to get
four identical hybrid orbitals with one unpaired electron in each
called sp3 orbitals.
◘ As the unpaired orbitals repel one another they are kept at a
maximum distance from one another.
◘ The geometry is tetrahedral
◘ Bond angle is 109.5◦.
◘ Each of the hybridised orbital has ¼ s and ¾ p character which
can then form four sigma bonds with electrons of other
monovalent atoms.
Application: Found in compounds with single bonds i.e, orbital
picture of methane or ethane, formation of water , alcohol,
amines , ether etc.

5. sp2 HYBRIDIZATION:
◘ One s and two p orbitals hybridize to give three sp2 hybridized orbitals with unpaired electrons each.
◘ These unpaired electrons repel each other , hence are kept at a maximum distance from each other.
◘ The bond angle is found to be 120◦.
◘ The orbitals have 1/3 s and 2/3 p character.
◘ The sp2 hybrid orbitals are slightly smaller than sp3 hybrid orbitals.
◘ The lobe of each orbital carries one electron which is shared with other monovalent atoms.
Application: Explains structure of compounds with double bond Ex:- Ethylene, aldehydes or ketones

6. sp HYBRIDIZATION:
◘ One s orbital and one p orbital hybridize to give two sp hybridized orbitals of equal energy.
◘ They are collinear.
◘ The remaining two unpaired p orbitals are perpendicular to one another and also to the plane of sp orbitals.
◘ The bond angle is found to be 180◦.
Application: Explain structure of compounds with carbon linked by a triple bond. Ex:- Acetylene.
Condition for hybridisation:
◘ Hybridization does not take place in isolated atoms.
◘ Hybrid orbitals are formed only during the time of bond formation.

7. SIGMA AND PI BONDS:
The molecular orbitals or bond orbitals formed by overlapping along their axes and have high electron density
along their internuclear axes.
The overlap is efficient and results in quite strong bonds.
These orbitals are called sigma orbitals and the bond is called sigma bond.
When half filled p orbitals forming molecular orbital are oriented in a direction perpendicular to the internuclear
axes ,they overlap laterally.
It consist two parts one lying above the plane and one below.
Such sideways overlapping is weak unlike sigma bonds.
These orbitals are called pi orbitals and the resulting bond is called pi bond.

8. According to Molecular Orbital theory a bond is formed between two atomic orbitals when they overlap each
other – greater the overlapping stronger is the bond formed.

9. CARBON – HYDROGEN SINGLE BOND:
The four sp3 hybrid atomic orbitals of carbon atoms overlap co
axially with 1s atomic orbitals of four hydrogen atoms to form four
s-sp3 sigma bonds which are strong and identical.
It has a tetrahedral geometry and bond angle of 109.28’.
CARBON- CARBON SINGLE BOND:
Formation of carbon-carbon single bond as in ethane involves
coaxial overlapping of two sp3 atomic orbitals.
Bond thus formed is a strong sigma bond with bond length 1.54A.
Bond is symmetrical about internuclear axes and is free to rotate
along its axis.

10. CARBON-CARBON DOUBLE BOND:
In this bond each carbon atom is bonded to three other atoms, two of
hydrogen and one of carbon.
In between carbon atoms there are two bonds , one is sigma bond by
linear overlap of sp2 hybridized orbitals from each carbon atom, and a
pi bond formed by formed by two unhybridized 2pz orbitals one from
each carbon atom.
As carbon atoms in pi bonds are more tightly held it has a shorter bond
length compared to that of sigma bond.
CARBON-CARBON TRIPLE BOND:
In this case out of the three bonds between two carbon atoms.
One is sigma bond formed by coaxial overlap of two sp orbitals one
from each carbon atom.
Other two are pi bonds which involve 2py and 2pz orbitals on each
carbon atom which are mutually perpendicular to sigma bond.

11. BOND LENGTH:
A covalent bond is formed by overlapping of their atomic orbitals. The distance that binds the atoms together is
called bond distance or bond length.
It ensures maximum stability to bond because at this distance the stabilizing force of overlapping of atomic
orbitals is balanced by repulsion between atomic nuclei.
Bond length is generally measured in angstrom(A) or picometer(pm) unit.
BOND ANGLE:
It is the angle between the union of two covalent bonds formed by mutual sharing of atomic orbital of an atom
with two neighboring atoms two form X-A-Y bond angle.
The bond angle is dependent on the type of hybrid orbitals involved in bond formation.
BOND ENERGY AND BOND DISSOCIATION ENERGY:
When atoms combine to form a molecule, energy is released or liberated. Similarly , when molecule dissociates
into atoms it must consume or take up an equivalent amount of energy. Energy consumed or liberated when a
covalent bond is broken or formed is known as bond dissociation energy(D).
Each bond has a characteristic value of D and it is a measure of strength of the bond, greater the bond
dissociation energy stronger is the bond. It depends on multiplicity of bond and type of hybridization.

12. LOCALIZED AND DELOCALIZED CHEMICAL BONDS:
 A covalent bond is formed by overlap of atomic or hybrid orbitals. When only two atoms share electrons
then the bond is said to be localized bond and such electrons are restricted to particular region are called
localized electrons. Localized electrons either belong to a single atom or are confined to a bond.
 But occasionally the shared electrons are not confined between the two atoms but can be spread over
three or more atoms, such electrons are called delocalized electrons and the corresponding bond is called
delocalized bond. Molecules having conjugated systems exhibit delocalization of electrons ,that is they
have delocalized chemical bonds.

13. HYDROGEN BOND:
In the compound Hydrogen fluoride, hydrogen is attached to highly electronegative fluorine by a polar covalent
bond. Hence hydrogen will have a positive charge over it. This positive charge over hydrogen in hydrogen fluoride
will be electronegatively attracted by negative charge on fluorine to form another molecule of hydrogen fluoride.
This electrostatic attraction between different molecules of hydrogen fluoride continues to form large molecular
aggregates. The bond in which hydrogen is linked to a highly electronegative atom fluorine by electrostatic
attraction is called hydrogen bond.
When hydrogen bonds are formed between the two or more molecules of same or different compounds
they are called intermolecular hydrogen bonds.
When hydrogen bonds are formed within atom of same molecule, they are called intramolecular hydrogen
bond.
Intermolecular hydrogen bonding Intramolecular hydrogen bond

14. Condition for hydrogen bonding:
Hydrogen bond in H-Z will be effective only when the hetero atom Z is highly electronegative and small in size. If
not the electrostatic attraction will be weak and hydrogen bond will not be effective.
Strength of hydrogen bond is of the order 8-40kJmol-1.
Effect of hydrogen bond on physical properties:
1. Effect on melting and boiling point.
2. Effect on solubility.
3. Effect on spectral charecteristics.
4. Effect on strength of carboxylic acids.
Importance:
1. Helps in study of physical properties.
2. Helps in study of precise shape and structure of proteins. Etc

15. POLAR AND NON- POLAR BONDS:
A covalent bond formed by the equal sharing of two electrons of opposite spins between two similar atoms in
said to be non – polar covalent bond.
If the bond formed between the dissimilar atoms of different electronegativity , the sharing is not equal and the
bond is called polar covalent bond.
A non polar covalent bond is one in which the bonded atom differ each other by less than 0.5 units of
electronegativity scale.
While the electronegativity range of polar bond ranges between 0.5 – 1.9.
IONIC OR ELECTROVALENT BOND:
An ionic bond is established as a result of the transference of electrons from one atom to other.
Elements having tendency to lose one or more electrons are called electropositive.
While elements with tendency to gain one or more electrons are called electronegative.
One or more electrons are transferred from former to latter and the atoms are covered into cations and anions.
As a result of the mutual electrostatic attraction between ions so formed, an ionic or electrovalent bond is
established.

17. COVALENT BOND:
Lewis suggested that there are atoms which attain inert gas configuration by sharing one or more electron pairs
with similar or dissimilar atoms.
Each atom contributes one electron to the electron pair has the Lewis electron pair bond a covalent bond. Thus
concept is named as Lewis- Langmuir concept.
Types: there are single double and triple bonded covalent bonds (multiple covalent bond)

18. REFERENCES:
 1. MORDERN ORGANIC CHEMISTRY BY M.K. JAIN AND S.C. SHARMA.
 2.TEXTBOOK ON ORGANIC CHEMISTRY BY K.S.MUKHERJEE