Krishna's
TEXT BOOK

ORGANIC

CHEMISTRY-I
Second Paper

B.Sc. First Year

st
(For B.Sc. I year Students of all Colleges affiliated to Universities in Uttar Pradesh)

As per U.P. Unified Syllabus (w.e.f. 2011-12)
By
Alok Bariyar
Ph.D. (CSIR fellow), M.Sc.(IIT Delhi), GATE., NET

Ex. Scientist (Nuclear)
Bhabha Atomic Research Centre,
Mumbai (Maharashtra)

R.P. Singh
M.Sc., Ph.D.

Babita Agrawal
M.Sc., Ph.D.

Head, Dep’t of Chemistry

Head, Dep’t of Chemistry

KNI of Physical Social Science,
Sultanpur (U.P.)

B.S.A. (P.G.) College,
Mathura (U.P.)

Ashish Dwivedi

Prashant Singh

M.Sc., Ph.D.

M.Sc., Ph.D.

Asst. Prof., Dep’t of Chemistry

Asst. Prof., Dep’t of Chemistry

Ganpat Sahai P.G. College,
Sultanpur (U.P.)

KNIPSS, Sultanpur (U.P.)

KRISHNA Prakashan Media (P) Ltd.
KRISHNA HOUSE, 11, Shivaji Road, Meerut-250 001 (U.P.), India


Jai Shri Radhey Shyam


Dedicated
to

Lord

Krishna
Authors & Publishers


P reface

W

e are happy to present this book entitled “Organic Chemistry-I”. It has been written
according to the latest U.P. Unified Syllabus to fulfil the requirement of B.Sc. Ist
year students of all Colleges & Universities in Uttar Pradesh.

The book is written with the following special features:
1.

It is written in a simple language so that all the students may understand it easily.

2.

It has an extensive and intensive coverage of all topics.

3.

In each Chapter, Solved Examples are given based on different Topics.


4.

The complete Syllabus has been divided into Seven Chapters under Four Units.

5.

Sufficient Numerical Problems, Subjective Questions and Objective type questions with
Hints & Solutions given at the end of each chapter will enable students to understand the concept .
First of all we want to express our sincere gratitude to Purnima Sinha, Dr. S.B.P Sinha,

Prof. J.C. Ahluwalia, Prof. N.K. Jha for their invaluable guidance, immense interest and
constant encouragement for the successful completion of the work. We are also thankful to
Bandana Bariyar, Ashish Bariyar, Abhishek Bariyar, Archi Bariyar & Aradhyaa Bariyar for
their kind help at many occasions.
We are extremely grateful to our respected and beloved Parents whose incessant
inspiration guided us to accomplish this work. We also express gratitude to our respective
Families for their moral support.
We are immensely thankful to Mr. S.K. Rastogi (Managing Director), Mr. Sugam Rastogi
(Executive Director), Mrs. Kanupriya Rastogi (Director) and entire team of Krishna Prakashan
Media (P) Ltd., for taking keen interest in this project and outstanding Management in getting
the book published.
The originality of the ideas is not claimed and criticism and suggestions are invited from the
Students, Teaching community and other Readers.

July 2015

–Authors

(iv)



Syllabus
ORGANIC CHEMISTRY-I
B.Sc. Ist Year; IInd Paper
UGC Model Curriculum (w.e.f. 2001) & U.P. UNIFIED Syllabus (w.e.f. 2011-12)

UNIT-I

M.M. 50

1.

Structure and Bonding: Hybridization, bond lengths and bond angles, bond energy, localized and
delocalized chemical bonding, Van der Waals interactions, inclusion compounds, clatherates, charge
transfer complexes, resonances, hyperconjugation, aromaticity, inductive and field effects, hydrogen
bonding.

2.

Mechanism of Organic Reactions: Curved arrow notation, drawing electron movements with allows,
half-headed and double-headed arrows, homolytic and heterolytic bond fission, Types of reagents –
electrophiles and nucleophiles, Types of organic reactions, Energy considerations. Reactive
intermediates – Carbocations, carbanions, free radicals, carbenes, arynes and nitrenes (with examples).
Assigning formal charges on intermediates and other ionic species. Methods of determination of
reaction mechanism (product analysis, intermediates, isotope effects, kinetic and stereochemical
studies).

3.

Alkanes and Cycloalkanes: IUPAC nomenclature of branched and unbranched alkanes, the alkyl

group, classification of carbon atom in alkanes, Isomerism in alkanes, sources methods of formation
(with special reference to Wurtz reaction, Kolbe reaction, Corey-House reaction and decarboxylation of
carboxylic acids), physical properties and chemical reactions of alkanes, Mechanism of free radical
halogenation of alkanes: orientation, reactivity and selectivity. Cycloalkanes – Nomenclature, methods
of formation, chemical reactions, Baeyer's strain theory and its limitations. Ring strain in small rings
(cyclopropane and cyclobutane), theory of strain less rings. The case of cyclopropane ring, banana
bonds.

UNIT-II
4.

Stereochemistry of Organic Compounds: Concept of isomerism, Types of isomerism; Optical
isomerism – elements of symmetry, molecular chirality, enantiomers, stereogenic center, optical
activity, properties of enantiomers, chiral and achiral molecules with two stereogenic centers,
disasteromers, threo and erythro diastereomers, meso compounds, resolution of enantionmer,
inversion, retention and recemization. Relative and absolute configuration, sequence rules, D & L and R
& S systems of nomenclature. Geometric isomerism – determination of configuration of geometric
isomers, E & Z system of nomenclature, geometric isomerism in oximes and alicyclic compounds.
Conformational isomerism – conformational analysis of ethane and n-butane; conformations of
cyclohexane, axial and equatorial bonds, conformation of mono substituted cyclohexane derivatives,
Newman projection and Sawhorse formulae, Fischer and flying wedge formulae, Difference between
configuration and conformation.

5.

Alkenes, Cycloalkenes, Dienes and Alkynes: Nomenclature of alkenes, methods of formation,
mechanisms of dehydration of alcohols and dehydrohalogenation of alkyl halids, regioselectivity

UNIT-III
(v)



in alcohol dehydration, The Saytzeff rule, Hoffmann elimination, physical properties and relative
stabilities of alkenes. Chemical reactions of alkenes – mechanism involved in hydrogenation,
electrophilic and free radical additions, Markownikoff's rule, hydroboration- oxidation,
oxymercuration-reduction. Epoxidation, ozonolysis, hydration, hydroxylation and oxidation with
KMnO4, Polymerization of alkenes, Substitution at the allylic and vinylic positions of alkenes,
Industrial applications of ethylene and propene. Methods of formation, conformation and chemical
reactions of cycloalkenes; Nomenclature and classification of dienes : isolated, conjugated and
cumulated dienes, Structure of allenes and butadiene, methods of formation, polymerization, chemical
reaction – 1, 2 and 1, 4 additions, Diels-Alder reaction. Nomenclature, structure and bonding in
alkynes, Methods of formation, Chemical reactions of alkynes, acidity of alkynes, Mechanism of
electrophilic and nucleophilic addition reactions, hydroboration-oxidation, metal-ammonia
reductions, oxidation and polymerization.

UNIT-IV
6.

Arenes and Aromaticity: Nomenclature of benzene derivatives, The aryl group, Aromatic nucleus
and side chain, Structure of benzene; molecular formula and kekule structure, stability and carboncarbon bond lengths of benzene, resonance structure, MO picture. Aromaticity: The Huckle rule,
aromatic ions. Aromatic electrophilic substitution – general pattern of the mechanism, role of s and p
complexes, Mechanism of nitration, halogenation, sulphonation, mercuration and Friedel-Crafts
reaction. Energy profile diagrams. Activating and deactivating substituents, orientation and ortho/para
ratio, Side chain reactions of benzene derivatives, Birch reduction; Methods of formation and chemical
reactions of alkylbenzenes, alkynylbenzenes and biphenyl, naphthalene and Anthracene;

7.

Alkyl and Aryl Halides: Nomenclature and classes of alkyl halides, methods of formation, chemical
reactions, Mechanisms of nucleophilic substitution reactions of alkyl halides, SN2 and SN1 reactions

with energy profile diagrams; Polyhalogen compounds : Chloroform, carbon tetrachloride; Methods
of formation of aryl halides, nuclear and side chain reactions; The addition-elimination and the
elimination-addition mechanisms of nucleophilc aromatic substitution reactions; Relative reactivities
of alkyl halides vs. allyl, vingl and aryl halides, Synthesis and uses of DDT and BHC.

(vi)


Detailed Contents
Unit - I
Chapter 1

Structure and Bonding

O–01-62

1.
2.
3.

Introduction
03
Localised and Delocalized Chemical Bonding
Characteristics of Covalent Bond

$

Solved Examples

4.

5.
6.
7.
8.

Delocalized Chemical Bonding and Resonance
Hybridization
18
Electron Displacement Effect
29
Electromeric Effect
33
Strength of Acids and Bases
42

$
$
$

Exercise
55
Answers
61
Hints & Solutions

Chapter 2

03
07


10

13

62

Mechanism of Organic Reactions

1.

Fundamental Concepts of Organic Reaction

$

Solved Examples

2.
3.

Various Reaction Intermediates
76
Methods of Determination of Reaction Mechanism

$
$
$

Exercise
99
Answers

102
Hints & Solutions

Chapter 3

O–63-102

63

66

96

102

Alkanes & Cycloalkanes

1.
2.
3.

Alkanes
103
Cycloalkanes
116
Stability of Rings and Ring Strain

$
$
$

$

Solved Examples
Exercise
130
Answers
133
Hints & Solutions

O–103-134

124

126

134

Unit - II
Chapter 4
1.

Stereochemistry of Organic Compounds

Introduction

135
(vii)

O–135-200



2.

Structural Isomerism

$

Solved Examples

3.
4.
5.

Stereoisomerism
Geometrical Isomerism
Conformational Isomerism

$
$
$

Exercise
186
Answers
194
Hints & Solutions

136
140


140
159
173

195

Unit - III
Chapter 5

Alkenes, Cycloalkenes Dienes and Alkynes

1.

Alkenes

$

Solved Examples

2.
3.
4.

Cycloalkene (or Cycloolefin)
Dienes
238
Alkynes
245

$

$
$

Exercise
266
Answers
273
Hints & Solutions

O–201-274

201
211

233

273

Unit - IV
Chapter 6

Arenes and Aromaticity

O–275-344

1.

Arenes

$


Solved Examples

2.
3.
4.
5.
6.
7.
8.

Structure of Benzene
278
Aromaticity and Huckel's Rule
283
Electrophilic Aromatic Substitution Reactions
287
Disubstitution in Benzene Ring and Theory of Substituent Effect
Fused or Condensed Aromatic Hydrocarbons
312
Anthracene
322
Biphenyls
329

$
$
$

Exercise

331
Answers
341
Hints & Solutions

Chapter 7

275
277

299

341

Alkyl and Aryl Halide

1.
2.

Introduction
Alkyl Halides

$

Solved Examples

3.
4.
5.
6.

7.

Aliphatic Nucleophilic Substitution
354
Polyhalogen Compounds
367
Aryl Halides
370
Synthesis and Uses of DDT
382
Synthesis and Uses of BHC
383

$
$
$

Exercise
391
Answers
395
Hints & Solutions

O–345-396

345
345
348

396


❍❍❍
(viii)


Book-2
Organic C hemistry-I
Unit-I
Chapter 1 : Structure and Bonding
Chapter 2 : Mechanism of Organic Reactions
Chapter 3 : Alkanes and Cycloalkanes

Unit-II
Chapter 4 : Stereochemistry of Organic Compounds

Unit-III
Chapter 5 : Alkenes, Cycloalkenes, Dienes and Alkynes

Unit-IV
Chapter 6 : Arenes and Aromaticity
Chapter 7 : Alkyl and Aryl Halide


UnitO-3
-I

C HAPTER

1


Structure and Bonding

1. Introduction
Before starting any discussion on organic chemistry, first let us make it very clear that the fundamentals
involving bond cleavage, various intermediates and their reactivities, types of reagents, field effects etc form
the backbone of entire organic chemistry. These concepts will be useful in understanding the acidic or basic
behaviour of many organic compounds, types of various organic reactions, mechanism of various reactions
and their other aspects.

2. Localised and Delocalized Chemical Bonding
2.1 Types of Bond
As stated a chemical bond is an attraction between atoms. This attraction may be seen as the result of
different behaviors of the outermost electrons of atoms. Although all of these behaviors merge into each other
seamlessly in various bonding situations so that there is no clear line to be drawn between them, customarily
the chemical bonds are classified into different types.
One more point has to be reiterated before discussing the classification of chemical bond. Atom may attain a
stable electronic configuration in three different ways by losing electron, by gaining or by sharing electron.
Moreover, elements may be divided into following three types depending upon their electronegativity as:
1.

Electropositive elements: Those elements whose atoms give up one or more electron fairly readily.

2.

Electronegative elements: Which will accept electron.

3.

Elements which have little tendency to lose or gain electrons.


Three different types of bond may be formed depending on the electropositive or electronegative character of
atom involved.
Electropositive element + Electropositive element → Metallic bond
Electropositive element + Electronegative element → Ionic bond
Electronegative element + Electronegative element → Covalent bond


O-4

2.2 Covalent Bond
In the simplest view of a so-called 'covalent' bond, one or more electrons (often a pair of electrons) are drawn
into the space between the two atomic nuclei. Here the negatively charged electrons are attracted to the
positive charges of both nuclei, instead of just their own. This overcomes the repulsion between the two
positively charged nuclei of the two atoms, and so this overwhelming attraction holds the two nuclei in a fixed
configuration of equilibrium, even though they will still vibrate at equilibrium position. Thus, covalent
bonding involves sharing of electrons in which the positively charged nuclei of two or more atoms
simultaneously attract the negatively charged electrons that are being shared between them. These bonds
exist between two particular same or different atoms, and have a direction in space, allowing them to be
shown as single connecting lines between atoms in drawings. For Example, Two chlorine atoms react to form
a Cl2 molecule
Cl + Cl

Cl

Cl

Each chlorine atom gives a share of one of its electrons to other atom. A pair of electrons is shared equally
between both atoms and each atom now has eight electrons in its outer shell (stable octet). In a similar way, a
molecule of tetra chloromethane CCl4 is made up of one carbon and four chlorine atoms.
C + 4C


Cl
Cl C Cl
Cl

The carbon atom is short off four electrons so as to have noble gas structure. Consequently, it forms four
bonds with the chlorine atoms which themselves are short of one electron so they each form one bond by
sharing electrons. In this way, both carbon and all four chlorine atoms attain a noble gas structure.
N + 3[H ]

H N H
H

A molecule of ammonia (NH3 ) is made up of one nitrogen and three hydrogen atoms. Other examples of
covalent bonds include water (with two covalent bonds) and hydrogen fluoride (one covalent bond and three
lone pairs).
H O ,
H

H F

In a polar covalent bond, one or more electrons are unequally shared between two nuclei. Covalent bonds
often result in the formation of small collections of better-connected atoms called molecules. These molecules
in solid and liquid state are bound to other molecules by intermolecular forces that are often much weaker
than the covalent bonds that hold the molecules internally together. Such weak intermolecular bonds give
organic molecular substances, such as waxes and oils, their soft character, and their low melting points. When
covalent bonds link long chains of atoms in large molecules, however (as in polymers such as nylon), or when
covalent bonds extend in networks through solids that are not composed of discrete molecules (such as
diamond or quartz or the silicate minerals in many types of rock) then the structures that result may be both
strong and tough, Also, the melting points of such covalent polymers and networks increase greatly.



O-5

2.3 Types of Covalent Bonds Sigma (σ) and Pi (π) Bonds
Depending upon the type of overlapping, the covalent bonds are mainly of two types.
1.

Sigma (σ) bond: When a bond is formed between two atoms by the overlap of their atomic orbitals
along the internuclear axis the resulting bond is called sigma(σ) bond. Such type of overlap is also
known as end to end or head on overlap. It is a strong bond and cylindrically symmetrical. The
overlapping along the internuclear axis can take place in any of the following ways:
(i)

s—s overlapping: This type of overlapping takes place between atoms having half filled
s–orbitals in their outer most energy shell. For example, in the formation of hydrogen molecule,
1s orbital of one hydrogen atom overlaps with 1s orbital of other hydrogen atom thus forming a
sigma bond.

+

+
1s arbital of
H–atom

1s arbital of
H–atom

+


s–s overlap

Fig. 1: Overlap of s orbital

(ii)

s— p overlapping: In this case, half filled s–orbital of one atom overlaps with the half filled
p–orbtial of another atom. A simple example of this type is the formation of hydrogen fluoride.
Here 1s orbitals of hydrogen overlaps with 2pz orbital of fluorine.

2pz orbital
of F–atom

1-s orbital
of H–atom

+

+

+

s–p overlap

Fig. 2: Overlap of s–p orbital

(iii) p—p overlapping: This type of overlapping occurs when p–orbital of one atom overlaps with
the p–orbital of the other as in case of fluorine molecule. The molecule of fluorine is produced by
the overlapping of 2pz orbitals of the two fluorine atoms.
+


+
2pz orbital
of F–atom

2pz orbital
of F-atom

+
p–p overlap

Fig. 3: Overlap of p–p orbital

2.

Pi (π) bond: Pi (π)bond is formed by lateral or sidewise overlapping of p orbitals. Sideways overlap
means overlapping of p orbitals in a direction perpendicular to the internuclear axis. A π bond is not
formed between two bonded atoms unless the two are held together with a σ −bond. It is relatively a
weaker bond since the electrons are not strongly attracted by the nuclei of bonding atoms. For example,


O-6
(i)

In case of oxygen molecule (each oxygen atom having electronic configuration,
1s2 2s2 2p2x 2p1y 2p1z ), the two atoms are held together by one σ-bond and one π-bond.

+

Two Half–filled

p–orbitals of
O–atom

Two Half–filled
p–orbitals of
O–atom

O2 molecule
(O=O)

Fig. 4: Formation of O2 molecule

(ii)

In the molecule of nitrogen both nitrogen atoms are held together by one σ-bond and 2 π-bonds.
Nitrogen atom has an electronic configuration 1s2 2s2 2p1x 2p1y 2p1z .

+

N–atom

N–atom

N2 molecule
–
(N=N)
2

Fig. 5: Formation of nitrogen


It is important to remember that the 's' orbitals can only form σ-bonds, whereas the p, d & f
orbitals can form both σ and π-bonds.

2.4 Multiple Bonding
When two atoms share a single pair of electrons, the bond is referred to as a single bond. Atoms can also
share two or three pairs of electrons in the aptly named double and triple bonds. The first bond between two
atoms is called the σ (sigma) bond. All subsequent bonds are referred to as π (pi) bonds. In Lewis structures,
multiple bonds are depicted by two or three lines between the bonded atoms. The bond order of a covalent
interaction between two atoms is the number of electron pairs that are shared between them. Single bonds
have a bond order of 1, double bonds 2, and triple bonds 3. Bond order is directly related to bond strength
and bond length. Higher order bonds are stronger and shorter, while lower order bonds are weaker and
longer. The Lewis structures for some common molecules involving multiple covalent bonds can be found
below.


O-7
H

H
C

C

H

H

C

N


C

C
C
H
acetylene

H

O

H
Formaldehyde

H
H
Formaldimine

Ethylene
(ethere)
H

H

H

–
+
C

O
carbon
monoxide

N
C
hydrogen
cyanide

Fig. 6: Lewis structures of molecules with multiple bonds

2.5 Co–ordinate Bond
A covalent bond results from sharing of a pair of electrons between two atoms, where each atom contributes
one electron to the bond. It is also possible to have an electron pair bond where both the electrons come from
one of the two binding atoms and there is no contribution from the other atom. Such bonds are called
co–ordinate bonds or dative bonds. So, co–ordinate bond is a special type of covalent bond in which both the
bonded electrons come from one of the two binding atoms. One common example is formation of
ammonium ion. Even though the ammonia molecule has a stable electronic configuration it can react with a
hydrogen ion (H+) by donating a lone pair of electrons from N atom to H+ ion forming the ammonium ion
NH4+ .
H
H

N

+

H
+ [H+]


+

or H N H

H N H

H

H

H

H

Covalent bonds are usually shown as a straight line joining the two atoms, and co–ordinate bonds as arrows
indicating which atom is donating the electron. Similarly, ammonia donates its lone pair to boron trifluoride
and by this means the boron atom attains noble gas configuration.

H

H

F

H N

B

H


F

F

H
+

N

B

F

F

H

F

In a similar way a molecule of BF3 can form a co–ordinate bond by accepting a lone pair from a F − ion.
–
F

+

F
B F
F

F

F

B

–
F

F

3. Characteristics of Covalent Bond
Some important characteristics of covalent bonds like bond length, bond angle and bond energy are
discussed below:


O-8

3.1 Bond Length
In molecular geometry, bond length or bond distance is the average distance between nuclei of two
bonded atoms which may be same or different in a molecule.

It is a transferable property of a bond between atoms of fixed types, relatively independent of the rest of the
molecule. Bond length is related to bond order, when more electrons participate in bond formation the bond
will get shorter. Bond length is also inversely related to bond strength and the bond dissociation energy, as all
other things being equal a stronger bond will be shorter. In a bond between two identical atoms half the bond
distance is equal to the covalent radius. Bond lengths are measured in the solid phase by means of X-ray
diffraction, or approximated in the gas phase by microwave spectroscopy. A set of two atoms sharing a bond
is unique going from one molecule to the next. For example the carbon to hydrogen bond in methane is
different from that in methyl chloride. It is however possible to make generalizations when the general
structure is the same.
The actual bond length between two atoms in a molecule depends on such factors as the orbital hybridization

and the electronic and steric nature of the substituents. The carbon-carbon bond length in diamond is 154 pm
which is also the largest bond length that exists for ordinary carbon covalent bonds.
There are compounds in which shorter than average carbon–carbon bonds distances are also possible.
Alkenes and alkynes have bond lengths of respectively 133 and 120 pm due to increased s-character of the
sigma bond and presence of π bond.

153.5 pm

C2H6

133.9 pm

C2H4

120.3 pm

C2H2

Fig. 8: Bond length in different hydrocarbons


O-9
In benzene all C-C bonds have the same length, 139 pm, due to resonance.
Tabel 1: Bond lengths in organic compounds
C—H

Length (pm)

C—C


Length(pm)

Multiple–bonds

Length (pm)

sp3—H

110

sp3—sp3

154

Benzene

140

sp2—H

109

sp3—sp2

150

Alkene

133.9


sp—H

108

sp2—sp2

147

Allene

120.3

sp3—sp

146

sp2—sp

143

sp—sp

137

3.2 Bond Angle
Molecular geometry is the three-dimensional arrangement of the atoms that constitute a molecule. It
determines several properties of a substance including its reactivity, polarity, phase of matter, color,
magnetism, and biological and physical activities as well. Molecular geometries can be specified in terms of
bond lengths, bond angles and torsional angles. A bond angle is the angle formed between three atoms
across at least two bonds. For four atoms bonded together in a chain, the torsional angle is the angle between

the plane formed by the first three atoms and the plane formed by the last three atoms. As a matter of fact, the
angles between bonds that an atom forms also depend on the rest of molecule, albeit weakly. Some of the
molecules alongwith bond angles are shown below,

109° 47
120°

180°
Linear

Trigonal
planar

Tetrahedral

90°
90°

120°

Trigonal
bipyramidal

90°

Octahedral

Fig. 9: Different molecular shapes along with bond angles



O-10

Solved Examples
Example1: Predict all bond angles in the following molecules.

(i) CH3Cl

(ii) CH3CN

(iii) CH3COOH

Solution: (i) The Lewis structure of methyl chloride is:
H
H C

Cl

H

In the Lewis structure of CH3Cl carbon is surrounded by four regions of high electron density, each of which
forms a single bond. Based on the VSEPR model, we predict a tetrahedral distribution of electron clouds
around carbon, H - C - H and H - C - Cl bond angles of 109.5°, and a tetrahedral shape for the molecule. Note
the use of doted lines to represent a bond projecting behind the plane of the paper and a solid wedge to
represent a bond projecting forward from the plane of the paper.

(ii) The Lewis structure of acetonitrile, CH3CN is:
H
H C

C


N

H

The methyl group, CH3-, is tetrahedral. The carbon of the -CN group is in the middle of a straight line
stretching from the carbon of the methyl group through the nitrogen.
109.5° H
H C C N
H 180°

(ii) The Lewis structure of acetic acid is:
H O
H C
H

C

O H


O-11
Both the carbon bonded to three hydrogens and the oxygen bonded to carbon and hydrogen are centers of
tetrahedral structures. The central carbon will have 120° bond angles.

The geometry around the first carbon is tetrahedral, around the second carbon atom is trigonal planar, and
around the oxygen is bent.

3.3 Bond Energy
In chemistry, bond energy (E) is the measure of bond strength in a chemical bond. It is the energy required

to break one mole of molecules into their individual atoms. For example, the carbon-hydrogen bond energy
in methane E(C-H) is the enthalpy change involved with breaking up one molecule of methane into a carbon
atom and 4 hydrogen radicals divided by 4. Bond energy (E) must not be confused with bond-dissociation
energy. As such, bond energy is an average of different bond dissociation energies of the same type of bonds.
The same bond can appear in different molecules, but it will have a different bond energy in each molecule as
the other bonds in the molecule will affect the bond energy of the specific bond. So the bond energy of C-H in
methane is slightly different than the bond energy of C-H in ethane. One can calculate a more general bond
energy by finding the average of the bond energies of a specific bond in different molecules to get the average
bond energy.
Table 2: Average bond energies (kj/mol)
Single Bonds

Multiple Bonds

H—H

432

N—H

391

I–I

149

C=C

614


H—F

565

N—N

160

I–Cl

208

C≡C

839

H—Cl

427

N—F

272

I–Br

175

O=O


495

H—Br

363

N—Cl

200

S–H

347

C = O*

745

H—I

295

N—Br

243

S–F

327


C≡O

1072

C—H

413

N—O

201

S–Cl

253

N=O

607

C—C

347

O—H

467

S–Br


218

N=N

418

C—N

305

O—O

146

S–S

266

N≡N

941

C—O

358

O—F

190


Si–Si

340

C≡N

891

C—F

485

O—Cl

203

Si–H

393

C≡N

615

C—Cl

339

O—I


234


O-12
C—Br

276

F—F

154

Si–C

360

C—I

240

F—Cl

253

Si–O

452

C—S


259

F—Br

237

Cl—Cl

239

Cl—Br

218

Br—Br

193

Some of the characteristics of bond energy values are:
1.

Average bond energy values are not as accurate as a molecule specific bond-dissociation energies.

2.

Double bonds are higher energy bonds in comparison to a single bond (but not necessarily 2-fold
higher).

3.


Triple bonds are even higher energy bonds than double and single bonds (but not necessarily 3-fold
higher).

Example 2: (i) What is the definition of bond energy? When is energy released and absorbed?

(ii) If the bond energy for H-Cl is 431 kJ/mol. What is the overall bond energy of 2 moles of
HCl?
(iii) Using the bond energies given in the chart above, find the enthalpy change for: the
decomposition of water
2H2O(g)→2H2+O2(g)
Is the reaction written above exothermic or endothermic? Explain.
(iv) Which bond in list below has the highest and lowest bond energy? H-H, H-O, H-I, H-F.
Solutions (i) Bond energy is the energy required to break a bond that exists between two atoms. Energy is
given off when the bond is broken, but is absorbed when a new bond is created.

(ii) Simply multiply the average bond energy of H-Cl by 2. This leaves one with 862 kJ/mol (using the
table).
(iii) The enthalpy change deals with breaking two mole of O-H bonds and the formation of 1 mole of O-O
bonds and two moles of H-H bonds.The sum of the energies required to break the bonds on the
reactants side is 4×460 kJ/mol = 1840 kJ/mol.The sum of the energies released to form the bonds on
the products side is
2 moles of H-H bonds = 2 x 436.4 kJ/mol = 872.8 kJ/mol
1 moles of O–O bond – 1 x 498.7 kJ/mol = 498.7 kJ/mol
The released energy = 872.8 kJ/mol + 498.7 kJ/mol = 1371.5 kJ/mol. Total energy difference is 1840
kJ/mol - 1371.5 kJ/mol = 469 kJ/mol, which indicates that the reaction is endothermic and that 469 kJ
of heat is needed to be supplied to carry out this reaction.


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(iv) H-F has the highest bond energy since the difference in electronegativity is the greatest. However, the

H-I bond is the lowest bond energy (not due to the electronegativity difference, but due to the greater
size of the I atom).

3.4 The Shape of Molecules
The three dimensional shape or configuration of a molecule is an important characteristics. This shape is
dependent on the preferred spatial orientation of covalent bonds to atoms having two or more bonding
partners. Three dimensional configurations are best viewed with the aid of models. In order to represent such
configurations on a two-dimensional surface (paper, blackboard or screen), one can normally use
perspective drawings in which the direction of a bond is specified by the line connecting the bonded
atoms. In most cases the focus of configuration is a carbon atom, so the lines specifying bond directions will
originate there. As defined in the diagram, a simple straight line represents a bond lying approximately in the
surface plane. The two bonds to substituents A in the structure below are of this kind,
normal bond

A
A

C

D
B

wedge bond
or

}hatched bond
dashed bond

A wedge shaped bond is directed in front of this plane (thick end toward the viewer), as shown by the bond to
substituent B; and a hatched bond is directed in back of the plane (away from the viewer), as shown by the

bond to substituent D. Some texts and other sources may use a dashed bond in the same manner as we have
defined the hatched bond, but this can be confusing because the dashed bond is often used to represent a
covalent bond that is partially formed or partially broken.

4. Delocalized Chemical Bonding and Resonance
In chemistry, resonance is a way of describing delocalized electrons within certain molecules or polyatomic
ions where the bonding cannot be expressed by one single Lewis formula. So, resonance is a bonding
behaviour in which a molecule or ion with such delocalized electrons is represented by several contributing
structures also known as resonating structures or canonical forms. Each contributing structure can be
represented by a Lewis structure, with only an integer number of covalent bonds between each pair of atoms
within the structure. Several Lewis structures are used collectively to describe the actual molecular structure.
However these individual contributors are not real structures. So, these cannot be observed in the actual
resonance-stabilized molecule; the molecule does not oscillate back and forth between the contributing
structures, as might be assumed from the word "resonance". The actual structure is an approximate
intermediate between the various canonical forms, but its overall energy is lower than each of the
contributors. This intermediate form between different contributing structures is called a resonance
hybrid.The point of importance is that the contributing structures differ only in the position of electrons, not
in the position of nuclei.


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4.1 Resonance Energy Concept along with Development
Resonance is a key of valence bond theory. Electron delocalization lowers the potential energy of the
substance and thus makes it more stable than any of the contributing structures. The difference between the
potential energy of the actual structure and that of the contributing structure with the lowest potential energy
is called the resonance energy or delocalization energy. Resonance is different from tautomerism and
conformational isomerism, which involve the formation of isomers, thus the rearrangement of the nuclear
positions.


4.2 Characteristics of Resonance
Molecules and ions with resonance have certain basic characteristics.
O–

O
–O

C
O

–O

–O

C

C
O

O

O–

Fig. 10: Contributing structures of the carbonate ion

In diagrams, contributing structures are typically separated by double-headed arrows (←→ ). The arrow
should not be confused with the right and left pointing equilibrium arrow ( ).
[S=C=N

S–C N]


Fig. 11: Contributing structures of the thiocyanate ion

All structures together may be enclosed in large square brackets, to indicate they picture one single molecule
or ion, not different species in a chemical equilibrium. Alternatively to the use of resonance structures in
diagrams, a hybrid diagram can be used. In a hybrid diagram, π bond that are involved in resonance are
usually pictured as curves or dashed lines, indicating that these are partial rather than normal complete π
bonds. In benzene and other aromatic rings, the delocalized π-electrons are sometimes pictured as a solid
circle. Some of the important characteristics are :
1.

They can be represented by several Lewis formulas, called "contributing structures", "resonance
structures" or "canonical forms". However, the real structure is not a rapid interconversion of
contributing structures. To represent the intermediate, a resonance hybrid is used instead.

2.

The contributing structures are not isomers. They differ only in the position of electrons, not in the
position of nuclei.

3.

Each resonating form must have the same number of valence electrons and thus the same total charge,
and the same number of unpaired electrons, if any.

4.

Bonds that have different bond orders in different contributing structures do not have typical bond
lengths. Measurements reveal intermediate bond lengths.


5.

The real structure has a lower total potential energy than each of the contributing structures would have.
This means that it is more stable than each separate contributing structure would be.The gap of the
potential energies of the resonance hybrid and the most stable resonating structure is known as
resonance energy, as stated earlier.


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4.3 Resonance Hybrids
The actual structure of a molecule in the normal quantum state has the lowest possible value of total energy.
This structure is called the "resonance hybrid" of that molecule. The resonance hybrid is the approximate
intermediate of the contributing structures, but the overall energy is lower than each of the contributors, due
to the resonance energy. Any molecule or ion exists in only one form - the resonance hybrid. It does not jump
back and forth between its resonance contributors- looking like one this moment and like another the next
moment.

4.4 Stability of Resonating Forms
One contributing structure may resemble the actual molecule more than another in the sense of energy and
stability. Structures with a low value of potential energy are more stable than those with high values and
resemble the actual structure more. The most stable contributing structures are called major contributors, so
more is the stability of a canonical form, more is its contribution towards the resonance hybrid. Energetically
unfavourable and therefore less probable structures are minor contributors. The stability of a canonical form
depends upon the following factors,
1.

The atoms of the structure must obey as much as possible the octet rule, so that 8 valence electrons must
be present around each atom rather than having deficiencies or surplus.


2.

More is the number of covalent bonds, higher is the stability.

3.

A major contributer is the one that carry a minimum of charged atoms. If unlike charges are present their
separation must be least while for like charges the separation must be maximum.

4.

In case of presence of negative charge, if any, it must be present on the more electronegative atoms and
positive charge, if any, on the most electropositive. So, in the following case, structure II is more stable
than I because the negative charge is placed on more electronegative O atom. Similarly positive
charges, if present, are best occupied on atoms of low electronegativity.

CH2–C–CH3

CH2

C–CH3

O

O

I

II


5.

The greater the number of contributing structures, the more stable the molecule. This is because the
more states at lower energy are available to the electrons in a particular molecule, the more stable the
electrons are. Also the more volume electrons can occupy, the more stable the molecule is. It can be
understood by borrowing a concept of physics, which states that charge dispersed is directly
proportional to stability. Here, electrons can be termed as charged bodies and the more volume they
occupy, more the charge gets dispersed ultimately leading to stability.

6.

Equivalent contributors contribute equally to the actual structure; those with low potential energy (the major
contributors) contribute more to the resonance hybrid than the less stable minor contributors. Especially
when there is more than one major contributor, the resonance stabilization is high.

4.5 Van der Waals Interactions
In chemistry, the Van der Waals force or Van der Waals interaction is the sum of the attractive or repulsive
forces between molecules or between parts of the same molecule other than those due to covalent bonds, or
the electrostatic interaction of ions with one another, with neutral molecules, or with charged molecules.


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Van der Waals forces include attractions and repulsions between atoms, molecules, and surfaces, as well as
other intermolecular forces. As such, Van der Waals forces define many properties of organic compounds,
including their solubility in polar and non-polar solvents etc.For example,in low molecular weight alcohols,
the hydrogen-bonding properties of the polar hydroxyl group dominate the weaker Van der Waals
interactions. In higher molecular weight alcohols, the properties of the nonpolar hydrocarbon chain(s)
dominate and define the solubility. Van der Waals forces quickly vanish at longer distances between
interacting molecules.
Van der Waals forces are relatively weak compared to covalent bonds which is quite expected of them being

the intermolecular forces and not a type of chemical bond. However, these forces play a fundamental role in
chemistry and its different offshoots like supramolecular chemistry, polymer science, nanotechnology, and
surface science. All Van der Waals forces are anisotropic except those between two noble gas atoms. It
means that the magnitude of these forces depend on the relative orientation of the molecules. The induction
and dispersion interactions are always attractive, irrespective of orientation, but the electrostatic interaction
changes sign upon rotation of the molecules. That is, the electrostatic force can be attractive or repulsive,
depending on the mutual orientation of the molecules. The main characteristics are:
1.

They are weaker than normal covalent or ionic bonds.

2.

Van der Waals forces are additive and cannot be saturated.

3.

They have no directional characteristics.

4.

They are all short - range forces and hence only interactions between nearest need to be considered
instead of all the particles. The greater is the attraction if the molecules are closer due to Van der Waals
forces.

5.

Van der Waals forces are independent of temperature except dipole - dipole interactions.

4.5.1 Types of Van Der Waals Forces

Van der Waals forces include a number of interactions. These are discussed below,
1.

Dipole-dipole interaction: A force between two permanent dipoles is known as dipole-dipole
interaction or Keesom force. It can be diagrammatically shown below,
δ+

δ–

δ–

δ+

attraction
Fig. 12: Dipole–dipole interaction

Dipole-Dipole interactions result when two polar molecules approach each other in space. When this occurs,
the partially negative portion of one of the polar molecules is attracted to the partially positive portion of the
second polar molecule. This type of interaction between molecules accounts for many physically and
biologically significant phenomena. An example is shown below:
δ+ δ–
H—Cl

–

δ+ δ
H—Cl

Fig. 13:Dipole–dipole interaction in HCl



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2.

Dipole - induced dipole interaction: A force between a permanent dipole and a corresponding
induced dipole is also known as dipole - induced dipole or Debye force. This type of attractive
interaction also depends on the presence of a polar molecule. However, the second participating
molecule need not be polar as shown below:

+

+

H2O

Xenon

Fig. 14: Example of dipole induced dipole interaction

In the dipole-induced-dipole interaction, the presence of the partial charges of the polar molecule causes a
polarization, or distortion, of the electron distribution of the other molecule. As a result of this distortion, the
second molecule acquires regions of partial positive and negative charge, and thus it becomes polar. The
partial charges so formed behave just like those of a permanently polar molecule and interact favourably with
their counterparts in the polar molecule that originally induced them. Hence, the two molecules attract as
shown below:

δ–

δ+


+

Spherical atom with no dipole,
The dot indicates the location
of the nucleus

Upon approach of a charged ion,
electrons in the atom respond and the
atom develops a dipole.

Fig. 15: Dipole–induced dipole interaction

This interaction also contributes to the intermolecular forces that are responsible for the condensation of
hydrogen chloride gas.
3.

Induced dipole-induced dipole interaction : It is a force between two instantaneously induced
dipoles also known as London dispersion force. This type of interaction acts between all types of
molecule, polar or not. It is the principal force responsible for the existence of the condensed phases of
certain molecular substances, such as benzene, other hydrocarbons, bromine, and the solid elements
phosphorus (which consists of tetrahedral P4 molecules) and sulfur (which consists of crown-shaped S8
molecules). The interaction is called the dispersion interaction or, less commonly, the
induced-dipole-induced-dipole interaction. Two nonpolar molecules of argon are considered near each
other as shown below,


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Ar


Ar

δ– δ+

δ– δ+

Fig. 16: Example of induced–dipole induced dipole interaction

Although there are no permanent partial charges on either molecule, the electron density can be thought of as
ceaselessly fluctuating. As a result of these fluctuations, regions of equal and opposite partial charge arise in
one of the molecules and give rise to a transient dipole. This transient dipole can induce a dipole in the
neighbouring molecule, which then interacts with the original transient dipole as shown here,
–

+

–

+

–

+
+
Repulse

–

Attract


Fig. 17: Induced dipole induced dipole interaction

Although the latter continuously flickers from one direction to another (with an average of zero dipole
overall), the induced dipole follows it, and the two correlated dipoles either attract or repel with one another.

5. Hybridization
120°

109.5°
A

180°
A

A

In chemistry, hybridization is one of the landmark concept of chemical bonding explaining a number of
properties of covalent compounds. It refers to the concept of mixing atomic orbitals into new hybrid orbitals
with different energies, shapes, etc., than the component atomic orbitals suitable for the pairing of electrons
to form chemical bonds in valence bond theory. Hybrid orbitals are very useful in the explanation of
molecular geometry and atomic bonding properties. Although sometimes taught together with the valence
shell electron-pair repulsion (VSEPR) theory, valence bond and hybridization are in fact not related to the
VSEPR model. It is the process of intermixing of atomic orbitals of same atom, having almost similar energies,
followed by redistribution of energies to form new orbitals of identical energies and sizes. The new orbitals
formed are called hybrid orbitals. The number of hybrid orbitals formed is always equal to the number of pure
atomic orbitals employed for hybridization.


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It is not known whether hybridization actually takes place or not but it is a concept which is used to explain

certain observed properties of molecules. The following points are to be remembered about hybridization,
1.

Orbitals belonging to the same atom or ion having almost similar energies get hybridized.

2.

Number of hybrid orbitals is equal to the number of pure atomic orbitals taking part in hybridization.

3.

The hybridization takes place to produce equivalent hybrid orbitals which are degenerate and which
give maximum symmetry.

4.

Hybrid orbitals are always involved in head on overlap, so the type of bonding resulted is always sigma
(σ) bond.

5.1 Important Types of Hybridization
In organic chemistry the various types of hybridization encountered are - sp3, sp2 and sp. These are discussed
in detail.

Fig. 18: Shape of sp3 hybrid lobes

1.

sp3 hybridization: In this type of hybridization, one 's' and three 'p' orbitals of the same value of n
(principle quantum number) mix up to form four sp3 hybridized orbitals. The mixing of orbitals is shown
below,