(Krishnas text book) alok bariyar, r p singh, babita agrawal, babita agrawal, prashant singh organic chemistry i krishna prakashan media (2015)

Structure and Bonding: Hybridization, bond lengths and bond angles, bond energy, localized and delocalized chemical bonding, Van der Waals interactions, inclusion compounds, clatherates

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

KRISHNA Prakashan Media (P) Ltd.

Krishna's

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)

TEXT BOOK

ORGANIC CHEMISTRY-I

By

Alok Bariyar

Ph.D (CSIR fellow), M.Sc.(IIT Delhi), GATE., NET

Ex Scientist (Nuclear)

Bhabha Atomic Research Centre, Mumbai (Maharashtra)

Babita Agrawal

M.Sc., Ph.D.

Head, Dep’t of Chemistry

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

R.P Singh

M.Sc., Ph.D.

Head, Dep’t of Chemistry

KNI of Physical Social Science,

Sultanpur (U.P.)

Prashant Singh

M.Sc., Ph.D.

Asst Prof., Dep’t of Chemistry

KNIPSS, Sultanpur (U.P.)

Ashish Dwivedi

M.Sc., Ph.D.

Asst Prof., Dep’t of Chemistry

Ganpat Sahai P.G College,

Sultanpur (U.P.)

Lord Krishna

e are happy to present this book entitled “Organic Chemistry-I” It has been written

st

according to the latest U.P Unified Syllabus to fulfil the requirement of B.Sc I 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

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.

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.

UNIT-III

5 Alkenes, Cycloalkenes, Dienes and Alkynes: Nomenclature of alkenes, methods of formation,

mechanisms of dehydration of alcohols and dehydrohalogenation of alkyl halids, regioselectivity

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 KMnO , Polymerization of alkenes, Substitution at the allylic and vinylic positions of alkenes, 4Industrial 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

carbon-carbon 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, SN and SN 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.

Unit - I

1 Introduction 03

2 Localised and Delocalized Chemical Bonding 03

3 Characteristics of Covalent Bond 07

$ Solved Examples 10

4 Delocalized Chemical Bonding and Resonance 13 5 Hybridization 18 6 Electron Displacement Effect 29 7 Electromeric Effect 33 8 Strength of Acids and Bases 42

$ Exercise 55

$ Answers 61

$ Hints & Solutions 62

Chapter 2 Mechanism of Organic Reactions O–63-102 1 Fundamental Concepts of Organic Reaction 63 $ Solved Examples 66

2 Various Reaction Intermediates 76 3 Methods of Determination of Reaction Mechanism 96

$ Exercise 99

$ Answers 102

$ Hints & Solutions 102

Chapter 3 Alkanes & Cycloalkanes O–103-134 1 Alkanes 103 2 Cycloalkanes 116 3 Stability of Rings and Ring Strain 124 $ Solved Examples 126

$ Exercise 130

$ Answers 133

$ Hints & Solutions 134

Unit - II

1 Introduction 135

(vii)

Detailed Contents

$ Solved Examples 140

3 Stereoisomerism 140 4 Geometrical Isomerism 159 5 Conformational Isomerism 173 $ Exercise 186

$ Answers 194

$ Hints & Solutions 195

Unit - III Chapter 5 Alkenes, Cycloalkenes Dienes and Alkynes O–201-274 1 Alkenes 201 $ Solved Examples 211

2 Cycloalkene (or Cycloolefin) 233

3 Dienes 238

4 Alkynes 245

$ Exercise 266

$ Answers 273

$ Hints & Solutions 273

Unit - IV Chapter 6 Arenes and Aromaticity O–275-344 1 Arenes 275

$ Solved Examples 277

2 Structure of Benzene 278

3 Aromaticity and Huckel's Rule 283

4 Electrophilic Aromatic Substitution Reactions 287

5 Disubstitution in Benzene Ring and Theory of Substituent Effect 299

6 Fused or Condensed Aromatic Hydrocarbons 312 7 Anthracene 322

8 Biphenyls 329

$ Exercise 331

$ Answers 341

$ Hints & Solutions 341

Chapter 7 Alkyl and Aryl Halide O–345-396 1 Introduction 345 2 Alkyl Halides 345 $ Solved Examples 348

3 Aliphatic Nucleophilic Substitution 354 4. Polyhalogen Compounds 367

5 Aryl Halides 370

6. Synthesis and Uses of DDT 382

7 Synthesis and Uses of BHC 383

$ Exercise 391

$ Answers 395

$ Hints & Solutions 396

❍❍❍

(viii)

Unit-I

Unit-II Unit-III Unit-IV

Chapter 1 : Structure and Bonding

Chapter 2 : Mechanism of Organic Reactions

Chapter 3 : Alkanes and Cycloalkanes

Chapter 4 : Stereochemistry of Organic Compounds

Chapter 5 : Alkenes, Cycloalkenes, Dienes and Alkynes

Chapter 6 : Arenes and Aromaticity

Chapter 7 : Alkyl and Aryl Halide

1 Introduction

Before starting any discussion on organic chemistry, first let us make it very clear that the fundamentalsinvolving bond cleavage, various intermediates and their reactivities, types of reagents, field effects etc formthe backbone of entire organic chemistry These concepts will be useful in understanding the acidic or basicbehaviour of many organic compounds, types of various organic reactions, mechanism of various reactionsand 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 ofdifferent 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, customarilythe chemical bonds are classified into different types

One more point has to be reiterated before discussing the classification of chemical bond Atom may attain astable 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-3

1 Structure and Bonding

2.2 Covalent Bond

In the simplest view of a so-called 'covalent' bond, one or more electrons (often a pair of electrons) are drawninto the space between the two atomic nuclei Here the negatively charged electrons are attracted to thepositive charges of both nuclei, instead of just their own This overcomes the repulsion between the twopositively 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, covalentbonding involves sharing of electrons in which the positively charged nuclei of two or more atomssimultaneously attract the negatively charged electrons that are being shared between them These bondsexist between two particular same or different atoms, and have a direction in space, allowing them to beshown as single connecting lines between atoms in drawings For Example, Two chlorine atoms react to form

a Cl2 molecule

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

The carbon atom is short off four electrons so as to have noble gas structure Consequently, it forms fourbonds with the chlorine atoms which themselves are short of one electron so they each form one bond bysharing electrons In this way, both carbon and all four chlorine atoms attain a noble gas structure

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

In a polar covalent bond, one or more electrons are unequally shared between two nuclei Covalent bondsoften 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 weakerthan the covalent bonds that hold the molecules internally together Such weak intermolecular bonds giveorganic 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 asdiamond or quartz or the silicate minerals in many types of rock) then the structures that result may be bothstrong and tough, Also, the melting points of such covalent polymers and networks increase greatly

Cl + Cl Cl Cl

C + 4 C Cl C

ClClCl

N + 3[H ] H N H

H

H O , H FH

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 orbitalsalong the internuclear axis the resulting bond is called sigma(σ) bond Such type of overlap is alsoknown as end to end or head on overlap It is a strong bond and cylindrically symmetrical Theoverlapping 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 filleds–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 asigma bond

(ii) s— p overlapping: In this case, half filled s–orbital of one atom overlaps with the half filledp–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

(iii) p—p overlapping: This type of overlapping occurs when p–orbital of one atom overlaps withthe 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

2 Pi (π) bond: Pi (π)bond is formed by lateral or sidewise overlapping of p orbitals Sideways overlapmeans 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,

1s arbital of H–atom

1s arbital of H–atom

(i) In case of oxygen molecule (each oxygen atom having electronic configuration, 1s 2s 2p 2p 2p2 2 2x 1y 1z), the two atoms are held together by one σ-bond and one π-bond.

(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 1s 2s 2p 2p 2p2 2 1x 1y 1z

It is important to remember that the 's' orbitals can only form σ-bonds, whereas the p, d & forbitals 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 alsoshare two or three pairs of electrons in the aptly named double and triple bonds The first bond between twoatoms 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 covalentinteraction between two atoms is the number of electron pairs that are shared between them Single bondshave a bond order of 1, double bonds 2, and triple bonds 3 Bond order is directly related to bond strengthand bond length Higher order bonds are stronger and shorter, while lower order bonds are weaker andlonger The Lewis structures for some common molecules involving multiple covalent bonds can be foundbelow

O molecule2(O=O)

Fig 4: Formation of O molecule 2

2.5 Co–ordinate Bond

A covalent bond results from sharing of a pair of electrons between two atoms, where each atom contributesone electron to the bond It is also possible to have an electron pair bond where both the electrons come fromone of the two binding atoms and there is no contribution from the other atom Such bonds are calledco–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 ofammonium ion Even though the ammonia molecule has a stable electronic configuration it can react with ahydrogen ion (H+) by donating a lone pair of electrons from N atom to H+ ion forming the ammonium ion

NH4+

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

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

3 Characteristics of Covalent Bond

Some important characteristics of covalent bonds like bond length, bond angle and bond energy arediscussed below:

+orH

H

HH

H

HN+

Formaldimine Formaldehyde

cyanide

carbonmonoxide

Fig 6: Lewis structures of molecules with multiple bonds

FF

F

FBF

–

F

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 themolecule Bond length is related to bond order, when more electrons participate in bond formation the bondwill get shorter Bond length is also inversely related to bond strength and the bond dissociation energy, as allother things being equal a stronger bond will be shorter In a bond between two identical atoms half the bonddistance is equal to the covalent radius Bond lengths are measured in the solid phase by means of X-raydiffraction, 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 isdifferent from that in methyl chloride It is however possible to make generalizations when the generalstructure is the same

The actual bond length between two atoms in a molecule depends on such factors as the orbital hybridizationand 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 thesigma bond and presence of π bond

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)

3.2 Bond Angle

Molecular geometry is the three-dimensional arrangement of the atoms that constitute a molecule Itdetermines 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 betweenthe plane formed by the first three atoms and the plane formed by the last three atoms As a matter of fact, theangles between bonds that an atom forms also depend on the rest of molecule, albeit weakly Some of themolecules alongwith bond angles are shown below,

Fig 9: Different molecular shapes along with bond angles

120°

90°

90°

90°

Example1: Pre dict all bond an gles in the fol low ing mol e cules.

(i) CH 3 Cl (ii) CH 3 CN (iii) CH 3 COOH

So lu tion: (i) The Lewis structure of methyl chloride is:

In the Lewis structure of CH3Cl carbon is surrounded by four regions of high electron density, each of whichforms a single bond Based on the VSEPR model, we predict a tetrahedral distribution of electron cloudsaround 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 torepresent a bond projecting forward from the plane of the paper

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

The methyl group, CH3-, is tetrahedral The carbon of the -CN group is in the middle of a straight linestretching from the carbon of the methyl group through the nitrogen

(ii) The Lewis structure of acetic acid is:

HH

H

ClC

HH

109.5°

180°

N

HH

H

CC

O

O H

S olved E xamples

Both the carbon bonded to three hydrogens and the oxygen bonded to carbon and hydrogen are centers oftetrahedral 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, andaround 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 carbonatom and 4 hydrogen radicals divided by 4 Bond energy (E) must not be confused with bond-dissociationenergy 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 inmethane is slightly different than the bond energy of C-H in ethane One can calculate a more general bondenergy 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)

C—Br 276 F—F 154 Si–C 360

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-foldhigher)

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

Ex am ple 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

2H 2 O(g)→2H 2 +O 2 (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.

So lu tions (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 thereactants side is 4×460 kJ/mol = 1840 kJ/mol.The sum of the energies released to form the bonds onthe 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 1840kJ/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

(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 greatersize 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 bondingpartners Three dimensional configurations are best viewed with the aid of models In order to represent suchconfigurations 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 bondedatoms In most cases the focus of configuration is a carbon atom, so the lines specifying bond directions willoriginate 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,

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 di rected 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 man ner as we have

de fined the hatched bond, but this can be con fus ing be cause the dashed bond is of ten used to rep re sent a

co va lent bond that is par tially formed or par tially 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 atomswithin 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 actualresonance-stabilized molecule; the molecule does not oscillate back and forth between the contributingstructures, as might be assumed from the word "resonance" The actual structure is an approximateintermediate 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

CA

BDA

normal bondwedge bond

dashed bondhatched bond

or}

4.1 Resonance Energy Concept along with Development

Resonance is a key of valence bond theory Electron delocalization lowers the potential energy of thesubstance and thus makes it more stable than any of the contributing structures The difference between thepotential 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 andconformational isomerism, which involve the formation of isomers, thus the rearrangement of the nuclearpositions

4.2 Characteristics of Resonance

Molecules and ions with resonance have certain basic characteristics

In diagrams, contributing structures are typically separated by double-headed arrows (← → ) The arrowshould not be confused with the right and left pointing equilibrium arrow ( )

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 indiagrams, a hybrid diagram can be used In a hybrid diagram, π bond that are involved in resonance areusually pictured as curves or dashed lines, indicating that these are partial rather than normal complete πbonds In benzene and other aromatic rings, the delocalized π-electrons are sometimes pictured as a solidcircle Some of the important characteristics are :

1 They can be represented by several Lewis formulas, called "contributing structures", "resonancestructures" or "canonical forms" However, the real structure is not a rapid interconversion ofcontributing 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 theposition 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 bondlengths 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 thepotential energies of the resonance hybrid and the most stable resonating structure is known asresonance energy, as stated earlier

S–C N]

[S=C=N

Fig 11: Contributing structures of the thiocyanate ion

CO

OO

O

OO–

–

CO

O O ––

Fig 10: Contributing structures of the carbonate ion

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 approximateintermediate 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 jumpback and forth between its resonance contributors- looking like one this moment and like another the nextmoment

4.4 Stability of Resonating Forms

One contributing structure may resemble the actual molecule more than another in the sense of energy andstability Structures with a low value of potential energy are more stable than those with high values andresemble the actual structure more The most stable contributing structures are called major contributors, somore is the stability of a canonical form, more is its contribution towards the resonance hybrid Energeticallyunfavourable and therefore less probable structures are minor contributors The stability of a canonical formdepends 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 andpositive charge, if any, on the most electropositive So, in the following case, structure II is more stablethan I because the negative charge is placed on more electronegative O atom Similarly positivecharges, if present, are best occupied on atoms of low electronegativity

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

6 Equivalent contributors contribute equally to the actual structure; those with low potential energy (the majorcontributors) contribute more to the resonance hybrid than the less stable minor contributors Especiallywhen 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 repulsiveforces between molecules or between parts of the same molecule other than those due to covalent bonds, orthe electrostatic interaction of ions with one another, with neutral molecules, or with charged molecules

CH –C–CH2 3 CH2 C–CH3

I II

Van der Waals forces include attractions and repulsions between atoms, molecules, and surfaces, as well asother 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 Waalsinteractions 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 betweeninteracting molecules

Van der Waals forces are relatively weak compared to covalent bonds which is quite expected of them beingthe intermolecular forces and not a type of chemical bond However, these forces play a fundamental role inchemistry and its different offshoots like supramolecular chemistry, polymer science, nanotechnology, andsurface science All Van der Waals forces are anisotropic except those between two noble gas atoms Itmeans that the magnitude of these forces depend on the relative orientation of the molecules The inductionand dispersion interactions are always attractive, irrespective of orientation, but the electrostatic interactionchanges 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 consideredinstead of all the particles The greater is the attraction if the molecules are closer due to Van der Waalsforces

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-dipoleinteraction or Keesom force It can be diagrammatically shown below,

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 thesecond polar molecule This type of interaction between molecules accounts for many physically andbiologically significant phenomena An example is shown below:

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

In the dipole-induced-dipole interaction, the presence of the partial charges of the polar molecule causes apolarization, or distortion, of the electron distribution of the other molecule As a result of this distortion, thesecond molecule acquires regions of partial positive and negative charge, and thus it becomes polar Thepartial 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 asshown below:

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 induceddipoles also known as London dispersion force This type of interaction acts between all types ofmolecule, polar or not It is the principal force responsible for the existence of the condensed phases ofcertain molecular substances, such as benzene, other hydrocarbons, bromine, and the solid elementsphosphorus (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, theinduced-dipole-induced-dipole interaction Two nonpolar molecules of argon are considered near each other as shown below,

Fig 14: Example of dipole induced dipole interaction

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

+

δ

–

δ

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 inone of the molecules and give rise to a transient dipole This transient dipole can induce a dipole in theneighbouring molecule, which then interacts with the original transient dipole as shown here,

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

5 Hybridization

In chemistry, hybridization is one of the landmark concept of chemical bonding explaining a number ofproperties of covalent compounds It refers to the concept of mixing atomic orbitals into new hybrid orbitalswith 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 ofmolecular geometry and atomic bonding properties Although sometimes taught together with the valenceshell electron-pair repulsion (VSEPR) theory, valence bond and hybridization are in fact not related to theVSEPR 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 orbitalsformed are called hybrid orbitals The number of hybrid orbitals formed is always equal to the number of pure atomic orbitals employed for hybridization

A

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

1 sp 3 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 shownbelow,

3 Fig 18: Shape of sp hybrid lobes

The sp3 hybrid orbitals have 25% 's' character and 75% 'p' character These orbitals orient themselvestowards the corners of a regular tetrahedron The angle between the orbitals is 109°28' The commonestexample is CH4 as shown below,

As such, in any organic compound if the atom is containing 4σ bonds or 3σ bonds plus one lone pair ofelectrons or 2σ bonds plus 2 lone pairs of electrons, then the the type of hybridization undergone by the atom

is sp3

2 sp 2 hybridization: This is a second type of hybridization involving pure one 's' and two 'p' atomicorbitals of the same principle quantum number These pure atomic orbitals mix together to form three

sp2 hybridized orbitals The mixing of orbitals can be shown as,

Fig 20: Shape of methane

The sp2 hybrid orbitals have 33.3% 's' character and 66.6% 'p' character These

orbitals orient themselves towards the corners of an equilateral triangle, as

shown above It is important to remember that one p orbital perpendicular to

the molecular plane is unhybridised An inorganic example is BF3,where the

central boron atom is sp2 hybridized

In cases like ethene (C2H4), the bonding can be shown as,

Fig 24: Bond line notation of C H 2 4

H p

2

sp H p

H p

Fig 25: Hybrid orbital arrangement of C H 2 4

2

Three sphybrid orbitals

2

Three sphybrid orbitals in one plane

2

Fig 22: Formation of sp hybrid orbitals

These two p-orbitals overlap in sideways manner to give rise to a double bond Both and framework bonding

is shown

It is always important to remember that when a bond is formed with head-to-head overlap, it is called a sigma(σ) bond Similarly,when a bond is formed with sideways overlap, it is called a pi (π) bond As such, in anyorganic compound if the atom is containing 3σ bonds or 2σ bonds plus one lone pair of electrons, then thethe type of hybridization undergone by the atom is sp2

3 sp hybridization: In this type, one 's' and one 'p' orbital of the same principle quantum number mix toform two sp hybridized orbitals, for example in BeCl2, Be is sp hybridized The mixing of s and p orbitalscan be shown as,

Fig 27: Shape of sp hybrid lobes

(a)C H σ–bonded framework 2 4 (b)C H π bonding 2 4

Fig 26: σ and π bond arrangement of C H 2 4

They possess 50% 's' and 50% 'p' character In cases like ethyne C2H2, the central carbon atom is sphybridized Ethyne has the following Lewis structure.

Hybridization wise, each carbon atom is sp hybridized having two perpendicular p orbitals with two unpairedelectrons present on each carbon atom These overlap in sideways fashion and produce a triple bond All theoverlaps are shown below,

As such, in any organic compound if the atom is containing 2σ bonds or 1σ bonds plus one lone pair ofelectrons, then the type of hybridization undergone by the atom is sp

5.2 Inclusion Compounds

In host-guest chemistry, the definition of inclusion compounds is very broad, extending to channels formed

between molecules in a crystal lattice in which guest molecules can fit As such, an in clu sion com pound is

a com plex com prised of two chem i cal spe cies, one chem i cal com pound is the "host" which gen er ally forms acav ity An other is the "guest" com pound which are lo cated in the cav i ties of the host If the spaces in the hostlat tice are en closed on all sides so that the guest spe cies is 'trapped' as in a cage, the com pound is known as aclathrate In this type of mo lec u lar en cap su la tion a guest mol e cule is ac tu ally trapped inside anothermolecule

σ-bond

π-bondπ-bond

5.3 Clathrate Compounds

A clathrate is a chemical substance consisting of a lattice that traps or contains molecules The word

clathrate is derived from the Latin clatratus meaning with a lattice Traditionally, clathrate compounds arepolymeric and completely envelop the guest molecule, but in modern usage clathrates also include

host-guest complexes and inclusion compounds According to IUPAC convention, clathrates are

Inclusion compounds in which the guest molecule is in a cage formed by the host molecule or by a lattice ofhost molecules

Traditionally clathrate compounds refer to polymeric hosts containing molecular guests More recently, theterm refers to many molecular hosts, including some inorganic polymers such as zeolites Many clathrates arederived from an organic hydrogen-bonded frameworks These frameworks are prepared from molecules that

"self-associate" by multiple hydrogen-bonding interactions The most famous clathrates are methaneclathrates where the hydrogen-bonded framework is contributed by water and the guest molecules aremethane Large amounts of methane naturally frozen in this form exist both in permafrost formations andunder the ocean sea-bed Other hydrogen-bonded networks are derived from hydroquinone, urea, andthiourea A much studied host molecule is Dianin's compound which is shown below,

5.4 Charge Transfer Complexes

Charge-transfer complexes exist in many types of molecules, inorganic as well as organic, and in solids,liquids, and solutions A well-known example is the iodine starch complex formed by iodine when combinedwith starch, which exhibits an intense blue charge-transfer band A charge-transfer complex or CT complex

or electron-donor-acceptor complex is an association of two or more molecules, or of different parts of

one large molecule, in which a fraction of electronic charge is transferred between the molecular entities Theresulting electrostatic attraction provides a stabilizing force for the molecular complex The source moleculefrom which the charge is transferred is called the electron donor and the receiving species is called theelectron acceptor The nature of the attraction in a charge-transfer complex is not a stable chemical bond,and is thus much weaker than covalent forces Many such complexes can undergo an electronic transitioninto an excited electronic state The excitation energy of this transition occurs very frequently in the visibleregion of the electro-magnetic spectrum, which produces the characteristics intense color for thesecomplexes These optical absorption bands are often referred to as charge-transfer bands (CT bands) Optical spectroscopy is a powerful technique to characterize charge-transfer bands

Fig 31: Dianin's compound

5.5 Hydrogen Bonding

Hydrogen bond is a weak force of attraction formed between an H atom and highly electronegative atom like

O, N or F either of the same molecule or of a different molecule but to which it is not directly attached.Moreover the H atom itself should be attached to a highly electronegative atom like O, N or F HF isconsidered as an example The hydrogen fluorine bond in hydrogen fluoride is polar covalent bond asfluorine is a strongly electronegative element As a result fluorine acquires a partial negative charge andhydrogen acquires a partial positive charge

H F

δ+ δ−

−The lone pair on the fluorine atom in another molecule of hydrogen fluoride will attract the positive charge on hydrogen in a molecule of hydrogen fluoride electrostatically This bond between hydrogen and fluorine ofdifferent molecules is known as hydrogen bond This type of linkage is represented by dotted lines

Here, it is important to remember that hydrogen bonding is not a type of chemical bond rather it is anintermolecular force of attraction

5.5.1 Conditions for Hydrogen Bonding

1 The molecule must contain a highly electronegative atom linked to hydrogen atom

2 The size of electronegative atom should be small These conditions are met only by F, O and N atoms.Although Cl has the same electronegativity as nitrogen, it does not frequently form effective hydrogen bond This

is because of its larger size than that of N with the result that its electrostatic attractions are weak

5.5.2 Types of Hydrogen Bonding

Generally the hydrogen bonds are classified into two types -

OO

O

OO

OO

OO

O

carbonyl groups suckelectron density awayfrom the acceptor unit

donar

acceptor

nitrogen atoms have a higher electron densitythan the carbonyl groupsand lend it to theacceptor

1 Intermolecular hydrogen bonding

2 Intramolecular hydrogen bonding

1 Intermolecular hydrogen bonding: In this type of hydrogen bonding, the two or more than twomolecules of same or different compounds combine together to give a polymeric aggregate Some ofthe examples are,

2 Intramolecular hydrogen bonding: In this type, hydrogen bonding occurs within two atoms of thesame molecule This type of hydrogen bonding results into ring formation commonly known aschelation Organic compounds frequently show this type of H bonding Some of the examples are,

5.5.3 Effects of Hydrogen Bonding

1 Boiling point and melting point: The high boiling point of water can be explained on the basis ofhydrogen bonding.Water molecules associate through extensive network of hydrogen bonding andthus require more energy to break the intermolecular forces Similarly we can explain the higher boilingpoint of alcohols, amines, amides etc Moreover, H2O is a liquid but H2S is a gas because H2O can formhydrogen bond but HS cannot

O

o–nitrophenol

Up till now those examples are considered where boiling point is increased due to hydrogen bonding.Normally, this happens in the case of intermolecular hydrogen bonding However, there are certaincases where the boiling point is decreased due to hydrogen bonding It happens in the case ofintramolecular hydrogen bonding For example o–hydroxy, nitro, carbonyl, carboxylic or chlorocompounds have lower melting and boiling points than the respective meta or para isomers Theexplanation is intramolecular hydrogen bonding taking place between the two groups placed ortho toeach other It prevents the association of the molecule with other neighbouring molecules As a result,the intermolecular forces of attraction decreases which eventually leads to decreased melting andboiling points.

On the other hand, in case of meta– and para–isomers, intramoelcular hydrogen bonding is not possible because of the bigger size of the ring formed Rather, these molecules indulge into intermolecularhydrogen bonding which causes association This association of molecules eventually leads to higher

mp and bp for meta and para isomers

2 Water solubility: A substance is said to be soluble in water if it is capable of forming hydrogenbonding with water molecules For example alcohols are soluble in water due to hydrogen bonding

It is important to note that when alkyl groups in alcohols are bulkier having four or more carbon atoms,the alkyl group predominates over OH group So, hydrogen bond forming ability is subdued with theresult that the solubility of such alcohols in water decreases On the contrary side, when a compoundhas a large ratio of –OH groups to hydrocarbon group, it will be highly water soluble Thus sugar, certain starches and polyvinyl alcohol are quite soluble in water Solubility of other compounds (organic orinorganic) having one or more than one –OH group can be explained on the basis of hydrogenbonding Detergents are also soluble in water due to solubilizing effect of hydroxyl groups

It is important to note that while the intermolecular hydrogen bonding increases solubility of thecompound in water, the intramolecular hydrogen bonding decreases This is due to fact that theformation of intramolecular hydrogen bond prevents hydrogen bonding of the compound with water It eventually reduces solubility of the compound in water

For Example: o–nitrophenol is less sol u ble in wa ter then p–nitrophenol be cause due to intramolecular

hy dro gen bond ing –OH group is not avail able to form hy dro gen bond with water

H – O H – O H – O H – O

RH

o–nitrophenolO

o–chlorophenol

o–hydroxybenzoic acid

O – H

Intramolecular Hydrogen bondingCl

OHO

HO

N

OO

NH

3 Volatile nature: A compound having high boiling point is less volatile and vice versa Asintermolecular hydrogen bonding increases the boiling point so it can be concluded that volatile nature

of a compound decreases if intermolecular hydrogen bonding is present On the other side, the volatility increases if intramolecular hydrogen bonding is present

5.6 Aromaticity

In organic chemistry, aromaticity is formally used to describe an unusually stable nature of some planar or

flat rings of atoms such that the ring may be homonuclear as in case of benzene or hetronuclear as in case ofpyridine

These structures contain a number of double bonds that interact with each other according to certain rules As

a result of their being highly stable, such rings tend to form easily, and once formed, tend to be difficult tobreak in chemical reactions Since one of the most commonly-encountered aromatic system of compounds

in organic chemistry is based on derivatives of benzene which is commonly found in petroleum, the word

"aromatic" is occasionally used to refer informally to benzene derivatives, and this is how it was first defined.Nevertheless, many non-benzene aromatic compounds do exist as well In living organisms, for example, themost common aromatic rings are the double-ringed bases in RNA and DNA, and interestingly benzene andits derivatives are too rare in biology Huckel's rule helps us to decide that a given compound would bearomatic or not by observing its structure According to this rule, for a compound to be aromatic, it mustsatisfy the following conditions,

1 There must be cyclic delocalization in the compound

2 The closed loop formed due to cyclic delocalization must have a total of (4n + 2)π electrons where n is

an integer having values 0, 1, 2, 3…… The (4n + 2)π electrons is called Hukel number

The first condition of cyclic delocalization takes into account the conditions required to show resonance sothat the compound must have conjugation and the compound must be planar or nearly planar If both theabove conditions are fulfilled, then the compound would be aromatic and if any of the above condition fails,then it is non-aromatic

C

CCCCC

HH

H

H

HH

C

CCCNC

H

HH

HH

O–H

NO

is not available for intermolecular bonding

First of all, let us take the simplest example of benzene As it is evident from the structure that there is

conjugation in the ring leading to cyclic delocalization Moreover,the closed loop thus formed

would have 6π electrons which corresponds to n=1 in the Huckel's rule Thus, benzene is said to be

aromatic compound The same would be true for phenol and nitrobenzene also

The rule is discussed in details in chapter VI

6 Electron Displacement Effect

In organic chemistry, there are some important mechanisms of electron displacement like field and inductiveeffect, electromeric effect, mesomeric effect and hyperconjugation These effects are discussed below one byone

6.1 Field and Inductive Effect

Inductive Effect is an experimentally observed effect of the transfer of charge - negative or positive, through achain of atoms in a molecule

The permanent dipole induced in one bond by another is called inductive effect The electron cloud in aσ-bond between two unlike atoms is not uniform and is slightly displaced towards the more electronegative of the two atoms This causes a permanent state of bond polarization, where the more electronegative atom has

a slight negative charge (δ-) and the other atom has a slight positive charge (δ+) If the electronegative atom is then joined to a chain of atoms, usually carbon, the positive charge is relayed to the other atoms in the chain.This is the electron-withdrawing inductive effect, also known as the -I effect Some groups, such as the alkylgroup, are less electron-withdrawing than hydrogen and are therefore considered as electron-releasing This

is electron releasing character and is indicated by the +I effect In short, alkyl groups tend to give electrons,leading to induction effect

Explanation

First let us look at the C–C bond in ethane This C–C bond has no polarity because it connects two similar atoms.However, the C–C bond in ethyl chloride is polarized by the presence of the electronegative chlorine atom Thispolarization is actually the sum of three effects In the first of these, the Cl atom have been deprived of some of itselectron density by the higher electronegativity of Cl Secondly, the electron deficiency of C1 is partiallycompensated by drawing the C–C electrons closer to itself, resulting in polarization of this bond and a slightpositive charge on the C2 atom Thirdly, the polarization of the C–C bond causes a (slight) polarization of the threemethyl C–H bonds The effect of C1 on C2 is less than the effect of Cl on C1 atom

2

6.2 Field Effect

In addition to such inductive effect operating through the σ–bonds in a compound, an analogous effect canalso operate either through the space surrounding the molecule or in solution via the molecules of solventthat surround it which is called field effect A point of distinction between the two effects is that the inductiveeffect depends only on the nature of bonds while the field effect depends on the geometry of the molecule.However, in many cases, it is not possible to distinguish inductive effect with this field effect But in general,reference to an inductive effect is assumed to include any such field effect

6.3 Characteristics of I Effect

Some important characteristics of inductive effect are;

1 It is a permanent electron displacement effect It is because no attacking reagent (electrophile ornucleophile)is required to cause the electron displacement

2 It always causes partial charge separation as shown below,

3 It decreases with increases of distance

4 It does not involve delocalization of π-electrons Rather,it involves displacement of σ-electrons only

6.4 Types of Inductive Effect

The inductive effect is divided into two types depending on their strength of electron withdrawing or electronreleasing nature with respect to hydrogen

1 Negative inductive effect (-I): The electron withdrawing nature of groups or atoms is called asnegative inductive effect It is indicated by -I Following are the examples of groups in the decreasingorder of their -I effect:

NH3+ > NO2 > CN > SO3H > CHO > CO > COOH > COCl > CONH2 > F > Cl > Br > I > OH > OR >

NH2 > C6H5 > H

2 Positive inductive effect (+I): It refers to the electron releasing nature of the groups or atoms and isdenoted by +I Following are the examples of groups in the decreasing order of their +I effect C(CH3)3 > CH(CH3)2 > CH2CH3 > CH3 > H

For measurements of relative inductive effects, hydrogen is chosen as reference point in the molecule CR3

-H If the H atom in this molecule is replaced by X (an atom or group of atoms), the electron density in the CR3part of the molecule is less than in CR3 - H, then X is said to have a -I effect or electron withdrawing or electron attracting inductive effect If the electron density in the CR3 part is greater than in CR3 - H, then X is said tohave a +I effect or electron releasing or electron repelling inductive effect Deuterium is also electrondonating with respect to hydrogen Generally speaking, all alkyl groups exhibit +I effect when they areattached to an unsaturated or trivalent carbon or other atom Keeping other things equal, atoms with spbonding generally have a greater electron withdrawing ability than those with sp2 bonding, which in turn

have more electron-withdrawing power than with sp3 bonding This reasonably accounts for the fact thatarylic, vinylic and alkynyl groups exhibit -I effect.

Ex am ple 3: Why alkyl groups are show ing pos i tive in duc tive ef fect?

Solution: Though the C-H bond is prac ti cally con sid ered as non-po lar, there is par tial pos i tive charge on

hy dro gen atom and par tial neg a tive charge on car bon atom There fore each hy dro gen atom acts as elec tron

do nat ing group This cu mu la tive do na tion turns the alkyl moi ety into an elec tron do nat ing group

6.5 Applications of Inductive Effect

Some of the prominent applications of inductive effect are discussed below

1 Stability of carbonium ions: The stability of carbonium ions increases with increase in number ofalkyl groups due to their +I effect The alkyl groups release electrons to carbon, bearing positive chargeand thus stabilize the ion The order of stability of carbonium ions is:

2 Stability of free radicals: In the same way the stability of free radicals increases with increase in thenumber of alkyl groups Thus the stability of different free radicals is:

3 Stability of carbanions: However the stability of carbanions decreases with increase in the number

of alkyl groups since the electron donating alkyl groups destabilize the carbanions by increasing theelectron density.Thus the order of stability of carbanions is:

R

+

RR

R

R C+

H

HH

a secondarycarbocation

a tertiarycarbocation

CH

least stable Primary radical Secondary radical Tertiary radical

4 Acidic strength of carboxylic acids and phenols: The electron withdrawing groups (-I) decreasethe negative charge on the carboxylate ion and thus by stabilizing it Hence the acidic strength increaseswhen -I groups are present However the +I groups decrease the acidic strength For example, (i) The acidic strength increases with increase in the number of electron withdrawing Fluorineatoms as shown below

CH3COOH < CH2FCOOH < CHF2COOH < CF3COOH

(ii) Formic acid is stronger acid than acetic acid since the -CH3 group destabilizes the carboxylateion Generally in gaseous medium, the order of basic strength is,

(iii) On the same lines, the acidic strength of phenols increases when-I groups are present on the ring e.g The p-nitrophenol is stronger acid than phenol since the -NO2 group is a -I group andwithdraws electron density Whereas the para-cresol is weaker acid than phenol since the -CH3group shows positive (+I) inductive effect

5 Basic strength of amines: The electron donating groups like alkyl groups increase the basic strength

of amines whereas the electron withdrawing groups like aryl groups decrease the basic nature.Therefore alkyl amines are stronger Lewis bases than ammonia, whereas aryl amines are weaker thanammonia Thus the order of basic strength of alkyl and aryl amines with respect to ammonia is :CH3NH2

> NH3 > C6H5NH2 Generally in gaseous medium, the order of basic strength for primary, secondaryand tertiary amines with respect to ammonia is,

Ex am ple 4: The in duc tive ef fect can be best de scribed as:

(i) The conjugation of σ-bonding orbital with the adjacent π-orbital

(ii) The ability of atom or group to cause bond polarization

(iii) The transfer of lone pair of electrons from more electronegative atom to lesser electronegative atom in a molecule

(iv) All of the above.

Solution. (iii)

Ex am ple 5: Which of the fol low ing state ment is in cor rect about the in duc tive ef fect ?

(i) It is a permanent effect

(ii) It decreases with increases of distance

(iii) It involves delocalization of π-electrons

(iv) It involves displacement of σ-electrons.

as the reagent is removed, the polarized molecule will come back to the original state as shown below,

Since A has lost its shared pair of electrons, it acquires a positive charge while B has gained the electron pair,thus it acquires a negative charge.

The most common example illustrating the electromeric effect is the reaction of an alkene with Br2 in CCl4

In this reaction, when the attacking reagent (bromine) approaches alkene, the temporary polarizationdevelops on the alkene with C2 atom gaining a negative charge and Cl atom acquiring positive charge as itcan be compensated by the +I effect of R group The alkenes being electron rich compounds, due to thepresence of π electron cloud, are attacked by the electrophile (Br+) to give a cyclic bromonium ion Here, theformation of cyclic bromonium ion as intermediate is possible because bromine is of considerably large sizehaving lone pairs to be bonded to both the carbons simultaneously The cyclic bromonium ion is thenattacked by Br− from the top (as lower side is already blocked) whereby the three membered ring is cleaved

by trans opening giving vicinal dibromide as the product

7.1 Types of Electromeric Effect

Like inductive effect, electromeric effect is also of two types - +E and -E effect, as discussed below in details

+E Effect: If the attacking species is an electrophile, the π electrons are transferred towards the positivelycharged atom This is the +E effect

An example is the protonation of ethene When the H+ comes near the double bond, the bond is polarizedtowards the proton

-E Effect: If the attacking reagent is a nucleophile, the electrons are transferred away from the attackingreagent and into the π system This is the -E effect and it is being shown below,

C=CR

Br– +

cyclic bromonium ion vicinal dibromide

H +E effect

H C=CH2 2 + H+ H C–CH2 2

+

δ δ–

+ H+ H C–CH2 + 3