chapter basic conceptsnewagepublishers.com/samplechapter/001836.pdf · · 2011-09-29for diatomic...

The outlook towards organic chemistry has undergone great changes in recent years. Besidessynthesis and the determination of the structure of molecules, the study, both chemical andkinetic, of organic reactions has been given considerable attention with a view to understandingthe mechanism of reactions. Organic chemists continue to be engaged in the search for newreagents and new reactions. Yet, simultaneously, old reactions are being reinvestigated in orderto establish their mechanisms. This has been possible by new experimental, analytical andspectroscopic techniques. Lewis first presented the concept of covalent bond in chemistry. Thecovalent bond was subsequently classified into polar and non-polar, depending on theelectronegativities of the bonded atoms. A knowledge of the nature and strength of bonds isessential for the chemical investigation of organic molecules. Moreover, the properties of amolecule are always influenced by its structure. A brief description of certain basic conceptshas been attempted in this chapter.

1.1 BOND DISSOCIATION ENERGY

It is a well known fact that when a covalent bond is formed between two atoms, a certainamount of energy is released. For instance, when hydrogen atoms combine to form hydrogenmolecules, 104 kcal of heat is evolved for each mol of hydrogen formed. The bond dissociationenergy instead, is the energy required (∆H° at 25°C and 1 atm pressure) in cleaving a bond intotwo radicals. In the case of hydrogen molecule, a certain amount of energy is expended insplitting the molecule, i.e.,

H — H (g) 2H•(g), ∆H ° = 104.0 kcal/mol

The dissociation energy of the H—H bond is thus 104 kcal/mol. Such a process is anendothermic process, therefore, ∆H° will be a positive value. This information is obtained bymolecular spectroscopy, thermal methods and electron impact studies.1 Walsh argued theoreticallythat the strength of a covalent bond increased with increase in the difference of electronegativityof the bonded atoms. This is apparent from the values listed in Table 1.1 (cf. carbon—halogenbonds).

The bond dissociation energy is not equal to the bond energy. The definition of bondenergy is though less precise, since it is the average value of the energy required to break abond between two atoms in a molecule. Let us, for instance, consider a water molecule whichcontains two H — O bonds. It is possible to measure the bond dissociation energy for theremoval of one and then the second H atom from H2O. The result is computed as follows:

H — O — H (g) O — H•(g) + H

•(g), ∆H ° = D

1 = 119.7 kcal/mol

O — H•(g) O

•(g) + H

•(g), ∆H ° = D2 = 101.5 kcal/mol

1

BASIC CONCEPTS

CHAPTER

1

2 ORGANIC REACTION MECHANISMS

It is apparent that the bond dissociation energy for the two steps is different but fromHess’s Law one can state that their sum is equal to the energy of atomization (conversion ofa molecule to its atoms) of the water molecule, i.e., 221.20 kcal/mol.

H2O (g) O

•(g) + 2H

•(g), D

1 + D

2 = 221.20 kcal/mol

In water molecule there are two O—H bonds present that are equivalent and energy valuecan be assigned to each bond which is equal to the average value of D

1 and D

2.

1

2

(D1 + D

2) 110.6 kcal/mol.

This is the bond energy and is often referred to as the mean theoretical bond energy andshould be distinguished from the bond dissociation energy.

For diatomic molecules such as H2, Cl

2, HCl, etc. bond energies and bond dissociation

energies are equal. The bond dissociation energy for the same molecule may differ from moleculeto molecule (see Table 1.1). The higher the bond energy, the stronger is the bond. Also increasings-character shortens a bond so bond strength increases with s-character.

In a chemical reaction a new bond is usually formed while the old one is broken. It is,therefore, important to know the bond dissociation energy, which is a measure of the strengthof a bond, to understand a chemical reaction. This may also reflect on the nature of thereaction.

TABLE 1.1: Bond Dissociation Energies (kcal/mol)

Bond Bond Dissociation Bond Bond DissociationEnergy (∆∆∆∆∆H°) Energy (∆∆∆∆∆H°)

106.0 98.1

104.2 136

83.1 103

38.4 88

33.2 71

36.6 50.3

58.0 129.3

46.1 69.7

36.1 108

42.2 84

51.3 70

50.9 56

76.0 81

102.4 81

104 81

98

9478

85

90 70

85.7 85

69.1

69.3

Except for the values of H—O, D—D and Na—Cl, other values are taken from S W Benson, J. Chem. Educ., 42, 502

(1965). The values of H—O, D—D and Na—Cl are from K S Pitzer, J. Am. Chem. Soc., 70 2140 (1948).

D — D

H — H

C — C

N — N

O — O

F — F

Cl — Cl

Br — Br

I — I

Si — Si

P — P

S — S

Si — H

H — O

CH3 — H

CH3CH2 — H

(CH3)2C — H|H

(CH3)3C — H

Si — C

Si — Br

C — Si

Na — Cl

H — F

H — Cl

H — Br

H — I

I — Cl

Si — F

C — N

C — F

C — Cl

C — Br

C — I

CH3 — Cl

CH3CH2 — Cl

(CH3)2C — Cl|H

(CH3)3C — Cl

C6H5CH2 — H

C6H5CH2CH2 — H

CH2=CHCH2 — H

BASIC CONCEPTS 3

It is noticed from the above Table that C — H bond dissociation energies for ethane(98 kcal/mol) and for toluene (85 kcal/mol) are considerably different and smaller for toluene.The obvious reason for this difference is that the benzyl radical produced by bond dissociationin toluene is substantially stabilized by resonance.

1.2 ELECTRONEGATIVITY

A fundamental property of a chemical bond is polarity. The polarity of a bond depends on thenature of the atoms linking that bond. The preferred tendency shown by an atom of a covalentbond to attract the shared pair of electrons towards itself is termed as a measure ofelectronegativity of that atom. This unequal attraction of electrons causes a charge separationbetween the atoms and produces a dipole, indicated by a charge notation (δ+

or δ– ) and shown

for a bond between two atoms, i.e., HCl molecule:

Hδ+— Clδ–

The partial negative charge on the chlorine atom indicates that the chlorine atom has agreater tendency to attract the shared electron pair towards it. As a result the electron densityof the bonding electrons lies more towards Cl atom. Chlorine is more electronegative thanhydrogen. A complete transfer of electrons takes place in the case of ionic bonds. We often useelectronegativities as a guide to predict whether a given bond will be polar or not and also thedirection of its dipole moment.

Electronegativity values have been measured from a consideration of bond distances,bond energies, dipole moments, and the data for some elements are listed in Table 1.2.

TABLE 1.2: Electronegativity Values of Different Elements

H

2.1

Li Be B C N O F

1.0 1.5 2.0 2.5 3.0 3.5 4.0

Na Mg Al Si P S Cl

0.9 1.2 1.5 1.8 2.1 2.5 3.0

K Ca Ge As Se Br

0.8 1.0 1.7 2.0 2.4 2.8

Cs Ba

0.7 0.9

Source: L Pauling, The Nature of the Chemical Bond, 3rd ed., Cornell University Press, Ithaca, (1960).

From these values the following conclusions may be drawn:1. On going from top to bottom in a group, as the atomic size of the element increases

the electronegativity decreases.2. In a given period, electronegativity increases with increasing atomic number.3. The more non-metallic an element, the greater is the electronegativity. The value of

electronegativity for fluorine is the largest.The electronegativity of carbon is almost similar to that of hydrogen, as a result C—H

bond is usually considered as non-polar. As we will study in Chapter 4, nucleophilicity decreaseswith increasing electronegativity of the attacking atom. The electronegativity order for hybridizedcarbon orbitals is sp > sp2 > sp3.


1.3 DIPOLE MOMENT

In the case of non-polar bonds, the electric charge is equally distributed between similar atomsforming the bond. On the other hand, in case of polar bonds there is a charge displacementtowards the more electronegative atom (causing a charge separation and making the bondpolar). When the centers of positive and negative charges do not coincide, the bond maypossess a dipole moment designated by the symbol µ. The dipole moment is equal to theproduct of the charge (e) and the distance (d) of separation between the atoms as shown inequation (1.1). A molecule is considered to possess dipole moment when the two bonded atoms

µµµµµ = e × d …(1.1)

differ in electronegativity. Since the electronic charge is 10–10 e.s.u. and the distance betweenatoms is in angstrom units (10–8 cm), the dipole moment is obtained in Debye units,D = 4.8 × 10–18 e.s.u. It is not possible to measure the dipole moment of an individual bondwithin the molecule. We can measure only the total dipole moment of the molecule. Dipolemoment values for some compounds are listed in Table 1.3.

TABLE 1.3: Dipole Moments of Some Common Organic and Inorganic Compounds(Values Determined in the Vapor State Unless Otherwise Mentioned)

Compound µµµµµ Compound µµµµµ(in D) (in D)

C6H5CH3 0.37 CH3I 1.64

C6H5Si(CH3)3 0.44* CH3OH 1.69

BrCl 0.57 p-ClC6H4Si(CH3)3 1.70

HCl 0.65 C2H5OH 1.74$

HBr 0.79 CH3Br 1.79

H2S 0.93 CH3F 1.81

CH3COOH 0.94†, † CH3Cl 1.83

HCOOH 1.06†, † H2O 1.84

HCl 1.08 p-CH3C6H4Cl 1.95†

C6H5COOH 1.13†, ‡ p-CH3C6H4Br 1.96†

CH3Cl 1.22$ ClCH2COOH 1.97†,$

C2H5OC2H5 1.22$ Oy

C6H5OH 1.40 HCH 2.23Oy

NH3 1.47 CH3CCH3 2.76$

C6H5NH2 1.48 HCN 2.93

C6H5Br 1.55$ CH3CN 3.94

C6H5Cl 1.58$ C6H5NO2 4.01$

SO2 1.61 p-CH3C6H4NO2 4.42†

Source: J W Smith, Electric Dipole Moments, Butterworths Scientific Publications, London,84–308, (1955). *J D Roberts et al., J. Am. Chem. Soc., 71, 2923 (1949); †In benzene solution; ‡Values for the dimeric

molecule; $In solution.

BASIC CONCEPTS 5

The dipole moment of a molecule will vary depending on the other atoms present in thevicinity. It is a vector quantity. In polyatomic molecules the dipole moment of the molecule asa whole is equal to the vector sum of the individual bond moments. It is thus expected thata molecule can be non-polar but still contain polarized bonds. This is possible if the bonds areso directed that the dipole moments of the individual bonds cancel out, i.e., their vector sumis zero. In carbon tetrachloride, for instance, the C — Cl bond possesses bond moment of value2.3 D. Since the bonds are symmetrical the dipole moment is cancelled out. Carbon tetrachloridetherefore, has a net zero dipole moment and is non-polar. Chloroform, on the other hand, hasa value of 1.22 D and is polar. Unshared pair of electrons make large contributions to the dipolemoment of water and ammonia. This is explained by the polarity of the lone-pair of electronspresent on O and N. A lone pair contributes to a large moment directed away from the centralatom.

O N N

H H FH H F

µ = 1.84 D µ = 1.5 D µ = 0.2 DH F

Water molecule has a dipole moment of 1.84 D. This proves that it is not a linear molecule.Contribution to the net dipole moment arises from individual bond moments and the lone-pairas shown above. Ammonia has a higher dipole moment than NF

3. This is due to the fact that

in NH3 the N—H bonds and the lone-pair moments reinforce each other while in NF

3 they

oppose each other.

Molecules containing dative bonds are expected to possess large dipole moments since theformation of such bonds requires a marked separation of charges. Because of small differencesbetween the electronegativity of C and H, alkanes have very small dipole moments.

A knowledge of dipole moment is of considerable importance in the study of moleculargeometry of molecules. For example, geometrical isomers can be identified unambiguously,from dipole moment data.

1.4 HYDROGEN BOND

Hydrogen bond is a weak bond formed by the attraction of a hydrogen on one atom withanother atom bearing an unshared electron pair. In other words, the hydrogen atom forms abridge between two electronegative atoms. Thus hydrogen can participate in hydrogen bondingwith O, N or F. (organic molecules do not contain H—F bonds). The hydorgen bond is denoted

O — H ····· O O — H ····· N

by a dotted line. This bond is much weaker than the covalent bond and energy of such a bondin a molecule ranges from 2–10 kcal/mol. Compounds like alcohols, acids, water, ammonia, etc.display hydrogen bonding.

The existence of this bond has been established by various techniques and molecularproperties. In n.m.r., for instance, a hydrogen-bonded hydroxyl group shows a downfield shiftof its proton. Similarly, i.r. studies reveal that the stretching vibration of the O—H bond


decreases in the presence of a hydrogen bond. Varied properties like mp, bp, heat of dilutionand solubility point etc. are due to the existence of hydrogen bond.

Two types of hydrogen bonds have been recognized, viz., intramolecular (within the samemolecule) and intermolecular (between two or more molecules). Both types markedly influencethe properties of molecules.2, 3 Hydrogen bonding plays an important role in many chemicaland biochemical processes as well.

Intramolecular hydrogen bonds are considered to influence the acidity of certain acids.Thus, salicylic acid is about eighteen times more stronger than benzoic acid whereasp-hydroxybenzoic acid is half as acidic as benzoic acid. The large acidity of salicylic acid is

COOH

Ka 6.25 × 10–5

2.9 × 10–5

COOH

OH

105 × 10–5

COOH

OH

6 × 10–2

COOH

OHHO

attributed to the intramolecular hydrogen bonding which is capable of stabilizing the salicylateion.

OH

C O

OH

+ H O2

O OH

C O

–

+ H O3+

The hydrogen atom of the hydroxyl group is sufficiently close to the oxygen atom of thesalicylate ion for it to interact electrostatically with the formation of a six-membered ring. Thisis also sometimes referred to as chellation. p-Hydroxybenzoic acid, on the other hand, is unableto enter into such a ring formation because of the remoteness of the p-OH group. The introductionof a second OH group in the ortho position results in even greater increase in the value of thedissociation constant of the acid.3 Because of H-bonding, 2,6-dihydroxybenzoic acid is strongerthan even phosphoric acid.

In many 2-substituted heterocyclic acids the hydrogen bond has the opposite effect, i.e.,

it reduces the strength of the acid. The dissociation constant, for instance, of picolinic acid (1)is decreased because of the H-bonding taking place between the hydrogen atom of the carboxylgroup and the ring nitrogen atom. The hydrogen bonding stabilizes the acid because the– COOH group acts as the proton donor and decreases its acidity.4

N

H

C

O

O

(1)

BASIC CONCEPTS 7

Between keto and enol forms of a molecule, the former is considered more stable. But theenol form of β-diketones and β-keto esters are stabilized by intramolecular H-bonding and arestable.

O O O O

H H

H

Enol/Keto > 50

O O O O

OCH3 OCH3

H

Enol/Keto > 29

Intermolecular hydrogen bonding also affects several additional properties such as boilingpoint, solubility, vapor pressure of solvents and dielectric constant. As an example considerethyl alcohol (bp 78°C) and dimethyl ether (bp –24°C) which have the same molecular formulaC

2H

6O but different structures. The former, however, has a higher boiling point than the latter.

This difference is explained by the presence of intermolecular H-bond formation, also calledassociation, of ethyl alcohol in the following manner (2):

O O OH H H

C H2 5

H

HO

H

C H2 5

OH C5 2 H

OHH

(2)

Hydrogen bonding is thus expected to cause an elevation of boiling point. Hydrogenbonding exists in H

2O but is absent in H

2S, H

2Se and H

2Te. Nitrobenzene is less soluble in

water than phenol or aniline though nitrobenzene has a large dipole moment, µ= 4.01 D.Presumably nitrobenzene does not form hydrogen bonds with water as do the other two.

It has been reported that hydrogen bonds exist in carbanions,5 between OH and π electrons,as in allylic phenols. Hydrogen bonds occupy a significant place in biologically importantmolecules such as nucleic acids and proteins.6 A large quantity of water is retained by livingcells in plants and animals and most of it is attached to proteins by hydrogen-bonds.


1.5 HYPERCONJUGATION (Interaction Between σσσσσ- and πππππ-Systems)

Baker and Nathan, during the reaction of organic halides with pyridine, observed the followingrelative rates:

R CH Br 2 + R CH2 N+

+ Br–

R H –CH3 –CH CH2 3 –CH(CH )3 2 –C(CH )3 3

Relative rates 1 1.66 1.48 1.34 1.35

N

The rate was largest when R was a methyl group. Similar results were obtained by otherworkers for the solvolysis of p-alkyl substituted benzyl halides, the order of reactivity being:CH

3 > CH

2CH

3 > CH(CH

3)2 > C(CH

3)3. This is the reverse of what would be expected if only

the inductive effects were involved.These workers explained the results by suggesting that the methyl group can donate

electrons through a no-bond resonance or hyperconjugation. In resonance the overlap of π andlone-pair electrons is involved. In contrast, in hyperconjugation structures are written involvingoverlap of σ electrons with the adjacent π-bond as shown for toluene. The methyl groupcomprises three σ bonds and can be conjugated with the ring in the following manner:

CH H

H

C–

C C CH H H HH+

H+

H+

H+

H H H H

–

–

–

Hyperconjugative structures of toluene

A methyl group adjacent to a carbon—carbon double bond acts as an electron donor.A methyl group is a better electron donor than any of the other alkyl groups by virtue of

the fact that it has more hyperconjugative structures and can thus stabilize the transition state.Similar structures for the t-butyl group would involve carbon—carbon bond electrons whichare less important. Hyperconjugation thus seems quite valid in explaining the above differencesin reaction rates. A critical analysis of hyperconjugation has been provided by Dewar.7

Like resonance, the concept of hyperconjugation affords adequate physical interpretationsfor the behavior of molecules in chemical reactions. As we have seen above, it plays animportant role in reactions involving positive transition states. Hyperconjugative effects alsoaccount for the higher resonance energies of compounds, for instance toluene has a resonanceenergy (37.5 kcal/mol) which is 1.5 kcal/mol higher than that of benzene (36 kcal/mol). Thedifference is ascribed to hyperconjugation in toluene which makes it more stable. Dipole momentdata support the same general view. Benzene has zero dipole moment whereas toluene possessesan appreciable value (0.37D). The basicity of di-and tri-methyl substituted benzenes is greaterthan toluene for the same reason.

The conjugative ability of a three-membered ring is less than that of a multiple bond. Thisis explained on the basis that the dipole moment values of cyclopropyl ketones are betweenthose of alkyl and vinyl ketones, i.e., a cyclopropyl ketone is less polarized than a vinyl ketone.

BASIC CONCEPTS 9

(CH ) CH3 2 CH

CH3 CH3CH3

C CCO OO

µ = 2.7 D µ = 2.98 Dµ = 2.84 D

CH2

Further support for hyperconjugation comes from the carbon— carbon bond distance data.The carbon— carbon bond distance in propene is 1.488 Å compared to 1.543 Å in ethane andthe C=C bond distance is 1.353 Å compared to 1.334 Å in ethylene. It is reasoned that theC

2— C

3 bond acquires some double bond character and the C

1— C

2 bond a single bond character

because of hyperconjugation. This shows an electron density transfer from C—H to theπ−system with a net strengthening of the bond between C2— C3 but a weakening of theC

1— C

2 π bond.

Hyperconjugative structures of propene

H HC C

H+

H+

H H

3 3CH CH

1.448 1.353

2 2CH2 CH2

–1 1–

There is little doubt about the general usefulness of the concept of hyperconjugation toexplain some characteristics of certain molecules. But it is still a controversial concept and isused selectively because certain reactions such as aromatic electrophilic substitutions areinconsistent with this phenomenon.

1.6 INDUCTIVE AND FIELD EFFECTS

A methyl group, when attached to a π-electron system acts as electron donor. The electronsbeing transmitted in this case by conjugation. There is, however, another way of transmissionof electrons to the reaction site from a distant substituent. It is a polar effect and takes placewhen the substituent attached to the carbon chain is electron-donating or electron-withdrawingin nature. In the classical sense this takes place via the inductive effect in which the electricaleffects of the group are propagated in a chain by the successive polarization of carbon— carbonσ bonds. This occurs by virtue of the difference in the electronegativity of the bonded atomsin the chain.

δδδ+ δδ+ δ+ δ–

C C C X

The argument for this operation is that the atom X of the C — X bond, which is polar(X may be F, Cl or any other electronegative group), acquires a formal negative charge becausethe electron pair is shared more by the electronegative atom. This induces polarization amongthe other bonds, and makes the carbon chain to attain a positive charge. This effect diminishesas the length of the carbon chain increases. The resonance and inductive effects may notnecessarily operate in the same direction. In the ground state these effects are permanent andare manifested in many properties of the molecules. One of the most ideal cases for correlationof inductive effects is the solvolysis of 4-substituted bicyclo[2.2.2]octa-1-yl-brosylate (3) inacetic acid at 75°C. The relative rates are:


R

OBS(3)

where

O–

OBS Br S

O

O

R – CH3

– CH2CH

3– CH(CH

3)2

– C(CH3)3

10–5 k sec–1 3.38 4.04 4.75 6.24

The reaction rates increase progressively from methyl to t-butyl group. This increase wasattributed8 to the inductive electron release by the alkyl groups. The t-butyl group is the mosteffective because it releases electrons more efficiently than any other group in this series. Itmust be noticed that inductive effect of the alkyl groups is opposite to that of the hyperconjugationeffect where the methyl group is more electron-releasing. The acidic and basic properties ofreagents have been attributed to the electron-releasing or withdrawing inductive effects of thesubstituents.

The inductive effect of substituents has a significant effect on the p.m.r. of molecules.Halogens, for instance, are electron-withdrawing, therefore neighboring hydrogen atoms suffera decrease in electron density and as a result their nuclei are deshielded. The order of thiseffect parallels the change of electronegativity F > Cl > Br > I of the halogens. The proton thusresonates at a lower field in the p.m.r.

Recently modified models have been proposed for the inductive effects. From a molecularorbital approach, Pople and Gordon9 have suggested that the inductive effect of the electronegativegroups results in charge alteration rather than steady charge decay. According to this, the βcarbon atom attains a negative charge, the γ carbon atom attains a positive charge, and so on.Dewar and Marchand,10 on the other hand, have concluded that the inductive effect is, in fact,a direct field effect acting across the space by successive polarization of the intermediate bonds.

δδ+ δδ– δ+ δ–

C C C Fγ β α

Polar effects may be directly transmitted through space or through charge-dipole or dipole-dipole interactions rather than along a carbon chain. Such effects are often called the fieldeffects. These are long-range polar interactions, i.e., those at more than two carbon— carbonbond lengths away. Both the inductive and field effects operate in the same direction. It is,therefore, difficult to separate these two effects. However, an attempt has been made in thisrespect. The ionization constants for instance, of polycyclic acids of type (4) were measured11

and the values of pKCl

and pKCOOCH3

were found to be higher than the value of pKH. Electron-

withdrawing groups are found to decrease the acidity. On the inductive model an increase in

acid strength should have been expected. The higher the pKa, weaker the acid. The conjugative

and hydrogen-bonding effects can be precluded in this system and the inductive electron-

BASIC CONCEPTS 11

X COOH

(4)

X = H, pKH

= 6.04X = Cl, pK

Cl= 6.25

X = COOCH3, pK

COOCH3= 6.20

withdrawing effects of these two groups are also minimal. The difference in the pK values isexplained only by invoking the field effects, i.e., the polar Cl or COOCH

3 groups destabilize the

carboxylate ion through space. The Cl group has a better destabilizing effect. Nitration oftriptycene has been explained in terms of the inductive effect.12 The deuterium exchange rateof fluorobenzenes has been interpreted by a simultaneous operation of both the field and theinductive effects. It thus follows that a clear-cut separation of the inductive and field effectshas to await further work.

1.7 STERIC EFFECTS

The steric effects arise from the interactions between atoms present in a molecule that are notbonded to each other. They are also called non-bonded interactions. This happens when bulkygroups are located in the vicinity of the reaction site and the reaction is hindered resulting inthe diminished rate of a reaction. Such a hindrance to reaction was first noted by Kehrman andthis phenomenon was named steric hindrance by Meyer. A vivid illustration of steric blockingof a reaction with the resultant decrease in reactivity, is provided by the nitration(HNO

3 + H

2SO

4) of alkyl substituted benzenes. Nitration, which presumably involves the NO

2+

ion, yields the ortho- and para-nitro derivatives. The o/p ratio for the following hydrocarbonsin nitration was found to be:

CH H

H

1.57o p/ ratio

H H

CH3

C

0.93 0.48

CH CH3

CH3

0.21

CCH3 CH3

CH3

A drop in the o/p ratio indicates a marked decrease in the formation of the ortho isomeras one approaches the t-butyl group. This decrease is interpreted in terms of the steadilyincreasing steric requirements of the alkyl groups. The size of the substituent group hinders theapproach of the NO+

2 ion to the ortho position.

Another example that illustrates a decreased reactivity of a molecule involves thedissociation of carboxylic acids. For example, the dissociation constant value in 50%H

2O — 50% CH

3OH of acetic acid is almost 25 times higher than that of methyl iso-butyl

neopentylacetic acid (5).

CH3 (CH ) CCH3 3 2C COH C

O C(CH )3 3

CH3

O

OH

(5)

1.1 × 10–7

2.7 × 10–6

Ka


*H K Hall, Jr. and A E Shekeil, J. Org. Chem., 45, 5325 (1980).

It is apparent that in the latter, large alkyl groups hinder the release of a proton. The roleof steric effects in organic chemistry has been reviewed by Newman.13

The steric effects do not always decrease the rate of a reaction but can often accelerateit. This happens when the reactants in the ground state are compressed and strain is relievedin attaining the transition state. This phenomenon is termed steric acceleration and has beenobserved in the reactions of flexible aliphatic and alicyclic systems.14 The formation of isobutylenein the solvolysis of t-butyl chloride has been argued on the basis of relief of strain from theintermediate t-butylcarbocation. In the rate determining step of the reaction a rehybridizationat the reaction center from sp3 to sp2 occurs. By the same token three and four-membered ringcompounds form open-chain products in deamination and solvolysis reactions requiring lowenergies of activation. Opening of a small ring involves the release of ring-strain and an enhancedrate. Brown and Hammer,15 however, maintain that such a relief to strain may not be possiblein rigid ring systems. In addition to the effect on reaction rates, the steric factors have implicationsin organic synthesis as well. They reduce resonance in molecules (steric inhibition of resonance),play a dominant role in preventing the dimerization of aroxyl radicals and explain the opticalisomerism of diphenyl derivatives.

1.8 BREDT’S RULE

This is an empirical rule deduced from experimental observations by Bredt. According to thisrule, in small bicyclic compounds a C=C bond cannot exist to the bridgehead atoms. Thuselimination of bromine from bicyclo[2.2.1]hexane (6) gives no bicyclo[2.2.1]hex-1-ene (7).

Br

t-BuO k– +

t-BuOH

(6) (7)

Compound (7) is highly strained molecule and is incapable of existence* (note the doublebond is cis in the six-membered ring).

Introduction of a carbon—carbon double bond to the bridgehead compresses the moleculeinto a planar conformation because of the planar geometry of the sp2 carbons being formed. Asa result the transition state energy for elimination required is considerably high and a smallbicyclic system cannot accommodate the resultant strain. Therefore, the alkene (7) cannotaccommodate a C=C bond and thus very unstable. However, bridgehead alkenes should bestable when the rings are large enough to accommodate the strain. How large must the ringsystems be before they defy the Bredt’s rule. Fawcett16 defined the limits of the rule andsuggested that the sum (S) of the numbers of atoms S in the bridge of a bicyclic compounddetermines whether the ring system is stable or not. He tentatively gave a numerical value of9 or more to S, i.e., S ≥ 9.

The synthesis of molecules which violates Bredt’s rule has been a challenge to chemists.Eventually in 1967 the synthesis of bicyclo[3.3.1]non-1-ene (9), from a β-lactone (8) which hasa value of S = 7, was reported simultaneously by two groups of workers.17, 18 Since thiscompound is stable, the value of S has to be lowered for defining stability, i.e., S ≥ 7 for stablering systems.

BASIC CONCEPTS 13

OO

a b c

(8) (9)

c

ba

Wiseman and Pletcher17 have suggested that the strain of a bridgehead is closely relatedto the strain in trans-cycloalkenes. Therefore, a bicyclic system with a bridgehead double bondshould be isolated, provided the double bond is trans in the larger of the two rings in whichit is endocyclic and has at least eight carbon atoms. In cycloalkene (9), the three rings systemsab, bc and ac are present. The double bond is trans and endocyclic in the ring ac which alsocontains eight atoms and this ring system is thus possible.

Several other bicyclic bridgehead alkenes (10 to 12) with S = 7 have subsequently beenprepared.

(10) (11) (12)

The Bredt’s rule has proved to be of immense value in mechanistic interpretations and inthe assignment of structures. The rule has been successfully employed to account for thefailure of β-keto acids containing the carboxyl group at the bridgehead to decarboxylate; foralkyl halides which resist elimination; alcohols which do not dehydrate; and for certain ketoneswhich do not exchange their α-hydrogen for deuterium. Thus bicyclo[2.2.1]heptan-7-one-1-carboxylic acid (13),19 (with S value of 5) resists decarboxylation even on heating to 500°C. Thereason being that a β-keto acid is believed to decarboxylate through an enol and in the case

COOH

500°CNo reaction

(13)

O

of (13) this involves the formation of a bridgehead enol which is impossible. The bicyclo[5.3.1]undecan-11-one-1-carboxylic acid (14) on the other hand, which has a value of S = 9,decarboxylates rather easily even on heating only at 132°C because the ring system can supporta double bond at the bridgehead.

O

COOH

(14)

The rule applies to heterocyclic systems as well.


The cyclic amide (15) shows basic properties (pKa = 5.3) unlike the open-chain

amide.20 This is attributed to the inhibition of a dipolar structure which would unfavorably puta carbon—nitrogen double bond at the bridgehead, and it will violate Bredt’s rule.

N N

CH3 CH3

CH3 CH3O O–

+(15)

For the same reason the following quinclidinone (16) will form an oxime.

N

O

(16)

The question arises whether or not the rule applies to systems containing atoms otherthan carbon and nitrogen, such as sulfur or silicon. Compounds containing such atoms seemto defy this rule as in the case of bicyclotrisulfone (17). Doering and Levy prepared (17) whichdissolves in aqueous sodium bicarbonate solution indicating that it is acidic in nature. The

O2S SO2SSO2

SO2 SO2SO2

H

– –O

O

(17) (17a) (17b)

SO2 SO2

–H+

etc.

anion formed can be stabilized by resonance as shown in structures (17a) and (17b). Thestructurally analogous compound, bicyclo[2.2.2]octan-2,6,7-trione (18), on the other hand, showsno acidity21 because the corresponding anion cannot be similarly stabilized by resonance.

CO CC

H

(18)

O O

Thus a carbon-sulfur double bond at a bridgehead is possible. Sulfur belongs to the VIBgroup of the periodic table and has the following electronic configuration:

1s 2s 2p 3s 3p

In the divalent state, the 3py and 3p

z orbitals are filled by electron sharing but there are

five unoccupied d orbitals. It has been argued22 that a sulfur atom utilizes the d orbitals to forma bond, described as the pπ–dπ bond.

BASIC CONCEPTS 15

1.9 AROMATICITY

We often say that certain organic molecules are aromatic in character whereas others are not.The simple substance benzene, for instance, is a classical example of the former class ofcompounds. Benzene is planar and the π-electrons are delocalized and it is stable. Benzene thushas aromatic character also known as aromaticity.

Cyclooctatetraene, on the other hand, is non-aromatic because it lacks at least these twofundamental properties even though it contains alternating single and double bonds. It is foundimpossible to construct an unstrained planar model of cyclooctatetraene. Actuallycyclooctatetraene behaves like an unsaturated compound and possesses a non-planar “tub-like”configuration.

Planar Tub shaped

Cyclooctatetraene

Cyclobutadiene (19) might be expected to be aromatic because like benzene, it also hasresonance structures. However, there is no effective overlap of p orbitals required for aromaticity.Cyclobutadiene is thus not aromatic rather it is antiaromatic. Apparently there must be someother way to predict the aromaticity23 of compounds.

(19) (20)

A method to define aromaticity was developed by Hückel24 using the molecular orbitaltheory. He found that the relative stability of conjugated monocyclic compounds depended onthe number of electrons in the system. His calculations have been translated into the form ofan expression called the Hückel’s rule. This rule states that monocyclic, conjugated and coplanar

systems that possess (4n + 2)π electrons (where n is an integer) are aromatic. The benzenoidhydrocarbons fit all criteria of aromaticity, i.e., they have high resonance energies. Applicationsof the PMO method also leads to the prediction that the benzenoid hydrocarbons should beespecially stable. Since the inception of this rule the aromaticity of a large number of compoundshas been predicted.25 We shall discuss several examples in the following pages.

n = 0: The simple case of 2π electron system where n = 0 is ethylene but this moleculeis non-cyclic and also non-aromatic. Hückel’s rule predicts that monocyclic ions with (4n + 2) πelectrons will be aromatic, but those with 4nπ electrons will not. The cyclopropenium cation

(21) is cyclic and complies with the Hückel’s rule. The overlap of the three equivalent p-orbitals(21a) delocalizes the π electrons over the three carbon atoms (21b) to give a stabilized cation.Thus aromaticity is not restricted to neutral molecules but a number of ions are aromatic aswell. In contrast to (21) the cyclopropenyl anion is non-stabilized by resonance and thus non-aromatic.


H H HH H H

HH H

Cl

AgBF4

BF4–

H

H

H

+

+

(21) (21a) (21b)

+

+ +

+– –

–

n = 1: When n = 1, the cyclic system should have 6π electrons. This is true in the caseof a large number of compounds like benzene, phenol, pyridine, cyclopentadiene anion, etc.These are aromatic and possess high energies of resonance.

Cyclopentadienyl anion has six π electrons, a negative charge is considered equal to twoelectrons and is aromatic. It is stable like other carbanions. It can be prepared by the abstractionof an acidic proton from cyclopentadiene using a strong base.

HH

H

(CH ) O3 3–

k+

–

+ (CH ) COH3 3

In cyclopentadiene there is no continuous ring of p-orbitals as one – CH2 group is sp3

hybridized. Deprotonation of – CH2 group leaves an orbital occupied by a pair of electrons. This

orbital can rehybridize to a p-orbital and a ring containing 6π electrons results. The anion nowobeys the Hückel’s rule and it is aromatic.

H

H

(CH ) CO3 3–

H + (CH ) COH3 3

–

Benzene is a cyclic ring with a continuous ring of overlapping p-orbitals. There are 6πelectrons in benzene. It is a (4n + 2) π system and complies with the Hückel’s rule. It is anunusually a stable compound.

The Hückel’s rule predicts that the cycloheptatrienylium ion (22) would be stable andaromatic. This ion, also called the tropylium cation, was prepared by Doering and Knox by thethermal dissociation of the corresponding bromide and, as predicted, is stabilized by resonance.

+H

(22) (22a)

+

Ions with 10π electrons have also been prepared and studied. Dianions of hydrocarbonsare difficult to make but some of them can be easily prepared because they are stable andaromatic. Katz and Coworkers26 prepared the cyclooctatetraene dianion (23) as a solid by

BASIC CONCEPTS 17

treating cyclooctatetraene with potassium metal in tetrahydrofuran. This provides anotherexcellent example for the verification of the Hückel’s rule. This also supports the view thataromatic stabilization leads to usually stable hydrocarbon anions.

+ 2KTHF

Reflux

––2

(23a)

–+ 2K

+

(23)

n = 2: For n = 2, the system would have 10π electrons and expected to be aromatic.Cyclodecapentaene (24) itself could not be prepared because of the large steric interaction

12

3

4

56

HH

(24) (25)

between the internal hydrogen atoms at C1 and C

6. With the prediction that an addition

methylene group at 1,6-position may make the molecule planar and strain free, Vogel and

Coworkers27 synthesized 1,6-methanocyclodecapentaene (25) which is a conjugated 10π electronsystem and it behaves like aromatic compounds. In this molecule, the steric problem has beenavoided with slightly less planarity in the π-system.

n = 3: It is interesting to note that Hückel’s rule applies equally well to certain polycycliccompounds such as pyrocyclene (26), naphthacene (27) and coronene (28).

(26) (27) (28)

Sondheimer28 has synthesized a new series of conjugated cyclic hydrocarbons which arecalled annulenes. The name annulene was also suggested as a general name for monocycliccompounds containing alternating system of single and double bonds. The ring size of anannulene is indicated by a number in brackets. According to this nomenclature, benzenebecomes [6] annulene and cyclooctatetraene as [8] annulene. In these molecules the numberof π electrons is equal to the number of carbon atoms in the ring. Large annulenes can attainplanar configurations and are aromatic. Thus cyclooctadecanonane or [18] annulene [29] contains18 π electrons and obeys Hückel’s rule. It is a planar molecule and the n.m.r. spectrum indicatesthat it has a ring current characteristic of aromatic compounds. The resonance energy of (29)from combustion data has been measured to be 100 kcal/mol.


HH

H H

H

H

H H

H

H

H

H

HH

H

H

H

H

(29)

Cyclotetradecaheptaene or [14] annulene (30) contains 14π electrons and is aromatic.29

However, it has been argued that this hydrocarbon must have low resonance energy becausethe ring is a bit distorted from planarity. However, the dehydro compound [14] annulene (31)also containing 14π electrons is planar and aromatic.30

H H

H

H

HH

H

H

H

(30)

H

HHH

H H

H

H

H

H

HHH H

HH

(31)

H H

The term dehydro implies compounds with greater degree of unsaturation. The triplebond in these compounds (31) is considered to constitute two π electrons for the Hückel’s rule.

There are, however, compounds which do not contain the benzene ring but are stillaromatic. Such compounds are known as non-benzenoid aromatics.31 Examples of such systemsare azulene (32), N-phenylsydnone (33), pyrrole (34), ferrocene (35), tropolone (36) etc. Thesecompounds behave like normal aromatic systems.

–O

O

N

C H6 5

N+N

H

–

Fe2+

OH

O

(32) (33) (34) (35) (36)

–

Despite the wide utility of the Hückel’s rule a number of systems are known which donot contain (4n + 2) π electrons but are still aromatic. Well-known examples of these anomalouscases are heptalene (37), pyrene (38), acenaphthene (39), etc., and they possess considerabledelocalization energies. A set of rules has been proposed to distinguish between the aromaticand non-aromatic natures of such polycyclic systems.

BASIC CONCEPTS 19

(37) (38) (39)

Cyclic conjugated systems in which electrons are delocalized are aromatic. Breslow32 hasdefined a new term called anti-aromaticity to describe conjugated, cyclic and planar systemsin which overlap of p-orbitals is destabilizing. They are thermodynamically less stable thantheir corresponding acyclic counterparts. This can be best illustrated by the ions (40) to (42)all of which contain 4nπ electrons and oxirene (43) and cyclobutadiene. All these species aredestabilized by resonance.

C H6 5

–

C H6 5

C H6 5

(40)

+

(41)

–

(42)

O

(43)

Cyclooctatetraene is not antiaromatic, though it has a continuous cycle of 4nπ electrons,it is not planar. Therefore, it is nonaromatic.

Recently33 a conformational criterion for aromaticity and anti-aromaticity has also beensuggested.

1.10 THE HAMMETT EQUATION

Groups located in a benzene ring or other π-systems have considerable influence on the propertiesof the molecule. A substituent can function as a neighboring group, an electron-donor oracceptor by resonance and inductive effect or alternatively offer steric effect. These effects cancause changes in the reactivity of molecules. For instance, nitration of phenol in dil nitric acidyields a mixture of isomeric o- and p-nitrophenols. This is a result of the increase in electrondensity on the phenol ring due to conjugative effects. We would like to know as to what partof the influence is due to each of the effects mentioned above. This was first accomplished byHammett.34 He used the ionization of benzoic acid in aqueous solution at 25°C as the referencereaction and then measured the effect of substituents placed at the m- or p-position on theobserved ionization constant of these substituted benzoic acids.

+ H O2

+ H O2

+ H O3+

+ H O3+

KO

KR R

COOH COO–

COO–

COOH


The basic assumptions were that the steric effect of the m- or p-substituents was absentand the influence on the ionization was exclusively due to the electronic nature of the groupsR. Hammett thus obtained a quantitative substituent effect denoted by σ

R (equation 1.2) for the

effect of a given substituent (R) on the ionization of benzoic acids.

log K

K0

F

HGI

KJ = σR

...(1.2)

where, K is the ionization constant of substituted benzoic acidK

0 is the ionization constant of benzoic acid

σ is a constant characteristic of the group RIn order to extend this treatment to other reactions of benzene derivatives such as ester

hydrolysis and hundreds of other related reactions an important assumption was further madethat the effect of a substituent on the reaction rate of any other reaction is proportional to itseffect on benzoic acid ionization, i.e.,

logk

k0

F

HGI

KJAny reaction

= ρ logK

K 0

F

HGI

KJBenzoic acid

...(1.3)

Equation (1.3) can be written as:

logk

k0

F

HGI

KJ= ρσ ...(1.4)

Equation (1.4) holds good for other reaction series as well.where σ is the substituent constant and

ρ is the reaction constantand k is the rate for any substituted benzene derivative undergoing the reaction

k0 is the rate of reaction for the unsubstituted compound undergoing the same reaction,such as alkaline hydrolysis of m- and p-substituted methyl benzoates.

In this form equation (1.4) is called the Hammett equation.

For the Hammett equation to be correct, we must use the values of k and k0 under

identical experimental conditions (i.e., solvent, temperature, concentration, etc.). The Hammettequation is correct because of the presence of linear relationship between the free energies (ofreaction or activation) for different series of reactions.

The expression for the rate of the reactions has the form shown in equations (1.5) and(1.6) according to the theory of absolute reaction rates (see chapter 2), for the two reactions.

log k = log –.

KT

h

G

RT

FHG

IKJ

≠∆

2303…(1.5)

and log k0

= log –.

KT

h

G

RT

∆0

2303

≠

…(1.6)

or log k – log k0

= –. .

∆ ∆G

RT

G

RT

≠ ≠

+2303 2303

0

or logk

k0

= –. .

∆ ∆G

RT

G

RT

≠ ≠

+2303 2303

0 …(1.7)

BASIC CONCEPTS 21

Compare equations (1.4) and (1.7), we get:

–.

∆GRT

≠

2303= −

≠∆GRT

0

2303. + ρσ

or ∆G≠

= ∆G0

≠ – 2.303RT ρσ …(1.8)

∆G0 is the free energy of activation of the reaction of the unsubstituted compound. Similar

to equation (1.8) we can write an equation (1.9) for the free energy of another reaction of asubstituted compound characterized by a different value of the constant:

∆G≠′ = ∆G

0

≠ ′ – 2.303 RT ρ′σ′ …(1.9)

Equations (1.8) and (1.9) can be written, to give (1.10) and (1.11) respectively:

Divide (1.8) by ρ∆G≠

ρ=

∆G0

≠

ρ – 2.303 RT σ …(1.10)

Divide (1.9) by ρ′∆G≠′

′ρ=

∆G0

≠′

′ρ – 2.303 RT σ′ …(1.11)

Subtract these two equations, to obtain (1.12)

∆ ∆G G≠ ≠′

′ρ ρ– =

∆ ∆G G0 0

≠ ≠′

′ρ ρ– …(1.12)

Multiply (1.12) by ρ

∆G≠ –

ρρ′

∆G0

≠′ = ∆G0

≠ = constant …(1.13)

Equation (1.13) exhibits a linear relationship between the free energies of activation of thetwo reactions. In other words, this relationship implies that for any reaction series the valuesof the free energies of activation which correspond to a change in reacting molecules, arelinearly related to the values of the free energy of activation of any other reaction seriescharacterized by the variations in the same parameter.

The Hammett equation describes the influence of polar meta- or para-substituents on theside chain reactions of benzene derivatives. This equation, however, is not applicable to theinfluence of ortho substituents because they exert steric effects and to aliphatic derivativesbecause the twisting of the carbon chain may also act sterically. A plot of log k/k

0 against σ is

linear with a slope of ρ. The substituent constant σσσσσ is defined by equation (1.2)

σ = logk

k0

…(1.2)

where the k’s refer to the ionization constants of benzoic acids. Hammett equation is obeyedwith moderate precision for polar reactions. Equation (1.2) measures the polar effect of thesubstituent relative to hydrogen and is independent of the nature of the reaction. Both theinductive and mesomeric effects are contained in this equation. The values of the substituentconstants σ for selected groups are compiled in Table 1.4. These are based on the ionizationdata of benzoic acids.

σσσσσ constant

The constant σ is defined by equation (1.2). It is a measure of its capacity to perturb theenvironment electronically. According to the standard conditions σ = 0, for the unsubstituted


benzoic acid. Also when R = H, then ρ = 1 for the ionization of benzoic acid in water. Theelectron density at the reaction center either increases or decreases depending on the natureof the substituent. It is clear from Table 1.4 that σ values for some groups are negative whilefor others they are positive. A negative value for –NH

2, –OH, alkyl and others implies an

increase in the electron density at the reaction center. A positive value for – CN, –NO2, halo

groups indicates electron deficiency. These values thus can be used as a measure of the degreeof electron-release or electron-withdrawal by the groups attached to the benzene ring. In otherwords, they reflect the degree of change of electron density at the reaction center.

ρρρρρ constant

The slope of the Hammett plot gives ρ, the reaction constant. The reaction constant ρ measuresthe susceptibility of a reaction to the electronic effects and is reaction independent. Themagnitude of ρ gives a measure of the degree to which the reaction responds to substituents.From the magnitude and sign of ρ one can make a correct conclusion on the type of transformationof the initial substrate in the limiting stage of the process. This information offers no possibilityfor elucidating the detailed mechanism, however, it is a useful constant. A reaction involvinga positive charge in the transition state will be aided by electron-donating groups and the valueof ρ will be negative. This point may be illustrated by the following example:

acetoneRC

6H

4N(CH

3)2 + CH

3I RC

6H

4N

+(CH

3)3I–

35°C

The value of ρ for this reaction has been observed to be –2.56. A negative value of ρindicates that the reaction is aided by electron-donation and vice versa. The sign of ρ is inaccordance with an S

N2 displacement involving an activated complex of the following type.

The nitrogen atom must bear a substantial positive charge in the transition state.

R C H (CH ) N6 4 3 2 C

H H

I

H

d+ d–

A reaction, on the other hand, involving a decrease in the positive charge or increase inthe negative charge in the transition state will be facilitated by electron-withdrawing substituentsand the value of ρ is positive. The magnitude of the value of ρ indicates that the reaction issensitive to the polar effects of the substituents and also provides information about the natureof the transition state involved in the reaction. A very low value of ρ indicates that the reactionmay be insensitive to the effect of groups. Rates of many reactions have been correlated withthe Hammett equation and for many others can be predicted. It is difficult to predict ρ fromexperimental conditions since it depends on a number of factors such as the solvent, the natureof the leaving group, etc. Interposition of a methylene group between the reaction center andthe aromatic ring decreases the value of ρ because the polar effects now have to be propagatedthrough an additional bond.35

Certain extended equations have been proposed to the original Hammett equation. Jaffe,for instance, investigated the additive nature of more than one substituent placed on thearomatic ring. He found that the σ values for various groups can be added and the followingrelation, [equation (1.14)] holds good:

log k/k0

= ρ σ∑ …(1.14)

. .

BASIC CONCEPTS 23

where Σσ means the sum of the σ values of all the groups. For compounds containing more thanone benzene ring, equation (1.15) is employed to correlate the results.

log k/k0

= nρσ …(1.15)

TABLE 1.4: Substituent Constant Values for Selected Substituents

Group amσ a

pσ +bmσ

+bpσ c

lσ

– O–

– 0.708 – 0.519

– NMe2 – 0.211 – 0.600 0.10

– N(CH3)2 – 0.161 – 0.660 0.10

– Si(CH3)3 – 0.121 – 0.072 0.011 – 0.021

– C(CH3)3 – 0.120 – 0.197 – 0.059 – 0.256

– OCH3 – 0.115 – 0.268 0.047 – 0.778 0.25

– CH3 – 0.069 – 0.170 – 0.066 – 0.311

– C6H5 – 0.06 – 0.01 – 0.109 – 0.179 0.10

– C2H5 – 0.043 – 0.151 – 0.064 – 0.295

– OH – 0.002 – 0.357 0.25

– H 0 0 0 0 0

– N2+d

1.7 1.8

– C(CN)e3 1.00 0.01

– N(CH3)+3 0.904 0.859 0.359 0.408 0.92

– NO2 0.710 0.778 0.674 0.790 0.63

– CN 0.678 0.628 0.562 0.659 0.56

– SO2CH3 0.647 0.728

– SO2NH2 0.46 0.57

– CF3 0.42 0.54

– Br 0.39 0.232 0.405 0.150 0.45

– Cl 0.373 0.227 0.399 0.114 0.47

– COOH 0.355 0.265 0.421

– CHO 0.355 0.261

– I 0.352 0.276 0.359 0.138 0.38

– F 0.337 0.062 0.352 – 0.073 0.52

– COOCH3 0.315 0.489

– COCH3 0.306 0.516 0.28

– C ≡ CHf

0.205 0.233 0.334 0.179

– C ≡ C–C6H5g

0.140 0.165

– C(NO2)3h 0.82

a. H H Jaffe, Chem. Rev., 53, 191 (1953). b. Y Okamoto, T Inukai and H C Brown, J. Am. Chem. Soc., 80, 4979 (1958).c. C D Ritchie and W F Sager, Progress Phys. Org. Chem., 2, 323 (1964). d. E S Lewis and M D Johnson, J. Am. Chem.Soc., 81, 2070 (1959). e. J K Williams, E L Martin and W A Shippard, J. Org. Chem., 31, 919 (1966). f. J A Lardgrebeand R H Rynbrandt, J. Org. Chem., 31, 2585 (1966). g. J Kochi and G Hammond, J. Am. Chem. Soc., 75, 3452 (1953).

h. J Hine and W C Baile, Jr., J. Org. Chem., 26, 2098 (1961).

In rigid aliphatic systems, such as 4-substitutedbicyclo[2.2.2]octan-1-carboxylic acid (44),the substituents obey the Hammett equation though with a different set of σ values, defined


by σ1. The σ

1 value represents the electrical effects of a substituent bonded to an sp3 hybridized

carbon atom because the effect is now relayed through σ electrons. Values of σ1 for some groups

are also given in Table (1.4).

COOH

R

(44)

COOH

R

The Hammett equation proved to be the most successful quantitative correlation betweenthe structure of a compound and the equilibria or the rates of the reactions. However, deviationsfrom the equation are also observed in several cases. It was found that non-linear plots of thelogarithm of the rate constant against σ were obtained in reactions such as chlorination andnitration of substituted benzenes and the reaction of benzyl halides with amines i.e., in whichan electron-donating group can mesomerically supply π-electrons and thus has a rate-acceleratingeffect on the rate of the reaction. The rate constants for the solvolysis of meta-substitutedphenyldimethylcarbinyl chlorides give a linear plot against σ constants, but the para-substituentsdeviate from linearity. Probably the most important reason for this deviation is considerableresonance interaction (as shown below) between an electron-donating substituent and thereaction center.

CCH3 CH3

ClCH3 CH3

C C

R R R

+CH3 CH3

+

Different σ values are, therefore, needed to correlate the reactivity of the substituent insuch reactions, Brown and co-workers36 proposed a new set of substituent constants designatedσ+ based on the solvolysis of phenyldimethylcarbinyl chloride in 90% acetone-water at 25°Cas the reference reaction. A reaction constant ρ was derived (equal to – 4.54) from a study ofthe m-substituted chlorides. It was thought that resonance contribution for m-substituentsshould be negligible and the usual σ

m values could be employed for correlation of the rate

constant of the above reaction. Therefore, σ+ values were then assigned to other substituentsR from experimental results of k

R i.e., the rate constant for the substituents R for the solvolysis

of p-substituted cumyl chlorides using, equation (1.16).

logk

k

R

0

F

HGI

KJ= – 4.54σ+ …(1.16)

where k0 is the value of the rate constant for the parent compound. The value of σ+ are based

on rate and not equilibrium constants. The modified Hammett equation is thus expressed inthe following form, (equation 1.17).

logk

k0

= ρσ+ …(1.17)

BASIC CONCEPTS 25

The values of σ+ for some representative substituents are listed in Table 1.4. It is clearfrom Table 1.4 that σ+

p differ considerably from σ

p for highly electron-donating substituents

which reflect a higher degree of resonance between the substituent and the positivelycharged reaction center. The correlation of rate data with σ+ values has also been obtained(see Table 1.5) in a number of cases.

TABLE 1.5: Correlation of Reaction Rates with the Substituent Constants

Reaction Correlation ρρρρρ Interpretation Sourcewith

Hydrolysis of para-substit- σ – 3.6 Electron-donating groups 1uted methyl benzoates, 95% accelerate reaction,H2SO4, 25°C acyl-oxygen fission

Hydrolysis of 4-substituted σ + 2.6 Electron-withdrawing 22,6-dimethylbenzoates, groups accelerate reaction,60 percent dioxane, acyl-oxygen fission40% water, at 82° –174°C

Hydrolysis of 4-substituted σ+ – 3.22 Electron-donating groups 32,6-dimethylbenzoates, accelerate the reactionH2SO4, at 25°C

Hydration of styrenes, σ+ – 3.44 Transition state resembles a 4HClO4, at 25°C carbocation intermediate

Thermal rearrangement of σ+ – 0.43 Electron-release aids the 5arylpropargyl ethers reaction

Claisen rearrangement of σ+ – 0.40 Transition state involves a 6substituted cinnamoyl depletion of electronsp-tolylethers, at 180°C

Benzylic hydrogen abstraction σ – 0.75 Transition state stabilized 7from substituted toluenes in by resonance effectsbenzene at 39.6°C

Benzylic hydrogen abstraction σ+ – 1.46 A positively charged 8from substituted toluenes by transition statetrichloromethyl radical, at 150°C

Base catalyzed oxidation of σ + 1.58 Transition state is 9benzaldehyde with negatively charged.permanganate

Permanganate oxidation of σ+ – 4.3 Large electron deficiency in 10furfurals, H2O at 25°C the transition state

Isomerization of cis-cinnamic σ + 1.30 Oxidation faster with 11acids, H2SO4 at 25°C electron releasing groups

Bromination of acetanilides, σ+ – 12.1 The substituents are capable 12Br2, HOAc-H2O at 25°C of greater resonance

interactions with the ring

Elimination reaction of 2,2, σ + 10.3 Substantial negative charge 132-trifluoroethanesulfonates, develops on the sulfur atom80 percent ethanol at 25°C in the transition state

Contd....


Dissociation of conjugate σ + 0.59 p-Substituents show no 14acids of benzyl phenyl resonance exaltationketones, ethanol at 25°C

Reaction of p-substituted σ – 0.49 Slight development of 15styrenes with 9-borabicyclo positive charge on the(3.3.1) nonane α-carbon of styrene.

Source: (1) B M Wepster et al., Rec., Trav., Chim., 88, 301 (1969). (2) H L Goering et al., J. Am. Chem. Soc., 76, 787(1954). (3) M L Bender and M C Chen, J. Am. Chem. Soc., 85, 37 (1963). (4) W M Schubert et al., J. Am. Chem. Soc.,86, 4729 (1969). (5) M Harfenist and E Thom., J Org. Chem., 37, 841 (1972). (6) W N White and W K Fife, J. Am. Chem.Soc., 83, 3846 (1961). (7) R D Gillion and B F Ward. Jr., J. Am. Chem. Soc., 87, 3944 (1965). (8) E S Huyser, J. Am. Chem.Soc., 82, 394 (1960). (9) K. B. Wiberg and F. Freeman, J, Org. Chem., 65, 573 (2000). (10) D S Noyce and H S Avarbock,J. Am. Chem. Soc., 84, 1644 (1962). (11) F Freeman et al., J. Org. Chem., 35, 982 (1970). (12) H C Brown and M Dubeck,J. Am. Chem. Soc., 82, 1939 (1960). (13) R K Grossland, Jr. et al., J. Am. Chem. Soc., 93, 4217 (1971). (14) A Fischer,et al., J. Am. Chem. Soc., 83, 4208 (1961). (15) L C Vishwakarma and A Fry, J. Org. Chem., 45, 5306 (1980).

It is noticed from Table 1.5 that the reaction involving carbocations usually give largenegative values of ρ and the reactions are facilitated by electron-release.

The Hammett plot (between pkAH

versus σ) is also curved for the dissociation of substitutedphenols. Those substituents which have electron-withdrawing properties mesomerically do notlie on the straight line. In other words certain groups have an added acid-enhancing effect onphenol than they have on the ionization of benzoic acids. Therefore, for such reactions inwhich the reaction center is conjugated with p-substituents which are electron-withdrawing anew constant σ– was defined based on the following reference reaction, equation (1.18).

OH

R R

+ H O2

O–

+ H O3+

...(1.18)

Values of σ– for certain groups are given in Table 1.6.

TABLE 1.6: σσσσσ– Values for Some Substituents

Substituent σσσσσ–

– C6H5 0.08

– C ≡ CH 0.52

– COOH 0.78

– COOC2H5 0.74

– COCH3 0.84

– CHO 1.04

– CN 0.88

– NO2 1.27

Source: O Exner, “A Critical Compilation of Substituent Constants”, Chap X of Correlation Analysis in Chemistry: Recent

Advances, (Eds.) N B Chapman and J Slater, Plenum Press, N.Y. (1978).

The σ– values for groups like –NO2, – CN and – COCH

3 are larger than σ. Thus a better

correlation of the pkAH

of anilinium cations is obtained with σ– rather than σ.

BASIC CONCEPTS 27

REFERENCES

1. S W Benson, J. Chem. Educ., 42, 502 (1965).

2. G Pimental and A McClellan, The Hydrogen Bond, W H Freeman & Co., San Francisco (1959).

3. I D Sadekov et al., Russ. Chem. Rev., 39, 179 (1970).

4. J H Blanch, J. Chem. Soc., (B), 1937 (1966).

5. L L Ferstanding, J. Am. Chem. Soc., 84, 3553 (1962).

6. A McClellan, J. Am. Chem. Educ., 44, 547 (1967).

7. M J S Dewar, Hyperconjugation, The Ronald Press Co., New York (1962).

8. P Von R Schleyer and C W Woodworth, J. Am. Chem. Soc., 90, 6528 (1968).

9. J A Pople and M Gordon, J. Am. Chem. Soc., 89, 4253 (1967).

10. M J S Dewar and A P Marchand, J. Am. Chem. Soc., 88, 354 (1966); and previous papers.

11. R Golden and L M Stock, J. Am. Chem. Soc., 88, 5928 (1966); L M Stock, J. Chem. Educ., 49,

400 (1972).

12. B H Klanderman and W C Perkins, J. Org. Chem., 34, 630 (1969).

13. M S Newman, (Ed.) Steric Effects in Organic Chemistry, John Wiley and Sons, Inc., New York(1963).

14. P von R Schleyer and E Wilkott, Tetrahedron Lett., 2845 (1967); P von R Schleyer et al., J. Am.Chem. Soc., 88, 4475 (166); P von R Schleyer, J. Am. Chem Soc., 86, 1854, 18566 (1964).

15. H C Brown and W J Hammer, J. Am. Chem. Soc., 89, 6378 (1967).

16. E W Fawcett, Chem. Rev., 47, 219 (1950). Also see G Kobrich, Angew. Chem. Int. Edn., (Engl.)12, 4644 (1973).

17. J R Wiseman and W A Pletcher, J. Am. Chem. Soc., 92, 956 (1970); J R Wiseman, J. Am. Chem.Soc., 89, 5967 (1967).

18. J A Marshall and H F Faubl, J. Am. Chem. Soc., 89, 5965 (1967).

19. C F H Allen et al., J. Org. Chem. 27, 1447 (1962).

20. H Pracejus, Ber., 92, 988 (1959); H K Hall and A E Shakeil, J. Org. Chem., 45, 5325 (1980).

21. W Theilacker and E Wagner, Ann., 604, 125 (1963).

22. C C Price and S Oae, Sulfur Bonding, the Ronald Press Co., New York (1962).

23. R J Garratt, Aromaticity, McGraw Hill Book Co. (UK) Ltd., London (1971).

24. G M Badger, Aromatic Character and Aromaticity, Cambridge University Press, London (1969);H F Vulpin, Russ. Chem. Rev., 29, 129 (1960).

25. T J Katz, J. Am. Chem. Soc., 82, 3784 (1960); H L Strauss et al., J. Am. Chem. Soc., 85, 2360(1963).

26. E Vogel et al., Angew Chem., 76, 786 (1964); E Vogel and W A Boll, Angew. Chem., 76, 784(1964).

27. F Sondheimer, Acc. Chem. Res., 5, 81 (1972).

28. F Sondheimer and Y Gaoni, J. Am. Chem. Soc., 82, 5765 (1960).

29. Y Gaoni and F Sondheimer, J. Am. Chem. Soc., 86, 521 (1964).

30. J P Snyder (Ed.), Non-benzenoid Aromatics, Academic Press, New York (1969).

31. R Breslow, J. Pure Appl Chem., 28, 111 (1971); R Breslow, J. Am. Chem. Soc., 89, 4383 (1967).

32. D J Raber et al., J. Org. Chem., 53, 2117 (1988).

33. C D Johnson, “The Hammett Equation”, Cambridge University Press, London (1973).

34. A Fisher et al., J. Am. Chem. Soc., 83, 4208 (1961).

35. C W McGraw et al., J. Am. Chem. Soc., 77, 3037 (1955). For a review see L M Stock and H C

Brown, Adv. Phys. Org. Chem., 11, 35 (1963).


REVIEW PROBLEMS

1.1. Offer reasons for the following observations:

(a) Nitrobenzene has a larger dipole moment than aniline and phenol but is much less solublein water.

(b) Tropolone does not form 2, 4-dinitrophenylhydrazone.

(c) Cyclobutadiene is unstable.

(d) Ionization constant (Ka) of acetic acid is 2.7 × 10–6 in 50 percent methanol-water but that of

triethylacetic acid is 0.36 × 10–6.

(e) The σ value for F is larger than for Cl.

(f) Phosphorus pentachloride has zero dipole moment whereas phosphorus trichloride has adefinite value.

(g) Methoxy group (– OCH3) has σp value of – 0.268.

(h) Deuteration of methyl group of acetophenone lowers the basicity of the molecule.

(i) Treatment of cycloheptatriene with potassium metal dispersed in benzene does not yield thecorresponding cycloheptatrienide anion.

1.2. Predict the approximate magnitude and sign of ρ for the following reactions and draw structuresfor the transition state in support of your answer.

(a)

OCH3

+ CH COCl3

OCH3

AlCl3

C H Cl , 25°C2 4 2

COCH3

(b)

N(CH )3 2

+ CH I3

N (CH ) I+ –

3 3

CH OH, 65°C3

(c)

CHO

+ HCN

CH(OH)CN

90% ethanol, 20°C

(d)

CN

+ + H S OH2–

C

C H OH2 5

NHHS

Heat

(e)

COOH

+ CH OH3

COOCH3

H+

25°C

(f)

CH(OH)CH COOH2 CH

H , H O+

2

45°C

CHCOOH

BASIC CONCEPTS 29

(g)

CH NHCOCH2 3

+ Cl2

CH N(Cl)COCH2 3

CH COOH3

180°C

(h)

CH F2

+ OH–

CH OH2

+ F–

1.3. The following three nitrophenols have the melting points indicated below each compound. Discuss

this progressive change in their melting points.

OH OH OH

NO2

NO2

NO2

45°C 96°C 114°Cmp

1.4. Write the products of the following reactions. If no reaction takes place write NR.

(a)

COOH

O

HOOC COOH

OHOOC

200°C(b) 245°C

COOH

O

O

(c) 260°C

COOH

O (d) 145°C

COOHO

1.5. (a) Which compound of the following pairs is a stronger acid?

COOH

COOH

COOH

COOHor

(C H ) CH6 5 3 or triptycene

(CH ) N CH3 3 2

+COOHor(CH ) N3 3

+COOH

(b) Which compound of the following pair is a stronger base?

CH3 CH3

orN N

CH3CH3

CH3CH3

NH2 NH2

CH3

CH3

CH3

CH3

NO2 NO2

or


(c) Explain the difference in the magnitude of dipole moments of the following compounds:

(CH ) N3 2 (CH ) N3 2NO2 CF3 N(CH )3 2

µ = 6.89 D µ = 4.62 D µ = 1.50 D

(d) Which of the following compounds do you expect to volatalize easily?

OH

NH NHCH CH

N Nor

(e) Explain the difference between the dipole moments of the following two isomeric compounds:

O N2 O N2

C CH H

C CNO2

NO2H

H

µ = 7.38 D µ = 0.50 D

1.6. The value of the rate constant k for the solvolysis of an unsubstituted organic halide is

1.6 × 10–4 s–1 at 40°C and ρ is + 3.78. Calculate the rate constant of the reaction for a

p-bromosubstituted compound if σp– Br is 0.232.

1.7. Could the following two enantiomers be isolated?

HOOC HOOC

Br Br

Br Br

COOH COOHand

1.8. Benzenesulfonyl chloride and substituted benzoate anions react in CH3OH at 25°C. The reaction

follows a second order kinetics and the following rate constants were obtained.

C H SO Cl + O6 5 2–

C H SO6 5 2C C

O Ox x

O

X σσσσσx 103 k2/l/mol/sec

p – OCH3 – 0.28 6.58

p – CH3 – 1.170 5.75

m – CH3 – 0.06 5.33

H 0 5.18

p – Cl 0.22 4.27

p – Br 0.23 4.04

m – Br 0.39 3.80

m – NO2 0.71 2.91

p – NO2 0.78 2.78

Calculate ρ for this reaction. Compare this value for the same reaction when substituted acetate

anion is replaced by substituted anilines (ρ = – 2.15).

[ J. Org. Chem., 45, 4966 (1980)]

BASIC CONCEPTS 31

1.9. Answer the following and give reasons.

(a) Would you expect compound (A) to show a reduced reactivity towards acetylation?

H

O

N

(A)

(b) Which of the two (B) or (C) will have slower rate of isomerization?

C H6 5 C H6 5

C H6 5 C H6 5

H H

C C

O O

C H6 5 C H6 5or

(B) (C)

(c) Which effect should determine the acidity of the acids (D) and (E) ? Which is more acidic?

Cl

Cl

C C

C C

COOH COOH

and

(D) (E)

1.10. Explain the difference in pKa values of the following acids:

N COOH

pKa 4.40 4.60 5.10

N COOH

CH3

N COOH

CH3

1.11. The dipole moment of the following compound has been estimated to be 5.6D. Explain.

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