2. Delocalized Chemical Bonding and Electronic Effects – Advanced Organic Chemistry


Delocalized Chemical Bonding and Electronic Effects


By the end of this chapter you should be familiar with

  • Interaction between orbitals over many bonds, stabilization by the sharing of electrons over molecules.
  • Molecular shape and structure that determine reactivity and represent one aspect of structure by curly arrows.
  • The effects that delocalized electrons have on the reactivity of organic compounds.
  • Resonance, resonance effects, hyperconjugation tautomerism, inductive, electromeric, and steric effect and hydrogen bonding, and their role on chemical reactivity.

Electrons that are restricted to a particular region are called localized electrons. The bonding in most of the organic compounds can be described by a single Lewis structure. However, in several other compounds a Lewis structure does not represent a correct representation for a molecule or ion. In these compounds bonding orbitals are not restricted to two atoms, but are spread out over more than two. This type of bonding is said to be delocalized. Delocalized electrons are shared by more than two atoms. Many organic compounds contain delocalized electrons. The properties, such as colour of molecules depend on the conjugation of electrons. How is this π framework responsible for the unexpected stability of benzene and other aromatic compounds? This chapter describes resonance, resonance effects, hyperconjugation tautomerism, inductive, electromeric and steric effects and hydrogen bonding, and their role on chemical reactivity on molecules.


A compound with delocalized electrons is said to have resonance; that is, resonating structures vibrate rapidly between localized structures. The approximate structure using localized electrons is called a resonance contributor, a resonance structure, a contributing resonance structure, or a canonical structure. The actual structure drawn using delocalized electrons is called a resonance hybrid, which is the real structure. In drawing canonical forms and in deriving the true structures from them, we are guided by certain rules:

  1. All the canonical forms must be bonafide Lewis structures, which exist only in one’s imagination.
  2. The position of the nuclei must be the same in all the structures, that is, resonance structures differ only in electron distribution; all the atoms stay in the same places.
  3. All atoms taking part in the resonance must lie in a plane, or nearly so.
  4. All canonical forms must have the same number of unpaired electrons.
  5. The energy of the actual molecule is lower than that of any contributing form.
  6. All canonical forms do not contribute equally to the real molecule. Each form contributes in proportion to its stability.

The structures of a large number of organic compounds can be written with the help of simple bond diagrams; e.g. ethylene as C2H4, acetylene as C2H2, etc. There are, however, many compounds for which simple bond diagrams do not accurately describe the molecules. One such example is benzene, molecular formula C6H6. Structure I gives the impression that it is a cyclic compound of six carbon atoms containing three single and three double bonds. If this were so, you would expect two values of carbon–carbon bond lengths, namely, one for single bonds (nearly 1.54 Å as in ethane) and the other for double bonds (nearly 1.33 Å as in ethylene). Experimental evidence (X–ray diffraction studies) shows that all the six carbon bonds in benzene are equal and have a length of 1.39 Å (which is in between 1.54 and 1.33 Å). Some examples of resonance contributors are as follows:

We know that although benzene is unsaturated it does not undergo addition reactions readily (characteristic of unsaturated compounds), but takes part in a substitution reaction (characteristic of saturated compounds). To account for the high degree of stability and unusual chemical properties, benzene is considered to be a “hybrid” of the following hypothetical structures I – V:

These structures are called resonance structures, contributors or canonical forms. The two ‘Kekule’ forms, I and II, are of lower energy (more stable, contributes 39% of the actual molecule) than the three ‘Dewat’ forms, III to V, which contribute 7.3% each. Structures I and II contribute more to the hybrid than III, IV or V; hence, the properties of benzene can be expected to resemble either I or II more closely than III, IV or V. Since I and II have the same energy, each would contribute to the hybrid by the same amount. The symbol of resonance, the double-headed arrow (↔), does not indicate equilibrium. The canonical structures I–V are not physically interconvertible but are hypothetical structures because of different electronic arrangement due to shift of electrons and in lone pairs within the molecule. The actual structure, drawn using localized electrons, is the resonance hybrid (VI).

The term hybrid has some analogy to biological hybridization. A hybrid produced by crossing two species has some of the characteristics of its parents but is distinctly different from both of them. However, the two species that are crossed are real, whereas in resonance, the contributors are hypothetical. Thus any canonical form alone does not represent the actual molecule, and a simple bond diagram cannot show the structure of the actual molecule. The following examples illustrate the applications of resonance.

In sodium formate the two C–O bonds have the same length (1.27 Å), whereas in formic acid the two C–O bonds lengths are different (1.23 and 1.36 Å). Normal C–O and C=O bond lengths are about 0.14 and 0.12 nm, respectively. Both C–O bonds in methanoate ion have a length in between, that is, 0.13 nm. The resonance structures for sodium formate and formic acid are shown below.

The contributors a and b of sodium formate have the same energy and hence contribute equally to the hybrid. The two C–O bonds in sodium formate have the same character and length. Sometimes it is represented by a non-classical structure c in which the dotted curved line indicates that the two C–O bonds have an identical nature, which is in between of a single bond and a double bond and in which the negative charge is not localized on any oxygen atom but is spread over. In a classical ion the charge is shown localized at any one atom, e.g. structures a or b. Formic acid has been represented by resonance structures d, e and f in which e and f have high energy, as these require separation of charges. Hence, e and f are expected to contribute only to a lesser extent than d to the resonance hybrid. Thus in formic acid, the two C–O bonds will have different characters (one as a single bond and the other as a double bond) and lengths.

The arrows show the shift of electron-pairs in π or σ molecular orbitals, though the framework remains unaltered in the above resonance structures. It is known that electrons are more mobile in π orbitals than in the σ orbitals. The shift of electron-pairs in the above structure has no physical reality. It only helps to show the spread or delocalization of the π electrons. Note that the direction of arrows is from electron-rich (negative charge or lone–pair) to electron-poor sites in the molecule. We can summarize the features that decrease or increase the predicted stability of a contributing resonance structures as follows:

Factors decreasing stability of contributing resonance structures

  1. An atom with an incomplete octet
  2. A negative charge that is not on the most electronegative atom or a positive charge that is not on the least electronegative atom.
  3. Charge separation

Factors increasing contribution of contributing resonance structures to resonance hybrid

  1. The greater the predicted stability of a resonance contributor, the more it contributes to the hybrid.
  2. The greater the number of relatively stable resonance contributors, the greater the resonance energy.
  3. The more nearly equivalent the resonance contributor, the greater the resonance energy.

Low Reactivity of Benzene If benzene is more stable than the alkenes, we expect benzene to be less reactive (more energy needed for reaction). For example,

  1. Hydrogenation: Alkene can be hydrogenated by H2 and Ni catalyst at room temperature while benzene requires 200°C and 200 KP.
  2. Bromination: No catalyst is needed for the reaction of Br2 with alkenes, but bromination of benzene needs a metal bromide.
  3. Polymerization: Phenylethene has both a benzene ring and an alkene but when it is polymerized with acids or radicals only the alkene reacts.

A compound with delocalized electrons is more stable than that in which electrons were localized. The extra stability a compound gains as a result of having delocalized electrons is called delocalization energy or resonance energy. That is, the difference in energy between the actual molecule and the Lewis structure of lowest energy is called resonance energy. Let us take a look at the resonance energy of benzene. The heats of hydrogenation of cyclohexatriene (hypothetical) and benzene, determined experimentally, are given here.

The enthalpy change on hydrogenation or on combustion of benzene is less exothermic than ‘expected’ in terms of a structure with three alkene C=C bonds. The heat evolved when hydrogen is added to cyclohexene (having one C=C bond) is −28.6 Kcal/mol. The expected value of the heat evolved when hydrogen is added to cyclohexatriene (hypothetical) (having three C=C bonds) should be 3 x 28.6 = −85.8 Kcal/mol, but the experimental value of the heat of hydrogenation for benzene, which is much smaller than that calculated for the hypothetical ‘cyclohexatriene’ molecule, is −49.8 Kcal/mol.

  1. Hydrogenation
  2. Combustion:

    C6H6 (1) + 15/2 O2 (g) → 6 CO2 (g) + 3H2O(1)
    ΔH = − 3313.8 KJ / mol;
    Expected value is −3473.4 KJ/mol

    Difference = 3473.4 − 3313.8 = 159.6 KJ/mol (38kcal/mol)
    360.4 − 209.2 = 151.2 kJ/mol (36kcal/mol)

This tells us that benzene is more stable (having a lower energy content) than the hypothetical molecule containing three isolated* C=C bonds by −85.8 − (−49.8) = −36.0 Kcal. This energy difference is called the resonance energy of one mole of benzene. It is understood that the benzene molecule is stabilized by resonance. Resonance energy** is the difference in energy between the real conjugated molecule and its imaginary unconjugated or hypothetical analogue. The greater the number of relatively stable resonance contributors, the greater the resonance energy. The more nearly equivalent the resonance contributors, the greater the resonance energy. A molecule is said to be stabilized by resonance if the amount of resonance energy is considerable.

Figure 2.1 Reaction coordinate diagrams for the hydrogenation of ‘cyclohexatriene’ and benzene

In the above reaction coordinate diagrams, the energy of the product will be the same in both cases because both reactions provide the same product. The only way to account for the difference in the heats of hydrogenation is for the reactants to have different stabilities. Since the experimental heat of hydrogenation for benzene is 36 Kcal/mol less than the heat of hydrogenation calculated for ‘cyclohexatriene’, benzene must be 36kcal/mol more stable than ‘cyclohexatriene’. Breaking the bonds in benzene is more endothermic than it would be for cyclohexatriene; the difference is used as a measure of the delocalization energy.


Can the phenomenon of resonance be used to explain the reactivity of compounds? The answer is Yes. The acidic and basic character and also the preferential reactivity centres in aromatic compounds can be explained using resonance. The electron density of the unshared pair does not reside entirely on the nitrogen, but is spread over the ring in aniline. This decrease in electron density at one position (and corresponding increase elsewhere) is called the resonance or mesomeric effect. The resonance effect depends upon the overlap of certain orbitals.

For example, carboxylic acids are stronger acids as compared to alcohols though each contains a hydroxy group attached to carbon. Let us see what happens when a base (a proton-acceptor) is present. By proton loss, carboxylic acids give carboxylate anions, which are stabilized by resonance; no such stabilization is possible for the alkoxide ions, which are formed by proton loss from alcohols. Since the two contributors to the hybrid in the carboxylate anion have the same energy, the stabilization due to resonance is considerable. The negative charge is spread over as shown in the non–classical structure of carboxylate anion.

Since the alkyl group R has carbon atoms, which are sp3 hybridized, it cannot participate in resonance. Resonance involves only electrons in p orbitals or lone pairs. Resonance stabilization of phenoxide ion also explains the greater acidity of phenols as compared to alcohols. The negative charge on oxygen in phenoxide ion is spread over the benzene ring imparting stability; hence, the equilibrium is shifted to the right. It is to be noted that in the resonance structures the negative charge is shown at alternate positions on the benzene ring.

Although resonance structures can be written for phenol itself, stabilization due to resonance is not much as the three contributing forms require separation of charges. Note that the hydroxy group is a −I group (that is, it withdraws electrons by inductive effect), but it is electron-donating by resonance. The inductive effect is small as compared to the resonance effect.

The lower pKa values of acetone and nitromethane (20 and 10.2, respectively) can be explained on the basis of resonance stabilization of their anions.

Aniline is less basic than aliphatic amines. The lone-pair electron density on nitrogen in aniline is reduced as it is spread over the benzene ring by resonance. In aliphatic amines (RNH2) no resonance is possible as the alkyl group contains sp3 hybridized carbon atoms. Moreover, alkyl groups are +I group in character and will increase the electron density on nitrogen, thus increasing the basicity of amines.

Amides are less basic than amines because the electron-pair on nitrogen in amides is conjugated to C=O group, and its density will be reduced due to resonance. The amide group is planar, that is, the first carbon atom of the R groups attached to the carbonyl group and nitrogen lie in a plane. Since the lone pair-electrons on nitrogen are delocalized into the carbonyl group, the C–N bond is strengthened.

Groups that withdraw electrons by resonance or conjugation are called −R groups. Some examples are nitro (NO2), cyano (CN), carbonyl (C=O), carboxylic (COOH), sulphonic (SO3H) groups, etc. Groups that donate electrons by resonance or conjugation are called +R groups. Some of these are hydroxy (−OH), amino (−NH2), alkoxy (−OR), halogens (−X) and alkyl amino (−NHR and −NR2.

We know that benzene and its derivatives undergo electrophilic substitution reactions that occur because of electrophiles (electron poor reagents). The application of resonance method can tell us about the reactivity of benzene derivatives in the above reactions.

The presence of −R groups on the benzene ring decreases reactivity, whereas the presence of +R groups increases the reactivity in electrophilic substitution reactions. Electron withdrawal from benzene ring would decrease the electron density in benzene, and hence, the attack on the electrophilic reagent would not take place readily; thus, −R groups deactivate the benzene ring. On the other hand, +R groups activate the benzene ring towards electrophilic substitution, as the electron donation will increase the electron density on the benzene ring. The NO2 group is a −R type. It creates positive centres on benzene carbons. The OH and NH2 groups are +R types and create negative centres on benzene carbons.

We can also predict the position of substitution on the benzene ring in monosubstituted benzene by following the resonance approach in which the nature of the group already present on the benzene ring directs the incoming group. There are two types of groups. One is ortho- and para-directing and the other only meta-directing. These positions are shown in a monosubstituted benzene derivative as:

If Y is a +R group, it will be an ortho- and para-directing type; such groups create negative charges at ortho (2 or 2’) and para (4) positions through which the attack on electrophile would occur. If Y were a −R group, it would create positive charges at ortho and para positions so that attack on electrophiles could occur preferably through meta (3 or 3’) positions; in this case, Y will be a meta-directing group. It is known that alkyl groups are ortho- and para-directing. Here a different approach can be made. Let us add an electrophile E+ at each of the three positions, that is, ortho, meta and para and write all possible resonance structures for the intermediates. Resonance effects of substituents help in explaining directive and rate-controlling factors in electrophilic aromatic substitution.

Structures a and b are specially stable and favoured as the alkyl group (Y = R) is a +I type; it will neutralize the positive charge on carbon most efficiently. The Alkyl group does not affect other structures, as we know that the inductive effect decreases with distance. Therefore, alkyl groups will be ortho- and para-directing. Similar treatment can also be done in cases of the other −R or +R types of groups.

Many examples are known where resonance is reduced or prevented because the atoms are sterically forced out of planarity. Bond lengths for the o- and p-nitro groups in picryl iodide are quite different. Distance a is 1.45 Å, whereas b is 1.35 Å. The obvious explanation is that the oxygens of the p-nitro group are in the plane of the ring and thus in resonance with it, so that b has a partial double-bond character, whereas the oxygens of the o-nitro groups are forced out of the plane by the large iodine atom.

Another example is 2,3-di-t-butylbutadiene, which is a stable, nonconjugated diene. The double bonds are not in the same plane but are forced out by the large t-butyl groups.


Table 2.1 Summary of the Main Electrophilic Substitutions on Benzene


By the field effect alone, the order of electron release for simple alkyl groups connected to an unsaturated system is tert-butyl > isopropyl > ethyl > methyl. However, Baker and Nathan in 1935 observed that the rates of reaction with pyridine of p-substituted benzyl bromides were about opposite than expected from electron release by the field effect, hence it is often referred to as the Baker-Nathan effect. The methyl-substituted compounds reacted fastest and the tert-butyl substituted compound reacted slowest. It is a special case of resonance in which a σ bond is broken and a new π bond is formed.

The molecular structures are unchanged in going from the gas phase into solution. A charged species is more stable if its charge is spread out (delocalized) over more than one atom. Hyperconjugation in the above case may be regarded as an overlap of the σ orbital of the C–H bond with the empty p-orbital, analogous to the π–π orbital overlap. Delocalization of electrons by overlap of a carbon–hydrogen or carbon–carbon σ bond orbital with an empty p-orbital is called hyperconjugation. It occurs only if the σ bond orbital and the empty p-orbital have the proper orientation. Muller and Mulliken, termed hyper conjugation in the ground state of neutral molecules as sacrificial hyperconjugation where canonical forms involve not only no-bond resonance, but also a charge separation not possessed by the main form.

In free radicals and carbocations the canonical forms display no more charge separation than the main form and is called isovalent hyperconjugation. Hence, with respect to this effect, CH3 is the strongest electron donor and t-butyl the weakest

The more resonance form like VIII can be written with the two other hydrogen atoms on the β carbon atom. In this representation, a shift in the electron density towards the carbon–carbon bond in VII has left the carbon–hydrogen bond with a partial no-bond character. This is also known as ‘no-bond resonance’. It is thought to play an important part in the stabilization of carbenium ions. One explanation has already been given to you based on the inductive effect (+I) of the alkyl groups. The other explanation is based on the relative stabilization of positive charge on carbon due to hyperconjugation (depending upon the number of β hydrogen atoms*).

Hence, the more the number of β hydrogen, greater will be the stabilization due to hyperconjugation. This principle can also been applied to explain the relative thermodynamic stability of methyl substituted ethylene.

The application of hyperconjugation is controversial, as a weaker double bond (=) is formed at the expense of a stronger single (−) bond. Hyperconjugation effects also account for the higher resonance energies of compounds. For example, toluene has resonance energy 37.5 Kcal/mol, which is 1.5 Kcal/mol higher than that of benzene (36 kcal/mol), which means that toluene is more stable than benzene. Dipole moment data support the same general view. Benzene has zero dipole moment whereas toluene possesses an appreciable value (0.37 D). The basicity of di- and trimethyl-substituted benzenes is greater than toluene for the same reason. Further support for hyperconjugation comes from the C-C bond distance data. The C-C bond distance in propene is slightly shorter than that in ethane and the C=C bond distance is 1.35 Å in ethene. The heat of hydrogenation of propene (126 KJ/mol) is less than that of ethene (137 KJ/mol). Hence, it is reasoned that C2-C3 bond requires a double bond and the C1-C2 bond a single bond character because of hyperconjugation.

2.5.1 Negative hyperconjugation

Negative hyperconjugation is a term that was coined to describe the stabilization of anionic species by σ-delocalization, involving σ*-orbitals as acceptors. An extreme case is the trifluoromethoxide anion, a highly reactive species, but stable enough in the solid state for the crystal structure to be determined. The C-O bond has almost the same length as ordinary C=O double bonds, whereas the C–F bonds are significantly longer than normal. The simplest explanation is that there is strong σ-delocalization of the non-bonding electrons on O into the antibonding orbitals of the C-F bonds. Negative hyperconjugation, anomeric effect and hyperconjugation are three terms that describe basically the same effects in different situations.


A ketone and an enol differ from each other only in the location of a double bond and in the location of hydrogen. Such isomers are called tautomers. Tautomers are isomers differing only in the positions of hydrogen atoms and electrons; the carbon skeletons are the same. The conversion of one tautomer into an other tautomer is called tautomerization. If an aldehyde or ketone possesses at least one hydrogen atom on the carbon atom adjacent to the carbonyl group, called the alpha (α) carbon, this hydrogen can migrate to the oxygen atom. As a result, a carbonyl compound with alpha hydrogen can exist in two isomeric forms, called tautomers. In the keto form the hydrogen is attached to the alpha carbon, while in the enol form it is attached to the carbonyl oxygen with the migration of the double bond. Keto–enol interconversion is also called keto–enol tautomerization or enolization. The tautomers of acetaldehyde are shown here.

Acetone exists in equilibrium almost wholly in its keto form, as the enol form is not stabilized through either resonance or the intramolecular hydrogen bonding.

The enol form is given that name because it is a combination of ene for the double bond and ol for the −OH (hydroxyl) group. The two isomers exist in equilibrium. In any such case, both tautomers are present, but in simple cases the keto form is much more stable than the enol form. For example, in acetaldehyde about 6 molecules out of every 10 million molecules are in the enol form at any given time. Nevertheless, the equilibrium always exists, and every molecule of acetaldehyde (as well as any other aldehyde or ketone with an alpha hydrogen) is converted to the enol form (and back again) several times per second. This is an important characteristic because a number of reactions of carbonyl compounds take place only through the enol forms. The fraction of enol tautomer is considerably greater for a β-diketone because its enol tautomer is stabilized by internal hydrogen bonding and by conjugation of carbon–carbon double bond with the second carbonyl group. For example, in the case of 2,4-pentanedione, 80 of every 100 molecules are in the enol form.

Ethyl acetoacetate (CH3COCH2COOC2H5) gives the reactions of a carbonyl (C=O) group, such as the formation of cyanohydrin with HCN, bisulphate addition compound with sodium bisulphate and phenyl hydrazone derivative with phenyl hydrazine. In addition to these, it behaves as an unsaturated alcohol (enol) as it evolves hydrogen on treatment with sodium, shows reddish–violet colouration with ferric chloride and decolourizes bromine water and adds diazomethane. These chemical properties of ethyl acetoacetate suggest that it exists in two forms: a saturated ketone A (keto form) and an unsaturated alcohol B (enol form), which is an example of prototropy that can be shown as:

The hybridization of carbon flanked by carbonyl and ester groups changes from sp3 in A to sp2 in B. These two forms exist together so that properties of both keto and enol are observed in the reactions of ethyl acetoacetate. Structures A and B are isomeric and differ in the relative position of hydrogen atoms and the position of the double bond. Since a 1,3–proton shift takes place in the above example from carbon atom to oxygen atom, this form of tautomerism is known as prototropy. The enol form is favoured in relatively non-polar solvents such as alkanes and carbon disulphide and also by dilution. In hydroxylic solvents such as water, methanol, ethanol and acetic acid etc., the keto form is favoured due to the formation of hydrogen bonds between the oxygen atom of the carbonyl group and the hydrogen atom of the solvent molecule (ZH).

On dilution, the intermolecular hydrogen bonds (between two different types of molecules like A and ZH are broken and, therefore, the percentage of enol form in the mixture increases. Two factors usually stabilize the enol form.

  1. Intramolecular hydrogen bonding: Hydrogen bond formation within the same molecule as shown here in the enol form of 1,3–diketones.
  2. Resonance: The enol form (phenol) of either 2,4–cyclohexadienone or 2,5–cyclohexadienone (keto forms) is stabilized by resonance to a greater extent than a similar stabilization of the keto forms, which require charge separation. Phenol is unusual in that its enol tautomer is more stable than its keto tautomer, because the enol tautomer is aromatic but the keto tautomer is not.

    Sometimes, because of resonance, the keto form is stabilized more as compared to the stabilization in the enol form due to intramolecular hydrogen bonding. This explains the larger percentage of the keto form as compared to the enol form of ethyl acetoacetate. The resonance effect is more pronounced in diethylmalonate where the keto form gets stabilized and hence only traces of enol are present.

In the liquid state, the percentage of enol forms in acetylacetone and benzoyl acetone is 80 and 89, respectively. In addition to intramolecular hydrogen bonding shown above for acetyl acetone, the enol form of benzoyl acetone gets further stabilization because of the conjugation of the carbon–carbon double bond with the phenyl group.

In a series of cis-fixed bicyclic diketones I, II and III the percentage enol varies from 100 to 80 and 1.4, respectively; in this series the hydrogen bond becomes obviously longer and weaker. For disubstituted 1,3-diketones (IV) the enol content increases rapidly as the size of the alkyl substituent is increased, when R is tert-butyl, the keto form is undetectable.

The keto-enol equilibrium has often been used as a measure of the resonance stabilization; for example, no enolization occurs in 3-cyclobutenones (V) and no ketone can be detected in phenol (VI). Fluoroglucinol (VII) chemically behaves more like a triketone than a triol; thus, it readily forms the trioxime (VIII) with hydroxylamine.

Sometimes, conformation plays a dominant role in deciding the extent of enolization in certain α-diketones. Biacetyl, for example, has very little enol content in contrast to 1,2-cyclopentanedione, which exists mostly in the enolic form. The carbonyl groups of the biacetyl are oriented in opposite directions so as to reduce dipole–dipole repulsion, whereas cyclopentane-l,2-dione (C) being a cyclic compound has no such choice and has to resort to enolization in order to avoid this repulsion.

These are some examples of keto–enol tautomerism. Besides these, some other systems also exhibit tautomerism.

2.6.1 Mechanism of Keto–Enol Interconversion

Keto–enol interconversion may happen in basic as well as acidic solution. The steps are reversed in the base- and acid-catalyzed reactions. In the base-catalyzed reaction, the first step is the removal of an α-proton and the second step is the protonation of oxygen. In the acid-catalyzed reaction, the first step is the protonation of oxygen. In the acid-catalyzed reaction, the first step is the protonation of oxygen and the second step is the removal of a α-proton.

2.6.2 Differences between Tautomerism and Resonance

  1. Tautomerism usually involves making and breaking of σ as well as π bond. While in resonance, only the electrons in π bonds or lone-pairs on heteroatoms shift; the σ framework is not disturbed. This difference results in a shift of an atom from one position to another in tautomerism. In resonance there is no shift of any atom.
  2. Tautomerism may involve a change in hybridization of atoms resulting in a change in shape of the molecule. In resonance there is no change in hybridization or the geometry.
  3. The two tautomeric forms exist together though the equilibrium may shift to either side with a change in condition. The tautomers have a physical reality whereas the resonance structures are imaginary.

The aliphatic nitro compounds exist in equilibrium with aci form. The aci-nitro form E, is much less stable than the nitro form D. This is due to resonance stabilization of the nitro form as against in nitroso case where such resonance is not possible.

In this example, a 1,3–shift of a proton from carbon in D to electronegative oxygen takes place. The nitro form is also known as the pseudo-acid form as it behaves as an acid in the presence of strong alkali. The aci-form is known as nitronic acid. It does not dissolve in weak bases like sodium carbonate, but is readily soluble in stronger bases like sodium hydroxide. Thus, the aci form behaves as a weak acid and is deprotonated in strong bases to give salt F. This disturbs the equilibrium leading to the formation of salt F when excess of base is used.

Acidification of the sodium salt of nitronic acid F regenerates the nitro form D, which can be seen by the appearance of oily drops through the aci form E.

Phenyl nitromethane is yellow oil. It is neutral to litmus. It slowly dissolves in alkali solution. On treatment with acetic acid it gives a colourless solid, mp 84°C, which is the aci-form. The aci form dissolves in sodium bicarbonate solution (behaves as an acid) and is slowly converted into the more stable nitro form.

A similar 1,3–proton shift occurs in the diazo-amino systems G and H and imine-enamine systems J and I.

Acidic hydrolysis of either G or H gives a mixture of phenol, p-toluidine, nitrogen, aniline and p-cresol. This supports the existence of equilibrium between G and H. The first three compounds are expected from G and the latter three from H.


In a polar covalent single bond, the electron pair forming the σ bond never lie exactly between the two atoms; they tend to be attracted towards the more electronegative atom of the two that creates a permanent polarization in the ground state of a molecule. Thus, inductive effect — a charge polarization through a series of σ bonds — causes a shift of electron density from the acidic site and thus stabilizes the anion formed by deprotonation. Because an inductive effect is transmitted through bonds, the effect is greater when transmission is through fewer bonds. Functional groups can be classified as electron-withdrawing (−I) or electron-donating (+I) groups relative to hydrogen. From the pKa values of acetic acid and chlorosubstituted acetic acids given in Table 2.1, it can be seen that substitution of hydrogen by chlorine atoms in CH3 group of acetic acid increases the acidic character.


Table 2.1 pKa Values of Acetic Acid and Chlorosubstituted Acetic Acids

Acetic acid CH3COOH 4.80
Monochloroacetic acid ClCH2COOH 2.86
Floroacetic acid FCH2COOH 2.66
Dichloroacetic acid Cl2CHCOOH 1.30
Trichloroacetic acid CI3CCOOH 0.65

In monochloroacetic acid, an electron-withdrawing chlorine atom has replaced one of the three hydrogen atoms in acetic acid. The electron-pair constituting the Cl–C bond is drawn close to the chlorine atom. This effect is transmitted through other atoms forming bonds to the OH bond in the acid, resulting in a shift of the electron-pair constituting the OH bond towards the oxygen atom. The outcome of this electron withdrawal by the chlorine atom (shown below) facilitates the departure of the proton (hydrogen atom without its electron) and thus increases the acidic character of chloroacetic acid as compared to acetic acid.

In the di- and trichloroacetic acids, the presence of the second and third chlorine atoms results in more electron withdrawal away from hydrogen of the O–H bond and would, therefore, increase the acidity of these compounds as compared to acetic acid. The other consideration would increase the acid character [compare (1) and (2) in Table 2.2]. As the distance of this group from the reaction site increases, its effect becomes less.


Table 2.2

Acids pKa
(1) CH3COOH 4.80
(2) (CH3)3+NCH2COOH 1.83
(3) (CH3)3+N(CH2)4COOH 4.27
(4) O2CCH2COOH 5.69
(5) O2C(CH2)4COOH 5.41
(6) HO2CCH2CO2H 2.83 (pKa1)

The presence of the electron-withdrawing carboxyl group (COOH) in place of hydrogen increases acidity [compare (6) and (1)]. The departure of the first proton from malonic acid (6) is easier as compared to that of the second proton (4). If the presence of a positively charged group increases the acidity of a molecule, a negatively charged group should decrease it [compare (1) and (4) above]. The negatively charged carboxylate anion (COO) withholds the proton and makes its departure difficult (electrostatic consideration). In polybasic acids the first ionization constant is always higher than the second or third.

Formic acid is more acidic as compared to acetic acid. The structural difference between the two is that a methyl group in acetic acid has replaced a hydrogen atom in formic acid. The methyl group has an electron-donating nature. When this effect is transmitted to the OH bond, the departure of a proton becomes more difficult in acetic acid than in formic acid.

The inductive effect progressively falls off with the increasing distance of the electron- attracting substituent from the carboxyl group. This is well exemplified by the significant difference in the pKa values of 2-, 3- and 4-chlorobutanoic acids (pKa 2.8, 4.0 and 4.5, respectively).

Some examples of the electron-donating groups (+I groups) and the electron-withdrawing groups (−I groups) are given in Table 2.3.


Table 2.3 Examples of Electron-donating Groups and Electron-withdrawing Groups

If the electron-donating effect of the methyl group decreases the acidity of acetic acid as compared to that of formic acid (described above), its presence should increase the basicity of a compound. This is observed when methyl groups replace the hydrogen atoms in ammonia to give methylamine and dimethylamine successively. The decreasing order of the basic character in water is:

We have seen earlier that ammonia and amines can act both as Bronsted bases (proton acceptors) and also Lewis bases (electron-pair donors), in which the amine behaves as Lewis bases towards the Lewis acid triethyl borane. In this reaction the decreasing order of basicity of amines is: methylamine > dimethylamine > trimethylamine.

Consider the following:

This change in the order of basic character is due to the large size of the acid (triethyl borane) as compared to a proton. This is caused by the steric effect, which will be explained later.

The electron-donating effect of alkyl groups has been used to explain the relative stability of carbenium ions. One of the ways of getting these ions is by ionization of alkyl halides in polar solvents:

The above equilibrium is shifted to the right when R is a tertiary ion rather than a secondary or a primary ion. Thus, the order of decreasing stability is: tertiary > secondary > primary. An example is

One possible explanation for the relative stability of the above ions is that a has three methyl groups (+I group) to stabilize the positive charge as compared to two methyl groups in b and only one methyl group in c.


This effect involves the transfer of the electron-pair towards one of the atoms constituting a double bond in the presence of an attacking reagent. A familiar example is the reaction of cyanide ion with carbonyl compounds, which can be shown as:

In this reaction, the electron-pair in the double bond is displaced towards the more electronegative oxygen atom (compared to carbon) as the reagent cyanide ion attacks the carbon of the C=O group. This effect is temporary and disappears when the attacking reagent is removed from the reaction site.


Effects that result from the size of substituents and the repulsion between them are called steric effects. Steric effects are space-filling effects. Steric hindrance is a consequence of repulsion between the electrons in all the filled orbitals of the alkyl substituents that refer to bulky groups at the site of the reaction, which makes it difficult for the reactant to approach each other, decreasing reactivity. Since a proton occupies a very small volume, steric effects do not play a significant role in proton transfer reactions. However, steric effects may indirectly influence acid or base strength by interfering in resonance stabilization. Since all the atoms of a delocalized system must lie in a plane, any steric factor that prevents atoms from assuming coplanarity reduces resonance and, therefore, destabilizes the whole system. For example, steric hindrance increases the base strength of N,N-dimethyl-o-toluidine (pKa 5.9) relative to N,N-dimethylaniline (pKa 5.1). The ortho-methyl group in N,N-dimethyl-o-toluidine compresses one of the methyl groups attached to the nitrogen in the conformation needed for maximum delocalization of the unshared pair of electrons on nitrogen with the aromatic ring; thus, coplanarity of the ring and the dimethylamino substituent is prevented.

We know that a trans-substituted alkene is more stable than the cis-isomer (energy difference about 1 Kcal/mol). The cis-2-butene has the two methyl groups too close to each other (steric hindrance) as compared to the trans-form in which the methyl groups are far apart. cis-Alkenes having larger groups than methyl have considerable repulsive interaction so as to destabilize the molecules in the trans-isomer.

The reactivity of a compound is dependent on the size or bulk of the groups present in the reactions. A representative example is the reaction between carboxylic acids and alcohols in the presence of a mineral acid to give esters and water, called esterifìcation.

It is clear that as the alkyl group (R’) in alcohol becomes more bulky (more methyl groups on carbon are attached to the hydroxy group), the approach of the attacking molecule becomes more difficult resulting in a decrease in the rate of the reaction. The rate of esterification also becomes slower as the R group in the carboxylic acid becomes more bulky. Thus, a decreasing rate of esterification is observed in the reaction of ethanol with the following carboxylic acids:

The reactivity depends on the size of the groups attached to the carbonyl carbon. When addition takes place, the substituent groups are pushed back closer to one another. Hence, with an increase in the bulk of the groups, a decrease in the reactivity is observed.

The relative rate of nitration at ortho-positon in toluene and tertiary butyl benzene are about 8:1. This decrease in reactivity is due to the steric effect (tertiary butyl group is more bulky than methyl).

In general, attack at ortho position will be hindered if a large sized substituent is present on the benzene ring. For the same substituent, the rate depends on the size of the substituting agent. The chlorination of tertiary butyl benzene occurs slowly at ortho-position; no bromination takes place. The shape of proteins is not a perfect α-helix. The steric effect of substituent groups influences the position of folds of peptide-linkage, causing distortion.


The strength of the hydrogen bonds is about 3–10 Kcal/mol. It is a much weaker bond as compared to a covalent bond (25-120 Kcal/mol). Two types of H-bonding have been recognized, namely, intramolecular (within the same molecule) and intermolecular (between two or more molecules). Both types markedly influence the properties of a molecule.

The interactions described so far are not limited to molecules of any specific composition. However, there is one important intermolecular interaction specific to molecules containing an oxygen, nitrogen, or fluorine atom that is attached to a hydrogen atom. This interaction is the hydrogen bond, an interaction of the form A–HB, where A and B are atoms of any of the three elements mentioned above, and the hydrogen atom lies on a straight line between the nuclei of A and B. A hydrogen bond is about 10 times as strong as the other interactions e.g., ion-dipole and dipole–dipole and, when present, it dominates all other types of intermolecular interaction. A hydrogen bond also differs from other inter-molecular attractions because it is localized between a hydrogen atom and a specific pair of electrons. Other intermolecular forces act between molecules as a whole rather than on specific atoms on adjacent molecules. Hydrogen bonding is responsible for the existence of water as a liquid at normal temperatures; because of its low molar mass, water would be expected to be a gas. The hydrogen bond is also responsible for the existence as solids of many organic molecules containing hydroxyl groups (−OH); the sugars glucose and sucrose are examples.

Many interpretations of the hydrogen bond have been proposed. One that fits into the general scheme of this article is to think of the A-H unit as being composed of an A atomic orbital and a hydrogen 1s orbital and to consider a lone pair of electrons on B as occupying a p orbital. When the three atoms are aligned, these three orbitals can form three molecular orbitals: one bonding, one largely nonbonding and one antibonding. There are four electrons to accommodate (two from the original A-H bond and two from the lone pair). They occupy the bonding and nonbonding orbitals, leaving the antibonding orbital vacant. Hence, the net effect is to lower the energy of the AHB grouping and thus to constitute an intermolecular bond. Once again, on encountering the hydrogen bond, one encounters a twist in the conventional attitude; the question raised by this interpretation is not why such a bond occurs but why it does not occur more generally. The explanation lies in the small size of the hydrogen atom, which enables the balance of energies in the molecular orbital scheme to be favourable to bonding.

Hydrogen bonding occurs to atoms other than nitrogen, oxygen and fluorine if they carry a negative charge and hence are rich in readily available electrons. Thus, hydrogen bonding is one of the principal mechanisms of hydration of anions in aqueous solution (the bonding of H2O molecules to the solute species) and contributes to the ability of water to act as a good solvent for ionic compounds. It also contributes to the hydration of organic compounds containing oxygen or nitrogen atoms and thus accounts for the much greater aqueous solubility of alcohols than hydrocarbons.

Hydrogen bonds are of great significance in determining the structure of biologically significant compounds, most notably proteins and deoxyribonucleic acid (DNA). An important feature of the structure of proteins (which are polypeptides or polymers formed from amino acids) is the existence of the peptide link, the group -CO-NH-, which appears between each pair of adjacent amino acids. This link provides an NH group that can form a hydrogen bond to a suitable acceptor atom and an oxygen atom, which can act as a suitable receptor. Therefore, a peptide link provides the two essential ingredients of a hydrogen bond. All polypeptides have one structure or the other and often have alternating regions of each. Since the properties and behaviour of an enzyme molecule (a particular class of polypeptides) are determined by its shape and, in particular, by the shape of the region where the molecule it acts on needs to attach, it follows that hydrogen bonds are centrally important to the functions of life. Hydrogen bonds are also responsible for the transmission of genetic information from one generation to another, for they are responsible for the specific keying together of cytosine with guanine and thymine with adenine moieties, which characterize the structure of the DNA double helix.

Hydrogen bonds are weak bonds formed from electron-rich atoms such as O, N or F to hydrogen atoms, also attached by normal bonds to the same sort of atoms.

Dimethyl ether and ethanol are isomers having the same number of carbon, hydrogen and oxygen atoms. The former is a gas at room temperature, whereas the latter is a liquid, bp 78°.

This difference in the physical state of the two compounds can be explained due to the association of several molecules of ethanol with each other. These are held together by hydrogen bonds, formed by the bonding of the hydrogen atom of one molecule with the oxygen atom of another. These hydrogen bonds are represented by broken or dotted lines, as in the following aggregation of several ethanol molecules:

Additional thermal energy is required to break such association (hydrogen bonds) for vapourization. Hence, those compounds that are capable of forming hydrogen bonds have a high boiling point. In these compounds, the relatively positive hydrogen of one molecule is attracted towards a relatively negative atom of another (in the case of ethanol this is the oxygen atom), resulting in aggregation of several molecules. Hydrogen bonds exist between the fairly acidic hydrogen atom of one molecule and an electronegative atom such as oxygen, nitrogen, fluorine, etc. Dimethyl ether does not have fairly acidic hydrogen, which can form a hydrogen bond with the oxygen atom of another molecule, and hence no association is possible. This explains the volatility (lower bp) of dimethyl ether as compared to ethanol. Liquid hydrogen fluoride is polymeric due to the presence of hydrogen bonds, as shown below:

Aliphatic amines, having lower molecular weights, are soluble in water to an appreciable degree. This property can be explained as being due to the formation of a hydrogen bond between an amine molecule and water as shown below:

Carboxylic acids have higher boiling points in comparison to esters although both are isomers.

  CH3COOH (b.p. 118°C) HCOOCH3 (b.p. 32°C)
  acetic acid methyl formate

Once again, esters do not have acidic hydrogen capable of forming hydrogen bonds, but carboxylic acids do. In the solid and liquid state carboxylic acids exist largely as cyclic dimers such as:

Amide molecules also form an aggregate in which the hydrogen atoms attached to nitrogen form hydrogen bonds with oxygen atoms of the carbonyl group. Such hydrogen bonds hold the peptide chain of a protein molecule coiled to form an alpha helix structure.

In the above example, the hydrogen bonds exist between one molecule and another and thus are intermolecular in nature. In many compounds, hydrogen bonds exist within the same molecule–intramolecular bonding. The physical properties and chemical reactivity in many cases can be explained due to the presence of intramolecular hydrogen bonding. A few examples follow:

Salicylic acid (o-hydroxy benzoic acid, pKa 3.0) is about eighteen times more acidic than benzoic acid, whereas p-hydroxy benzoic acid (pKa 4.6) is half as acidic as benzoic acid. This difference in acidity can be explained as the result of the salicylate anion is being stabilized by intramolecular hydrogen bonding. The hydrogen atom of the OH group is sufficiently close to the oxygen atom of the salicylate ion for it to interact electrostatically with the formation of a six-membered ring. This is also sometimes referred to as Chelation. No similar stabilization is possible for the benzoate anion.

Similarly maleic acid has a lower pKa value (1.92) for the first ionization constant than the trans-isomer, fumeric acid (pKa 3.0), because the electrostatic effect of one carbonyl dipole facilitates the departure of the first proton from the other carboxyl group that is in its proximity; this is coupled with the fact that the singly charged cis anion is stabilized by intramolecular hydrogen bonding.

ortho-Nitrophenol is more volatile and is less soluble in water than the para isomer. In the former, hydrogen of the phenolic group lies close to the oxygen of the nitro group, forming an intramolecular hydrogen bond, which is not possible in para-nitrophenol due to the large separation of phenolic and nitro groups. However, intermolecular hydrogen bonds can exist between two or more molecules of the para-isomer. Due to intramolecular hydrogen bonding, ortho-nitrophenol cannot form a hydrogen bond with water and hence is less soluble. The para-isomer, having intermolecular hydrogen bonds, which break on dilution, is capable of forming hydrogen bonds with water and hence is more soluble.

Polyhydroxy alcohols provide more than one site per molecule for H-bonding. The simplest diol, 1,2-ethanediol boils at l97°C. The lower diols are miscible with water. In the two examples given earlier, the intramolecular hydrogen bond formation leads to a six-membere ring, which is favoured amongst the cyclic system.

In many 2-substituted heterocyclic acids the H-bond has the opposite effect, that is, it reduces the strength of the acid. The H-bonding stabilizes the acid and decreases the dissociation constant.

The anions are stabilized in solution due to the formation of hydrogen bond with solvent molecules. Stabilization by solvent attraction is termed solvation. The hydrogen bonding solvents are called protic, such as water, alcohols and carboxylic acids. Those solvents that do not have hydrogens capable of forming hydrogen bonds are called aprotic, for example hydrocarbons, ether, acetone, acetonitrile, etc.

The existence of H-bonds has been established both by physical and chemical means. In NMR, an H-bonded OH group shows a downfield shift of its proton. IR spectroscopy has played a key role in the study of H-bonding. In dilute solution in a non-polar solvent like CCl4 (or in the gas phase) where association between molecules is minimal, EtOH shows an O–H stretching band at 3640 cm−1. As the concentration of EtOH is increased, a broader band at 3350 cm−1 gradually replaces this band. The bonding of hydrogen to the second oxygen weakens the O–H bond, and lowers the energy and hence the frequency of the vibration.

It has been observed that H-bonds exist in carbanions, between OH and π electrons, as in allylic protons, etc. Hydrogen bonds occupy a significant place in biologically important molecules such as nucleic acids and proteins. Living cells in plants and animals retain a large quantity of water and most of it is attached to proteins by hydrogen bonds.


The major problem in studying organic chemistry is how to bring some sort of order and system to the vast body of compounds, and their many reactions, that go to make up the subject. One useful generalization is that of functional groups—all compounds that contain a particular group (e.g. NH2) can be expected to have at least some chemical behaviour in common.

An analysis of the enormous range of organic reactions establishes that there are effectively only four different types of reactions: substitution, addition, elimination and rearrangement. The vital feature of these reactions is the breaking of existing bonds between atoms and the formation of new, different bonds. The shifting of electron-pairs from one position to another—which is what bond-breaking/bond-making entails—may be emphasized by using curly arrows to represent these processes.

A further helpful generalization is that there are substantially only three types of reagents that attack centres (very often carbon atoms) in organic compounds: nucleophiles (electron-rich reagents), electrophiles (electron-deficient reagents) and radicals (reagents having an unpaired electron in their outer shell). Finally, there are essentially only two effects that the rest of an organic molecule can exert on the behaviour of the particular bond that is undergoing attack: electronic effects and steric effects. Having established these important generalizations, we will now use them to review the broad spectrum of organic reactions in a systematic way.

Structural changes can bring about marked changes in acidic and basic behaviour. These arise due to inductive, resonance and steric effects and also because of hydrogen bonding. Inductive effects operate through σ bonds. The effects decrease rapidly as the distance of the acting group increases from the reaction site. A steric effect is operative due to the large size or bulk of the group. In this situation the attack on the reaction site is prevented. The steric requirement of a proton is usually negligible due to its small size. Resonance operates in conjugated systems. This can be seen in the reorganization of non-bonding and/or π-electron densities and by drawing arrows involving lone pair of electrons, negative charge or electrons in π bonds. Such resonance structures are very often used to represent the real molecule — a hybrid of the contributing resonance structures. The extent of stabilization of a molecule or species by resonance can be estimated in terms of resonance energy, which can be determined by physico-chemical methods.

  1. Draw resonance structures for the following species and indicate the most stable contributor in each case:
    1. CH2N2
    2. CO32−
    3. N2O
  2. By adding an electrophile E+ to (a) phenol and (b) nitrobenzene and writing resonance structures of the cations, show the ortho and para-directing nature of the OH group and the meta-directing nature of the NO2 group.
  3. 2, 3-Dimethyl-2-cyclohexenone and 2, 3-dimethyl-3-cyclohexenone are more readily interconvertible than 1, 2-dimethylcyclohexene and 2, 3-dimethylcyclohexene. Explain.
  4. Draw resonance structure of:
    1. chlorobenzene
    2. acetonitrile
    3. pyrrole
    4. allyl cation
    5. benzyl cation
    6. diphenylmethyl radical
    7. p-methoxy nitrobenzene
  5. Which compound in each of the following pairs would be more extensively enolized?
  6. Are the following pairs of structures resonance contributors or different compounds?
  7. Draw each structure below to show its internal hydrogen bond.
  8. Write the structure of a proton tautomer of each of the following:
    1. CH3NO2
  9. Offer explanations for the following:
    1. Compound I has the greater dipole moment than compound II.
    2. Which compound has the greater electron density on its nitrogen atom?
    3. Compound I is more stable than II.
    4. A methyl group bonded to benzene loses proton more readily than a methyl group bonded to cyclohexane.
    5. The direction of the dipole moment in fulvene and calcene.
    6. Pyridine is more basic than pyrrole.
    7. Tropone is more stable than 1,3,5-cycloheptatriene.
    8. Cyclobutadiene is unstable.
    9. Cyclopropenone is a stable compound while cyclopentadiene has not been prepared.
    10. Cyclooctatetraene reacts with two moles of potassium to yield a stable compound whose NMR spectrum indicates a single line. (k) n-Butanol has a much higher boiling point than its isomers isobutanol and diethyl ether.
    11. The second ionization of fumaric acid is 24 times less than the first ionization, whereas it is 24,000 times less in the case of maleic acid.
    12. Compound I is more stable than II.
    13. Why does 15% of acetylacetone exist as the enol tautomer in water and as 92% in hexane?
  1. Which of the following compounds would you expect to be most soluble in water?
    1. C2H5.O.C2H5
    2. C6H5Cl
    3. C6H5.NH2
    4. C6H5.CO.OC2H5
    5. CH3.CO2H
  2. Which of the following groups has an electron-withdrawing mesomeric effect?
    1. −CH2.CH3
    2. −Cl
    3. −OCH3
    4. −CN
    5. −NH.CH3
  3. Which compound is likely to have the highest boiling point?
    1. CH3.CH2.CH2.OH
    2. CH3.CO.CH2.CH3.
    3. CH3.O.CH2.CH3.
    4. CH3.CH2.CH2.SH
    5. CH3.CH2.CHO
  4. Which of the following would not have a dipole moment?

    (i) CCl4    (ii) CH2Cl2    (iii) trans-1,2-dichloroethene    (iv) ClC ≡ CCl

    1. (i) and (ii)
    2. (i), (ii) and (iii)
    3. (i), (iii) and (iv)
    4. (iii) and (iv)
  5. One or more of the following compounds are in their most stable tautomeric forms. Which are they?
  6. Which of the following series of canonical forms is incorrect?
  7. Which of the following compounds cannot exhibit keto-enol tautomerism?
  8. The inductive effects of the groups: −CH3, −CO2 , −Br, −NH3+ are respectively
    1. +I   −I   +I   +I
    2. +I   +I   −I   −I
    3. −I   −I   +I   +I
    4. −I   +I   −I   +I
  9. Indicate which of the following assertions correctly applies to the electromeric effect.
    1. It can either facilitate or hinder attack on a molecule by a reagent.
    2. It results in the presence of an electronegative group within a molecule.
    3. It results from the polarizability of systems containing π-orbitals.
    4. All molecules showing it have dipole moments.
    5. It makes paraffin hydrocarbons (alkanes) susceptible to attack by electrophiles.
  10. In benzene the overlapping between carbon–carbon is of the type
    1. p-p
    2. sp3-sp3
    3. sp2-sp2
    4. sp-sp
  11. In tautomerism
    1. a proton is moved around the molecule
    2. electrons are moved around the molecule
    3. no actual movement occurs
    4. shift of double bond occurs
  12. Enolization is catalyzed by
    1. acids only
    2. bases only
    3. acids and bases both