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Cycloadditions in weakly and highly organized aqueous mediaRispens, Taede

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Chapter 2

Solvent Effects on 1,3-Dipolar

Cycloadditions of Benzonitrile Oxide

The kinetics of 1,3-dipolar cycloadditions of benzonitrile oxide with a series of

N-substituted maleimides and with cyclopentene are reported for a wide range

of solvents and binary solvent mixtures. The results indicate the importance of

both solvent polarity and specific hydrogen-bond interactions in governing the rates of

the reactions.

The aforementioned reactions are examples for which these factors counteract, lead-

ing to a complex dependence of rate constants on the nature of the solvent.

2.1 Introduction

2.1.1 1,3-Dipolar Cycloadditions

Besides Diels-Alder (DA) reactions, 1,3-dipolar cycloadditions (DC) have been exten-

sively studied and recognized as an important class of cycloadditions.1,2 In recent years,

mechanistic studies of DC reactions regained interest, both experimentally3,4 and theo-

retically.5–9

DC reactions involve the cycloaddition of an unsaturated compound to a 1,3-dipole,

yielding a five-membered ring. The dipoles formally contain a positive and a nega-

tive charge, and are best represented by the resonance structures shown in Scheme 2.1.

Dipole moments range from 0.17 Debye for nitrous oxide to 3.6 Debye for N -methyl-

C-phenylnitrone (I and II in Scheme 2.1, resp.), and are the result of the partial can-

cellation of the dipole moments of two dominant resonance structures. Azides and

diazoalkanes have much smaller dipole moments than nitrones and nitrile oxides.

From a mechanistic point of view, DA and DC reactions are very similar.10–12 Thus,

formation of two new bonds is a concerted process; Woodward-Hoffmann selection rules

43

Chapter 2

a b c a b c ab

c ab

c

N N O N N O

CN

O

CH3

C6H5

HC

NO

CH3

C6H5

H

µ (Debye)

I

II

0.17

3.55

SCHEME 2.1.

apply; the reaction is characterized by an early transition state, and a large negative

volume of activation. Reactivity is for a large part governed by the highest occupied

molecular orbital (HOMO) and lowest unoccupied molecular orbital (LUMO) of the 1,3-

dipole and the dipolarophile. Frontier molecular orbital (FMO) theory is therefore often

used to rationalize trends in reactivity (Section 2.1.2).13,14 Depending on the nature of

the dipole and the dipolarophile, a reaction may be classified as having either normal

electron demand (NED) or inverse electron demand (IED) (Figure 2.1). In contrast to

DA reactions, where only few examples of a change from NED to IED are known15 in a

series of related reactions, there exist numerous examples of DC reactions showing this

behavior.1

DA reactions are often described as relatively insensitive towards the nature of the

solvent, but rate enhancements of the order of 102–103 upon going from n-hexane to

water are nevertheless commonly found. Rate constants of DC reactions are even less

dependent on the the solvent.1,11, 12, 16 Even rate constants in water hardly differ from

those in other solvents, and when accelerations are observed,3,10, 17 they are modest,

compared to those for DA reactions. Notably, DC reactions sometimes show a reverse

dependence of rate constant on the polarity of the medium, the reaction being slowed

in polar media.1,3, 4, 18 Intermediate cases are also known. For example, for the reaction

between phenyl diazomethane and norbornene rate constants vary only by a factor of

1.8 over a wide range of solvents (water not included).1 The inverse dependence of

rate constant on the polarity of the solvent is most pronounced for nitrones and nitrile

oxides as dipolarophiles. These compounds possess relatively high dipole moments,

which are (partially) lost in the course of the reaction. The latter accounts for the

inverse dependence. Rate constants plotted against ET(30) (a measure of the solvent

polarity, see below) usually show a fair correlation. However, rate constants in protic

solvents sometimes deviate from this trend.1 Lewis acids are also known to sometimes

44

1,3-Dipolar Cycloadditions

HOMO

LUMO

HOMO

LUMO

LUMO

HOMO

electron-poordienophile/dipolarophile

diene/1,3-dipoleelectron-richdienophile/dipolarophile

E

FIGURE 2.1: Schematic representation of the dominant MO interactions in a cy-

cloaddition reaction. Left: Normal electron demand cycloaddition. The dashed lines

indicate the situation where the dienophile/dipolarophile is activated by a Lewis acid.

Right: Inverse electron demand cycloaddition.

induce either accelerations or inhibitions of DC reactions19–21

Before introducing DC reactions with benzonitrile oxide (Section 2.1.4), two topics

need a brief introduction: FMO theory, used to rationalize ‘catalytic’ effects of hydro-

gen bonds (Section 2.1.2), and the ET(30) solvatochromic scale, used to correlate rate

constants with ‘solvent polarity’ (Section 2.1.3).

2.1.2 Frontier Molecular Orbital Theory

Frontier molecular orbital (FMO) theory is a simple model which can be used to predict

trends in reactivity.22,23 Despite its simplicity, the theory accounts for many observa-

tions. The theory is based on the notion that the FMOs of molecules contribute most

to the interaction energy, and that for reactions with an early transition state, as is the

case for many pericyclic additions, the FMOs of the reactants are a reasonable starting

point for approximating those of the transition state. Then, in a series of reactions,

the combination of reactants with the most favorable FMO interactions reacts fastest,

or leads to the dominant stereoisomer. Figure 2.1 shows how trends in the rates of

cycloadditions can be rationalized. The interactions between the HOMO and LUMO

of both reactants are of prime importance, and their relative energies determine which

interaction is dominant. For a typical DA reaction with an electron-poor dienophile,

the FMOs of the diene will be higher in energy than those of the dienophile, and the

dominant interaction will be HOMOdiene–LUMOdienophile (NED). It is proposed that the

standard Gibbs energy of activation correlates with the energy gap between the (dom-

inant) interacting HOMO and LUMO. If the double bond of the dienophile becomes

45

Chapter 2

FIGURE 2.2: Left : log(k2) vs 1/(EHO − ELU) for the DA reactions of cyclopen-

tadiene and 9,10-dimethylanthracene with some cyano-substituted ethylenes. Nice

correlations are found for these series. Right : A similar plot for the DA reaction of

a series of substituted dienes with tetracyanoethylene. In this case, merely scatter is

observed. (1 eV = 96.4 kJ mol−1. Taken from Ref. 24.)

more electron-poor, its FMOs are lowered in energy, the gap between HOMOdiene and

LUMOdienophile becomes smaller and the rate will be higher. Hence, the catalytic action

of a Lewis acid complexing to a dienophile is accounted for. Hydrogen bonding to a

carbonyl group next to a double bond may also lead to a decrease in the electron density

in the double bond. This hydrogen bonding is regarded as one of the main factors that

cause the large increase in rate of DA reactions in water. Due to its small size, water

can form multiple hydrogen bonds (involving more than one water molecule) with the

carbonyl group.

Within the framework of FMO theory, trends in reactivity (and selectivity) are

explained only on the basis of the interactions between the FMOs.25 In general, this is

a too restricted description; Figure 2.2. In particular, ring strain and steric interactions

may also seriously affect the rate.

FMO theory is often invoked to explain trends in regioselectivity and stereoselectivity

as well.14,23 However, for DC reactions, its predictive power is limited, as has been

discussed in the literature for several decades.7–9,14, 26–28 A nice overview of the predictive

power of FMO theory under various circumstances is given in Ref. 23.

2.1.3 Solvent Polarity and ET(30)-values

The polarity of a solvent is an intuitive and ill-defined property. Here, we regard the

polarity of a solvent as the capacity to solvate/solubilize molecules by means of different

46

1,3-Dipolar Cycloadditions

types of polar interactions, e.g. dipole-dipole, charge-dipole, hydrogen-bonding. Hydro-

gen bonding is a special case. For the reactions under study, hydrogen-bond donating

solvent molecules may have a specific effect; i.e. acting as a Lewis acid, catalyzing the

reaction by altering the FMOs (Section 2.1.2). Furthermore, there will be a non-specific

effect, i.e. the ability to stabilize polar molecules, or a polar activated complex, because

of the polar character of the hydrogen-bonding component.

One method for determining solvent polarity makes use of a solvatochromic dye, and

takes changes in the UV-Vis spectra (often the shift of a specific transition band) as

a measure of the polarity. For this purpose, betaine-30 (Reichardt’s dye) is commonly

used.29,30 The measure of the polarity in this case is the excitation energy corresponding

to the lowest-energy absorption band, expressed in kcal mol−1 and denoted by ET(30).

Betaine-30, bearing both a positive and a negative charge (Scheme 2.2), is sensitive

to different types of polar interactions. It is a relatively weak ‘hydrogen-bond donor’

(electron-pair acceptor), but a strong hydrogen-bond acceptor. The transition is illus-

trated in Scheme 2.2. In the excited state, electrons are redistributed such that the large

ground-state dipole moment is strongly reduced (µ changes from 14.7 to 6.0 Debye).31

Therefore, the excited state is much less polar than the ground state and hardly sus-

ceptible to polar interactions. Hence, the absorption maximum is very sensitive to the

polarity of the solvent.

Complications may arise in mixtures of solvents, if the betaine-30 is preferentially

solvated by one of the components. Now any property correlating with ET(30) (e.g.

log(k)) will only do so for these mixtures insofar preferential solvation to similar extents

occurs (e.g. of the reactants). This condition is of particular importance if one of the

solvents is capable of forming hydrogen bonds with the negatively charged oxygen.

N

O

N

O

hν

Betaine-30

SCHEME 2.2.

47

Chapter 2

HOMO

LUMO

HOMO

LUMO

LUMO

HOMO

electron-poordipolarophile

benzonitrileoxide

electron-richdipolarophileE (eV)

0

-10

-1/-0.5

-10/-9.4

2

-9

0

-11

FIGURE 2.3: Schematic representation of the FMOs for the reaction of benzonitrile

oxide with an electron-poor dipolarophile (left) and an electron-rich dipolarophile

(right). Solid lines depict the situation in e.g. n-hexane, dashed lines that in a protic

solvent. Data taken from Ref. 14, (italic from Ref. 26), HOMO/LUMO of 1 in protic

solvent based on data from Ref. 6.

2.1.4 1,3-Dipolar Cycloadditions of Benzonitrile Oxide

Benzonitrile oxide (1, Scheme 2.3), a very reactive 1,3-dipole, was first prepared in 1886

by Gabriel and Koppe,32 and then again in 1894 by Werner and Buss.33 Benzonitrile

oxide is often generated in situ, because it dimerizes quickly. In fact, it dimerizes

in solution so easily, that its reactivity was at first not recognized,34 but cycloadditions

proceed smoothly even with completely unactivated dipolarophiles such as ethene under

ordinary laboratory conditions.35 Cycloadditions with 1 were explored in the fifties

(overview given in Refs. 36); early mechanistic studies include Ref. 28 and 37.

Solvent effects on DC reactions with 1 and derivatives are remarkably small.1 Stud-

ies of solvent effects have often been brief and inconclusive about the different factors,

which control rates.17,38, 39 In a detailed kinetic study, the DC reactions between 1 and

several electron-rich and electron-poor dipolarophiles have been studied for a number

of solvents, including water, as well as for mixtures of ethanol and water.18 The dipo-

larophiles include cyclopentene, methyl vinyl ketone and N -methylmaleimide. Whereas

reactions involving an electron-rich dipolarophile are still 3–10 times faster in water than

in most organic solvents, reactions involving electron-poor dipolarophiles are slightly de-

cellerated. This difference was rationalized on the basis of FMO theory (Figure 2.3); 1

is a good hydrogen-bond acceptor, and its FMOs are lowered in energy when dissolved

48

1,3-Dipolar Cycloadditions

in a protic solvent. When reacting with the electron-rich dipolarophile cyclopentene,

the dominating interaction is LUMO1–HOMOcyclopentene. Consequently, the energy gap,

and hence the Gibbs energy of activation, is smaller in a protic solvent. In the case of

an electron-poor dipolarophile, the dominating interaction is LUMOdipolarophile–HOMO1.

FMO energies of both reactants are lowered in a protic solvent (the electron-poor dipo-

larophiles studied are also susceptible to hydrogen bond formation), but it was proposed

that this occurs more efficiently for 1, leading to a rate retardation.40 However, this ex-

planation does not account for the complicated dependence of the rate constants on

the solvent; for instance, k(n-hexane) ≈ k(ethanol) ≈ k(water) > k(1,4-dioxane) >

k(dichloromethane) ≈ k(2,2,2-trifluoroethanol) for the reaction of 1 with 2a. The fact

that, for electron-poor dipolarophiles, rate constants in n-hexane, ethanol and water are

nearly equal, contradicts the explanation of the lowering of the rate constant due to the

hydrogen-bond interactions with 1. A larger destabilization of the hydrophobic initial

state, relative to the less hydrophobic transition state (enforced hydrophobic interac-

tions), may explain why in water the rate is not much lower than in organic solvents,

but for ethanol, such a counteracting effect on the rate is not present. In summary, hy-

drogen bonding and hydrophobic interactions are important factors that influence rate

constants in water, but in general, solvent effects on DC reactions of benzonitrile oxide

(1) are only partially understood.

2.1.5 Rate Constants in Solvent Mixtures and Organized Reaction

Media

Solvent polarity, hydrogen-bond donating capacity, and differences in hydrophobic hy-

dration also affect rate constants in more complex media, in particular within micelles

and in water/alcohol mixtures. In the latter case a complicated dependence of rate

constant on the composition of the mixture in the water-rich regime is found for many

cycloadditions; Chapter 3. In the case of either a micelle or a vesicle, it is difficult to

unravel the different contributions to the rate constant, that may include, besides those

already mentioned, a (partial) mismatch in binding sites, and the asymmetric charge

distribution at the surface of charged micelles. Comparison with reactions in aqueous

salt solutions, solvent mixtures, or even a combination of these, may lead to a better

understanding of the reaction in the micellar (pseudo-)phase. A clear understanding of

what determines the rate constant in water/cosolvent mixtures is therefore a prerequisite

towards understanding reactivity in micellar solutions.

On the other hand, by varying their mole fractions, properties of individual solvents

can be gradually ‘mixed in’, and therefore effectively ‘varied’, in a solvent mixture,

allowing a detailed study of the influence, which various solvent properties can have

on reaction rates. Therefore, results from kinetic experiments in both aqueous and

49

Chapter 2

N

O

O

R

1 2 3

R = Et

n-Bu

(2a)

(2b)

+

CHNOH

C N OBleach

C N OON

NO

OR

(2c)CH2Ph

4 5

+C N OON

SCHEME 2.3.

non-aqueous solvent mixtures is presented in this chapter.

Rate constants for the reaction of 1 with N -methylmaleimide in several solvents

have been previously determined.18 The complicated results prompted a more detailed

study. In this chapter, an extensive study is presented of the influence of the medium

(pure solvents and mixtures of solvents) on rate constants of 1,3-dipolar cycloadditions of

benzonitrile oxide (1) with N -alkyl substituted maleimides (2a–c) and with cyclopentene

(4) (Scheme 2.3). Emphasis is placed on the complex interplay of different factors

that control the rate, in particular hydrogen bonding and polarity. In this regard, the

reactions of 1 with 2a–c are of particular interest, because both substrates are susceptible

to hydrogen-bond formation. The reaction of 1 with 4 is a nice reference, because 4

does not form significant hydrogen bonds.

2.2 Results and Discussion

2.2.1 The Solvent Dependence of the Rate Constant

In Figure 2.4, log(k2) is plotted against ET(30) for a wide range of solvents for the

reaction of 1 with 2a and with 4. The solvents roughly form two groups: protic and

aprotic solvents.

50

1,3-Dipolar Cycloadditions

30 35 40 45 50 55 60 65

-1.4

-1.2

-1.0

-0.8

-0.6

-0.4

-0.2

1-BuOH 2-BuOH

t -BuOH H

2 O

TFE

MeOH EtOH

1-PrOH

2-PrOH

CH 3 CN

CH 2 Cl

2

CHCl 3

THF

1,4-dioxane

CCl 4

n -hexane

log(

k 2 )

E T (30)

30 35 40 45 50 55 60 65

-1.0

-0.8

-0.6

-0.4

-0.2

0.0

n -hexane

CCl 4

1,4-dioxane

2-BuOH

2-PrOH

1-PrOH

EtOH MeOH

CH 3 CN CHCl

3

CH 2 Cl

2

TFE

H 2 O

2 +

log(

k 2 )

E T (30)

FIGURE 2.4: Values of log(k2) of the reaction of 2a (left) and of 4 (right) with 1

vs the ET(30) values29 of various solvents41 at 25 ◦C.42 (closed circles are values for

N -methyl- rather than N -ethylmaleimide.18) k2 has units M−1 s−1.

First, when considering the group of apolar solvents (with values of ET(30) below ca.

40), log(k2) decreases roughly linearly with ET(30). This pattern is indicative of a polar

initial state (1) and a less polar transition state, in which the charge separation of the

1,3-dipole has partly disappeared. Note that DA reactions are almost invariably faster in

a more polar solvent. In the activated complex, some charge separation developes, that

is stablilized by polar interactions. One reason why rates of DC reactions (proceeding

via an analogous mechanism) are so weakly dependent on the solvent, is that this charge

separation is mediated by the partial disappearance of the 1,3-dipole, leading to only

small accelerations, or even — as for the reactions of 1 with 2a and 4 — to a decrease

in rate upon going to a more polar solvent.

For solvents, where ET(30) > 40, the kinetics of the reactions is complicated. A

comparison of the reactions between 1 and 2a and between 1 and 4 sheds some light on

this phenomenon. The latter reaction is classified as IED. Therefore, hydrogen bonding

to 1 is favorable (Section 2.1.2, Figure 2.3), Furthermore, 4 is not capable of forming

hydrogen bonds. Compared to the aprotic solvents, alcohols show larger rate constants

for the reaction of 1 with 4. The more polar the alcohol, the smaller the rate constant,

though (except for TFE). This trend may be the continuation of the effect of the polarity

(see below; Figure 2.6). Two solvents stand out. 2,2,2-Trifluoroethanol (TFE), a very

potent hydrogen-bond donor, causes a larger rate constant than the other alcohols. This

is also true for water. In addition, a further acceleration may be ascribed to enforced

hydrophobic interactions. Thus both solvent polarity and the hydrogen bond-donating

51

Chapter 2

TABLE 2.1: Rate constants for the reaction of 1 with 2a and 4 in various solventsa

at 25 ◦C.

Solvent ET(30)b kc2,2a kd

2,4 Solvent ET(30)b kc2,2a kd

2,4

n-hexane 30.9 0.330 0.333 2-BuOH 47.1 0.308 0.334

CCla4 32.5 0.210 0.255 2-PrOH 48.4 0.289 0.274

1,4-dioxane 36.0 0.12e 0.170 1-BuOH 49.7 0.310 n.d.

THF 37.4 0.100 n.d. 1-PrOH 50.7 0.320 0.259

chloroform 39.1 0.059 0.127 EtOH 51.9 0.22e 0.265

CH2Cl2 41.1 0.07e 0.120 MeOH 55.5 0.196 0.229

t-BuOH 43.3 0.254 n.d. TFE 59.8 0.061 0.380

CH3CN 45.6 0.10 0.124 water 61.3 0.350 0.978a) Abbreviations are explained in the notes.41 b) Values from 29. Units are kcal mol−1. c)

Units M−1 s−1. d) Units 10−2 M−1 s−1. e) Value for N -methylmaleimide.18

capacity of the solvent affect the reaction rates and, interestingly, in this particular case

in opposite directions. (The hydrogen bond-donating capacity of course also contributes

to the solvent polarity.) For most DA reactions, both factors increase the rate of the

reaction. For DC reactions, it appears that usually these two factors either enhance or

diminish the rate constants. The DC reaction of 4 with 1 is an example, where these

factors are opposed, and therefore generate a much more complex dependence of rate

constant on solvent.

For the reaction of 1 with 2a, hydrogen bonding is also possible for the dipolarophile,

which introduces further complexity. In the absence of hydrogen-bond donors, the same

dependence of rate constant on solvent polarity is found as for the reaction of 1 with 4.

Again, hydrogen-bond donating solvents produce an additional rate increase. But now,

the rate constant in TFE is much lower than that in other alcohols.

The two simplest explanations for this pattern are: (i) the reaction is still mainly

IED; hydrogen bonding occurs both to 1 and 2a, affecting 1 more than 2a. Only in

case of TFE, hydrogen bonding is more efficient to 2a; (ii) the reaction is mainly NED,

hydrogen bonding also occurs both to 1 and 2a, but affects 2a more than 1. Only in

case of TFE, hydrogen bonding to 1 supersedes hydrogen bonding to 2a. The kinetic

data can be understood using either explanation, but UV-VIS spectra of 1 in different

solvents support the latter explanation (Figure 2.5). The maximum in the absorption

band of 1 shifts from 252 nm in n-hexane to 248 nm in acetonitrile (cyanomethane),

to 250 nm in 1-propanol, 243 nm in water, and 240.5 nm in TFE, indicating both a

relatively weak interaction with 1-propanol and an efficient interaction with TFE (even

more efficient than with water). The energy of the transition (ET(1)) plotted against

ET(30) clarifies this pattern. For 1 + 4, the rate constant for TFE deviates positively

52

1,3-Dipolar Cycloadditions

220 240 260 280

0.0

0.1

0.2

0.3

0.4

0.5

30 40 50 60 112

114

116

118

120

E T

( 1 )(

kcal

/mol

)

water

TFE

1-propanol

acetonitrile

chloroform

n -hexane

E T (30) (kcal/mol)

n -hexane

chloroform 1-propanol

acetonitrile

water TFE

wavelength (nm)

A

FIGURE 2.5: UV spectra of 1 in various solvents at 25 ◦C. The inset shows the

corresponding transition energies ((ET(1) plotted against ET(30)).

from the trend found among the alcohols, whereas for 1 + 2a, a negative deviation is

found (Figure 2.6). This pattern is also in line with a relatively efficient binding of TFE

to 1, and supports the NED mechanism for the reaction of 1 with 2a.

A sharp deviation from a general trend in a plot of log(k) against solvent polarity

(Figure 2.6) may indicate a change in mechanism. In this case, one may perhaps regard

the introduction of hydrogen bonds as a change in mechanism. A similar (but reversed)

pattern was found for the reaction of phenyl azide with norbornene, both experimen-

tally10 and theoretically.5 Nevertheless, the deviation remains an unusual observation.

An alternative explanation, based on a change from a concerted mechanism to a mech-

anism involving a zwitterionic intermediate, may be rejected on several grounds: (i) the

dependence on the solvent polarity should be much larger for such a mechanism;16 (ii)

Hammett ρ values for the reaction of 1 with electron-poor styrenes (in CCl4)37 and with

acrylonitrile (in water)39 are small and similar, contradicting a (change to a) zwitteri-

onic mechanism; (iii) activation entropies are large and negative over the full range of

solvents (see below) and are characteristic of a concerted reaction mechanism.

The effect of solvent polarity is extrapolated from the range of solvents with ET(30)-

values between 30–40 (Figure 2.6), to illustrate the divergence from this trend for the

more polar solvents (ET(30)-values > 40). The ET(30)-scale is based on one parame-

ter, that includes hydrogen-bond donor capacity, hydrogen-bond acceptor capacity, and

53

Chapter 2

30 35 40 45 50 55 60 65

-3.0

-2.8

-2.6

-2.4

-2.2

-2.0

-1.8

-1.6

-1.4

-1.2

-1.0

-0.8

-0.6

-0.4

-0.2

30 35 40 45 50 55 60 65

-1.0

-0.5

0.0

2-BuOH TFE

MeOH

H 2 O

1-BuOH 2-BuOH

t -BuOH H 2 O

TFE

MeOH EtOH

1-PrOH

2-PrOH

CH 3 CN

CH 2 Cl

2

CHCl 3

THF

1,4-dioxane

CCl 4

n -hexane

log(

k 2 )

E T (30)

FIGURE 2.6: Data from Figure 2.4, with various trends indicated. Main plot: 1 + 2a;

inset: 1 + 4.

polarizablility/dipolar interactions. The contribution of the hydrogen-bond donating

capacity of the solvent in just providing a more polar reaction environment could be

smaller than indicated by ET(30), as betaine-30 is rather sensitive to these interactions.

In fact, the fair correlation of log(k) with ET(30), including alcohols, found for many

other DA and DC reactions, may be the result of solvent polarity (including hydrogen

bonds) playing a smaller role in determining the rate in alcohols than estimated with

ET(30), together with hydrogen bonds inducing (catalyze/inhibit) additional effects,

which work in the same direction as the solvent polarity. This is supported by many

other cases, in which a difference in slope is found among the alcohols, and is consis-

tent with the idea, that in general both non-specific (polarity) and specific (‘catalytic’

hydrogen bonds) solvation is important.

Acetonitrile is the odd one out, in particular for the reaction of 1 with 2a. A similar

pattern was found for a related reaction.17 A specific accelerating effect of aprotic dipolar

54

1,3-Dipolar Cycloadditions

TABLE 2.2: Isobaric activation parameters for 1 + 2a and 2b at 25 ◦C.

Solvent ∆‡G◦ (kJ mol−1)a ∆‡H◦ (kJ mol−1)b −T∆‡S◦ (kJ mol−1)b

1+2a

n-hexane 75.8 42.5 33.3

chloroform 80.0 51.8 28.2

1-propanol 75.9 42.7 33.1

trifluoroethanol 79.9 34.4 45.5

water 75.7 50.2 25.5

1+2b

n-hexane 75.5 40.8 34.7

1-propanol 75.2 40.5 34.8

trifluoroethanol 79.5 33.0 46.4

water 74.6 48.0 26.6a) Standard error < 0.1 kJ mol−1. b) Standard error 1.5–2 kJ mol−1 for 2a; 1 kJ mol−1 for1b.

solvents has been suggested,43 but the effect is far from general (in case of 1 + 4, the

rate constant for acetonitrile only slightly deviates. Examples where the effect is absent

are given in Ref. 1, 3, and 16).

Several multiparameter analyses have also been undertaken (Chapter 3), using the

Abraham-Kamlett-Taft model, extended with the solvent parameter Sp. These models

discern different aspects of polarity; e.g. the hydrogen-bond donor capacity (α). Usually,

decent correlations are found for DA reactions, but for the DC reactions descibed in this

chapter, no satisfactory fit was obtained.

Isobaric Activation Parameters

Isobaric activation parameters (∆‡G◦, ∆‡H◦, and ∆‡S◦) for the reactions of 1 with 2a

and 2b have been determined in different solvents (Table 2.2). Upon going from n-

hexane to chloroform (trichloromethane), the rate decreases, and the accompanying

increase in activation enthalpy is in line with a stabilization of the polar initial state

(with respect to the transition state) by the more polar chloroform. For the reaction

in 1-propanol, ∆‡H◦ and ∆‡S◦ are the same as those in n-hexane. When compared to

chloroform, the enthalpy of activation is decreased, as a result of hydrogen bonding (to

the dipolarophiles), in such a way that the FMO interaction energy is lowered. Note

that in both cases, the changes in ∆‡H◦ and ∆‡S◦ are more dramatic than the changes

in ∆‡G◦, but in large part compensating. This indicates that differences in solvation play

an important role, although the overall rate constant need not be affected to a large

extent, because of this compensating behavior.

55

Chapter 2

In TFE, the rate constants of these reactions are low, which has been explained

in terms of FMO theory.18 The activation parameters reveal that the decrease in rate

constant is entirely due to a less favorable entropy of activation, which seems hard to

reconcile with a larger difference in energies between the HOMO and LUMO of the

reactants.44 Instead, the high solvent polarity may be responsible for the low rate

constant, but this is expected to lead to an increase in ∆‡H◦ as well. Moreover, for

the reaction of 1 with 4 no corresponding decrease in rate constant is found. Yet

another explanation is that the activated complex is strongly solvated (more strongly

than the reactants are), but that relatively few of these configurations exist, leading to

a reduced chance of the formation of transition state structures. This explanation is

highly speculative, but in line with the large negative entropy of activation in TFE. Of

course, this effect ultimately may be present together with the lowering of the HOMO of

1. As mentioned, it is possible that the changes in desolvation cause (large) differences

in ∆‡H◦ and T∆‡S◦ that nearly compensate each other.

The activation parameters for the reactions in water are more in line with expecta-

tion, with relatively small negative entropies of activation and relatively large enthalpies

of activation. Water and aqueous mixtures are discussed in Section 2.2.3.

Activation parameters for the reactions of 1 with 2a and 2b follow the same pattern.

Any differences in solvation (hydration) due to the presence of a larger alkyl substituent

apparently do not affect the activation process. The n-butyl tail may still be too small

to fold back, or have any interaction with 1, and therefore acts as an inert group.

2.2.2 Mixtures of Two Solvents

To investigate further the complex kinetic behavior of these reactions, rate constants

for the reaction of 1 with 2a were determined in the solvent mixtures chloroform/CCl4,

chloroform/TFE, chloroform/1-propanol, and 1-propanol/TFE; Figures 2.7 and 2.8. A

comparison with ET(30)-values in these mixtures provides information concerning the

influence of the solvent polarity on the reaction. The ET(30)-values already offer clues

concerning the nature of the mixtures. In all cases, ET(30) shows a linear dependence

on the composition, except for a small range of composition. In case of chloroform/TFE,

the formation of strong hydrogen bonds between TFE and betaine-30 is definitely re-

sponsible for the sharp increase in ET(30) at low mole fractions of TFE. In mixtures of

acetonitrile with propanol46 or chloroform with ethanol47 the same pattern was observed,

with a dependence of ET(30) on the mole fraction resembling a binding curve at low

amounts of propanol or ethanol, and a more gradual, linear dependence once ‘satura-

tion’ has been reached. Strong hydrogen bonds between the alcohol and the negatively

charged oxygen of betaine-30 are responsible for this pattern. In the other mixtures, a

strong preference for one of the solvents is not observed because (i) neither of the two

56

1,3-Dipolar Cycloadditions

a)

0.0 0.2 0.4 0.6 0.8 1.0

-1.2

-1.0

-0.8

-0.6

34 36 38 40 42

-1.2

-1.0

-0.8

-0.6

log(

k )

E T (30) (kcal/mol)

0.0 0.2 0.4 0.6 0.8 1.0

34

36

38

40

42

E T

(30)

(kc

al/m

ol)

mole fraction CHCl 3

log(

k 2 )

mole fraction CHCl 3

b)

0.0 0.2 0.4 0.6 0.8 1.0

-1.2

-1.0

-0.8

-0.6

-0.4

42 44 46 48 50 52 -1.5

-1.0

-0.5

log( k 2 )

E T (30)

0.0 0.2 0.4 0.6 0.8 1.0

42

44

46

48

50

52

E T

(30)

(kc

al/m

ol)

log(

k 2 )

mole fraction 1-PrOH

c)

0.0 0.2 0.4 0.6 0.8 1.0

-1.5

-1.4

-1.3

-1.2

0.0 0.2 0.4 0.6 0.8 1.0 40

45

50

55

60

E T

(30)

(kc

al/m

ol)

log(

k 2 )

mole fraction TFE

FIGURE 2.7: Logarithms of the rate constants of the reaction of 2a with 1 in mixtures

of chloroform with a) CCl4, b) 1-propanol and c) TFE at 25 ◦C.45 Insets show the

corresponding ET(30)-values and log(k2) vs ET(30). k2 has units M−1 s−1.

57

Chapter 2

0.0 0.2 0.4 0.6 0.8 1.0

-1.2

-1.0

-0.8

-0.6

-0.4

50 52 54 56 58 60 62 -1.5

-1.0

-0.5

log( k 2 )

E T (30)

0.0 0.2 0.4 0.6 0.8 1.0

50

52

54

56

58

60

62

E T

(30)

(kc

al/m

ol)

mole fraction TFE

log(

k 2 )

FIGURE 2.8: Logarithms of the rate constants of the reaction of 2a with 1 in mixtures

of 1-propanol with TFE at 25 ◦C.45 Insets show the corresponding ET(30)-values and

log(k2) vs ET(30). k2 has units M−1 s−1.

solvents is a (strong) hydrogen bond donor, or (ii) both solvents are (strong) hydrogen

bond donors. The chloroform/CCl4 mixture is an interesting case, because chloroform

is not a strong hydrogen bond donor. Nevertheless, some kind of binding is observed.48

In comparison with betaine-30, 1 can be expected to form strong hydrogen bonds.

Also 2a is a good hydrogen bond acceptor, but probably to a lesser degree.

In chloroform/CCl4 mixtures, two linear relationships are observed in a plot of log(k2)

vs the composition, with slightly different slopes. No strong interactions (hydrogen

bonds) between solvents49 and reactants are anticipated, and the rate constant is gov-

erned primarily by the polarity of the medium. As reactants and betaine-30 are prefer-

entially solvated to different degrees, a correlation between log(k2) and ET(30) over the

full range of composition is not observed.

In chloroform/1-propanol mixtures, a good hydrogen bond donor is present (1-

propanol).49 The FMOs of both reactants are affected by the formation of hydrogen

bonds, and a net accelerating effect results. A plot of log(k2) vs composition reveals

that no strong binding of 1-propanol to either reactant occurs, because log(k2) grad-

ually increases towards the value of log(k2) in pure 1-propanol. The increased slope

for x1-PrOH > 0.3 may result from additional hydrogen bonding with 2a starting to

be involved in the activation process. Solvent polarity plays a role as well, but the

(additional) ‘catalytic’ effect of 1-propanol is more prominent.

In mixtures of chloroform and TFE rate constants are lower than in either pure

solvents, passing through a minimum at xTFE = 0.5. ET(30)-values are indicative of

58

1,3-Dipolar Cycloadditions

0.0 0.2 0.4 0.6 0.8 1.0 -1.0

-0.9

-0.8

-0.7

-0.6

-0.5

-0.4

log(

k 2 )

mole fraction TFE

0.0 0.2 0.4 0.6 0.8 1.0

114

115

116

117

118

119 E

T (30) (kcal/m

ol)

E T

( 1 )

(kca

l/mol

)

40

45

50

55

60

mole fraction TFE

FIGURE 2.9: Left : values of log(k2) for the reaction of 1 with 4 at 25 ◦C in mixtures

of chloroform and TFE. k2 has units M−1 s−1. Right : corresponding values of ET(1)

(�), and ET(30) (◦).

a preferential solvation of betaine-30 by TFE, and the same may well be the case for

1. However, ET(1) indicates that although a significant degree of preferential solvation

of 1 occurs, the preferential solvation is not as dramatic as in case of betaine-30 (Fig-

ure 2.9). For the sake of comparison, the reaction of 1 with 4 was also studied using

chloroform/TFE mixtures (Figure 2.9). A nearly linear dependence of log(k2) on the

mole fraction of TFE was found, with only a small deviation in the chloroform-rich end.

This pattern rules out any irregular effects of TFE on 1. We therefore conclude that

initially the catalytic effect of TFE on 2a is small, but gains importance at higher mole

fractions of TFE, analogous to mixtures of chloroform with 1-propanol.

These observations lead to the following explanation: the rate-accelerating effect

induced by hydrogen bonds depend on the solvent and are different for the different re-

actants. TFE interacts with 1 efficiently and induces catalytic effects that enhance (4)

or reduce (2a) the rate constant. The interactions are not so strong that extensive pref-

erential solvation of 1 by TFE occurs in mixtures of chloroform and TFE. 1-Propanol

interacts with 1 less efficiently, inducing smaller effects. By contrast, 1-propanol inter-

acts with 2a efficiently, and an overall accelerating effect is found for 1-propanol. In

mixtures of chloroform with TFE, the rate initially drops (1 + 2a) because of hydrogen

bonding of TFE to 1. However, at higher mole fractions of TFE, the rate constant

again increases, most likely because (additional) hydrogen bonds between TFE and 2a

are involved.

Compared to 1-propanol, the inhibition in TFE due to the hydrogen bonds formed to

1 seems to dominate. However, note that in the hypothetical case that only the polarity

would affect the rate constant (as illustrated in Figure 2.6) the rate constant would be

(much) lower. The catalytic effect (hydrogen bonds with 2a) therefore still contributes

59

Chapter 2

more than the inhibition (hydrogen bonds with 1). The inhibition is just more efficient

with TFE than with other alcohols.

In mixtures of 1-propanol with TFE, the many possibilities for TFE to form hydrogen

bonds to 1-propanol will reduce peculiarities in solvating 1 or 2a, and a linear dependence

of log(k2) on xTFE is found for the complete solvent composition range.

2.2.3 Water and Aqueous Solvent Mixtures

Rate constants for the reactions of 1 with 2a–c in water and 1-propanol are listed in

Table 2.3. The rate constants show a slight increase with increasing tail length R (Et,

n-Bu, Bz) in 1-propanol, and a more pronounced increase in water. In other words, the

rate constants in water, compared to 1-propanol (kw/kalcohol) increase with increasing hy-

drophobicity of the dipolarophile. This pattern has also been observed for DA reactions

of 2,3-dimethylbutadiene with N -alkylmaleimides,50 and could be due to an additional

hydrophobic interaction between the reactants, lowering the standard Gibbs energy of

activation (but see also Section 2.2.1, Activation Parameters). Another, extreme exam-

ple of this phenomenon is observed for the DC reactions of C,N -diphenylnitrone with

dimethyl fumarate and dibutyl fumarate, for which kw/kalcohol are 12 and 108, respec-

tively.3 Inspection of Figure 2.6 reveals that, compared to the trend found among the

alcohols, the rate constant in water shows a positive deviation. This small effect may be

attributed to ‘enforced hydrophobic interactions.’ Such a small contribution is to be ex-

pected with two polar reactants. From elaborate calculations on DA reactions involving

the hydrophobic cyclopentadiene, the contribution of enforced hydrophobic interactions

was estimated to be a factor of 5–6 in rate.51 In this case, the contribution is estimated

to be a factor of 2–3. For the reaction of 1 with 4, the effect appears larger, which can

be attributed to the fully apolar character of 4. A more detailed discussion is given in

Chapter 3.

TABLE 2.3: Rate constants in for the reaction of 1 with 2a–c in various media at

25 ◦C.

2a 2b 2c

Medium k2 (M−1s−1) k/kw k2 (M−1s−1) k/kw k2 (M−1s−1) k/kw

water 0.35 1 0.55 1 0.73 1

1-propanol 0.30 0.86 0.37 0.67 0.45 0.59

60

1,3-Dipolar Cycloadditions

0 10 20 30 40 50

0.0

0.2

0.4

0.6

0.8

1.0

1.2

k 2 (M

-1 s -1

)

[H 2 O] (M)

FIGURE 2.10: Rates in water/1-propanol mixtures (squares) and water/2-methyl-2-

propanol mixtures (circles), for the reaction of 1a with 2a (closed symbols) and 2b

(open symbols).

Rate Constants in Aqueous Alcohol Mixtures

In mixtures of water and 1-propanol or 2-methyl-2-propanol, with [alcohol] . 2 M, the

reactions of 1 with 2a and 2b are slightly accelerated, compared to water (Figure 2.10),

leading to a maximum in rate constant, at around 45 M of water. Similar maxima have

been observed for many cycloadditions, although there are examples where the effect is

absent (Chapter 3, Table 3.3). Plotted versus the mole fraction of water, the positions of

the maxima seem to depend on the hydrophobicity of the alcohol. However, Figure 2.10

reveals that the maximum appears at roughly equal concentrations of water. The small

size of a water molecule may lead to correlations with what turns out to be the size of

the cosolvent molecules, if the mole fraction scale is used for aqueous mixtures. The

usefullness of using the mole fraction scale for aqueous solvent mixtures may therefore

be questioned.

There is also a shallow minimum in rate constant, around 15 M of water. In this

concentration range, the hydrogen-bond network of water is completely disrupted and

hydrophobic effects do not play a role. Compared to 1-propanol, addition of a small

amount of water increases the hydrogen-bond donating capacity of the solvent mix-

ture, which for these reactions is both favorable (activation of the dipolarophile) and

unfavorable (deactivation/stabilization of the 1,3-dipole), and a net unfavorable effect

results.18 In methanol/water mixtures, the minimum is absent, and the rate constant

61

Chapter 2

a)

0 10 20 30 40 50

30

40

50

activ

atio

n pa

ram

eter

s (k

J m

ol -1

)

[H 2 O] (M)

b)

0 10 20 30 40 50

30

40

50

activ

atio

n pa

ram

eter

s (k

J m

ol -1

)

[H 2 O] (M)

FIGURE 2.11: Activation parameters for the reaction of 1 with a) 2a and b) 2b in

water/1-propanol mixtures at 25 ◦C: ∆‡G◦ − 30 (�), ∆‡H◦ (�), −T∆‡S◦ (◦).

62

1,3-Dipolar Cycloadditions

monotonously increases upon adding water, up to 40 M (Chapter 3). Once more, an

unusually complex dependence of rate constant on the nature of the reaction medium

is found.

Isobaric Activation Parameters in Mixtures of Water and 1-Propanol

Isobaric activation parameters have also been determined for the reactions of 1 with 2a

and 2b in water/1-propanol mixtures (Figure 2.11). Up to 40 M of water, the varia-

tion in activation parameters is small. ∆‡H◦ slightly decreases and −T∆‡S◦ increases

accordingly. In the water-rich mixtures, −T∆‡S◦ suddenly drops significantly, accompa-

nied by a similar increase in ∆‡H◦. In the water-rich regime, typical hydrophobic effects

come into play, and −T∆‡S◦ drops (initial state destabilization by HH). Whereas at

around 40 M of water, the solvent mixture behaves as a ‘normal’ polar solvent, at higher

water concentrations the characteristic features of water at room temperature (i.e. sol-

vation of apolar compounds accompanied by a large unfavorable entropy term) become

apparent. The patterns in the activation parameters are similar for 2a and 2b, as was

found among the pure solvents, but significant differences are observed for ∆‡H◦ and

∆‡S◦ in the concentration range where hydrophobic effects are important. The alkyl

substituent in the dipolarophile clearly has an influence on the activation process in

these aqueous mixtures. The butyl tail could, for example, induce a larger preference

of the dipolarophile for being solvated by propanol. These differences in solvation are

again ‘innocent’ and are not reflected in the rate constant. Nevertheless, these data

show that despite the absence of large effects of water on the rate constants of these

reactions, typical ‘aqueous’ behavior is occurring in water-rich mixtures.

2.3 Conclusions

A systematic study of solvent effects on the reactions of benzonitrile oxide (1) with N -

alkylmaleimides (2) and cyclopentene (4) has provided additional insight into the factors

that determine the rates of cycloadditions in different media. This study emphasizes

the importance of both polarity and hydrogen-bond donating capacity: (i) differences

in charge distributions of reactants and activated complex cause polar interactions with

the solvent (including hydrogen bonds) to enlarge or reduce the energy gap between

both states; (ii) a hydrogen bond formed between reactant and solvent (hydrogen-bond

donor) affects the HOMO and LUMO of the reactant, similar to the impact of a Lewis-

acid catalyst, and this specific effect can either accelerate or inhibit the reaction. For the

reactions presented in this chapter these two factors oppose each other, which partially

explains the very modest solvent effects on these reactions, and leads to the complex

dependence of rate constants on solvent.

63

Chapter 2

In the case of the reactions of the N -alkylmaleimides with benzonitrile oxide, hy-

drogen bonding (with corresponding changes in the FMOs) to both reactants occurs,

with opposite effects: hydrogen bonding to the dipolarophiles is favorable, but hydrogen

bonding to the benzonitrile oxide unfavorable. This pattern can be rationalized with

FMO theory, assuming the reaction has normal electron demand (NED).

In mixtures of solvents, of which only one is a hydrogen-bond donor, log(k) varies

gradually with the composition (often linearly), indicating the absence of significant

preferential solvation of the reactants due to hydrogen-bond interactions. The specific,

catalytic effects of hydrogen bonds are not accompanied by strong binding or complex-

ation of the hydrogen-bond donating solvent to the reactants.

Activation parameters reveal significant differences in solvation in different solvents,

that are not reflected in the rate constants.

In (highly) aqueous mixtures, hydrophobic effects are important, but this does not

lead to large increases in rate, because the contribution of these effects is modest, and

these effects are counteracted by other factors. This pattern contrasts with common

DA reactions, where polarity, hydrogen-bond donor capacity and enforced hydrophobic

interactions work together and can result in impressive rate accelerations in water.

2.4 Experimental

2.4.1 Materials

N -n-Butylmaleimide (2b)50 has been synthesized previously. The ET(30) probe was

kindly provided by prof. dr. Chr. Reichardt. All other materials were obtained from

commercial suppliers, and were of the highest purity available. Solvents were either

analytical grade or distilled. Acetonitrile was run over basic aluminium oxide prior to

use. Cyclopentene (4) was distilled before use.

2.4.2 Kinetic Experiments

The procedure, described in the literature,18 where 1 is generated in situ in a

CH2Cl2/bleach two-phase system, and small aliquots of the organic layer transferred

to the reaction mixture, was found to lead to poor kinetics for aqueous solutions, be-

cause of solubility problems. Instead, the preparation of 1 was performed by dissolving

benzaldoxime in a bleach/1-propanol mixture in a test tube and shaking this tube for a

few seconds. After the addition of sodium chloride a two-phase system quickly emerged,

and 0.5-1 µL of the organic layer was transferred to a quartz cuvet, which contained the

reaction mixture with the dipolarophile. This method led to excellent kinetics, and was

used for most kinetic measurements described in this chapter. No differences in rate

64

1,3-Dipolar Cycloadditions

constants were found in aprotic solvents, using 1-propanol rather than dichloromethane.

Kinetic measurements were performed using UV-VIS spectroscopy (Perkin Elmer λ5

or λ12 spectrophotometer). The dipolarophile was used in excess, and reactions were

monitored at 273 nm. The reactions were followed for at least four half-lives and pseudo-

first-order rate constants were obtained using a fitting program. Typical conditions were:

[dipolarophile]= 1–10 mM, [1,3-dipole]= ca. 0.025–0.05 mM. Activation parameters were

calculated from 4–5 rate constants in the temperature range of 20–40 ◦C.

UV-VIS spectra of 1 and ET(30) were recorded on a Perkin Elmer λ5 spectropho-

tometer at 25 ◦C. ET(30) values were calculated from the longest wavelength charge-

transfer absorption band of the dye, as ET(30) (kcal mol−1) = 28591/λmax (nm);29 ET(1)

values were calculated accordingly from its longest wavelength absorption band. The

solvatochromic dye was added to a known volume of solution or solvent mixture by

injecting a few microliters of a stock solution in ethanol (ET(30)) or of the 1-propanol

layer of the 1-propanol/bleach two-phase system in which the 1,3-dipole was generated

(see above).

References and Notes

[1] Padwa, A. 1,3-Dipolar Cycloaddition Chemistry; Wiley, New York, 1984.

[2] Eastman, C. J. In Advances in Heterocyclic Chemistry; Vol. 60; Academic Press: San

Diego, 1994.

[3] Gholami, M. R., Yangheh, A. H. Int. J. Chem. Kinet. 2001, 33, 118–123.

[4] Elender, K., Riebel, P., Weber, A., Sauer, J. Tetrahedron 2000, 56, 4261–4265.

[5] Pekasky, M. P., Jorgensen, W. L. Faraday Discuss. 1998, 110, 379–389.

[6] Domingo, L. R. Eur. J. Org. Chem. 2000, 2265–2272.

[7] Hu, Y., Houk, K. N. Tetrahedron 2000, 56, 8239–8243.

[8] Mendez, F., Tamariz, J., Geerlings, P. J. Phys. Chem. A 1998, 102, 6292–6296.

[9] Rastelli, A., Gandolfi, R., Amade, M. S. J. Org. Chem. 1998, 63, 7425–7436.

[10] Wijnen, J. W. Cycloadditions in Aqueous Media, Ph.D. thesis, University of Groningen,

1997.

[11] Huisgen, R., Fisera, L., Giera, H., Sustmann, R. J. Am. Chem. Soc. 1995, 117, 9671.

[12] Huisgen, R. Pure Appl. Chem. 1980, 52, 2283.

[13] Houk, K. N., Sims, J., Watts, C. R., Luskus, L. J. J. Am. Chem. Soc. 1973, 95,

7287–7301.

[14] Houk, K. N., Sims, J., R. E. Duke, J., Strozier, R. W., George, J. K. J. Am. Chem. Soc.

1973, 95, 7301–7315.

65

Chapter 2

[15] Konovalov, A. I., Samuilov, Y. D., Slepova, L. F., Breus, V. A. Zh. Org. Khimii, Eng.

Translation 1973, 9, 2539–2541 (2519–2521).

[16] Huisgen, R. Pure App. Chem. 1980, 52, 2283–2302.

[17] Araki, Y. I. K., Shiraishi, S. Bull. Chem. Soc. Jpn. 1991, 64, 3079–3083.

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8801–8805.

[19] Seerden, J.-P. G. Ph.D. thesis, Catholic University of Nijmegen, 1995.

[20] Curran, D. P., Kim, B. H., Piyasena, H. P., Loncharich, R. J., Houk, K. N. J. Org.

Chem. 1987, 52, 2137.

[21] Kanemasa, S., Nishiuchi, M., Wada, E. Tetrahedron Lett. 1992, 33, 1357–1360.

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[23] Anh, N. T., Maurel, F. New. J. Chem. 1997, 21, 861–871.

[24] Sauer, J., Sustmann, R. Angew. Chem. Int. Ed. Engl. 1980, 19, 779–807.

[25] This is possible because of the early transition state.

[26] Fisera, L., Sauter, F., Frohlich, J., Feng, Y., Ertl, P., Mereiter, K. Monatsh. Chem.

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[27] Toma, L., Quandrelli, P., Perrini, G., Gandolfi, R., Valentin, C. D., Corsaro, A.,

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[29] Reichardt, C. Chem. Rev. 1994, 94, 2319–2358.

[30] In a recent study, a hydrophilic alternative to betaine-30 has been described: Herodes,

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66

1,3-Dipolar Cycloadditions

[40] The relative energies of the FMOs of 1 and electron-poor dipolarophiles are such, that

both LUMO1–HOMOdipolarophile and HOMO1–LUMOdipolarophile interactions may

significantly contribute. The focus on only one of these HOMO–LUMO interactions may

therefore not be justified, although this simplification was sufficient to interpret the data

presented in this chapter.

[41] CCl4: tetrachloromethane; THF: tetrahydrofuran; CHCl3: chloroform

(trichloromethane); CH2Cl2: dichloromethane; CH3CN: acetonitrile (cyanomethane);

t-BuOH: 2-methyl-2-propanol; 2-BuOH: 2-butanol; 1-BuOH: 1-butanol; 1-PrOH:

1-propanol; 2-PrOH: 2-propanol; EtOH: ethanol; MeOH: methanol; TFE:

2,2,2-trifluoroethanol.

[42] DMSO is left out, as side reactions interfered with the cycloaddition. The choice of

solvents is further limited by the requirement of being able to monitor the reaction at

273 nm.

[43] Kadaba, P. K. Synthesis 1973, 71–84.

[44] Note that for ordinary Diels-Alder reactions, a larger hydrogen-bond donating capacity

of a solvent leads to an increase in rate because of TS stabilisation and that this is

reflected in a decrease in the enthalpy of activation. In terms of FMO theory: the

LUMO of the dienophile is lowered in energy because of (stronger) hydrogen bonding

and the energy-gap with the HOMO of the diene is decreased.

[45] In all cases, the plots hardly differ when converted to a molar scale, as in all cases, the

molar volumes of the solvents are similar and the extent of nonideal mixing limited. (In

all cases, the concentrations of solvent 2 were 0, 0.2, 0.4, 0.6, 0.8, and 1 × molarity of

pure solvent 2.).

[46] Elias, H., Gumbel, G., Neitzel, S., Volz, H. Fres. Z. Anal. Chem. 1981, 306, 240–244.

[47] Balakrishnan, S., Easteal, A. J. Aust. J. Chem. 1981, 34, 933–941.

[48] Maksimovic, Z. B., Reichardt, C., Spiric, A. Z. Anal. Chem. 1974, 270, 100–104.

[49] The hydrogen bond donor capacity α is 0.78 for 1-propanol, 0.93 for methanol, 1.17 for

water, and 0.44 for chloroform. This indicates, that chloroform has an ability to form

hydrogen bonds, but to a much lesser extent than any alcohol.

[50] Meijer, A., Otto, S., Engberts, J. B. F. N. J. Org. Chem. 1998, 63, 8989–8944.

[51] Chandrasekhar, J., Shariffskul, S., Jorgensen, W. L. J. Phys. Chem. B 2002, 106,

8078–8085.

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Chapter 2

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