University of Groningen
Cycloadditions in weakly and highly organized aqueous media
Rispens, Taede
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Rispens, T. (2004). Cycloadditions in weakly and highly organized aqueous media: kinetic studies of
cycloadditions in aqueous solvent mixtures and surfactant solutions Groningen: s.n.
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Chapter 2
Solvent Effects on 1,3-Dipolar
Cycloadditions of Benzonitrile Oxide
he 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.
T
The aforementioned reactions are examples for which these factors counteract, leading 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 extensively 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 theoretically.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 negative 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-methylC-phenylnitrone (I and II in Scheme 2.1, resp.), and are the result of the partial cancellation 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
a
c
b
b
b
b
c
a
c
a
c
µ (Debye)
N
I
N
O
N
N
CH3
H
II
C6H5
O
0.17
CH3
H
N
C
O
N
C
O
3.55
C6H5
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 1,3dipole 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 E T (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
LUMO
LUMO
LUMO
HOMO
E
HOMO
HOMO
electron-poor
dienophile/
dipolarophile
diene/1,3-dipole
electron-rich
dienophile/
dipolarophile
FIGURE 2.1: Schematic representation of the dominant MO interactions in a cycloaddition 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 hydrogen bonds (Section 2.1.2), and the E T (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 observations. 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 (dominant) 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 cyclopentadiene 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. Hydrogen 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 E T (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 illustrated 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 susceptible 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 E T (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
hν
N
O
O
Betaine-30
SCHEME 2.2.
47
Chapter 2
2
LUMO
0
0
-1/-0.5
LUMO
E (eV)
electron-poor
dipolarophile
LUMO
benzonitrile
oxide
electron-rich
dipolarophile
-9
-10/-9.4
-10
-11
HOMO
HOMO
HOMO
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 Studies 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 dipolarophiles 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 decellerated. 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 dipolarophiles 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 explanation 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 interactions), 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, hydrogen 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 hydration 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
CH
NOH
Bleach
C N O
O
N
C N O +
O
N R
O
O
N
O
1
2
R
3
(2a)
R = Et
n-Bu
(2b)
CH2Ph
(2c)
N
C N O +
4
O
5
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
-0.2
0.0
1-BuOH
2-BuOH
1-PrOH
t-BuOH
2-PrOH
EtOH
MeOH
n-hexane
log(k2)
-0.6
H2O
H2O
-0.2
2 + log(k2)
-0.4
CCl4
-0.8
1,4-dioxane
CH3CN
-1.0
THF
-0.4
2-BuOH
n-hexane
-0.6
2-PrOH EtOH
MeOH
CCl4
CH2Cl2
-0.8
1,4-dioxane CH Cl
2
2
TFE
-1.2
CHCl3
TFE
1-PrOH
CHCl3
-1.0
CH3CN
-1.4
30
35
40
45
50
55
60
65
30
ET(30)
35
40
45
50
55
60
65
ET(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 E T (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
E T (30)b
n-hexane
30.9
a
CCl4
32.5
1,4-dioxane
36.0
THF
37.4
chloroform
39.1
CH2 Cl2
41.1
t-BuOH
43.3
CH3 CN
45.6
c
k2,2a
0.330
0.210
0.12e
0.100
0.059
0.07e
0.254
0.10
d
k2,4
0.333
0.255
0.170
n.d.
0.127
0.120
n.d.
0.124
Solvent E T (30)b
2-BuOH
47.1
2-PrOH
48.4
1-BuOH
49.7
1-PrOH
50.7
EtOH
51.9
MeOH
55.5
TFE
59.8
water
61.3
c
k2,2a
0.308
0.289
0.310
0.320
0.22e
0.196
0.061
0.350
d
k2,4
0.334
0.274
n.d.
0.259
0.265
0.229
0.380
0.978
a) 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
0.5
acetonitrile
water
TFE
0.4
1-propanol
chloroform
TFE
water
118
E (1)(kcal/mol)
n-hexane
120
116
acetonitrile
1-propanol
T
0.3
chloroform
A
114
n -hexane
112
0.2
30
40
50
E T (30) (kcal/mol)
60
0.1
0.0
220
240
260
280
wavelength (nm)
FIGURE 2.5: UV spectra of 1 in various solvents at 25 ◦ C. The inset shows the
corresponding transition energies ((ET (1) plotted against E T (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 experimentally10 and theoretically.5 Nevertheless, the deviation remains an unusual observation.
An alternative explanation, based on a change from a concerted mechanism to a mechanism 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) zwitterionic 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 E T (30)values between 30–40 (Figure 2.6), to illustrate the divergence from this trend for the
more polar solvents (E T (30)-values > 40). The E T (30)-scale is based on one parameter, that includes hydrogen-bond donor capacity, hydrogen-bond acceptor capacity, and
53
Chapter 2
-0.2
-0.4
2-BuOH 1-BuOH
n-hexane
1-PrOH
t-BuOH
EtOH
CCl4
-0.8
1,4-dioxane
MeOH
CH3CN
-1.0
THF
CH2Cl2
-1.2
TFE
CHCl3
-1.4
log(k2)
H2O
2-PrOH
-0.6
-1.6
-1.8
-2.0
0.0
-2.2
H2 O
2-BuOH
-2.4
TFE
-0.5
MeOH
-2.6
-1.0
-2.8
30
35
40
45
50
55
60
65
-3.0
30
35
40
45
50
55
60
65
ET(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 E T (30), as betaine-30 is rather sensitive to these interactions.
In fact, the fair correlation of log(k) with E T (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
E T (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 consistent 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
n-hexane
chloroform
1-propanol
trifluoroethanol
water
75.8
80.0
75.9
79.9
75.7
n-hexane
1-propanol
trifluoroethanol
water
75.5
75.2
79.5
74.6
∆‡ H ◦ (kJ mol−1 )b
1+2a
42.5
51.8
42.7
34.4
50.2
1+2b
40.8
40.5
33.0
48.0
−T ∆‡ S ◦ (kJ mol−1 )b
33.3
28.2
33.1
45.5
25.5
34.7
34.8
46.4
26.6
a) Standard error < 0.1 kJ mol−1 . b) Standard error 1.5–2 kJ mol−1 for 2a; 1 kJ mol−1 for
1b.
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 nhexane 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 expectation, 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 E T (30)-values in these mixtures provides information concerning the
influence of the solvent polarity on the reaction. The E T (30)-values already offer clues
concerning the nature of the mixtures. In all cases, E T (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 responsible for the sharp increase in E T (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 E T (30) on the mole fraction resembling a binding curve at low
amounts of propanol or ethanol, and a more gradual, linear dependence once ‘saturation’ 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
-0.6
ET(30) (kcal/mol)
42
-0.8
38
36
34
-0.6
0.0
0.2
0.4
0.6
0.8
mole fraction CHCl3
1.0
-0.8
-1.0
log(k)
log(k2)
40
-1.0
-1.2
-1.2
34
0.0
36
38
40
ET(30) (kcal/mol)
0.2
42
0.4
0.6
0.8
1.0
mole fraction CHCl3
a)
52
ET(30) (kcal/mol)
-0.4
-0.6
50
48
46
44
42
0.0
-0.8
0.2
0.4
0.6
0.8
log( k 2)
1.0
log(k2)
-0.5
-1.0
-1.0
-1.2
-1.5
42
0.0
0.2
0.4
44
46
0.6
48
E T(30)
0.8
50
52
1.0
mole fraction 1-PrOH
b)
-1.2
ET(30) (kcal/mol)
60
log(k2)
-1.3
55
50
45
40
0.0
0.2
0.4
0.6
0.8
1.0
-1.4
-1.5
0.0
0.2
0.4
0.6
0.8
1.0
mole fraction TFE
c)
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
62
-0.4
ET(30) (kcal/mol)
60
-0.6
log(k2)
56
54
52
50
log( k 2 )
0.0
-0.5
-0.8
58
0.2
0.4
0.6
0.8
1.0
-1.0
-1.0
-1.5
50
-1.2
0.0
0.2
52
54
56
58
E T (30)
60
0.4
62
0.6
0.8
1.0
mole fraction TFE
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 governed primarily by the polarity of the medium. As reactants and betaine-30 are preferentially solvated to different degrees, a correlation between log(k2 ) and E T (30) over the
full range of composition is not observed.
In chloroform/1-propanol mixtures, a good hydrogen bond donor is present (1propanol).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 ) gradually 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. E T (30)-values are indicative of
58
1,3-Dipolar Cycloadditions
-0.4
119
-0.5
118
60
ET(1) (kcal/mol)
log(k2)
-0.7
-0.8
55
117
116
50
115
45
ET(30) (kcal/mol)
-0.6
-0.9
114
-1.0
0.0
40
0.2
0.4
0.6
0.8
1.0
0.0
0.2
0.4
0.6
0.8
1.0
mole fraction TFE
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 (Figure 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 reactants. 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 preferential 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 interacts 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 hydrophobicity 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 example 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, respectively.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 expected 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
Medium
k2 (M s )
water
0.35
1-propanol
0.30
−1 −1
k/kw
1
0.86
2b
k2 (M s )
0.55
0.37
−1 −1
60
k/kw
1
0.67
2c
k2 (M s )
0.73
0.45
−1 −1
k/kw
1
0.59
1,3-Dipolar Cycloadditions
1.2
1.0
k2 (M-1s-1)
0.8
0.6
0.4
0.2
0.0
0
10
20
30
40
50
[H2O] (M)
FIGURE 2.10: Rates in water/1-propanol mixtures (squares) and water/2-methyl-2propanol 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 mixture, 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
activation parameters (kJ mol-1)
50
40
30
0
10
20
30
40
50
40
50
[H2O] (M)
a)
activation parameters (kJ mol-1)
50
40
30
0
10
20
30
[H2O] (M)
b)
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 variation in activation parameters is small. ∆‡ H ◦ slightly decreases and −T ∆‡ S ◦ increases
accordingly. In the water-rich mixtures, −T ∆‡ S ◦ suddenly drops significantly, accompanied 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. solvation 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 Nalkylmaleimides (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 Lewisacid 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, hydrogen 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 complexation 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 E T (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
CH2 Cl2 /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, because 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 pseudofirst-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 E T (30) were recorded on a Perkin Elmer λ5 spectrophotometer at 25 ◦ C. E T (30) values were calculated from the longest wavelength chargetransfer absorption band of the dye, as E T (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 (E T (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] Méndez, F., Tamariz, J., Geerlings, P. J. Phys. Chem. A 1998, 102, 6292–6296.
[9] Rastelli, A., Gandolfi, R., Amadé, 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.
[18] van Mersbergen, D., Wijnen, J. W., Engberts, J. B. F. N. J. Org. Chem. 1998, 63,
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.
[22] Houk, K. N. Acc. Chem. Res. 1975, 8, 361–369.
[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] Fišera, L., Sauter, F., Fröhlich, J., Feng, Y., Ertl, P., Mereiter, K. Monatsh. Chem.
1994, 125, 553–563.
[27] Toma, L., Quandrelli, P., Perrini, G., Gandolfi, R., Valentin, C. D., Corsaro, A.,
Caramella, P. Tetrahedron 2000, 56, 4299–4309.
[28] Christl, M., Huisgen, R. Tetrahedron Lett. 1968, 9, 5209–5213.
[29] Reichardt, C. Chem. Rev. 1994, 94, 2319–2358.
[30] In a recent study, a hydrophilic alternative to betaine-30 has been described: Herodes,
K., Koppel, J., Reichardt, C., Koppel, I. A. J. Phys. Org. Chem. 2003, 16, 626–632.
[31] Reichardt, C. Solvents and Solvent Effects in Organic Chemistry; VCH: Weinheim,
1988.
[32] Gabriel, S., Koppe, M. Ber. Dtsch. Chem. Ges. 1886, 19, 1145–1148.
[33] Werner, A., Buss, H. Ber. Dtsch. Chem. Ges. 1894, 27, 2193.
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[37] Dondoni, A., Barbaro, G. J. Chem. Soc., Perkin Trans. 2 1973, 1769–1773.
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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); CH2 Cl2 : dichloromethane; CH3 CN: 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] Maksimović, Z. B., Reichardt, C., Spirić, 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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