Ureas and amides as dipolar aprotic solvents in highly basic media. The dependence of kinetic basicity on solvent composition

Article abstract of DOI:10.1039/a808435a

The basicity of dipolar aprotic solvent water HO systems with amides and ureas as the organic component has been studied kinetically because previous information is not available. excluding some H values measured for aqueous dimethylformamide (DMF) and tetramethylurea (TMU). It was found that the increase in basicity with the mole fraction of organic component is at least of the same magnitude as in aqueous dimethyl sulfoxide (DMSO). For instance, in the detritiation of chloroform-t the slopes of the plots log(k₂/mol ¹ dm³ s ¹) vs. x(urea) varied between 11.4 14.6 (as compared to 11.0 in aqueous DMSO) when TMU and cyclic ureas. 1.3-dimethylimidazolidin-2-one (DMI) and 1.3-dimethyl-3.4.5.6-tetrahydropyrimidin-2(1H)-one (DMPU). were used as the organic component in solvent mixture. In aqueous TMU acidity functions H were extrapolated from kinetic results using linear free energy correlations. Agreement with literature values was evident. This method was also used to extrapolate the H values in aqueous DMPU. On the basis of present work aqueous ureas can be recommended as solvents in highly basic media. The utility of amides. dimethylformamide and dimethylacetamide. is limited by their instability in basic water solutions.

Full text of DOI:10.1039/a808435a

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Ureas and amides as dipolar aprotic solvents in highly basic media.
The dependence of kinetic basicity on solvent composition
Alpo Kankaanperä,* Pirketta Scharlin,* Ilona Kuusisto, Riitta Kallio and Emma Bernoulli
Department of Chemistry, University of Turku, FIN-20014, Turku, Finland
Received (in Cambridge) 30th October 1998, Accepted 14th December 1998
The basicity of dipolar aprotic solvent–water–HOϪ systems with amides and ureas as the organic component has
been studied kinetically because previous information is not available, excluding some HϪ values measured for
aqueous dimethylformamide (DMF) and tetramethylurea (TMU). It was found that the increase in basicity with the
mole fraction of organic component is at least of the same magnitude as in aqueous dimethyl sulfoxide (DMSO). For
instance, in the detritiation of chloroform-t the slopes of the plots log(k2/molϪ1 dm3 sϪ1) vs. x(urea) varied between
11.4–14.6 (as compared to 11.0 in aqueous DMSO) when TMU and cyclic ureas, 1,3-dimethylimidazolidin-2-one
(DMI) and 1,3-dimethyl-3,4,5,6-tetrahydropyrimidin-2(1H)-one (DMPU), were used as the organic component in
solvent mixture. In aqueous TMU acidity functions HϪ were extrapolated from kinetic results using linear free energy
correlations. Agreement with literature values was evident. This method was also used to extrapolate the HϪ values
in aqueous DMPU. On the basis of present work aqueous ureas can be recommended as solvents in highly basic
media. The utility of amides, dimethylformamide and dimethylacetamide, is limited by their instability in basic
water solutions.
fraction of amide or urea is between 0.1 and 0.3, depending on
Introduction
the organic component. These maxima suggest12 that water
Water–dipolar aprotic solvent–HOϪ systems are very useful
forms stable complexes with ureas and amides. The measured
highly basic media because their basicity can be changed
viscosities of aqueous TMU,12 DMPU,13 DMF17 and DMA16
gradually by increasing the amount of organic component
are also in accordance with this complex formation as the
in the solvent.1a As an additional advance, the solubility of
maximum value is observed when the mole fraction of amide or
organic compounds increases simultaneously. For instance,
urea is in the above mentioned range. Negative excess volumes
dimethyl sulfoxide (DMSO), sulfolane, acetonitrile, N,N-
of aqueous TMU,12 DMI,18 DMPU,13,14 DMF19 and DMA16
dimethylformamide (DMF), hexamethylphosphoric triamide
reveal that interactions between water and amide (or urea) are
(HMPA), pyridine, and tetramethylurea (TMU) have been
more significant than water–water or amide–amide (or urea–
urea) interactions. As a result of the complex formation
between water and urea (or amide) desolvation of hydroxide
used as the organic component and acidity functions, HϪ, for
these systems are available.1a,2–7 The highest basicity is obtained
in aqueous HMPA6,8 but its suitability for use is reduced by
ion may be exceptionally high when water is replaced by urea or
its carcinogenic properties. Cyclic ureas, 1,3-dimethylimid-
amide. A considerable increase in basicity is thus expected
azolidin-2-one (DMI, I) and 1,3-dimethyl-3,4,5,6-tetra-
hydropyrimidin-2(1H)-one (DMPU, II), which can replace
in these solvents. The maxima in densities and viscosities are
especially high in water–DMPU mixtures.13,14 Consequently, a
HMPA as the solvent, have been studied intensively since the
significant increase in basicity is expected with x(DMPU).
eighties.9 However, information on the basicity of the water–
Some values for the acidity function HϪ have previously been
cyclic urea (I) or (II)–HOϪ system is not available. Therefore,
measured for aqueous DMF4 and aqueous TMU.6
we studied in this work the suitability of DMI and DMPU
in highly basic media. The basicity of some acyclic amide
In this work the basicities of different amide and urea sys-
tems were determined kinetically. The detritiation of carbon
and urea analogs, N,N-dimethylformamide HCON(CH3)2,
acids [eqns. (1) and (2)] was chosen as the model reaction in
N,N-dimethylacetamide (DMA) CH3CON(CH3)2, and tetra-
methylurea (CH3)2NCON(CH3)2, was also studied for
the measurements. Eqn. (3) can be obtained for the rate of
comparison.
k1
RT ϩ HOϪ
RϪؒTOH
RϪؒTOH
(1)
(2)
kϪ1
k2
RϪ ϩ TOH
N
N
N
N
H3C
CH3
H3C
CH3
ν = [k1k2/(kϪ1 ϩ k2)][RT][HOϪ] = kobs[RT][HOϪ] (3)
O
O
I
II
detritiation by applying the steady state approximation to the
hydrogen bonded complex RϪؒTOH.1b In the detritiation of
those carbon acids which produce localized carbanions
The amides and ureas mentioned above are sufficiently
dipolar for water–aprotic solvent–HOϪ mixtures with the
following relative permittivities at 298 K: 37.60 (DMI),10
36.12 (DMPU),10 37.0 (DMF),11 37.78 (DMA)11 and 23.45
(TMU).11 Relative permittivities of aqueous TMU12 and
DMPU13 decrease gradually from 78.39 to that of pure urea. In
the density values of aqueous TMU,12 DMPU,13,14 DMF15
and DMA16 at 298.15 K a maximum is observed when the mole
kϪ1 ӷ k2 and consequently kobs = (k1/kϪ1)k2 = Kk2 in eqn. (3).
When delocalized carbanions are formed in the detritiation
k2 ӷ kϪ1, which means that in eqn. (3) kobs = k1. Thus, depen-
ding on the mechanism of detritiation, the observed rate
coefficient is proportional either to the equilibrium constant of
the ionization reaction or to the rate coefficient of this ioniz-
J. Chem. Soc., Perkin Trans. 2, 1999, 169–174
169

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ation. Because both of these constants are dependent on the
medium basicity, the measured log (kobs/molϪ1 dm3 sϪ1) values
give information on changes in solvent basicity. As previously,20
the detritiation of chloroform-t CTCl3 (1) was chosen as the
model reaction. To show that the order in basicity is indepen-
dent of the carbon acid used in the kinetic measurements, the
rate coefficients of the detritiation of 1,1,1-trifluoro-2-bromo-
2-chloroethane F3CCBrClT (2, halotane-t) were determined
in aqueous DMA and DMPU. In addition, 1,1-difluoro-2,2-
dichloroethyl methyl ether CH3OCF2CCl2T (3, methoxy-
flurane-t) was submitted to kinetic studies in aqueous cyclic
ureas to extend the measurements to mixtures with x(amide or
urea) at least 0.5. In aqueous ureas some measurements were
performed with acetophenone-t PhCOCH2T (4) as the model
carbon acid.
chloroethane (halotane) and acetophenone. The labelling of
chloroform and acetophenone has been described previously.24
Methoxyflurane and halothane were labelled by the same
method.
Detritiation of carbon acids
The performance of kinetic measurements has been described
previously.20 In the reaction mixture the concentration of
tetramethyl ammonium hydroxide was 0.011 mol dmϪ3. The
second order rate coefficients were found to be independent
of hydroxide ion concentration. The base was added just before
the reaction was started to exclude the possible decomposition
of amide or urea. The detritiation of chloroform-t and
acetophenone-t was studied at 298.15 K. With chloroform-t
measurements were performed at four temperatures and the
rate coefficients at 298.15 K were calculated on the basis of
the Arrhenius equation. The detritiation of halotane-t and
methoxyflurane-t was studied mainly at 278.15 K. In some
mixtures measurements were performed at 3 or 4 temperatures.
In dilute aqueous solutions the HOϪ promoted hydrolysis of
amides21,22 takes place with attack of the hydroxide ion on the
acyl carbon. The reaction produces amines and carboxylate
ions, lowering the hydroxide ion concentration in aqueous
amide or urea. If this base-promoted reaction is significant, the
utility of the amide (or urea)–water system as a highly basic
medium is questionable. The instability of the dimethyl-
formamide–water system containing hydroxide ion is evident
on the basis of previous data.23 If the instability has not been
taken into account when kinetic measurements have been
performed in aqueous DMF, the validity of rate coefficients
is questionable due to the uncertainty of the hydroxide ion
concentration of the solution. Therefore, we studied sys-
tematically the effect of the possible hydrolytic decomposition
on hydroxide ion concentration in aqueous amide or urea.
Decomposition of aqueous amide and urea mixtures in basic
solutions
The hydrolytic decomposition of amides and ureas was studied
at 298.15 K in mixtures of x(amide or urea) ≈ 0.1. In water–
DMF and water–DMA solvent mixtures the composition
was also varied. In kinetic runs the concentration of sodium
hydroxide was 0.02 mol dmϪ3. The hydrolytic decomposition
was followed with a Metrohm 644 conductometer equipped
with a thermostated reaction cell. To prevent the absorption of
carbon dioxide, nitrogen was passed over the solution during
the reaction. As the hydrolytic decomposition of DMF and
DMA was relatively fast, the progress of the reaction was
followed directly in the reaction cell. For tetramethylurea and
cyclic ureas 10 cm3 samples were taken from the reaction
mixture and quenched with 10 cm3 of ice-cold water to stop the
reaction. Hydroxide ion concentration in these samples was
determined conductometrically by titrating with 2 mol dmϪ3
aqueous hydrogen chloride. In titration the acid solution was
added with an Agla semimicro syringe. The consumption of
aqueous hydrogen chloride was proportional to the progress of
the reaction.
Experimental
Materials
Dipolar aprotic solvents used in this work were TMU (Merck,
99%), DMI (Aldrich-Chemie, 98% and Fluka 99.5%), DMPU
(Aldrich-Chemie, 98%), DMF (Merck 99.5%) and DMA
(Fluka, 99%). The solvents were distilled at reduced pressure
over calcium hydride. The middle fraction was collected. The
products were stored in brown bottles to prevent possible
photochemical decomposition reactions. Laboratory reverse-
osmosis water was degassed before use. Mixtures of water with
amides and ureas were prepared by weight.
Results
The following tritiated carbon acids were used in kinetic
measurements: chloroform, 1,1-difluoro-2,2-dichloroethyl
methyl ether (methoxyflurane), 1,1,1-trifluoro-2-bromo-2-
Hydrolytic decomposition of ureas and amides
First-order rate coefficients for the hydrolytic decomposition
of amides and ureas are given in Table 1. Dimethylformamide
and dimethylacetamide were studied in several water–amide
mixtures. The rate coefficients in aqueous DMF are in accor-
dance with those of Buncel and Symons23 (Table 1). A
maximum is observed at x(DMF) ≈ 0.2. In aqueous DMA the
dependence of rate coefficients on solvent composition remains
obscure on the basis of present data.
The stability of aqueous TMU, DMI and DMPU was
studied only in mixtures with x(organic component) = 0.1
(Table 1). The rate coefficients for the hydrolytic decomposition
give half lives of 16 d (TMU), 5 d (DMI) and 10 d (DMPU)
which are markedly higher than those for DMA (4 h) and DMF
(9 min) in the same conditions.
Table 1 Hydrolytic decomposition of different amides and ureas in
their mixtures with water at 298.15 K. The initial concentration of
sodium hydroxide was 0.02 mol dmϪ3
Amide or urea
DMF
x (amide or urea)
k/10Ϫ4 sϪ1
Reference
0.076
0.091
0.291
0.303
0.499
0.536
0.806
9.9
13.6
18.5
19.2
5.8
6.1
2.3
23
This work
23
This work
This work
23
23
DMA
0.095
0.192
0.313
0.499
0.56
0.63
0.53
0.81
This work
This work
This work
This work
Detritiation of different carbon acids in aqueous ureas or amides
Kinetic results for the detritiation of chloroform-t (1), halotane-t
(2), methoxyflurane-t (3) and acetophenone-t (4) are presented
in Table 2. In the detritiation of chloroform-t the measurements
at 298.15 K could be extended to solutions with x(amide or
urea) < 0.2. When the logarithms of the rate coefficients in
water–DMPU and water–DMA mixtures are plotted against
the mole fraction of organic component, a linear correlation is
DMI
0.100
0.100
0.100
0.015
This work
This work
This work
DMPU
TMU
0.0081
0.0049
170
J. Chem. Soc., Perkin Trans. 2, 1999, 169–174

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Table 2 Second order rate coefficients for the detritiation of ST in water–amide and water–urea mixtures: (1) chloroform-t, (2) halotane-t,
(3) methoxyflurane-t and (4) acetophenone-t. 0.011 mol dmϪ3 tetramethylammonium hydroxide has been used as the catalyst
ST
T/K
Amide or urea
DMF
x (amide or urea)
/molϪ1 dm3 sϪ1
αc
k2
1
298.15
0.000
0.172a
0.0486
0.0500
0.0966
0.157
0.406 0.005
0.433 0.013
1.43 0.03
6.50 0.25
11.07 0.10d
11.3 0.3
DMA
TMU
0.000
0.172a
0.0483
0.0952
0.140
0.583 0.003
1.87 0.02
6.45 0.17
0.000
0.172a
0.0273
0.0492
0.102
0.324 0.004
0.596 0.005
2.90 0.18b
8.92 0.20
0.143
0.188
0.000
33.6 1.7b
12.32 0.18
DMI
0.172a
0.0149
0.0267
0.0495
0.0512
0.0955
0.1009
0.142
0.202 0.002
0.265 0.001
0.483 0.005
0.565 0.035b
1.54 0.12b
1.64 0.15b
5.21 0.06
0.193
22.8 0.5b
11.4 0.3d
DMPU
DMA
0.000
0.172a
0.0250
0.0496
0.0763
0.0972
0.190
0.383 0.005
0.931 0.006
2.39 0.03
4.69 0.01b
100 4b
14.60 0.13
2
278.15
0.000
0.0235 0.0020b
0.0572 0.0013
0.214 0.004
0.124 0.001
0.680 0.014
2.20 0.08
0.0369
0.0667
0.0735
0.109
0.146
0.154
2.65 0.04
13.8 0.4
18.3 0.5
DMPU
DMI
0.000
0.0235 0.0020b
0.0856 0.0015
0.314 0.003
0.730 0.013
2.72 0.05
0.0256
0.0614
0.0795
0.111
3
278.15
0.000
0.0977
0.192
0.265
0.387
0.488
0.000
0.0966
0.193
0.262
0.372
0.459
0.503
0.000080 0.000003b
0.000348 0.000005
0.00328 0.00007
0.0160 0.0002
0.187 0.016b
2.22 0.03
9.59 0.18d
DMPU
0.000080 0.000003b
0.000704 0.000012
0.0055 0.0006
0.0243 0.0002
0.311 0.010b
2.02 0.05
5.0 1.0
9.56 0.05
4
298.15
TMU
DMI
0.000
0.0518
0.112
0.197
0.377
0.573
0.000
0.0529
0.111
0.192
0.379
0.570
0.664
0.000
0.0509
0.101
0.190
0.375
0.550
0.00574a
0.00669 0.00005
0.01010 0.00014
0.0224 0.0003
0.119 0.002
0.740 0.005
0.00574a
4.043 0.004e
0.00488 0.00008
0.00690 0.00007
0.0148 0.0004
0.0981 0.0017
0.682 0.001
1.89 0.02
4.41 0.04e
4.62 0.10
DMPU
0.00574a
0.00926 0.00006
0.0161 0.0001
0.0352 0.0005
0.255 0.004
2.08 0.02
a See ref. 20. b Calculated from the measurements at different temperatures. c Slope of the plot log (k2/molϪ1 dm3 Ϫ1) vs. x(amide or urea). d Points
s
x(DMF) = x(DMI) = 0.000 are not included in the calculation of slopes. e Only the points with x(urea) higher than 0.1 are included in the calculation of
slopes.
J. Chem. Soc., Perkin Trans. 2, 1999, 169–174
171

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Fig. 2 Logarithms of the second-order rate coefficients in the detriti-
ation of methoxyflurane-t (3) at 278.15 K and acetophenone-t (4) at
298.15 K. Notations: Water–DMPU (square), water–DMI (triangle),
methoxyflurane-t (solid symbol), and acetophenone-t (open symbol).
Fig. 1 Logarithms of the second-order rate coefficients in the detriti-
ation of chloroform-t (1) at 298.15 K and halotane-t (2) at 278.15 K.
Notations: Water–DMPU (square), water–DMA (circle), chloroform-t
(open symbol), and halotane-t (solid symbol).
DMA > DMF. As the half lives of urea systems are at least
several days, ureas are stable enough to be used as highly basic
media. The situation is quite different with aqueous amides as
their half lives in hydrolytic decomposition are of the magni-
tude of minutes or hours on the basis of the kinetic data in
Table 1. If, for instance, aqueous dimethylformamide or
dimethylacetamide are used in kinetic measurements, the base
must be added to the reaction mixture just before the reaction is
started to eliminate the decrease in hydroxide ion concen-
tration. If, in addition, the reactions studied are relatively fast,
the change in hydroxide ion concentration remains neglible
during the reaction. Extrapolation to zero time can also be used
as previously in the determination of the equilibrium basicity
of aqueous DMF.4
observed (Fig. 1). In the detritiation of chloroform-t in
water–DMF, water–TMU, and water–DMI mixtures a linear
dependence also exists with the exception of the points
x(DMF) = x(DMI) = 0.
The detritiation of halotane-t was studied at 278.15 K in
aqueous DMA and DMPU. These measurements could be
extended to mixtures with x(organic component) < 0.15. A
linear correlation between logarithms of the rate coefficients
and the mole fraction of urea or amide is evident also in this
reaction (Fig. 1). The slopes of these plots, however, are
markedly higher than the corresponding slopes in the detriti-
ation of chloroform-t.
To extend the kinetic measurements to mixtures in which the
mole fraction of urea is higher than above, the detritiation
of methoxyflurane-t was studied in aqueous DMI and DMPU.
At 278.15 K even mixtures with x(urea) ≈ 0.5 could be used in
kinetic measurements. There is a linear correlation between
log(k2/molϪ1 dm3 sϪ1) and x(urea) (Fig. 2) also in this range,
except the point x(DMI) = 0. In addition, these two plots are
parallel within experimental errors.
The detritiation of acetophenone-t was studied in aqueous
ureas. These measurements could be extended even to mixtures
with x(urea) > 0.5. As shown in Fig. 2, the plots of log(k2/molϪ1
dm3 sϪ1) vs. x(DMI or DMPU) are linear and parallel if in
aqueous DMI the range x(DMI) < 0.1 is excluded. In aqueous
tetramethylurea a linear correlation is also evident on the
grounds of kinetic data in Table 2 when x(TMU) > 0.1.
Kinetic basicity of aqueous amides and ureas
As shown previously,20 in aqueous dimethyl sulfoxide the
logarithms of the rate coefficients for detritiation of different
carbon acids depend linearly on the equilibrium basicity HϪ of
the solvent. Therefore, the slopes of the plots log(k2/molϪ1 dm3
s
Ϫ1) vs. x(urea or amide) in Table 2 give also information on the
increase in basicity. For the detritiation of chloroform-t these
slopes vary between 11.07 and 14.60 with the highest slope in
water–DMPU mixtures. In aqueous DMSO the corresponding
slope is 11.0 0.2 as obtained from previous kinetic data.20
On the basis of the detritiation of chloroform-t it is obvious
that in aqueous amide and urea systems the increase in basicity
is slightly higher than in water–DMSO mixtures. From the
previous data8 the slope 21.3 0.8 can be calculated for the
detritiation of chloroform-t in aqueous HMPA. This indicates
that in the water–HMPA system the increase in basicity with
the mole fraction of the dipolar aprotic solvent is significantly
higher than in aqueous DMPU. However, taking into account
the issue of safety at work, aqueous DMPU is very capable of
being used as the solvent in highly basic medium.
Discussion
Stability of aqueous amides and ureas in basic solutions
Based on the kinetic results in Table 1, the stability of the
water–amide or urea system containing sodium hydroxide as
the base decreases in the order TMU > DMPU > DMI ӷ
172
J. Chem. Soc., Perkin Trans. 2, 1999, 169–174

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In the detritiation of halotane-t the slopes of the plots
log(k2/molϪ1 dm3 sϪ1) vs. x(DMA or DMPU) are also presented
in Table 2. In aqueous DMPU the slope is higher than in the
corresponding DMA system by a factor of 1.33 which means
that the increase in basicity is significantly higher in cyclic
urea. In the detritiation of chloroform-t the corresponding
factor 1.29 is equal within experimental error. Kinetic basicity
seems thus to be independent of the carbon acid used in the
measurements. In addition, the slopes in Table 2 reveal that the
rate increase as a function of x(amide or urea) is much higher
in the detritiation of halotane-t than in the detritiation of
chloroform-t, the slopes differing by factors of 1.22 (DMA)
and 1.25 (DMPU). The higher rate increase in the detritiation
of halotane-t can be discussed in the terms of the Brønsted and
Bjerrum eqn. (4) where γ(HOϪ), γ(ST) and γ‡ are the activity
to solvent basicity is mainly due to the hydroxide-ion-like
transition state which means that the activity coefficient ratio
γ(HOϪ)/γ‡ in eqn. (4) increases only slightly with the mole
fraction of organic component in the reaction mixture.20 In
aqueous DMI the rate coefficients even decrease slightly when
x(DMI) < 0.1. Thus it is evident that in these mixtures changes
in activity coefficients γ(ST) must be taken into account in
addition to those of charged species in accordance with kinetic
results in aqueous sulfolane.25
Estimation of equilibrium basicities from kinetic results
As seen above, changes in basicity of aqueous amides and ureas
can be discussed solely on the basis of rate coefficients for
detritiation reactions. If, however, comparison with equilibrium
basicities HϪ is desired, linear free energy correlations can be
used for this purpose. In the following we discuss only the cases
k = k0{γ(HOϪ)γ(ST)/γ‡}
(4)
where deviations from the linearity of log (k2/molϪ1 dm3 sϪ1
)
vs. x(amide or urea) are not observed. In addition, since the
detritiation of chloroform-t and halotane-t could be studied
only in mixtures with a very low x(urea or amide), HϪ values
were not extrapolated from those data with the exception
of chloroform-t in aqueous TMU. It has been observed
previously20 that in the detritiation of different carbon acids
in aqueous DMSO the logarithms of rate coefficients depend
linearly on both HϪ and x(DMSO) of the solvent.7 These linear
free energy correlations can be applied also when the solvent
system for the detritiation reaction is changed from aqueous
DMSO to aqueous urea or amide.
The detritiation of chloroform-t in aqueous TMU is
discussed first because equilibrium basicities of this system
are available for comparison.6 Using previous data20 for the
detritiation of chloroform-t in aqueous DMSO, relation (5)
is obtained. From the kinetic data presented in Table 2 for
the detritiation of chloroform-t in aqueous TMU the linear
correlation (6) was observed. Eqns. (5) and (6) give the relation
coefficients of hydroxide ion, substrate and transition state of
the reaction as compared with that in water. In both reactions
changes in activity coefficient of hydroxide ion are equal. In
addition, solvent effects on neutral substrate, chloroform or
halotane, can be excluded as compared to those of charged
species, hydroxide ion and transition state of the reaction.
The observed difference in the susceptibility to solvent basicity
must be due to the changes in the activity coefficients of the
transition states. It can be assumed that in the transition state
of the detritiation of halotane-t the degree of triton transfer is
higher than in the case of chloroform-t which means that in the
detritiation of halotane-t the transition state is less hydroxide-
ion-like than in the detritiation of chloroform-t. As changes
in activity coefficients of hydroxide ion and transition state of
the reaction with x(urea or amide) are parallel, they compen-
sate for each other partially depending on the structure of
the transition state. For halotane-t this compensation may be
lower and therefore a higher increase in rate coefficients is
expected.
log (k2/molϪ1 dm3 sϪ1) =
In the detritiation of methoxyflurane-t the slopes of the plots
log (k2/molϪ1 dm3
s
Ϫ1) vs. x(urea), 9.59 0.18 (DMI) and
(1.076 0.023)HϪ Ϫ (13.73 0.29) (5)
9.56 0.05 (DMPU) (Table 2) are equal within experimental
errors and thus the increase in basicity is also equal. This result
is well-founded because the cyclic ureas studied are structurally
closely related. Therefore, at first sight, it seems surprising that
in the detritiation of chloroform-t in aqueous DMI the slope
of the plot log (k2/molϪ1 dm3 sϪ1) vs. x(urea) is lower than in
aqueous DMPU: 14.60 (DMPU) and 11.4 (DMI). Kinetic
measurements for the detritiation of chloroform-t have, how-
ever, mainly been performed in the range x(DMI) = 0–0.1 where
the rate increase is exceptionally low also on grounds of kinetic
results for methoxyflurane so that the point x(DMI) = 0 has
been excluded in Fig. 2.
In the detritiation of the haloalkanes studied the suscepti-
bility of reaction rate to solvent basicity can be compared,
for instance, on the basis of kinetic data in aqueous DMPU
(Table 2). The slopes of the plots (18.3, 14.60 and 9.56) decrease
in the series halotane, chloroform and methoxyflurane. As
discussed above, this means that the hydroxide ion like structure
of the transition state increases in the same series.
To obtain additional information on the basicity of aqueous
acyclic and cyclic ureas, the detritiation of acetophenone-t was
studied in these solvent systems. As the susceptibility of this
reaction to solvent basicity is relatively low the measurements
could be extended even to a mole fraction of about 0.7. These
results also indicate that the increase in basicity is similar
in aqueous DMI and DMPU as the slopes of the plots
log (k2/molϪ1 dm3 sϪ1) vs. x(DMI or DMPU), 4.41 0.04 and
4.62 0.10, are equal within experimental errors. In aqueous
TMU the rate increase is slightly lower, as shown by the slope
4.043 0.004. All these slopes are, however, significantly lower
than in the detritiation of haloalkanes. This lower susceptibility
log (k2/molϪ1 dm3 sϪ1) =
(12.32 0.17)x(TMU) Ϫ (0.803 0.019) (6)
between x(TMU) and HϪ of the solvent. Kinetic measurements
for the detritiation of chloroform-t were performed only in
mixtures with x(TMU) < 0.188 and the equilibrium basicities6
have been determined after x(TMU) > 0.1844. Therefore, it
is reasonable to extrapolate the HϪ values from eqns. (5) and
(6) to the mixture with x(TMU) = 0.1844 where equilibrium
and kinetic measurements overlap. The HϪ value 14.11 (Table 3)
obtained by kinetic method is, within experimental errors, equal
to the equilibrium HϪ value 14.19 determined by Bowden and
Prasannan.6 Thus the method used in the present work seems to
give reliable results.
For the detritiation of acetophenone-t the linear relationship
(7) has been obtained in aqueous DMSO at 298.15 K.20 From
the kinetic results presented in Table 2 for the detritiation of
acetophenone-t in aqueous DMPU the linear correlation (8)
can be derived. The equality of eqns. (7) and (8) again gives a
log (k2/molϪ1 dm3 sϪ1) =
(0.433 0.004)HϪ Ϫ (7.48 0.06) (7)
log (k2/molϪ1 dm3 sϪ1) =
(4.62 0.10)x(DMPU) Ϫ (2.27 0.03) (8)
relation between HϪ and solvent composition. The HϪ values
shown in Table 3 were calculated in the mole fraction range
0–0.6 where kinetic measurements have been performed. These
J. Chem. Soc., Perkin Trans. 2, 1999, 169–174
173

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Table 3 HϪ values of aqueous TMU and DMPU calculated from
kinetic results using a linear free energy correlation. For comparison,
equilibrium basicities of aqueous TMU6 and DMSO7 are given in this
table. The concentration of tetramethylammonium hydroxide is 0.011
mol dmϪ3
278.15 K, the linear correlation (10) is obtained. Again, the left
sides in eqns. (9) and (10) are equal and thus a correlation
between HϪ and x(DMPU) is also obtained on the basis of
these kinetic results. The calculated HϪ values, given in Table 3,
are almost equal with those obtained from the kinetic data for
the detritiation of acetophenone-t. This agreement confirms the
applicability of the extrapolation method.
Organic
x(org.
component) HϪ
component
Method
TMU
0.1844
0.1844
14.19 Equilibrium6
14.11 Kinetic,
References
detritiation of chloroform-t
1 J. R. Jones, The Ionization of Carbon Acids, Academic Press,
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2 K. Bowden, Chem. Rev., 1966, 119.
3 D. Dolman and R. Stewart, Can. J. Chem., 1967, 45, 911.
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185.
DMPU
0.0
0.1
0.2
0.3
0.4
0.5
0.6
12.01 Kinetic,
13.08 detritiation of acetophenone-t
14.14
15.21
16.28
17.35
18.41
7 R. A. Cox and R. Stewart, J. Am. Chem. Soc., 1976, 98, 488.
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12 C. Okpala, A. Guiseppi-Elie and D. M. Maharajh, J. Chem. Eng.
Data, 1980, 25, 384.
DMPU
DMSO
0.0
0.1
0.2
0.3
0.4
0.5
12.08 Kinetic,
13.11 detritiation of methoxyflurane-t
14.13
15.15
16.17
17.20
0.0
0.1
0.2
0.3
0.4
0.5
0.6
12.04 Equilibrium7
13.04
14.24
15.37
16.34
17.44
18.49
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743.
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basicities, which refer to hydroxide ion concentration 0.011 mol
dmϪ3, can be compared with the equilibrium HϪ values of
aqueous DMSO7 with the same hydroxide ion concentration
(Table 3). The basicities of these two systems are almost equal.
21 N. S. Isaacs, Physical Organic Chemistry, Longman, Belfast, 1987,
pp. 484.
22 R. S. Brown, A. J. Bennett and H. Slebocka-Tilk, Acc. Chem. Res.,
1992, 25, 481.
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24 A. Kankaanperä, L. Oinonen and P. Salomaa, Acta Chem. Scand.,
Ser. A, 1977, 31, 551.
log (k2/molϪ1 dm3 sϪ1) =
(9.56 0.05) x(DMPU) Ϫ (4.094 0.015) (9)
log (k2/molϪ1 dm3 sϪ1) =
(0.936 0.012)HϪ Ϫ (15.41 0.27) (10)
25 A. Kankaanperä, P. Scharlin, M. Lahti, M.-L. Ekholm, L. Hautala
and K. Ranta, Acta Chem. Scand., 1998, 52, 1071.
26 A. Kankaanperä and M. Suovanen, unpublished results.
Kinetic results for the detritiation of methoxyflurane-t in
aqueous DMPU (Table 2) can also be used to extrapolate HϪ
values for this solvent system. At 278.15 K the free energy
correlation (9) was found. Based on the kinetic results26 for
the detritiation of methoxyflurane-t in aqueous DMSO at
Paper 8/08435A
174
J. Chem. Soc., Perkin Trans. 2, 1999, 169–174