Kinetics of hydrolysis of 1-benzoyl-1,2,4-triazole in aqueous solution as a function of temperature near the temperature of maximum density, and the isochoric controversy

Article abstract of DOI:10.1039/b210246c

At temperatures above and below the temperature of maximum density, TMD, for water at ambient pressure, pairs of temperatures exist at which the molar volumes of water are equal. First-order rate constants for the pH-independent hydrolysis of 1-benzoyl-1,2,4-triazole in aqueous solution at pairs of such isochoric temperatures show no unique features. Taken together with previously published kinetic data for the hydrolysis of a range of simple organic solutes in both water and D₂O near their respective TMDs, we conclude that special significance in the context of rates of chemical reactions in aqueous solutions should not be attached to the isochoric condition.

Full text of DOI:10.1039/b210246c

View Article Online / Journal Homepage / Table of Contents for this issue
Kinetics of hydrolysis of 1-benzoyl-1,2,4-triazole in aqueous
solution as a function of temperature near the temperature
of maximum density, and the isochoric controversy
a
b
b
c
Michael J. Blandamer,* Niklaas J. Buurma, Jan B. F. N. Engberts and João C. R. Reis
a
Department of Chemistry, University of Leicester, Leicester, UK LE1 7RH.
E-mail: mjb@le.ac.uk
b
Physical Organic Chemistry Unit, Stratingh Institute, University of Groningen, Nijenborgh 4,
747 AG Groningen, The Netherlands. E-mail: J.B.F.N.Engberts@chem.rug.nl
Departmento de Quimica e Bioquimica, Centro de Electroquímica e Cinética,
Faculdade de Ciências, Universidade de Lisboa, P-1749-016 Lisboa, Portugal.
E-mail: jreis@fc.ul.pt
9
c
Received 22nd October 2002, Accepted 20th December 2002
First published as an Advance Article on the web 28th January 2003
At temperatures above and below the temperature of maximum density, TMD, for water at ambient pressure, pairs of
temperatures exist at which the molar volumes of water are equal. First-order rate constants for the pH-independent
hydrolysis of 1-benzoyl-1,2,4-triazole in aqueous solution at pairs of such isochoric temperatures show no unique
features. Taken together with previously published kinetic data for the hydrolysis of a range of simple organic solutes
in both water and D O near their respective TMDs, we conclude that special significance in the context of rates of
2
chemical reactions in aqueous solutions should not be attached to the isochoric condition.
atures would yield pairs of isochoric rate constants. The sugges-
tion was not taken up. However, another controversy had arisen
Introduction
1
In 1935 Evans and Polanyi suggested that isochoric activation
concerning chemical reactions in aqueous solutions at low tem-
parameters for chemical reactions in aqueous solution might be
more mechanistically informative than conventional isobaric
activation parameters; i.e. [∂ln(k)/∂T ] rather than [∂ln(k)/∂T ]
16
peratures. Hills and Viana reported that rate constants for
hydrolysis of benzyl chloride (see above) in very dilute aqueous
solutions increased with increase in temperature above the
TMD of water and, significantly, increased with decrease in
v
p
where k is the rate constant for spontaneous chemical
2,3
reaction.
This proposal sparked enormous interest in
17
temperature below the TMD. Albery and Curran were unable
4
several subject areas including ionic transport and chemical
to confirm the pattern in rate constants reported by Hills and
5
equilibria in solution.
16
Viana.
With reference to chemical reactions in dilute aqueous solu-
tion the isochoric standard internal energy of activation ∆ U
is related to the isobaric standard enthalpy of activation ∆ H
Therefore we turned our attention to the suggestion made by
≠
o
v
o
p
8
Haak et al. concerning rates of reaction at isochoric pairs of
≠
temperatures. For this study to be meaningful, the studied reac-
tion had to satisfy at least three criteria:
at temperature T and the standard volume of activation
≠
o
*
*
∆
V
using equation (1) where αpl and κT l are, respectively,
(
(
(
i) the rate of chemical reaction can be precisely measured;
ii) the mechanism of reaction is well understood, and
iii) the rate of reaction is sensitive to temperature.
the isobaric expansibilities and isothermal compressibilities of
water (l).
These three criteria are satisfied by the spontaneous first-
order hydrolysis of 1-benzoyl-1,2,4-triazole in aqueous solu-
tion. The mechanism of hydrolysis involves water-catalysed
∆≠
U
o
= ∆ H
≠
o
Ϫ T [αpl/κT l]∆ V
*
*
≠
o
(1)
v
p
6
Baliga and Whalley noted that the dependence on solvent
mixture composition of ∆ U
for solvolysis of benzyl chloride in ethanol–water
mixtures at 298.15 K. A similar pattern was reported by Baliga
and Whalley for the hydrolysis of 2-chloro-2-methylpropane
in the same mixture at 273.15 K. Nevertheless the significance
of isochoric parameters in the context of both kinetic and
equilibrium parameters has been extensively debated.
Much of the debate centred on the isochoric condition and
the answer to the question “which volume is held constant?”.
With reference to equation (1) αpl and κT l depend on temper-
ature. Then the volume identified by subscript V on ∆ U
8
nucleophilic attack at the carbonyl group.
≠
o
v
is less complicated than that for
The dependence of rate constants on temperature near the
TMD of water is examined in terms of the dependence of the
molar volume of water on temperature reported by Kell and
≠
o
p
∆
H
7
18
19
Whalley and summarised by Kell. We show that the ‘iso-
choric condition’ does not reveal novel features concerning the
hydrolysis reaction in aqueous solution. We draw attention to
published kinetic data which support our conclusion but which,
in the debate over isochoric kinetics, were overlooked.
8
9
10–15
9
*
*
≠
o
v
is
Results
dependent on temperature. Further, the molar volumes of bin-
ary liquid mixtures at fixed T and p (cf. kinetic data in reference
First-order rate constants for the hydrolysis of 1-benzoyl-1,2,4-
triazole in aqueous solution (Scheme 1) are reported at 12
different temperatures between 276.15 and 278.15 K; Table 1.
The difference between the temperatures in the series is gener-
ally between 0.1 and 0.3 K. At three temperatures, 276.25,
276.45 and 277.25 K, repeat experiments yielded rate constants
which agree to better than ±0.42ꢀ. The resulting plot (Fig. 1)
≠
o
7
) depend on composition so that in the context of ∆ U
v
, the
volume identified by subscript V is not constant across a range
8
of solvent mixtures. In 1985, Haak et al. noted that above and
below the temperature of maximum density (TMD) of water
there are pairs of temperatures where the molar volumes of
water at ambient pressure are equal. Hence rates of reaction for
chemical reactions in dilute aqueous solution at such temper-
Ϫ1
Ϫ1
Ϫ1
of (k/s ) against (T /K) is linear. The dependences of (k/s )
7
20
O r g . B i o m o l . C h e m . , 2 0 0 3 , 1, 7 2 0 – 7 2 3
T h i s j o u r n a l i s © T h e R o y a l S o c i e t y o f C h e m i s t r y 2 0 0 3

View Article Online
Table 1 First-order rate constants for the spontaneous hydrolysis
of 1-benzoyl-1,2,4-triazole as a function of temperature in aqueous
a
solution at temperatures close to the TMD
Ϫ4 Ϫ1
T /K
k/10
s
2
2
2
2
2
2
2
2
2
2
2
2
76.15
3.91
76.25
76.45
76.65
76.95
77.05
77.25
77.35
77.45
77.65
77.85
78.15
3.92; 3.93; 3.95
3.93; 3.96; 3.97
4.04
4.11
4.16; 4.26
4.20; 4.22; 4.25
4.29
4.29; 4.32
4.37
4.51
4.59
Fig. 2 Dependence of rate constant for the hydrolysis of 1-benzoyl-
,2,4-triazole and molar volume of the aqueous solution on
temperature. The dotted line marks the TMD.
1
a
Ϫ4 Ϫ1
At 298.15 K, k = 21.2 × 10
s .
Scheme 1
Fig. 3 Dependence of rate constant for the hydrolysis of 1-benzoyl-
2
1
2
(
,2,4-triazole on (T -TMD) over the temperature range 276.15 ≤ T /K ≤
78.15. The solid line shows the calculated rate constant using equation
1) with θ = 276.85 K, k(θ) = 4.106 × 10
Ϫ4 Ϫ1
≠
o
s
and ∆ H (θ) = 49.8 kJ
Ϫ1
mol
.
temperature range covered by the results in Table 1 are too
small and so the experimental data were analysed using two
≠
o
parameters k(θ) and ∆ H (θ). Hence (Fig. 3) with θ = 276.85 K,
Ϫ4 Ϫ1
≠
o
k(θ)/10
1
s
equals 4.106 ± 0.006 and ∆ H (θ) equals 49.8 ±
Ϫ1
21
.6 kJ mol , in agreement with previous results.
Discussion
The careful and detailed investigation reported here was
prompted by the idea that isochoric rate constants for reactions
in solution might contribute to the problem of understanding
the role of solvents in controlling activation parameters for
reactions in solution. In general terms, activation parameters
for a given reaction in aqueous solution are determined in part
by solute–solvent interactions between both initial and trans-
ition states and neighbouring solvent molecules. Solvent–solute
intermolecular distances and solvent organisation around
solutes are functions of temperature. Therefore the case was
Fig. 1 Dependence of ln(k/T) on reciprocal temperature for the
hydrolysis of 1-benzoyl-1,2,4-triazole in aqueous solution. The dotted
line marks the TMD. The inset plot shows the dependence in relation to
2
1
the previously reported rate constant at 298.2 K.
*
l
and molar volume V
(l) on temperature are compared in Fig. 2.
These plots show no novel features as the temperature increases
from below to above the TMD which is estimated to be 276.85
K; see Experimental section. The dependence of rate constants
20
on temperature was analysed using equation (2) where θ is a
1
reference temperature.
made that by keeping the volume constant these distances
remain constant when the temperature is changed. The expect-
ation was therefore that as one increases the temperature the
influence of changes in solute–solvent distances are minimal
under isochoric conditions. Hence by ‘removing’ this influence
one can direct attention to the changes taking place in the
reactants on going from initial to the transitions states. The
(
2)
9
question which ‘volume is constant?’ was raised.
The importance of the experimental results presented here is
that they illustrate how
The analysis yields estimates of rate constant at k(θ), the
≠
o
standard enthalpy of activation ∆ H (θ) and the standard
≠
o
isobaric heat capacity of activation ∆ C
p
(θ) at reference tem-
≠
o
perature θ. A statistically significant estimate of ∆ C
p
(θ)
requires kinetic data over an extended temperature range. The
O r g . B i o m o l . C h e m . , 2 0 0 3 , 1, 7 2 0 – 7 2 3
721

View Article Online
can be calculated with respect to isochoric conditions. Interest-
ingly there are pairs of isochoric data points on the plot shown
in Fig. 1. All the data points, both above and below the TMD,
show that the rate of reaction is insensitive to the presence
of the TMD. In fact, this pattern is shown by published
kinetic data for hydrolysis reactions in aqueous solutions
as reported by Robertson and co-workers. These data sets
also include kinetic data for solvolysis of a range of com-
The cell compartment was fitted with a hole in the cover
through which wires and tubing could be led into the sealed cell
3
compartment. Through this hole 1.5 cm of solution in the
sample cell was withdrawn. Between 4 and 8 µL of a stock
3
solution containing 5 mg of 1-benzoyl-1,2,4-triazole in 1 cm
cyanomethane was injected into the cuvette. The solution in the
cuvette was withdrawn and re-injected several times. Finally the
solution was injected into the cuvette very slowly in order to
prevent the formation of air bubbles. After the reaction had
been initiated the temperature of the cell was monitored. Only
after the temperature had returned to the required temperature
were data points recorded of absorbance and time. During a
given kinetic run the absorbances were measured during a brief
exposure (1 second in each minute) of the cell to the incident
UV–visible beam of light. Poor kinetic data resulted if the
solution was subjected to continuous radiation.
22–30
22,23,25,31
pounds in water
and in D O,
the TMD of the latter
2
being 284.35 K. The compounds included methyl trifluoro-
2
2
ethanoate, succinic2 and phthalic anhydride,
dimethyl-
3
sulfamoyl chloride,
diethyl and methyl(ethyl)sulfamoyl
24
25
chloride, methanesulfonyl and benzenesulfonyl chloride,
2
7
methyl chlorosulfate and ethyl chlorosulfate,
2-chloro-
ethyl(methyl) sulfide, methyl methanesulfonate and 4-substi-
26
28
29
30
tuted benzenesulfonyl chlorides. Moelwyn-Hughes et al.
reported rate constants for the solvolysis of 2-chloro-2 methyl-
propane in aqueous solution at intervals of 1 K from 273.15
to 298.15 K. Novel features in the kinetic data can neither be
discerned from isobaric pairs of temperatures nor from around
the TMD.
The TMD of the solutions used in the experiments was
slightly lower than the TMD of water at ambient pressure. The
lowering of the TMD was caused by the use of cyanomethane
in the reaction medium injected as stock solution of the
substrate. In a typical experiment 6 ± 2 µl of a stock solution in
3
We note that most of these data were published before the
cyanomethane was injected together with 2.75 cm of water (l)
16
claim by Hills and Viana was published concerning ‘negative
activation energies’. In other words published data was avail-
able to show that the TMD of water plays no real part in
determining rate constants for solvolytic reactions. In fact,
the interest of the senior author Robertson on many of these
resulting in a 0.126 ± 0.040 molꢀ aqueous solution of
34
cyanomethane. The TMD of water (l) is lowered by 0.28 ±
0.09 K to 276.85 ± 0.09 K. In the following we assume that the
molar volume of water ϩ cyanomethane binary liquid mixture
at ambient pressure is a quadratic function of the temperature
around the TMD. Nevertheless the conclusions presented
here are not affected by setting the TMD to either 276.85 or
277.13 K.
22–31
publications
was directed towards determining the isobaric
heat capacity of activation for solvolytic reactions in water and
26
D O; e.g. 2-chloro-2-methylpropane (aq).
2
These comments do not, of course, detract from the
1
importance of the original proposal by Evans and Polanyi. We
simply suggest that in many cases where isochoric activation
parameters have been reported the proper meaning of the term
Acknowledgements
MJB thanks the receipt of an RSC Journals Grant enabling
a visit to the University of Lisbon. MJB also thanks the
University of Lisbon for their hospitality. JCRR is grateful to
Fundação para a Ciência e a Tecnologia for a research grant.
‘isochoric’ has not been recognised. However, we have
shown that in the important case of aqueous solutions no
unique feature emerges in the kinetic data for hydrolysis
reactions in the region of the TMD.
References
Experimental
1
M. G. Evans and M. Polanyi, Trans. Faraday Soc., 1935, 31,
75.
M. G. Evans and M. Polyani, Trans Faraday Soc., 1937, 33, 448.
8
Materials
2
1
-Benzoyl-1,2,4-triazole was synthesised according to published
3 D. M. Newitt and A. Wassermann, J. Chem. Soc., 1940, 735.
4 See, for example, G. J. Hills, P. J. Ovenden and D. R. Whitehouse,
Discuss. Faraday Soc., 1965, 39, 207 and references therein.
21,32,33
procedures.
distillation unit.
Water was distilled twice from an all-quartz
5
D. A. Lown, H. R. Thirsk and Lord Wynne-Jones, Trans. Faraday
Soc., 1970, 66, 51.
B. T. Baliga and E. Whalley, J. Phys. Chem., 1967, 71, 1166.
B. T. Baliga and E. Whalley, Can. J. Chem., 1970, 48, 528.
J. R. Haak, J. B. F. N. Engberts and M. J. Blandamer, J. Am. Chem.
Soc., 1985, 107, 6031.
Kinetics
6
7
8
The spontaneous hydrolysis of 1-benzoyl-1,2,4-triazole was
followed using a Shimadzu Diode-array spectrophotometer,
absorbances being recorded between 200 and 400 nm. The cell
compartment of the spectrophotometer was thermostatted to a
preset temperature using a Haake thermostatting unit equipped
with a Pt100 electrode for direct temperature control of the cell
block in the spectrophotometer. The temperature was checked
using a copper/constantin thermocouple and found to deviate
by no more than 0.05 K from the preset temperature during
the kinetic experiments. Experience showed that control of
temperature was improved by using ethanol rather than water
as a bath liquid. We attribute the problems with using water to
the formation of ice in the refrigeration unit.
9
M. J. Blandamer, J. Burgess, B. Clark and J. M. W. Scott, J. Chem.
Soc., Faraday Trans. 1, 1984, 80, 3359.
10 P. G. Wright, J. Chem. Soc., Faraday Trans. 1, 1986, 82, 2557.
1 L. M. P. C. Albuquerque and J. C. R. Reis, J. Chem. Soc., Faraday
Trans. 1, 1989, 85, 207.
1
1
1
2 E. Whalley, J. Chem. Soc., Faraday Trans. 1, 1987, 83, 2901.
3 M. J. Blandamer, J. Burgess, H. J. Cowles, I. M. Horn, J. B. F. N.
Engberts, S. A. Galema and C. D. Hubbard, J. Chem. Soc., Faraday
Trans. 1, 1989, 85, 3733.
14 J. B. F. N. Engberts, J. R. Haak, M. J. Blandamer, J. Burgess and
H. J. Cowles, J. Chem. Soc., Perkin Trans. 2, 1990, 1059.
5 L. M. P. C. Albuquerque and J. C. R. Reis, J. Chem. Soc., Faraday
Trans., 1991, 87, 1553.
1
A quartz cuvette, path length 1cm, contained approx.
1
1
6 G. Hills and C. A. N. Viana, Nature, 1971, 229, 194.
7 W. J. Albery and J. S. Curran, J. Chem. Soc., Chem. Commun., 1972,
425.
3
2
.75 cm of water for which the pH had been adjusted to 3.8 ±
0
.3 using HCl (aq). The cuvette had a stopper fitted with two
small holes. The cuvette was thermostatted overnight. The cell
compartment was sealed with cling film so that the air-flow in
the spectrophotometer did not affect the temperature of the cell
compartment. The cell compartment was continually flushed
with dry pre-cooled air. The latter was important in order to
avoid condensation of water vapour on the cell walls.
18 G. S. Kell and E. Whalley, Philos. Trans. R. Soc. London, A, 1965,
2
58, 565.
1
2
9 G. S. Kell, J. Chem. Eng. Data, 1967, 12, 66.
0 E. C. W. Clarke and D. N. Glew, Trans. Faraday Soc., 1966, 62,
5
39.
2
1 W. Karzijn and J. B. F. N. Engberts, Tetrahedron Lett., 1978, 19,
1787.
7
22
O r g . B i o m o l . C h e m . , 2 0 0 3 , 1, 7 2 0 – 7 2 3

View Article Online
2
2
2 B. Rossall and R. E. Robertson, Can. J. Chem., 1975, 53, 869.
28 M. J. Blandamer, H. S. Golinkin and R. E. Robertson, J. Am. Chem.
3 E. C. F. Ko and R. E. Robertson, J. Am. Chem. Soc., 1972, 94,
Soc., 1969, 91, 2678.
5
73.
29 P. W. C. Barnard and R. E. Robertson, Can. J. Chem., 1961, 39, 881.
30 R. E. Robertson and B. Rossall, Can. J. Chem., 1971, 49, 1441.
31 B. Rossall and R. E. Robertson, Can. J. Chem., 1971, 49, 1451.
32 H. A. Staab, M. Luking and F. H. Durr, Chem. Ber., 1962, 95, 1275.
33 H. J. Mooij, J. B. F. F. N. Engberts and M. Charton, Recl. Trav.
Chim. Pays-Bas, 1988, 107, 185.
2
2
4 E. C. F. Ko and R. E. Robertson, Can. J. Chem., 1972, 50, 946.
5 R. E. Robertson, B. Rossall, S. E. Sugamori and L. Treindl, Can. J.
Chem., 1969, 47, 4199.
6 E. A. Moelwyn-Hughes, R. E. Robertson and S. Sugamori, J. Chem.
Soc., 1965, 1965.
2
2
7 E. C. F. Ko and R. E. Robertson, Can. J. Chem., 1972, 50, 434.
34 G. Wada and S. Umeda, Bull. Chem. Soc. Jpn., 1962, 35, 1797.
O r g . B i o m o l . C h e m . , 2 0 0 3 , 1, 7 2 0 – 7 2 3
723