Spontaneous hydrolysis of ethyl formate: Isobaric activation parameters

Article abstract of DOI:10.1002/(SICI)1097-4601(2000)32:1<67::AID-JCK8>3.0.CO;2-M

Ethyl formate undergoes spontaneous autocatalytic hydrolysis via water catalyzed (neutral), as well as hydrogen-ion catalyzed mechanisms. The activation parameters for the neutral reaction are ΔH = 91±8 kJ/mol and ΔS = -48±8 J K^-1/mol. This result is in contrast to the values reported in the hydrolyses of the more polar activated haloesters. The specific rate of neutral hydrolysis of ethyl acetate can also be predicted.

Full text of DOI:10.1002/(SICI)1097-4601(2000)32:1<67::AID-JCK8>3.0.CO;2-M

Spontaneous Hydrolysis of
Ethyl Formate: Isobaric
Activation Parameters
J. F. MATA-SEGREDA
School of Chemistry, University of Costa Rica, 2060 Costa Rica
Received 3 May 1999; accepted 17 September 1999
ABSTRACT: Ethyl formate undergoes spontaneous autocatalytic hydrolysis via water catalyzed
(neutral), as well as hydrogen-ion catalyzed mechanisms. The activation parameters for the
Ϫ
1
neutral reaction are ⌬H‡ ϭ 91 Ϯ 8 kJ/mol and ⌬S‡ ϭ Ϫ48 Ϯ 8 J K /mol. This result is in
contrast to the values reported in the hydrolyses of the more polar activated haloesters. The
specific rate of neutral hydrolysis of ethyl acetate can also be predicted. ᭧ 2000 John Wiley &
Sons, Inc. Int J Chem Kinet 32: 67–71, 2000
INTRODUCTION
fer of protons. The commonly accepted minimal struc-
ture for the TS complex is the following:
Esters with electron-withdrawing substituents in either
the acyl or alkyl moieties undergo hydrolysis in the
absence of added acid or base. This is a special case
of general base catalysis by water itself, the so-called
neutral or water-catalyzed hydrolysis of esters [1–3].
The neutral hydrolyses of carboxylic anhydrides
and activated amides and esters take place with: (i)
high solvent deuterium kinetic isotope effects in the
range from 2 to 4, (ii) large rate-decreasing salt effects,
(iii) very negative entropies of activation in the range
O
H
H
O
H
O
H
C
ORЈ
R
in which a second H2O molecule acts as the general
base, abstracting a proton from the nucleophilic H2O.
There is experimental [16,18] and theoretical [21,22]
basis to consider a cyclic atomic array made of one
RCO-X molecule and three water molecules, as a more
complete picture of the TS complex.
Esters of formic acid also undergo neutral hydrol-
ysis, at faster rates compared with the corresponding
acetates. The reason for this rate differential is doubt-
lessly the result of decreased steric retardation and
electron release. Furthermore, hyperconjugation ex-
erted by the methyl group in acetate esters is absent in
formates. Thus, the electrophilic character of the car-
boxyl carbon atom in formates is greater, resulting in
their increased hydrolytic fragility.
1
from –0.05 to Ϫ0.2 kJ KϪ /mol, and (iv) large nega-
tive volumes of activation from –9 to –22 cm3/mol.
The low entropies are counterbalanced by low en-
thalpies of activation (from 25 kJ/mol to 80 kJ/mol),
these latter being the result of the high electrophilic
character of the carboxyl carbon atom bound to elec-
tron-withdrawing substituents.
These kinetic features suggest transition state (TS)
structures in which many solvent molecules are im-
mobilized and subjected to electrostriction associated
with the solvation of developing charges and the trans-
Ethyl formate undergoes this spontaneous process
through two parallel mechanistic pathways: water ca-
talysis and specific hydrogenion catalysis. The second
Correspondence to: J. F. Mata-Segreda
᭧ 2000 John Wiley & Sons, Inc.
CCC 0538-8066/00/010067-05

68
MATA-SEGREDA
process arises from the accumulation of hydrogen ion
in the medium, derived from the ionization of the rel-
atively strong formic acid (pKa ϭ 3.751 at 25ЊC)
formed as one of the reaction products. (See equations
that follow.)
ln{[HCO2H]ϱ /([HCO2H]ϱ Ϫ [HCO2H])}
t
ϭ k1t ϩ k2 ͐0 [Hϩ]dt (2)
From the point of view of statistical analysis, Eq.
(2) is a multilinear regression equation, where time
and the cumulative hydrogenion concentration (͐ [Hϩ]
dt) are the independent variables. Yamasaki et al. have
published a method for the calculation of rate con-
stants from “linear” rate equations similar to the case
at hand [4]. However, the much lower number of mea-
surements allowed by the experimental procedure and
the uncertainty associated with the cumulative hydro-
genion concentration restrict the quality of the regres-
sion coefficients from this numerical method (see
later).
HCO2Et ϩ H2O !: HCO2H ϩ EtOH
Ϫ
HCO2H ϭ Hϩ ϩ HCO2
ϩ
HCO2Et ϩ H2O 9
d[HCO2H]/dt ϭ k1 [HCO2Et]
ϩ k2 [Hϩ] [HCO2Et] (1)
9
H: HCO2H ϩ EtOH
During the early stages of the reaction, the slower
water-catalyzed hydrolysis route is the main kinetic
contributor. As Hϩ accumulates, the second mecha-
nism becomes more significant. The reaction progress
curve of this process is thus expected to be of sig-
moidal shape. This kinetic scheme is similar to the
autocatalyzed reduction of aqueous toluidine blue by
phenyl hydrazine reported by Jonnalagadda and Nattar
[23].
Thus, the following “chemical” approach was
taken. In separate runs, the values of k2 were deter-
mined in the presence of dilute HCl (0.05 – 0.10 mol/
dm3). The second-order rate constants were calculated
by dividing the observed pseudo-first-order specific
rates by the corresponding values of [HCl].
The k1 values were determined as the slopes in Eq.
(3b):
It can be safely hypothesized that the water-cata-
lyzed hydrolysis of alkyl formates should involve a
much higher ⌬H‡ and a less negative ⌬S‡ than the
values observed for haloesters and anhydrides, due to
the absence of polar substituents.
g(t) ϭ ln[HCO2H]ϱ Ϫ ([HCO2H]ϱ
t
Ϫ [HCO2H]) Ϫ k2 ͐0 [Hϩ]dt
(3a)
(3b)
g(t) ϭ k1t
This account presents the isobaric activation pa-
rameters, at normal pressure, for the water-catalyzed
hydrolysis of ethyl formate and a discussion of the
kinetic features of this autocatalytic reaction.
The cumulative hydrogenion concentration was ob-
tained by fitting the [Hϩ] – time data pairs to a cubic
polynomion and subsequent analytical integration.
The cubic fit was empirically found to provide the
smallest variance, in all cases. The [Hϩ] values were
calculated from the pKa of formic acid, which is al-
most temperature independent in the thermal range
(e.g., 30–50ЊC) of this study [1 and tables in CRC
Handbookof Chemistry and Physics ].
Materials
Ethyl formate (Eastman) was distilled from CaO to
eliminate both free EtOH and HCO2H formed on
standing on the shelf.
RESULTS AND DISCUSSION
Activation Parameters
Kinetic Measurements and Data Treatment
The experiments were done by titration of HCO2H
formed with NaOH/bromocresol green (pH range 3.8–
5.4), in runs made by mixing 0.10 dm3 of distilled
water with 3.0 cm3 of the ester. The amount of sam-
pling was always 3.00 cm3 aliquots, titrated against
0.1 mol/dm3 NaOH. This procedure ensures a maxi-
mum error of 1% in the early titrations. In all experi-
ments, more than ten data points were collected (90%
or more of the total extent of reaction).
The acid-catalyzed hydrolysis of HCO2Et shows no
extraordinary features. The second-order specific rates
in Table I yield the values of 60 Ϯ 2 kJ/mol for the
enthalpy barrier and an entropy of activation of
1
–89 Ϯ 7 J KϪ /mol. This result agrees with the results
1
from the literature [2]: 60.5 kJ/mol and –86 J KϪ /
mol.
Figure 1 shows the reaction progress curves for the
spontaneous hydrolysis of HCO2Et at different tem-
peratures. The sigmoidal curves tend to become more
hyperbolical as the temperature increases. This behav-
The raw data were fitted to the integrated form of
Eq. (1):

SPONTANEOUS HYDROLYSIS OF ETHYL FORMATE
69
Figure 1 Reaction progress curves for the spontaneous hydrolysis of ethyl formate in water solvent.
ior agrees qualitatively with the expectation that the
water-catalyzed reaction must proceed with a high en-
thalpy of activation, as stated in the hypothesis. Re-
actions with high-energy barriers are greatly acceler-
ated by temperature, compared with reactions of lower
⌬H‡. At higher temperatures, faster reaction via the
water-catalyzed mechanism results in earlier accu-
mulation of Hϩ in the medium and the earlier predom-
inance of the faster Hϩ-catalyzed reaction.
Table I gives the k1 and activation parameters for
the water-catalyzed hydrolysis of HCO2Et. The un-
certainty limits in k1 are in the range from 3% up to
13%, higher than what is customary for solution ki-
netics. The algorithm used involves the unavoidable
build up and spread of errors in the data trios: time,
[HCO2H] and ͐ [Hϩ] dt. Nevertheless, the plot ln(k1/
T) vs . 1/T gives an acceptable linear correlation y ϭ
(18 Ϯ 3) – (11 Ϯ 1) ϫ 103 x, with r ϭ 0.98, p Ͻ
0.005.
Quantum mechanical calculations carried out by
Wolfe et al. [22] yielded a value of 66 kJ/mol for the
water-catalyzed hydration of formaldehyde in water
solvent. This substrate has a more electrophilic car-
bonyl carbon atom than the ester carbonyl. As it
should be expected, this enthalpy barrier is lower than
the experimental value of ⌬H‡ ϭ 91 Ϯ 8 kJ/mol for
the water catalyzed hydrolysis of HCO2Et.
In contrast, the entropies of activation for these two
1
reactions must be similar; a value of –51 J KϪ /mol
1
for the hydration of CH2O and –48 Ϯ 8 J KϪ /mol for
the hydrolysis of HCO2Et.
Multiple simultaneous fitting of the data to Eq. (2)
gives regression coefficients that agree with the spe-
cific rates in Table I, but only within an order of mag-
nitude. Nevertheless, analysis of variance (ANOVA)
indicates that both “independent variables” in Eq.
(2) are statistically significant at p ϭ 0.01. The correla-
tion equation for the multifit rates (x) and the chem
rates (y) obtained from Eq. (3 b) is y ϭ (1.2 Ϯ 0.7) ϩ
(0.8 Ϯ 0.2) x, r ϭ 0.91, and p ϭ 0.95 (multifit values
not shown).
The failure in the attempt to produce rate constants
from the linear treatment of Eq. (2) is due to the rel-
ative scarcity of data points (10–15 points) and the
higher uncertainty of the ͐ [Hϩ] dt values, compared
with time alone. Both variables are weighted equally,
Table I Isobaric Activation Parameters at Normal
Pressure for the Spontaneous Hydrolysis of Ethyl
Formate in Water Solvent
Ϫ
1
Ϫ
Ϫ
Temperature/K
105 k1/s
103 k2/s 1 mol 1 dm3
303.2
310.2
312.8
318.4
323.0
0.78 Ϯ 0.06
2.1 Ϯ 0.1
3.1 Ϯ 0.4
5.8 Ϯ 0.2
6.7 Ϯ 0.3
91 Ϯ 8
5.01 Ϯ 0.03
5.23 Ϯ 0.05
9.9 Ϯ 0.1
17 Ϯ 1
—
60 Ϯ 3
Ϫ
1
⌬H‡/kJ mol
Ϫ
Ϫ1
⌬S‡/J K 1 mol
Ϫ48 Ϯ 8
Ϫ89 Ϯ 7

70
MATA-SEGREDA
The hydrolysis of HCO2Et at 30ЊC was studied in
the presence of 1.0 mol/dm3 KCl to assess the salt
effect. The result was 106 k1 ϭ 6.0 Ϯ 0.7 sϪ1 and 103
k2 ϭ 5.56 Ϯ 0.07 sϪ molϪ dm3. The effect of 1.0
mol/dm3 ionic strength is a 30% depression in k1. The
magnitude of the effect is similar to that observed in
the neutral hydrolyses of CHF2CO2Et [7], CF3COSEt
[8], and acetic anhydride [9]. Thus, the k1 value of
Stefanidis and Jencks corrected for the inhibitory salt
1
1
effect should be 2.7 ϫ 10Ϫ7 sϪ .
1
Figure 2 Isergonic plot for the water-catalyzed hydrolyses
of haloesters, carboxylic anhydrides and amides in water
solvent.
According to the result of this work, the k1 extra-
polated down to 25ЊC is 1.9 ϫ 10Ϫ sϪ , a value that
is seven times larger than the one observed by the au-
thors cited. The mismatch is equivalent to a deviation
of 4% in ⌬G‡, an amount lower than the experimental
uncertainty associated to the activation parameters.
It is then concluded that, in spite of the degree of
uncertainty in the k1 values and the corresponding ac-
tivation parameters, the experiment supports the stated
hypothesis.
6
1
though their degrees of accuracy are not the same. No
further discussion is devoted to this calculation
method, because it produced unreliable results.
As noted earlier, an enthalpy-entropy correlation
can be observed for the neutral hydrolysis of esters,
amides, and carboxylic anhydrides. There is experi-
mental evidence indicating that for all these substrates,
solvent and TS complexes interact essentially by only
a single physical mechanism. Figure 2 shows activa-
tion figures for a group of haloesters, amides, and car-
boxylic anhydrides [2,3,14–17]; the data pair found
for ethyl formate in this work fits well within the set
of data. The isokinetic temperature observed is 276 Ϯ
30 K (r ϭ 0.89, p Ͻ 0.005). This value is in the range
associated with solvation effects in aqueous media
(isokinetic temperatures around 300 K), which arise
from both electronic as well as steric interactions [19].
It is interesting to note that the isokinetic temper-
ature for the acid-catalyzed hydrolysis of alkyl acetates
and thioacetates is 330 K, in 62% aqueous acetone
[19]. The similarity between the isokinetic tempera-
tures for the neutral and acid-catalyzed reactions is
somewhat to be expected because in the words of
Kirby [2]: “Since a protonated ester may be regarded
as a reactive ester bearing a strongly electron-with-
drawing substituent on the carbonyl oxygen atom it is
not unreasonable to suppose that they will react with
water similarly to esters with strongly electron-with-
drawing substituents in the other two positions.”
The water-catalyzed hydrolysis of ethyl formate
was studied, among other aryl and alkyl formates, by
Stefanidis and Jencks [5]. These authors reported k1 ϭ
Extrapolation to Ethyl Acetate
The dynamic features of TS complexes contribute an
important aspect of mechanistic chemistry, not only
for the study of simple reactions, but also for the un-
derstanding of enzymic reactions, be it aimed at basic
bioorganic studies or design of drugs. Simple models,
such as the present, are useful in setting limits to the
intrinsic stereodynamic features of chemical reactions
in solution.
The quantitative features of the neutral hydrolysis
of HCO2Et indicate that protonation of the neutral TS
complex decreases the enthalpy of activation by about
30 kJ/mol, and that the higher degree of solvent elec-
trostriction associated to the protonated complex is ev-
idenced by a ⌬S‡ value 40 J KϪ /mol more negative
1
than the value for the neutral hydrolytic reaction.
The thermodynamic activation requirements should
be much more stringent for the water-catalyzed hy-
drolyses of alkyl carboxylates other than for formates
or esters bearing electron-withdrawing substituents
vicinal to the carboxyl moiety. One obvious case is
ethyl acetate. A half-life of 89 years was estimated by
Skrabal and Zahorka in 1929 [cited in reference 2] for
this reaction at 25ЊC in water solvent (k1 ϭ 10Ϫ9.6 sϪ ).
1
The ⌬H‡ for the hydrogen-ion catalyzed hydrolysis
of CH3CO2Et is equal to 68 kJ/mol [10]. For the neu-
tral reaction, the corresponding ⌬H‡ can be expected
to be just around 0.10 MJ/mol.
7
1
2.1 ϫ 10Ϫ sϪ for HCO2Et at 25ЊC in 1.0 mol/dm3
KCl. This value agrees (within a factor of 2) with
7
1
4.2 ϫ 10Ϫ sϪ , found by Sawyer and Kirsch [6] for
the hydrolysis of HCO2Me under the same conditions.
The k1’s obtained in this work seem much higher, at
first glance, than what would be expected from the
results of these authors.
⌬S‡ can be estimated by following this rationale.
The entropies of the TS complexes in the alkaline hy-
drolyses of acetates are more negative than for formate
1
hydrolysis by about 21 J KϪ /mol [11]. This result has

SPONTANEOUS HYDROLYSIS OF ETHYL FORMATE
71
2. Kirby, J. In: Comprehensive Chemical Kinetics, Vol.
10; Bamford, C. H.; Tipper, C. F. H. Eds.; Elsevier:
Amsterdam, 1972.
3. Jencks, W. P. Catalysis in Chemistry and Enzymology;
Dover: New York, 1987.
4. Yamasaki, Y.; Watanabe, A.; Kakuda, T.; Tokue, I. Int
J Chem Kinet 1998, 30, 47.
5. Stefanidis, D.; Jencks, W. P. J Am Chem Soc 1993, 115,
6045.
6. Sawyer, C. B.; Kirsch, J. F. J Am Chem Soc 1973, 95,
7375.
7. Jencks, W. P.; Carriuolo, J. J Am Chem Soc 1961, 83,
1743.
8. Fedor, L. R.; Bruice, T. C. J Am Chem Soc 1965, 87,
4138.
9. Bunton, C. A.; Fuller, N. A.; Perry, S. G.; Pitman, I. H.
J Chem Soc 1962, 4478.
10. The ⌬H‡ for the acid-catalyzed hydrolysis of ethyl ac-
etate was checked in this laboratory, and a value of 68
kJ/mol resulted, in good agreement with the value of 66
kJ/mol in the NBS Tables for Chemical Kinetics.
11. Thornton, E. K.; Thornton, E. R. In: Transition States
of Biochemical Processes; Gandour, R. D.; Schowen,
R. L. Eds.; Plenum Press: New York, 1978; p 34.
12. Frost, A. A.; Pearson, R. G. Kinetics and Mechanism,
2nd ed.; John Wiley & Sons: New York, 1963; pp 327–
334.
been interpreted in terms of freezing of the rotational
degree of freedom of the methyl group in acetates.
This rigidity may be due to a greater degree of solvent
immobilization for acetate TS as compared with for-
mate TS. Basically, the same result is observed when
comparing the saponification of ethyl acetate and ␥-
butyrolactone [12], because there is no rotating methyl
group in the lactone ring.
Use of Eyring’s equation with ⌬H‡ ϭ 105 Ϯ 5 kJ/
1
1
mol and ⌬S‡ ϭ Ϫ69 J KϪ /mol (Ϫ48 J KϪ /mol – 21
J KϪ /mol) yields k1 ϭ 10Ϫ
sϪ , a number in
1
9.2
Ϯ
0.8
1
agreement with the estimate of Skrabal.
A similar result is obtained on consideration of the
effect of replacing a hydrogen atom by a methyl group
in the hydration of H2CO and CH3CHO, and of
HCO2Me and CH3CO2Me to give the corresponding
tetrahedral gem-diols. This structural effect has been
estimated to be ␦⌬GЊ Ϸ ␦⌬G‡ ϭ 20 kJ/mol [13,20–
22]. The specific rate for the neutral hydrolysis of
CH3CO2Et at 25ЊC can then be calculated by correct-
ing the value for HCO2Et by the magnitude of the
stereoelectronic effect due to the H/CH3 replacement:
1
k1(CH3CO2Et) ϭ 1.9 ϫ 10Ϫ6 sϪ
ϫ exp (Ϫ20 ϫ 103/RT)
13. Guthrie, J. P. J Am Chem Soc 1973, 95, 6999.
14. Bunton, C. A. Fendler, J. H. J Am Chem Soc 1965, 30,
1365.
1
k1(CH3CO2Et) ϭ 10Ϫ9.2 sϪ
15. Talbot, R. J. E. In: Comprehensive Chemical Kinetics,
Vol. 10; Bamford, C. H.; Tipper, C. F. H. Eds.; Elsevier:
Amsterdam, 1972.
16a. Davis, K. R.; Hogg, J. L. J Org Chem 1983, 48, 1041;
16b. Gopalakrishnan, G.; Hogg, J. L. J Org Chem
1984, 49, 3161.
The results obtained for the spontaneous catalyzed
hydrolysis of HCO2Et are consistent with the behavior
of the much less reactive acetate analogue. This con-
sistency is an indirect way to check on the validity of
the rate constants and activation parameters deter-
mined in this study.
17. Haak, J. R.; Engberts, J. B. F. N. J Am Chem Soc 1986,
108, 1705.
18. Venkatasubban, K. S.; Bush, M.; Ross, E.; Schultz, M.;
Garza, O. J Org Chem 1998, 63, 6115.
The author wishes to express his gratitude to Prof. William
P. Jencks and to Dr. Dimitros Stefanidis for their collegial
assistance, concerning the magnitude of the rate constants
in this work.
19a. Leffler, J. E.; Grunwald, E. Rates and Equilibria of
Organic Reactions; John Wiley & Sons: New York,
1963; 19b. Schowen, R. L. J Pharm Sci 1967, 56, 931.
20. Critchlow J. E. J Chem Soc Faraday Trans 1 1972, 68,
1774.
21. Williams, I. H.; Spangler, D.; Femec, D. A.; Maggiora,
G. M.; Schowen, R. L. J Am Chem Soc 1983, 105, 31.
22. Wolfe, S.; Kim, C. K.; Yang, K.; Weinberg, N.; Shi, Z.
J Am Chem Soc 1995, 117, 4240.
23. Jonnalagadda, S. B.; Nattar, K. Int J Chem Kinet 1999,
31, 83.
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