Phenylureas. Part 2. Mechanism of the acid hydrolysis of phenylureas

Article abstract of DOI:10.1039/b008535i

The mechanism of the hydrolytic decomposition of phenylureas in acid media is investigated. It includes, in part, knowledge already present in the literature. Over the investigated pH range the occurrence of a rate maximum in the pH curves due to the strongly reduced water activity at higher acid strengths is observed. An addition-elimination mechanism with rate-determining attack of water at the N-protonated substrate is proposed. The reversion of the substituent influence on the reaction rate with increasing acidity of the reaction medium points to a change of the hydrolytic decomposition mechanism in strongly acidic media.

Full text of DOI:10.1039/b008535i

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Phenylureas. Part 2.1 Mechanism of the acid hydrolysis of
phenylureas
Robert Laudien*† and Rolf Mitzner
2
Institute of Physical and Theoretical Chemistry, University of Potsdam, Am Neuen Palais 10,
14469 Potsdam, Germany
Received (in Cambridge, UK) 23rd October 2000, Accepted 27th July 2001
First published as an Advance Article on the web 5th October 2001
The mechanism of the hydrolytic decomposition of phenylureas in acid media is investigated. It includes, in part,
knowledge already present in the literature. Over the investigated pH range the occurrence of a rate maximum in the
pH curves due to the strongly reduced water activity at higher acid strengths is observed. An addition–elimination
mechanism with rate-determining attack of water at the N-protonated substrate is proposed. The reversion of the
substituent influence on the reaction rate with increasing acidity of the reaction medium points to a change of the
hydrolytic decomposition mechanism in strongly acidic media.
Kallies and Mitzner18 deduce from quantum-mechanical
Introduction
calculations that the protonation of the carbonyl oxygen of acyl
The hydrolysis of carboxylic acid esters and amides in acid
compounds leads to an increase of resonance at the reaction
media mainly proceeds according to an addition–elimination
centre in addition to the catalytically acting rise of electro-
mechanism.2–8 The carbonyl oxygen of the substrate is proton-
philicity of the carbonyl carbon. The increase of the resonance
ated in a preequilibrium, by which the electrophilicity of the
is due to the stronger involvement of the non-bonding electrons
carbonyl carbon is increased and, consequently, the nucleo-
of the alcoholic oxygen with the aminic nitrogen, which has
philic attack of water is facilitated thereby forming a short-lived
to be cancelled in the case of addition of a nucleophile. This
tetrahedral intermediate. In the case of carboxylic acid amides,
negative catalytic effect towards a nucleophilic attack at the
protonation of the nitrogen is considered already to proceed
carbonyl group increases in the order ester, amide, urea because
in moderately acid solutions along with the protonation of the
of the increasing resonance stabilization of the O-protonated
carbonyl oxygen. This can be deduced from theoretical con-
substrate and can account for the change to an elimination–
siderations of the difference in the pKa values of O- and N-
addition mechanism for the acid hydrolytic decomposition of
protonated amides.9,10
ureas in contrast to esters and amides.
For the acid hydrolysis of alkylureas in solutions of low
and medium acidity an elimination–addition mechanism
via isocyanate is postulated. The decomposition of urea in
weakly acid media proceeds without acid catalysis via a rate-
determining intramolecular proton transfer and subsequent
dissociation of the transition state11,12 (Scheme 1). The reaction
According to Giffney and O’Connor,19,20 for phenylureas
the elimination of the amine from the unprotonated or N-
protonated urea is promoted by an intermolecular proton
transfer via water (Scheme 2). An N-protonation besides the
Scheme 2
O-protonation is already considered to be possible in relatively
weakly acid solutions. The existence of a six-membered cyclic
complex between the N-protonated urea and water, however,
Scheme 1
could not be confirmed through quantum-mechanical calcu-
is decelerated with increasing acid content in media of low
lations by Lee and co-workers.21
acidity due to a reduction in the concentration of the more
The aim of the present investigations was to reexamine
the mechanism of the acid hydrolysis of phenylureas. Mainly
kinetic investigations were carried out with respect to structure–
reactivity relations, specific acid catalysis and deuterium
reactive unprotonated urea species.13 In high acidity ranges of
the acid (ca. 70% w/w H2SO4), however, a maximum in the rate
constant occurs, which is attributed to a changed reaction
course—bimolecular attack of water at the diprotonated urea
solvent isotope effects, the experimental results of Giffney and
species.14 The existence of such a species could be established
O’Connor19,20 being in part included in the discussion.
for urea15,16 and several substituted methylureas17 in strongly
acid solutions by 1H- and 15N-NMR spectroscopic investi-
gations.
Results and discussion
The hydrolysis of phenylureas, variably substituted both at
the phenyl ring and at the nitrogen of the leaving group, was
studied kinetically in acidic water–methanol solutions (9 : 1) at
† Current address: Research Institute of Electronics, Shizuoka
University, 3-5-1 Johoku, Hamamatsu 432-8011, Japan. E-mail:
rlrober@ipc.shizuoka.ac.jp
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J. Chem. Soc., Perkin Trans. 2, 2001, 2230–2232
This journal is © The Royal Society of Chemistry 2001
DOI: 10.1039/b008535i

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Table 1 Rate constants (in sϪ1) for the hydrolysis of phenylureas in 0.1 M H2SO4 at 80 and 90 ЊC
Structure 1 Structure 2
Structure 3
No.
X
105 k1
105 k2
No.
X
105 k1
105 k2
No.
R, RЈ
105 k1
105 k2
1a
1b
1c
1d
1e
1f
1g
1h
1i
m-NO2
m-CF3
m-Cl
p-Cl
m-OH
0.490
0.659
0.653
0.689
0.821
0.881
0.988
0.960
0.970
1.009
1.192
1.381
1.745
2.025
2.178
2.358
2.524
2.854
3.256
3.136
3.191
3.051
3.728
4.341
2a
2b
2c
2d
2e
2f
3,4-Cl
p-Cl
3-Cl, 4-OCH3
3-Cl, 4-CH3
p-OC6H4Cl
H
2.732
3.362
3.556
3.615
3.391
4.149
4.758
9.555
12.27
11.91
13.20
13.47
15.66
17.04
3a
3b
3c
3d
2f
3e
3f
C4H9, H
C2H5, H
CH(CH3)2, H
CH(CH3)2
H
0.272
0.296
0.478
7.212
4.149
6.397
11.30
15.30
15.76
52.64
207.9
1.077
1.191
1.862
24.84
15.66
22.92
38.24
56.08
58.00
173.6
63.09c
H
–C5H10–b
C3H7
3,4-OCH3
–OCH2O–a
p-CH3
p-OCH3
2,5-OCH3
o-OCH3
2g
p-CH(CH3)2
3g
3h
3i
C4H9
C2H5
CH(CH3)C2H5
CH(CH3)2
1k
1l
1m
3k
a 3,4-Methylenedioxy. b Piperidino. c Rate constant at 70 ЊC.
Table 2 Rate constants (in sϪ1) for the hydrolysis of phenylureas at
various H2SO4 concentrations
[H2SO4]/mol lϪ1 1f, 90 ЊC, 105 k
2f, 90 ЊC, 104 k
3k, 80 ЊC, 103 k
0.001
0.002
0.005
0.01
0.02
0.05
0.1
0.2
0.3
0.5
0.9
1.716
0.746
0.678
0.912
1.356
1.650
1.869
2.023
2.053
1.996
1.998
2.211
2.431
2.772
3.234
3.934
4.405
5.028
5.651
1.292
1.519
1.627
1.720
1.747
1.730
1.690
1.632
1.440
1.730
Scheme 3
80 and 90 ЊC. The investigated phenylureas have the following
basic structures.
Table 1 contains pseudo first-order rate constants for the
hydrolysis of the phenylureas in 0.1 M H2SO4 at 80 ЊC (k1) and
90 ЊC (k2). In Table 2 rate constants for the hydrolysis of three
phenylureas at various hydronium ion concentrations are listed.
All pseudo first-order rate constants are given in the unit sϪ1
.
The rate of hydrolysis of phenylureas is not directly pro-
portional to the hydronium ion concentration, but rises with
the acid strength only in solutions of low acidity, whereas
it passes through a maximum and decreases again in stronger
acid solutions in accordance with the results of Giffney and
O’Connor.19,20 This result clearly indicates the importance for
the reaction course of water activity, which is appreciably
reduced at relatively high concentrations of sulfuric acid.
A possible reaction mechanism is shown in Scheme 3. In the
course of the hydrolytic decomposition the substrate is first
protonated followed by a rate-determining attack by water. The
tetrahedral intermediate then decomposes to an amine and
a phenylcarbamic acid which again is decarboxylated very
quickly under acidic conditions to form the corresponding
aniline.
A Hammett plot for the hydrolysis of substituted N-phenyl-
ureas in 0.1 M H2SO4 at 90 ЊC (Fig. 1) results in a small negative
reaction constant ρ of Ϫ0.33 (R2 = 0.952). A similar reaction
constant of Ϫ0.38 (R2 = 0.972) is obtained for the hydrolysis of
substituted N-phenyl-NЈ,NЈ-dimethylureas at 90 ЊC. The values
for the rate data at 80 ЊC are Ϫ0.36 (R2 = 0.962) and Ϫ0.34
(R2 = 0.961), respectively. In the Hammett plots the Hammett
Fig. 1 Hammett plot for the hydrolysis of substituted N-phenylureas
in 0.1 M H2SO4 at 90 ЊC.
constant σϪ was applied because of the direct resonance inter-
action between the ϪM substituents and the ϩM reaction
center leading to an increased stabilisation of the negative
charge at the reaction centre. That is, the rate of the hydrolysis
of phenylureas in 0.1 M H2SO4 is increased by electron-
donating substituents and decreased by electron-attracting
substituents on the aryl group. Electron-donating substituents
increase the electron density in the molecule and thus facilitate
the protonation of the nitrogen of the alkylamine group which
greatly supports the nucleophilic attack of water at the carbonyl
carbon.
In a similar manner the dependence of the hydrolysis rate
on the basicity of the aminic leaving group can be interpreted.
The hydrolysis rate basically increases with rising basicity of
the amine leaving group. Stronger basic alkylamine groups
facilitate the protonation of the alkyl–NH group.
According to the data of Giffney and O’Connor,19,20 the
influence of the substituents at the phenyl ring on the hydrolysis
rate changes with increasing acidity of the reaction medium.
J. Chem. Soc., Perkin Trans. 2, 2001, 2230–2232
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Because of these low isotope effects, a proton transfer from or
to water in a rate-determining step in the reaction course can be
excluded.
Experimental
For experimental details see Part 1 of this series.1
Acknowledgements
This work was supported by the Innovationskolleg INK 16/A1-1
of the Deutsche Forschungsgemeinschaft.
Fig. 2 Plot of reaction constants ρ versus H2SO4 concentration for
the hydrolysis of substituted N-phenylureas at 101 ЊC (data from ref. 19
and 20).
References
1 R. Laudien and R. Mitzner, J. Chem. Soc., Perkin Trans. 2, 2001,
preceding paper (DOI: 10.1039/b008532o).
2 D. P. N. Satchell and R. S. Satchell, in The Chemistry of Acid
Derivatives, ed. S. Patai, Wiley, Chichester–New York–Brisbane–
Toronto–Singapore, 1992, vol. 2, p. 747.
3 H. R. Christen and F. Vögtle, Organische Chemie – Von den
Grundlagen zur Forschung, 2nd edn., Salle & Sauerländer, Frankfurt
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4 P. Sykes, Reaktionsmechanismen der Organischen Chemie, 9th edn.,
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6 W. P. Jencks, Catalysis in Chemistry and Enzymology, McGraw-Hill,
New York, 1987, p. 463.
Giffney and O’Connor measured the hydrolysis rates of N-
phenylureas, variably substituted at the phenyl ring, at various
sulfuric acid concentrations in the range 0.058 to 9.15 mol lϪ1 at
101 ЊC. Hammett plots of these data show that the substituent
influence, as above, is found only up to a concentration of
around 1 mol lϪ1, while it is reversed in increasingly acidic solu-
tions. From a concentration of around 4 mol lϪ1 the reaction
constant remains positive over the whole substituent range.
The changing dependence of the reaction constant ρ on the
sulfuric acid concentration for the hydrolysis of substituted
N-phenylureas is shown in Fig. 2.
This change of the substituent influence on the hydrolysis
rate with increasing acid strength points to a changed reaction
mechanism which is clearly due to the reduced water activity in
high acidity ranges. This new mechanism is very likely to be
an A1 mechanism as shown in Scheme 4 in which the aminic
7 S. L. Johnson, Adv. Phys. Org. Chem., 1967, 5, 237.
8 M. L. Bender, Chem. Rev., 1960, 60, 53.
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10 R. B. Homer and C. D. Johnson, in The Chemistry of Amides,
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Scheme 4
leaving group is eliminated without prior formation of a tetra-
hedral addition complex with water.
For the hydrolysis of 1f and 2f, deuterium solvent isotope
effects have been determined. For 1f kH/kD is 1.15, for 2f 0.98.
As the rate constants are affected by an error rate of ca. 3%, an
uncertainty of 6% maximum results for the isotope effects.
2232
J. Chem. Soc., Perkin Trans. 2, 2001, 2230–2232