Kinetics and mechanism of the mineral acid catalysed reactions of hydroxamic acids

Article abstract of DOI:10.1524/zpch.218.5.563.30505

There is currently a great deal of interest in the chemistry of hydroxamic acids. In recent years we have been studying the synthesis, structure and nucleophilicities of hydroxamic acids. This paper reports a kinetic study of reactivity of some hydroxamic acids RC(O)·N(OH)R′; R = C ₆H₅·CH=CH, R′ = 4-CH₃· C₆H₄ [p-tolylcinnamo hydroxamic acid]; R = C ₆H₅, R′ 4-CH₃·C₆H ₄ [p-tolylbenzo hydroxamic acid]; R = C₆H₅, R′ = 2-CH₃·C₆H₄ [o-tolylbenzo hydroxamic acid] in aqueous mineral acids (HCl, HClO₄ and H ₂SO₄). The rate data of hydrolysis reaction revealed that the reactivity/stability sequence of the compounds is generally p-TBHA > o-TBHA > p-TCHA. An excess acidity analysis reveals that the reaction proceeds by nucleophilic attack of water molecule on the protonated substrate.

Full text of DOI:10.1524/zpch.218.5.563.30505

Z. Phys. Chem. 218 (2004) 563–573
by Oldenbourg Wissenschaftsverlag, München
Kinetics and Mechanism of the Mineral Acid
Catalysed Reactions of Hydroxamic Acids
∗
By Kallol K. Ghosh , Jyoti Vaidya, and Daliya Sinha
School of Studies in Chemistry, Pt. Ravishankar Shukla University, Raipur (C.G.), India
(Received December 15, 2003; accepted in revised form February 25, 2004)
Hydroxamic Acid / Acid Catalysis / Excess Acidity / Hydrolysis
There is currently a great deal of interest in the chemistry of hydroxamic acids. In
recent years we have been studying the synthesis, structure and nucleophilicities of
hydroxamic acids. This paper reports a kinetic study of reactivity of some hydroxamic
acids RC(O)·N(OH)Rꢀ; R = C6H5 ·CH=CH, Rꢀ = 4-CH3 ·C6H4 [p-tolylcinnamo hydrox-
amic acid]; R = C6H5, Rꢀ = 4-CH3 ·C6H4 [p-tolylbenzo hydroxamic acid]; R = C6H5,
Rꢀ = 2-CH3 ·C6H4 [o-tolylbenzo hydroxamic acid] in aqueous mineral acids (HCl, HClO4
and H2SO4). The rate data of hydrolysis reaction revealed that the reactivity/stability
sequence of the compounds is generally p-TBHA > o-TBHA > p-TCHA. An excess
acidity analysis reveals that the reaction proceeds by nucleophilic attack of water
molecule on the protonated substrate.
1. Introduction
Recently, hydroxamic acids have received much attention due to their nu-
merous biological, medicinal and analytical applications [1–8]. These acids
contain the simplest and smallest structure (NH··O=C) that winds the DNA
and RNA protein (secondary structure) into an α-helix skeleton. Many bio-
chemical and chemical reactions involve at least one proton transfer, and the
most of the common mechanism of catalysis of these reactions, by enzymes
or in the laboratory, involves some sort of facilitation of this proton transfer.
Classical physical organic chemists studies over the last 20 years yielded some
important clues about the reactivity of hydroxamic acids in acid and basic solu-
tions [9–13]. Structural modifications of hydroxamic acids are often performed
in order to improve the properties of these compounds for research in medicine,
biochemistry and analytical chemistry.
*
Corresponding author. E-mail: kallolkghosh@yahoo.com

564
K. K. Ghosh et al.
Scheme 1.
The present paper reports a kinetic study of the some hydroxamic acids viz.
(Scheme 1) p-tolyl cinnamo hydroxamic acid (p-TCHA), p-tolyl benzohy-
droxamic acid (p-TBHA) and o-tolylbenzohydroxamic acid (o-TBHA) in min-
eral acids.
The Cox–Yates excess acidity method [14] was used to analyse the rate
data obtained. This method has been found to be very useful in obtaining
mechanistic information in strong acid media concerning a wide variety of hy-
+
+
represents the ‘extra’
drolysis reactions. The excess acidity X = fB fH / fB
*
*
H
acidity of the medium due to its non-ideal nature (14c). It manifests the use-
ful property of being zero in the standard state of unit activity coefficient, the
same hypothetical ideal 1 M acid solution that is the standard state for pH
measurements. The novel feature of this investigation is that it provides a judi-
cious and rational kinetic basis for the choice of an organic analytical reagent
from the family of hydroxamic acids. From analytical considerations it will
be very rewarding to kinetically investigate the hydrolytic stabilities of metal-
hydroxamic acid complexes particularly in concentrated acid solutions. The
effect of mineral acid concentration on the colour intensity of a system is like
wise examined empirically. Kinetic data for such systems will provide a ratio-
nal basis for the choice of optimum range of acid concentration for maximum
colour development.
2. Results and discussion
All the reactions followed pseudo-first order kinetics:
d[HA]
−
= K[HA][H+] = kψ[HA]
dt

Kinetics and Mechanism of the Mineral Acid Catalysed Reactions . . .
565
Table 1. Pseudo first-order rate constants for the hydrolysis of hydroxamic acids in HCl,
Temp. = 55 ◦C, 20% (v/v) 1-4 dioxane.
HCl
X
kψ105/s−1
[M]
p-TBHA
o-TBHA
p-TCHA
1.00
1.45
2.00
2.90
4.00
5.00
5.80
7.00
0.24
0.33
0.46
0.74
1.03
1.34
1.65
2.09
3.00
4.50
5.50
7.00
11.0
15.1
18.7
23.0
1.70
2.10
2.70
3.58
5.00
6.70
7.90
9.60
0.60
0.90
1.45
2.20
3.50
4.50
5.79
7.45
Table 2. Pseudo first-order rate constants for the hydrolysis of hydroxamic acids in
HClO4, Temp. = 55 ◦C, 20% (v/v) 1-4 dioxane.
HClO4
X
kψ105/s−1
[M]
p-TBHA
o-TBHA
p-TCHA
1.00
1.45
2.00
2.90
4.00
5.00
5.80
7.00
0.24
0.33
0.49
0.86
1.31
1.80
2.34
2.84
5.40
6.50
7.20
9.20
11.0
13.0
14.6
16.1
0.90
1.80
2.30
3.10
3.80
4.50
5.10
6.00
0.50
1.29
2.00
3.40
4.90
5.20
5.10
7.10
The observed pseudo-first order rate constants of p-TCHA, p-TBHA and
o-TBHA for the catalytic effect of hydrochloric, sulfuric and perchloric acids
are given in Tables 1–3. In order to explain the differences in the catalytic
effeciencies of the different acids and to illustrate the rate–acidity profiles,
these data have been plotted in Figs. 1–3 against X (excess acidity) of H2SO4,
HCl and HClO4. In the weakly acidic solutions, linear plots were obtained.
From the intercept it could be inferred that hydroxamic acid did not react
with water in the absence of catalyst to a measurable extent at 55 ◦C. Gen-
erally hydroxamic acids are susceptible to acid and to base catalysis, but
their spontaneous hydrolysis contribute relatively little to the overall reac-
tion rate. At higher acid regions (H2SO4) the dependence of p-TCHA is
characterised by an initial rate increase with the acid concentration, pass-
ing through a maxima, followed by a rate decrease with further increasing
acid concentration. This non-linearity in parts because the equilibrium be-
tween the reactants and the protonated species of the rate determining step

566
K. K. Ghosh et al.
Table 3. Pseudo first-order rate constants for the hydrolysis of hydroxamic acids in
H2SO4, Temp. = 55 ◦C, 20% (v/v) 1-4 dioxane.
H2SO4
X
kψ 105/s−1
[M]
p-TBHA
o-TBHA
p-TCHA
1.00
1.45
2.00
2.90
4.00
5.00
5.80
7.00
0.24
0.34
0.48
0.79
1.14
1.51
1.90
2.35
3.00
4.10
6.10
7.90
15.9
24.2
31.4
40.4
1.40
1.93
2.50
3.80
5.90
7.80
9.10
27.00
3.20
3.80
5.70
5.90
4.40
3.60
3.50
6.00
Fig. 1. Rate acidity profiles for the hydrolysis of hydroxamic acids in HCl using 20% (v/v)
1-4 dioxane.
does not correspond to a simple protonation, but also to addition of water
molecule.
Several interesting feature has been observed. Firstly, the rate of hy-
drolysis of hydroxamic acids i.e., p-TCHA is very slow in comparison
with other hydroxamic acids i.e. p-TBHA and o-TBHA. The rate constant
data indicate that the relative reactivities of the hydroxamic acids toward
the hydrolysis are p-TBHA > o-TBHA > p-TCHA. There is no absolute
measure of comparison of relative hydrolytic stability. Different hydroxamic
acids have different nature of rate–acid concentration profiles. This order
of reactivities is rationalized in terms of resonance, steric hindrance and

Kinetics and Mechanism of the Mineral Acid Catalysed Reactions . . .
567
Fig. 2. Rate acidity profiles for the hydrolysis of hydroxamic acids in HClO4 using 20%
(v/v) 1-4 dioxane.
Fig. 3. Rate acidity profiles for the hydrolysis of hydroxamic acids in H2SO4 using 20%
(v/v) 1-4 dioxane.
inductive effects. Thus from kinetic studies it is concluded that p-TCHA
has the highest hydrolytic stabilities and is therefore, best suited for sol-
vent extraction and spectrophotometric determinations in concentrated acidic
solutions.
2.1 Excess acidity analysis
The Cox–Yates excess acidity method [14] is the most widely used method
for analysing reaction kinetics and determining mechanisms in aqueous min-
eral acids media. Most studies of carboxylic acid derivatives i.e. amides, esters

568
K. K. Ghosh et al.
and anilides involve the following three types of acid catalysed hydrolysis
processes:
A-1 mechanism:
K
+
SH
S+H+ ꢀ SH+
k
0
SH+ −→ A+ −→ products .
slow
fast
This mechanism represents a rate determining unimolecular reaction, with
a medium independent rate constant k0, of a protonated substrate SH+ formed
from the substrate S in a fast per-equilibrium proton transfer.
A-SE-2 mechanism:
k
1
S+H+ −→ SH+ −→ products .
slow
fast
This represents rate-determining proton transfer to the substrate, with subse-
quent reaction of SH+ being fast.
A-2 mechanism:
K
+
SH
S+H+ ꢀ SH+
fast
k
2
SH+ +n Nu (H2O) −→ products .
This is like the A-1 mechanism except that the protonated substrate, instead of
reacting alone, reacts with n number of water molecules (nucleophile) in the
rate determining step.
These three possible hydrolysis mechanisms (A-1, A-2 and A-SE-2) can
be easily distinguished and analysed by this method. The excess acidity repre-
sents the difference between the actual solution acidity and the stoichiometric
acid concentration [15]. The following three equations for A-1, A-SE-2 and A-2
mechanism can be obtained.
k0
#
*
+
log kψ −log CH = log
+m m X
(1)
(2)
(3)
+
KSH
#
*
+
log kψ −log CH = log k1 +m m X
log kψ −log CH −n log aNu = log
k2
#
*
+
+m m X .
+
KSH
+
In these equations kψ is observed rate constant as a function of acidity, CH
is the molar proton concentration, X is the excess acidity of the medium,
Nu stands for the nucleophile (normally water) and n is the number of water
molecules.

Kinetics and Mechanism of the Mineral Acid Catalysed Reactions . . .
569
Fig. 4. Cox–Yates plots for the A-2 mechanism of hydrolysis of hydroxamic acids at 55 ◦C
in HClO4.
These equations have been extensively tested for the determination of re-
action mechanism and work well. The excess acidity methods works by exam-
+
ining plots of log kψ −log CH against X (figures not shown). These plots are
curved, indicating A-2 mechanism. If it were A-1 or A-SE-2, the plots would be
straight line. Once a mechanism is identified as being A-2, the species reacting
with the protonated substrate at the transition state (Nu) can be positively iden-
+
tified by plotting log kψ −log CH −n log aNu against X, trying different Nu,
until linearity is achieved. Linear plots were obtained with n = 2 (2 log aH O).
2
Some representative excess acidity plots for p-TCHA, p-TBHA and o-TBHA
in mineral acids are shown in Figs. 4–6.
The two water molecules must be involved in bond breaking/making pro-
cesses at the transition state, since the excess acidity method separates, kinet-
ically involved water from solvation water (which shows up in m or m#).
*
One water molecule is a nucleophile, attacking the carbonyl oxygen, the sec-
ond is an acid/base transferring protons between others. Table 4 gives the slope
#
*
(m m ) values for all the hydroxamic acids. These values are consistent with
*
A-2 mechanism [16]. The m for carbonyl oxygen protonation [14] is 0.6 or
less & for A-2 reactions m# = 1, thus an overall slope against X of 0.6 or less
should result for Eq. (3) [Table 4].
The X and aH values employed in these correlations were taken from
O
2
literature [14a].
2.2 Product identification
Hydroxamic acids hydrolyse to give respective hydroxylamine (N-o-tolyl, and
N-p-tolyl) and carboxylic acids (benzoic and cinnamic acids). The carboxylic

570
K. K. Ghosh et al.
Fig. 5. Cox–Yates plots for the A-2 mechanism of hydrolysis of hydroxamic acids at 55 ◦C
in H2SO4.
Fig. 6. Cox–Yates plots for the A-2 mechanism of hydrolysis of hydroxamic acids at 55 ◦C
in HCl.
acids were separated and analysed by m.p., UV and IR spectra. The substi-
tuted hydroxylamines were unstable and decomposed into a variety of complex
products.
2.3 Mechanism
On the basis of above discussion and results obtained from the excess acidity
analysis the most plausible mechanism (Scheme 2) appears to involve a rapid

Kinetics and Mechanism of the Mineral Acid Catalysed Reactions . . .
571
Table 4. Excess acidity correlation for the hydrolysis of hydroxamic acids, Temp. = 55 ◦C,
20% (v/v) 1-4 dioxane.
p-TBHA
o-TBHA
p-TCHA
No. of Slope Correlation No. of Slope Correlation No. of Slope Correlation
points
coefficient points
coefficient points
coefficient
HCl
H2SO4
HClO4
8
8
7
0.426
0.896
0.396
0.985
0.995
0.985
8
8
8
0.368
0.893
0.544
0.985
0.987
0.980
8
8
8
0.527
0.389
0.633
0.995
0.905
0.970
Scheme 2.
protonation equilibrium of hydroxamic acid followed by a slow reaction of
water with protonated form, (II), leading to the transition complex, III, and
subsequent hydrolysis resulting in the fission of N-acyl bond giving corres-
ponding carboxylic acid and protonated hydroxylamine as the products. It is
now a well recognized fact that the carbonyl oxygen is the predominant basic
center [17, 18] in hydroxamic acids.
3. Experimental
All the hydroxamic acids were prepared by benzoylation freshly prepared and
purified by literature procedure [19]. All other chemicals were of analytical
reagent grade.

572
K. K. Ghosh et al.
3.1 Rate measurements
For each kinetic run two reaction vessels were used. One of this contained
appropriate volume of acid (catalyst) and water, the other one contained the
hydroxamic acid. After thermostating for about 30 min the acid solution was
transferred to the reaction vessel containing hydroxamic acid. After shaking
the content of the reaction vessel an aliquot (2 ml) of the reaction mixture
was withdrawn into 10 ml volumetric flask containing 2 ml of ferric chlo-
ride. Double purpose, quenching of the reaction and colour development, was
thus served. The volume of the coloured solution (4 ml) was made to 10 ml
with water and its absorbance was measured at 500–520 nm using a ref-
erence solution containing 2 ml of same ferric chloride in 10 ml of water.
The decrease of hydroxamic acid concentration was followed spectrophoto-
metrically using its reaction with ferric ion. The kinetic runs were studied
generally upto one to two half-lives. For measuring absorbance Systronics
UV-VIS spectrophotometer type 108 and Unicam UV2 300 spectrophotometer
were used.
For a pseudo first-order reaction a plot of log absorbance vs. t (where ab-
sorbance is the concentration of reacting species at time t) will give a straight
line of slope, kψ/2.303 (where kψ is the pseudo first-order rate constant
for the reaction). As Beer’s law is applicable to all the ferric hydroxamic
acids complexes, the concentration of reacting species is proportional to the
absorbance A. [log A = (log a− x)]. To obtain the rate constant kψ, log(a− x)
was plotted against time t, and from the slope of the plot, kψ was determined
the fact that a straight line was obtained for all the plots of log(a − x) vs. t
measured in this investigation is, in itself, an indication that the reactions are all
first-order in substrate. The initial concentration of the hydroxamic acid in re-
action mixture is about 7.0×10−3 M. The experimental errors in the respective
runs were generally less than 1.0% and the reproducibility of the rate constants
was within 1.5%.
Acknowledgement
Financial support of this work from Pt. Ravishankar Shukla University, Raipur
is greatly acknowledged. The authors are grateful to Professor G. L. Mundhara,
Head for providing facilities.
References
1. C. lndiani, A. Feis, B. D. Howes, M. P. Marzocchi, and G. J. Smulevich, J. Am.
Chem. Soc. 122 (2000) 7368.
2. D. E. Robinson and M. W. Holladay, Org. Lett. 2 (2000) 2777.
3. M. Takayanagi, T. Flessner, and C. H. Wong, J. Org. Chem. 65 (2000) 3811.
4. L. Wang and O. Phanstiel IV, J. Org. Chem. 65 (2000) 1442.
5. S. Hashimoto, T. Inui, and Y. Nakamura, Chem. Pharm. Bull. 48 (2000) 603.

Kinetics and Mechanism of the Mineral Acid Catalysed Reactions . . .
573
6. E. Farkas, H. Csoka, and I. Toth, Dalton Trans. (2003) 1645.
7. El yazal and Y. P. Pang, J. Phys. Chem. B 104 (2000) 6499.
8. J. Schraml, Appl. Organometal. Chem. 14 (2000) 604.
9. D. C. Berndt and I. E. Ward, J. Org. Chem. 43 (1978) 13.
10. K. K. Ghosh, S. K. Patle, P. Sharma, and S. K. Rajput, Bull. Chem. Soc. Jpn. 76
(2003) 283.
11. K. K. Ghosh, P. Tamrakar, and S. K. Rajput, J. Org. Chem. 64 (1999) 3053.
12. K. K. Ghosh, Indian J. Chem. B 368 (1997) 1089.
13. K. K. Ghosh and S. Ghosh, J. Org. Chem. 59 (1994) 1369; K. K. Ghosh and
P. Tamrakar, Z. Phys. Chem. 217 (2003) 1153.
14. (a) R. A. Cox, Adv. Phys. Org. Chem. 1 (2000) 1. (b) R. A. Cox, Acc. Chem. Res.
20 (1987) 27. (c) R. A. Cox and K. Yates, J. Am. Chem. Soc. 100 (1978) 3861.
15. R. A. Cox, Can. J. Chem. 74 (1996) 1774.
16. R. A. Cox and K. Yates, Can. J. Chem. 62 (1984) 1613.
17. M. Saldyka and Z. Mielke, J. Phys. Chem. A 107 (2003) 2448.
18. A. Bagno, G. Scorrano, and R. A. More O’Ferrall, Rev. Chem. Int. 7 (1987) 713.
19. Y. K. Agrawal and S. G. Tandon, J. Chem. Eng. Data 16 (1971) 371.