Hydrolysis of N-phenylalkanesulfinamides in aqueous mineral acids

Article abstract of DOI:10.1080/10426507.2010.508691

The acid-catalyzed hydrolysis of N-phenylalkanesulfinamides (RSONHPh; 1, R = ⁱPr; 2, R = ^tBu; 3, R = 1-adamantyl) has been studied in aqueous mineral acids. Hydrolysis was found to proceed via a slow spontaneous (uncatalyzed) pathway, an A-2 (bimolecular) acid-catalysis pathway, and an acid-dependent nucleophilic catalysis pathway, the last of which predominates in hydrobromic and hydrochloric acid solutions. A mechanistic switch over from A-2 to A-1 was detected for compounds 2 and 3 in concentrated sulfuric acid. Order of catalytic activity, effect of added salts, Arrhenius parameters, kinetic solvent isotope, and solvent effects are all consistent with the proposed mechanisms. Copyright Taylor & Francis Group, LLC.

Full text of DOI:10.1080/10426507.2010.508691

This article was downloaded by: [Moskow State Univ Bibliote]
On: 20 December 2013, At: 06:13
Publisher: Taylor & Francis
Informa Ltd Registered in England and Wales Registered Number: 1072954 Registered
office: Mortimer House, 37-41 Mortimer Street, London W1T 3JH, UK
Phosphorus, Sulfur, and Silicon and the
Related Elements
Publication details, including instructions for authors and
subscription information:
http://www.tandfonline.com/loi/gpss20
Hydrolysis of N-
Phenylalkanesulfinamides in Aqueous
Mineral Acids
Mrityunjoy Datta a , Alan J. Buglass a & John G. Tillett b
a Department of Chemistry , KAIST , Daejeon, Republic of Korea
b Centre for Continuing Education , University of Essex , Colchester,
United Kingdom
Published online: 07 Mar 2011.
To cite this article: Mrityunjoy Datta , Alan J. Buglass & John G. Tillett (2011) Hydrolysis of N-
Phenylalkanesulfinamides in Aqueous Mineral Acids, Phosphorus, Sulfur, and Silicon and the Related
Elements, 186:3, 565-573, DOI: 10.1080/10426507.2010.508691
To link to this article: http://dx.doi.org/10.1080/10426507.2010.508691
PLEASE SCROLL DOWN FOR ARTICLE
Taylor & Francis makes every effort to ensure the accuracy of all the information (the
“Content”) contained in the publications on our platform. However, Taylor & Francis,
our agents, and our licensors make no representations or warranties whatsoever as to
the accuracy, completeness, or suitability for any purpose of the Content. Any opinions
and views expressed in this publication are the opinions and views of the authors,
and are not the views of or endorsed by Taylor & Francis. The accuracy of the Content
should not be relied upon and should be independently verified with primary sources
of information. Taylor and Francis shall not be liable for any losses, actions, claims,
proceedings, demands, costs, expenses, damages, and other liabilities whatsoever or
howsoever caused arising directly or indirectly in connection with, in relation to or arising
out of the use of the Content.
This article may be used for research, teaching, and private study purposes. Any
substantial or systematic reproduction, redistribution, reselling, loan, sub-licensing,
systematic supply, or distribution in any form to anyone is expressly forbidden. Terms &
Conditions of access and use can be found at http://www.tandfonline.com/page/terms-
and-conditions

Phosphorus, Sulfur, and Silicon, 186:565–573, 2011
Copyright ꢀC Taylor & Francis Group, LLC
ISSN: 1042-6507 print / 1563-5325 online
DOI: 10.1080/10426507.2010.508691
HYDROLYSIS OF N-PHENYLALKANESULFINAMIDES
IN AQUEOUS MINERAL ACIDS
Mrityunjoy Datta,1 Alan J. Buglass,1 and John G. Tillett2
1Department of Chemistry, KAIST, Daejeon, Republic of Korea
2Centre for Continuing Education, University of Essex, Colchester,
United Kingdom
GRAPHICAL ABSTRACT
Abstract The acid-catalyzed hydrolysis of N-phenylalkanesulfinamides (RSONHPh; 1, R =
t
iPr; 2, R = Bu; 3, R = 1-adamantyl) has been studied in aqueous mineral acids. Hydrolysis
was found to proceed via a slow spontaneous (uncatalyzed) pathway, an A-2 (bimolecular)
acid-catalysis pathway, and an acid-dependent nucleophilic catalysis pathway, the last of which
predominates in hydrobromic and hydrochloric acid solutions. A mechanistic switch over from
A-2 to A-1 was detected for compounds 2 and 3 in concentrated sulfuric acid. Order of catalytic
activity, effect of added salts, Arrhenius parameters, kinetic solvent isotope, and solvent effects
are all consistent with the proposed mechanisms.
Keywords Acid-catalyzed hydrolysis; acid-dependent nucleophilic catalysis; alkanesulfi-
namides; reaction mechanism
INTRODUCTION
Nucleophilic substitution at sulfur has been observed for a number of sulfinyl com-
pounds, hydrolysis being typical of these reactions.1 Both acid-catalyzed and hydrogen
ion–dependent, nucleophile-catalyzed hydrolysis and related reactions have been reported
for sulfoxides,2,3 sulfite esters,4,5 sulfinyl sulfones,6 sulfinates,7 and arenesulfinamides.8–10
Wagner et al. compared the rates of neutral (spontaneous) hydrolysis of a series of
cyclic sulfinamides with that of N,N-dimethylmethanesulfinamide at pD 7.7 in D2O.11 The
Received 7 June 2010; accepted 6 July 2010.
Financial support from KAIST is gratefully acknowledged.
Address correspondence to Alan J. Buglass, Department of Chemistry, KAIST, Daejeon 305-701, Republic
of Korea. E-mail: alan buglass@yahoo.com
565

566
M. DATTA ET AL.
latter was unexpectedly more reactive than its cyclic counterparts, and its rate of hydrolysis
was too fast to measure in all but mildly acid conditions.
For the hydrolysis of N-arylarenesulfinamides in aqueous solutions of hydrochloric
and hydrobromic acids, parallel acid–catalyzed and acid-dependent nucleophilic mech-
anisms were proposed as the principal pathways.8 In perchloric and sulfuric acids, the
mechanism was considered to be of the A-2 type, while in halogen acid solutions, the
rate-limiting step involves formation of a reactive sulfinyl intermediate, which rapidly de-
composes to the products. While no spontaneous reaction could be detected at 25◦C, at
higher temperatures a minor nucleophilic-catalyzed spontaneous reaction was detected.
To date, however, apart from the brief study of N,N-dialkylalkanesufinamides,11 there
is no detailed account of the acid-catalyzed hydrolysis of alkanesulfinamides in the litera-
ture. We now report on the hydrolysis of N-phenylalkanesulfinamides (1–3) (RSONHPh;
i
t
1, R = Pr; 2, R = Bu; 3, R = 1-adamantyl) in aqueous solutions of mineral acids.
RESULTS AND DISCUSSION
Figure 1 shows the kinetic profile of alkanesulfinamide 1, while Table 1 displays the
pseudo first-order rate constants (k1) for the hydrolysis of 2 and 3 in aqueous mineral acids
of different concentrations. All three compounds show similar kinetic profiles, which are
broadly similar to those observed previously for the hydrolysis of N-arylarenesulfinamides,
at relatively low acid concentration.8
However, there are several distinct differences. First, compounds 1–3 are less reactive
toward aqueous mineral acids than the N-arylarenesulfinamides,8 whose rates of hydrolysis
at 25◦C had to be studied by a stopped-flow spectrophotometric method. Even so, in
our study, the hydrolysis of N-phenylisopropanesulfinamide (1) in acid solution is so fast
that rate data was only obtained at very low concentrations of mineral acids. Simpler N-
phenylalkanesulfinamides (R = Me or Et) are even more reactive, and their reactions in
aqueous mineral acids could not be followed by conventional methods.
Second, unlike their arene counterparts, 1–3 exhibit slow, but significant, spontaneous
reaction in pure water. Although a spontaneous reaction was previously observed by Wagner
et al. for the hydrolysis of MeSONMe2 in D2O,11 compounds 1–3 are somewhat more
Figure 1 Pseudo first-order rate constant/acid concentration profile for the hydrolysis of 1 at 44.8◦C.

HYDROLYSIS OF N-PHENYLALKANESULFINAMIDES
567
Table 1 Values of 104 k1 (s−1) for the hydrolysis of 2 and 3 at 44.8◦C
[Acid]/M
1.00
2.00
3.00
4.00
5.00
6.00
2, HClO4
2, H2SO4
3, HClO4
3, H2SO4
2.71
7.95
1.70
4.70∗
—
23.3
4.80
14.0
14.0
58.7
8.33
34.3
23.9
118
16.2
63.4
43.4
203
28.8
107
66.8
321
—
158
[Acid]/M
7.00
8.00
9.00
10.00
11.00
12.00
2, H2SO4
3, H2SO4
433
188
462
200
450
193
455
209
507
275
692
362
[Acid]/M
0.200
0.300
0.400
0.500
0.600
0.800
2, HCl
2, HBr
3, HCl
3, HBr
7.71
24.5
5.15
20.0
—
49.3
—
—
88.8
—
45.3
—
29.2
—
—
418
—
112
—
83.3
330
41.4
72
184
104 ko (s−1) values are: 0.41 (2) and 0.38 (3).
∗In D2SO4/D2O, 104k1 = 8.11 s−1
.
reactive than this compound [104ko (s−1) for 1, 2, and 3 is 0.46, 0.41, and 0.38 at 44.8◦C in
H2O, respectively, compared to 0.28 for MeSONMe2 at 55.5◦C in D2O].
Third, the plots of k1 versus acid concentration for the hydrolysis of 2 and 3 in
H2SO4 are both characterized by a rate maximum occurring at ∼7.5 M-H2SO4, followed
by a minimum and then a final increase.
The catalytic order of acids for the hydrolysis of 1–3, namely HBr > HCl >>
H2SO4 > HClO4, clearly reflects the relatively high nucleophilicity of Br− and Cl−. This
was confirmed by observing the pronounced effect of added LiBr and LiCl on the rate of
hydrolysis of 2 and 3 in 1.00 M HClO4 (Table 2). By comparison, added LiClO4 resulted
in only minor increases in rate, as can be seen from Table 2. Also, the relative order and
magnitude of the nucleophilic attack at sulfinyl sulfur for the hydrolysis of 2 and 3 in
aqueous HBr and HCl (k(Br−)/k(Cl−) ∼ 4.0 for 2 and ∼ 3.0 for 3 are similar to those
observed for N-arylarenesulfinamides (∼4.8)8 and sulfinyl sulfones (∼5.0).6
The order of catalytic power is typical of hydrolysis reactions that proceed via a
bimolecular rate-limiting step, in sharp contrast to A-1 type reactions, such as the ring-
opening hydrolysis of aromatic cyclic sulfamates,12 where the catalytic order is HClO4 >
H2SO4. Additionally, sulfuric acid is typically a better catalyst for A-2 reactions than
Table 2 Effect of added salts on the rate of hydrolysis (104k1 s−1) of 2 and 3 in 1.00 M HClO4 at 44.8◦C
[Salt]/M
Substrate
0.05M-LiBr
0.10 M-LiBr
0.10 M-LiCl
0.20 M-LiCl
0.50 M-LiCl
1.00 M-LiClO4
2∗
3#
39.6
33.8
74.7
63.9
23.6
—
40.3
32.4
113.4
89.1
3.95
2.91
∗104k1 (s−1) in 1.00 HClO4 = 2.71.
#104k1 (s−1) in 1.00 HClO4 = 1.70.

568
M. DATTA ET AL.
perchloric acid. Such behavior has been attributed by Bunton et al. to the differential
stabilization of initial states and transition states by hydrogen sulfate and perchlorate
anions.13,14
Asefi and Tillett8 suggested a general rate equation for the hydrolysis of N-
arylarenesulfinamides in acidic solutions containing halide ions (Eq. 1).
k1 = ko + koꢁ [X−] + kH+[H+] + kHX[H+][X−]
(1)
Here, the terms represent (a) spontaneous (or uncatalyzed) reaction of the neutral species
with water, (b) a nucleophile-catalyzed spontaneous reaction, (c) an acid-catalyzed reaction,
and (d) hydrogen ion-dependent nucleophilic catalysis, respectively. Each term could be
subject to a specific salt effect.
The effect of added lithium salts on the spontaneous reaction of 3 (ko = 0.38 × 10−4
s
−1) was very slight. Thus, the pseudo first order rate constants for the spontaneous reaction
were 0.43 × 10−4 s−1 and 0.45 × 10−4 s−1 with added 0.25 M LiClO4 and 0.50 M LiClO4,
respectively, and 0.46 × 10−4 s−1 and 0.50 × 10−4 s−1 with added 0.25 M LiCl and 0.50
M LiCl, respectively. The differences between these rate constants are of the same order as
the uncertainty of measurement ( 0.05 × 10−4 s−1), and hence for the hydrolysis of 1–3,
the second term can be considered as being unimportant.
Of the other terms of Eq. (1), ko makes only a minor contribution: except at very
low acid concentrations, kH+ is prominent for hydrolysis in HClO4 and H2SO4, and kHX
dominates hydrolysis in HBr and HCl solutions (as supported by the data in Table 2),
especially at higher acid concentrations.
Figure 2 shows a plot of k1 versus acid concentration for the hydrolysis of 2 in
HCl at constant chloride ion concentration (1.50 M, achieved by the presence of LiCl),
superimposed on data from Table 1. The linear plot at constant chloride ion concentration
confirms the contribution of acid-catalysis (kH+) to the overall pseudo first order rate
constant for this compound in HCl.
Another distinct feature is the rate maximum observed in the plots of k1 versus acid
concentration for the hydrolysis of each of 2 and 3 in sulfuric acid, as shown in Table 1
and Figure 3. Similar behavior has been reported for the hydrolysis of many carboxylic
acid derivatives, such as amides15 and esters.16,17 Such maxima can arise for a variety of
Figure 2 Hydrolysis of 2 at 44.8◦C in HCl at constant chloride concentration ([Cl−] = 1.50 M).

HYDROLYSIS OF N-PHENYLALKANESULFINAMIDES
569
Figure 3 Rate profiles for hydrolysis of 2 and 3 in sulfuric acid at 44.8◦C.
reasons, including decrease in the activity of water or superposition of a negative salt effect
on an acid-catalyzed reaction or by extensive conversion of the substrate into its conjugate
acid (“extensive substrate protonation”).15 Since the UV spectra of 3 (at 28.0◦C) in aqueous
1.00 M and 9.00 M H2SO4 were found to be very similar (allowing for spectral changes
caused by the significant reaction in the 9.00 M medium, even at this temperature), then
the latter effect is unlikely to be the cause of this maximum.
Further confirmation of the A-2–like character of the hydrolysis of N-
phenylalkanesulfinamides in dilute sulfuric acid comes from the entropy of activation of 3
in 3.00 M H2SO4 (−48 J K−1 mol−1) (Tables 3 and 4), which falls within the range usually
associated with a bimolecular mechanism.18 The value of the kinetic solvent isotope effect
k1(D2O)/k1(H2O) in 1.00 M D2SO4/H2SO4 for 3 (1.72) also supports this view.19
Table 5 compares the pseudo first-order rate constants (k1) for the hydrolysis of 3 in
aqueous acid and in corresponding acidified 50% acetonitrile/water (v:v). The lower value of
k1 for hydrolysis in HClO4 in the partially organic solvent is consistent with a bimolecular
A-2 type mechanism, reflecting the lower activity of water in the rate-limiting step in
partially organic solvent. On the other hand, the higher values of k1 in 50% MeCN/water
Table 3 Values of 104k1 (s−1) for the hydrolysis of 3 at different temperatures
[Acid]/M
39.8◦C
44.8◦C
51.0◦C
58.3◦C
3.00 M H2SO4
[Acid]/M
20.2
34.2◦C
34.3
39.6◦C
59.6
44.8◦C
108
50.5◦C
11.0 M H2SO4
77.3
159
275
596

570
M. DATTA ET AL.
Table 4 Arrhenius parameters of the hydrolysis of 3
Compound
Acid
[Acid]/M
ꢀH=/kJ mol−1
ꢀS=/J K−1 mol−1
3
3
H2SO4
H2SO4
3.00
11.00
75
99
3
4
−48
7
+46 12
solutions of HBr and HCl are likewise consistent with a bimolecular halide ion catalysis
mechanism, reflecting the higher nucleophilicity of halide ions in partially organic solvent.
In acidic solutions, the most important hydrolysis pathways for 1–3 involve attack
of water or nucleophile on the conjugate acid. Protonation can occur on either oxygen
or nitrogen. NMR evidence for the site of protonation is equivocal,20,21 while related
alkylation and X-ray studies suggest O-protonation.22,23 We therefore adopt O-protonation
for the dominant form, although this must be in equilibrium with the minor N-protonated
form to enable the reaction to occur (Schemes 1 and 2).
The key question regarding nucleophilic substitution at sulfur continues to be whether
the reaction occurs via a classic SN2-type (direct displacement) mechanism involving a
trigonal bipyramidal transition state or via a two-step addition-followed-by-elimination
mechanism, involving a discrete trigonal bipyramidal intermediate. An earlier attempt to
detect the latter by 18O-exchange during the alkaline hydrolysis of arenesulfinamides was
unsuccessful.23 More recently, however, Okuyama et al. have reported 18O-exchange and
a break in the pH-rate profile for the hydrolysis of N-arylbenzenesulfinamides in very
dilute perchloric acid, which they suggest can be accounted for by an intermediate that
can undergo pseudorotation.9 To date, however, no direct evidence has been reported to
distinguish between these possibilities for the hydrolysis of alkanesulfinamides.
The present data can be accommodated by a modified version (Scheme 1) of an
earlier scheme.8
Scheme 1
The final rate increase observed for the hydrolyses of 2 and 3 in concentrated sulfuric
acid (see Table 1 and Figure 3) is similar to that observed for the hydrolysis of certain
carboxylate esters, where it is typically associated with a changeover from an A-2 to an
A-1 mechanism.17 This suggests that, in high concentrations of sulfuric acid, where the

HYDROLYSIS OF N-PHENYLALKANESULFINAMIDES
571
Table 5 Values of 104k1 (s−1) for the hydrolysis of 3 in water and 50% acetonitrile/water (v:v) media at 44.8◦C
1.00 M-HClO4
0.50 M-HCl
0.60 M-HBr
Water
1.70
0.67
29.2
60.0
184
422
50% MeCN/water (v:v)
Scheme 2
activity of water is considerably diminished, an additional reaction pathway is available
to the conjugate acids of 2 and 3 (Scheme 2). Here, rate-limiting decomposition of the
conjugate acid occurs to give a cationic species that is stabilized by the high polarity of the
concentrated sulfuric acid medium.
From these preliminary results, we propose a mechanistic switch from A-2 to A-1 for
the hydrolysis of 2 and 3 in concentrated H2SO4 solutions, which is supported by the much
more positive value of entropy of activation (ꢀS= = +46 J K−1 mol−1) for the hydrolysis
of 3 in 11.0 M H2SO4) (Table 4). Recently, a similar mechanistic changeover has been
suggested for the hydrolysis of arenesulfinamides with a phthalimido leaving group.10
EXPERIMENTAL
Synthesis of Compounds 1–3: General Procedure
Tertiarybutanesulfinyl chloride was obtained from Sigma Aldrich, whereas adaman-
tanesulfinyl chloride and isopropanesulfinyl chloride were prepared by the methods of
Stetter et al.25 and Youn and Herrmann,26 respectively.
Compounds 1–3 were prepared by the method of Stetter et al.,25 using alkanesulfinyl
chloride (5 mmol) and aniline (10 mmol) in anhydrous ether or dry chloroform (30 mL).
After completion of the reaction (monitored using TLC), the white amine salt was filtered
off and the solvent was removed under reduced pressure. Column chromatography (silica
gel; dichloromethane or ethyl acetate:dichloromethane, 1:9) was used to purify the residue.
N-Phenylpropane-2-sulfinamide (1). White crystals. Yield: 95%. Mp:
57–59◦C. FTIR (KBr) (cm−1) 3021, 1601, 1541, 1520, 1497, 1467, 1374, 1275, 1246,
1
1236, 1192, 1073, 1027, 867. H NMR (400 MHz, CDCl3, ppm with respect to TMS) δ
7.24–7.19 (m, 2H), 7.00–6.95 (m, 3H), 6.44 (bs, 1H), 3.05–2.98 (m, 1H), 1.33 (d, J =
6.9 Hz, 3H), 1.29 (J = 6.9 Hz, 3H). 13C NMR (100 MHz, CDCl3, p.p.m. with respect to
TMS) δ 141.7, 129.3, 122.9, 118.2, 54.5, 15.8, 15.4. EIMS m/z (%) 183 (M+, 49), 141
(93), 140 (M+—iPr, 33), 93 (MH+—iPrSO, 100), 92 (88), 91 (11), 78 (83).
2-Methyl-N-phenylpropane-2-sulfinamide (2). White crystals. Yield: 96%.
Mp: 103–104◦C. FTIR (KBr) (cm−1) 3015, 2599, 2330, 1496, 1469, 1420, 1370, 1274,
1068, 1026, 888, 859. 1H NMR (400 MHz, CDCl3, ppm with respect to TMS) δ 7.62–7.22
(m, 2H), 7.00–6.98 (m, 3H), 5.48 (bs, 1H), 1.30 (s, 9H). 13C NMR (100 MHz, CDCl3, ppm

572
M. DATTA ET AL.
with respect to TMS) δ 142.0, 129.4, 123.0, 118.4, 56.5, 22.4. EIMS m/z (%) 197 (M+,
18), 141 (100), 140 (M+—tBu, 28), 105 (M+—PhNH, 86), 92 (M+—tBuSO, 72), 78 (77),
57 (93).
The structure of 2 was confirmed by X-ray crystallography, using a crystal produced
by evaporation of solvent at room temperature from a dichloromethane solution.27
N-Phenyladamantane-1-sulfinamide (3). White crystals. Yield: 97%. Mp:
154–155◦C. FTIR (KBr) (cm−1) 3179, 2908, 2851, 1595, 1488, 1450, 1285, 1228, 1175,
1063, 1034, 877. 1H NMR (400 MHz, CDCl3, ppm with respect to TMS) δ 7.26–7.22 (m,
2H), 7.01–6.97 (m, 3H), 5.43 (bs, 1H), 2.18 (s, 3H), 1.92 (dd, J = 11.8, 22.8 Hz, 6H), 1.74
(dd, J = 12.2, 23.4 Hz, 6H). 13C NMR (100 MHz, CDCl3, ppm with respect to TMS) δ
142.4, 129.2, 122.4, 117.9, 58.2, 36.3, 34.6, 28.5. EIMS m/z (%) 276 (MH+, 39), 275 (M+,
85), 259 (M+—16), 228 (18), 227 (M+—SO, 75), 136 (59), 135 (M+—PhNHSO, 100),
107 (28), 93 (MH+—adamantaneSO, 66), 79 (61).
The structure of 3 was confirmed by X-ray crystallography, using a crystal produced
by evaporation of solvent at room temperature from a dichloromethane solution.28
Product Authentication
Identification of hydrolysis reaction products of N-arylarenesulfinamides, namely
sulfinic acid and amine, has been demonstrated previously by Biasotti and Andersen24
(alkaline hydrolysis) and by Kutuk et al.10 (acid hydrolysis).
In this article, however, the authentic alkanesulfinic acids were unavailable for com-
parison, and in any case they will have λmax of around 200 nm, making them difficult to
distinguish from the anilinium salt by UV spectroscopy. Hence we decided to check the
presence of aniline in the product mixture by TLC. Compound 3 (20 mg) was dissolved in
a 40% methanol/water solution of 2.0 M-HCl (5.0 mL) and kept at 44.8◦C for 18 h. After
cooling to room temperature, the solution was neutralized with aqueous 10.0 M-NaOH
(1.0 mL). TLC analysis (silica gel, CH2Cl2, with UV visualization) confirmed the presence
of aniline (Rf = 0.38). Compound 2 gave a similar result.
Kinetic Measurements
The rates of hydrolysis of 1–3 were measured by following the change in ab-
sorbance at 230 nm, corresponding to disappearance of the substrate, using a JASCO V-530
UV/visible spectrophotometer with a thermostatted cell holder, held at 44.8 0.1◦C, unless
stated otherwise. All determinations were carried at least in duplicate. For all compounds,
0.1 M stock solutions in spectroscopic grade methanol were prepared. An aliquot (4 µL)
of the relevant stock solution was added by syringe to 4.0 mL of the reaction solution
contained in a cuvette and already equilibrated at the desired temperature. The reactions
were monitored over at least three half-lives, and “infinity” values were measured in all
cases, except for the slowest reactions. Pseudo first-order rate constants were calculated
from plots of ln(A − A∞) against time, where A is absorbance at time t and A∞ is the
absorbance at “infinity,” except for the slowest reactions, in which case the method of
Guggenheim was used.29 All acid reaction solutions were prepared from analytical grade
concentrated acids, using deionized water. Acetonitrile was of HPLC quality and all salts
were of analytical grade quality.
The linearities of pseudo first-order plots were good, with R2 better than 0.999, except
for slower and spontaneous reactions (R2 better than 0.995). Reproducibility, measured as

HYDROLYSIS OF N-PHENYLALKANESULFINAMIDES
573
% relative standard deviation (%RSD) for n = 3, was better than 1.5%, except for slower
and spontaneous reactions (better than 5%).
REFERENCES
1. Tillett, J. G. Chem. Rev. 1976, 76, 747–772.
2. Mislow, K.; Siemens, T.; Melillo, J. T.; Tierney, A. L. J. Am. Chem. Soc. 1964, 86, 1452–1453.
3. Landini, D.; Modena, G.; Scorrano, G.; Taddei, F. J. Am. Chem. Soc. 1969, 94, 6703–6707.
4. Bunton, C. A.; de la Mare, P. D. B.; Tillett, J. G. J. Chem. Soc., 1958, 4754–4760; 1959,
1766–1774.
5. Bunton, C. A.; Schwerin, G. J. Org. Chem. 1966, 31, 842–848.
6. Kice, J. L.; Guaraldi, G. J. Am. Chem. Soc. 1967, 89, 4113–4121.
7. Bunton, C. A.; Hendy, B. N. J. Chem. Soc. 1962, 2562–2567.
8. Asefi, H.; Tillett, J. G. J. Chem. Soc. Perkin Trans 2, 1979, 1579–1582.
9. Okuyama, T.; Lee, J. P.; Ohnishi, K. J. Am. Chem. Soc. 1994, 116, 6480–6481.
10. Kutuk, H.; Bekdemir, Y.; Soydas, Y. J. Phys. Org. Chem. 2001, 14, 224–228.
11. Wagner, B. J.; Doi, J. T.; Musker, W. K. J. Org. Chem. 1990, 55, 5940–5945.
12. Bekdemir, Y.; Kutuk, H.; Celik, S.; Tillett, J. G. Phosphorus, Sulfur Silicon Relat. Elem. 2002,
177, 973–980.
13. Bunton, C. A.; Fendler, J. A. J. Org. Chem. 1966, 31, 3764–3771.
14. Bunton, C. A.; Crabtree, J. H.; Robinson, L. J. Am. Chem. Soc. 1968, 90, 1258–1265.
15. Edwards, J. T.; Meacock, S. C. R. J. Chem. Soc. 1957, 2000–2007.
16. Yates, K.; McClelland, R. A. J. Am. Chem. Soc. 1967, 89, 2686–2692.
17. Yates, K. Acc. Chem. Res. 1971, 4, 136–144.
18. Schaleger, L. L.; Long, F. A. Adv. Phys. Org. Chem. 1963, 1, 1–33.
19. Long, F. A.; Pritchard, J. G. J. Am. Chem. Soc. 1956, 78, 2663–2667.
20. Bagno, A.; Eustace, S. J.; Johansson, L.; Scorrano, G. J.Org. Chem. 1994, 59, 232–233.
21. Bujnicki, B.; Drabowicz, J.; Mikołajczyk, M.; Kolbe´, A.; Stefaniak, L. J. Org. Chem. 1996, 61,
7593–7596.
22. Minato, H.; Yamaguchi, K.; Okuma, K.; Kobayashi, M. Bull. Chem. Soc. Japan 1976, 49,
2590–2593.
23. Pickersgill, F.; Marchington, A. P.; Thornton-Pett, M.; Rayner, C. M. J. Chem. Soc. Chem.
Commun. 1995, 647–648.
24. Biasotti, J. B.; Andersen, K. K. J. Am. Chem. Soc. 1971, 93, 1178–1182.
25. Stetter, H.; Krause, M.; Last, W.-D. Chem Ber. 1969, 102, 3357–3363.
26. Youn, J.-H.; Herrmann, R. Tetrahedron Lett. 1986, 27, 1493–1494.
27. Datta, M.; Buglass, A. J.; Elsegood, M. R. J. Acta Cryst. 2009, E65, o2034.
28. Datta, M.; Buglass, A. J.; Hong, C. S.; Lim, J. H. Acta Cryst. 2008, E64, o1393.
29. Guggenheim, E. A. Phil. Mag., 1926, 2, 538–543.