Kinetics and mechanism of N-Boc cleavage: Evidence of a second-order dependence upon acid concentration

Article abstract of DOI:10.1021/jo101767h

The kinetics of the HCl-catalyzed deprotection of the Boc-protected amine, thioester 2 to liberate AZD3409 1 have been studied in a mixture of toluene and propan-2-ol. The reaction rate was found to exhibit a second-order dependence upon the HCl concentration. This behavior was found to have a degree of generality as the deprotection of a second Boc-protected amine, tosylate 3 to yield amine 4 using HCl, sulfuric acid, and methane sulfonic acid showed the same kinetic dependence. In contrast the deprotection of tosylate 3 with trifluoroacetic acid required a large excess of acid to obtain a reasonable rate of reaction and showed an inverse kinetic dependence upon the trifluoroacetate concentration. These observations are rationalized mechanistically in terms of a general acid-catalyzed separation of a reversibly formed ion-molecule pair arising from the fragmentation of the protonated tert-butyl carbamate.

Full text of DOI:10.1021/jo101767h

pubs.acs.org/joc
Kinetics and Mechanism of N-Boc Cleavage: Evidence of a Second-Order
Dependence upon Acid Concentration
Ian W. Ashworth,* Brian G. Cox, and Brian Meyrick
Pharmaceutical Development, AstraZeneca R&D, S41/18 PR&D Building, Silk Road Business Park,
Charter Way, Macclesfield, Cheshire, SK10 2NA, United Kingdom
ian.ashworth@astrazeneca.com
Received September 8, 2010
The kinetics of the HCl-catalyzed deprotection of the Boc-protected amine, thioester 2 to liberate
AZD3409 1 have been studied in a mixture of toluene and propan-2-ol. The reaction rate was found
to exhibit a second-order dependence upon the HCl concentration. This behavior was found to have a
degree of generality as the deprotection of a second Boc-protected amine, tosylate 3 to yield amine 4
using HCl, sulfuric acid, and methane sulfonic acid showed the same kinetic dependence. In contrast
the deprotection of tosylate 3 with trifluoroacetic acid required a large excess of acid to obtain a
reasonable rate of reaction and showed an inverse kinetic dependence upon the trifluoroacetate
concentration. These observations are rationalized mechanistically in terms of a general acid-
catalyzed separation of a reversibly formed ion-molecule pair arising from the fragmentation of
the protonated tert-butyl carbamate.
Introduction
the removal of a Boc protecting group from thioester 2, isopropyl
N-{5-({[(2S,4S)-1-(tert-butoxycarbonyl)-4-(pyridine-3-ylcar-
bonyl)thiopyrrolidin-2-yl]methyl}amino)-2-[2-(4-fluorophenyl)-
ethyl]benzoyl}-L-methioninate, with HCl (5 M) in propan-2-ol
(IPA) (Scheme 1).4 During the development of this process the
performance of the reaction was found to be extremely sensitive
to the HCl usage. A kinetic investigation was undertaken to
understand this sensitivity with a view to defining robust operat-
ing conditions. Additional studies of the deprotection of tosylate
3, tert-butyl 4-[(4-methylbenzene)sulfonyloxymethyl]piperidine-
1-carboxylate, to yield amine 4, 4-[(4-methylbenzene)sulfonyl-
oxymethyl]piperidine (Scheme 2), were undertaken to test the
generality of the surprising findings of the initial investigation.
The tert-butoxycarbonyl (Boc) protecting group is widely
used synthetically for the protection of amines and alcohols.1
Two methods are commonly used to deprotect Boc-protected
amines: either a large quantity of an acid such as trifluoroacetic
acid or a relatively smaller quantity of a stronger acid such as
hydrochloric acid.1,2
AZD3409 1, isopropyl N-{5-({[(2S,4S)-4-(pyridine-3-ylcar-
bonyl)thiopyrrolidin-2-yl]methyl}amino)-2-[2-(4-fluorophenyl)-
ethyl]benzoyl}-L-methioninate, is a farnesyl transferase inhibitor
undergoing investigation for the treatment of breast cancer and
other tumors.3 The final stage in the synthetic sequence involves
(1) (a) Carpino, L. A. Acc. Chem. Res. 1973, 6, 191–8. (b) Greene, T. W.;
Wuts, P. G. M. Protective Groups in Organic Synthesis,3rd ed.; Wiley: New
Results and Discussion
ꢀ
York, 1999; p 518. (c) Kocienski, P. Protecting Groups, 3rd ed.; Thieme: New
York, 2004; p 461.
Kinetic Measurements. Reactions were profiled by HPLC
at a range of temperatures and acid concentrations in toluene
IPA (57% v/v). Raw HPLC peak area versus time data were
normalized by using the toluene peak as an internal stan-
dard, corrected for any difference in response factor between
(2) (a) Mehta, A.; Jaohari, R.; Benson, T. J.; Douglas, K. T. Tetrahedron
Lett. 1992, 33, 5441–5444. (b) Lin, L. S.; Lanza, T., Jr.; de Laszlo, S. E.;
Truong, Q.; Kamenecka, T.; Hagmann, W. K. Tetrahedron Lett. 2000, 41,
7013–7016. (c) Gibson, F. S.; Bergmeier, S. C.; Rapoport, H. J. Org. Chem.
1994, 59, 3216–3218. (d) Strazzolini, P.; Melloni, T.; Giumanini, A. G.
Tetrahedron 2001, 57, 9033–9043.
(3) (a) Stephens, T. C.; Wardleworth, M. J.; Matusiak, Z. S.; Ashton,
S. E.; Hancox, U. J.; Bate, M.; Ferguson, R.; Boyle, T. Proc. Am. Assoc.
Cancer. Res. 2003, 44, R4870. (b) Bell, I. M. J. Med. Chem. 2004, 47, 1869–
1878.
(4) Abbas, S.; Ferris, L.; Norton, A. K.; Powell, L.; Robinson, G. E.;
Siedlecki, P.; Southworth, R. J.; Stark, A.; Williams, E. G. Org. Process Res.
Dev. 2008, 12, 202–212.
DOI: 10.1021/jo101767h
r
Published on Web 11/10/2010
J. Org. Chem. 2010, 75, 8117–8125 8117
2010 American Chemical Society

Ashworth et al.
JOCArticle
SCHEME 1. Deprotection of Thioester 25
SCHEME 2. Test Reaction System Used to Further Investigate
the Removal of the tert-Butoxycarbonyl Protecting Group
the starting Boc-protected amine and the product and con-
verted into concentration versus time data. Repetitive ana-
lyses of reaction samples that had been prepared for analysis
demonstrated that the samples were stable for significantly
longer times than were required to complete the analysis.
The work was carried out under a nitrogen atmosphere in
a 100 mL jacketed reactor fitted with an overhead stirrer
(Hastelloy pitched blade turbine), using an external circulat-
ing bath to maintain the reaction temperature. In a typical
experiment tosylate 3 (3.70 g, 10.0 mmol) was suspended in
toluene (12.89 g, 14.8 mL) and IPA (6.3 g, 8.0 mL) and the
mixture was stirred until the tosylate dissolved. The solution
was heated to 50 °C and a 5 M solution of HCl in IPA (3.0
mL, 15.0 mmol) was charged via a syringe. A 50 μL sample
was withdrawn, diluted into 1:1 acetonitrile:water (10 mL)
and analyzed by HPLC. The experiment was maintained at
50 °C with periodic sampling for HPLC analysis. Typical
concentration versus time profiles are shown in Figure 1.
Deprotection of Thioester 2. Deprotections of thioester 2
were studied with 3.2, 4, and 5 molar equiv of HCl in 57% v/v
toluene IPA at 50 °C. Experiments were also undertaken at
30 °C with 5 and 6 molar equiv of HCl. The results are shown
in Figures 1 and 2 and clearly show the conversion of
thioester 2 to AZD3409 1 without any evidence of a kineti-
cally significant intermediate and that the reaction rate
increases as the HCl concentration increases.
FIGURE 1. Influence of HCl usage on the deprotection of thioester
2 at 50 °C: 5 equiv of HCl thioester 2 (0) and AZD3409 1 (9), 4 equiv
of HCl 2 (]) and 1 ([), and 3.2 equiv of HCl 2 (4) and 1 (2).
the reaction is fast.7 For the reaction to be acid catalyzed the
liberated tert-butyl cation must undergo further reaction to
liberate a proton. In reality the deprotection reaction will
consume an equivalent of acid through the protonation of
the product amine. From this scheme it follows that the rate
law for the deprotection reaction, assuming that little pro-
tonation of the Boc group occurs, is given by eqs 1 and 2.8
d½Pꢀ
dt
d½Bocꢀ
dt
d½Hþꢀ
dt
k1½BocꢀT½Hþꢀ
ðK1 þ ½HþꢀÞ
¼ -
¼ -
¼
ꢁ kobs½BocꢀT½Hþꢀ
ð1Þ
ð2Þ
The widely accepted mechanism for the acid-catalyzed
deprotection of a Boc-protected amine is a rapid pre-equi-
librium protonation of the Boc group, followed by a rate-
limiting fragmentation of the resultant protonated inter-
mediate 5 (Scheme 3).6 In this mechanism it is assumed that
the breakdown of the carbamic acid 6 initially produced by
k1
kobs
¼
K1
Simultaneous fitting of the experimental concentration
versus time data to this rate law (eq 1) was undertaken using
Micromath Scientist at each of the temperatures investigated.9
An initial [HCl] was calculated at each acid usage based on the
(5) The scheme depicts the isolated final product of the deprotection
reaction. The actual product prior to the basic workup is expected to be the
trication arising from the protonation of all of the basic sites (nitrogens)
within the molecule.
(6) (a) Losse, G.; Zeilder, D.; Grieshaber, T. Liebigs Ann. Chem. 1968,
715, 196–203. (b) Hudson, R. F.; Searle, R. J. G.; Mancuso, A. Helv. Chim.
Acta 1967, 50, 997–1002. (c) Lundt, B. F.; Johansen, N. L.; Vølund, A.;
Markussen, J. Int. J. Pept. Protein Res. 1978, 12, 258–268.
(7) Caplow, M. J. Am. Chem. Soc. 1968, 90, 6795–6803.
(8) See the Supporting Information for the derivation of eq 1.
(9) Micromath Scientist, Version 2.0, from Micromath, St. Louis, Mis-
souri. A sample model based on eq 1 for fitting three data sets simultaneously
is given in the Supporting Information.
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FIGURE 2. Influence of HCl usage on the deprotection of thioester
2 at 30 °C: 6 equiv of HCl thioester 2 (0) and AZD3409 1 (9) and
5 equiv of HCl 2 (]) and 1 ([).
FIGURE 3. Best-fit plot for the acid-catalyzed deprotection of
thioester 2 at 50 °C based on the expected mechanism: 5 equiv of
HCl thioester 2 (0) and AZD3409 1 (9), 4 equiv of HCl 2 (]) and 1
([), and 3.2 equiv of HCl 2 (4) and 1 (2).
SCHEME 3. Expected Mechanism for the Acid-Catalyzed
Deprotection of a Boc Protected Amine
plot deviated from linearity was suggestive of a higher order
dependence. Consideration of the initial-rate data in terms of a
second-order acid dependence (Figure 4) strongly suggested
that the rate law for the deprotection reaction obeyed third-
order kinetics, eq 3.
d½Pꢀ
dt
d½Bocꢀ
dt
d½HClꢀ
dt
2
¼ -
¼ -
¼ kobs½BocꢀT½HClꢀ
ð3Þ
Fitting of the experimental data to this rate law was
successful as shown by the best fit lines in Figures 1 and 2,
giving rate constants of 5.7 ( 0.4 ꢂ 10-3 (50 °C) and 5.4 (
0.2 ꢂ 10-4 M-2 s-1 (30 °C).
A brief investigation of the fate of the liberated tert-butyl
cation was undertaken by GCMS. It was found that the main
tert-butyl-derived product was the ether 7 derived from the
trapping of the tert-butyl cation by solvent (IPA).10 Approxi-
mately 10% (based on GC area percent) of the tert-butyl
group underwent trapping by chloride to generate tert-butyl
chloride 8, which is expected to slowly solvolyze under the
reaction conditions.11 Traces of isobutene 9, arising from the
elimination of a proton, were also seen (Scheme 5). No
evidence was seen of the tert-butyl derivative of thioester 2
arising from the trapping of the tert-butyl cation by the
product amine.
SCHEME 4. Protonation of Thioester 2
These results broadly support the hypothesis that the
catalytic proton is regenerated by the onward reaction of
the tert-butyl cation liberated during the deprotection. The
formation of tert-butyl chloride represents a loss of acid term
in the system, which will be most significant kinetically when
close to stoichiometric acid is used for the deprotection. The
inclusion of this pathway as an additional HCl loss term
within the model used to fit the experimental data led to only
a very slight improvement in the goodness of fit. The simple
assumption that 2 molar equiv of acid had been consumed in
protonating the basic groups within thioester 2 (Scheme 4).
The model also adjusted the [HCl] over the course of the
reaction to allow for the protonation of the product amine. It
was clear from the best-fit plots (for example, see Figure 3) that
the experimental data were not described by the expected rate
law.
Reanalysis of the kinetic profiles obtained at 50 °C gave the
initial rate of reaction as a function of the initial free [HCl] and
demonstrated that the initial rate of reaction did not show a
first-order dependence upon the acid concentration (see the
Supporting Information, Figure S1). The manner in which the
(10) Dias, E. L.; Hettenbach, K. W.; am Ende, D. J. Org. Process Res.
Dev. 2005, 9, 39–44.
(11) Grunwald, E.; Winstein, S. J. Am. Chem. Soc. 1948, 70, 846–54.
J. Org. Chem. Vol. 75, No. 23, 2010 8119

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FIGURE 5. Best-fit plots for the deprotection of tosylate 3 with
HCl at 50 °C: 1.5 equiv of tosylate 3 (0) amine 4 (9) and 1.1 equiv of
3 (]) 4 ([).
FIGURE 4. Dependence of the initial rates of thioester 2 deprotec-
tion upon the initial [HCl]2 at 50 °C. The initial [HCl] used allows for
the HCl consumed by the protonation of the basic sites within 2.
SCHEME 5. Fate of the tert-Butyl Group Liberated by the
Deprotection of Thioester 2
TABLE 1. Temperature Dependence of the Third-Order Rate Con-
stants for the Deprotection of Tosylate 3 with HCl in 57% v/v Toluene IPA
[HCl] (M)
temp (°C)
104kobs (M-2 s-1
8.83 ( 0.02a
)
0.582
0.427
0.582
0.582
40
50
50
60
22.8 ( 0.2b
75.0 ( 0.3
aThe quoted errors are the 95% confidence limits of the best-fit rate
constants arising from the fitting. bRate constant determined by simul-
taneously fitting the data from both of the experiments conducted at
50 °C.
catalyst. The deprotection with MSA was also studied in IPA
with no toluene.
Initial experiments at 40, 50, and 60 °C with 1.5 molar
equiv of HCl and 50 °C with 1.1 molar equiv of HCl gave
clean reactions to liberate amine 4. Fitting of the experi-
mental concentration versus time data to a model based on
eq 3 was successful (Figure 5)12 to give the observed third-
order rate constants listed in Table 1. This suggested that the
unexpected second-order kinetic dependence upon the [HCl]
concentration observed in the deprotection of thioester 2
may be more general. The temperature dependence of the
observed rate constants obeyed the Arrhenius equation
(Figure 6) with an activation energy of 93 ( 29 kJmol-1
model, eq 3, was therefore used to generate the quoted rate
constants.
The manufacturing process for AZD3409 1 used a 0.2
molar equiv excess of HCl over the stoichiometric quantity
required to protonate all the basic groups within AZD3409 1.
A clear explanation for the extreme sensitivity of the process
performance to the HCl charge used for the reaction was
therefore provided by the observed second-order depen-
dence in the deprotection kinetics of thioester 2 upon the
[HCl]. Natural errors in charging reagents at scale could lead
to a situation where the actual excess of HCl was lower than
desired, leading to a very slow reaction that would take an
unacceptably long time to meet the end of reaction criteria.
Simulation of the process based on the observed rate law and
experimental rate constants was undertaken to determine the
size of the small increase in the HCl charge required to ensure
robust process performance.
and a pre-exponential factor, A, of 2.5 ꢂ 1012 M-2 s-1
.
Deprotections with H2SO4 (1.8 molar equiv) and MSA
(1.5 molar equiv) at 50 °C behaved similarly to the HCl-
catalyzed reaction. Fitting of the experimental concentration
versus time profiles for these deprotections to a model based
on the experimentally determined rate law (eq 3, substituting
the appropriate acid) was successful as shown in Figure 7.
The best-fit rate constants are tabulated (Table 2) and
compared to the rate constant for the HCl-mediated depro-
tection. Interestingly, the rate constants vary remarkably
Deprotection of Tosylate 3. The reactions of tosylate 3 were
initially studied with the HCl in toluene IPA (57% v/v)
system used for the deprotection of thioester 2. Further
studies were then undertaken with H2SO4, methane sulfonic
acid (MSA), trifluoroacetic acid (TFA), trichloroacetic acid,
dichloroacetic acid, chloroacetic acid, and acetic acid as the
(12) See Figure S2 in the Supporting Information for the best-fit plots at
40 and 60 °C.
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TABLE 2. Third-Order Rate Constants for the Deprotection of Tosy-
late 3 with Different Acids in 57% v/v Toluene IPA at 50 °C
acid
usage (molar equiv)
103kobs (M-2 s-1
)
H2SO4
HCl
MSA
1.8
1.1 and 1.5
1.5
1.5
3.50 ( 0.02a
2.83 ( 0.02
1.60 ( 0.03
1.58 ( 0.10
MSAb
aThe quoted errors are the 95% confidence limits of the best-fit rate
constants arising from the fitting. bRate constant for the deprotection of
3 in IPA.
FIGURE 6. Arrhenius plot for the HCl-catalyzed deprotection of
tosylate 3: rate constant k in M-2 s-1
.
FIGURE 8. Best-fit plot for the deprotection of tosylate 3 with 10
molar equiv ofTFA at 50 °C: rate µ [3][TFA]2, tosylate 3 (]) and
amine 4 ([).
which took 250 min to reach over 90% conversion. A slow
observable reaction was obtained when 10 molar equiv of TFA
was used, which took 5 days to give over 95% conversion at
50 °C. The concentration versus time profile for this reaction
does not fit to the same rate law (eq 3) as the deprotections of
tosylate3catalyzed by H2SO4, HCl, and MSA (see Figure 8) or
to a rate law with a first-order acid dependence eq 1. Deprotec-
tions catalyzed by 10 molar equiv of trichloroacetic, dichlor-
oacetic, chloroacetic, and acetic acid at 50 °C were even slower
with degradation of the tosylate 3 and amine 4 occurring on a
similar time scale to the desired reaction.
Deprotection Mechanism. It is clear from the experimental
data that the kinetic behavior of the deprotection reactions
of thioester 2 and tosylate 3 catalyzed by H2SO4, HCl, and
MSA is not adequately described by the expected mechanism
(Scheme 3), in which a pre-equilibrium protonation is followed
by a rate-limiting fragmentation. Assuming that these acids are
strong (fully dissociated), then the simplest explanation for the
observed second order dependence of these deprotections
upon the [HX] (X- = Cl-, CH3SO3-, HSO4-) is that a second
proton is incorporated into the molecule before or during the
rate-determining transition state. Two simple mechanisms can
be written that include this second proton:
FIGURE 7. Best-fit plots for the deprotection of tosylate 3 with
H2SO4 and MSA at 50 °C: 1.8 equiv of H2SO4, tosylate 3 (0), and
amine 4 (9) and 1.5 equiv of MSA, 3 (]) and 4 ([).
little as the acid is changed within this series. This relative
insensitivity to the acid used makes the behavior observed
with acids such as TFA all the more surprising (see below). In
the case of MSA the rate constant for the deprotection
reaction also appears to show little sensitivity upon changing
the solvent from 57% v/v toluene IPA to IPA.13
The deprotection of tosylate 3 with 1.5 molar equiv of
TFA was attempted at 50 °C and abandoned after a week, as
little reaction had occurred. This represents a dramatic
change in rate relative to the MSA-catalyzed deprotection,
The first of these is an extension of the expected mechan-
ism and involves the formation of a highly reactive diproto-
nated intermediate 10, which then fragments to give a
protonated carbamic acid 11 that looses CO2 to give the
(13) See Figure S3 in the Supporting Information for a comparison of the
best-fit plots for the MSA-mediated deprotection in IPA and 57% v/v
toluene IPA.
J. Org. Chem. Vol. 75, No. 23, 2010 8121

Ashworth et al.
SCHEME 8. Mechanism for N-Boc cleavage via an Acid-
JOCArticle
SCHEME 6. Mechanism for N-Boc Cleavage via a Diproto-
nated Intermediate
Mediated Separation of an Ion-Molecule Pair
SCHEME 7. Mechanism for N-Boc Cleavage via a Bimolecular
Acid-Catalyzed Fragmentation of a Protonated Intermediate
(Scheme 7). The rate law for this mechanism, eq 6, may be
derived in an analogous manner to eq 4 and simplified to give
a rate law, eq 7, with the same form as the experimentally
observed rate law. While this mechanism avoids the genera-
tion of a relatively implausible intermediate, it does lead to
the generation of two high-energy species in the rate-deter-
mining step, a protonated carbamic acid 11 and a tert-butyl
cation, which make it less than satisfactory.
2
d½Pꢀ
dt
k3½Bocꢀ ½Hþꢀ
ðK1 þ T½HþꢀÞ
¼
ð6Þ
2
d½Pꢀ k3½BocꢀT½Hþꢀ
ꢁ
ð7Þ
dt
K1
protonated product amine (Scheme 6). Derivation of the rate
law for a reaction according to this mechanism (treating the
protonations as rapid pre-equilibria) gives eq 4, which may
be simplified (eq 5) based on the assumption that very little of
the diprotonated intermediate is formed (i.e., K1K2 is the
dominant term in the denominator). This rate law has the
same form as the experimentally observed rate law. How-
ever, experimental studies of alkyl carbamates in super acidic
media at low temperatures have failed to find any evidence of
diprotonated intermediates of the form of 10,14 although
protonated carbamic acids similar to 11 have been observed
by NMR at low temperatures under strongly acidic
conditions.15 It is therefore unlikely that this mechanism is
operating under the significantly less acidic deprotection
conditions.
A third mechanistic possibility also exists,16 which is also a
modification of the expected mechanism. In this proposal the
protonated carbamate 5 undergoes a reversible fragmenta-
tion to give an ion-molecule pair 12 that rapidly recombines
to regenerate the protonated carbamate 5. The progress of
the deprotection reaction is driven by a rate-limiting proton-
ation of the carbamic acid within the ion-molecule pair,
which drives the separation of the ion-molecule pair to give
the protonated carbamic acid 11 and a tert-butyl cation. This
is analogous to the behavior observed in the acid-catalyzed
deprotection of trityl-protected amines.17 The rate law for
this mechanism may be derived based on the assumption of a
rapid pre-equilibrium protonation (eq 8). Simplification
is possible if it is assumed that little of the protonated
d½Pꢀ
dt
d½Bocꢀ
dt
d½Hþꢀ
¼ -
¼ -
dt
(16) The authors would like to acknowledge the referees’ helpful sugges-
tions with respect to this mechanism.
(17) (a) Bleasdale, C.; Golding, B. T.; Lee, W. H.; Maskill, H.; Risebor-
ꢀ
ough, J.; Smits, E. J. Chem. Soc., Chem. Commun. 1994, 93–94. (b) Moises,
C. L.; Ibrahim, D.; Maskill, H. J. Chem. Soc., Perkin Trans. 2 2001, 1748–
1752.
(18) Clare, B. W.; Cook, D.; Ko, E. C. F.; Mac, Y. C.; Parker, A. J. J. Am.
Chem. Soc. 1966, 88, 1911–1916.
(19) Kolthoff, I. M.; Chantooni, M. K., Jr. J. Phys. Chem. 1979, 83, 468–
474.
2
k2½BocꢀT½Hþꢀ
¼
ð4Þ
ð5Þ
2
ðK1K2 þ K1½Hþꢀ þ ½Hþꢀ Þ
2
d½Pꢀ k2½BocꢀT½Hþꢀ
ꢁ
dt
K1K2
(20) Kanning, E. W.; Bobalek, E. G.; Byrne, J. B. J. Am. Chem. Soc. 1943,
65, 1111–1116.
(21) Perrin, D. D. Ionisation Constants of Inorganic Acids and Bases in
Aqueous Solution, IUPAC Chemical Data Series No. 29, 2nd ed.; Pergamon:
Oxford, UK, 1982.
(22) Serjeant, E. P.; Dempsey, B. Ionisation Constants of Organic Acids in
Aqueous Solution, IUPAC Chemical Data Series No. 23, Pergamon: Oxford,
UK, 1979.
(23) Chantooni, M. K., Jr.; Kolthoff, I. M. Anal. Chem. 1979, 51, 133–
140.
The second potential mechanism again builds on the
expected mechanism and involves a bimolecular acid-cata-
lyzed fragmentation of the protonated intermediate 5 to give
the same protonated carbamic acid 11 as the first mechanism
(14) Olah, G. A.; Heiner, T.; Rasul, G.; Prakash, G. K. S. J. Org. Chem.
1998, 63, 7993–7998.
(15) Olah, G. A.; Calin, M. J. Am. Chem. Soc. 1968, 90, 401–404.
8122 J. Org. Chem. Vol. 75, No. 23, 2010

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SCHEME 9. Mechanism for N-Boc Cleavage via a General Acid-Catalyzed Separation of an Ion-Molecule Pair
SCHEME 10. Mechanism for N-Boc Cleavage via a General Acid-Catalyzed Separation of an Ion-Molecule Pair Resulting from the
Fragmentation of an Intermediate Ion Pair
TABLE 3. Comparison of Catalyst pKa Values in Methanol, IPA, and
Water
carbamate 5 is formed and that the collapse of the
ion-molecule pair 12 back to the protonated carbamic acid
is rapid compared to product formation. This gives rise to a
rate law with the same form as the experimentally observed
rate law (eq 9).
acid
pKa MeOH
pKa IPA
3.1b
pKa water
HCl
H2SO4
MSA
TFA
CCl3CO2H
CHCl2CO2H
CH2ClCO2H
CH3CO2Ha
1.23a
1.5c
-6.3a
-3.6d
-1.8e
0.5e
2
d½Pꢀ
dt
k1k4½BocꢀT½Hþꢀ
0.5e
¼
ð8Þ
6.4a
7.7a
9.6a
7.8f
9.2f
1.29a
2.86a
4.76a
ðk- 1 þ k4½HþꢀÞðK1 þ ½HþꢀÞ
11.3f
2
d½Pꢀ k1k4½BocꢀT½Hþꢀ
aSee ref 18. bSee ref 19. cSee ref 20. dSee ref 21. eSee ref 22. fSee ref 23.
ꢁ
ð9Þ
dt
k- 1K1
Differentiation between the three mechanisms discussed is
not possible based on the available information, as all are
described by the experimentally determined rate law. Addi-
tionally the mechanisms give rise to composite observed rate
constants, which makes any mechanistic analysis based on
activation parameters derived from the temperature depen-
dence of the rate constants problematic. Of the three, the
pathway containing an ion-molecule pair is the most ap-
pealing, as it does not require the formation of a species that
has not been experimentally observed and has a reasonable
precedent.17
dissociated). Table 3 lists the acid strengths in methanol,
IPA, and water of several of the catalytic acids used in this
study. All become less acidic on transfer to methanol, and
even more so in IPA. In particular, pKa(HCl) = 3.1 in IPA,
and hence in the concentrated solutions used for the kinetic
studies, it will only be partially dissociated.
It is possible to write an analogous mechanism to that in
Scheme 8 for catalysis by a weak acid (i.e., the acid is not fully
dissociated), which makes the rate-limiting protonation to
drive the separation of the ion-molecule pair a case of
general acid catalysis, as shown in Scheme 9.
The mechanism depicted in Scheme 8 and eq 8 predicts
that the rate constant for all three acids should be equal, but,
while similar, they show a trend to lower values as the
strength of the acids decreases from H2SO4 to MSA.
Furthermore, in the solvents used (IPA/toluene or IPA) the
acids are expected to be relatively weak (i.e., not fully
Such a mechanism leads to a rate law, eq 10, which may be
simplified by assuming that only a small amount of the
protonated intermediate is generated (i.e., K3[X-] . [HX])
and that the separation of the ion-molecule pair is slow
relative to its collapse back to the protonated carbamate
to give eq 11. This has an inverse dependence upon the
J. Org. Chem. Vol. 75, No. 23, 2010 8123

Ashworth et al.
JOCArticle
conjugate base of the catalytic acid, X-, and is therefore not
consistent with the rate law, eq 3, observed with HCl, H2SO4,
and MSA.
2
d½Pꢀ
dt
k1k5½BocꢀT½HXꢀ
¼
ð10Þ
ðk- 1 þ k5½HXꢀÞðK3½X- ꢀ þ ½HXꢀÞ
2
d½Pꢀ k1k5½BocꢀT½HXꢀ
ꢁ
ð11Þ
dt
k- 1K3½X-
ꢀ
It is known, however, that ion-pair association constants
for salts in IPA are very high (typically >103).19 This
suggests a modification to Scheme 9 in which the initial
protonated intermediate exists as an ion-pair 13 rather than
free ions, giving rise to Scheme 10. The rate law (eq 12) based
on this scheme is in agreement with the observed rate law,
eq 3. This mechanism is also consistent with the modest
variation in observed rate constants for HCl, H2SO4, and
MSA. A high free concentration of the conjugate base of the
catalytic acid does not build up as the reaction proceeds due
to ion-pairing with the protonated amine produced by the
deprotection reaction. In the case of thioester 2 the proton-
ation of the basic sites within the molecule is also likely to
lead to ion-pairing rather than free ions.
FIGURE 9. Best-fit plot for the TFA-catalyzed deprotection of
tosylate 3 at 50 °C: rate µ [3][TFAH]2/[TFA-], tosylate 3 (]) and
amine 4 ([).
2
trying to develop efficient deprotection reactions this higher
order kinetic dependence has significance, as it means the
reaction tail will be extremely sensitive to the level of acid
remaining toward the end of the reaction. Reaction rates among
the stronger acids studied exhibit an order consistent with their
expected order of acidities: H2SO4 > HCl > MSA. The dif-
ferences are, however, only modest covering a range of a factor
of 2. Weaker acids, such as TFA, and especially other, weaker
carboxylic acids, require much larger concentrations and excesses
in order to achieve practical reaction times. They appear to
exhibit similar behavior, but it is modified by a reciprocal depen-
dence upon the concentration of the conjugate base of the acid.
Mechanistically, a scheme consisting of a general acid-
catalyzed separation of an ion-molecule pair arising from
d½Pꢀ k6k7½BocꢀT½HXꢀ
ꢁ
ð12Þ
dt
k- 6K4
Finally, we consider the TFA-catalyzed deprotection of
tosylate 3. The reaction is extremely slow compared with the
HCl-catalyzed reaction, for example, and even at 50 °C
required a large excess (10 equiv) to achieve a reasonable
conversion over 60 h. The high excess of TFA used in this
experiment was such as to make it a significant part of the
solvent system. Additionally it effectively meant that
[TFAH] did not change during the reaction, meaning that
it was not possible to distinguish effectively between rate
laws with zero-, first-, and second-order dependencies upon
[TFAH]. The analysis in Figure 8, however, shows that a
simple rate law of the form of eq 3 was not able to provide an
effective fit to the data, with the reaction becoming progressively
slower than expected as it proceeded. A significantly better fit
was achieved if an inverse dependence upon [TFA-] was
included as in eq 11 (see Figure 9), with a composite second-
order rate constant of 2.1 ( 0.2 ꢂ 10-7 M-1 s-1 at 50 °C. A rate
law with a second-order dependence upon [TFAH] was used, as
this had a degree of congruence with the behavior observed with
the other acids and also gave a marginally better fit than rate
laws with zero- or first-order dependencies upon [TFAH]. It is
possible that the high [TFAH] was sufficient to provide stabi-
lization to TFA- through H-bonding (homohydrogen-bond
formation), enabling TFA- to exist as a “free” ion and hence
adversely affect the pre-equilibrium protonation of tosylate 3 as
the reaction proceeds.
TABLE 4. HPLC Method for Thioester 2 and AZD3409 1
column
Luna C-18
5 cm ꢂ 2.0 mm, 2.0 μm
40 °C
temp
wavelength
flow
sample injection vol
run time
post time
265 nm
0.8 mL/min
3 μL
12 min
2 min
Gradient
% CH3CN
time (min)
% water
1% trifluoroacetic
acid soln
0
7
8
50
0
50
40
90
40
10
10
10
Conclusions
Sample Preparation: Take 50 μL of Reaction Mixture into a 10 mL
Volumetric Flask, Dilute to Volume with 1:1 Aqueous Acetonitrile
This paper has demonstrated that removal of the Boc group
routinely used for the protection of amines can exhibit a second-
order kinetic dependence upon the concentration of the acid
used to catalyze the reaction. The resultant third-order rate law
has a degree of generality as it has proved applicable to two
different substrates with three different acids. Naturally when
component
retention time (min)
AZD3409 1
toluene
thioester 2
1.4
1.7
5.0
8124 J. Org. Chem. Vol. 75, No. 23, 2010

Ashworth et al.
JOCArticle
TABLE 5. HPLC Method for Tosylate 3 and Amine 4
Tosylate 3, tert-butyl 4-[(4-methylbenzene)sulfonyloxymethyl]-
piperidine-1-carboxylate, was commercial material supplied by
Isochem. Thioester 2, isopropyl N-{5-({[(2S,4S)-1-(tert-butoxycar-
bonyl)-4-(pyridine-3-ylcarbonyl)thiopyrrolidin-2-yl]methyl}amino)-
2-[2-(4-fluorophenyl)ethyl]benzoyl}-L-methioninate and AZD3409
1, isopropyl N-{5-({[(2S,4S)-4-(pyridine-3-ylcarbonyl)thiopyrrolidin-
2-yl]methyl}amino)-2-[2-(4-fluorophenyl)ethyl]benzoyl}-L-methio-
ninate were prepared according to the previously described pro-
cedure.4 Its identity was confirmed by HPLC analysis (see Table 4)
against an authentic sample.
Preparation of 4-Piperidylmethyl 4-Methylbenzenesulfonate
(4) Methane Sulfonate Salt. A mixture of tosylate 3 (7.40 g, 1.0
molar equiv), toluene (48.1 mL), and propan-2-ol (3.7 mL) was
stirred and heated to 50 °C under a nitrogen atmosphere to give a
clear solution. Methane sulfonic acid (2.89 g, 1.5 molar equiv) was
charged and the resultant two-phase system was maintained at
50 °C overnight. Reaction completion was confirmed by HPLC
analysis (see Table 5). The reaction mass was cooled to 10 °C over
2 h and stirred at 10 °C for 2 h, after which the white crystalline
solid was isolated by filtration under reduced pressure, washed
with propan-2-ol (2 ꢂ 10 mL), and dried in vacuo at 40 °C. The
unoptimized yield of amine 4 as its methane sulfonate salt was 3.2 g
(43%): 1H NMR (400 MHz, DMSO-d6) δ (ppm) 1.31 (m, 2 H),
1.73 (d, J = 12.8 Hz, 2 H), 1.94 (m, 1 H), 2.35 (s, 3 H), 2.43 (s, 3 H),
2.84 (m, 2 H), 3.24 (d, J = 12.5 Hz, 2 H), 3.93 (d, J = 6.0 Hz, 2 H),
7.50 (d, J = 7.9 Hz, 2 H), 7.80 (d, J = 8.4 Hz, 2 H), 8.18 (br s, 1 H),
8.57 (br s, 1 H); 13C NMR (100.6 MHz, DMSO-d6) δ (ppm) 21.1,
24.4, 32.6, 39.7, 42.5, 73.5, 127.6, 130.2, 132.2, 145.0; MS (ES) m/z
(M þ Hþ) 270.1166 vs 270.1158 calcd for C13H20NO3S.
column
Zorbax SB C-18
5 cm ꢂ 2.0 mm, 1.8 μm
temp
wavelength
flow
sample injection vol
run time
post time
40 °C
256 nm
1.5 mL/min
3 μL
8 min
2 min
Gradient
% CH3CN
time (min)
% water
1% trifluoroacetic
acid solution
0
85
85
0
0
85
5
5
90
90
5
10
10
10
10
10
0.2
5.2
6.4
6.5
Sample Preparation: Take 50 μL of Reaction Mixture into a 10 mL
Volumetric Flask, Dilute to Volume with 1:1 Aqueous Acetonitrile
component
retention time (min)
amine 4
toluene
tosylate 3
2.95
4.4
5.1
the reversible fragmentation of the protonated Boc group is
proposed to rationalize the observed kinetic behavior.
Acknowledgment. We thank Dr. G. E. Robinson, Dr.
L. Ferris, Mrs. J. Cherryman, Dr. A. T. Bristow and Dr.
I. C. Jones for discussions and for help with synthetic and
analytical work.
Experimental Section
Materials. Acetonitrile and propan-2-ol (IPA) used in the
kinetic experiments were HPLC grade solvents, while the to-
luene was reagent grade. HCl in propan-2-ol was purchased as a
5.0 M solution from Acros and analyzed by titration prior to
use. Acetic acid, chloroacetic acid, dichloroacetic acid, methane
sulfonic acid, sulfuric acid, and trichloroacetic acid were of the
highest commercial quality available and were used without
further purification. HPLC grade trifluoroacetic acid was used
for both the HPLC analysis and kinetic experiments.
Supporting Information Available: Derivation of eq 1, sam-
ple Micromath Scientist model file, plot of initial rate vs [HCl], for
2, best-fit plot for the deprotection of 3 at 40 and 60 °C, best-fit plot
for the deprotection of 3 with MSA in IPA, and 1H and 13C NMR
spectra of compound 4. This material is available free of charge via
the Internet at http://pubs.acs.org.
J. Org. Chem. Vol. 75, No. 23, 2010 8125