Enhancement of the Oxidizing Power of an Oxoammonium Salt by Electronic Modification of a Distal Group

Article abstract of DOI:10.1021/acs.joc.7b01965

The multigram preparation and characterization of a novel TEMPO-based oxoammonium salt, 2,2,6,6-tetramethyl-4-(2,2,2-trifluoroacetamido)-1-oxopiperidinium tetrafluoroborate (5), and its corresponding nitroxide (4) are reported. The solubility profile of 5 in solvents commonly used for alcohol oxidations differs substantially from that of Bobbitt's salt, 4-acetamido-2,2,6,6-tetramethyl-1-oxopiperidinium tetrafluoroborate (1). The rates of oxidation of a representative series of primary, secondary, and benzylic alcohols by 1 and 5 in acetonitrile solvent at room temperature have been determined, and oxoammonium salt 5 has been found to oxidize alcohols more rapidly than does 1. The rate of oxidation of meta- and para-substituted benzylic alcohols by either 1 or 5 displays a strong linear correlation to Hammett parameters (r > 0.99) with slopes (ρ) of -2.7 and -2.8, respectively, indicating that the rate-limiting step in the oxidations involves hydride abstraction from the carbinol carbon of the alcohol substrate.

Full text of DOI:10.1021/acs.joc.7b01965

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Article
Enhancement of the Oxidizing Power of an Oxoammonium
Salt by Electronic Modification of a Distal Group
Kyle M Lambert, Zachary D. Stempel, Sadie M Kiendzior, Ashley L Bartelson, and William F Bailey
J. Org. Chem., Just Accepted Manuscript • DOI: 10.1021/acs.joc.7b01965 • Publication Date (Web): 02 Oct 2017
Downloaded from http://pubs.acs.org on October 5, 2017
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Enhancement of the Oxidizing Power of an Oxoammonium Salt by
Electronic Modification of a Distal Group
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Kyle M. Lambert,†,* Zachary D. Stempel,† Sadie M. Kiendzior,§ Ashley L. Bartelson,‡,* and
William F. Bailey†,*
†
Department of Chemistry, University of Connecticut, Storrs, Connecticut 06269-3060, United
States
§ Department of Chemistry, University of Saint Joseph, West Hartford, Connecticut 06117-2791,
United States
‡
Department of Chemistry, Seton Hill University, Greensburg, Pennsylvania 15601, United
States
O
O
HN
CH3
HN
CF3
VS
N
O
N
O
BF4
BF4
R
R
OH
R
O
O
1.7 – 3.2 times
faster
ABSTRACT: The multi-gram preparation and characterization of a novel TEMPO-based
oxoammonium salt, 2,2,6,6-tetramethyl-4-(2,2,2-trifluoroacetamido)-1-oxopiperidinium
tetrafluoroborate (5), and its corresponding nitroxide (4), are reported. The solubility profile of 5
in solvents commonly used for alcohol oxidations differs substantially from that of Bobbitt’s salt,
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4-acetamido-2,2,6,6-tetramethyl-1-oxopiperidinium tetrafluoroborate (1). The rates of oxidation
of a representative series of primary, secondary and benzylic alcohols by 1 and 5 in acetonitrile
solvent at room temperature have been determined and oxoammonium salt 5 has been found to
oxidize alcohols more rapidly than does 1. The rate of oxidation of meta and para-substituted
benzylic alcohols by either 1 or 5 displays a strong linear correlation to Hammett parameters (r >
0.99) with slopes (ρ) of –2.7 and –2.8, respectively, indicating that the rate-limiting step in the
oxidations involves hydride abstraction from the carbinol carbon of the alcohol substrate.
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INTRODUCTION
Stable nitroxide radicals have drawn considerable interest among chemists for use as spin
labels, antioxidants, and mediators in polymerization reactions.1 They have also found a growing
role in oxidation chemistry as precursors to oxoammonium cations.2 Such cations may be
isolated as stable salts, such as 4-acetamido-2,2,6,6-tetramethyl-1-oxopiperidinium
tetrafluoroborate (1, Scheme 1), more commonly known as “Bobbitt’s salt”,3 or they may be
generated from a stable nitroxide, such as 4-acetamido-TEMPO (2), in a catalytic cycle (Scheme
1) involving the corresponding hydroxylamine (3). A number of stable nitroxide species have
been identified,4 the most well-known of which is 2,2,6,6-tetramethylpiperidine 1-oxyl (TEMPO,
Figure 1). The stability of such radicals arises from the absence of an α–hydrogen on the carbons
adjacent to the nitrogen possessing the oxygen radical and, in the case of TEMPO and its
analogues, additional kinetic stabilization of the radical is achieved by the steric constraints
imposed by the four neighboring methyl groups.
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Scheme 1. Redox Cycle of Oxoammonium Salt Oxidants
5
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9
Nitroxide catalyst
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AcHN
N
O
2
1e– oxidation
1e– oxidation
BF4
O
AcHN
Oxoammonium
cation
AcHN
Hydroxylamine
N
N
OH
1
3
2e– oxidation
R
OH
R
O
In an effort to design better catalysts for alcohol oxidation, less hindered nitroxides, such
as those shown in Figure 1, have been prepared.5 While modification of the structural motif of
nitroxide catalysts has been a primary focus of catalyst design, a less commonly addressed factor
is the potential influence of remote substituents on the oxidizing capability of the oxoammonium
cations derived from nitroxides.
O
N
N
N
N
O
O
O
O
ABNO
KetoABNO
AZADO
TEMPO
Figure 1. Selected nitroxides.
Several reports have noted that when the steric difference between two nitroxide catalysts
is negligible, catalytic activity correlates with the nitroxide / oxoammonium cation redox
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potential and, in general, nitroxides bearing electron-withdrawing substituents were found to
have higher redox potentials. Kanai and co-workers, for example, reported that ketoABNO was
more effective than ABNO in the copper-catalyzed oxidation of secondary amines to imines and
they attributed this observation to the electron-withdrawing ability of the carbonyl group.6 More
recently, computational studies by Coote and Iwabuchi,7 as well as structure-activity
relationships reported by both Stahl8 and Sigman,9 demonstrated that electron-donating
substituents appended to a nitroxide decrease its redox potential while electron-withdrawing
substituents increase the potential. These results correlate well with experimental measurements
of nitroxide / oxoammonium cation redox potentials.
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As a result of our interest in the use of Bobbitt’s salt (1) for a variety of oxidations,10 we
made note of a 2014 paper by Szolcsányi and coworkers reporting that a TEMPO-based
nitroxide containing a 4-trifluoromethylacetamido group (4) exhibits a large difference in anodic
and cathodic peak potentials, almost double that of the 4-acetamido derivative (2).11 We were
interested in exploring whether the oxoammonium salt derived from this CF3 - containing
nitroxide (4) would oxidize alcohols more rapidly than does Bobbitt’s salt (1). As demonstrated
by the results presented below, this electronic modification of the amide moiety at the 4-position
of 1 leads to a significant increase in the rate of oxidation of alcohols.
RESULTS AND DISCUSSION
The 4-trifluoroacetamido oxoammonium salt 5 was readily prepared on a multi-gram
scale in good overall yield from 4-amino-2,2,6,6-tetramethylpiperidine and trifluoroacetic
anhydride as illustrated in Scheme 2 by modification of the procedure used for the preparation of
1.3 Oxoammonium salt 5 is isolated as a yellow, crystalline solid that melts with decomposition
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at 160°C. When left open to ambient air on the lab bench, samples of 5 turned white in color
over the course of a few days. This white solid was identified as 1-hydroxy-2,2,6,6-tetramethyl-
4-(2,2,2-trifluoroacetamido)piperidinium tetrafluoroborate (6), the protonated, reduced form of
5, by comparison to an authentic sample. Initially it was thought that 6 resulted from oxidation of
adventitious moisture by 5. Endo for example, has reported that that the oxoammonium bromide
salt derived 4-methoxy-TEMPO is reduced over the course of a few hours when dissolved in
basic water.12 However, 5 is stable for weeks when dissolved in water and sealed under argon;
thus, it is not likely that residual moisture is the source of decomposition of 5. Indeed, 5 can be
stored for periods of several months without degradation in a sealed plastic container under an
atmosphere of argon or nitrogen.
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Scheme
2.
Preparation
of
2,2,6,6-Tetramethyl-4-(2,2,2-trifluoroacetamido)-1-
oxopiperidinium tetrafluoroborate (5) and its Corresponding Nitroxide (4)
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O
O
5
6
7
8
NH2
HN
N
CF3
HN
CF3
1. cat. Na2WO4•2H2O
2. cat. EDTA
3. 6 eq H2O2
O
O
F3C
O
CF3
Na2CO3
N
H
9
N
H
ether
0°C, 1h
36h, RT
H2O
15 min, RT
H2O
H
H
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CF3COO
99%
O
O
HN
CF3
HN
CF3
1. 0.5 eq aq. HBF4
2. 0.5 eq aq. NaOCl
3. 1 eq NaBF4
N
O
N
O
RT, 4h
H2O
BF4
4
83%
5
73%
multi-gram scale
The solubilities of oxoammonium salts 1 and 5 in various solvents were determined and
the results are summarized in Table 1. On the whole, the 4-trifluoromethylacetamido derivative 5
is more soluble in polar organic solvents such as CH2Cl2, EtOAc, and CH3CN than is the parent
4-acetamido salt 1. In water, the situation is reversed; 1 is some 2.7 times more soluble than is 5.
With the relative solubility of the two oxoammonium salts in hand, it was of interest to evaluate
the relative rate of oxidation of a variety of alcohol substrates by each of the two salts.
Table 1. Solubility of Oxoammonium Salts 1 and 5 in Various Solvents.a
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O
O
5
6
HN
CH3
HN
CF3
7
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solventb
N
O
N
O
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BF4
BF4
1
5
cyclohexane
toluene
insoluble
0.10 (3.3 x 10–4 M)
insoluble
insoluble
0.10 (2.8 x 10–4 M)
diethyl ether
dichloromethane
ethyl acetate
acetonitrile
water
insoluble
0.50 (1.7 x 10–3 M )
0.75 (2.1 x 10–3 M)
0.85 (2.4 x10–3 M)
61.7 (0.17 M)
insoluble
36.0 (0.12 M)
70.4 (0.23 M)
30.4 (0.086 M)
a
Solubility is expressed as grams of solute per liter of solvent and, within
parentheses, as molarity. b All organic solvents were dried prior to use by either
distillation from a dark purple solution of Na/benzophenone or calcium hydride
or MgSO4.
The colorimetric properties of oxoammonium salts allows for determination of the rate of
disappearance of each oxoammonium salt during an oxidation by UV-Vis spectroscopy.
Although both salts 1 and 5 are yellow in color, each has a unique UV-Vis spectrum in
acetonitrile solvent as illustrated in Figures 2 and 3 for salts 1 and 5, respectively. There is,
however, substantial overlap in the spectra of the two salts. This overlap was dealt with, as
detailed in the Experimental Section, by employing the method described by Blanco and
coworkers13 with further simplification by application of a least-squares analysis of the spectra
using Excel Solver as described by Harris.14
The oxidation of an alcohol by an oxoammonium salt is clearly a second-order process.
However, the kinetic analysis was simplified by employing a 100-fold excess of alcohol relative
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to the oxoammonium salt. Under these conditions, the oxidations were found to be cleanly
pseudo-first order in each salt, and plots of ln [salt] vs. time (s) gave a linear fit with the slope =
– k. Rate studies were conducted in competition using both salts in approximately equal
concentration; accurately determined individual concentrations of each salt in CH3CN (~0.005
M) afforded a total oxoammonium salt concentration of ~0.01 M. The individual concentrations
of each salt were monitored during alcohol oxidations at wavelengths between 370 nm and 530
nm: both 1 and 5 exhibit local absorption maxima in this wavelength range (1 at 456 nm and 475
nm; 5 at 460 nm, and 479 nm). Neither the alcohol substrates used in the study nor the
protonated hydroxylamines (the HBF4 salt of 3) and (6), the co-products of the oxidation, absorb
within this range, ensuring minimal interference. Control experiments, detailed in the Supporting
Information, demonstrated that aldehydes or ketones were the only products generated in the
oxidations by either salt.
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Figure 2. UV-Vis spectrum of Bobbitt’s salt (1) in acetonitrile as function of concentration.
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Figure 3. UV-Vis spectrum of oxoammonium salt 5 in acetonitrile as a function of
concentration.
The Supporting Information includes UV-Vis spectra used in the kinetic analyses,
determination of molar absorptivities, as well as plots of ln ([1]) and ln ([5]) vs. time (s) for all
substrates investigated. A representative example of the kinetic data is shown below for the
oxidation of p-methoxybenzyl alcohol (Figure 4); the pseudo-first order rate constants for
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oxidation by 1 and 5 were determined from the slope of the linear least squares fit of these data
5
6
(Figure 5).
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Figure 4. Change in the UV-Vis spectra of equimolar quantities of both salts (1 and 5) during the
oxidation of p-methoxybenzyl alcohol.
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Figure 5. Rates of oxidation of p-methoxybenzyl alcohol by salts 1 and 5.
The individual pseudo-first order rate constants for oxidation of a representative series of
alcohols in CH3CN solvent by 1 and 5, as well as the half-times for the oxidations and the
relative rates of oxidation of each alcohol by 1 and 5, are summarized in Tables 2 and 3. The
estimated error in the rate constants is < 5%. In all cases, 5 was found to oxidize a given alcohol
at a rate generally 1.7 – 3.2 times faster than 1. Table 2 summarizes the data for various aliphatic
and unsaturated alcohols and, as expected,15 α,β-unsaturated or benzylic substrates are oxidized
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more quickly than aliphatic alcohols and secondary alcohols are oxidized faster than are
unactivated primary alcohols by either oxoammonium salt. Not surprisingly, there is a
considerable steric component that influences the rate of oxidation of aliphatic alcohols:
neopentyl alcohol is oxidized 5 times slower by 1 and 4 times slower by 5 than is 1-heptanol.
The oxidation of cinnamyl alcohol is something of an anomaly as it is apparently oxidized ~18
times faster by 5 than 1. The origin of this significant rate difference is currently unknown, but it
may simply be an artifact of the very rapid oxidation of the substrate by 5 having a half-time of
only ~3 s.
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Table 2. Rates of Oxidation of Various Aliphatic and Unsaturated Alcohols by 1 and 5 in
CH3CN Solvent.
O
O
HN
CF3
HN
CH3
Vs.
N
O
N
O
BF4
BF4
5
1
R
O
R
OH
rate of
rate of
relative
rate of
oxidation
(5/1)
half-time of
oxidation by
Bobbitt’s salt 1
half-time of
oxidation by
CF3 salt 5
oxidation by
Bobbitt’s salt 1
(k x 104 s-1)
oxidation by
CF3 salt 5
alcohol
(k x 104 s-1)
1-heptanol
2-hexanol
1.73
9.14
66.8 min
12.6 min
3.71
1.67
31.1 min
6.9 min
2.2
1.8
neopentyl
alcohol
0.385
5.0 h
0.965
2.0 h
2.5
cyclohexanol
cyclopentanol
allyl alcohol
cinnamyl
18.9
19.1
28.2
145
6.1 min
6.1 min
4.1 min
48 s
33.3
33.3
58.0
267
3.5 min
3.5 min
2.0 min
2.6 s
1.8
1.7
2.1
18
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alcohol
5
6
7
piperonyl
alcohol
365
19.0 s
980
7.1 s
2.7
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Table 3. Rates of Oxidation of Various Substituted Benzyl Alcohols by 1 and 5 in CH3CN
Solvent.
O
O
HN
CF3
HN
CH3
Vs.
N
O
N
O
BF4
BF4
O
OH
5
1
R
R
half-time of
oxidation by
Bobbitt’s salt
1
half-time of
oxidation
by CF3 salt
5
relative
rate of
oxidation
(5/1)
rate of oxidation
by Bobbitt’s salt
1 (k x104 s-1)
rate of oxidation by
CF3 salt 5
R
(k x 104 s-1)
p-OCH3
p-(t-Bu)
p-CH3
H
422
165
16.4 s
42.0 s
731
332
9.5 s
20.9 s
1.7
2.0
1.9
1.2
1.4
3.2
2.2
2.8
2.2
1.7
1.7
2.1
208
33.3 s
389
17.8 s
75.1
32.5
11.7
2.42
28.3
10.1
9.32
6.53
3.65
1.5 min
3.6 min
9.9 min
47.7 min
4.1 min
1.1 min
12.4 min
17.7 min
31.7 min
93.4
46.6
37.1
5.31
80.0
22.4
15.4
10.8
7.78
1.2 min
2.5 min
3.1 min
21.8 min
1.4 min
30.9 s
p-F
p-Br
p-CF3
m-OCH3
m-CH3
m-F
7.5 min
10.7 min
14.8 min
m-Cl
m-CF3
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The influence of substituents on the rate of oxidation of a representative sample of para-
and meta-substituted benzyl alcohols is summarized in Table 3. Cursory examination of the data
presented in Table 3 demonstrates that the rate of oxidation of a benzylic alcohol by either 1 or 5
is noticeably affected by the nature of the substituent. Notably, plots of the rate of oxidation of
substituted benzyl alcohols by 1 and by 5 vs. the appropriate Hammett parameters,16 derived
from the pKas of substituted benzoic acids, give strongly linear correlations (r = 0.99) as
illustrated in Figure 6. The slopes of these plots are essentially identical; ρ = –2.8 for oxidations
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by 1 and –2.7 for oxidations by 5. A negative ρ-value of this magnitude indicates the
development of a substantial positive charge during the rate-limiting step of the oxidation. This
result provides convincing experimental support for the mechanism proposed to account for the
oxidation of alcohols by an oxoammonium cation in neutral or slightly acidic media:15 namely,
the oxidation proceeds via removal of a hydride from the carbinol carbon of the alcohol resulting
in the development of a positive charge on that carbon.
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Figure 6. Hammett plot of the rates of oxidation (Table 3) of substituted benzyl alcohols by 1
and 5; kBA = rate constant for oxidation of benzyl alcohol and kX = rate constant for oxidation of
substituted benzyl alcohol.
To ensure that the relative rate differences between oxidations conducted using 1 and 5 in
CH3CN solution were not simply a result of solubility differences between the oxoammonium
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salts (Table 1), a number of oxidations were conducted in water in which 1 is more soluble than
5. The study was perforce limited to three fairly water-soluble secondary alcohols as over
oxidation to the carboxylic acid occurs when primary alcohols are oxidized under aqueous
conditions.10e In the event, the same results were obtained (Table 4): oxoammonium salt 5
oxidizes alcohols at a rate about twice that of Bobbitt’s salt (1). The rates of oxidation of the
secondary alcohols are somewhat slower in aqueous solution than when conducted in
acetonitrile.
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Table 4. Rates of Oxidation of Various Aliphatic Alcohols by 1 and 5 in Water.
O
O
HN
CF3
HN
CH3
Vs.
N
O
N
O
BF4
BF4
5
1
R
O
R
OH
rate of
oxidation by
half-time of
oxidation by
rate of
relative
rate of
oxidation
(5/1)
half-time of
oxidation by
CF3 salt 5
oxidation by
CF3 salt 5
alcohol
Bobbitt’s salt 1 Bobbitt’s salt
(k x 104 s-1)
(k x 104 s-1)
1
2-propanol
cyclohexanol
2-butanol
2.44
8.80
3.21
47.3 min
13.1 min
36.0 min
5.09
1.40
5.37
22.7 min
8.3 min
2.1
1.6
1.7
21.5 min
CONCLUSIONS
In summary, a seemingly minor electronic modification of Bobbitt’s salt (1), involving
replacement of the 4-acetamido group with a 4-trifluoromethylacetamido group, affords an
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oxoammonium salt (5) that oxidizes alcohols significantly more rapidly than does 1: alcohols are
oxidized 1.7 to 3.2 times more quickly by 5 than by 1. The enhanced oxidizing power of 5 vis-à-
vis 1 is attributable to the presence of the electron-withdrawing trifluoromethylacetamido group
at the 4-position of the salt. The rate of oxidation of substituted benzyl alcohols by either 1 or 5
exhibits a very strong linear correlation with the Hammett parameters giving a ρ of –2.8. This
observation provides experimental corroboration of the hydride-transfer mechanism proposed to
account for the oxidation of alcohols by an oxoammonium cation under neutral or slightly acidic
conditions.15 In broader perspective, the results presented above suggest that development of new
nitroxides, and their corresponding oxoammonium salts, might profitably benefit from
consideration of the activating effect of remote electron-withdrawing substituents.
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EXPERIMENTAL SECTION
General Procedures. All reactions were carried out in clean, dried glassware. NMR
spectra (1H, 13C, 19F) were recorded in either deuterated acetonitrile containing 1 v/v%
tetramethylsilane, a solution of 70% trifluoroacetic acid / 30% deuterated chloroform, or a
solution of 70% acetic acid / 30% deuterated chloroform, as noted below, on a 400 MHz NMR
spectrometer. Chemical shifts are reported in parts per million (ppm) relative to TMS at δ = 0.00.
1
1H and 13C NMR spectra were referenced at δ = 7.26 and 77.23, respectively, for the residual H
13
resonance of the solvent and the center line of the C adsorption for deuterated chloroform or at
δ = 0.00 for the 1H or 13C resonance of TMS. 19F NMR spectra were recorded in the appropriate
non-fluorinated solvent, as noted below, with fluorobenzene used as an internal standard (δ = –
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113.15), and chemical shifts are reported relative to CCl3F at δ = 0.00. HRMS molecular mass
determinations were performed on a TOF mass spectrometer; ionization methods are noted
below for individual compounds. Acetonitrile solvent was dried over calcium hydride and
distilled directly prior to use in the rate studies. Bobbitt’s salt (1) was prepared as previously
described3 and was dried overnight in a drying pistol at 56 °C (acetone, 30 mm) prior to use.
Melting points are not corrected.
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2,2,6,6-Tetramethyl-4-(2,2,2-trifluoroacetamido)piperidinium trifluoroacetate. To a
1.5-L Erlenmeyer flask equipped with a large stirbar was added 78.14 g (500 mmol) of 4-amino-
2,2,6,6-tetramethylpiperidine and 400 mL of dry diethyl ether. The solution was cooled to 0 °C
by placing the flask in an ice / water bath, stirring was initiated, and a solution of 126.0 g (600
mmol) of trifluoroacetic anhydride in 100 mL of dry diethyl ether was added dropwise over the
course of 1 h. Upon addition of the anhydride a white precipitate forms. The mixture was stirred
for an additional 1 h and then filtered through a Büchner funnel. The resulting off-white solid
was rinsed with a 200-mL portion of diethyl ether and then dried in a vacuum oven at 120 ºC to
afford 180.7 g (99% yield) of the title compound as a white powdered solid: mp 256 °C (dec),
[lit.17 mp 220 °C (dec)]; 1H NMR (400 MHz, CD3CN) δ 1.43 (s, 6H), 1.50 (s, 6H), 1.75 (t, J =
13.0 Hz, 2H), 1.96 (dd, J = 3.8 Hz, J = 13.0 Hz, 2H), 4.23–4.33 (m, 1H), 7.74 (d, J = 5.3 Hz,
1H); 13C NMR (100 MHz, CD3CN) δ 25.0, 30.5, 40.7, 42.8, 58.1, 117.0 (q, JC-F = 287.6 Hz),
118.1 (q, JC-F = 294.6 Hz), 157.4 (q, JC-F = 36.7 Hz), 161.2 (q, JC-F = 33.5 Hz); 19F NMR (376
MHz, CD3CN) δ –74.2 (3F), –74.7 (3F); HRMS (DART-TOF) m/z calcd for C11H20F3N2O [M –
CF3COO]+ 253.1528, found 253.1522.
2,2,6,6-Tetramethyl-4-(2,2,2-trifluoroacetamido)piperidine-1-oxyl (4). To a 2-L
beaker equipped with a large magnetic stir bar was added 176.0 g (0.48 mol) of 2,2,6,6-
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tetramethyl-4-(2,2,2-trifluoroacetamido)piperidinium trifluoroacetate and 600 mL of deionized
water. The slurry was stirred for 10 min and then 71.0 g (0.67 mol) of solid sodium carbonate
was added in portions over a period of 30 min.18 Following complete addition of the Na2CO3 the
pH of the mixture was ~ 9. At this point, 8.59 g (260 mmol) of sodium tungstate dihydrate and
12.07 g (290 mmol) of EDTA were added to the mixture followed by the slow addition over
period of 3 h of 360 mL (3.2 mol) of a 30 wt% aqueous solution of hydrogen peroxide. The
reaction mixture was stirred vigorously for 12 h at room temperature, an additional 120 mL (1.1
mol) of the hydrogen peroxide solution was then added dropwise over a period of 2 h, and the
mixture was then stirred at room temperature for an additional 24 h. The orange reaction mixture
was cooled to –10 °C in an ice/salt bath and filtered through a large Büchner funnel to give 107.5
g (83%) of the title compound as an orange solid. An analytical sample, prepared by
recrystallization from EtOAc, afforded a glistening, orange-red, crystalline solid that displayed
the following properties: mp 149.6–151.8 °C, (lit.17 mp 149–152 °C); HRMS (DART-TOF) m/z
calcd for C11H18F3N2O2 [M]+ 267.1320, found 267.1335.
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2,2,6,6-Tetramethyl-4-(2,2,2-trifluoroacetamido)-1-oxopiperidinium
tetrafluoroborate (5). To a 250-mL beaker equipped with a magnetic stir bar was charged with
20.0 g (74.9 mmol) of 2,2,6,6-tetramethyl-4-(2,2,2-trifluoroacetamido)piperidine-1-oxyl (4) and
50 mL of deionized water. The slurry was stirred at room temperature and 15.0 mL (7.50 g, 85.4
mmol) of a 50 wt% aqueous solution of tetrafluoroboric acid was added dropwise over the
course of 15 min. The mixture was stirred for ~ 0.5 h until the solution was a yellow color, at
which point, 33.8 mL (2.79 g, 37.4 mmol, NaOCl) of an 8.25 wt% solution of commercial bleach
was added dropwise to the vigorously stirred mixture over a period of 0.5 h. Following the
addition of bleach, 8.22 g (74.9 mmol) of sodium tetrafluoroborate was added and the reaction
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mixture was stirred for an additional 1 h. The reaction mixture was then cooled to –10 °C in an
ice / salt bath and filtered through a large Büchner funnel to afford a yellow solid. The solid was
dried in a drying pistol overnight [acetone (56 ºC, 30 mm)], to yield 19.30 g (73%) of the title
compound in greater than 90% purity; this material, containing traces sodium carbonate and
sodium bicarbonate, is sufficiently pure for most purposes. For the kinetic studies, the salt was
recrystallized as follows. The yellow solid was dissolved in 40 mL of water at 90 °C, the hot
solution was quickly filtered, cooled to room temperature and then to –10 °C in an ice / salt bath.
The cold slurry was filtered through a Büchner funnel and the resulting yellow solid was rinsed
with 75 mL of dry diethyl ether and dried overnight in a drying pistol [acetone (56 ºC, 30 mm)]
to yield 11.04 g (55%) of the title compound as a yellow crystalline solid that was powdered with
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a mortar and pestle and stored under argon in a plastic bottle: mp 160 °C (dec); H NMR (400
MHz, 70% CF3COOH / 30% CDCl3) δ 1.64 (s, 6H), 2.03 (s, 6H), 2.84 (apparent dd, J = 4.0 Hz,
13
J = 14.1 Hz, 2H), 2.97 (t, J = 13.1 Hz, 2H), 5.37–5.47 (m, 1H), 8.12 (d, J = 7.2 Hz, 1H); C
NMR (100 MHz, 70% CF3COOH / 30% CDCl3) δ 28.4, 32.4, 41.2, 43.3, 104.3, 115.5 (q, JC-F
=
285.4 Hz), 160.3 (q, J
= 39.6 Hz); 19F NMR (376 MHz, 70% CH3COOH / 30% CDCl3) δ –
C-F
75.5 (3F), –148.1 (4F); HRMS (DART-TOF) m/z calcd for C11H18F3N2O2 [M – BF4]+ 267.1320,
found 267.1336.
1-Hydroxy-2,2,6,6-tetramethyl-4-(2,2,2-trifluoroacetamido)piperidinium
tetrafluoroborate (6). To a 50-mL round bottomed flask was added 618 mg (1.70 mmol) of
2,2,6,6-tetramethyl-4-(2,2,2-trifluoroacetamido)-1-oxopiperidinium tetrafluoroborate (5), 364 mg
(6.3 mmol) of allyl alcohol and 30 mL of dichloromethane. The reaction mixture was stirred for
2 h at room temperature until the solution became colorless and gave a negative starch-KI test for
the presence of 5. Solvent and excess alcohol were removed under reduced pressure to afford
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537 mg (87%) of the title compound as a white solid: mp 170 °C (dec); H NMR (400 MHz,
CD3CN) δ 1.46 (s, 6H), 1.47 (s, 6H), 1.98 (t, J = 13.2 Hz, 2H), 2.21 (apparent dd, J = 2.4 Hz, J =
13.4 Hz, 2H), 4.25–4.35 (m, 1H), 7.66 (br s, 1H), 8.98 (br s, 1H); 13C NMR (100 MHz, CD3CN)
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δ 20.1, 28.0, 41.3, 41.7, 70.7, 116.9 (q, JC-F = 287.8 Hz), 157.5 (q, J
= 37.4 Hz); 19F NMR
C-F
(376 MHz, CD3CN) δ –74.8 (3F), –149.0 (4F); HRMS (DART-TOF) m/z calcd for
C11H20F3N2O2 [M – BF4]+ 269.1477, found 269.1463.
Rate Studies. Standard solutions of each oxoammonium salt were prepared in dry
CH3CN (0.00500 M for 1 and 0.00502 M for 5) and their UV-Vis spectra were recorded on a
Cary 60 spectrophotometer from Agilent Technologies (version 2.00) scanning from 550–350
nm at a rate of 4600 or 9600 nm/min. The baseline was corrected prior to obtaining readings.
The molar absorptivity of each salt was calculated from the average of five individual
determinations of the absorbance of these standard solutions at 10 nm intervals from 370–530
nm. This allowed for the calculation of an expected absorbance value for the mixture of two
salts. During a kinetic run, a least-squares minimization of the difference between the expected
absorbance and the experimentally measured absorbance at 17 individual wavelengths between
370 and 530 nm using an iterative procedure described by Harris14 and implemented using the
“Solver” function in Excel provided the concentration of each oxoammonium salt as a function
of time.
A typical experiment was conducted as follows. A dry 100-mL volumetric flask was
charged with 147.2 mg (0.4905 mmol) of Bobbitt’s salt (1), 168.3 mg (0.4753 mmol) of the 4-
trifluoromethylacetamido derivative (5), and dry CH3CN. The flask was sonicated for 2-3 min to
ensure the oxoammonium salts had completely dissolved and additional CH3CN was added to
give 100.0 mL of solution. The UV-Vis spectrum was taken of this stock solution to measure an
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initial absorbance reading. An accurately determined quantity of the alcohol substrate (~0.50
mmol) was added to a clean, dry 3.5 mL x 10 mm x 10 mm quartz cuvette, the cuvette was
placed in the UV-Vis spectrophotometer and an accurately measured 2.50 mL of the stock
solution was quickly added to the cuvette. The UV-Vis spectrum was recorded for a given period
of time (5.76 s – 300 s) depending on the substrate and rate of oxidation. The first UV-Vis
spectrum recorded was treated as the initial reading (t = 0 s). Any inconsistencies between the
initial absorbance readings of the oxidation and those of the stock solution were accounted for by
subtracting the two, and then subtracting this difference from all absorbance readings measured
to afford a corrected absorbance reading at each time (t). This corrected reading was used in the
least squares analysis described by Harris.14 UV-Vis spectra as well as the plots of ln ([1]) vs.
time (s) and ln ([5]) vs. time for all substrates investigated can be found in the Supporting
Information; the rate of oxidation was determined from the slope of the linear least squares fit of
these data.
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Solubility Studies (Table 1). The oxoammonium salts were recrystallized from water and
dried for 12 h in an Abderhalden (56 °C, 1 mm, CaCl2) prior to use. The solvent of interest was
measured into a volumetric flask and the oxoammonium salt was then added portion-wise to the
flask, while noting the amount added each time. Periodically, the flask was sonicated for a few s
to ensure complete dissolution at room temperature (23 °C). This process was repeated until no
more oxoammonium salt appeared to dissolve. The flask was left standing for 2–3 h, at which
point, if solid remained the maximum solubility was determined to be reached; if not, the process
was repeated. Solvent was then decanted into a tared round-bottomed flask, the solvent was
removed under reduced pressure, and the mass of oxoammonium salt was recorded. This mass
was divided by the volume of solvent to afford the solubility.
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ASSOCIATED CONTENT
Supporting Information
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UV-Vis spectra of all oxidations; plots of ln ([1]) vs. time (s) and ln ([5]) vs. time (s) for all
oxidations; NMR spectra of all products. This material is available free of charge via the Internet
at http://pubs.acs.org
AUTHOR INFORMATION
Corresponding Authors
* Email: kyle.lambert@uconn.edu
* Email: bartelson@setonhill.edu
* Email: william.bailey@uconn.edu
ACKNOWLEDGEMENTS
The authors wish to dedicate this work to Professor James M. Bobbitt with thanks for helpful
discussions and for being an excellent mentor. The work conducted at the University of
Connecticut was supported by grants from Procter and Gamble Pharmaceuticals, Mason, OH; the
University of Connecticut’s Office of Undergraduate Research (ZDS); and the National Science
Foundation (REU site grant 1359081; SMK). The authors would like to thank Drs. Jie He and
Jing Zhao for use of their UV-Vis spectrophotometers, and Dr. Christopher Kelly for helpful
discussions.
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REFERENCES
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(1)
(a) Kocherginsky, N.; Swartz, H. M. Nitroxide Spin Labels: Reactions in Biology and
Chemistry, CRC Press; New York, 1995. (b) Volodarsky, L. B.; Raznikov, V. A.;
Ovcharenko, V. J.: Synthetic Chemistry of Stable Nitroxides; CRC Press: Boca Raton,
FL, 2000. (c) Hawker, C, J.; Bosman, A. W.; Harth, E. Chem. Rev. 2001, 101, 3661–
3688. (d) Gigmes, D.; Marque, S. R. A. Nitroxide-Mediated Polymerization and its
Applications. In Encyclopedia of Radicals in Chemistry, Biology, and Materials, Vol. 4;
Chatgilialoglu, C.; Studer, A., Eds.; John Wiley and Sons: Hoboken, NJ, 2012; Vol. 4. (e)
Brückner, C.: Nitroxide-Catalyzed Alcohol Oxidations in Organic Synthesis. In Stable
Radicals: Fundamentals and Applied Aspects of Odd-Electron Compounds; Hicks, R. G.,
Ed.; John Wiley & Sons: West Sussex, 2010; pp 433-460. (f) Tebben, L.; Studer, A.
Angew.Chem. Int. Ed. 2011, 50, 5034–5068.
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(
2)
For reviews of oxoammonium salt oxidations see: (a) Leadbeater, N. E.; Bobbitt, J. M.
Aldrichimica Acta 2014, 47, 65–74. (b) Bobbitt, J. M.; Brückner, C.; Merbouh, N. Org.
React. (NY) 2010, 74, 103–424. (c) Merbouh, N.; Bobbitt, J. M.; Brückner, C. Org. Prep.
Proced. Int. 2004, 36, 1–31 (d) Bobbitt, J. M.; Flores, C. L. Heterocycles 1988, 27, 509–
533.
(3)
(4)
Bobbitt, J. M. J. Org. Chem. 1998, 63, 9367–9374.
For a reviews see: (a) Keana, J. F. W. Chem. Rev. 1978, 78, 37–67. (b) Naik, N.; Braslau,
R. Tetrahedron, 1998, 54, 667–696.
(5)
(a) Shibuya, M.; Tomizawa, M.; Suzuki, I.; Iwabuchi, Y. J. J. Am. Chem. Soc. 2006, 128,
8412–8413. (b) Iwabuchi, Y. J. Synth. Org. Chem. Jpn. 2008, 66, 1076–1084. (c)
Shibuya, M.; Tomizawa, M.; Sasano, Y.; Iwabuchi, Y. J. Org. Chem. 2009, 74, 4619–
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