Primary kinetic hydrogen isotope effects in deprotonations of a nitroalkane by intramolecular phenolate groups

Article abstract of DOI:10.1002/poc.1631

Rate constants and kinetic isotope effects have been determined for the formation of nitronate anions from the ethers 1-(2-methoxyphenyl)-2- nitropropane, 7(X=H, L=H and D) and 1-(2-methoxy-5-nitrophenyl)-2-nitropropane, 7(X=NO₂, L=H and D), and from the corresponding phenols, 1-(2-hydroxyphenyl)-2-nitropropane, 3(X=H, L=H and D), and 1-(2-hydroxy-5- nitrophenyl)-2-nitropropane, 3(X=NO₂, L=H and D), in aqueous basic medium. For the ethers 7, rates of deprotonation by hydroxide are comparable with those found for deprotonations of 2-nitropropane, with k^H/k ^D (25 °C)=7.7 and 7.8, respectively. In both the cases, the isotope effects are conventionally temperature dependent. For the corresponding phenols 3, conditions have been established under which the deprotonations of the nitroalkane are dominated by intramolecular deprotonation by the kinetically first-formed phenolate anion, with an estimated effective molarity EM ～ 250. For 3 (X=H, L=H or D), k^H/k^D (25 °C)=7.8, with E_a^D - E_a^H =6.9 kJ mol^-1 and A^H/A^D=0.5. For 3(X=NO₂, L=H or D), rates of intramolecular deprotonation are reduced 30-fold, and an elevated kinetic isotope effect is found (k^H/k^D (25 °C)=10.7). Activation parameters (E_a^D - E_a^H =17.8 kJ mol^-1 and A^H/A^D=0.008) are compatible with an enhanced tunnelling contribution to reactivity in the H-isotopomer. Copyright

Full text of DOI:10.1002/poc.1631

Research Article
Received: 18 June 2009,
Revised: 1 September 2009,
Accepted: 3 September 2009,
Published online in Wiley InterScience: 13 November 2009
(www.interscience.wiley.com) DOI 10.1002/poc.1631
Primary kinetic hydrogen isotope effects
in deprotonations of a nitroalkane by
intramolecular phenolate groups
Nicholas Backstroma, Neil A. Burtona and C. Ian F. Watta
*
Rate constants and kinetic isotope effects have been determined for the formation of nitronate anions from the ethers
1-(2-methoxyphenyl)-2-nitropropane, 7(X ¼ H, L ¼ H and D) and 1-(2-methoxy-5-nitrophenyl)-2-nitropropane,
7(X ¼ NO2, L ¼ H and D), and from the corresponding phenols, 1-(2-hydroxyphenyl)-2-nitropropane, 3(X ¼ H, L ¼ H
and D), and 1-(2-hydroxy-5-nitrophenyl)-2-nitropropane, 3(X ¼ NO2, L ¼ H and D), in aqueous basic medium. For the
ethers 7, rates of deprotonation by hydroxide are comparable with those found for deprotonations of 2-nitropropane,
with kH/kD (25 -C) ¼ 7.7 and 7.8, respectively. In both the cases, the isotope effects are conventionally temperature
dependent. For the corresponding phenols 3, conditions have been established under which the deprotonations of
the nitroalkane are dominated by intramolecular deprotonation by the kinetically first-formed phenolate anion, with
an estimated effective molarity EM ꢀ 250. For 3 (X ¼ H, L ¼ H or D), kH/kD (25 -C) ¼ 7.8, with EDa S EaH ¼ 6.9 kJ molS1 and
AH/AD ¼ 0.5. For 3(X ¼ NO2, L ¼ H or D), rates of intramolecular deprotonation are reduced 30-fold, and an elevated
kinetic isotope effect is found (kH/kD (25 -C) ¼ 10.7). Activation parameters (EaD S EHa ¼ 17.8 kJ molS1 and AH/
AD ¼ 0.008) are compatible with an enhanced tunnelling contribution to reactivity in the H-isotopomer. Copyright
ß 2009 John Wiley & Sons, Ltd.
Keywords: effective molarity; intramolecular deprotonation; nitroalkanes; primary kinetic hydrogen isotope effect;
temperature dependences
computational and experimental effort has been devoted to
discover how enzyme structures promote tunnelling.[3] Certainly,
INTRODUCTION
In this paper, we describe our continuing study on hydrogen
kinetic isotope effects (kies) in intramolecular deprotonations of
carbon acids. A fuller introduction may be found in our initial
papers on the subject.[1,2] In the present paper, we note that our
experiments have been prompted, firstly, by the very unusual kies
reported for some enzyme-catalysed deprotonations[3] and secondly,
by the dearth of comparable measurements in intramolecular
transfers in small molecules,[4] which might model the active-site
chemistry in those enzymes. (For a rationale of the use of
small molecule chemistry in the modelling of active sites, see
Reference [5,6].)
Particular enzymic examples are the rate-limiting deprotona-
tions of a Schiff base by an active site carboxylate in the oxidative
deamination of primary amines by dehydrogenases exploiting
tryptophan tryptophylquinone cofactors. The case of reaction of
methylamine catalysed by MADH,[7] 1, is shown in Fig. 1. In this
case kH/kD ¼ 16.8 at 25 8C, and this ratio is only weakly tempe-
rature dependent with an Arrhenius treatment giving EDa ꢁ EaH ¼
0.5 (ꢂ0.5) kJ molꢁ1 and AH/AD ¼ 13.5. Such behaviour is not
compatible with a kinetic isotope effect arising only from changes
in the isotopically induced ground state zero-point energy
differences,[8] and has been rationalised in terms of hydron
transfer mechanisms involving quantum mechanical tunnelling
of the transferring nuclei through the potential energy barrier.
Complete absence of temperature dependence, indeed, represents
an extreme case where tunnelling dominates the transfer processes
for both the proton and deuteron. Controversially, it has been
suggested that enzymes may have evolved to exploit tunnelling
mechanisms in maximising reaction efficiency,[9,10] and much
such large temperature independent isotope effects are rare in
‘small-molecule’ chemistry, but then measurements of kies for
intramolecular deprotonations of carbon acids are even rarer, so
that experiments with model compounds currently neither
support nor refute any postulate as to the special nature of the
enzyme-catalysed proton transfers.
In our first papers, we reported experimental and compu-
tational studies of the deprotonation of the nitroalkane 2(R ¼ Ph)
by an intramolecular carboxylate, arguing that the extensive
charge reorganisation involved in forming the nitronate anion
had some parallels with the extensive charge shifts accompany-
ing the proton transfer in the MADH-catalysed reaction. This small
molecule chemistry may also model that occurring in an
oxidative conversion of nitroalkanes to corresponding aldehydes
or ketones catalysed by the flavoenzyme, nitroalkane oxidase
(NAO), shown to involve deprotonation of an enzyme-bound
nitroalkane by the carboxylate group of an aspartate residue,[11]
with a remarkable 109-fold rate enhancement over the acetate
induced reaction. A kie, kH/kD (25 8C) ¼ 7.9, has been reported for
the NAO-catalysed reaction of 1,1-dideuterionitroethane,[12,13] but
with conventional Arrhenius parameters (EDa ꢁ EaH ¼ 5.6 kJ ꢃ molꢁ1
,
and AH/AD ¼ 0.89). In the event, both the kie for the intramolecular
*
Correspondence to: C. I. F. Watt, The School of Chemistry, University of
Manchester, Manchester M13 9PL UK.
E-mail: Ian.Watt@manchester.ac.uk
a
N. Backstrom, N. A. Burton, C. I. F. Watt
The School of Chemistry, University of Manchester, Manchester, UK
J. Phys. Org. Chem. 2010, 23 711–722
Copyright ß 2009 John Wiley & Sons, Ltd.

N. BACKSTROM, N. A. BURTON AND C. I. F. WATT
Figure 1. Deprotonations of carbon acids by carboxylate anions in the
active site of MADH and in small molecule model
transfer in 2(R ¼ Ph), kH/kD (25 8C) ¼ 5.8, and its Arrhenius
parameters (EDa ꢁ EHa ¼ 5.5 kJ molꢁ1, and AH/AD ¼ 0.63) were also
resolutely conventional, giving no indications of a major
tunnelling contribution to this simple model reaction.
Catalysis by the intramolecular carboxylate in 2 was not
efficient, with an effective molarity (EM) of only 13.7. Although
low EMs are not unusual in deprotonations of carbon acids, we
have been seeking molecules in which structure places and
retains the catalysing basic group in a more favourable position
with respect to the nitroalkane. Additionally, theory at any level
suggests that an energetic balance in the transfer favours
tunnelling, and this is not the case in the reactions of 2, with the
net transfer of proton from nitroalkane to carboxylate being
endothermic by ca 17.8 kJ molꢁ1 (based on the pKas of the
carboxylic acid and nitroalkane being 4.5 and 7.5 respectively).
We have, therefore, also sought molecules, which might permit a
systematic adjustment of the energetics of the intramolecular
reaction, as well as offering a more effective positioning of the
catalysing group.
Figure 2. The phenolic nitroalkanes and the anticipated sequence of
proton transfers in aqueous basic medium
Introduction of a nitro substituent on the benzene ring to
produce 3(X ¼ 4-NO2) is expected to enhance the acidity of the
phenolic hydroxyl group by a factor of ca 103 (pKa of
2-methyl-4-nitrophenol is 7.1[20]). The effect of the substitution
on the acidity of the nitroalkane residue is expected to be much
smaller and a likely upper limit is a factor of 5.5 (estimated using
sm ¼ 0.710 for the nitro group and r ¼ 1 for the equilibrium
deprotonation). Therefore, the pKas of the oxygen and carbon
acids should be much better balanced (DpKa < 0.5), and
comparison of the behaviours of 3(X ¼ H) and 3(X ¼ NO2) should
permit a simple examination of the effects of the net reaction
energetics on the efficiency of the intramolecular catalysis and
the kies. As far as possible, the behaviour of these phenolic
compounds is to be also compared with those of the corres-
ponding methyl ethers 7, so that the contributions of the
intramolecular base can be assessed.
RESULTS AND DISCUSSION
Experimental design
The compounds we have selected are 2-nitroalkanes carrying
substituted phenolic groups, 3, shown in Fig. 2. In the parent
compound 3(X ¼ H), the anticipated pKa s of phenol and
nitroalkane sites are 10.2 (the pKa of o-cresol) and 7.9,
respectively,[14,15] so that intramolecular proton transfer, conver-
sion of 4 to 5, remains energetically unbalanced, but now with an
anticipated exotherm of ca 14 kJ molꢁ1. In principle at least, this
raises difficulties with the observability of the process. However,
phenols are ‘normal’ acids and deprotonations of the phenolic
hydroxyl group by hydroxide in water, to generate the intra-
molecular phenolate, are expected to occur at rates controlled by
Preparative considerations and preliminary
reactivity studies
1-(o-Methoxyphenyl)-2-nitropropane,[21] 7(X ¼ H) is a distillable
liquid, available in two steps from o-anisaldehyde by Henry
condensation with nitroethane[22] to yield 1-(o-methoxyphenyl)-
2-nitropropene, then by selective reduction of the electron
deficient double bond with sodium borohydride (Scheme 1). Its
1H-NMR spectrum shows a 6-line multiplet (1H) at d4.95, associated
with the hydrogen attached a-to the nitro group, with coupling to
the methyl group and to the adjacent diastereotopic pair whose
signals appear as doublets of doublets at d3.33 and d3.09,
respectively. Nitration of 7(X ¼ H) proceeded with preference for
substitution para to the ethereal position to yield 7(X ¼ NO2),
showing similar signals for the methine and methylene hydro-
gens of the nitropropane side chain. Treatment of 7(X ¼ H) or
7(X ¼ NO2) with d4-methanol and a catalytic amount of pyridine
diffusion or solvent reorganisation (>1010 ꢁ1 sꢁ1),[16] while
M
deprotonations of nitroalkanes are slow, reflecting the extensive
reorganisations of charge, bonding, and solvation, which must
occur before the full stabilising effect of the delocalising nitro
group comes into play.[17–19] Despite the higher thermodynamic
acidity of the nitroalkane residue, the anticipated sequence of
proton transfers in aqueous medium at pH > 10.5 is, therefore,
that outlined in Fig. 2, with the final product being the di-anion 6.
We expect, therefore, to be able to observe the behaviour of the
phenolate anion and characterise the intramolecular proton
transfer involved in the conversion of phenoxide 4 to nitronates 5
and 6.
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Copyright ß 2009 John Wiley & Sons, Ltd.
J. Phys. Org. Chem. 2010, 23 711–722

PRIMARY KINETIC HYDROGEN ISOTOPE EFFECTS
Again, we associate the behaviour with the diffusion-controlled
formation of the nitrophenolate anion 4(X ¼ NO2) and then
slower formation of the nitronate anion, 6(X ¼ NO2).
For both 6(X ¼ H) and 6(X ¼ NO2), basic solutions were stable
for prolonged periods as long as they were protected from
atmospheric oxygen. Acidification to pH 4 regenerated the
originating nitroalkanes via the nitronic acids 9(X ¼ H or NO2). If,
however, solutions were not deoxygenated, and then protected
from molecular oxygen, these recoveries were not clean, with
ketonic products also being formed in variable amounts
(Scheme 2). A larger scale experiments with 3(X ¼ H) permitted
isolation of the known ketone 10(X ¼ H)[26,27] and indicated that
the new compounds were the products of an oxidative Nef
reaction similar to that encountered first in the reactions of the
4-nitro-4-butanoic acids 2(R ¼ Ph and R ¼ Me) by Lewis et al.,[28,29]
and confirmed in our earlier work.[1] The mechanism of this
process is not yet understood, but evidently it occurs readily in
the presence of a suitably placed intramolecular carboxylic acid
group, as in 2, or a phenolic hydroxyl as in 3(X ¼ H or X ¼ NO2).
Although we considered it unlikely that this degradation could
interfere with the rate measurements described below since they
involved the use of basic solutions, in all our detailed kinetic
studies, solutions were degassed, stored under nitrogen and then
the deprotonation reactions were run under nitrogen.
Scheme 1. Preparation 7 and 3 (X ¼ H and X ¼ NO2)
cleanly exchanged the hydrogen attached to the stereogenic
carbon a- to the nitro group for deuterium with loss of the
associated six-line multiplet and simplification of the signals from
the adjacent methyl group and diastereotopic pair. Extents of
deuteriation were readily estimated from 1H-NMR spectra by
integration of the residual methine signal, using the signal of the
non-exchanging hydrogens at the methyl group as an internal
reference. Demethylations with BBr3 in the case of 7(X ¼ H), or
BF3;SMe2 in the case of 7(X ¼ NO2),[23] yielded the required
phenolic nitroalkanes 3(X ¼ H, L ¼ H or D) and 3(X ¼ NO2, L ¼ H or
D). As anticipated from the established stability of nitroalkanes to
acidic medium,[24,25] there was no measurable loss of isotopic
label from deuteriated samples, provided that conditions remained
acidic throughout the reaction, isolation, and chromatography.
Methyl ether, 7 (X ¼ H) was only moderately soluble in water
but, at 30 8C, solutions could be obtained showing a peak at
273 nm(emax ¼ 1500) and end absorption at 220 nm(e ¼ 7000).
Addition of a few drops of sodium hydroxide solution to make the
solution ca 0.01 M in base induced a growth in absorbance at
Reaction kinetics
Rates of deprotonation in basic aqueous medium were first measured.
In all cases, reactions were followed by monitoring first order
absorbance growths at 235 nm. Table 1 shows the rate constants
for nitronate ion formation from both ethers and phenols.
For the ethers, 7(X ¼ H) and 7(X ¼ NO2), first order rate
constants are proportional to base concentration, with no
significant background reaction (see Eqns (1) and (2) respectively
and Figs 3A and B).
wavelengths below 260 nm with t ꢀ 3 min. For the optimum
½
balance between background absorbance and growth, the
changes were monitored at 235 nm and shown to be of the first
order. When solutions were acidified, the spectra slowly reverted
to those of the original nitroalkane, and we associate the growth
with the formation of the anticipated nitronate anion in 8(X ¼ H).
Methyl ether, 7(X ¼ NO2) was also sparingly soluble in pure water,
but solutions could be obtained showing a peak at 320 nm
(emax ¼ 6000) and a shoulder at 228 nm(e ¼ 4900). Addition of
kobs ¼ 0:356ðꢂ0:018Þ ꢄ ½NaOHꢅ þ 8:5 ðꢂ10:1Þ ꢄ 10ꢁ4 (1)
kobs ¼ 0:877ðꢂ0:006Þ ꢄ ½NaOHꢅ ꢁ 9:0 ðꢂ3:0Þ ꢄ 10ꢁ4 (2)
The second order rate constants might be compared with
those reported for similar deprotonations of 2-nitropropane[30]
base induced growth of the shoulder at 228 with t ꢀ 2 min, and
½
again, spectra reverted slowly to those of the original nitroalkane
on acidification of the solution.
Phenol 3(X ¼ H) was much more soluble in water and UV–Vis
spectra showed a shoulder at 209 nm (e ¼ 7050) and a peak at 273
(emax ¼ 1200). Addition of sodium hydroxide solution to make the
solution ca 0.01 M in base induced immediate bathochromic shift
and increase in extinction producing peaks at 292(emax ¼ 8700)
and 233(emax ¼ 1800) nm, and then a slower first order growth
with t ꢀ 5 seconds, most evident at 230 nm. We associate these
½
changes with the anticipated diffusion-controlled formation of
the phenolate 4(X ¼ H) and then slower formation of the
nitronate anion 6(X ¼ H). Nitrophenol 3(X ¼ NO2) exhibited a
similar pattern, with solutions in water showing a shoulder at
223 nm(e ¼ 9800) and a peak at 320 nm(emax ¼ 10,100). Addition
of sodium hydroxide solution induced immediate enhanced end
absorption, and replacement of the peak at 320 nm by a new
peak at 408 nm(emax ¼ 18,000), followed by slower first-order
Scheme 2. Regeneration of nitroalkanes on acidification, and compet-
ing oxidative degradations
growth with t ꢀ 2 min, to develop a distinct new peak at 223 nm.
½
J. Phys. Org. Chem. 2010, 23 711–722
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N. BACKSTROM, N. A. BURTON AND C. I. F. WATT
Table 1. Rates of nitronate ion formation from 7 and 3(X ¼ H, L ¼ H) and from 7 and 3(X ¼ NO2, L ¼ H) in aqueous basic medium at
30.1 8C
c
102 kobs (sꢁ1
)
[Naþ]/M
7(X ¼ H)a
3(X ¼ H)a
7(X ¼ NO2)a
3(X ¼ NO2)a
3(X ¼ NO2)b
0.02
0.04
0.06
0.08
0.760
1.52
2.31
2.87
23.3
23.7
23.4
23.6
1.66
3.39
5.19
6.91
1.11
1.37
1.80
2.05
0.794
0.797
0.799
a NaOH solution, m ¼ 0.1 using KCl;
b Sodium carbonate buffer at pH 10.2, m ¼ 0.08 with KCl.
c Rate constants are means of at least two separate determinations whose values agreed to within 2%.
(k2 ¼ 0.353 Mꢁ1 sꢁ1at 25 8C) and of 1-phenyl-2-nitropropane[31]
(k2 ¼ 1.74 Mꢁ1 sꢁ1 at 30 8C) to gauge the effect of introduction of
the phenyl group, and then of substituted phenyl groups at the
1-position the 2-nitropropane. A simple phenyl group has a rate
enhancing effect (ca 3-fold), but the combined electronic and
steric effects of the o-methoxy group in 7(X ¼ H) reduce the effect
of the phenyl substitution ca 5-fold. Interestingly, the second
order rate constant for deprotonation of 7(X ¼ NO2) by hydroxide
(0.877 Mꢁ1 sꢁ1) is a factor of 2.5 higher than for 7(X ¼ H), an
acceleration presumably reflecting an enhancement in the
acidity of nitroalkane grouping by the additional nitro group on
the benzene ring. A two-point Hammett plot yields r ¼ 0.51 for
the electronic effect of a remote substituent on the rate of
nitroalkane deprotonation.
Reactions of the phenol 3(X ¼ H) are faster (ca 8-fold in 0.08 M
NaOH) than the corresponding ether 7(X ¼ H), but dependence
on base concentration is much reduced (Eqn (3) and Fig. 3A).
Presumably, the change in the electronic nature of the substi-
tuent from an electron-withdrawing methoxy group to an
electron-donating phenolic oxyanion (sI ¼ 0.24 and ꢁ0.16,
respectively) is sufficient to reduce the reactivity in the
intermolecular deprotonation of the nitroalkane by hydroxide
by a factor of about 10. Importantly, the linear plot now has a
substantial intercept (Eqn (3) and Fig. 3A), consistent with the
contribution from the intramolecular deprotonation by the
oxyanion of the phenolate, with kintra ¼ 0.234 sꢁ1. From Eqn (3),
the intramolecular contribution to the formation of nitronate
from 3(X ¼ H) is >98% of the total reaction, even in 0.1 M sodium
hydroxide solution.
kobs ¼ 0:030ðꢂ0:040Þ ꢄ ½NaOHꢅ þ 2:34ðꢂ0:03Þ ꢄ 10ꢁ1
kobs ¼ 0:165ðꢂ0:018Þ ꢄ ½NaOHꢅ þ 7:7ðꢂ0:7Þ ꢄ 10ꢁ3
(3)
(4)
For the phenol, 3(X ¼ NO2), rates show a distinct dependence
(k2 ¼ 0.165 Mꢁ1sꢁ1) on base concentration (Eqn (4) and Fig. 3B)
and a reduced, but clearly discernable, background reaction
(7.7 ꢄ 10ꢁ3 sꢁ1), which we associate the intramolecular depro-
tonation by the phenolate grouping. Compared with 3(X ¼ H),
the slope is larger by at least 15-fold, and the intercept reduced
˝
30-fold and Bronsted treatment (using the estimated pKa s of the
phenols) yields b ¼ 0.48. Second order rate constants for depro-
tonations of simple secondary alkyl nitroalkanes by other
phenolate anions are not available, but those for deprotonations
of 2-nitropropane by hydroxide and by acetate have been
reported[32] and a Bronsted plot (2-points again) yields b ¼ 0.46
˝
for these oxyanionic bases so that there is a degree of
consistency. Both the enhanced slope and the reduced intercept
in the behaviour of 3(X ¼ NO2) compared with that of 3(X ¼ H) are
consistent with the reduced ability of the much less basic
intramolecular phenolic oxyanion in 3(X ¼ NO2) to compete with
the external hydroxide anion in the deprotonation of the
nitroalkane.
The behaviour of 3(X ¼ NO2) was also examined in an aqueous
carbonate buffer at pH 10.2. In this medium, both nitronate and
nitrophenolate anions are expected to be stable. Rates of
nitronate anion formation were independent of the buffer concen-
tration (Eqn (5) and Fig. 3B) and agreed, within experimental
uncertainty, with the extrapolation to zero base concentration for
the hydroxide-induced reaction. In the buffered medium, the
intramolecular deprotonation accounts for no less than 99.9% of
the total reactivity
kobs ¼ :00125ðꢂ0:00914Þ ꢄ ½Naþꢅ þ 7:92ðꢂ0:06Þ ꢄ 10ꢁ3 (5)
A measure of the effectiveness of the intramolecular catalysis
in phenols may be obtained by comparing the second order rate
constant for deprotonation of ethers 7(X ¼ H) and 7(X ¼ NO2) by
hydroxide with the first order rate constants for the intramole-
cular reaction of 3(X ¼ H) and in 3(X ¼ NO2), but allowance must
be made for the difference in the basicity between the hydroxide
and phenolate oxyanions, on the one hand, and of the
nitroalkanes on the other hand. Using the Bronsted b ¼ 0.46
Figure 3. The effects of variation of base concentration on rates on
nitronate formation. Plot A compares reactions of 7(X ¼ H) (^) and
3(X ¼ H) (*), in deprotonations by NaOH. Plot B compares the reactions
of 7(X ¼ NO2) (^) with NaOH with those of 3(X ¼ NO2) with NaOH (*) and
in carbonate buffer at pH 10.2 (*). Data are taken from Table 1
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Copyright ß 2009 John Wiley & Sons, Ltd.
J. Phys. Org. Chem. 2010, 23 711–722

PRIMARY KINETIC HYDROGEN ISOTOPE EFFECTS
calculated above, permits extraction of second order rate
constants for intermolecular deprotonation of the 2-nitroalkanes
residues in 7(X ¼ H) by phenolate, and in 7(X ¼ NO2) by
p-nitrophenolate as 9.0 ꢄ 10ꢁ4 and 3.3 ꢄ 10ꢁ5 Mꢁ1 sꢁ1, respect-
ively. Comparison with first order rate constants for the
intramolecular deprotonations of 3(X ¼ H) and 3(X ¼ NO2) gives
effective molarities for the phenolate groupings of 261 and
234 M, respectively. These are very crude estimates, since no
allowance has been made for the much smaller substituent
induced shifts in the acidity of the nitroalkane. Even as
order-of-magnitude values, these EMs are quite large for carbon
acid deprotonations,[33–35] which are only exceptionally significantly
larger than 10, and then in sterically compressed systems,[36,37]
where the full solvation of the intramolecular base is not possible.[38]
They are certainly larger than those found for the intramolecular
deprotonations (ca 14 M) of the nitroalkane by carboxylate anion
in the carboxylic acid 2 (R ¼ Ph).[1] In all these cases, the O-to-C
proton transfers occur in a six-centre array, and the improved
catalytic efficiency in the phenols must reflect the locking of one
of the bonds in that ring by its incorporation as a part of a
phenolic array. Additionally, inspection of molecular models and
empirical force field calculation suggests that, in the ground state
conformation of the phenols, the 1,2-carbon–carbon bond of the
nitroalkane lies orthogonally to the plane of the benzene ring,
and that the optimum staggered arrangement of substituents on
Table 2. Solvent isotope effects for deprotonations of
1-(2-methoxyphenyl)-2-nitropropane, 7 (X ¼ H, L ¼ H) and of
1-(2-methoxy-5-nitrophenyl)-2-nitropropane 7 (X ¼ NO2,
L ¼ H)
b,c
k2 (Mꢁ1 sꢁ1
Solvent
T (8C)
)
kH2O/kD2O at 25 8C
7(X ¼ H, L ¼ H)a
H2O
H2O
H2Ob
D2O
D2O
D2Ob
19.9
1.44 ꢄ 10ꢁ1
3.24 ꢄ 10ꢁ1
2.20 ꢄ 10ꢁ1
1.93 ꢄ 10ꢁ1
4.26 ꢄ 10ꢁ1
2.92 ꢄ 10ꢁ1
29.8
25.0
19.9
29.8
25.0
0.75(ꢂ0.04)
7(X ¼ NO2, L ¼ H)a
H2O
H2O
H2Ob
D2O
D2O
D2Ob
19.9
29.8
25.0
19.9
29.8
25.0
3.45 ꢄ 10ꢁ1
7.32 ꢄ 10ꢁ1
5.12 ꢄ 10ꢁ1
4.78 ꢄ 10ꢁ1
1.05 ꢄ 100
0.71(ꢂ0.04)
7.20 ꢄ 10ꢁ1
a Base concentrations were [NaOH] ¼ 0.00890 M and
[NaOD] ¼ 0.01285 M.
˚
that bond places the nitroalkane hydron (L) within ca 3.0 A of the
b Interpolated values.
c Rate constants are means of at least two separate determi-
nations whose values agreed to be within 2%.
phenolic oxygen (Fig. 4). Only minor adjustments of the dihedral
angles about the aryl-C1 and C1—C2 bonds are then required to
bring the anionic oxygen and nitroalkane hydrogen to Van der
˚
Waals contact (2.6 A).
Although effective molarities in lactonisations are much larger,
a similar trend is evident in the comparison of lactonisations of
derivatives of 5-hydroxypentanoic acid[39] and o-hydroxy-
hydrocinammic acid.[40] The similarity in the effective molarities
of the catalysing groups in 3(X ¼ H) and in 3(X ¼ NO2) suggests
that there is little enhancement in the catalytic efficiency
associated exclusively with balancing the pKa of the catalysing
and substrate groups.
Table 3. Solvent isotope effects for deprotonations of
1-(2-hydroxyphenyl)-2-nitropropane, 3 (X ¼ H, L ¼ H) and of
1-(2-hydroxy-5-nitrophenyl)-2-nitropropane 3 (X ¼ NO2, L ¼ H)
Solvent
T (8C)
kobs (sꢁ1
)
kH2O/kD2O at 25 8C
c
Solvent kinetic isotope effects, kH2O/kD2O, for the intermole-
cular and intermolecular deprotonations were determined to
probe for possible involvement of hydrons derived from the
water molecules in the deprotonations. The rate data are collected
in Tables 2 and 3. In the case of the methyl ethers, 7(X ¼ H, L ¼ H)
and 7(X ¼ NO2, X ¼ H), values of kH2O/kD2O(at 25 8C) are 0.75(ꢂ0.04)
and 0.71(ꢂ0.04), respectively. These values are only a little less
inverse than that reported for the deprotonations of
2-nitropropane by hydroxide in water (kH2O/kD2O ¼ 0.68),[30]
and, within experimental uncertainty, equal to the fractionation
factor for isotopic distribution between bulk water and water
hydrogen-bonded to hydroxide or deuteroxide (f ¼ 0.74).[41] The
rate ratios are, therefore, consistent with the isotopically induced
3(X ¼ H, L ¼ H)a
H2O
H2O
H2O
H2O
H2Od
D2O
D2O
D2O
D2Od
20.2
7.35 ꢄ 10ꢁ2
1.33 ꢄ 10ꢁ1
2.19 ꢄ 10ꢁ1
3.82 ꢄ 10ꢁ1
1.32 ꢄ 10ꢁ1
1.09 ꢄ 10ꢁ1
1.93 ꢄ 10ꢁ1
4.35 ꢄ 10ꢁ1
1.45 ꢄ 10ꢁ1
24.8
29.5
34.5
25.0
22.6
27.4
34.6
25.0
0.91(ꢂ0.05)
3(X ¼ NO2, L ¼ H)b
H2O
H2O
H2Od
D2O
D2O
D2Od
19.9
29.8
25.0
19.9
29.8
25.0
2.29 ꢄ 10ꢁ3
7.32 ꢄ 10ꢁ3
4.18 ꢄ 10ꢁ3
2.28 ꢄ 10ꢁ3
7.49 ꢄ 10ꢁ3
4.22 ꢄ 10ꢁ3
0.99(ꢂ0.05)
a [NaOH] ¼ 0.00890 M and [NaOD] ¼ 0.01285 M.
b Carbonate buffer, pH 10.2.
c Rate constants are means of at least two separate determi-
nations, whose values agreed to be within 2%.
d Interpolated values.
Figure 4. The ground state conformation of the phenolic nitroalkanes
J. Phys. Org. Chem. 2010, 23 711–722
Copyright ß 2009 John Wiley & Sons, Ltd.
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N. BACKSTROM, N. A. BURTON AND C. I. F. WATT
differences in the desolvation of the oxyanionic base prior to its
involvement in the abstraction of a proton from the nitroalkane.
Involvement of a solvent-derived hydron in covalent bonding
changes in the rate-limiting step would generate a larger ‘normal’
primary effect working against the inverse fractionation factor
and the values here, therefore, do not support any involvement of
water molecules as relays in the proton transfer.
again precede engagement in abstraction of the nitroalkane
hydrogen.
Primary kinetic isotope effects
Our intention was to measure kinetic isotope effects for the
deprotonations over as wide a temperature range as practical. For
the ethers 7(X ¼ H) and 7(X ¼ NO2), we have already noted that
solubility of these compounds in water is not high. Problems
arising from this were not severe in the preliminary experiments
described above, but it soon became evident that, with the
apparatus available to us, problems of low solubility and slow
rates of solution might cause problems if we continued to work
with pure water, especially at lower temperatures. Our kinetic
isotope measurements for these compounds were, therefore,
made using a mixed solvent, 2:1 w/w dimethoxyethane (DME):
water which, on a molar basis, is 71.24% water. This is a mixture
we used earlier in the deprotonations of nitroalkanes[1] and found
that solubilities and rates of solution are much improved while
the reactivity patterns found in the fully aqueous solutions are
not strongly modified. The rate data for the isotopomers of ethers
7(X ¼ H) and 7(X ¼ NO2) are collected in Tables 4 and 5.
Reactivities are enhanced by the change from water to 2:1
DME water by a factor of 1.5 in the case of 7(X ¼ H), and 2.3 in the
case of 7(X ¼ NO2), presumably reflecting reduced solvation of
the hydroxide of the by the more organic medium. Generally, the
kinetics and temperature dependences are well behaved, and we
find activation parameters similar to those reported by Amin and
Saunders[31] for deprotonations of 1-phenyl-2-nitropropane in
water and in aqueous alcoholic mixtures. In all cases, entropies of
activation are large and negative, as might be expected for a
bimolecular process, and our values for both the H-isotopomers
Interestingly, the solvent isotope effects for intramolecular
deprotonations in 3(X ¼ H, L ¼ H) and in 3(X ¼ NO2, X ¼ H) are
kH2O/kD2O ¼ 0.91 (ꢂ0.05) and kH2O/kD2O ¼ 0.99 (ꢂ0.05), respect-
ively, both significantly less inverse than those for the
intermolecular deprotonations. Initially, we speculated that this
reduction might signal some involvement of a solvent-derived
hydron in the process and are still not able to completely exclude
this possibility. We note, however, that there is a factor of ca 105 in
the basicity between hydroxide and phenoxide, and ca 107
between hydroxide and p-nitrophenoxide, and that inverse
relationships between the basicity and hydrogen-bond energies
have been demonstrated for phenolates, both in solution[42] and
in the gas phase.[43] A reduction in hydrogen-bond strength
should be associated with a reduction in depth and a widening of
the potential well describing the bond, both of which would
reduce isotopically-induced energy differentials. It does not seem
unreasonable, therefore, that the fractionation factor for isotopic
distribution between bulk water and water hydrogen-bonded to
hydroxide or deuteroxide (f ¼ 0.74) might be more inverse than
that for distribution between bulk water and water hydro-
gen-bonded to phenoxides. We suggest that the reduced ability
of the phenoxides as hydrogen-bond acceptors (compared with
hydroxide) is a more likely explanation of the less inverse solvent
isotope effects, which simply signal that desolvation of the less
strongly hydrogen-bonded intramolecular phenoxide anion must
Table 4. Rates of deprotonations of 1-(2-methoxyphenyl)-2-nitropropane, 7 (X ¼ H, L ¼ H or D) in 2:1 (w/w) DME: water mixture
b
T (8C)
kobs (sꢁ1
)
k2 (Mꢁ1 sꢁ1
)
Activation parameters
7(X ¼ H, L ¼ H) in DME: H2Oa
12.4
20.6
28.2
34.0
40.6
46.1
25.0c
1.17 ꢄ 10ꢁ2
2.33 ꢄ 10ꢁ2
4.33 ꢄ 10ꢁ2
6.66 ꢄ 10ꢁ2
1.10 ꢄ 10ꢁ1
1.63 ꢄ 10ꢁ1
1.24 ꢄ 10ꢁ1
2.48 ꢄ 10ꢁ1
4.60 ꢄ 10ꢁ1
7.09 ꢄ 10ꢁ1
1.17 ꢄ 100
DHz ¼ 56.2 (ꢂ0.3) kJ molꢁ1
DSz ¼ ꢁ63.4 (ꢂ0.8) J Kꢁ1 molꢁ1
DGz ¼ 75.6 (ꢂ0.5) kJ molꢁ1
ln A ¼ 22.9 (ꢂ0.2)
1.74 ꢄ 100
Ea ¼ 59.2 (ꢂ0.1) kJ molꢁ1
3.53(ꢂ0.04) ꢄ 10ꢁ1
7(X ¼ H, L ¼ D) in DME: H2Oa
12.6
20.6
25.5
30.4
31.5
37.9
43.3
52.0
25.0c
1.72 ꢄ 10ꢁ3
3.32 ꢄ 10ꢁ3
4.78 ꢄ 10ꢁ3
6.54 ꢄ 10ꢁ3
8.01 ꢄ 10ꢁ3
1.31 ꢄ 10ꢁ2
2.0 ꢄ 10ꢁ2
1.83 ꢄ 10ꢁ2
3.5 ꢄ 10ꢁ2
DHz ¼ 59.8 (ꢂ01.6) kJ molꢁ1
DSz ¼ ꢁ69.0 (ꢂ5.4) J Kꢁ1 molꢁ1
DGz ¼ 80.4 (ꢂ0.3.2) kJ molꢁ1
5.09 ꢄ 10ꢁ2
6.96 ꢄ 10ꢁ2
8.52 ꢄ 10ꢁ2
1.39 ꢄ 10ꢁ1
2.18 ꢄ 10ꢁ1
4.41 ꢄ 10ꢁ1
4.56 (ꢂ0.38) ꢄ 10ꢁ2
ln A ¼ 22.2 (ꢂ0.7)
Ea ¼ 62.4 (ꢂ1.7) kJ molꢁ1
4.14 ꢄ 10ꢁ2
a [NaOH] ¼ 0.0940 M.
b Rate constants are means of at least two separate determinations whose values agreed to within 2%.
c Interpolated values.
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Copyright ß 2009 John Wiley & Sons, Ltd.
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PRIMARY KINETIC HYDROGEN ISOTOPE EFFECTS
Table 5. Rates of deprotonations of 1-(2-methoxy-5-nitrophenyl)-2-nitropropane, 7 (X ¼ NO2, L ¼ H or D) in 2:1 (w/w) DME: water
mixture
T (8C)
bkobs (sꢁ1
)
k2 (Mꢁ1 sꢁ1
)
Activation parameters
7(X ¼ H, L ¼ H) in DME: H2Oa
11.2
20.5
29.7
39.1
2.25 ꢄ 10ꢁ3
4.64 ꢄ 10ꢁ3
9.84 ꢄ 10ꢁ3
2.07 ꢄ 10ꢁ2
4.31 ꢄ 10ꢁ1
8.89 ꢄ 10ꢁ1
1.88 ꢄ 100
3.96 ꢄ 100
DHz ¼ 56.4 (ꢂ1.8) kJ molꢁ1
DSz ¼ ꢁ53.6 (ꢂ4.5) J Kꢁ1 molꢁ1
DGz ¼ 72.3 (ꢂ2.7)kJ molꢁ1
ln A ¼ 24.7 (ꢂ0.6)
Ea ¼ 58.8 (ꢂ1.4) kJ molꢁ1
25.0c
2.59 (ꢂ0.05) ꢄ 100
7(X ¼ H, L ¼ D) in DME: H2Oa
11.2
20.5
29.7
39.1
2.45 ꢄ 10ꢁ4
6.13 ꢄ 10ꢁ4
1.28 ꢄ 10ꢁ3
3.10 ꢄ 10ꢁ3
5.26 ꢄ 10ꢁ2
1.17 ꢄ 10ꢁ1
2.43 ꢄ 10ꢁ1
5.93 ꢄ 10ꢁ1
DHz ¼ 61.0 (ꢂ2.6) kJ molꢁ1
DSz ¼ ꢁ54.9 (ꢂ5.4) J Kꢁ1 molꢁ1
DGz ¼ 77.3 (ꢂ8.6) kJ molꢁ1
ln A ¼ 24.5 (ꢂ1.1)
Ea ¼ 63.5 (ꢂ2.6) kJ molꢁ1
25.0c
3.31 (ꢂ0.23) ꢄ 10ꢁ1
a [NaOH] ¼ 0.00522 M.
b Rate constants are means of at least two separate determinations whose values agreed to within 2%.
c Interpolated values.
(ꢁ63.4 and ꢁ53.6 J Kꢁ1 molꢁ1) and D-isotopomers (ꢁ69.0 and
ꢁ55.0 J Kꢁ1 molꢁ1), and are comparable with the ꢁ57.4 J Kꢁ1 Mꢁ1
reported for deprotonation of 1-phenyl-2-nitropropane in water.
The kies and their temperature dependences are presented in
Table 8, and it is apparent that those for the ethers, kH/kD ¼ 7.1
(ꢂ0.4) and 7.6 (ꢂ0.6) at 25 8C, do not differ significantly, and are
within the range expected for rate ratios associated with loss of
ground state isotopic zero-point energy differences. They are also
similar to kies reported for deprotonation of 2-nitropropane[44]
(7.4) and 1-phenyl-2-nitropropane (7.5) by hydroxide in aqueous
medium.[31] For both the ethers, the dominant contributor to the
kies is different in activation energy, with ratios of pre-
exponential factors being well within the range (1/H2 < AH/
AD < H2)[8] associated with an absence of significant tunnelling
of either isotope.
Table 6. Rates of deprotonations of
1-(2-hydroxyphenyl)-2-nitropropane, 3 (X ¼ H, L ¼ H or D) in
water
Table 7. Rates for deprotonations of
1-(2-hydroxy-5-nitrophenyl)-2-nitropropane, 3 (X ¼ NO2, L ¼ H
b
T (8C)
kobs (sꢁ1
)
Activation parameters
or D)
3(X ¼ H, L ¼ H) in H2O; [NaOH] ¼ 0.100 M
4.28 ꢄ 10ꢁ2
7.35 ꢄ 10ꢁ2
1.33 ꢄ 10ꢁ1
2.19 ꢄ 10ꢁ1
3.82 ꢄ 10ꢁ1
6.15 ꢄ 10ꢁ1
DHz ¼ 83.9 (ꢂ0.9) kJ molꢁ1
DSz ¼ þ19.6 (ꢂ2.9) J Kꢁ1 molꢁ1
DGz ¼ 78.1 (ꢂ1.7) kJ molꢁ1
T (8C)
kobs (sꢁ1
)
Activation parameters
b
15.9
20.2
24.8
29.5
34.5
39.0
3(X ¼ NO2, L ¼ H) in H2Oa
30.0
40.0
45.0
55.0
8.80 ꢄ 10ꢁ3
2.47 ꢄ 10ꢁ2
4.00 ꢄ 10ꢁ2
1.03 ꢄ 10ꢁ1
DHz ¼ 81.3 (ꢂ0.3) kJ molꢁ1
DSz ¼ ꢁ15.3 (ꢂ0.8) J Kꢁ1 molꢁ1
DGz ¼ 85.9 (ꢂ0.5) kJ molꢁ1
ln A ¼ 27.53 (ꢂ0.09)
ln A ¼ 32.8 (ꢂ0.4)
Ea ¼ 86.4 (ꢂ0.9) kJ molꢁ1
25.0c 1.33(ꢂ0.03) ꢄ 10ꢁ1
Ea ¼ 81.3 (ꢂ0.3) kJ molꢁ1
3(X ¼ H, L ¼ D) in H2O; [NaOH] ¼ 0.100 M
25.0c
5.07 (ꢂ0.15) ꢄ 10ꢁ3
14.1
20.8
24.8
28.8
36.3
41.0
47.4
4.28 ꢄ 10ꢁ3
9.91 ꢄ 10ꢁ3
1.65 ꢄ 10ꢁ2
2.82 ꢄ 10ꢁ2
6.82 ꢄ 10ꢁ2
1.15 ꢄ 10ꢁ1
2.35 ꢄ 10ꢁ1
DHz ¼ 90.8 (ꢂ0.4) kJ molꢁ1
DSz ¼ þ25.8 (ꢂ1.3) J Kꢁ1 molꢁ1
DGz ¼ 83.1(ꢂ0.8) kJ molꢁ1
3(X ¼ NO2, L ¼ D) in H2Oa
30.0
40.0
45.0
55.0
9.06 ꢄ 10ꢁ4
3.16 ꢄ 10ꢁ3
5.80 ꢄ 10ꢁ3
1.81 ꢄ 10ꢁ2
DHz ¼ 99.1 (ꢂ0.1) kJ molꢁ1
DSz ¼ 24.2 (ꢂ0.3) J Kꢁ1 molꢁ1
DGz ¼ 91.9 (ꢂ0.2) kJ molꢁ1
ln A ¼ 32.31 (ꢂ0.03)
ln A ¼ 33.6 (ꢂ0.2)
Ea ¼ 93.3 (ꢂ0.4) kJ molꢁ1
Ea ¼ 99.1 (ꢂ0.1) kJ molꢁ1
25.0c 1.72 (ꢂ0.05) ꢄ 10ꢁ2
a [NaOH] ¼ 0.100 M.
25.0c
4.75 (ꢂ0.13) ꢄ 10ꢁ4
a Carbonate buffer ay pH 10.2.
b Rate constants are means of at least two separate determi-
nations whose values agreed to within 2%.
c Interpolated values.
b Rate constants are means of at least two separate determi-
nations whose values agreed to within 2%.
c Extrapolated values.
J. Phys. Org. Chem. 2010, 23 711–722
Copyright ß 2009 John Wiley & Sons, Ltd.
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N. BACKSTROM, N. A. BURTON AND C. I. F. WATT
Table 8. Summary of rates, primary kinetic isotope effects (25 8C) and associated activation parameters
Compounds
kH
kH/kD
Activation parameters
7(X ¼ H, X ¼ H and D)
3.53 ꢄ 10ꢁ1 Mꢁ1 sꢁ1
7.7(ꢂ0.7)
EDa ꢁ EHa ¼ 3.2 (ꢂ1.8) kJ molꢁ1
AH/AD ¼ 0.5 (ꢂ0.3)
7(X ¼ NO2, X ¼ H and D)
3(X ¼ H, X ¼ H and D)
3(X ¼ NO2, X ¼ H and D)
2.59 ꢄ 100 Mꢁ1 sꢁ1
1.33 ꢄ 10ꢁ1 sꢁ1
5.07 ꢄ 10ꢁ3 sꢁ1
7.8(ꢂ0.7)
7.8(ꢂ0.4)
10.7(ꢂ0.6)
EDa ꢁ EHa ¼ 4.3 (ꢂ4.9) kJ molꢁ1
AH/AD ¼ 1.2 (ꢂ1.4)
EDa ꢁ EHa ¼ 6.9 (ꢂ0.8) kJ molꢁ1
AH/AD ¼ 0.5 (ꢂ0.8)
EDa ꢁ EHa ¼ 17.8 (ꢂ0.4) kJ molꢁ1
AH/AD ¼ 0.008 (ꢂ0.004)
Figure 5. Diagramme of possible barrier relationships to account for the enhanced tunnel contribution in the reactions of 3(X ¼ NO2)
No similar problems with solubility were experienced for the
phenols 3(X ¼ H) and 3(X ¼ NO2. In the case of 3(X ¼ H, L ¼ H and
D), rates were determined in aqueous 0.1 M sodium hydroxide
and for 3(X ¼ NO2, L ¼ H and D), in the same aqueous sodium
carbonate buffer (pH ¼ 10.2) used earlier. As noted earlier, these
are conditions under which reactivity is dominated by the
intramolecular deprotonation.
The result is an elevated kie, kH/kD ¼ 10.9(ꢂ0.6), at 25 8C,
associated with an exaggerated difference in activation energies,
EDa ꢁ EHa ¼ 17.8 (ꢂ0.4) kJ molꢁ1, and much reduced ratio of
pre-exponential factors, AH/AD ¼ 0.008 (ꢂ0.001), reflecting, we
suggest, anomalous behaviour of 3(X ¼ NO2, L ¼ H).
Rate constants at range of temperatures, and calculated
activation parameters for 3(X ¼ H, L ¼ H and D) are presented in
Table 6, and compared in Table 8. For this intramolecular reaction,
entropies of activation are now positive and comparable (19.6
and 25.8 J Kꢁ1 molꢁ1, respectively), but the kie, kH/kD ¼ 7.8 (ꢂ0.4)
(Table 8) is not dissimilar to those found for intermolecular
deprotonations by hydroxide and, again, arises from the difference
in activation energies with closely comparable pre-exponential
factors.
CONCLUSIONS
The data presented here have no resemblance to the extreme
tunnelling situations encountered in the interesting enzyme-
catalysed reactions mentioned earlier.[7] However, the possibility
that tunnelling contributes to reactivity in all the compounds
studied must not be excluded. Kreevoy[45] has commented that
shallow tunnelling is likely to be wide spread, and points out that
it is as intrinsic to quantum mechanisms as zero-point energy
differentials. There is probably no completely unambiguous way
to demonstrate (or exclude) the phenomenon for reactions
around 3008K. The pattern we have found is consistent with a
small but measurable additional tunnelling contribution to the
reactivity of 3(X ¼ NO2, L ¼ H) only. Comparison of the behaviours
of the two isotopomeric pairs of phenolic nitroalkanes suggests
that its reactivity is enhanced by a factor of ca 1.5 by this
additional tunnelling contribution. Bell’s[46] truncated correction
The situation for the isotopomers 3(X ¼ NO2, L ¼ H and D) is
less simple (Tables 7 and 8). While Eyring treatment of the tem-
perature dependence of rates for the D-isotopomer 3(X ¼ NO2,
L ¼ D), yields a positive entropy of activation (24.2 J Kꢁ1 molꢁ1
)
closely comparable to those found for 3(X ¼ H, L and D), and a
similar pre-exponential factor, the dependences for the H-isotopomer,
3(X ¼ NO2, L ¼ H) reveal a negative entropy of activation (ꢁ15.3 J
Kꢁ1 molꢁ1), and a significantly reduced pre-exponential factor.
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Copyright ß 2009 John Wiley & Sons, Ltd.
J. Phys. Org. Chem. 2010, 23 711–722

PRIMARY KINETIC HYDROGEN ISOTOPE EFFECTS
for tunnelling through an inverted parabolic barrier then yields
nz ¼ 600i cmꢁ1 for the imaginary frequency.
applicable on the product side. Additionally, because of the
higher acidity of the phenolic hydroxyl group in 3(X ¼ NO2), the
internal hydrogen bonding in D may be more favourable than
with 3(X ¼ H), and although we cannot accurately place D on the
energy axis, we place D(X ¼ H) above D(X ¼ NO2) in our
diagramme.
These are qualitative arguments, but we believe that they have
some weight, and the point we wish to make in Fig. 5 is, that
irrespective of details, simple trigonometry demands a narrowing
of the barrier for the transfer in 3(X ¼ NO2). Computational
studies are in progress and these are expected to yield detail of
the barrier shape required to account for the experimental result
that an enhanced tunnelling contribution occurs in 3(X ¼ NO2,
L ¼ H) but not in 3(X ¼ NO2, L ¼ D).
The behaviour found in the intramolecular deprotonations
might usefully be compared to those in intermolecular
deprotonations of 2-nitropropane by alkyl pyridine bases in
water or aqueous alcoholic mixtures[47–49] which, in contrast to
deprotonations by hydroxide described in this work and
elsewhere, also exhibit elevated kies associated with exaggerated
activation energy differences and small ratios of pre-exponential
factors. These effects are most pronounced[44] when the
pyridines carry alkyl substituents at 2- and 6-positions. The
acidities of the pyridinium ions (6 < pKa < 7) are not badly
mismatched to that of 2-nitropropane. The reactions are at least
106-fold slower (DDGz ꢀ 34 kJ molꢁ1) than those with hydroxide,
(k < 6.7 ꢄ 10ꢁ7 Mꢁ1 sꢁ1), reflecting both the basicity change and
steric hindrance to approach of the basic nitrogen lone pair to the
nitroalkane hydrogen. The effects of steric hindrance remain a
matter of debate,[50] but enhanced light atom tunnelling has
been associated with these high symmetrical barriers, and
pressure dependences also suggest that exclusion of solvent
from the reaction volume contributes to the tunnelling, possibly
by reducing the mass of the transferring particle.[51]
EXPERIMENTAL SECTION
General methods
NMR spectra were recorded on Bru¨ker 300, 400 or 500 AMX
spectrometers. Chemical shifts are parts per million (ppm)
downfield from tetramethylsilane. Signal splittings are quoted in
Hz. Signals are listed as: singlet (s), doublet (d), triplet (t), quartet
(q), and multiplet (m) with width (W) to outside lines. Infrared
spectra were recorded on a Perkin-Elmer Spectrum FTIR and were
run as liquid films or as films evaporated from deuteriated
chloroform. UV–Vis spectra were recorded on a Cary 50 UV-visible
spectrophotometer. Mass spectra were recorded on a Kratos
MS25, Fisons VG Trio 2000 or on a Micromass platform. Modes of
ionisation are electron impact (EI), positive chemical ionisation
(þCI) using ammonia and fast atom bombardment (FAB), positive
electrospray (þES), and negative electrospray (ꢁES). Melting
points were recorded on a Kofler heated stage microscope and
are uncorrected. Thin layer chromatography (TLC) was carried out
on Polygram Sil G/UV254 0.25 mm silica gel plates with solvent
systems as indicated. Analytical gas chromatography was carried
out using: Perkin-Elmer capillary gas chromatography model
8310 with flame ionisation detection (FID). Analytical HPLC was
carried out using: Waters 510 HPLC pump, Bondclone C18
100 ꢄ 8 mm column, Perkin-Elmer LC480 Diode Array System
with detection at 255 nm and 285 nm premixed degassed solvent
systems 75:25 MeOH: H2O or as stated in the text, flow rate 2 ml/min.
In our two intramolecular reactions, change of the remote
substituent from X ¼ H to X ¼ NO2 induces only a 35-fold rate
reduction (for the D-isotopomers) corresponding to an increase
in free energy of activation of 8.8 kJ molꢁ1, an amount which, in
comparison to the behaviour described above, does seem to be
large enough to account for the observed tunnelling in
3(X ¼ NO2). A more detailed consideration of the course of the
intramolecular reactions, however, suggests an explanation
(Fig. 5).
We have already pointed out that the likely ground state
conformation of the reactants (A in Fig. 5) is one in which only
minor and energetically undemanding adjustments in dihedral
angles are necessary to bring the anionic oxygen within reacting
distance of the acidic C–H bond. Additionally, desolvation of the
anionic oxygen must precede proton abstraction and, if the pKa
shifts on moving from water to DMSO (10.1–18.1 for phenol and
7.2–10.8 for p-nitrophenol)[52] are taken as indicative, this step
(A ! B in Fig. 5) may cost as much as 45.6 kJ mol(1 in the case of
3(X ¼ H) and 20.5 kJ mol(1 in the case of 3(X ¼ NO2). Comparison
with the experimentally determined free energies of activation
for barriers for the deuteriated isotopomers (83.1 and 91.9 kJ molꢁ1
)
then suggests that the remaining parts of the free energies of
activation might differ by as much as 33.9 kJ molꢁ1, similar to the
differences in activation energy found in the intermolecular
deprotonations we have cited above. These are the parts of the
reaction associated with the actual bonding changes at the
migrating hydron, taking B via transition state C to D, a carbanion
in which charge remains largely localised[53] on carbon, but which
is hydrogen bonded to the phenolic hydroxyl group. The process
is then completed by charge and solvent reorganisation (not
shown) to yield a fully delocalised nitronate. This raising of barrier
is expected to be effective in promoting tunnelling if it is also
associated with a narrowing of the barrier and we suggest that,
because of the intramolecular nature of the process under
consideration and the mode of operation of the substituent, the
width of the barrier for the process in which bonding changes
occur at the migrating hydron must be the same in 3(X ¼ H and
3(X ¼ NO2). On the reactant side (A or B), the remote nitro group
should have no effect on conformational preferences at the
reacting group of atoms. Similar geometry considerations are
The compounds
1-(2-Methoxyphenyl)-2-nitropropene: 2-Methoxybenzaldehyde (13.5 g,
0.1 mol) and ammonium acetate (2.0 g) were dissolved in
nitroethane (100 ml), in an RB flask. The mixture was heated to
reflux overnight and monitoring by TLC then indicated full
conversion of reactant to a new product. Most of the nitroethane
was then removed by distillation at atmospheric pressure before
cooling and dilution with ethyl acetate. This organic mixture was
then washed with dilute aq. HCl, then with aq. sodium
bicarbonate solution and finally with saturated sodium sulphate
solution before drying over anhydrous magnesium sulphate.
Filtration and evaporation yielded an orange oil (16.2 g, 84%)
whose NMR spectrum indicated that it was the anticipated
alkene. A portion was recrystallised from methanol; mp 51–53 8C
(lit., 53[22]). dH (500 MHz, CDCl3) 8.15 (1H, s), 7.28 (1H, t, J ¼ 10.1),
7.16 (1H, d, J ¼ 10.1), 6.89 (1H, t, J ¼ 10.1), 6.82 (1H, d, J ¼ 10.1),
3.75(3H, s), and 2.25(3H,s); dC (75 MHz, CDCl3) 158.6, 147.9, 132.2,
130.4, 130.2, 121.8, 120.8, 111.2, 55.9, 14.4.
J. Phys. Org. Chem. 2010, 23 711–722
Copyright ß 2009 John Wiley & Sons, Ltd.
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N. BACKSTROM, N. A. BURTON AND C. I. F. WATT
1-(2-Methoxyphenyl)-2-nitropropane, 7(X ¼ H, L ¼ H): Sodium
hydroxide pellets (1.02 g) and sodium borohydride (1.02 g) were
dissolved in absolute ethanol (30 ml). In a separate flask,
1-(2-methoxyphenyl)-2-nitropropene (0.84 g) was dissolved in
absolute ethanol (20 ml) in a stirred cooled flask. Aliquots (ca
100 ml) of the sodium borohydride solution were then added, and
the reaction was monitored a few minutes after each addition by
loss of the broad absorption at 348 nm in the UV–Vis spectrum of
the mixture (5 ml of the mixture diluted in 5 mls of ethanol).
Reaction was complete after addition of 1.8 ml of borohydride
solution and 1 h at 0 8C. Any excess sodium borohydride was
destroyed by addition of acetone (2 ml). After stirring for 10 min,
solid carbon dioxide was added, followed by a few drops of
aqueous HCl to adjust the solution to pH 6 as judged by spotting
on wide-range pH paper. Saturated aqueous NaCl solution (40 ml)
was then added and the mixture was left undisturbed for 1 h
before extraction with ether. The combined ether layers were
dried over anhydrous Na2SO4, evaporated, and the resulting
yellow oil was subjected to bulb-to-bulb distillation (145–150 8C
at 1 mm Hg) to yield 1-(2-Methoxyphenyl)-2-nitropropane, as a
light yellow oil (0.61 g, 72%) which was a single component by
GLC and HPLC analysis. dH (500 MHz, CDCl3) 7.29 (1H, t, J ¼ 8.0),
7.11 (1H, d, J ¼ 7.8), 6.91 (2H, m), 4.95(1H, 6 lines, W ¼ 34.1),
3.87(3H, s), 3.30(1H, dd, J ¼ 13.6 and 7.4), 3.09(1H, dd, J ¼ 13.6 and
6.7 Hz); 1.56(3H, d, J ¼ 6.6); dC (75 MHz, CDCl3) 157.9, 131.4, 129.3,
124.4, 121.0, 110.7, 83.4, 55.5, 37.4, 19.8; nmax/cmꢁ1, 3064, 2936,
1539, 1439, 1238, 1124, 1021, 747; m/z (EI) 195(28%, Mþ., found
195.0898, calcd. for C10H13O3N, 195.0980), 151(69%), 148(100%),
120(82%), 115(57%), 91(100%), 84(44%), 77(39%), 49(81%).
2-Deuterio-1-(2-Methoxyphenyl)-2-nitropropane, 7(X ¼ H, L ¼ D):
1-(2-Methoxyphenyl)-2-nitropropane, 7(X ¼ H, L ¼ H), (1.02 g) was
dissolved in methanol-OD (5 ml) and pyridine (1 ml) added. The
reaction mixture was refluxed under nitrogen atmosphere for
purification was by column chromatography (silica, eluting with
ether: petrol 50:50 (v/v), containing 0.05% of formic acid) to yield
a near colourless oil (0.61 g, 89%).
For 3 (X ¼ H, L ¼ H), dH (500 MHz, CDCl3) 7.16 (1H, t, J ¼ 7.4),
7.09 (1H, d, J ¼ 7.4), 6.89 (1H, t, J ¼ 7.4), 6,76 (1H, d, J ¼ 7.4), 5.2(1H,
bs), 5.03(1H, 6 lines, W ¼ 34.2), 3.31(1H, dd, J ¼ 13.8 and 7.8);
3.11(1H, dd, J ¼ 13.8 and 6.3), 1.59(3H, d, J ¼ 6.6); dC (75 MHz,
CDCl3) 154.2, 131.9, 129.2, 122.6, 121.4, 115.8, 83.4, 36.7, 19.7;
nmax/cmꢁ1 3468, 2932, 1595, 1447, 1504; m/z(CI,NH3) 180(M-1),
146, 164, 131; (EI) 181, 134, 107; (ES ꢁve) 180 (100%, M-H, found
180.0661, calcd. for C9H10O3N, 180.0661)
For 3 (X ¼ H, L ¼ D), dH (500 MHz, CDCl3) 7.16 (1H, t, J ¼ 7.4),
7.09 (1H, d, J ¼ 7.4), 6.89 (1H, t, J ¼ 7.4), 6,76 (1H, d, J ¼ 7.4), 5.8(1H,
bs), 3.31(1H, d, J ¼ 13.8); 3.11(1H, d, J ¼ 13.8), 1.59(3H, s).
Integration of the H-NMR spectrum over the signal region
(d5.03) for the exchangeable site indicated deuterium incorp-
oration in excess of 97%.
1-(2-methoxy-5-nitrophenyl)-2-nitropropane, 7 (X ¼ NO2, L ¼ H
or D): 1-(2-Methoxyphenyl)-2-nitropropane (L ¼ H or D) (0.41 g,
2.1 mmol) was dissolved in glacial acetic acid (2 ml) and the
stirred solution cooled to 5 8C. A mixture of concentrated (70%)
nitric acid and acetic acid (2 ml dissolved in) was then added over
a period of 10 min, and the reaction was followed by tlc, which
indicated the formation of two new compounds. The solution
was allowed to warm to 20 8C and stirred for 4 h when it was
poured into ice-water (20 ml) and extracted with three portions
(5 ml each) of dichloromethane. The combined extracts were
dried over anhydrous magnesium sulfate and evaporated to yield
an oil containing residual acetic acid, which was pumped off
under high vacuum at 35 8C. Analysis of the resulting brown oil by
HPLC indicated the presence of two major components in ca 6:4
ratio and the major component was isolated by chromatography
on silica, eluting with 30:70 v/v ether: petrol containing 0.05% of
formic acid, to yield a light yellow oil (0.18 g, 0.79 mmol, 38%).
For 7 (X ¼ NO2, L ¼ H), dH (300 MHz, CDCl3) 8.10 (1H, dd, J ¼ 9.3
and 2.9), 7.92 (1H, d, J ¼ 2.9), 6.88 (1H, d, J ¼ 9.3), 4.85 (1H, 6 lines,
W ¼ 34.1), 3.89(3H, s), 3.22(1H, dd, J ¼ 13.8 and 8.5), 3.05(1H, dd,
J ¼ 13.8 and 5.8), 1.51(3H, d, J ¼ 6.6); dC (75 MHz, CDCl3) 162.5,
1
35 h, and the exchange was followed by H-NMR spectroscopy.
The solvent and pyridine were then removed under reduced
pressure and the residue was taken up in fresh methanol-OD
(2 ml) and pyridine (0.2 ml) added. The mixture was refluxed for
further 48 h before the solvent and pyridine were removed under
reduced pressure to leave a light yellow oil, purified by distillation
under vacuum. dH (300 MHz, CDCl3) 7.80 (1H, t, J ¼ 8.0), 7.11 (1H,
d, J ¼ 7.8), 6.91 (2H, m), 3.87(3H, s), 3.30(1H, d, J ¼ 13.6); 3.09(1H, d,
J ¼ 13.6); 1.57(3H, s); dC (75 MHz, CDCl3) 157.7, 131.2, 130.1, 124.2,
120.9, 110.6, 83.2, 55.5, 36.8, 19.2; nmax/cmꢁ1, 3064, 2939, 1602,
1575, 1495; m/z (EI) 196(26%, Mþ.), 152(63%), 149(100%),
121(80%), 115(57%), 91(100%). Integration of the H-NMR
spectrum over the signal region for the exchangeable site
(d4.95) indicated deuterium incorporation in excess of 97%.
1-(2-Hydroxyphenyl)-2-nitropropane, 3 (X ¼ H, L ¼ H and D):
1-(2-Methoxyphenyl)-2-nitropropane(X¼ H or D) (0.81 g, 4.1 mmol)
was dissolved in chloroform (5.0 ml), in a round bottom flask fitted
for inert atmosphere. The solution was cooled to ꢁ78 8C and stirred
under nitrogen. Boron tribromide (0.5 ml) was then added to the
solution by syringe, and the temperature of the reaction mixture
was allowed to rise to 20 8C and then to left undisturbed for 1 h.
The temperature was then reduced again to ꢁ80 8C and the
reaction was quenched by pouring the reaction mixture into cold
methanol (10.0 ml). The temperature was then allowed to rise to
room temperature, and the solvent was removed by evaporation
under reduced pressure. The residue was taken up in fresh
methanol (10 ml) and the solvent was removed again under
reduced pressure. This process was repeated, and tlc analysis
then indicated the presence of a single major product. Final
141.2, 126.5, 125.4, 125.2, 110.3, 82.2, 56.4, 36.1, 19.3; nmax/cmꢁ1
,
2942, 1546, 1512, 1336, 1260, 1100, and 1029; m/z(CI,NH3)
258(100%, M þ NH4), 193(71%); (EI) 240(8%, Mþ, found 240.0739,
calcd. for C10H12O5N2, 240.0741), 193 (100%) 166 (44%),
115(19%), 77(31%).
For 7 (X ¼ NO2, L ¼ D), dH (300 MHz, CDCl3) 8.10 (1H, dd, J ¼ 9.3
and 2.9), 7.92 (1H, d, J ¼ 2.9), 6.88 (1H, d, J ¼ 9.3), 3.89(3H, s),
3.21(1H, d, J ¼ 13.8), 3.04(1H, d, J ¼ 13.8), 1.50(3H, s). Integration
over of the H-NMR spectrum over the signal region (d4.85) for the
exchangeable site indicated deuterium incorporation in excess of
97%.
1-(2-Hydroxy-5-nitrophenyl)-2-nitropropane, 3 (X ¼ NO2, L ¼ H
and D): The methyl ethers were dissolved in chloroform (ca
30 mgs in 0.5 ml of chloroform) in a small RB flask with a well
fitting stopper. At room temperature, the reagent (BF3: SMe2,
0.5 ml) was added by a syringe in a single portion and the flask
stoppered and swirled to mix. The stoppered flask was then was
left undisturbed at 20 8C for 24 h with occasional swirling. A deep
red oil was deposited, which could be redissolved with a little
heating. After warming to ensure homogeneity, the liquid was
transferred by pipette to about 50 ml of methanol in an
ice-cooled RB flask. The flask was then placed on the rotavap
and the volume was reduced to about 3 ml of red liquid. This was
taken up in chloroform (20 ml) and partitioned with water. The
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Copyright ß 2009 John Wiley & Sons, Ltd.
J. Phys. Org. Chem. 2010, 23 711–722

PRIMARY KINETIC HYDROGEN ISOTOPE EFFECTS
water layer became deep red and the chloroform layer lost much
of its colour. The chloroform layer was separated, and shaken with
a second wash with water (about 5 ml) before being separated
again, and then dried over anhydrous magnesium sulphate. The
sample was applied to a short silica wad, and filtered by washing
with ether with a little formic acid (2 drops in 10 ml). Much of the
residual dark red colour was retained on the silica. Evaporation of
the eluate yielded an oil, which solidified slowly. Preparative TLC
on silica, eluting with CHCl3; MeOH; HCOOH 97.0; 2.8:0.2, by
volume, then recrystallisation from ether: hexane gave off-white
crystals as needles, mp 107–108 8C (17 mgs).
For 3 (X ¼ NO2, L ¼ H), dH (300 MHz, CDCl3) 7.99 (1H, dd, J ¼ 8.7
and 2.8), 7.96 (1H, d, J ¼ 2.8), 6.83 (1H, d, J ¼ 8.7), 6.79(1H,bs), 4.94
(1H, 6 lines, W ¼ 34.2); 3.25(1H, dd, J ¼ 14.2 and 8.5 Hz); 3.07(1H,
dd, J ¼ 14.2 and 5.7); 1.55(3H, d, J ¼ 6.6); dC (75 MHz, CDCl3) 159.9,
141.3, 127.2, 125.2, 123.4, 115.6, 82.3, 35.9, 19.4; nmax/cmꢁ1 3380,
3094, 2987, 1594, 1514, 1333, 1265; m/z (CI) 244(88%, M þ NHþ4 :),
227(9%, M þ Hþ), 211(100%), 194(91%), 177(32%), 148(41%); m/z
(ES ꢁve) 225 (100%, M-H, found 225.0506, calcd. for C9H9O5N2,
225.0511).
that had, after cleaning, been given a final rinse with a dilute
solution (0.01 M) of formic acid in water prior to drying and use. In
the absence of this precaution, solutions of deuteriated materials
could slowly lose the deuterium content. The carbonate buffer
solutions were prepared by dissolving sodium carbonate
(2.293 g) and sodium hydrogen carbonate (1.477 g) in water
(1000 ml). Solution pH was measured at 25 8C using a EDT pH
meter and FisherBrand glass electrode was calibrated against
standard buffers at pH 4 and pH 10 (Hydrion buffer capsules).
UV–Vis spectroscopy: UV–Vis investigations were carried out on
a Cary-50 BIO spectrometer fitted with a thermostatted block.
Temperatures of cells in the block were measured using a
platinum resistance thermometer dipping into a specially
constructed cell. In a typical experiment, the basic solvent,
2.5 ml, was pipetted into cell under nitrogen flow. The cell was
then stoppered and allowed to equilibrate to the set tempera-
ture. Reaction was initiated by addition of 10 ml of a solution of
the nitroalkane to ethanol (ca 0.3 M) and changes at the
appropriate wavelength (see text) were monitored, collecting
data for at least five half-lives. Final absorbances were less than
1.5 and typically absorbance changes were between 0.2 and 0.4
units. For the reaction kinetics, at least 50 points were collected
over at least four half-lives and data transferred to Microsoft
ExcelTM for treatment. Rate constants for H-isotopomers were
extracted by non-linear least squares fitting of Acalcd ¼ Ainf ꢁ DAeꢁkt
to the observed absorbances using Ainf, DA and k as adjustable
parameters using the Solver add-in. Uncertainties and statistics
associated with fitting were generated using the SolvStat macro
of Billo,[54] and correlation coefficients were >0.999 in all cases
and in most cases > 0.9999. For runs with deuteriated materials,
initial points showed deviation from first order behaviour,
ascribed to reactions of small residual amounts of undeuteriated
material. Omission of points from the first 5% of reaction again
yielded excellent fits and rate constants for the deuteriated
isotopomers. Fitting to a sum of two exponentials (five adjustable
parameters) over the whole range yielded rates for both
isotopomers and rate constants obtained in this way were
consistent with those obtained by fitting to single exponentials
with allowance for residual undeuteriated material, but with large
uncertainties for the minor component. For 1-(2-methoxyphenyl)-
2-nitropropane and 1-(2-hydroxyphenyl)-2-nitropropane, deproto-
nation rates were measured using a using a Hi-Tech SF 61
stopped-flow spectrofluorimeter, mixing equal volumes of
solutions of the nitroalkane(ca 5 ꢄ 10ꢁ4 M) and the appropriate
aqueous base.
For For 3 (X ¼ NO2, L ¼ D), dH (300 MHz, CDCl3) 8.00 (1H, dd,
J ¼ 8.7 and 2.8), 7.96 (1H, d, J ¼ 2.8), 6.82 (1H, d, J ¼ 8.7), 4.05(1H,
bs), 3.25(1H, d, J ¼ 14.2); 3.07(1H, d, J ¼ 14.2); 1.55(3H, s).
Integration over of the H-NMR spectrum over the signal region
(d4.94) for the exchangeable site indicated deuterium incorp-
oration in excess of 97%.
The Nef chemistry: 1-(2-Hydroxyphenyl)-2-nitropropane, 3
(X ¼ H, L ¼ H) (0.052 g) was dissolved in a mixture of methanol
and water (5 ml of 1:1 v/v). Sodium hydroxide solution was added
(1 ml of 1 M) and the solution was left undisturbed for 5 min. Air
was then bubbled through the basic solution and dilute
hydrochloric acid to make pH ¼ 0. The aerated solution was
then stirred for 5 min before extraction into chloroform and
drying over anhydrous magnesium sulfate. Examination by tlc on
silica, eluting with ether: petrol (1:1 v/v) showed the absence of
the recovered nitroalkane and a single major new material, more
polar than the nitroalkane. This with was isolated by preparative
TLC on silica as a colourless oil which crystallised on standing,
identified as 2-hydroxyphenylacetone, 10(X ¼ H), (lit., mp
63–64[26,27]). dH (400 MHz, CDCl3) 7.11 (1H, td, J ¼ 7.4 and 1.4),
7.02 (1H, dd, J ¼ 7.4 and 1.2), 6.83 (1H, t, J ¼ 7.5), 6,80 (1H, dd,
J ¼ 7.3 and 1.2), 6.2(1H, bs), 3.71 (2H, s), 2.28 (3H, s); dC (75 MHz,
CDCl3) 210.4, 155.1, 131.0, 129.1, 121.0, 120.9, 117.5, 46.5, and
29.9; nmax/cmꢁ1 3380, 3094, 2987, 1711; m/z (ES ꢁve) 149 (100%,
M-H), 127 (6%).
Reaction kinetics
Acknowledgements
Solutions: Water was distilled in an all-glass apparatus and
deoxygenated with nitrogen. Dimethoxyethane was stirred over
calcium hydride and then distilled under nitrogen from calcium
hydride. The mixed solvents (DME: water) were prepared by
weight.
We thank the University of Manchester for their support and
EPSRC (Grant number GR/580486/01) for a studentship (to Nick
Backstrom). We thank Raman Sharma and Michelle Thorley for
discussions throughout this work.
Stock solutions of aqueous solutions of NaOH in water were by
dilution of certified 1N concentrates (Fisher Chemical Company).
Solutions of NaOD in D2O were made up by dilution of 40% NaOD
in D2O (Aldrich) in D2O and molarities determined by titration
against standard hydrochloric acid. Stock solutions of NaOH in
mixed solvents were made up by weight using the tabulated
densities of NaOH and NaOD solutions and molarities checked by
titration against standard hydrochloric acid. Stock solutions of
nitroalkanes were made up and stored prior to use in glassware
REFERENCES
[1] N. Backstrom, N. A. Burton, S. Turega, C. I. F. Watt, J. Phys. Org. Chem.
2008, 21(7–8), 603–613.
[2] R. Sharma, M. Thorley, J. P. McNamara, C. I. F. Watt, N. A. Burton, Phys.
Chem. Chem. Phys. 2008, 10, 2475–2487.
[3] For an extensive modern introduction to this area: see Hydrogen
Transfer Reactions, Vols 1–4 (Eds: J. T., Hynes, J. P., Klinman, H.,
J. Phys. Org. Chem. 2010, 23 711–722
Copyright ß 2009 John Wiley & Sons, Ltd.
www.interscience.wiley.com/journal/poc

N. BACKSTROM, N. A. BURTON AND C. I. F. WATT
Limbach, R. L. Schowen), Wiley-VCH, Weinheim 2007. ISBN
978-3-527-30777-7
[27] C. F. Carvalho, V. M. Sargent, J. Chem. Soc., Perkin Trans. 1 1984,
1605–1612.
[4] T. L. Amyes, J. P. Richards, in Hydrogen-Transfer Reactions, Vol. 3 (Ed.: J.
T. Hynes ), Wiley-VCH, 2007, pp. 949–973.
[5] A. J. Kirby, Angew. Chem., Int Ed. Engl. 1996, 35, 707.
[6] A. J. Kirby, Hydrogen-Transfer Reactions, Vol. 3 (Ed.: J. T. Hynes),
Wiley-VCH, Weinheim 2007, pp. 975–1012.
[7] M. J. Sutcliffe, L. Masgrau, A. Roujeinikova, L. O. Johannissen, P. Hothi,
J. Basran, K. E. Raraghan, A. J. Mulholland, D. Leys, N. J. Scrutton, Phil.
Trans. Roy. Soc. B 2006, 361, 1375.
[8] M. E. Schneider, M. J. Stern, J. Am. Chem. Soc. 1972, 94, 1517.
[9] Y. Cha, C. J. Murray, J. P. Klinman, Science 1989, 243, 1325.
[10] Y. Cha, C. J. Murray, J. P. Klinman, Science 1989, 244, 244.
[11] P. F. Fitzpatrick, M. Allen, A. Nagpal, M. P. Valley, Arch. Biochem.
Biophys. 2005, 433, 157–1165.
[28] H. Wilson, E. S. Lewis, J. Am. Chem. Soc. 1972, 94, 2283.
[29] H. Wilson, J. D. Caldwell, E. S. Lewis, J. Org. Chem. 1973, 38, 564.
[30] V. Gold, S. Grist, J. Chem. Soc., Perkin Trans. 2 1972, 89–8995.
[31] M. Amin, W. H. Saunders, J. Phys. Org. Chem. 1993, 6, 393–398.
[32] C. F. Bernasconi, M. Panda, M. W. Stronach, J. Am. Chem. Soc. 1995,
117, 9206.
[33] A. J. Kirby, Adv. Phys. Org. Chem. 1980, 17, 183–1278.
[34] A. J. Kirby, Angew. Chem. Int. Ed. Engl. 1996, 35, 707–724.
[35] A. J. Kirby, Acc. Chem. Res. 1997, 30, 290–296.
[36] E. Hartwell, D. R. W. Hodgson, A. J. Kirby, J. Am. Chem. Soc. 2000, 122,
9326–9327.
[37] N. Assad, J. E. Davies, D. R. W. Hodgson, A. J. Kirby, L. van Vliet, L.
Ottavi, J. Phys. Org. Chem. 2005, 18, 101–109.
[12] M. P. Valley, S. E. Tichy, P. F. Fitzpatrick, J. Am. Chem. Soc. 2005, 127,
2062–2066.
[13] D. T. Major, D. M. York, J. Gao, J. Am. Chem. Soc. 2005, 127,
16374–16375.
[14] H. F. Gilbert, J. Am. Chem. Soc. 1980, 102, 7059.
[15] D. Turnbull, S. J. Maron, J. Am. Chem. Soc. 1943, 65, 212.
[16] M. Eigen, Angew. Chem., Int. Ed. Engl. 1964, 3, 1–72.
[17] W. J. Albery, C. F. Bernasconi, A. J. Kresge, J. Phys. Org. Chem. 1988, 1,
29.
[38] R. Pascal, J. Phys. Org. Chem. 2002, 15, 566–569.
[39] B. L. Capon, S. T. McDowell, W. V. Raftery, J. Chem. Soc., Perkin Trans. 2
1973, 1118.
[40] S. Milstein, L. A. Cohen, J. Am. Chem. Soc. 1972, 94, 9158–9165.
[41] R. Alvarez, R. L. Schowen, Isotopes in Organic Chemistry, Vol. 7, Chap. 1
(Eds: E. Buncel, C. C. Lee), Elsevier, Amsterdam 1989.
[42] I. M. Kolthoff, M. K. Chantooni, J. Am. Chem. Soc. 1969, 91, 4621.
[43] J. B. Cumming, M. A. French, P. Kebarle, J. Am. Chem. Soc. 1977, 99,
6999–7003.
[18] C. F. Bernasconi, Adv. Phys. Org. Chem. 1992, 27, 119; and references
therein.
[44] R. P. Bell, D. M. Goodall, Proc. Roy. Soc. (London) Ser. A 1966, 294(1438),
273–297.
[19] J. P. Guthrie, J. Am. Chem. Soc. 1997, 119, 1151.
[20] V. Boehmer, R. Wamsser, H. Kaemmerer, Monatsh. Chem. 1973, 104,
1315–1324.
[45] Y. Kim, M. M. Kreevoy, J. Am. Chem. Soc. 1992, 114, 7116.
[46] R. P. Bell, The Tunnel Effect in Chemistry, Chapman and Hall, New York,
1980.
[21] K. Bailey, Can. J. Chem. 1971, 49, 3143–3151.
[22] Y. Kawai, Y. Inaba, N. Tokitoh, Tetrahedron Asymmetry 2001, 12,
309–318.
[23] M. T. Konieczny, W. Maciewski, W. Konieczny, Synthesis 2005, 10, 1575.
[24] I. Erden, J. R. Keeffe, F. P. Xu, J. B. Zheng, J. Am. Chem. Soc. 1993,
115(21), 9834–9835.
[25] N. C. Deno, R. W. Gaugler, T. Spencer, J. Am. Chem. Soc. 1968, 31,
1968–1969.
[26] A. Lattanzi, A. Senatore, A. Massa, A. Scettri, J. Org. Chem. 2003, 68,
3691–3694;
[47] E. S. Lewis, J. D. Allen, J. Am. Chem. Soc., 1967, 89, 2322.
[48] E. S. Lewis, L. H. Funderburk, J. Am. Chem. Soc. 1968, 90, 4337.
[49] E. S. Lewis, J. K. Robinson, J. Am. Chem. Soc. 1968, 90, 4337.
[50] W. Galezowski, W. Grzeskowiak, A. Jarcewski, J. Chem. Soc., Perkin
Trans. 2 1998, 1607–1611.
[51] N. S. Isaacs, K. Javaid, J. Chem. Soc., Perkin Trans. 2 1979, 1583–1586.
[52] F. G. Bordwell, Acc. Chem. Res. 1988, 21, 456.
[53] F. G. Bordwell, J. E. Bartmess, J. A. Hautala, J. Org. Chem. 1978, 43,
3107.
[54] E. J. Billo, Excel for Chemists, 2nd edn, Wiley-VCH, New York, 2001.
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Copyright ß 2009 John Wiley & Sons, Ltd.
J. Phys. Org. Chem. 2010, 23 711–722