Reaction kinetics in ionic liquids as studied by pulse radiolysis: Redox reactions in the solvents methyltributylammonium bis(trifluoromethylsulfonyl)imide and N-butylpyridinium tetrafluoroborate

Article abstract of DOI:10.1021/jp013808u

Rate constants for several reduction and oxidation reactions were determined by pulse radiolysis in three ionic liquids and compared with rate constants in other solvents. Radiolysis of the ionic liquids methyltributylammonium bis(trifluoromethylsulfonyl)imide (R₄NNTf₂), N-butylpyridinium tetrafiuoroborate (BuPyBF₄), and N-butyl-4-methylpyridinium hexafluorophosphate (BuPicPF₆) leads to formation of solvated electrons and organic radicals. In R₄NNTf₂ the solvated electrons do not react rapidly with the solvent and were reacted with several solutes, including CCl₄, benzophenone, and quinones. In the pyridinium ionic liquids the solvated electrons react with the pyridinium moiety to produce a pyridinyl radical, which, in turn, can transfer an electron to various acceptors. The rate constant for reduction of duroquinone by the benzophenone ketyl radical in R₄NNTf₂ (k = 2 × 10⁷ L mol^-1 s^-1) is much lower than that measured in water (k = 2 × 10⁹ L mol^-1 s^-1) due to the high viscosity of the ionic liquid. Rate constants for electron transfer from the solvent-derived butylpyridinyl radicals in BuPyBF₄ and BuPicPF₆ to several compounds (k of the order of 10⁸ L mol^-1 s^-1) also are lower than those measured in water and in 2-PrOH but are significantly higher than the diffusion-controlled rate constants estimated from the viscosity, suggesting an electron hopping mechanism through solvent cations. Electron transfer between methyl viologen and quinones takes place 3 or 4 orders of magnitude more slowly in BuPyBF₄ than in water or 2-PrOH and the direction of the electron transfer is solvent dependent. The driving force reverses direction on going from water to 2-PrOH and is intermediate in the ionic liquid. Radiolysis of ionic liquid solutions containing CCl₄ and O₂ leads to formation of CCl₃O₂^. radicals, which oxidize chlorpromazine (ClPz) with rate constants near 1 × 10⁷ L mol^-1 s^-1, i.e., much lower than in aqueous solutions and close to rate constants in alcohols. On the other hand, the experimental rate constants in the ionic liquids and in water are close to the respective diffusion-controlled limits while the values in alcohols are much slower than diffusion-controlled.

Full text of DOI:10.1021/jp013808u

J. Phys. Chem. A 2002, 106, 3139-3147
3139
Reaction Kinetics in Ionic Liquids as Studied by Pulse Radiolysis: Redox Reactions in the
Solvents Methyltributylammonium Bis(trifluoromethylsulfonyl)imide and
N-Butylpyridinium Tetrafluoroborate
D. Behar and P. Neta*
Physical and Chemical Properties DiVision, National Institute of Standards and Technology,
Gaithersburg, Maryland 20899-8381
Carl Schultheisz
Polymers DiVision, National Institute of Standards and Technology, Gaithersburg, Maryland 20899-8544
ReceiVed: October 15, 2001; In Final Form: January 15, 2002
Rate constants for several reduction and oxidation reactions were determined by pulse radiolysis in three
ionic liquids and compared with rate constants in other solvents. Radiolysis of the ionic liquids methyltribu-
tylammonium bis(trifluoromethylsulfonyl)imide (R4NNTf2), N-butylpyridinium tetrafluoroborate (BuPyBF4),
and N-butyl-4-methylpyridinium hexafluorophosphate (BuPicPF6) leads to formation of solvated electrons
and organic radicals. In R4NNTf2 the solvated electrons do not react rapidly with the solvent and were reacted
with several solutes, including CCl4, benzophenone, and quinones. In the pyridinium ionic liquids the solvated
electrons react with the pyridinium moiety to produce a pyridinyl radical, which, in turn, can transfer an
electron to various acceptors. The rate constant for reduction of duroquinone by the benzophenone ketyl
radical in R4NNTf2 (k ) 2 × 107 L mol-1 s-1) is much lower than that measured in water (k ) 2 × 109 L
mol-1 s-1) due to the high viscosity of the ionic liquid. Rate constants for electron transfer from the solvent-
derived butylpyridinyl radicals in BuPyBF4 and BuPicPF6 to several compounds (k of the order of 108 L
mol-1 s-1) also are lower than those measured in water and in 2-PrOH but are significantly higher than the
diffusion-controlled rate constants estimated from the viscosity, suggesting an electron hopping mechanism
through solvent cations. Electron transfer between methyl viologen and quinones takes place 3 or 4 orders of
magnitude more slowly in BuPyBF4 than in water or 2-PrOH and the direction of the electron transfer is
solvent dependent. The driving force reverses direction on going from water to 2-PrOH and is intermediate
•
in the ionic liquid. Radiolysis of ionic liquid solutions containing CCl4 and O2 leads to formation of CCl3O2
radicals, which oxidize chlorpromazine (ClPz) with rate constants near 1 × 107 L mol-1 s-1, i.e., much lower
than in aqueous solutions and close to rate constants in alcohols. On the other hand, the experimental rate
constants in the ionic liquids and in water are close to the respective diffusion-controlled limits while the
values in alcohols are much slower than diffusion-controlled.
Introduction
respect to electron-transfer reactions to other compounds. In
this study we utilize two other ionic liquids, one that does not
react with solvated electrons, thus permitting reduction of solutes
by esolv-, and one that does react with solvated electrons but
produces a radical that can reduce other solutes. The first ionic
liquid is based on a tetraalkylammonium cation, which is known
to be unreactive toward eaq-, and is a salt of the asymmetric
cation (n-C4H9)3N(CH3)+ and the anion (CF3SO2)2N- (abbrevi-
ated as R4NNTf2). The second type is based on the N-
Room-temperature ionic liquids serve as good solvents for
various thermal and electrochemical reactions, are nonvolatile
and nonflammable, and have been proposed as “green solvents”
for various processes.1 Recently we have compared reaction rate
constants in ionic liquids with those in other solvents by studying
•
the reactions of CCl3O2 radicals with chlorpromazine and
Trolox in 1-butyl-3-methylimidazolium hexafluorophosphate
(BMIPF6) and tetrafluoroborate (BMIBF4).2 The rate constants
were found to be close to those measured in nonpolar organic
solvents and much lower than those determined in aqueous
solutions. The low rate constants in BMIPF6 were only partially
due to the high viscosity of the solutions and were interpreted
to indicate a high degree of ion-association in this ionic liquid.
On the other hand, the activation energy was found to be closer
to that in water than in alcohols. The results were suggested to
be due to specific solvation of the reactants by different
components of the ionic liquid and the need for proton transfer
to occur along with the electron transfer in this particular
oxidation reaction. Reduction reactions could not be studied in
BMIPF6 because BMI+ reacts very rapidly with solvated
electrons and the BMI• radical produced is very unreactive with
-
alkylpyridinium cation, which is known to react with eaq
rapidly, but produces a pyridinyl radical that can transfer an
electron to many solutes. Several ionic liquids of the second
type were prepared, with N-butylpyridinium (BuPy) and N-butyl-
-
4-methylpyridinium (BuPic) as the cations and with PF6 and
BF4- as the anions. Rate constants were measured in these ionic
liquids for several oxidation and reduction reactions.
Experimental Section3
Methyltributylammonium bis(trifluoromethylsulfonyl)imide
(R4NNTf2) was prepared by reacting equimolar amounts of
methyltributylammonium chloride (MeBu3N+Cl-) with lithium
bis(trifluoromethylsulfonyl)imide ((CF3SO2)2N-Li+) in aqueous
10.1021/jp013808u CCC: $22.00 © 2002 American Chemical Society
Published on Web 03/08/2002

3140 J. Phys. Chem. A, Vol. 106, No. 13, 2002
Behar et al.
solutions at room temperature. The viscous ionic liquid separated
from the aqueous phase. The product was purified by repeated
extractions of the LiCl and the unreacted materials with water
and then dried at 70 °C under vacuum (yield 88%).
(1.75 ( 0.10) µm K-1, and a correction has been applied to the
data taken above room temperature to account for this change
in the test geometry. The reported viscosity is the mean of 3-10
values determined at different shear rates, and the relative
standard uncertainty of the mean viscosity was less than 5%,
except for BuPyBF4 with added water. For those fluids, the
viscosity decreases rapidly with added water, and the relative
standard uncertainty of the mean viscosity increased from 2%
to 10% with increasing water volume fraction. Measurements
were carried out with neat ionic liquids and after addition of
1% CCl4, 1% TEA, or other major additives as used in the
kinetic measurements. The liquids were equilibrated with air
in the rheometer.
N-Butylpyridinium chloride was prepared by adding 1.1 equiv
of 1-chlorobutane (HPLC grade, purity >99.5%) to 1 equiv of
pyridine (Biotech grade, purity >99.9%, packaged under N2)
and gently refluxing for 4 days under a dry atmosphere. The
solid product was precipitated after cooling, filtered and washed
several times with ethyl acetate in a dry atmosphere (because it
is highly hygroscopic), and then dried at 90 °C under vacuum
(yield 70%). The N-butylpyridinium tetrafluroborate (BuPyBF4)
ionic liquid was prepared by mixing 1 equiv of N-butylpyri-
dinium chloride with 1.5 equiv of NH4BF4 in acetonitrile. The
compounds are not soluble, but after mixing for 5 days most of
the NH4Cl precipitated. This was filtered and the acetonitrile
was removed from the liquid phase by distillation under vacuum.
The ionic liquid product was passed through a column of neutral
alumina to remove opacity due to remaining NH4Cl and then
dried in a vacuum at 100 °C (yield >95%).
N-Butyl-4-methylpyridinium bromide was prepared by adding
1.1 equiv of 1-bromobutane (purity >99%) to 1 equivalent of
4-picoline (freshly distilled under N2) in ethanol solution and
gently refluxing for 24 h. The solid product precipitated after
cooling, was filtered and washed several times with ethyl acetate
in a dry atmosphere, and then dried at 90 °C under vacuum
(yield 50%). The PF6 salts of the pyridinium ions are not liquid
at room temperature. Their melting points are 38-39 °C for
the N-butyl-4-picolinium and about 67 °C for the N-butylpyri-
dinium. They were prepared by mixing equimolar amounts of
the halide salt with KPF6, each dissolved in water. The organic
PF6 salts immediately precipitated and were filtered, washed
several times with ethyl acetate, and dried in a vacuum at 100
°C (yield 86%).
Results and Discussion
In general, radiolysis of a liquid leads initially to production
of solvated electrons and solvent radical cations. The solvated
electrons may react with the solvent and, if this reaction is slow,
they may react with solutes. The radical cations may oxidize
solutes or may undergo deprotonation or fragmentation and the
products then react with solutes differently. Geminate recom-
bination of the solvated electrons and radical cations can produce
excited species, which may undergo fragmentation to form
radicals. Despite this complexity it is often possible to observe
clean reduction or oxidation reactions in many solvents due to
rapid conversion of primary species and selective reactions.
Some of the mechanisms may be predicted on the basis of
reaction rates in other solvents. For the ionic liquids under study,
-
-
it is known that the BF4 and PF6 anions and tetrabutylam-
monium cations do not react with solvated electrons, whereas
the pyridine ring is quite reactive.6 In the next section we discuss
the radiolysis of each ionic liquid more specifically.
Radiolysis of the Neat Ionic Liquids. R4NNTf2. Radiolysis
of neat R4NNTf2 leads to production of solvated electrons,
protons, radical cations, and neutral radicals.
Various compounds were used to probe reduction and
oxidation reactions: benzophenone (BP), benzoquinone (BQ),
duroquinone (DQ, tetramethylbenzoquinone), 9,10-anthraquinone-
2-sulfonate (AQS-), methyl viologen (MV2+, 1,1′-dimethyl-4,4′-
bipyridinium dichloride), p-nitroacetophenone (p-NAP), p-ni-
trobenzoic acid (p-NBA), chlorpromazine (ClPz, 2-chloro-10-
(3-dimethylaminopropyl)phenothiazine), N,N,N′,N′-tetramethyl-
p-phenylenediamine (TMPD), and Trolox (6-hydroxy-2,5,7,8-
tetramethylchroman-2-carboxylic acid, a Vitamin E analogue).
These compounds, the starting materials for the syntheses, and
the various solvents and additives were of the purest grade
available from Aldrich, Baker, or Mallinckrodt. Water was
purified with a Millipore Super-Q system.
(R4N)+(NTf2)- Df esolv-, H+, R4N•2+, NTf2 , R• (1)
•
The solvated electrons do not react with R4N+.6 Their reactivity
with NTf2- is unknown and was, therefore, examined in aqueous
solutions. We found no effect of up to 27 mmol L-1 LiN(SO2-
CF3)2 on the lifetime of eaq- in deoxygenated aqueous alcohol
solutions at pH 10.1,7 from which we estimate an upper limit
for the rate constant of NTf2- with eaq- of k < 1 × 106 L mol-1
s-1. Therefore, the solvated electrons in this ionic liquid are
expected to react with the protons and with other reactive
solutes, as demonstrated in a later section.
Reaction kinetics in ionic liquids were determined by pulse
radiolysis. Experiments were carried out with (0.1 to 1.5) µs
pulses of 6 MeV electrons from a Varian linear accelerator;
other details were as described before.4 The dose per pulse was
determined by thiocyanate dosimetry.5 Unless stated otherwise,
the measurements were performed at room temperature, (22 (
2) °C. Rate constants and molar absorption coefficients are
reported with their estimated standard uncertainties.
The oxidized species formed in reaction 1, and excited species
formed by geminate recombination, may undergo fragmentation
or deprotonation. Concerning the quaternary ammonium ion, it
is known that hydrogen abstraction reactions by H• atoms and
•OH radicals are very slow for (CH3)4N+ but the rate constant
increases with chain length8 and it appears that only the
hydrogens at the R-position are highly deactivated by the N+
•
site. The rate constant for reaction of OH radicals with (n-
Viscosity measurements at different temperatures were carried
out in steady shear between 50 mm diameter parallel plates using
a Rheometric Scientific, Inc. ARES rheometer. Tests were
performed at shear rates between 0.1 and 10 s-1 (referenced to
the outer radius of the plates), and none of the fluids tested
showed any signs of non-Newtonian behavior. Small gaps of
0.7 mm or less were used to minimize the sample size and the
rate of rotation needed to achieve higher shear rates. The change
in the gap with thermal expansion has been measured to be
C4H9)4N+ was reported to be very high, 5 × 109 L mol-1 s-1 8
.
Therefore, R4N•2+ is likely to deprotonate at one of the more
remote C-H bonds of the butyl groups to form alkyl radicals
with a tetraalkylammonium moiety (•C4H8NR3+). Fragmentation
of an excited state may lead to C-C bond rupture as well and
thus to alkyl radicals with and without the ammonium moiety.
Concerning the species formed from the bis(trifluoromethyl-
sulfonyl)imide anion, we examined the reaction of this anion
•
with OH radicals in N2O-saturated aqueous solutions at pH 7

Reaction Kinetics in Ionic Liquids
J. Phys. Chem. A, Vol. 106, No. 13, 2002 3141
and detected a transient species with a weak absorption at 330
nm (ꢀ ) 450 L mol-1 cm-1, assuming G ) 0.58 µmol J-1
)
formed with a rate constant of ≈1.2 × 107 L mol-1 s-1. The
-
•
most likely reaction is oxidation of NTf2 with OH to form
• 9
NTf2 . This product may have a finite lifetime and may react
by accepting an electron or a hydrogen atom.10 In addition,
•
•
fragmentation of excited states may form CF3 or CF3SO2
fragments. The various fragments formed from the anion and
the cation are represented by R• in reaction 1. Most of these
radicals are neither oxidants nor reductants but some of them
can react with O2 to form oxidizing peroxyl radicals.
Pulse radiolysis of neat R4NNTf2 produced species with only
a weak tail of absorption from 300 to 400 nm. This absorption
is due to a mixture of the radicals discussed above. We cannot
identify these radicals from the present results, but we can
examine the outcome of their reactions with solutes. Radiolytic
reduction and oxidation of various reactants dissolved in R4-
NNTf2 are discussed in later sections.
BuPyBF4 and BuPicPF6. Radiolysis of N-butylpyridinium
tetrafluoroborate (BuPyBF4) and N-butyl-4-methylpyridinium
hexafluorophosphate (BuPicPF6) leads to production of solvated
electrons, protons, radical cations, and neutral radicals, e.g.,
(BuPy)+(BF4)- Df esolv-, H+, (BuPy)•2+, BF4 (2)
•
•
(BuPy)•2+ and BF4• (or BuPic•2+ and PF6 ) may act as oxidants
or undergo fragmentation. (BuPy)•2+ is likely to lose a proton
from the butyl group to form a pyridylalkyl radical.
Figure 1. Transient absorption spectra monitored by pulse radiolysis
of (a) deoxygenated aqueous solution containing 0.1 mol L-1 t-BuOH
and 1 mmol L-1 BuPyBF4, recorded 5 µs after the pulse, and (b)
deoxygenated neat BuPyBF4, recorded 10 µs after the pulse.
(C4H9Py)•2+ f •C4H8Py+ + H+
(3)
The most likely product is the benzylic radical, although other
isomers cannot be ruled out. Such radicals are not oxidants but
can react with O2 to form oxidizing peroxyl radicals.
of these compounds by solvated electrons and by organic
radicals derived from the solvent. From the absorption maxima
of the radicals produced (≈ 410 and 420 nm, respectively) we
conclude that the semiquinone radicals were protonated under
these conditions,11 the proton being produced by the radiolysis
(reaction 1).
•C4H8Py+ + O2 f •OOC4H8Py+
(4)
The solvated electrons formed in reaction 2 react very rapidly
with the pyridinium cation to form the neutral butylpyridinyl
radical.
BQ + esolv- + H+ f BQH•
(6)
BuPy+ + eaq- f BuPy•
(5)
Addition of 0.07 mol L-1 triethylamine (TEA), to capture the
protons, led to production of the semiquinone radical anions,
with absorption maxima at 420 and 445 nm for BQ and DQ,
respectively.11
We measured the rate constant for this reaction in deoxygenated
aqueous solutions containing 0.2 mol L-1 MeOH7 at pH 11 and
found k5 ) (4 ( 1) × 1010 L mol-1 s-1, in line with rate
constants for similar compounds.6 The spectrum of the N-
butylpyridinyl radical, BuPy•, was recorded by pulse radiolysis
of a deoxygenated aqueous solution containing 0.1 mol L-1
t-BuOH7 and shows a weak absorption at 340 nm (ꢀ ) 850 L
mol-1 cm-1, assuming G ) 0.29 µmol J-1) (Figure 1a). The
spectrum recorded upon irradiation of neat BuPyBF4 shows a
similar weak peak (Figure 1b), as expected if the main absorbing
species is the same pyridinyl radical. This radical is expected
to transfer an electron to compounds with higher electron
affinity, such as quinones, as demonstrated in the next section.
Radiolytic Reduction in the Ionic Liquids. In this section
we discuss the radiolytic reduction of quinones, benzophenone,
methyl viologen, and nitroaromatic compounds in various ionic
liquids, the electron transfer from the benzophenone ketyl
radicals to duroquinone, and the electron transfer between
quinones and methyl viologen, and compare the rate constants
with values measured in other solvents.
BQ + esolv- f BQ•-
(7)
The yields of both radical anions were much higher in the
presence of TEA because TEA quenches excited states and
oxidizing radicals to produce additional reducing radicals. BQ•-
was produced in two distinct steps; approximately half the
absorbance was formed immediately after the pulse (k > 1 ×
108 L mol-1 s-1) and the other half within 20 µs (k ≈ 1 × 107
L mol-1 s-1). The fast step is due to reaction 7 as well as
reaction 8 of BQ with the reducing radical derived from TEA.
BQ + (C2H5)2NC˙ HCH3 f BQ•- + (C2H5)2N+)CHCH3
(8)
Both reactions 7 and 8 are known to have very high rate
constants in aqueous solutions (k > 3 × 109 L mol-1 s-1).6,12
The slower step may be due to reaction of BQ with various
alkyl radicals formed from the solvent. Such reactions have been
observed in aqueous solutions and found to occur via addition
Reduction Reactions in R4NNTf2. Irradiation of deoxygenated
R4NNTf2 containing BQ or DQ (≈ 6 mmol L-1) led to reduction

3142 J. Phys. Chem. A, Vol. 106, No. 13, 2002
Behar et al.
TABLE 1: Rate Constant for Reduction of DQ by BPH•
of the radical to the ring of BQ and subsequent electron transfer
from the adduct to another BQ.13
(Reaction 11) in Various Solvents at Room Temperature
solvent
k
exp, L mol-1 s-1
k
diff,a L mol-1 s-1
BQ + R• f R-BQ•
(9)
R4NNTf2
2-PrOH
H2O/2-PrOH (3/1)
glycerol
(2.0 ( 0.3) × 107
(1.7 ( 0.2) × 108
(2.0 ( 0.4) × 109
(4 ( 1) × 106
1.5 × 107
∼3 × 109
∼3 × 109 b
∼5 × 106
R-BQ• + BQ f [R-BQ+] + BQ•-
(10)
a The values of the diffusion-controlled rate constants, kdiff, were
estimated from the viscosity by using eq 12 and are considered to be
lower limits. They may be higher by up to 50% depending on the exact
shapes of the reactants (Edwards, J. T. J. Chem. Ed. 1970, 47, 261).
bThis value was estimated by taking the viscosity of water with 25 v
% 2-PrOH as 0.002 Pa s, based on values given in the literature
(Landolt-Bo1rnstein. Numerical Data and Functional Relationships in
Science and Technology; Springer-Verlag: Berlin, 1969; Vol. 2, Part
5a, pp 366-367) for water/EtOH and water/1-PrOH mixtures (the
viscosity of water/alcohol mixtures is higher than that of either solvent
alone).
A slightly different pattern was observed in the reduction of
duroquinone. The DQ•- radical anion was produced in three
distinct steps, a very rapid step due to reaction with solvated
electrons, a slower step (k ≈ 6 × 107 L mol-1 s-1) presumably
due to reaction with the TEA radical, and a much slower step
(k < 2 × 105 L mol-1 s-1) probably due to reactions parallel to
reactions 9 and 10. The increase of absorbance produced in the
last step was very small, and the process was probably
incomplete because of radical-radical reactions. Therefore, its
observed rate constant is considered to be an upper limit. If
this last step indeed involves addition of radicals to the ring (as
in reaction 9), it is not surprising that its rate constant is much
lower for DQ than for BQ because of the steric hindrance of
the four methyl groups on the ring.
Reduction of benzophenone (0.04 to 0.09 mol L-1) in
deoxygenated R4NNTf2 solutions containing TEA produced
initially the radical anion BP•-, which has an absorption peak
in water at 610 nm14,15 and in the ionic liquid at ≈ 700 nm.
This species was converted within 10 µs into BPH•, which has
an absorption peak at 545 nm both in water14,15 and in the ionic
liquid. After this rapid initial production of the radical, a slower
increase in absorbance, with k ≈ 1 × 106 L mol-1 s-1, was
observed, probably due to reduction of BP by the TEA radical.
The finding that BP•- undergoes protonation under these
conditions while BQ•- and DQ•- are not protonated is in line
with the higher pKa value in aqueous solution of BPH• (9.2)14,15
as compared with those of the semiquinones (reactions 4 and
5).11 From the measured absorbance of BPH• at 545 nm we
can estimate the yield of solvated electrons in this ionic liquid
by assuming that the molar absorption coefficient of BPH• in
R4NNTf2 is similar to that in aqueous solutions (5.5 × 103 L
mol-1 cm-1).15 By using thiocyanate dosimetry and correcting
for the difference in electron density between the aqueous
thiocyanate solution and the ionic liquid (646/556), we estimate
that the radiolytic yield of solvated electrons captured by BP
in the absence of TEA is 0.9 × 10-7 mol J-1 and approximately
twice as high as that in the presence of TEA.
Figure 2. Viscosities of the neat ionic liquids at different temperatures
(between 20 °C and 75 °C), shown as Arrhenius plots, for R4NNTf2
(O), BuPyBF4 (b), BuPicPF6 (2), and BMIPF6 (4). Sample numerical
values (in Pa s) are: R4NNTf2: 0.75 at 20 °C, 0.52 at 25 °C.
BuPyBF4: 0.17 at 20 °C, 0.13 at 25 °C. BuPicPF6: 0.14 at 45 °C.
These measured values should be compared to the diffusion-
controlled rate constants, which can be estimated from the
viscosities.
The viscosity of the neat ionic liquid was determined at
various temperatures and the results are presented in the form
of an Arrhenius plot (Figure 2). From the viscosity we can
estimate the diffusion-controlled rate constant (kdiff) by using
the approximation,16
Electron transfer from BPH• to DQ (reaction 11) was
examined in R4NNTf2 solutions containing 70 mmol L-1 TEA,
90 mmol L-1 BP, and various concentrations of DQ between
0.15 and 1.2 mmol L-1
.
BPH• + DQ f BP + H+ + DQ•-
(11)
kdiff ) 8000RT/3η
(12)
The rate constant was determined at room temperature by
following the formation of DQ•- at 445 nm and found to be
(2.0 ( 0.3) × 107 L mol-1 s-1. Similar measurements were
carried out in alcohol and aqueous solutions for comparison
(Table 1). The rate constant determined in 2-PrOH solutions
containing similar concentrations of TEA, BP, and DQ was (1.7
( 0.2) × 108 L mol-1 s-1 and the rate constant in aqueous
solutions, containing 25% 2-PrOH to permit sufficient solubility,
was (2.2 ( 0.3) × 109 L mol-1 s-1. Thus, the measured rate
constant of this electron-transfer reaction in the ionic liquid is
much lower than that in alcohol or aqueous solutions. This effect
must be due in part to the high viscosity of the ionic liquid.
For comparison, we measured the same rate constant in glycerol
where R is the gas constant (8.3144 J K-1 mol-1), T the absolute
temperature, and η the viscosity (in Pa s). The values of kdiff
for the respective conditions (Table 1) were estimated from the
viscosity measured under those conditions (i.e., with solutions
containing all the main components, not just with neat solvents).
It is seen that the rate constant for reaction 11 measured in R4-
NNTf2 is approximately the same as the estimated kdiff. The
measured values in glycerol and in water/2-PrOH are also close
to the corresponding value of kdiff. On the other hand, the
measured rate constant in 2-PrOH is ∼20 times lower than kdiff
.
This comparison suggests that R4NNTf2 behaves in this case
as a polar solvent, which facilitates electron transfer, similar to
water or glycerol and more polar than 2-PrOH.
and found it to be even lower, (4 ( 1) × 106 L mol-1 s-1
.

Reaction Kinetics in Ionic Liquids
J. Phys. Chem. A, Vol. 106, No. 13, 2002 3143
those of water and 2-PrOH were taken from the literature. In
water and 2-PrOH, the electron-transfer rate constants are
somewhat lower than the values of kdiff, as observed before for
many other reactions. In contrast, all the rate constants measured
in the ionic liquids are about an order of magnitude higher than
the calculated kdiff. This unexpected observation may be
explained by an electron-transfer mechanism that involves
hopping of the electron between solvent cations so that the
unpaired electron of BuPy• can reach the organic substrate
without the need for complete diffusion of a specific BuPy•
radical to the substrate. This is particularly facilitated when the
neutral substrate is solvated by the pyridinium cations, which
can channel the electron. The positively charged methyl viologen
-
is expected to be solvated partly by the BF4 anions and thus
the rate of electron transfer, although higher than kdiff, is not as
high as with the neutral solutes (Table 2).
The measured electron-transfer rate constants in the ionic
liquids were found to increase with temperature as shown by
Arrhenius plots (Figure 3) and their activation energies are
summarized in Table 2. For comparison, we calculated kdiff at
the different temperatures and derived the activation energy for
Figure 3. Arrhenius plots for the rate constants for reduction by the
pyridinyl radicals in the ionic liquids. In BuPyBF4: BuPy• + DQ (b),
BuPy• + MV2+ (O), BuPy• + p-NBA (0). In BuPicPF6: BuPic• +
DQ (2).
k
diff: 43.6 kJ mol-1 for BuPicPF6 and 37.6 kJ mol-1 for
BuPyBF4. These values are very close to the activation energies
for the reactions of DQ and p-NBA in both ionic liquids (Table
2). This finding is in line with the hopping mechanism suggested
above if we assume that increased diffusion rate (a) facilitates
encounters between solvent cations by faster repositioning of
the solvent ions, thus enhancing the electron hopping, and (b)
increases the probability of encounters between the radical and
the solute, thus requiring fewer hopping steps. The activation
energy for the reaction of MV2+ is slightly lower due to its
solvation with anions as discussed above.
Reduction Reactions in BuPicPF6 and BuPyBF4. Solutes in
these ionic liquids are not reduced by the radiolytically produced
-
solvated electrons since esolv is rapidly scavenged by the
pyridinium cations (reaction 5). Therefore, the main reducing
species in these ionic liquids is the butylpyridinyl radical. This
radical does not reduce benzophenone (an upper limit of k < 1
× 106 L mol-1 s-1 was estimated in 2-PrOH/water solutions),
and we could not measure the rate of reduction of quinones by
the benzophenone ketyl radical in these ionic liquids as we did
in R4NNTf2. We measured, however, the rate constants for
reduction of various compounds by the butylpyridinyl radical
(e.g., reaction 13).
The radiolytic yield for production of MV•+ was determined
by measuring the absorbance of this radical at 605 nm in
deoxygenated BuPyBF4 solutions containing 0.8 mmol L-1
MV2+ and using aqueous thiocyanate dosimetry.5 By assuming
that the molar absorption coefficient of MV•+ in BuPyBF4 is
BuPy• + DQ f BuPy+ + DQ•-
(13)
similar to that in water (1.3 × 104 L mol-1 cm-1 17
)
and
Rate constants were measured for the reduction of duro-
quinone, methyl viologen, p-nitrobenzoic acid, and p-nitroac-
etophenone by the butylpyridinyl radical in these ionic liquids.
Since the absorption of BuPy• is weak, the rate constants for
its reactions with these reactants were determined by following
the buildup of absorption of the radicals produced from the other
correcting for the differences in electron densities between water
and BuPyBF4 (668/556) we find G(MV•+) ) (1.8 ( 0.2) ×
10-7 mol J-1. The radiolytic yield in BuPyBF4 was also
determined by measuring the yield of reduction of p-NAP and
assuming that the absorption coefficient of its radical anion is
similar to that in water. The molar absorption coefficient of
p-NAP•- in aqueous solutions was reported to be 17600 L mol-1
cm-1 at 350 nm and 2900 L mol-1 cm-1 at 545 nm.14 We
monitored the absorbance at 545 nm to avoid interference by
the absorbance of the solvent radicals. To ascertain that p-NAP•-
does not undergo protonation to form p-NAPH•, which has lower
absorbance, we added TEOA as a base. From the absorbance
at 545 nm observed with deoxygenated BuPyBF4 containing
0.06 mol L-1 TEOA and 0.6 mmol L-1 p-NAP, we estimate
reactants: 445 nm for DQ,11 400 and 605 nm for MV2+ 17
, 330
nm for p-NBA,18 and 370 nm for p-NAP.14 For each reactant,
rate constants were measured at several concentrations and
several temperatures, between 22 and 75 °C. Reactions in
BuPicPF6 were measured only between 43 and 75 °C since the
melting point of this ionic liquid is 38-39 °C. Second-order
rate constants were derived from the concentration dependence
of kobs. The values at the different temperatures give good
Arrhenius plots (Figure 3) and the results are summarized in
Table 2.
the total yield of radicals G ) (4.8 ( 0.5) × 10-7 mol J-1
.
This value is much higher than that determined with MV2+ in
the absence of TEOA and indicates that most of the initial
oxidizing species are converted into reducing radicals by the
TEOA, as discussed above for TEA. Therefore, p-NAP is
reduced by both BuPy• and by the radicals derived from TEOA.
The rate constants for the same reactions were also determined
in aqueous solutions and the rate constants are compared in
Table 2. In all cases, the rate constants are about an order of
magnitude higher in water than in the ionic liquid. In the case
of DQ the rate constant was also measured in 2-PrOH and found
to have an intermediate value between those in water and in
the ionic liquid.
Table 2 also shows the values of the diffusion-controlled rate
constants estimated from the viscosities by using eq 12. The
viscosities of the ionic liquids were measured (Figure 2) and
BuPy• + p-NAP f BuPy+ + p-NAP•-
(HOCH2CH2)2NC˙ HCH2OH + p-NAP f
(14)
(HOCH2CH2)2N+dCHCH2OH + p-NAP•- (15)

3144 J. Phys. Chem. A, Vol. 106, No. 13, 2002
Behar et al.
TABLE 2: Rate Constants for Reduction Reactions in Various Solvents
reaction
solvent
T, °C
k
exp, L mol-1 s-1
Ea,a kJ mol-1
log A
k
diff,b L mol-1 s-1
BuPic• + DQ
BuPicPF6
BuPyBF4
BuPyBF4
BuPyBF4
2-PrOHd
H2Oe
43
24
23
24
22
22
22
22
22
(4.1 ( 0.6) × 108
(4.4 ( 0.7) × 108
(2.3 ( 0.7) × 108
(4.7 ( 0.8) × 108
(1.3 ( 0.3) × 109
(3.3 ( 0.5) × 109
(3.1 ( 0.5) × 109
(4.0 ( 0.5) × 109
(5.4 ( 0.8) × 109
42
15.6
14.8
13.6
15.6
5 × 107
5 × 107
5 × 107
5 × 107
3 × 109
7 × 109
7 × 109
7 × 109
7 × 109
BuPy• + DQ
34.9
29.5
39.8
BuPy• + MV2+
BuPy• + p-NBAc
BuPic• + DQ
BuPic• + DQ
BuPy• + DQ
H2Oe
BuPy• + MV2+
H2Of
BuPy• + p-NBAc
H2Oe
a Activation energies have been determined with an overall estimated standard uncertainties of less than (10%. b See footnote a in Table 1.
c The rate constant in the ionic liquid was determined with p-nitrobenzoic acid but in water with p-nitrobenzoate anion since the acid is insoluble
in water. d 1% TEA was added as a base to prevent protonation of DQ•- into DQH•. e 0.2 mol L-1 t-BuOH was added as an OH radical scavenger.
f 1 mol L-1 t-BuOH was added as an OH radical scavenger.
The observed formation of p-NAP•-, however, appeared to take
TABLE 3: Rate Constants for Electron Transfer between
Quinones and Methyl Viologen
place in one step, with a rate constant of (2.2 ( 0.4) × 108 L
reaction
solvent
T, °C kexp, L mol-1 s-1
mol-1 s-1
.
MV•+ + DQ
MV•+ + DQ
MV•+ + DQ
MV•+ + DQ
MV•+ + DQ
MV•+ + DQ
MV•+ + DQ
MV•+ + DQ
DQ•- + MV2+
BuPyBF4
25 (1.4 ( 0.3) × 105
25 (7.1 ( 1.1) × 105
25 (3.0 ( 0.5) × 106
25 (2.0 ( 0.3) × 107
25 (1.5 ( 0.3) × 108
25 (5.5 ( 0.8) × 108
25 (1.1 ( 0.1) × 109
25 (3.3 ( 0.5) × 109
25 (1.7 ( 0.4) × 109
22 (4.0 ( 0.8) × 106
22 (6.1 ( 0.9) × 109
In an attempt to clarify the role of TEOA in the reduction of
p-NAP, we conducted experiments in aqueous solutions. The
rate constant for reduction of p-NAP by the BuPy• radical in
deoxygenated aqueous solutions containing 0.2 mol L-1 t-
BuOH, 11 mmol L-1 BuPyBF4, 5 mmol L-1 phosphate buffer
at pH 7, and between 0.07 and 0.2 mmol L-1 p-NAP, but no
TEOA, was determined to be k14 ) (6.2 ( 0.9) × 109 L mol-1
s-1. In the absence of t-BuOH but with 0.07 mol L-1 TEOA in
N2O-saturated aqueous solutions (to convert eaq into OH so
that all the primary species are converted into TEOA radicals),
at pH 9.4, the reduction of p-NAP took place with k15 ) (3.2
( 0.6) × 109 L mol-1 s-1, i.e., about half the value of k14.19 In
the ionic liquid, reactions 14 and 15 took place in parallel and
it was not possible to distinguish two steps of reduction.
Electron Transfer between Methyl Viologen and Quinones.
In addition to the electron transfer from the solvent butylpy-
ridinyl radicals to various solutes, we studied secondary electron
transfer from a solute radical to a different solute. First, we
measured the rate constant for reaction 16 by pulse radiolysis
of deoxygenated BuPyBF4 solutions containing MV2+ (0.7 to
4 mmol L-1) and duroquinone (0.05 to 5 mmol L-1).
BuPyBF4, 2 v % H2O
BuPyBF4, 5 v % H2O
BuPyBF4, 9 v % H2O
BuPyBF4, 19 v % H2O
BuPyBF4, 30 v % H2O
BuPyBF4, 50 v % H2O
H2O, 10 v % 2-PrOH
2-PrOH, 5 v % TEOA
AQS•2- + MV2+ BuPyBF4, 2 v % H2Oa
-
•
MV•+ + AQS-
H2O, 2 v % 2-PrOHb
AQS•2- + MV2+ 2-PrOH, 15 v % watera 22 (1.5 ( 0.3) × 109
a A small amount of water was necessary to permit sufficient
dissolution of AQS-. 2-PrOH was used as a scavenger for H• and
b
•OH radicals; the solution was neutral, unbuffered.
MV•+ + DQ h MV2+ + DQ•-
(16)
From the rate of decay of the MV•+ absorption at 400 nm as a
function of DQ concentration we derive k16 ) (1.4 ( 0.2) ×
105 L mol-1 s-1. This rate constant is 2 orders of magnitude
lower than kdiff, unlike the rate constant for reaction 13, which
was an order of magnitude higher than kdiff. For comparison,
we also determined the rate constant for reaction 16 in aqueous
solutions containing 10% 2-PrOH and found k16 ) (3.3 ( 0.5)
× 109 L mol-1 s-1, i.e., k16 is 4 orders of magnitude lower in
the ionic liquid than in water. We then measured k16 in several
BuPyBF4/water mixtures. The results (Table 3) show a large
increase in rate constant upon addition of small amounts of water
and then a slower approach to the aqueous solution value (Figure
4). From comparing with the values of kdiff estimated from the
viscosities (Figure 4), it is seen that in predominantly aqueous
solutions the experimental k16 is about 50% of kdiff but as the
fraction of ionic liquid approaches 1 the value of k16 decreases
below 1% of kdiff. This comparison suggests that additional
factors contribute to the decrease in k16, most likely a change
in the equilibrium constant K16 due to the change in the nature
of the solvent. It is known that the reduction potential of MV2+
is considerably more positive in alcohol solutions than in
water,20 while the reduction potentials of quinones are more
negative in alcohols.21 It is also known that in aqueous solutions
Figure 4. Rate constant for electron transfer from MV•+ to DQ
measured in different BuPyBF4/water mixtures (b) and the correspond-
ing values of kdiff (O) estimated from the measured viscosities by using
eq 12.
reaction 16 has a driving force of about -0.21 V,22 but in
2-PrOH the reaction is driven in the opposite direction.23 To
confirm this trend we followed reaction 16 in 2-PrOH solutions,
and indeed, we found it to proceed to the left, with k-16 ) (1.7
( 0.4) × 109 L mol-1 s-1 (Table 3). We also attempted to
determine the equilibrium constant for reaction 16 in 2-PrOH
but could only estimate K > 104 at the most extreme concentra-
tion ratio that was experimentally practical. Thus the driving
force for reaction 16 changes from -0.21 V in water to at least
0.2 V in 2-PrOH. Therefore, we suggest that the driving force
in BuPyBF4 is intermediate between these values and thus the

Reaction Kinetics in Ionic Liquids
J. Phys. Chem. A, Vol. 106, No. 13, 2002 3145
reaction occurs much more slowly. It is also possible that
reaction 16 in neat BuPyBF4 is near equilibrium and that it is
pulled to the right due to protonation of DQ•-. Irradiation of
the solution produces protons (reaction 2), which are not
neutralized under these conditions and can react with DQ•- to
form DQH•. Since DQH• has a shorter lifetime than DQ•- and
cannot transfer an electron to MV2+, reaction 16 can be driven
to the right. In confirmation of this suggestion we found that
addition of 2 v% pyridine to neutralize part of the protons led
to reversal of the direction of reaction 16. Electron transfer from
DQ•- to MV2+ was observed in the presence of pyridine, with
a rate constant k-16 ≈ 106 L mol-1 s-1, and the decay of MV•+
occurred much more slowly than in the absence of pyridine.
Thus reaction 16 proceeds in BuPyBF4 in the same direction
as in 2-PrOH and in opposite direction to that in water.
This reversal of the reaction driving force was demonstrated
more clearly in similar experiments with MV2+ and AQS-.
total yield of oxidation under air was only 0.9 × 10-7 mol J-1 25
.
The different oxidation yields of these two compounds indicate
that some radicals formed in this solvent can oxidize TMPD
(E ) 0.26 V vs NHE)22 but do not react rapidly with ClPz (E
) 0.8 V).22 Alkylperoxyl radicals are known to exhibit such
behavior and they are produced in the present system when
irradiated under air. The radicals produced in deoxygenated
solutions include Tf2N•, which is probably the strong oxidant
that leads to the rapid oxidation step, and various fragments
that are weaker oxidants. In oxygenated solutions, it is possible
•
that the peroxyl radicals, such as CF3O2• and R3N+CR2O2 , are
formed and these can oxidize ClPz..
Irradiation of various solvents containing CCl4 and O2 is
known to produce CCl3O2• radicals, which are strong oxidants.
•
This is due to the fact that CCl4 can form CCl3 radicals when
it reacts with solvated electrons as well as with various organic
radicals. Recently2 we have studied the rate constant for
•
oxidation of ClPz by CCl3O2 radicals in another ionic liquid,
MV•+ + AQS- h MV2+ + AQS•2-
(17)
BMIPF6.
•
CCl3O2 + ClPz f CCl3O2- + ClPz•+
(18)
In aqueous solutions reaction 17 proceeds to the right with k17
) (6.1 ( 0.9) × 109 L mol-1 s-1 while in 2-PrOH the direction
is reversed23 and k-17 ) (1.5 ( 0.3) × 109 L mol-1 s-1 (Table
3). In BuPyBF4 the reaction proceeds to the left as in 2-PrOH,
with k-17 ) (4.0 ( 0.8) × 106 L mol-1 s-1 (Table 3). Upon
addition of 2 v% pyridine, to minimize the extent of protonation
of AQS•2- to AQSH•-, k-17 increased to (2.7 ( 0.5) × 107 L
mol-1 s-1. These results clearly indicate the reversal of the
driving force upon going from water to BuPyBF4 and emphasize
the similarity between the ionic liquid and the alcohol.
Most of the reactions discussed in this study involve at least
one uncharged species and thus their rates are not affected by
the ionic strength of the medium to any appreciable extent. Only
reaction 17 involves two charged species in both directions. The
reverse direction, involving two doubly charged species with
opposite signs, is decelerated at increased ionic strength.
Although k-17 measured in the ionic liquid is much smaller than
that measured in 2-PrOH, this decrease is due to at least in part
to changes in reduction potentials and it is not possible to assess
whether the ionic liquid exerts the ionic strength effect expected
from ≈ 3 mol L-1 salt concentration.
In the present study we determined the rate constant for the
same reaction in R4NNTf2 by pulse radiolysis of oxygenated
solutions containing 1% CCl4 and various concentrations of
ClPz. The rate of formation of ClPz•+ was monitored at 525
nm as a function of ClPz concentration, and the value of k18
was derived from a linear plot of kobs vs [ClPz]. However, unlike
the earlier case of BMIPF6, where oxidation of ClPz took place
to a significant extent only in the presence of CCl4 and O2, in
R4NNTf2 there was oxidation of ClPz even in the absence of
CCl4, as discussed above. Therefore, we measured the rate
constant for oxidation of ClPz in oxygenated solutions of R4-
NNTf2 both in the presence and absence of CCl4. The absor-
bance observed in the presence of CCl4 was about twice as high
as that observed in the absence of CCl4, but the rate constants
were similar, and it was not possible to determine different rate
constants for different processes. This is not surprising if the
solvent derived radicals include the strong oxidants26 CF3O2
•
and R3N+CH(R)O2 . By assuming that the rate constant
•
measured in the presence of CCl4 (4.3 × 106 L mol-1 s-1) is
•
Radiolytic Oxidation in the Ionic Liquids. In this section
we discuss the oxidation of TMPD and ClPz in the neat ionic
liquids, from which we estimate the total oxidation yields, and
the rate constants for oxidation of ClPz and Trolox by the
the weighted average for the reactions with CCl3O2 and with
the other radicals (measured separately as 3.7 × 106 L mol-1
s-1), we estimate that the rate constant for reaction 18 is 5 ×
106 L mol-1 s-1
.
•
CCl3O2 radicals produced in the presence of CCl4.
We also measured the rate constants as a function of
temperature from 21 to 75 °C and found slightly different
temperature dependences in the presence and absence of CCl4.
The activation energies are summarized in Table 4. The
experimental rate constants (kexp) were then corrected for the
effect of the diffusion-controlled limit (kdiff) to derive the
activation-controlled rate constants (kact).16
Oxidation Reactions in R4NNTf2. The radiolytic yield of
oxidation in R4NNTf2 was determined by using TMPD and
ClPz. These compounds are oxidized to produce long-lived
radical cations, which have strong absorption peaks with known
molar absorption coefficients in various solvents. Pulse radi-
olysis of deoxygenated R4NNTf2 containing 6 mmol L-1 TMPD
produced TMPD•+ with two absorption peaks at 565 and 613
nm. The absorbance at 565 nm was produced in two steps, one
part was formed within 1 µs after the pulse and an additional
absorbance was formed within 100 µs, indicating the existence
of at least two oxidizing species. The radiolytic yield of the
fast oxidation step was 1.1 × 10-7 mol J-1, and the total yield
kact-1 ) kexp-1 - kdiff
(19)
-1
The values of kdiff as a function of temperature were estimated
from the viscosity measured at different temperatures (Figure
2) according to eq 12. Arrhenius plots for the calculated kact
values are shown in Figure 5 and the results from these plots
are also summarized in Table 4. It is seen that the experimental
rate constants are approximately the same near 43 °C and that
the Arrhenius plots diverge from there, giving an activation
energy and a preexponential factor that are lower in the presence
of CCl4 than those in its absence. Although the activation energy
for reaction 18 derived in this complex manner is approximate,
was 1.7 × 10-7 mol J-1 24
.
When a similar solution was
irradiated under air, an additional formation step took place over
2 ms and the total yield increased to 2.1 × 10-7 mol J-1. This
increase is due to oxidation of TMPD by peroxyl radicals.
Similar experiments with ClPz gave lower yields. The initial
step of rapid formation of ClPz•+, monitored at 525 nm, was
very small, corresponding to only 0.3 × 10-7 mol J-1, and the

3146 J. Phys. Chem. A, Vol. 106, No. 13, 2002
Behar et al.
TABLE 4: Rate Constants for Oxidation Reactions in Ionic Liquids
reaction
solvent
T, °C
kexp L mol-1 s-1
Eaa
log A
kdiffb L mol-1 s-1
kact L mol-1 s-1
Eaa
log A
CCl3O2• + ClPzc
R4NNTf2
R4NNTf2
BuPicPF6
BuPyBF4
21
21
45
23
22
22
22
22
23
25
22
(4.3 ( 0.6) × 106
(3.7 ( 0.6) × 106
(1.9 ( 0.3) × 107
(2.4 ( 0.4) × 107
(1.2 ( 0.2) × 107
(2.9 ( 0.4) × 107
(9.6 ( 1.6) × 106
(6 ( 1) × 106
29
34
30.8
25.2
30.3
11.8
12.7
12.3
11.9
12.4
1.1 × 107
9.2 × 106
5.2 × 107
4.8 × 107
2.2 × 107
6.1 × 107
6.1 × 107
6.7 × 107
4.8 × 107
7 × 109
7.2 × 106
6.2 × 106
3.0 × 107
4.8 × 107
2.6 × 107
5.5 × 107
1.1 × 107
6.6 × 106
1.1 × 107
22.6
28.8
25.0
18.2
22.7
10.8
11.9
11.6
10.9
11.4
RO2• + ClPzd
CCl3O2• + ClPz
CCl3O2• + ClPz
CCl3O2• + ClPz
CCl3O2• + ClPz
CCl3O2• + Trolox
CCl3O2• + Trolox
CCl3O2• + Trolox
CCl3O2• + ClPz
CCl3O2• + ClPz
e
BMIPF6
BMIPF6
BMIPF6
BMIBF4
e,f
e,f
e
BuPyBF4
(9 ( 1) × 106
24.2
23.7
8.7i
11.1
13.2
9.4i
22.8
11.0
H2Og
1.2 × 109
2-PrOH
3.1 × 107 h
3 × 109
a Activation energies (in kJ mol-1) have been determined with an overall estimated standard uncertainties of less than (10%. b Estimated from
the viscosities that were measured with identical solutions. c This reaction took place in parallel with reactions of solvent derived peroxyl radicals
with ClPz (see text) and the processes could not be separated. d Reaction involving several solvent radicals but excluding CCl3O• (see text). e Results
from ref 2. f These solutions also contained 10 v% 2-PrOH. g Solution contained 10% 2-PrOH; results from ref 27. h From ref 28. i From ref 27 for
2-PrOH solutions containing 9% water.
s-1. The diffusion-controlled rate constants estimated from the
viscosities measured under the same conditions are about twice
as high, as listed in Table 4, or possibly as much as three times
as high (see footnote a in Table 1). The values of kact, estimated
from kexp and kdiff, are higher than kexp by a factor of 1.5-2 and
remain close to 1 × 107 L mol-1 s-1. These values are clearly
lower than the rate constant measured for this reaction in
water27,28 and close to the values measured in alcohols.28 On
the other hand, the activation energies are closer to those
measured in aqueous solutions and higher than those in alcohols
(Table 4). On this basis, we cannot equate the solvent effect of
the ionic liquids on reaction 18 with that of either alcohols or
water. Furthermore, if we compare the rate constants for reaction
18 in the different solvents not on an absolute basis but as a
fraction of the diffusion-controlled limit, we find that this
fraction is ≈0.5 in the ionic liquids, ≈0.2 in water, and only
≈0.01 in 2-PrOH (Table 4) and other less polar organic
solvents.28 This comparison indicates that the correction of kexp
to derive kact is small for water and negligible for alcohols but
very large for the ionic liquids, where kexp and kdiff are close.
•
Figure 5. Arrhenius plots for the rate constants for reaction of CCl3O2
As a result, the calculated values of kact in the ionic liquids are
radicals with ClPz in ionic liquids. The solid symbols are for the
extremely sensitive to the exact value of kdiff. The above
measured kexp and the open symbols are for the calculated kact. The
comparison shows, however, that encounters between reactants
in water and ionic liquids are much more likely to result in
reaction than encounters in alcohols and nonpolar organic
solvents. Possibly, encounters in ionic liquids take place in more
organized microenvironments, where the reactants remain in the
same domain for longer periods and have a higher probability
solvents are BuPyBF4 (90), ByPicPF6 (bO), R4NNTf2 with CCl4 (24),
and R4NNTf2 without CCl4 (13).
its value is within the range of the activation energies determined
in a more direct manner in the other ionic liquids (see below).
Oxidation Reactions in BuPicPF6 and BuPyBF4. The radi-
olytic oxidation yield in BuPyBF4 was determined by using
TMPD solutions, as described above with R4NNTf2. The yield
was 1.5 × 10-7 mol J-1, slightly lower than the yield in R4-
NNTf2.
•
of reacting. The reactions of CCl3O2 radicals have been
suggested to involve an electron transfer and a proton transfer
occurring in concert,29 so that the insipient CCl3O2 anion is
-
immediately converted to the more stable CCl3O2H. If the proton
is supplied from the ammonium tail of ClPz (or from the
phenolic OH of Trolox), lengthening the time of encounter
between the reactants will also increase the probability for the
proton transfer.
The rate constants for oxidation of ClPz and Trolox by the
CCl3O2 radicals were determined in oxygenated ionic liquid
•
solutions containing 1% CCl4. No significant oxidation was
detected in these ionic liquids in the absence of CCl4. In all
cases examined the rate constants at room temperature (Table
4) are close to 107 L mol-1 s-1. The rate constants were also
measured at different temperatures to derive the activation
energies (Table 4). The experimental data were then corrected
for the effect of kdiff to derive kact. Arrhenius plots of kact are
shown in Figure 5, and the results calculated from these plots
are also summarized in Table 4.
The rate constants for oxidation of Trolox in the ionic liquids
(Table 4) are slightly lower than those for oxidation of ClPz,
as noted before for the same reactions in other solvents.27,28
This is due to the fact that Trolox is a neutral phenol and only
the deprotonated form undergoes rapid oxidation. Therefore,
the solvent effect on reaction with Trolox includes as an
additional parameter the ability of the solvent to remove a proton
from Trolox. This point was discussed in detail previously.28
Oxidation Rate Constants. The experimental rate constants
•
for oxidation of ClPz by CCl3O2 radicals in the various ionic
liquids (Table 4) are all within a factor of 2 of 1 × 107 L mol-1

Reaction Kinetics in Ionic Liquids
J. Phys. Chem. A, Vol. 106, No. 13, 2002 3147
References and Notes
Summary and Conclusions
(1) Welton, T. Chem. ReV. 1999, 99, 2071, and references therein.
Brennecke, J. F.; Maginn, E. J. AIChE J. 2001, 47, 2384.
(2) Behar, D.; Gonzalez, C.; Neta, P. J. Phys. Chem. A 2001, 105,.
(3) The mention of commercial equipment or material does not imply
recognition or endorsement by the National Institute of Standards and
Technology, nor does it imply that the material or equipment identified are
necessarily the best available for the purpose.
(4) Grodkowski, J.; Neta, P. J. Phys. Chem. A 2000, 104, 4475.
(5) Schuler, R. H.; Patterson, L. K.; Janata, E. J. Phys. Chem. 1980,
84, 2088.
Pulse radiolysis of ionic liquids (BMIPF6, BMIBF4, BuPicPF6,
BuPyBF4, and R4NNTf2) permits production of a variety of
radicals and measurement of absolute rate constants for reduc-
tion and oxidation of various molecules. The experimental rate
constants for oxidation of ClPz and Trolox by CCl3O2• radicals
in the ionic liquids (Table 4) are much lower than rate constants
for the same reactions in aqueous solutions and closer to rate
constants measured in alcohols. This comparison suggests that
the ionic liquids do not behave as highly polar solvents.
However, from comparison of the experimental rate constants
with the diffusion-controlled limits estimated from the viscosities
of the solvents, it is concluded that encounters between reactants
in ionic liquids are more fruitful than encounters in water and
much more than those in nonpolar solvents.
(6) Buxton, G. V.; Greenstock, C. L.; Helman, W. P.; Ross, A. B. J.
Phys. Chem. Ref. Data 1988, 17, 513.
(7) Radiolysis of water produces hydrated electrons along with hydroxyl
radicals and small amounts of hydrogen atoms. The alcohol serves as a
scavenger for the OH radicals and permits observation of the reaction of
the hydrated electrons with other solutes. The use of high pH is for
preventing the reaction of hydrated electrons with protons that are also
formed in the radiolysis.
(8) Bobrowski, K. J. Phys. Chem. 1980, 84, 3524.
Solvated electrons that are radiolytically produced in R4NNTf2
do not react rapidly with the solvent and can react with solutes.
Thus it was possible to reduce benzophenone to its ketyl radical
and to determine the rate constant for electron transfer from
this radical to duroquinone. Although the experimental rate
constant is relatively low, it is close to the diffusion-controlled
limit in this solvent. The rate constants for the same reaction in
water and in glycerol are also close to the respective diffusion-
controlled limits, but the value is much lower in 2-PrOH. In
contrast with R4NNTf2, solvated electrons produced in imida-
zolium or pyridinium ionic liquids are rapidly scavenged by
the solvent cations. In the case of BMI+ the electrons are
completely trapped, as discussed before,2 but with BuPy+ and
BuPic+ it is possible to follow the electron transfer from the
solvent radicals to various solutes. Surprisingly, several such
reductions occur with rate constant that are significantly higher
than the diffusion-controlled limit, unlike the rate constants for
the same reactions in water and 2-PrOH, which are slower than
the diffusion-controlled limit. This very fast reaction is inter-
preted as an electron hopping mechanism, whereby the electron
reaches its final destination via intervening pyridinium groups
without requiring the diffusion of a specific radical to approach
a reactant molecule.
(9) (CF3SO2)2NH is a strong acid, with pKa ) 1.7 (Foropoulos, J., Jr.;
DesMarteau, D. D. Inorg. Chem. 1984, 33, 3720).
(10) By comparison with the behavior of the succinimidyl radical: Lind,
J.; Jonsson, M.; Eriksen, T. E.; Merenyi, G. J. Phys. Chem. 1993, 97, 1610.
Lind, J.; Jonsson, M.; Shen, X.; Eriksen, T. E.; Merenyi, G.; Eberson, L. J.
Am. Chem. Soc. 1993, 115, 3503.
(11) Rao, P. S.; Hayon, E. J. Phys. Chem. 1973, 77, 2274.
(12) Neta, P.; Grodkowski, J.; Ross, A. B. J. Phys. Chem. Ref. Data
1996, 25, 709.
(13) Veltwisch, D.; Asmus, K.-D. J. Chem. Soc., Perkin Trans. 2 1982,
1147.
(14) Adams, G. E.; Willson, R. L. J. Chem. Soc., Faraday Trans. 1
1973, 69, 719.
(15) Hayon, E.; Ibata, T.; Lichtin, N. N.; Simic, M. J. Phys. Chem. 1972,
76, 2072.
(16) Espenson, J. H. Chemical Kinetics and Reaction Mechanisms, 2nd
ed.; McGraw-Hill: New York, 1995; pp 201-202.
(17) Watanabe, T.; Honda, K. J. Phys. Chem. 1982, 86, 2617.
(18) Neta, P.; Simic, M. G.; Hoffman, M. Z. J. Phys. Chem. 1976, 80,
2018.
(19) In the aqueous TEOA solutions, when the concentration of p-NAP
was reduced below 0.3 mmol L-1, the kinetic traces indicated two formation
steps. The slower step became more pronounced as the concentration of
p-NAP was decreased and its rate constant was estimated to be ∼4 × 108
L mol-1 s-1. We interpret this observation by considering that two radicals
can be formed from TEOA, (HOCH2CH2)2NC˙ HCH2OH and (HOCH2CH2)2-
NCH2C˙ HOH, both reducing p-NAP with very similar rate constants, but
the second radical can undergo base-catalyzed elimination of (HOCH2CH2)2-
•
NH to yield CH2CHO (Gilbert, B. C.; Larkin, J. P.; Norman, R. O. C. J.
Ionic liquids also affect the reduction potentials of certain
couples, e.g., by preferential solvation of the reduced or oxidized
species as compared with the solvation of these species in other
solvents. It has been demonstrated that the equilibrium between
the MV2+/MV•+ couple and the AQS-/AQS•2- couple reverses
direction upon going from water to 2-PrOH.23 We confirm these
findings and show that the equilibrium between the MV2+/MV•+
couple and the DQ/DQ•- couple also reverses direction upon
going from water to 2-PrOH. In all cases the electron transfer
rate constants are fairly high. On the other hand, these electron
transfer reactions are much slower in BuPyBF4 and the direction
of the equilibrium is the same as in 2-PrOH and opposite that
in water. The slowness of the electron transfer is partly due to
the high viscosity and partly due to a lower driving force for
the reaction, leading to a very large solvent effect on the rate
constant. The change in driving force, i.e., the shift in equilib-
rium 16 and 17, appears to be due to enhanced solubility of
species with fewer negative or positive charges. This may be
taken as another indication that this ionic liquid is highly
associated, thus solubilizing neutral species better than ionic
species. The effect of the microenvironement on the reduction
potentials of the reactants is demonstrated by the finding that
addition of small amounts of water to the ionic liquid has a
very large effect on the electron transfer rate constant.
Chem. Soc., Perkin Trans. 2 1972, 794. Steenken, S. J. Phys. Chem. 1979,
83, 595) while the first radical cannot do so. As the concentration of p-NAP
is decreased the rate of reduction is accordingly decreased, permitting a
more substantial fraction of the second radicals to undergo elimination. This
results in a decrease in the contribution of the fast step and an increase in
the contribution of the slow step. Since •CH2CHO is not a reducing radical,
the slow step probably involves reaction with p-NAP via addition, as do
various other radicals (Jagannadham, V.; Steenken, S. J. Am. Chem. Soc.
1984, 106, 6542; 1988, 110, 2188).
(20) Ledwith, A. In Biochemical Mechanisms of Paraquat Toxicity;
Autor, A. P., Ed.; Academic Press: New York, 1977; p 21.
(21) Wardman, P.; Clarke, E. D. Br. J. Cancer 1987, 5, Suppl. 8, 172.
(22) Wardman, P. J. Phys. Chem. Ref. Data 1989, 18, 1637.
(23) Pal, H.; Mukherjee, T. J. Indian Chem. Soc. 1993, 70, 409.
(24) Calculated from thiocyanate dosimetry by correcting for the
difference in electron density of the solutions and assuming that the molar
absorption of TMPD•+ is 1.2 × 104 L mol-1 cm-1, similar to that determined
in CCl4 solutions by Burrows, H. D.; Greatorex, D.; Kemp, T. J. J. Phys.
Chem. 1972, 76, 20.
(25) Calculated by assuming that the molar absorption coefficient of
ClPz•+ in MtBANTf2 is similar to that in water (1.07 × 104 L mol-1 cm-1):
Pelizzetti, E.; Meisel, D.; Mulac, W. A.; Neta, P. J. Am. Chem. Soc. 1979,
101, 6954.
(26) Neta, P.; Huie, R. E.; Mosseri, S.; Shastri, L. V.; Mittal, J. P.;
Maruthamuthu, P.; Steenken, S. J. Phys. Chem. 1989, 93, 4099.
(27) Alfassi, Z. B.; Huie, R. E.; Kumar, M.; Neta, P. J. Phys. Chem.
1992, 96, 767.
(28) Alfassi, Z. B.; Huie, R. E.; Neta, P. J. Phys. Chem. 1993, 97, 7253.
(29) Neta, P.; Huie, R. E.; Maruthamuthu, P.; Steenken, S. J. Phys.
Chem. 1989, 93, 7654.