PAPER www.rsc.org/greenchem | Green Chemistry
Kinetics and solvent effects in the synthesis of ionic liquids: imidazolium
Jay C. Schleicher† and Aaron M. Scurto*
Received 19th May 2008, Accepted 5th February 2009
First published as an Advance Article on the web 26th February 2009
DOI: 10.1039/b808364a
Ionic liquids (ILs) are being considered as a promising class of potentially
environmentally-friendly (“green”) solvents and materials for use in a variety of applications.
However, ionic liquids are conventionally synthesized by batch, without known kinetics, in
non-sustainable solvents. For ILs to be a truly “green” technology for widespread use, they must
themselves be made in a correspondingly benign manner for low cost, as enabled by process
development. This investigation will illustrate the kinetics and large solvent effects in the synthesis
of 1-hexyl-3-methyl-imidazolium bromide in 10 different solvents: acetone, acetonitrile,
2-butanone, chlorobenzene, dichloromethane, dimethyl sulfoxide (DMSO), ethyl formate, ethyl
lactate, methanol, and cyclopentanone. The kinetic rate constant for the synthesis in DMSO is
over an order-of-magnitude larger than that in methanol. While the kinetic rate of these type of
SN2 reactions is generally known to increase with solvent “polarity”, multi-parameter solvent
descriptors, e.g. of Kamlet and Taft, are required to quantify these effects in a Linear Solvation
Energy Relationship. These relationships are used with environmental and toxicity databases, such
as the Rowan Solvent Selection Table, to rapidly optimize the solvent for favorable kinetics and
minimal human and environmental impact.
1. Introduction
Ionic liquids (ILs) have been touted as the next great class
of environmentally-friendly solvents due to their molecularly
“tunable” properties and lack of volatility. New applications
are being developed at a rapid pace. ILs have shown promise
in enhancing catalytic activity, selectivity, stability and ability
for recycling in various catalyzed reaction systems, such as
hydroformylation reactions,1,2 and even in a wide range of
biocatalytic transformations.3 ILs have a number of uses in
various separation process. For instance, ILs can “break” a
number of azeotropes;4 separate gases,5 including refrigerant
gases;6 desulfurize diesel fuel;7 dissolve and process cellulose and
other carbohydrates;8 are involved commercially in a biphasic
acid scavenging processes;9 etc. In the field of analytical chem-
istry, ILs can be used as stationary phases in chromatography
and other separation and detection techniques.10 ILs have
a long history in the field of electrochemistry11 with more
recent examples in electro-nanomaterial technologies.12 Active
pharmaceutical ingredients (APIs) have been re-formulated as
ionic liquids themselves13 to overcome difficulties with solid-
state polymorphic forms and other processing challenges, etc.
Despite all of the numerous chemistries and applications
possible, reports of the synthesis of ionic liquids often include the
very solvents that they will purportedly replace. Solvents such
Department of Chemical and Petroleum Engineering, Department of
Chemistry, and NSF-ERC Center for Environmentally Beneficial
Catalysis, University of Kansas, Lawrence, KS 66045, USA.
E-mail: ascurto@ku.edu; Fax: +1 (785) 864 4967; Tel: +1 (785) 864
4947
† Current address: GE, Inc., 1233 West Loop South, Houston, TX
77027, USA
as dichloromethane, 1,1,1-trichloroethane, petroleum ether,
toluene, etc. have all been used in their synthesis.14–18 Moreover,
ILs are often too costly to be utilized as an alternative solvent
in many large-scale industrial processes.19 This is primarily due
to small batch production and nearly non-existent kinetic and
thermodynamic data of their synthesis that has resulted in little
emphasis on reaction engineering and process intensification in
the literature. For ionic liquids to be truly “green” and to be used
ubiquitously, they must be made in a corresponding benign way
in potentially large quantities and for low cost.
1.1. Reactions: solvent or solvent-less?
Is a solvent necessary for the synthesis of ionic liquids? Would
not a neat reaction be preferred over even an environmentally-
benign solvent? Typically, quaternization reactions are highly
exothermic reactions. It has been reported that the heat of reac-
tion for 1-methylimidazole and 1-bromobutane is -96 kJ mol-1
by Waterkamp et al.20 They estimated the adiabatic temperature
for a run-away reaction for this system to be 48 ◦C and
greater.20 This amount of heat release can cause a number of
safety concerns along with poor quality product. Burrell et al.21
investigated synthesis of larger scale quantities (~1 kg) of ILs
and warn: “Caution: [this] reaction is exothermic and cooling is
advisable for large scale reactions”. In a controlled experiment
in our laboratory, when 200 mL of 1-methylimidazole were
mixed at room temperature with 200 mL of 1-bromoethane
in a 500 mL round-bottom flask, the mixture began to over-
boil within approximately 15 minutes. However, a solvent would
help dissipate and manage much of the heat generated on both
a laboratory scale and larger.
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In addition,many ILs or their intermediates can often be quite
viscous.22,23 Thus, as the reaction proceeds, uniform mixing for
heat andmass transport issuesmaybecomea concern.24 Byusing
solvents, the viscosity can be kept relatively low.25 Moreover,
some ILs or their intermediates are actually higher melting-
point solids, which would require much different processing
techniques than for liquid solutions. For instance, the liquid-
phase reaction between 1-methylimidazole and 1-bromoethane,
mentioned above for its exothermicity, forms a solid compound
at room-temperature and melts to a viscous liquid at 76 ◦C.26
By using a solvent, the product can be kept in a relatively low-
viscosity liquid solution that can be more easily processed. As
will be shown in the following sections, the reaction kinetics can
actually be faster in some solvents than in the neat reaction.
1.2. Overview
The current study quantitatively examines the solvent effects
on the kinetics in the synthesis of a model ionic liquid,
1-hexyl-3-methylimidazolium bromide [HMIm][Br] in a wide
variety of conventional and low-toxicity organic solvents. This
contribution will demonstrate that safer/more-benign solvents
can be utilized for producing ILs, while maintaining a high rate
of reaction. Solvent selection was determined by understanding
the effect of polarity on the kinetic rate constant through
correlation with Kamlet Taft (KT) polarity parameters27 in a
Linear Solvation Energy Relationship (LSER). These corre-
lations, along with solvent toxicity and environmental impact
data, then enabled the rapid optimization of solvents for both
productivity and low impact. These results will allow larger-scale
production of ILs, which will ultimately decrease their cost.
2. Background
2.1. Ionic liquid synthesis
ILs aremost commonly synthesized by a quaternization reaction
of a substituted amine (or phosphine, N-heterocycles, etc.)
with alkyl halides28–32 (see Fig. 1), which is followed by anion
exchange if necessary. Throughout the quaternization reaction,
two neutral reactants form oppositely charged ions through
a polar transition state. The rate of reaction is influenced
heavily by the “polarity” of the reaction mixture or solvent.33,34
This was first demonstrated by Menshutkin over a century
ago while studying nucleophilic substitution reactions (SN2)
between amines with haloalkanes in 23 solvents.27,35 As will be
discussed below, solvent “polarity” can be further described by
several different features. Over the past century, a large number
of studies have been performed examining the solvent effects
on the transition state in the formation of ionic compounds.
This investigation will illustrate the large difference in reaction
kinetics in different solvents for the synthesis of 1-alkyl-3-
methyl-imidazolium ionic salts and liquids.
2.2. IL process development
From the literature, only a few studies have examined reaction
engineering or process intensification for the production of
ILs. The issue of heat removal from a neat reaction may be
aided by proper reactor engineering solutions. For instance,
Waterkamp et al.20 have used micro-reactors to produce 1-butyl-
3-methyl-imidazolium bromide, as micro-reactors have high
surface area for heat removal. Minnich et al.36 investigated the
kinetics of producing 1-ethyl-3-methylimidazolium ethylsulfate
for use in micro-reactors. However, micro-reactors can only
be used for liquids and require adequate pumping for the
high-viscosity liquefied product. Varma and Namboodiri37 and
Deetlefs and Seddon38 have shown that ILs can be synthesized
more rapidly using microwave radiation which decreases the
relative reaction times.Leveque et al.39 have found that ultrasonic
irradiation is a useful tool for making a large number of ILs
in a one-pot synthesis method. Grosse-Bo¨wing and Jess have
reported the bimolecular kinetic constants for the production
of 1-butyl-3-methyl-imidazolium chloride.40 Recently, they have
described some of the important reactor engineering properties
for the neat synthesis of 1-ethyl-3-methyl-imidazolium ethylsul-
fate ionic liquids.41
2.3. “Green”/sustainable solvent selection
Many factors are involved in the selection of an “ap-
propriate” solvent42 for syntheses based upon principles of
“green”/sustainable chemistry and engineering.43,44 For the
production of any chemical, high kinetic rates are preferred.
Facile and low-energy separations are needed to purify the ionic
liquid. The human and environmental impact, especially of the
solvent, is an important aspect in designing safer and more
sustainable processes. However, as will be shown below, some
of the more benign solvents may yield slow reaction rates and
energy-intensive separations processes. Thus, a balance must be
made between the inherent toxicity for a given solvent with the
desired rates of reaction to produce ILs, and any subsequent
processing. Life-cycle assessment (LCA)45–47 is a methodology
to determine the effect of various process parameters on the
required energy, environmental impact, and profitability for
the process from cradle-to-grave. Kralisch et al.48,49 employed
a modified LCA to analyze the synthesis of several ionic liquids
and catalytic processes in them, based upon preliminary or
estimated data. For instance, the energy inputs to create the
starting materials for imidazolium based ILs are larger than
for pyridinium and quaternary ammonium ILs. Increasing
the temperature of the synthesis initially decreases the overall
impact, but reaches a minimum below 100 ◦C. While all of the
aforementioned aspects of sustainable processes are important,
this contribution will focus on the solvents of ionic liquid
synthesis, efficiency of the reaction kinetics, and address the
issue of separations.
Fig. 1 Reaction between 1-methylimidazole and 1-bromohexane forming 1-hexyl-3-methylimidazolium bromide [HMIm][Br].
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There are many different ways that solvents can be classified
as to their effect on humans and the environment. The US Food
andDrugAdministration (FDA) typically uses a system of three
classes: 1 through 3; in addition to Generally-Regarded-as-Safe
(GRAS) solvents.50–52 Under these guidelines any Class 1 solvent
shall not be used in the production of pharmaceuticals as they
are known or suspected human carcinogens. Class 2 solvents
have been deemed toxic and should be limited in the production
of pharmaceuticals. However, the effects from exposure are
reversible and acceptable limits have been established (permitted
daily exposure (PDE)). Class 3 solvents are accepted in the
production of pharmaceuticals and have low toxicity. GRAS
solvents are preferred solvents, which are accepted in the
production of pharmaceuticals, and some have been deemed
safe as food additives.
While the FDA guidelines are a good indication of human
toxicity, they do not include aspects of environmental impact
such as aquatic toxicity, smog emissions, biodegradation poten-
tial, global warming potential, etc. In addition, the weighting of
these different measures of toxicity and environmental impact is
also an important issue. Onemethodwhich takes a large number
of health and environmental concerns into consideration is
the Rowan Solvent Selection Table (RSST).53,54 Originally, the
various indicators were weighted to yield an objective function,
called the Pharmaceutical Index, which combines human toxi-
city indices, i.e. ingestion, inhalation, and carcinogen toxicities,
with environmental effects, e.g. aquatic toxicity, soil absorption,
ozone depletion, smog formation, etc. While the weighting
of these different parameters can be discussed, debated and
modified, the results do give a good indication of the relative
effect on people and the environment. The RSST will be used
extensively in this study.
Another useful guide for determining solvent selection is
the GlaxoSmithKline (GSK) solvent selection table, which is
based on the “International Conference on Harmonization of
Technical Requirements for Registration of Pharmaceuticals for
Human Use” (ICH guidelines).55 In this solvent selection table
there are nine assessments which are examined for solvent se-
lection, and include: incineration, recyclability, volatile organic
compounds (VOC’s), bio-treatment, environmental impact on
air, environmental impact of water, health hazards, exposure
potential, and safety hazards.55 Each assessment is equally
weighed in the ranking of solvents.
While the solvents with the lowest impact on humans and the
environment could be chosen at the outset, the kinetics of the
synthesis and subsequent separations may actually result in a
much worse net effect. In order to expedite solvent screening,
a method has been utilized here to correlate and predict
solvent performance (kinetics) based upon linear solvation
energy relationship (LSER) regression using Kamlet Taft (KT)
parameters for the individual solvents.
2.4. Kamlet Taft polarity scales and LSER
Rates of reaction in the production of ionic liquids are highly
dependent on the solvent media. A method to quantitatively
correlate and predict the kinetics based upon the properties
of the solvent, such as polarity, would be highly useful. The
term “polarity” embodies a number of different concepts,
including dipole moment, dielectric constant, hydrogen bond
accepting ability, polarizability, etc. While one-parameter scales
for polarity, such as the ET (30) scale27,56,57 can approximate
qualitative trends, they often cannot be used to quantitatively
correlate reaction rates. Kamlet Taft (KT) parameters differen-
tiate various aspects of “polarity”, viz. acidity (a), basicity (b),
and dipolarity/polarizability (p*). Acidity, a, is a measure of
the solvent’s ability to donate a proton in a solvent-to-solute
hydrogen bond,58 b is the measure of the solvent’s ability to
accept a proton in a solvent-to-solute hydrogen bond,59 and
p* is a measure of the solvent’s ability to stabilize a charge or
dipole.60,61 The KT parameters of the solvent can be used to
correlate and predict kinetic rate constants in different solvents
using a Linear Solvation Energy Relationship (LSER). The
LSER method regresses parameters to correlate the kinetic
rate constants, k, with the solvent-dependent physicochemical
properties: a, b, and p*:27
ln k = ln k0 + aa + bb + p(p* - dd) (1)
The regressed coefficients, a, b, p, and d will indicate the
magnitude and direction (positive or negative) the polarity
parameter contributes to the kinetic rate. d is the polarizability
correction term which is equal to 0.0 for non-chlorinated
solvents, 0.5 for polychlorinated solvents, and 1.0 for aromatic
solvents.62
3. Results and discussion
3.1. Solvent effects on kinetics
Initially, this study focused on five traditional organic solvents
for the reaction of 1-methylimidazole with 1-bromohexane to
form the ionic liquid [HMIm][Br]. The initial solvents (acetoni-
trile, acetone, methanol, dichloromethane, and chlorobenzene)
have a wide range of polarity and different levels of toxicity
and environmental impact factors that will be discussed below.
For each reaction, the mole ratio of reactants to solvent was
maintained at 1 : 1 : 20 to avoid concentration effects on the
bulk polarity from the reactant and/or product. Each reaction
was conducted at three different temperatures: 25 ◦C, 40 ◦C,
and 60 ◦C, yielding insight to the transition state and activation
parameters. For the chlorobenzene system, the mixture splits
into two phases (IL-rich and reactant/solvent-rich) after ~6%
conversion. While the developing reaction may occur in either
phase, 1-bromohexane is relatively insoluble (immiscible) in the
ionic liquid. Thus, little reaction is believed to occur in the IL
phase and the overall rate reported here is also the rate in the
solvent phase. The change of concentration with time for the
synthesis of [HMIm][Br] in acetonitrile at 40 ◦C is shown in
Fig. 3 as an example. All rates of reaction at 40 ◦Cwere regressed
non-linearly assuming 2nd order kinetics and are presented in
Table 1, along with the KT parameters and ET (30) values; for
reference, the KT parameters of the reactants and IL product
are also given. It was anticipated from the well-known work
of Menshutkin35 and general SN2 reactions that the kinetic rate
would increase with increasing polarity. From the five initial
solvents, the rate of reaction is the greatest in acetonitrile, and is
more than one order ofmagnitude higher than that ofmethanol.
Methanol was the slowest despite being considered one of the
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Table 1 Rate constants at 40 ◦C, KT parameters and ET (30) values for all solvents examineda
k ¥ 106/M-1 s-1 Kamlet Taft parameters
Solvent 40 ◦C a b p* ET (30)/kcal mol-1
Dimethyl sulfoxide 77.89 ± 1.72 -0.013 ± 0.003 0.724 ± 0.009 1.032 ± 0.004 45.11 ± 0.02
Acetonitrile 21.56 ± 0.21 0.230 ± 0.009 0.376 ± 0.012 0.787 ± 0.012 45.62 ± 0.02
Cyclopentanone 15.11 ± 0.11 -0.085 ± 0.005 0.565 ± 0.004 0.748 ± 0.003 39.85 ± 0.01
Acetone 12.67 ± 0.06 0.110 ± 0.002 0.523 ± 0.012 0.715 ± 0.002 42.58 ± 0.03
2-Butanone 11.56 ± 0.08 0.053 ± 0.004 0.568 ± 0.004 0.675 ± 0.002 41.06 ± 0.06
Dichloromethane 8.47 ± 0.11 0.042 ± 0.003 -0.020 ± 0.014 0.790 ± 0.004 40.88 ± 0.02
Ethyl formate 7.97 ± 0.14 0.094 ± 0.035 0.412 ± 0.075 0.570 ± 0.042 40.19 ± 0.11
Chlorobenzeneb 3.64 ± 0.11b 0.051 ± 0.004 0.080 ± 0.009 0.624 ± 0.004 36.91 ± 0.02
Ethyl lactate 2.86 ± 0.06 0.642 ± 0.004 0.633 ± 0.010 0.689 ± 0.002 51.01 ± 0.04
Methanol 2.03 ± 0.08 0.909 ± 0.006 0.629 ± 0.009 0.697 ± 0.006 55.53 ± 0.04
1-Methylimidazole — 0.232 ± 0.012 0.712 ± 0.016 0.961 ± 0.014 44.85 ± 0.01
1-Bromohexane — 0.014 ± 0.07 -0.009 ± 0.011 0.500 ± 0.01 37.9 ± 0.70
[HMIm][Br] — 0.453 ± 0.069 0.562 ± 0.066 0.983 ± 0.037 50.49 ± 0.18
a All rates of reaction at 40 ◦C were conducted at a 1 : 1 : 20 mole ratio (1-methylimidazole : 1-bromohexane : solvent). b Mixture split into two
phases during reaction and the reported kinetic constants assume that the reaction does not occur in the IL-rich phase due to poor solubility of
1-bromohexane.
Fig. 2 Transition state for the reaction between 1-methylimidazole and 1-bromohexane.
most “polar.” Thus, simple heuristics of increasing polarity
to increase kinetic rate is not always qualitatively valid. The
reaction in acetonitrile is even greater than in the neat reaction
without solvent. The natural log of k was fit to the LSER
coefficients based on the KT parameters (acidity a, basicity
b, and dipolarizability p*):
ln k = -62.08 - 3.79a + 20.89b + 56.36(p* - 0.23d)
R2 = 0.99 (2)
Fig. 3 Concentration versus time for the formation of [HMIm][Br] in
acetonitrile at 40 ◦C.
From the regression it is seen that the p* parameter (dipolar-
ity/polarizability) has the largest positive effect on the reaction
rate, followed by the b parameter (basicity). The a parameter
(acidity) has a negative effect on the rate of reaction. This rela-
tionship of 5 solvents with a wide variety of polarity then allows
us to optimize the rate at least qualitatively: solvents with large b
and p* and a small a should be chosen, i.e. high dipolarity and
hydrogen bond acceptor/electron donor capability, and small
hydrogen bond donating/electron accepting ability.
The general effects of polarity on the synthesis of imi-
dazolium based ILs as determined from the solvent subset
can now be used to aid the choice of other solvents with
lower toxicity and environmental impact. From eqn (2) and
KT parameters, five additional solvents were selected among
Class 3 and GRAS solvents: ethyl formate, ethyl lactate (also a
bio-renewable solvent63,64), dimethyl sulfoxide, 2-butanone, and
cyclopentanone. The results for the 10 solvents are presented
in Table 1 along with the experimentally measured KT and
ET (30) parameters. A majority of the new solvents did not have
a complete set of parameters based on one set of solvatochromic
probes.
A single parameter correlation for all ten solvents using
the ET (30) scale is inadequate at qualitative and quantitative
correlation, as shown in Fig. 4. The ET (30) scale only embodies
one aspect of polarity important in this reaction. This result
was also seen by Abraham et al.65 who worked with a trimethy-
lamine/p-nitrobenzyl chloride system. A similar conclusion was
obtained by Skrzypczak and Neta34 studying the reaction of
1,2-dimethylimidazole and benzylbromide.
The LSER analysis, now for all 10 solvents, was regressed
against the kinetic rate constant:
ln k= -14.72 - 2.07a + 0.07b + 4.99(p* - 0.20d) R2 = 0.95 (3)
and the results are given in Fig. 5. The rates of reactions are
heavily influenced by the solvent choice, and can be predicted
fairly accurately using KT parameters in a LSER regression.
This phenomena can qualitatively be explained by the solvation
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Fig. 4 Correlation of kinetic constant, k, with the ET (30) polarity scale
for all solvents examined.
Fig. 5 LSER (eqn (7)) results for the 10 solvents used in this study at
40 ◦C.
scheme of Hughes and Ingold.27,66–68 For the alkylation of
1-methylimidazole with haloalkanes, two neutral reactants form
a transition state of significant charge separation, followed by
full charge separation in the product cation and anion (Fig. 2).
As seen in Table 1 for the polar aprotic solvents, an increase
in the solvent polarity (p* and ET (30)) results in an increase
in the rate of reaction. Following Hughes and Ingold’s66,68
interpretation, the reaction is dominated by the degree of charge
existing at the transition state. Solvents with higher dipolarity
and basic character can stabilize better the charged transition
state. However, polar protic solvents retard the reaction rate
despite their high dipolarity. They form hydrogen bonds with
the lone pair of electrons on the nitrogen of 1-methylimidazole,
thus inhibiting nucleophilic attack on the haloalkane.
These effect of polarity on the kinetics may also be viewed
by analysis of the Arrhenius69 and Eyring70,71 parameters. The
Eyring equation is given by:
k
k T
h
H
RT
S
R
=
-
Ê
Ë
Á
Á
ˆ
¯
˜
˜
Ê
Ë
Á
Á
ˆ
¯
˜
˜
B exp exp
D D‡ ‡
(4)
where kB is the Boltzmann’s constant, T is temperature, h is the
Plank’s constant, R is the gas rate constant, DH‡ is the enthalpy
of activation, and DS‡ is the entropy of activation.
The Arrhenius equation is given by:
k k
E
RT=
-
Ê
ËÁ
ˆ
¯˜
0
a
exp (5)
where Ea is the activation energy and k0 is the pre-exponential
term.
Table 2 summarizes the parameters for each of the models
determined from the kinetic constants at 25 ◦C, 40 ◦C, and
60 ◦C, unless otherwise specified. The values for DH‡ are similar
between the two major solvent types: polar aprotic solvents and
Table 2 Rates of reaction and kinetic parameters
k ¥ 106/M-1 s-1
Solvent 25 ◦C 40 ◦C 60 ◦C k0 ¥ 10-6/M-1 s-1 Ea/kJ mol-1 DH‡/kJ mol-1 DS‡/J mol-1 K-1
DMSO 22.22 ± 0.11 77.89 ± 1.72 322.31 ± 3.53a 2.51 63.05 60.43 -131.19
Acetonitrile 6.03 ± 0.14 21.56 ± 0.21 110.64 ± 1.42 6.51 68.73 66.11 -123.26
Cyclopentanone 3.75 ± 0.03 15.11 ± 0.11 76.11 ± 1.72 10.4 71.00 68.38 -119.37
Acetone 3.69 ± 0.17 12.67 ± 0.06 63.67 ± 0.61 2.20 67.26 64.64 -132.30
2-Butanone 2.69 ± 0.03 11.56 ± 0.08 53.75 ± 0.28 6.25 70.50 67.88 -123.60
Dichloromethane 2.28 ± 0.11 8.47 ± 0.11 —b 1.85 67.98 65.44 -133.45
Ethyl formate 1.61 ± 0.03 7.97 ± 0.14 —b 507 82.75 80.21 -86.78
Chlorobenzenec 1.11 ± 0.06 3.64 ± 0.11 24.64 ± 1.39 7.51 73.43 70.81 -122.07
Ethyl lactate 0.53 ± 0.03 2.86 ± 0.06 20.28 ± 0.28 636 86.05 83.43 -85.18
Methanol 0.42 ± 0.03 2.03 ± 0.08 17.14 ± 0.11 987 87.85 85.24 -81.52
Neatc ,d 4.53 ± 0.04 17.63 ± 0.06 106.34 ± 13.2 51.2 74.57 71.95 -106.12
a Reaction conducted at 1 : 1 : 80 mole ratio 1-methylimidazole : 1-bromohexane : dimethyl sulfoxide due to high exothermicity of more concentrated
solutions. b Exceeds boiling point of solvent. c Mixture split into two phases during reaction and the reported kinetic constants assume that the
reaction does not occur in the IL-rich phase due to poor solubility of 1-bromohexane. d 1-methylimidazole : 1-bromohexane = 1 : 1 by mole.
Fig. 6 1H NMR chemical shifts for reactants and product.
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polar protic solvents (methanol, ethyl lactate). The values of
DH‡ for the polar protic solvents are approximately 20 kJ mol-1
greater than the polar aprotic solvents. The kinetic parameters
in ethyl formate (EF) are similar to those of the polar protic
solvents, despite EF having a low acidity (a parameter). Ethyl
formate is susceptible to water, base, acid, and salt catalyzed
hydrolysis and decomposition72–76 to products with higher a
values, such as ethanol, formic acid, etc. Haberfield et al.77
suggest, for general SN2 reactions, that the difference between
the DH‡ for protic solvents may be due to stabilization of
the reactants relative to aprotic solvents, thus increasing the
energy difference between the reactants and the transition state.
Additional energywould be needed to break the hydrogen bonds
between the solvent and 1-methylimidazole in order for the
reaction to occur. Hydrogen bond strengths can vary between
2–20 kJ mol-1,78 which is of the same order as the difference
in DH‡ between the polar protic and aprotic solvent classes.
Moreover, the DS‡ for the polar protic solvents is approximately
40 J mol-1 K-1 more (less negative) than the polar aprotic
solvents. Hydrogen bonding between the 1-methylimidazole and
the solvent would induce a more-ordered ground state, i.e. lower
entropy. Thus, the degree of order in the transition state is closer
to the ground state with protic solvents.
The current example uses 1-bromohexane as the alkylation
agent. Many applications with ionic liquids, e.g. catalysis, are
adversely affected by the presence of halide and its removal after
anion-exchange is often tedious. As such, several alternative
halide-free alkylation techniques have emerged. Bonhoˆte et al.79
have used ethyl triflate and ethyl trifluoroacetates for the
alkylation of 1-methylimidazole. Ue et al.80 have used the alky-
lation of 1-ethylimidazole with dimethyl carbonate. Lectercq
et al.81 have used a complex mechanism involving the reaction
between tetrahydrofuran or 1,4-dioxane, with triflic anhydride
to form a diester that is then reacted with an N-substituted
imidazole. Holbrey et al.16 used an alkylation technique between
dimethyl sulfate or diethyl sulfate with N-substituted imidazole
to form imidazolium alkyl sulfate ILs. Yoshizawa et al.82 have
synthesized a number of zwitterionic type ILs by reacting
imidazoleswith 1,3-propanesultone. The kinetic rates and effects
of solvents of several of these alternative alkylations are under
investigation.
3.2. Diffusion limited kinetic constant
The diffusion-limited rate constant was calculated for com-
parison to the intrinsic kinetic rate constants. The binary
diffusion coefficients for 1-methylimidazole and 1-bromohexane
in dimethyl sulfoxide (DMSO) have been determined at 25 ◦C
in another study in our laboratory83 and is illustrated in Table 3.
DMSO was chosen as a model solvent as it simultaneously
produces the largest kinetic rate of synthesis and is the most
viscous solvent studied. Thus, the reaction in DMSO should
be the nearest to the diffusional limited regime. As shown
in the table, 1-methylimidazole diffuses slightly faster than
1-bromohexane.
The analysis for the diffusion-limited rate constant for a
reaction in a homogeneous medium is given by:78
kd = 4pR* D NA (6)
Table 3 Calculating the kd in DMSO for the reaction between
1-methylimidazole and 1-bromohexane
Da (1-methylimidazole)a 6.10 ¥ 10-10 m2 s-1
Db (1-bromohexane)a 5.90 ¥ 10-10 m2 s-1
D = (Da + Db) 1.20 ¥ 10-9 m2 s-1
Viscosity (DMSO)b 1.987 ¥ 10-3 kg m-1 s-1
Ra 1.80 ¥ 10-10 m
Rb 1.86 ¥ 10-10 m
R* = (Ra + Rb)/2 1.83 ¥ 10-10 m
kd 1.66 ¥ 109 M-1 s-1
k (at 25 ◦C) 2.22 ¥ 10-5 M-1 s-1
a Da andDb are binary diffusion coefficients in DMSO at 25 ◦C from ref.
83. b Taken from ref. 92.
where kd is the diffusion-controlled rate constant, R* is the
distance at which the reaction occurs (assumed to be the average
distance between the reactants hydrodynamic radii), D is the
sumof the reactants diffusion coefficients at the concentration of
the reaction, and NA is Avogadro’s constant. Using the Stokes–
Einstein equation, the hydrodynamic radii for eachmolecule can
be calculated by:78
R
k T
D
R
k T
Da
B
a
b
B
b
= =
6 6ph ph
and (7)
whereRa andRb are the hydrodynamic radii, kB is the Boltzmann
constant, T is the temperature in K, h is the viscosity of the me-
dia, and Da and Db are the diffusion coefficient for each species
in the solvent media. The equations above assume a hard sphere
model for the hydrodynamic radius and yield an approximation
of the relative rate in a diffusion-limited regime. The viscosity
of the media is assumed to be the viscosity of pure DMSO
and the diffusion coefficient in dilute conditions is assumed to
representative of the actual reaction. Details of the analysis are
given in Table 3. The kinetic constant in a diffusion-controlled
regime, kd, is estimated to be 1.66 ¥ 109 M-1 s-1, which is 14 orders
of magnitude greater than the reported intrinsic rate constant.
The speed of these SN2 reactions requiresmore than just the reac-
tants coming into proximity as suggested by the solvent effects
discussed above. As expected, the kinetics rates reported here
truly reflect the intrinsic kinetics and indicate that methods to
increase mass transfer within the solution, such as stirring, etc.,
are more important for heat transfer issues than mass transfer.
3.3. Green/sustainable solvent selection
Table 4 lists the solvents in order of decreasing rate of reaction for
1-bromohexane with 1-methylimidazole (DMSO is the fastest
and methanol the slowest). For solvents that have not been
examined in the Rowan Solvent Table, the closest related
solvent was used: methyl formate for ethyl formate, methyl
lactate for ethyl lactate, cyclohexanone for cyclopentanone, and
o-dichlorobenzene for chlorobenzene. According to the index,
the top 5 most benign solvents of the 10 studied are: DMSO >
methyl(ethyl) lactate > acetone > methanol > methyl (ethyl)
formate. The GSK solvent selection guide yielded similar con-
clusions with the top 3 benign solvent choices being: DMSO >
acetone > methanol. However, the best solvents for the highest
kinetic rates are: DMSO > acetonitrile > cyclopentanone >
acetone > 2-butanone. On first inspection, DMSO may seem
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Table 4 Abbreviated results adapted from the RSST for determining the pharmaceutical index
Solvent Indexa
Inhalation
TLV (ppm)b
Ingestion/
mg kg-1 ratb
Carcinogen
[0–5]b
Aquatic/
mg L-1 fishb
Soil adsorption
coefficientb
Smog
formationb
Oct/water partitioning
coefficientb
DMSO 0.42 1000 14 500 0 66 901 0.65 0.00 -1.35
Acetonitrile 3.21 40 3800 0 4111 0.65 0.00 -0.34
Cyclohexanonec 5.37 50 1800 3 754 1.18 0.53 0.81
Acetone 2.15 500 5800 1 6967 0.30 0.18 -0.24
2-Butanone 4.00 200 2737 2 3173 0.58 0.51 0.29
Dichloromethane 5.36 50 1600 3 428 1.38 0.03 1.34
Methyl formatec 3.11 100 475 1 6260 0.33 0.00 0.03
Methyl lactatec 1.57 1000 2000 0 29 583 0.00 0.00 -0.67
Methanol 2.52 200 5628 0 8403 0.00 0.21 -0.77
a Pharmaceutical index weighting. b Values taken from ref. 53,54. c Analogues of solvents used in the kinetic analysis.
Table 5 Boiling points, latent heats, sensible heats and total energy obtained for energy analysis
Solvent
Molecular
weight/g mol-1
Boiling point,
Tb/◦Ca
DHvap(Tb)a/
kJ mol-1
Cp at 25 ◦Ca/
J g-1 K-1
Sensible heatb/
kJ mol-1
Total heat/
kJ mol-1
Dimethyl sulfoxide 78.13 189.0 43.1 1.958 22.79 65.89
Acetonitrile 41.05 81.7 29.75 2.229 3.84 33.59
Cyclopentanone 98.14 130.6 36.35 1.840 16.43 52.78
Acetone 58.08 56.1 29.1 2.175 2.06 31.16
2-butanone 72.1 79.6 31.3 2.201 6.35 37.65
Dichloromethane 84.93 40.0 28.06 1.192 0.0 28.06
Ethyl formate 74.08 54.4 29.91 2.015 2.09 32.00
Chlorobenzene 147.01 131.7 35.19 1.334 17.85 53.04
Ethyl lactate 118.13 154.5 52.5c 2.150 28.70 81.20
Methanol 32.04 64.6 35.21 2.531 2.00 37.21
a Taken from ref. 92. b Sensible heat = Cp(Tb -40 ◦C). c Ref. 93.
the ideal solvent to produce [HMIm][Br]: the highest reaction
rate and lowest toxicity/environmental impact.However, what is
not considered here is the energy to separate the solvent from the
IL using distillation/evaporation. Consideration of the boiling
point and heat of vaporization may yield a peripheral energy
analysis. In actual distillation/evaporation, one would need
to consider the phase equilibrium thermodynamics (activity
coefficients, etc.) to design the separation train and energy
requirements; these are often more than simply the sensible heat
and latent heat of vaporization.55 Table 5 lists the boiling points,
heats of vaporization, sensible heat, and the total energy for the
solvents investigated. As seen in Table 5, DMSO’s boiling point
and heat of vaporization (thus energy requirement) is quite high
compared to the other solvents in the table. This excess energy
would result in higher energy usage and, thus, more pollution,
which in turn would worsen the actual measure of toxicity and
environmental impact of the solvent. Based on comparison
of the reaction rates, the solvents’ Index, and a scan of the
energy requirements of separation, acetone appears to possess
an optimal combination of properties for use, at least on the
bench-scale, if not on an industrial scale. In addition, acetone is
relatively inexpensive, can be purchased on a large scale, and can
be produced by bio-renewable methods (ABE fermentation).84
Other means of separation should also be considered for a
complete analysis. These techniques range from liquid extraction
tomore recent techniques using low tomoderate pressure ofCO2
to induce phase splitting.85–87 CO2 has been found to induce a
broad range of IL–solvent mixtures to split into an IL-rich and
solvent-rich phase that can be decanted, or, at higher pressures,
extracted by near- or supercritical CO2. Scurto88 indicates that
depending on the needed purity, the energy needed to power
a compressor for a CO2 separation process may be competitive
with the energy (heat) requirements for evaporation/distillation,
even with relatively high volatility solvents such as methanol.
4. Experimental
4.1. Kinetic measurements
The reaction of 1-methylimidazole and haloalkanes with differ-
ent solvents was conducted at three different temperatures using
a multi-well reactor block from Chemglass, Inc. (model number
CG-1991–03) which holds 16 standard 20 mL scintillation vials.
The temperature and stirring of the reactor block was main-
tained using an IKAMAG RET basic hotplate equipped with
an ETS-D4 fuzzy logic temperature controller that maintained
the temperature to ±1 ◦C. Reactants were carefully weighed
to 0.01 mg in 20 mL scintillation vials and a magnetic stir
bar was placed in each vial. The vials were then placed in
the reactor block and samples were drawn from each vial and
placed inNMR tubes containing deuterated chloroform at room
temperature. The samples were quickly analyzed using a Bruker
400 MHz Nuclear Magnetic Resonance (NMR) spectrometer.
Typically, the time duration between sample extraction and
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NMR analysis was less than 5 minutes; further conversion
during this timewas negligible due to the relatively slower kinetic
rates at ambient conditions, and the more dilute concentrations
after adding deuterated chloroform. The conversion over time
was determined by following the disappearance of the reactant
and appearance of the product peaks on the NMR spectrum.
The methyl peaks for the reactant, 1-methylimidazole and the
corresponding methyl peaks for [HMIm][Br] at dH ~3.65 ppm
and dH ~4.16 ppm, respectively (see Fig. 6), are integrated to
determine the fractional conversion, X , using the following
equation:
X
C C
C
N
N
N
N N
I
I I
=
-
= =
+
=
+
M M
M
IL
M
IL
M IL
H
H H
Im Im
Im Im Im
.
.
0
0 0
4 16
4 16 3
d
d d .65
(8)
where N is the number of moles of each species at any point in
time,N0 is initial amounts of limiting reactant, and I is the peak
area at each of the chemical shifts being analyzed. Alternatively,
the conversion could be determined based on the difference in
the methylene peak next to the bromine (dH 3.40 ppm) or the
imidazolium ring (dH 4.37 ppm); see Fig. 6. When the individual
solvent peaks overlap with these peaks in the NMR spectra,
the peaks at dH 7.02 ppm and dH 10.28 ppm were used for
calculating conversion, which correspond to the hydrogen in
position number 5 on the 1-methylimidazole and the hydrogen
in position number 2 on the [HMIm][Br]; see Fig. 6. For further
details of the NMR technique, see Schleicher.89 The accuracy of
this method to determine X has been estimated to ±1%.
The reaction rates were confirmed to be 2nd order/
bimolecular according to the expression:
r r r
C
t
kC CIL M BrHex
M
M BrHex= - = - = -
∂
∂
Ê
ËÁ
ˆ
¯˜
=Im
Im
Im (9)
or alternatively:90
r r C
X
t
k C C X C C XIL M M M
0
M
0
BrHex
0
M
0= - =
∂
∂
Ê
ËÁ
ˆ
¯˜
= - -Im Im Im Im Im( )( )
0
(10)
where r is the reaction rate based on component i, k the kinetic
constant, C0i the initial concentrations of the components
(molarity), and X is the fractional conversion (eqn (8)).
When the initial concentrations of the two reactants are equal
(1 : 1 stoichiometry), eqn (10) becomes:
r r C
X
t
k C X
X
C X
kt
IL M M M
0
M
0
= - =
∂
∂
Ê
ËÁ
ˆ
¯˜
= ( ) -
-
=
Im Im Im
Im
( )
( )
0 2 21
1
(11)
While the kinetic constant could be obtained by graphing the
experimental data to a linearized form of eqn (11), this often
introduces unnecessary errors or undue emphasis on certain
time regimes. A non-linear regression technique is used here in
the software SigmaPlot 2000.
4.2. Kamlet Taft measurements
Many different solvatochromic probes can be used to determine
the KT parameters. Different sets of dyes produce slightly
different results and, thus, comparison with other studies should
be made only with similar dye sets.58–60,91 In this study, the sol-
vatochromic probes: N,N-diethyl-4-nitroaniline, 4-nitroaniline,
and Reichardt’s dye, were used to calculate the three solvent
parameters a, b, and p*. All KT parameters and ET (30) values
were obtained from solutions with the appropriate dyes using a
Varian Cary 300 Bio Ultra violet-Visible (UV-Vis) Spectropho-
tometer, with a dual cell Peltier accessory temperature controller.
The temperature was maintained at the standard 25 ◦C. The
wavelengths of maximum absorption of the dyes are related to
the KT parameters using the standard formulas.27 The LSER
coefficients were regressed using the non-linear optimization
techniques in the software: Sigma-Plot 2000 version 6.0.
4.3. Materials
Reagents: 1-methylimidazole (>99%), 1-chlorohexane (95%),
1-iodohexane (>98%), and 1-bromopropane (99%) were ob-
tained from Acros Organics, while 1-bromohexane (>99%),
1-bromopentane (99%), 2-bromopentane (95%), 1-bromo-3-
methylbutane (96%), 2-bromo-2-methylbutane (95%), and 1-
bromodecane (98%) were obtained from Sigma Aldrich. Sol-
vents: acetonitrile (>99.9%), acetone (>99.9%), methanol
(>99.9%), chlorobenzene (99.9%), dichloromethane (99.8%),
dimethyl sulfoxide (>99%), cyclopentanone (>99%), ethyl lac-
tate (>98%), 2-butanone (>99.7%), and cyclohexane (>99.9%)
were all obtained from Sigma Aldrich, while ethyl formate
(>98%) was purchased from Acros Organics. Solvatochromic
probes: 4-nitroaniline (>99%) was purchased from Sigma
Aldrich, N,N-diethyl-4-nitroaniline (97%) was purchased from
Oakwood Products Inc., and Reichardt’s Dye (>90%) was
purchased from Fluka. All starting materials were distilled and
kept under argon gas prior to use. All solvents were dried
using 3 A˚ or 4 A˚ molecular sieves. The solvatochromic probes
Reichardt’sDye, 4-nitroaniline, andN,N-diethyl-4-nitronaniline
were used as received.
5. Conclusion
The kinetic rate constants for the reaction of 1-bromohexane
with 1-methylimidazole have been determined in 10 solvents
at 25 ◦C, 40 ◦C, and 60 ◦C. Kamlet–Taft parameters in a
LSER regression can quantitatively correlate the kinetics of
reaction with the parameters of the solvent. For imidazolium
based ILs, reaction rates increase with solvents containing
high dipolarity/polarizability and basicity with low acidity.
Estimation of the diffusion-limited rate constant confirms that
the measured reactions rates represent the intrinsic kinetics.
Using the kinetic data, the toxicological and environmental data,
and the volatility of the solvents, a method has been proposed to
determine the optimal solvents for the synthesis of ILs. Acetone
has been shown to possess a number of optimal attributes for
the production of ionic liquids.
Acknowledgements
This material is based upon work supported by the USNational
Science Foundation under Grant No. CBET-0626313. The
author (AMS) appreciates the support of the DuPont Young
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Professor Award. Ms Silvia Nwosu is kindly thanked for some
experimental assistance.
References
1 A. Ahosseini, W. Ren and A. M. Scurto, Chim. Oggi, 2007, 25,
40–42.
2 M. Haumann and A. Riisager, Chem. Rev., 2008.
3 R. A. Sheldon, R.M. Lu,M. J. Sorgedrager, F. v. Rantwijk andK. R.
Seddon, Green Chem., 2002, 4, 147–151.
4 C. Jork,M. Seiler, Y.-A. Beste andW. Arlt, J. Chem. Eng. Data, 2004,
49, 852–857.
5 J. L. Anderson, J. K. Dixon and J. F. Brennecke, Acc. Chem. Res.,
2007, 40, 1208–1216.
6 M. B. Shiflett and A. Yokozeki, Chim. Oggi, 2006, 24, 28–30.
7 A. Bo¨smann, L. Datsevich, A. Jess, A. Lauter, C. Schmitz and P.
Wasserscheid, Chem. Commun., 2001, 2494–2495.
8 S. Murugesan and R. J. Linhardt, Curr. Org. Synth., 2005, 2, 437–
451.
9 R. D. Rogers, and K. R. Seddon, Ionic Liquids IIIB: Fundamentals,
Progress, Challenges, and Opportunities: Transformations and Pro-
cesses, 2005.
10 M. Koel, Crit. Rev. Anal. Chem., 2005, 35, 177–192.
11 R. J. Gale and R. A. Osteryoung, J. Electrochem. Soc., 1980, 127,
2167–2172.
12 J. N. Barisci, G. G.Wallace, D. R.MacFarlane andR.H. Baughman,
Electrochem. Commun., 2004, 6, 22–27.
13 W. L. Hough, M. Smiglak, H. Rodriguez, R. P. Swatloski, S. K.
Spear, D. T. Daly, J. Pernak, J. E. Grisel, R. D. Carliss, M. D.
Soutullo, J.H.Davis andR.D.Rogers,NewJ.Chem., 2007, 31, 1429–
1436.
14 P. Bonhoˆte, A. P. Dias, N. Papageorgiou, K. Kalyanasundaram and
M. Gratzel, Inorg. Chem., 1996, 35, 1168–1178.
15 S. V. Dzyuba and R. A. Bartsch, PhysChemPhys, 2002, 3, 161–167.
16 J. D. Holbrey, W. M. Reichert, R. P. Swatloski, G. A. Broker, W. R.
Pitner, K. R. Seddon and R. D. Rogers, Green Chem., 2002, 5, 407–
413.
17 D. R. MacFarlane, J. Golding, S. Forsyth, M. Forsyth and G. B.
Deacon, Chem. Commun., 2001, 1430–1431.
18 J. S. Wilkes, J. A. Levisky, R. A. Wilson and C. L. Hussey, Inorg.
Chem., 1982, 21, 1263–1264.
19 J. F. Brennecke and E. J. Maginn, AICHE J., 2001, 47, 2384–2389.
20 D. A. Waterkamp, M. Heiland, M. Schluter, J. C. Sauvageau,
T. Beyersdorff and J. Thoming, Green Chem., 2007, 9, 1084–
1090.
21 A. K. Burrell, R. E. Del Sesto, S. N. Baker, T. M. McClesky and
G. A. Baker, Green Chem., 2007, 9, 809–809.
22 K. R. Harris, M. Kanakubo and L. A. Woolf, J. Chem. Eng. Data,
2006, 51, 1161–1167.
23 J. M. Crosthwaite, M. J. Muldoon, J. K. Dixon, J. L. Anderson and
J. F. Brennecke, J. Chem. Thermodyn., 2005, 37, 559–568.
24 K. Kunkel and G. Maas, Eur. J. Org. Chem., 2007, 3746–3757.
25 E. J. Gonzalez, B. Gonzalez, N. Calvar and A. Dominguez, J. Chem.
Eng. Data, 2007, 52, 1641–1648.
26 E. A. Turner, C. C. Pye and R. D. Singer, J. Phys. Chem. A, 2003,
107, 2277–2288.
27 C. Reichardt, Solvent and Solvent Effects in Organic Chemistry 3edn.,
Wiley-VCH, Weinheim, Germany, 2003.
28 L.W.Deady andD. C. Stillman,Aust. J. Chem., 1976, 29, 1745–1748.
29 Y. Gao and J. M. Shreeve, Synthesis, 2004, 1072–1082.
30 M. M. Kabachnik, Y. A. Khomutova and I. P. Beletskaya,
Russ. J. Org. Chem., 1999, 35, 26–27.
31 J. Ropponen, M. Lahtinen, S. Busi, M. Nissinen, E. Kolehmainen
and K. Rissanen, New J. Chem., 2004, 28, 1426–1430.
32 J. Sun, D. R. MacFarlane and M. Forsyth, Sol. State Ionics, 1997, 3,
356–362.
33 T. Nevecna and V. Bekarek, Acta Univ. Palack. Olomuc. Fac. Rerum
Natur. Chemica XXXI, 1992, 108, 33–37.
34 A. Skrzypczak and P. Neta, Int. J. Chem. Kinet., 2004, 36, 253–258.
35 N. Menshutkin, Z. Phys. Chem, 1890, 5, 589.
36 C. B. Minnich, L. Ku¨pper, M. A. Liauw and L. Greiner, Catal.
Today, 2007, 126, 191–195.
37 R. S. Varma and V. V. Namboodiri, Chem. Commun., 2001, 7, 643–
644.
38 M. Deetlefs and K. R. Seddon, Green Chem., 2003, 5, 181–186.
39 J.-M. Leveque, G. Cravotto, L. Boffa, W. Bonrath and M. Draye,
Synlett, 2007, 2065–2068.
40 A. Grosse-Bo¨wing and A. Jess, Green Chem., 2005, 7, 230–235.
41 A.Grosse-Bo¨wing andA. Jess,Chem. Eng. Sci., 2007, 62, 1760–1769.
42 P. T. Anastas, Clean Solvents, 2002, 819, 1–9.
43 P. Anastas, and J. C. Warner, Green Chemistry: Theory and Practice,
Oxford University Press, New York, USA, 1998.
44 P. T. Anastas, Environ. Sci. Technol., 2003, 37, 423a.
45 G. Koller, U. Fischer and K. Hungerbu¨hler, Ind. Eng. Chem. Res.,
2000, 39, 960–972.
46 R. L. Lankey and P. T. Anastas, Ind. Eng. Chem. Res., 2002, 41,
4498–4502.
47 W. McDonough, M. Braungart, P. T. Anastas and J. B. Zimmerman,
Environ. Sci. Technol., 2003, 37, 434a–441a.
48 D. Kralisch, D. Reinhardt and G. Kreisel, Green Chem., 2007, 9,
1308–1318.
49 D. Kralisch, A. Stark, S. Korsten, G. Kreisel and B. Ondruschka,
Green Chem., 2005, 7, 301–309.
50 P. G. Jessop, R. Stanley, R. A. Brown, C. A. Eckert, C. L.
Liotta, T. T. Ngo and P. Pollet, Green Chem., 2003, 5, 123–
128.
51 U.S. FDA, 2007, http://vm.cfsan.fda.gov/~dms/eafus.html.
52 U.S. FDA, CFSAN/Office of Food Additive Safety, 2006,
http://www.cfsan.fda.gov/~dms/opa-appa.html#ftnE.
53 Rowan University, http://www.rowan.edu/greenengineering, ac-
cessed: January 20, 2008.
54 C. S. Slater and M. J. Savelski, J. Environ. Sci. Health Part A, 2007,
42, 1595–1605.
55 Corporate Environment and Safety Product Stewardship Guide: Sol-
vent Selection Guidelines, GlaxoSmithKline(GSK) pharmaceuticals,
2003.
56 M. J. Muldoon, C. M. Gordon and I. R. Dunkin, J. Chem. Soc.
Perkin Trans. 2, 2001, 433–435.
57 C. Reichardt, Green Chem., 2005, 7, 339–351.
58 M. J. Kamlet and R. W. Taft, J. Am. Chem. Soc., 1976, 2886–2894.
59 M. J. Kamlet and R. W. Taft, J. Am. Chem. Soc., 1976, 377–383.
60 M. J. Kamlet, J. L. Abboud and R. W. Taft, J. Am. Chem. Soc., 1977,
99, 6027–6037.
61 C. Laurence, P. Nicolet andM. T. Dalati, J. Phys. Chem., 1994, 5807–
5816.
62 M. J. Kamlet, J. M. Abboud, M. H. Abraham and R.W. Taft, J. Org.
Chem., 1983, 2877–2887.
63 J. McLaren, J. Chem. Technol. Biotechnol., 2000, 927–932.
64 D. Wilke, Appl. Microbiol. Biotechnol., 1999, 135–145.
65 M. H. Abraham, Prog. Phys. Org. Chem., 1974, 11, 1–87.
66 E. D. Hughes, Trans. Faraday Soc., 1941, 37, 603–637.
67 E. D. Hughes, M. L. Dhar, C. K. Ingold, A. M. M. Mandour, G. A.
Maw and L. I. Woolf, J. Chem. Soc., 1948, 2093–2119.
68 E. D. Hughes and C. K. Ingold, J. Chem. Soc., 1935, 244–255.
69 Arrhenius, Z. Physik. Chem., 1889, 4, 226.
70 H. Erying, Chem. Rev., 1935, 17, 65–77.
71 H. Eyring, J. Chem. Phys., 1935, 3, 107–115.
72 U. A. Chaudry and P. L. A. Popelier, J. Phys. Chem. A, 2003, 107,
4578–4582.
73 J. F. Mata-Segreda, Int. J. Chem. Kinet., 2000, 32, 67–71.
74 D. Stefanidis and W. P. Jencks, J. Am. Chem. Soc., 1993, 115, 6045–
6050.
75 L. Singh, R. T. Singh and R. C. Jha, J. Indian Chem. Soc., 1981, 58,
966–969.
76 A. B. Manning, J. Chem. Soc., 1921, 119, 2079–2087.
77 P. Haberfield, A. Nudelman, A. Bloom, R. Bloom and H. Ginsberg,
J. Org. Chem., 1971, 36, 1792–1795.
78 P. Atkins, Physical Chemistry 6th Edition, W.H. Freeman and
Company, New York, 1998.
79 P. Bonhote, A.-P. Dias, N. Papageirgiou, K. Kalyanasundaram and
M. Gratzel, Inorg. Chem., 1996, 35, 1168–1178.
80 M. Ue, M. Takeda, T. Takahashi and M. Takehara, Electrochem.
Solid-State Lett., 2002, 5, A119–A121.
81 L. Leclercq, I. Suisse, G. Nowogrocki and F. Agbossou-Niedercorn,
Green Chem., 2007, 9, 1097–1103.
82 M. Yoshizawa, M. Hirao, K. Ito-Akita and H. Ohno, J. Mater.
Chem., 2001, 11, 1057–1062.
83 J. Schleicher, A. Ahosseini and A. M. Scurto, manuscript in prepara-
tion, 2008.
702 | Green Chem., 2009, 11, 694–703 This journal is © The Royal Society of Chemistry 2009
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ity
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n 
17
/1
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17
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8.
 
View Article Online
84 V. V. Zverlov, O. Berezina, G. A. Velikodvotskaya and
W. H. Schwarz, Appl. Microbiol. Biotechnol, 2006, 71, 587–
597.
85 S. N. V. K. Aki, A. M. Scurto and J. F. Brennecke, Ind. Eng. Chem.
Res., 2006, 45, 5574–5585.
86 A.M. Scurto, S. N. V. K. Aki and J. F. Brennecke, J. Am. Chem. Soc.,
2002, 124, 10276–10277.
87 A. M. Scurto, S. N. V. K. Aki and J. F. Brennecke, Chem. Commun.,
2003, 572–573.
88 A. M. Scurto, Ph.D. Dissertation, University of Notre Dame, 2002.
89 J. Schleicher, MS Thesis, University of Kansas, 2007.
90 O. Levenspiel, Chemical Reaction Engineering, Wiley, New York,
USA, 1998.
91 M. J. Kamlet, T. N. Hall, J. Boykin and R. W. Taft, J. Org. Chem.,
1979, 44, 2599–2604.
92 D. Lide, CRC Handbook of Chemistry and Physics 88th edn., CRC
Press, Boca Raton, 2007.
93 M.Temprado and J. S. Chickos,Thermochim. Acta, 2005, 435, 49–56.
This journal is © The Royal Society of Chemistry 2009 Green Chem., 2009, 11, 694–703 | 703
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