rXXXX American Chemical Society A dx.doi.org/10.1021/cr1003248 | Chem. Rev. XXXX, XXX, 000–000
REVIEW
pubs.acs.org/CR
Room-Temperature Ionic Liquids: Solvents for Synthesis and Catalysis. 2
Jason P. Hallett and Tom Welton*
Department of Chemistry, Imperial College London, South Kensington Campus, London SW7 2AZ, United Kingdom
bS Supporting Information
CONTENTS
1. Introduction A
2. Preparation C
3. Handling D
4. Physical Properties E
5. Solvent-Solute Interactions in Ionic Liquids F
5.1. Water F
5.2. Polarity H
5.2.1. Dielectric Constant H
5.2.2. Hildebrand Solubility Parameter, δH I
5.2.3. Single-Molecule Spectroscopic Probes I
5.2.4. Multiparameter Scales L
5.2.5. Antagonistic Behavior in Hydrogen
Bonding N
5.2.6. Apparent Discrepancies between Polarity
Measurements in Ionic Liquids N
5.3. Infinite Dilution Activity Coefficients O
5.3.1. Activity Coefficients and Solute-Ionic Li-
quid Interactions O
5.3.2. Modeling of γ¥ Values S
5.3.3. Summary S
6. Ionic Liquid Effects on Stoichiometric Chemical
Reactions T
6.1. Energy Transfer T
6.2. Electron Transfer Reactions T
6.3. Acid-Base Reactions U
6.4. Substitutions X
6.4.1. Type a SN2 Y
-
þ R—X Reaction X
6.4.2. Type b SN2 Y þ R—X Reaction Z
6.4.3. Type c SN2 Y þ [R—X]
þ Reaction AB
6.4.4. Type d SN2 Y
-
þ [R—X]þ Reaction AB
6.4.5. Type e SN1 R—X Reaction AC
6.4.6. Type f SN1 [R—X]
þ Reaction AC
6.4.7. Ionizing vs Dissociating Solvents AC
6.4.8. Nucleophilic Aromatic Substitutions AD
6.5. Elimination AD
6.6. Additions AE
6.6.1. Electrophilic Addition AE
6.6.2. Nucleophilic Addition AF
6.6.3. 1,4-Conjugate Additions AF
6.6.4. Diels-Alder Reactions AF
6.6.5. Oxidative Addition AH
6.6.6. Metal Alkyl and Aryl Addition Reactions AH
6.7. Acid-Catalyzed Reactions AI
6.7.1. Electrophilic Aromatic Substitutions AI
6.7.2. Isobutane Alkylation AL
6.7.3. Esterification AM
6.7.4. Diels-Alder Reactions AO
6.7.5. Dehydration AP
6.7.6. Beckmann Rearrangement AP
6.7.7. Ring Closing of Isonitrosoacetanilides AQ
6.7.8. Mannich Reaction AQ
6.7.9. Glycosylation AQ
6.7.10. Summary AQ
6.8. Base/Nucleophile-Catalyzed Reactions AQ
6.8.1. Aldol Reaction AQ
6.8.2. Morita-Baylis-Hillman Reaction AS
6.8.3. Knoevenagel Reaction AS
6.8.4. Keto-Enol Tautomerization AT
6.8.5. Nucleophilic Substitutions AU
6.8.6. Esterification/Acetylation of Alcohols AU
6.8.7. Acid Scavenging AV
6.8.8. Summary AV
7. Transition-Metal-Catalyzed Reactions AV
7.1. Ionic Liquids as Catalyst Immobilizers AW
7.2. Hydrogenations AW
7.3. Oxidations AZ
7.4. Carbonylations and Hydroformylations BB
7.5. Dimerizations BD
7.6. Palladium-Catalyzed Coupling Reactions BD
7.7. Olefin Metathesis BE
7.8. Summary BE
8. Overview BE
Author Information BE
Biographies BF
Acknowledgment BF
References BF
1. INTRODUCTION
In 1999 Chemical Reviews published its first review on the
subject of “Room-Temperature Ionic Liquids: Solvents for
Received: September 23, 2010

B dx.doi.org/10.1021/cr1003248 |Chem. Rev. XXXX, XXX, 000–000
Chemical Reviews
REVIEW
Synthesis and Catalysis”.1 Then the field was dominated by a
small number of specialist research groups, and it would have
been very unlikely to find a general synthetic chemist considering
whether an ionic liquid might be the best possible solvent in
which to try a reaction. At that time there were very few chemists
who had even heard of ionic liquids. Today, the same chemist is
likely to be well aware of the existence of ionic liquids and to
know that they are a useful addition to the range of solvents that
they might use. There has been an explosion of interest in ionic
liquids. There have been well over 6000 papers published in the
past 10 years with the phrase “ionic liquid(s)” in the title, and just
over half of these have been on the subject of chemical synthesis
and/or catalysis in one way or another. Whereas in 1999 it was
possible to be truly comprehensive with just 123 references, now
it is an impossible task. Therefore, it is important for us to first
define the scope of this review.
First, the question of the definition of an “ionic liquid” has been the
subject ofmuch debate over the past 10 years.Here we reproduce the
note on nomenclature that was used in the original review;1 we
believe that it is still the best approach to this thorny issue: “Room-
temperature ionic liquid, nonaqueous ionic liquid, molten salt, liquid
organic salt, and fused salt have all been used to describe salts in the
liquid phase. With the increase in electronic databases, the use of
keywords as search tools is becoming ever more important. While
authors are free to choose any name that they wish for their systems, I
would suggest that they at least include the term ionic liquid in
keyword lists. In this paper, I allow the term ionic liquid to imply that
the salt is low melting.”
There are a number of systems for which it is fair to say that
they are not composed solely of ions when pure that have been
described, such as “deep eutectic solvents” 2 and some protic
ionic liquids (depending upon the acid-base equilibrium
constant),3 including “Olah’s reagent”.4 These will be included
in this review when their use meets the criteria below.
This review is designed to update and complement the 1999
paper, with the primary literature surveyed up to the end of 2009,
and will follow most of the same format. However, this time we will
not separate the halogenoaluminate ionic liquids for special atten-
tion. Ionic liquids can be used inmanyways. The focus of this review
is their use as solvents in synthesis and catalysis. In 1999 it was
exciting that a small number of organic and catalytic reactions had
been conducted in ionic liquids. The primary conclusion that
emerges from there being such a vast literature of the use of ionic
liquids (ILs) in synthesis and catalysis is that, providing there is no
direct reaction between the ions of the IL and the starting materials,
most reactions will “go” in most ionic liquids. With the benefit of
hindsight, it now seems that we were naïve to have found this to be
anything other than what we should have expected. Consequently,
we will not review papers where this is the sole conclusion of the
work. Much of the interest in ionic liquids has centered on their
possible use as “green” alternatives to volatile organic solvents.5 This
claim usually rests on the fact that ionic liquids are generally
nonvolatile under ambient conditions. Hence, exposure risk to ionic
liquids ismuch lower than it is for a volatile solvent, and they have no
damaging atmospheric photochemistry. This nonvolatility also leads
to most ionic liquids being nonflammable under ambient
conditions.6 These properties were known at the time of the 1999
review and were used to justify interest in the use of ionic liquids
then. Hence, we will not include studies in which the advantage
claimed for the use of ionic liquids over other solvents rests on these
properties alone. There have beenmany reviews of ionic liquids that
can be referred to for information that is outside the scope of this
paper. Some of these have focused on particular applications, e.g.,
analysis,7 biocatalysis,8 electrochemical devices,9 or engineering
fluids.10 Others have concentrated on particular subgroups of ionic
liquids, e.g., protic ionic liquids3 or task-specific ionic liquids.11
Catalysis has been particularly well reviewed very recently.12 Two
special journal issues appeared in 2007 containing reviews covering a
wide range of subjects in ionic liquid research.13 Finally, three books
are available.14 There is also an online, free to users, ionic liquid
database.15 We will not seek to simply rereview these. We will
instead concentrate on those studies that have sought to understand
how ionic liquids can affect the reactivity of solute materials (e.g.,
yields, rates, and/or selectivities).
As the number of ionic liquids and the number of research groups
working in ionic liquids has proliferated the number of different
abbreviations for ionic liquid, ions have done likewise. This has been
particularly so for ionic liquid cations. Most abbreviations used in the
literature derive from the name of the ion; e.g., [bmim]þ and
[BMIM]þ both refer to the 1-butyl-3-methylimidazolium cation.
However, ambiguities rapidly arise; e.g., should [pmim]þ refer to
the 1-propyl-3-methylimidazolium cation or to the 1-pentyl-3-methy-
limidazolium cation? To circumvent these problems, we have chosen
to use an alphanumeric system to describe the alkyl chains, with an
alphabetic abbreviation for the charged center (Figure 1). In this
system the 1-butyl-3-methylimidazolium cation becomes [C4C1im]
þ.
If the alkyl chain is not linear, this can be noted; e.g., the 1-tert-butyl-3-
methylimidazolium cation becomes [tC4C1im]
þ. If the side chain is
functionalized, then the type and position of the functional group are
noted; e.g., [(HO)4C4C1im]
þ denotes that there is an alcohol group
on the terminal carbon of the butyl chain. In-chain inclusion of non-
carbon atoms can also be shown; e.g., [(C1OC2)C1im]
þ shows the
Figure 1. Some commonly used cations for ionic liquids.

C dx.doi.org/10.1021/cr1003248 |Chem. Rev. XXXX, XXX, 000–000
Chemical Reviews
REVIEW
presence of an ethereal oxygen in the chain between the second and
terminal carbon atoms.Alkyl chains are assumed tobe saturated unless
noted otherwise; e.g., [(C1dC2)C1im]
þ represents 1-allyl-3-methyl-
imidazolium cation. For cyclic cations, alkylation is assumed to be on
the heteroatom(s) unless otherwise indicated by giving the numerical
position of the alkylation on the ring such that 1-butyl-2,3-dimethy-
limidazoliumbecomes [C4C1C1
2im]þ.Where alkyl chains arepresent
in anions, the same system will be used (Figure 2); e.g., [C4OSO3]
-
denotes butyl sulfate. No system is perfect, but this one gives the least
ambiguity and is the most systematic of those currently in use.
2. PREPARATION
The past 10 years has seen a proliferation of the number of
ionic liquids that have been prepared. However, it remains true
that the vast majority of ionic liquid cations are based upon
alkylated amines, with a smaller number of phosphonium and
sulfonium salts used. Consequently, although details vary, there
are common themes that reappear in the synthesis of ionic
liquids: first alkylation of an amine/phosphine/sulfide to give an
intermediate salt, followed by anion exchange to give the ionic
liquid. Most ionic liquids cannot be distilled in large quantities.
Hence, the ability of synthetic protocols to give high-purity ionic
liquids has been an important focus for their selection.16
Impurities in ionic liquids can come fromanumberof sources. The
first is that they are contained in the startingmaterial. Consequently, it
is necessary to purify these before use, typically by distillation from a
suitable drying agent in an inert atmosphere. However, when very
pure “spectroscopic grade” ionic liquids have been required, prewash-
ing of the starting materials has also been used.17,18
Once the startingmaterials have been prepared, the alkylation step
can then be performed. Throughout this step air andmoisture should
be rigorously excluded. This nucleophilic substitution is highly
exothermic, and this can lead to runaway reactions.19,20 Overheating
of the reaction solution also leads to greater formation of elimination
products from the competing reaction (Scheme 1). Hence, the
reaction is usually performed in solution, the reactants are usually
added slowly to prevent hot spots, and there has been a trend in
academic laboratories over the years to reduce the temperature and to
extend the time of the reaction. A kinetic investigation of the reactions
of 1-bromohexane with 1-methylimidazole in 10 different solvents,
followed by a linear solvation energy relationship (LSER) analysis,
showed that the rate was greatest in dipolar non-hydrogen bond
donor solvents.21 The corollary of this is, of course, that the same rate
of reaction can be achieved at lower temperature.
Microwave heating has also been used for these syntheses.22
Although this allows reactions to be conducted very quickly, the
difficulty in controlling the reaction conditions, especially the
generation of hot spots, leads to variable quality of the products.23
Ultrasound has also been used at this stage of the synthesis, with
excellent yields being reached more quickly and at lower tempera-
tures than with conventional heating and stirring.24 Here, the
authors also reported that the product salts were purer than those
made by the conventional method. These results probably arise
from themore efficientmixing that is achievedwith ultrasound, both
leading to faster reaction and preventing the formation of hot spots,
particularly later in the reaction when the mixture becomes more
viscous. [C4C1im]Br has been prepared at a rate of up to 9.3 kg/day
in a continuous-flow microreactor, which was shown to be an
appreciable improvement over a batch system.19
Most of the halide salts of cations used in ionic liquids are solid
at room temperature, allowing purification of the intermediate
salts by recrystallization, most often from acetonitrile; ethyl
acetate can be added to aid precipitation of the salt and removal
of unreacted starting materials. Finally the salt is dried in vacuo.
The ionic liquid can then be prepared by metathesis of the
halide salt with ametal or ammonium salt or the conjugate acid of
the required anion (Scheme 2). For hydrophobic ionic liquids
this can be done in aqueous solution,25 with the product
separating during the reaction. The aqueous solution can then
be washed with CH2Cl2 and the ionic liquid isolated from the
CH2Cl2 solution to increase the overall yield. Of course, the ionic
liquid is not dry at this stage, and any remaining water is removed
Scheme 1
Figure 2. Some commonly used anions for ionic liquids.

D dx.doi.org/10.1021/cr1003248 |Chem. Rev. XXXX, XXX, 000–000
Chemical Reviews
REVIEW
in vacuo. This is usually conducted at elevated temperatures, but
care is required because this can lead to decomposition of the
ionic liquid.6 For hydrophilic ionic liquids the metathesis is
usually performed in a water-immiscible organic solvent.26 The
resulting mixture is then filtered and the filtrate washed with
water to remove any residual halide salt. The greater the
miscibility of the product ionic liquid with water, the less effective
this process is, leading to either low yields or halide contamina-
tion of the ionic liquid.27 The application of ultrasound for the
metathesis reaction of [NH4][PF6] or [NH4][BF4] with
[C4C1im]Cl in acetone has been shown to lead to less colored
ionic liquids in shorter time than the same reaction with stirring
(500 rpm at room temperature).28 The organic solvent is then
removed in vacuo. If all of the steps in the synthesis have been
performed well, the ionic liquid is then ready for drying and use.
However, the ionic liquids are often colored at this stage or
contain other impurities, such as unreacted starting materials.
These are typically removed by the action of sorbents such as
activated charcoal, alumina (usually neutral or acidic), or silica.29
However, concern has recently been expressed that this can leave
small amounts of the sorbent material in the ionic liquid.30
The general preparations of ionic liquid as described above form
the vast majority of the syntheses described in the literature. The
details of the syntheses vary between ionic liquids and research
groups and over time and are not always reported in detail. With the
exception of the points noted above, it is not at all clear which of the
variations reported do provide a synthetic advantage and which are
merely the consequences of local custom. For instance, when
working on chloroaluminate ionic liquids in the 1980s, one of us
noticed that the use of acetone in the cleaning of glassware gave
colored ionic liquids, and we have never used it since, preferring
instead to use ethanol, regardless of which ionic liquids we are
working with at the time.
The metathesis step and subsequent washings and separations
cannot always be performed to give an ionic liquid of the required
standard, particularly for ionic liquids with more basic anions, so a
number of techniques have arisen to avoid this. When the hydroxide
salt is available, as with [(C4)4P]OH, the ionic liquid can be prepared
by direct neutralizationof an acid.31When it is not available, a solution
of it can be generated using an ion exchange column. This has been
used for the synthesis of 20 ionic liquids prepared from amino acids.32
An aqueous solution of [C2C1im]Cl was passed down a column of
Amberlite IRA400OH resin to generate a dilute aqueous solution of
[C2C1im]OH, which was then neutralized with an aqueous solution
of the amino acid. The water was then removed by freeze-drying.
Aqueous solutions of [C2C1im]OH have also been prepared by
membrane electrodialysis.33,34
Direct alkylation of an amine/phosphine can also be used to
avoid the need for a halide intermediate. For instance, methyl triflate
has been reacted with a stoichiometric amount of the 1-alkylimida-
zole in 1,1,1-trichloromethane to yield the [CnC1im][OTf] ionic
liquid.35 This is an exothermic reaction, and methyl triflate is
sensitive to moisture, so the reaction must be carried out under
anhydrous conditions with cooling of the reaction mixture. Other
sulfonate ionic liquids have been prepared by the direct reaction of
alkanesulfonate esters with substituted imidazoles to yield
[CnC1im][CmSO3] (n = 1 or 4; m = 1, 2, or 4).
36 Several of these
were solid at room temperature and could be recrystallized from
acetone. The authors then used the [C4C1im][C1SO3] salt to
prepare [C4C1im]X {where X = [BF4], [PF6], [PF3(CF2CF3)3],
[OTf], or [NTf2]} by metathesis in water followed by extraction
with CH2Cl2 to give Cl
--free ionic liquids.
The first alkyl sulfate ionic liquid, [C4C1im][C8OSO3], was
prepared by metathesis of [C4C1im]Cl and Na[C8OSO3] to give a
less toxic and “greener” ionic liquid than those available at the time.37
However, dialkyl sulfates and trialkyl phosphates can be used as direct
alkylating agents to give alkyl sulfate or dialkyl phosphate salts,
respectively.38-40 Again care is required when using these direct
alkylation routes to avoid overheating of the reaction mixture and
the formation of colored or otherwise impure ionic liquids.
[C2C1im][CmOSO3] (m = 1 or 2) have been converted into salts
with longer alkyl chains on the sulfate anions by transesterificationwith
a suitable alcohol.41 However, this requires the presence of an acid
catalyst that can be difficult to remove entirely from the ionic liquid.
In an attempt to alkylate 1-methylimidazole with dimethyl
carbonate to give [C1C1im][C1CO3], Rogers et al. formed the
zwitterionic 1,2-dimethylimidazolium 2-carboxylate.42 Reaction of
this salt with strong acid can lead to decarboxylation and formation
of the dimethylimidazolium salt of the conjugate base of the acid.43
However, the selectivity problems that arise with this route are
significant and difficult to control. Attempts are continuing to
alleviate these problems, such as using this route to first generate
[C1C1im][HCO3], which itself can be used as an intermediate for
the formation of other salts.44 Although promising, this route
requires more development before it can be widely applied to the
synthesis of a variety of 1,3-dialkylimidazolium salts.
The avoidance of impurities arising from themetathesis reactions
has also been attempted by the synthesis of volatile intermediates.
1,3-Dialkylimidazolium salts have been treated with strong base to
give the 1,3-dialkylimidazolylidenes, which were then distilled and
added to acid HX to give [CnCmim]X ionic liquids.
45 This
methodology was subsequently extended to 1,2,3-trialkylimidazo-
lium and trialkylmethylphosphonium salts.46 While this technique
does avoid halide impurities, it is difficult to perform the addition of
the acid precisely and acid impurities can remain in the final ionic
liquid.
The importance of Brønsted acids as catalysts in chemical
synthesis (see below) led to the preparation of ionic liquids with
the cation bearing a -SO3H functional group as the first of the so-
called “task-specific ionic liquids”.47 The reactions of N-butylimida-
zole or triphenylphosphine with 1,4-butane or 1,3-propane sultone
led to the zwitterions [(-O3S)
nCnC1im
þ] or [(-O3S)
nCnPPh3
þ]
(n = 3 or 4), respectively.48 These are then followed by the addition
of a strong acid, which protonates the sulfonate group to yield the
product ionic liquid. Another group of Brønsted acidic ionic liquids
are the protic ionic liquids.3 These are prepared by simple neu-
tralization of the appropriate base with a suitably strong acid and
subsequent drying.49,50 Of course, BASF’s BASIL (biphasic acid
scavenging utilizing ionic liquids) process involves the formation of
a protic ionic liquid during a proton abstraction reaction.51
3. HANDLING
When considering how a material should be handled, two
aspects need to be considered, these being possible detrimental
effects on thematerial itself and hazards that thematerial poses to
those exposed to it.
Scheme 2

E dx.doi.org/10.1021/cr1003248 |Chem. Rev. XXXX, XXX, 000–000
Chemical Reviews
REVIEW
It can fairly be said that it was the publication of a paper titled
“Air and water stable 1-ethyl-3-methylimidazolium based ionic
liquids” that paved the way for the widespread application of
ionic liquids in synthesis and catalysis.25 However, it also led
many to believe that there are no problems associated with the
use of these ionic liquids in the air. Of course, the likely reactivity
of an ionic liquid with the air, principally oxygen and water, is a
chemical phenomenon and depends upon the identity of the
component ions of the ionic liquid. There is no evidence that
being a component of an ionic liquid makes any given ion more
reactive to oxygen or water, but there is equally no evidence that
it makes these ions any less reactive. Unfortunately, some ionic
liquids, such as those containing the ions [PF6]
-, [SbF6]
-,
[BF4]
-, and [CnSO4]
-, which were originally described as air
and moisture stable, undergo hydrolysis.52 In many cases the
toxic and corrosive HF is formed.
Even when the ionic liquid components do not react with
moisture, failure to protect the ionic liquids from the atmosphere
will lead to considerable concentrations of water. The hydrophilicity
of the ionic liquids is largely determined by the nature of the anion of
the ionic liquid; themore basic the anion, the greater the solubility of
water, with a secondary cation effect of shorter alkyl chains or
substitution with hydrophilic functional groups leading to greater
water solubility.26,53 However, even the ionic liquids with the least
miscibility with water will contain considerable concentrations if not
handled under anhydrous conditions,54 and this can lead to dramatic
changes in both physical and chemical properties of the solution.55
Oxygen has proven to be far less problematic. This is because ionic
liquids are rarely made using easily oxidized ions, the solubility of O2
in most ionic liquids is relatively low,56 and it is easily removed by
applying even a moderate vacuum.
Of course, although it is important for anyone working with ionic
liquids to know how to protect these during handling, it is more
important that they are aware of any risks to themselves from the
ionic liquids. Alongside the “green” label for ionic liquids came a
widely held misconception that all ionic liquids were nontoxic. It is
only recently that the study of the toxicity of ionic liquids has come to
the fore, and most of this has been applied to gaining an under-
standing of (eco)toxicity, rather than being directly about risk
assessment for those working with ionic liquids.57 Furthermore, no
single ionic liquid has been studied so completely that it can be
claimed to be fully understood with respect to its toxicity. Although
several studies have shown a link between the length of alkyl chains
on the cations and toxicity, none have used enough ionic liquids to,
by themselves, be sure of all structure-toxicity relationships.
Attempts have been made to draw results from several studies
together to derive sufficiently robust quantitative structure-property
relationships (QSPRs) to give some predictive capacity for ionic
liquid toxicity. Luis et al.58 used their QSPR to develop a group
contributionmethod for estimating the toxicity of ionic liquids to the
aquatic organism Vibrio fischeri. They showed that the contribution
to an ionic liquid’s toxicity increased in the order pyrrolidinium <
imidazolium < pyridinium and with increasing length of the alkyl
substituents. A subsequent study of V. fischeri and Daphnia magna
showed up to four important factors.59 The effect of changing the
anion of the ionic liquid was pronounced, but not systematic. For the
cations, the presence of an aromatic ring and then the number of
heteroatoms in the ring increased the toxicity of the ionic liquid,
which is a result different from that of the earlier study. It was shown
again that increasing the alkyl chain on the cation increased the
toxicity of the ionic liquid, but furthermethylationof the ring reduced
its toxicity. The link between toxicity and alkyl chain length was also
observed in another QSPR study of the human Caco-2 cell line.60
This cell line is interesting in that it is a human epithelial [derived
from a colon (large intestine) carcinoma] cell line that has been
proposed to be a good model for human toxicology studies.
The increasing toxicity of ionic liquids with the increased
length of alkyl substituents on cations has been explicitly, if
imperfectly, related to the lipophilicity of the resultant ionic
liquid, as described either by the octanol/water partition
coefficient61 or by the retention behavior in HPLC.62 Ionic
liquids with greater lipophilicity are more able to interfere with
biological membranes, which can in turn lead either directly to
cell breakdown or to the passage of toxic substances to the cell
interior and hence greater toxicity.
The claim that ionic liquids are safer solvents has never really
rested on a belief that they are composed of nontoxic ions. Rather it
has been based on the fact that their insignificant vapor pressures
mean that accidental exposure to ionic liquids is far less likely than for
volatile organic carbons. Transdermal exposure remains a risk.
[C4C1im]Cl has been found to be acutely orally toxic in female
Fischer 344 rats, with an estimatedLD50 of 550mg/kg.
63 In the same
study, the transdermal toxicity was shown to be more complex.
When administered as a concentrated aqueous solution, there were
no signs of gross toxicity (LD50 > 2000 mg/kg), whereas when
administered in a solution in N,N-dimethylformamide (DMF), an
LD50 of between 800 and 2000 mg/kg was found in the female rats,
but the LD50 was greater than 2000 mg/kg in male rats. Clearly the
ability of the salt to cross the skin is an important component of its
transdermal toxicity, and this can be increased by the presence of
cosolvents. It must be remembered that [C4C1im]Cl is a solid at
room temperature and could not be administered as a pure liquid.
Furthermore, [C4C1im]Cl is highly hydroscopic and is difficult to
extract from aqueous solution. It has been proposed that suitably
lipophilic ions can be used to deliver pharmaceutical counterions
transdermally.64 This suggests that ionic liquids that are liquid when
in contact with skin and particularly those that are lipophilic aremore
likely to lead to transdermal toxicity.
4. PHYSICAL PROPERTIES
There are many liquid properties that are important to the
performance of a reaction solvent; these include the liquid range,
heat capacities, viscosities, etc. These properties for ionic liquids
have been very well reviewed elsewhere and are not detailed
here.13-15,65 These properties are controlled by the selection of
both the cation and anion. This has led to the concept of ionic
liquids being “designer solvents”.66 However, for this to be
achieved requires not just a post hoc rationalization of the ionic
liquids’ properties, but the ability to predict them.
Most people who have worked with ionic liquids for any length of
time will have experienced the frustration of making a salt that seems
to be little different from a known room-temperature ionic liquid only
to discover that it has a melting point much higher than expected.
Consequently, the search for the ability to predict the melting points
of organic salts has become an important activity. The QSPR
approach has had some limited success in this respect. The first
attempt at this was to correlate and predict the melting points of
organic salts based on the quaternary ammonium cation.67 Moderate
correlations were found for a set of 75 tetraalkylammoniumbromides
and for a set of 34 (hydroxyalkyl)trialkylammonium bromides.
Descriptors used in the correlations were analyzed to determine
which structural features led to lowermelting points. These were that
asymmetry, due to one or two moderately long chains (e.g., octyl)

F dx.doi.org/10.1021/cr1003248 |Chem. Rev. XXXX, XXX, 000–000
Chemical Reviews
REVIEW
with two or three shorter chains (e.g., butyl), gives lower melting
points; branching on the longer chains was also predicted to lead to
lower melting points, as long as two or more bonds separate the
branch points. Since only bromide salts were studied, nothing could
be said about the effect of structural change on the anion. However, it
was noted that more complex anions would lead to lower melting
points.
In a more wide ranging study,68 QSPR modeling of the melting
point of a total of 717 nitrogen-containing bromide salts was
performed using a variety of different techniques and descriptors.
It was noted that although the neural network methods used were
marginally superior, the differences between the predictive perfor-
mances of the various methods were small. However, it was
concluded that the deviation of predicted and calculated results
was sufficiently large that the models could only give general trends
rather than accurate predictions of the melting point of a candidate
ionic liquid. Much of the problem with generating these QSPRs is
with the experimental data. There is the well-known problem that
many ionic liquids are glass-forming materials. Consequently, their
melting points are difficult to determine, can be dependent upon
heating/cooling rates, and give a very broad range of results. Other
possible effects arise from many salts having more than one
polymorph,69 each with different melting points. Also, if more than
one polymorph is present in the solid state, this will lead to a lower
melting point eutectic. Consequently, achieving the necessary level of
control over the experimental conditions to achieve more precise
melting point data is unrealistic, and it is unlikely that these data, or
precise predictions based upon these data, will be generally available
for a wide range of salts. However, notwithstanding the complexities
described above, it has been possible to generate some reasonable
predictions for small, closely related groups of ionic liquids using
parameters such as calculated molecular volumes and/or quantum
mechanical calculations.70,71
The upper operating limit of an ionic liquid is given by its thermal
decomposition. As a chemical rather than physical phenomenon,
decomposition temperatures for ionic liquids are expected to be
difficult to predict. This is particularly true given that the kinetics of
decomposition reactions need to be considered as well as the
thermodynamics. Angell et al.50,72 found that the temperature at
which the total vapor pressure of a protic ionic liquid reaches 1 atm
(boiling) can be predicted by the difference in the pKa values, in
dilute aqueous solution, of the acid and base from which the ionic
liquid was made (ΔpKa). The larger the ΔpKa, the higher the
boiling temperature. The onset of boiling was also shown to be over
a small temperature range and so can be used as an estimation of the
upper limit of the useful liquid range of the protic ionic liquids. This
concept has subsequently been used to design some highly ther-
mally stable protic ionic liquids.73
An attempt to predict the upper end of the liquid range and the
thermal decomposition of other ionic liquids has also been made.74
Quantum chemical calculations were used to calculate activation
energies to predict decomposition mechanisms and rates. It was
found that ionic liquids with sufficiently nucleophilic anions decom-
posed primarily though SN2 dealkylation of the cation by the anion
and that rate constants correlated well with experimental decom-
position temperatures. Interestingly, the predicted rate constants
gave decomposition temperatures approximately 100 C below the
commonly quoted experimental values, which is in good agreement
with studies that have looked at long-term stabilities of ionic liquids.6
Ionic liquids with the [NTf2]
- anion were found to decompose via
the elimination of SO2 from the anion, again in agreement with
experimental data.75
Molecular volume data have been used to predict a number of
physical properties of ionic liquids, namely, densities,76-78 sur-
face tensions,76,79 and viscosities and conductivities.77 Given the
potential importance of molecular volume data for predicting
physical properties of ionic liquids, it is useful that they have also
been the subject of prediction using a variety of methods.71,80
5. SOLVENT-SOLUTE INTERACTIONS IN IONIC LIQUIDS
5.1. Water
The ubiquitous presence of water in even themost well controlled
systems singles it out for special consideration as a solute species. At
the time of the 1999 review, interest was concentrated on water’s
reactions with chloroaluminate species and the consequent protic
species present in these ionic liquids and their reactivity and super-
acidity. Since then, concern has moved to the effects that even small
amounts ofwater canhaveon theproperties of thewide rangeof ionic
liquids that do not react with water, such as electrical conductivity,81
viscosity,55 surface tension,82 and how the ionic liquid interacts with
other solutes,83 even for ionic liquids often described as hydrophobic.
In fact, all ionic liquids that have been described to date are
hygroscopic and will absorb water from their surroundings.
Early investigations of watermiscibility found that the ionic liquids
[CnC1im][PF6] (n= 4, 6, or 8) formed biphasicmixtures with water
whereas [CnC1im]Cl (n = 4, 6, or 8) did not.
84 Furthermore, it was
shown that the miscibility of the [CnC1im][PF6] ionic liquids with
water decreased as the alkyl chain length increased. The reactivity of
the [PF6]
- ion with water was not considered. [CnC1im][BF4] (n =
2-5) has been found to be fully miscible with water at room
temperature, whereas for [CnC1im][BF4] (n = 6-10) two layers
were formed.27 The temperature-dependent behavior of this phe-
nomenon has been used to provide temperature-reversible ionic
liquid-water two-phase/one-phase systems for catalysis.85 The
miscibility of [PF6]
- ionic liquids with water has been increased
by preparing imidazolium cations with ether-functionalized alkyl
chains.86 For [C4C1im]
þ ionic liquids the macroscopic phase
behavior of ionic liquid-water mixtures has been demonstrated to
vary with the anion such that Cl-, Br-, [OTf]-, and [BF4]
- give
ionic liquids that mix with water in all compositions, whereas
[C(CN)3]
-, [C1OC2OC2OSO3]
-, [PF6]
-, and [NTf2]
- will
phase separate.53 It has also been shown that by using an anion with
a suitably long perfluorinated alkyl chain, e.g., nonafluorobutanesul-
fonate, [FC4SO3]
-, can give an ionic liquid that forms two layerswith
water, even with protic cations, e.g., [C3C2C1NH]
þ.87 As well as
phase separation, complex phase behaviors, such as micelle88 and
gel89 formation, have been observed for ionic liquid-watermixtures.
Octanol-water partition coefficients, KOW, offer a methodology
to quantify the hydrophobicity of a compound. These have been
measured for a number of imidazolium ionic liquids.90 For anions in
[C4C1im]
þ ionic liquids, KOW was found to increase in the order
[OTf]- < [BF4]
- < Br- < [NO3]
- < Cl- < [PF6]
- < [NTf2]
-,
suggesting that hydrogen bonding to the anion is a significant
contribution to the hydrophilicity of the ionic liquid. For [NTf2]
-
ionic liquids, KOW was found to increase with the alkyl chain length,
but not with the number of substitutions, of the di- or trisubstituted
imidazolium cations. These observations led to the conclusion that
hydrogen bonding from the C2-Hof the cation was not a significant
contribution to the hydrophilicity of the ionic liquid. This is in direct
contrast to the result of a calorimetric study using 1-propanol to study
[C4C1C1
2im]Cl and [C2C1im]Cl in water.
91 However, it should be
noted that these experiments were conducted with very different
concentration ranges of ionic liquid in the water.

G dx.doi.org/10.1021/cr1003248 |Chem. Rev. XXXX, XXX, 000–000
Chemical Reviews
REVIEW
A puzzling set of results regarding the lack of reactivity of water in
some ionic liquid processes have been reported. This is in spite of the
fact that hydrolysis reactions in ionic liquids have also been clearly
demonstrated.92 Water has even been shown to have an increased
nucleophilicity in some reactions (see below).93 This lack of expected
reactivity of water in ionic liquids was first observed in studies of the
oxidation of alcohols using ruthenium94 or palladium95 catalysts. It is
well-known that the presence ofwater often leads to significant further
oxidation of the product aldehydes to acids when primary alcohols are
used. This is usually dealt with by the addition of molecular sieves to
the reaction mixture. However, in the hydrophilic ionic liquid
[C4C1im][BF4], this was not necessary, and even upon the addition
of water, no overoxidation was seen until the amount of water present
exceeded that of ionic liquid.94The same lack of overoxidationhas also
been seen in the TEMPO-CuCl-catalyzed (TEMPO = 2,2,6,6-
tetramethylpiperidine-1-oxyl) oxidation of alcohols to aldehydes.96
PCl3 and POCl3 have been found to be significantly more hydro-
lytically stable in the ionic liquids [C4C1im][NTf2], [C4C1im]
[BF4], [C4C1pyrr][NTf2], and [C2C1im][OTf] than in molecular
solvents.97 In the [NTf2]
- ionic liquids PCl3 could be stored for days
in contact with the air without any hydrolysis. In a recent study of the
precipitation of the antimalarial compound artemisinin from ionic
liquids by the addition of water as an antisolvent, it was found that the
cybotactic region around the arteminisin was more effectively dehy-
drated by a hydrophilic ionic liquid, [(HO)2C2(C1)2NH][C2CO2],
than the hydrophobic [(HO)2C2(C1)3N][NTf2].
98 It is known that
ionic liquids can interact stronglywithwater and that it is possible that
hydrated and anhydrous domains can exist in these mixtures (see
below), but it is not at all clear how these domains contribute to the
effects on the chemistry of water dissolved in ionic liquids. Ionic
liquids have even been used as drying agents, such as in the
dehydration and breaking of ethanol-water or tetrahydrofuan
(THF)-water azeotropes in extractive distillation.99
In an attempt to understand the molecular basis for these
phenomena, the IR spectra of water in a range of ionic liquids,
[C4C1im]X {X = [PF6], [SbF6], [BF4], [ClO4], [OTf], [NTf2],
[CF3CO2], or [NO3]}, were investigated.
26 The spectra showed that
the water dissolved in the [C4C1im]X {X = [PF6], [SbF6], [BF4],
[ClO4], [OTf], or [NTf2]} ionic liquids was in a symmetric 2:1
complexboundtotwoanionsviaH-bonding(Figure3).In[C4C1im]X
{X = [CF3CO2] or [NO3]} the spectra of the water were found to
indicatetheexistenceofwateraggregatesorevendroplets.Thestrength
of the hydrogen bonding from the water to the anion was found to
increase in the order [PF6]
- < [SbF6]
- < [BF4]
- < [ClO4]
- <
[OTf]- < [NTf2]
- < [CF3CO2]
- < [NO3]
-, which agrees well
with other measures of the hydrogen bond acceptor abilities of
these ions (see below). This was found to also correlate with the
concentrations of water in the ionic liquids in equilibriumwith the air.
In no case in this study was there spectroscopic evidence for an
interaction between the water and the cation of the ionic liquid.
Similar results have been found in other IR studies.100 In fact,
this is so well accepted that the IR spectra of dissolved D2O have
even been suggested as polarity probes for ionic liquids.101
Variable-concentration experiments in [C4C1im][BF4] have
shown that, at higher concentrations of water than used in the
initial study above, self-association of the water occurs in an
anion-water-water-anion chain (Figure 3),102 with the appli-
cation of pressure to some degree reversing this.103
A recent study of the IR spectra of D2O in [C2C1im][BF4]
focused more on the changes occurring in the interactions
between the ionic liquid’s ions.104 At low concentrations the
water begins to break down the three-dimensional structure of
the ionic liquid, which then goes on to form ionic clusters as the
concentration of water increases until eventually ion pairs form,
which are the dominant species in the aqueous solution (10 mol
% ionic liquid). It would seem that the breakup of the ionic liquid
into clusters and the self-association of water molecules are
occurring in similar concentration ranges. This and other 2D
IR experiments have also indicated that some hydrogen bonding
between the imidazolium ring protons and the water can also
occur, also at higher concentrations of water than used in the
original study.106,105
1H NMR spectra of various concentrations of water in
[C4C1im][BF4] have also been studied.
106 Through-space cou-
pling experiments showed that water comes between the anions
and the imidazolium ring of the cations, thus loosening the
former tight structure and leading to a greater distance between
neighboring imidazolium rings. At the same time the formation
of hydrophobic domains, where the butyl groups of neighboring
cations become closer, was observed. As the concentration of
water was increased, distinct interactions between the cation
protons and the water could be seen, indicating the formation of
hydrogen bonds between them.
The crystal structure of the hemihydrate [C10C3C1C1N]
Br
3
0.5H2O contains isolated water molecules that are hydrogen
bonded to two bromide ions (Br----HOH---Br-) in the same
way as proposed from the IR studies.107 In the monohydrate
[C10C3C1C1N]Br
3
H2O, containing twice as much water, cyclic
[(H2O)Br]2
2- dimers (Figure 3) are seen.108 In the crystal struc-
tures of [C18C1im]Cl
3
H2O, [C14C1im]Cl
3
H2O (polymorph 1),
and [C18C1C1
2im]Cl
3
H2O similar [(H2O)Cl]2
2- dimers are
seen, but in these structures the water is also hydrogen bonded to
the imidazolium ring protons via the water oxygen atoms.109
However, a second polymorph of [C14C1im]Cl
3
H2O contained
the water hydrogen bonded to the Cl- ions in an infinite chain,
again with imidazolium ring CH---OH2 hydrogen bonds.
Molecular dynamics (MD) simulations of water in [C1C1im]X
{X = Cl or [PF6]} showed remarkably little difference in their
microscopic properties, in spite of the macroscopic difference of
[C1C1im]Cl being miscible with water at all concentrations and
[C1C1im][PF6] forming abiphasicmixture.
110 Itwas shown from the
radial and spatial distribution functions around the water molecules
that themain interactions with the solvent were twoH-bonds formed
to two anions. The nearest cation was also close by, but its precise
position resulted from the proximity of the water to the anions with
which the cation was interacting rather than a direct interaction
between the cation and the water. As the concentration of the water
was increased, self-association of the water began to be seen, with
water clusters and a percolating network of water molecules forming.
Subsequent simulations ofwater-[CnC1im]X {n=2or 4;X= [BF4]
or [NTf2]} mixtures have led to similar conclusions,
111 although a
Figure 3. Some observed ionic liquid-water hydrogen-bonded struc-
tures: symmetric 2:1 anion-water, anion-water-water-anion chain,
and anion-water-anion-water cyclic dimer.

H dx.doi.org/10.1021/cr1003248 |Chem. Rev. XXXX, XXX, 000–000
Chemical Reviews
REVIEW
more recent study has reported that a cation-water hydrogen bond
can be found in water-[C4C1im][BF4] mixtures.
112 This study also
concluded that the higher concentration clusters of water were
composed of linear chains of water molecules. Finally, quantum
mechanical (QM) calculations have shown that both cation-water
and anion-water hydrogen bonds can be important in determining
the structures found in [C2C1im]X {X=Cl, Br, [BF4], or [PF6]}.
113
A Car-Parrinello QM-MD simulation of a single [C2C1im]Cl ion
pair dissolved in 60 water molecules found that the ions remain
somewhat associated, with theCl- ion lying in the vicinity of theC2-
H of the imidazolium ring, even in this relatively dilute system.114
Combining all of the above, a picture in which at low concentra-
tions water is dispersed throughout the ionic liquid as isolated
molecules hydrogen bonded to the ionic liquids’ anionswith aweaker
interaction with the cations emerges. This leads to the beginnings of
the separation of the oppositely charged ions. As the concentration of
water increases in water-miscible ionic liquids, it can begin to self-
associate in dimeric structures. For ionic liquids that do not mix
fully with water, phase separation occurs at or before this stage. As
the concentration of water increases still further, these dimeric
structures begin to form chains of molecules, which can percolate
through the ionic liquids’ structures and cause them to break up into
first large, but then smaller, ionic clusters. As the amount of water
increases still further, the ionic liquids dissolve to become ion pairs in
aqueous solution and eventually free hydrated ions. The precise
concentrations at which the different behaviors emerge are very
dependent on the anion and to a lesser extent the cation of the ionic
liquid.
5.2. Polarity
Mostmoderndiscussions of solvents rely on the concept of solvent
polarity. Qualitative ideas of polarity are based upon observations
such as “like dissolves like” and arewell accepted andunderstood.The
currently accepted definition of polarity is that it is the sum of all
possible (specific and nonspecific) intermolecular interactions be-
tween the solvent and any potential solute, excluding those interac-
tions leading to definite chemical changes (reactions) of the
solute.115,116 This includesCoulombic interactions, the various dipole
interactions, hydrogen-bonding interactions, and electron pair
acceptor-electron pair donor acid-base interactions and is both
a physical and a chemical phenomenon. Hence, there can be no
simple single measure of all of these interactions. There are many
empirical solvent polarity scales117 that attempt to give quantita-
tive estimates of solvent polarity, several of which have been
applied to ionic liquids. None of these scales are perfect, and all
are more or less sensitive to different aspects of polarity. Care is
therefore required, since attempts to compare polaritiesmeasured
by different techniques that are sensitive to different properties of
the solvent often only lead to confusion. The test of a solvent
polarity scale is its usefulness in explaining and/or predicting
changes in solute properties when dissolved in different solvents.
There is no useful concept of “right” or “wrong”when comparing
these scales; rather the idea that a particular scale is more or less
appropriate in a given circumstance ismore helpful. However, it is
possible to draw some tentative general conclusions.
All ionic liquids are not the same; different combinations of
anions and cations lead to solvents with different polarities. No ionic
liquids have shown themselves to be “superpolar”; regardless of the
method of assessing their polarities, values for ionic liquids come
within the range of molecular solvents. Most general measures of
overall polarity place ionic liquids in the range of the short to
medium alkyl chain length alcohols.
5.2.1. Dielectric Constant. The dielectric constant (εr) of the
pure liquid is by far themost commonly used polarity scale. Dielectric
constants have been measured for most molecular liquids and are
widely available in reference texts, including most commercial
catalogues. Molecular solvents with εr < 9 are considered to be
nonpolar, those with 9 < εr < 15 are moderately polar, and values in
the range 15 < εr < 30 characterize the solvent as polar, but not
exceptionally so. Liquids with εr in excess of 50 are considered to be
highly polar. The ubiquity of εr leads many to treat it as if it is
synonymous with polarity. It is not; it is a physical phenomenon that
is used as an empirical polarity scale. Estimates of static dielectric
constants of ionic liquids have beenderived frommicrowavedielectric
spectroscopy (e.g., Table 1).118 Most values found ranged from 9 to
15 depending upon the ionic liquid, characterizing these as moder-
ately polar liquids. However, values in excess of 30 were found for
some protic ionic liquids.118f The dielectric constant of the ionic
liquids appears to be related to the ability of the ionic liquids to enter
into hydrogen bond networks. Hence, for anions the dielectric
constant varies such that [HCO2]
- > [C2OSO4]
-
≈ [NO3]
-
.
[OTf]- > [BF4]
- > [NTf2]
-
≈ [PF6]
-, which is the order of the
decreasing basicity of the ions, and for cations the order is
[(HO)2C2NH3]
þ > [C2NH3]
þ > [CnCmim]
þ. The dielectric
constant was also found to decrease as the length of the alkyl chain
on the cations increased. These effects are similar to the behaviors
seen for molecular solvents.
The static dielectric constant has also been estimated from
measurements of the speeds of sound, densities, and heats of
vaporization (Table 1).119 The speeds of sound and densities provide
the basis for the estimation of the cohesive energy densities (CEDs)
of the ionic liquids, while the heats of vaporization give the internal
pressures (Pi). Together these values then provide the basis for the εr
calculations. Given how different this method of deriving εr is from
dielectric spectroscopy, it is remarkable how well the results agree.
However, it should be noted that the only ionic liquid for which a
significant difference was found using the two techniques is
Table 1. Estimated Static Dielectric Constants for Ionic
Liquids
ionic liquid
εr (dielectric
spectroscopy)118
εr (Pi þ
CED)119
εr (other
methods)
[C1NH3][HCO2] 40.3
[C2NH3][HCO2] 30.3
[C4NH3][HCO2] 29.2
[(HO)2C2NH3][HCO2] 57.3
[C2C1im][C2OSO4] 27.9 13.5 37-39
101
[C2NH3][NO3] 26.2
[C2C1im][OTf] 15.2 15.8
[C4C1im][OTf] 13.2 13.5
[C2C1im][BF4] 12.8 14.8
[C4C1im][BF4] 11.7 12.9
[C4C1im][PF6] 11.4 14.0 11.4
120
[C6C1im][PF6] 8.9 11.1
[C2C1im][NTf2] 12.25 11.5 <10
35
14-16101
[C3C1im][NTf2] 11.80 10.6
[C4C1im][NTf2] 11.52 9.4
[C5C1im][NTf2] 11.45 7.9
[C2C1C1
2im][NTf2] 11.45
[C2C1C1
2im][NTf2] 11.6

I dx.doi.org/10.1021/cr1003248 |Chem. Rev. XXXX, XXX, 000–000
Chemical Reviews
REVIEW
[C2C1im][C2OSO4], which is the only ionic liquid from the more
polar sets according to microwave spectroscopy that has been tested
by both techniques. The trend that, for a given anion, changing the
length of the alkyl chain on the imidazolium cation decreases the εr is
seen. However, while a general trend of more basic ions giving higher
εr values is seen, the precise ordering {[OTF]
- > [PF6]
-
≈ [BF4]
-
> [NTf2]
-
} is different from that seenwithmicrowave spectroscopy.
The dielectric constants of some ionic liquids have also been
estimated from measurements of the spectra of a number of probes.
The band shape of the emission spectrum of pyrenecarboxaldehyde
was used to give an upper limit of εr < 10 for [C2C1im][NTf2].
35 By
comparing the spectrum of pyrene with that of 6-propionyl-2-(N,N-
dimethylamino)naphthalene (PRODAN), Baker et al. deduced a
value of εr = 11.4 for [C4C1im][PF6],
120 the same as that derived
from microwave dielectric spectroscopy. The excellent agreement of
these completely independentmeasurements suggests that the values
derived from microwave spectroscopy can indeed be relied upon, or
at least that these measurements are recording the same behaviors.
Good agreements have also been found with values derived
from infrared spectra of water and D2O for [C2C1im][NTf2],
although again there was greater deviation for the values for
[C2C1im][C2OSO4].
101
5.2.2. Hildebrand Solubility Parameter, δH. For any
species to be dissolved in a solvent, the molecules or ions of the
solvent must first be separated. Similarly, any process that involves a
significant change in the molar volume of the solute(s) will involve a
change in the separations of the solvent species. The Hildebrand
solubility parameter, δH, attempts to measure the energy required to
do this:115
ðδHÞ
2
¼
ΔUv
Vm

ΔHv - RT
Vm
ð1Þ
where Vm is the molar volume of the solvent andΔUv and ΔHv are
themolar energy and enthalpy of vaporization to a gas of zero pressure.
It was for many years believed that ionic liquids were com-
pletely nonvolatile so direct access to the Hildebrand parameter
would not be possible. However, this is now known to have been
a false assumption.121 Since this realization, the enthalpies of
vaporization have been measured for ionic liquids and δH values
derived.122 For a range of imidazolium, pyrrolidinium, and
ammonium ionic liquids, including some with alcohol-functio-
nalized chains, values for δH in the range 16.3-26.5 (J cm
-3)1/2
were found, which are of the same order as those found for short
to medium chain length alcohols. The values were found to be
dependent upon the anion of the ionic liquid, with the values for
the bis(heptafluoroethylsulfonyl)imide {[beti]-} ionic liquids
being lower than those of the [NTf2]
- ion, and to decrease with
increasing chain length of the alkyl chains on the cations.
Where comparison is possible, the δH values derived from
enthalpies of vaporization are slightly lower than those derived from
surface tension measurements123 and significantly lower than those
derived from viscosities124,125 or chemical reactivity.126 For molec-
ular solvents a gas of zero pressure consists of isolatedmolecules and
is well approximated to under usual experimental conditions.
However, it has been shown that for ionic liquids it is likely that
any gas phase is composed of ion pairs and higher order clusters.127
It is likely that this would lead to an underestimate for δH.
5.2.3. Single-Molecule Spectroscopic Probes. There are
many polarity scales that are based upon the measurement of the
spectrum of a single probe molecule.115,117 This approach has the
advantage that the measurements are often simple and can be made
quickly using readily available equipment. However, while the
response of the probe to the solvent in which it is dissolved is
determined by all possible solvent-probe interactions, there is no
reason for there to be an equal contribution from all of these for all
probes. Hence, it is important to be as aware as possible of which
interactions are likely to have a strong effect on the particular probe
used andwhich have a lesser effect. Since ionic liquids are composed
of both anions and cations, either of whichmay preferentially solvate
a particular probe molecule, this requirement is even greater than is
usually the case for molecular solvents. In many papers one of these
scales is reported for a small number of ionic liquids and then used as
a basis tomake statements about their overall polarity.None of these
polarity scales are capable of delivering this kind of information,
except in the vaguest sense. Finally,many of the probes used for such
measurements are used in very low concentrations and may be
susceptible to preferential solvation by impurities in the ionic liquids,
the presence of which can lead to distorted results. However, these
problems do not mean that these data are without use, quite the
reverse; it just means that care is required when comparing results.
5.2.3.1. Electronic Absorption Spectra. The first solvatochro-
mic dye to be used with a number of ionic liquids was Nile red
(Figure 4).128 The ionic liquids used were composed of [CnC1im]
þ
cations with [NO2]
-, [NO3]
-, [BF4]
-, [PF6]
-, and [NTf2]
-
anions. The values of the energy of the electronic transition,ENR, did
not vary greatly and characterized the ionic liquids as being
moderately polar liquids, in broad agreement with the εr values
reported above.
The effect of hydrogen bond donation to Nile red from the
cation of the ionic liquid can be seen when comparing the ENR
values of imidazolium salts protonated on one of the ring
nitrogen atoms129 with that of the equivalent methyl-substituted
salt {e.g., ENR for [C4C1im][BF4] = 217.2 kJ mol
-1 and ENR for
[C1Him][BF4] = 212.5 kJ mol
-1
}. This is also seen when
comparing the value for [(C10)3C1P][NO3] (ENR = 225.8
kJ mol-1) with that of [C8C1im][NO3] (ENR = 217.4 kJ mol
-1).130
This strong cation effect arises due to the ease of hydrogen bond
donation to Nile red. Functionalization of the side chain of the
imidazolium cation has also been found to give an increase in the
polarity of the resultant ionic liquid, regardless of whether a
hydrogen bond donor {[(HO)2C2C1im][NTf2], ENR = 211.9 kJ
mol-1} or hydrogen bond acceptor {[(C1OC2)C1im][NTf2],
ENR = 213.5 kJ mol
-1
} is used.131
The anion also has an effect on theENRvalue of the ionic liquids.
128
TheENR values for several 1-ethylimidazolium ionic liquids have been
compared to the pKa of the conjugate acid of the anion of the ionic
liquid.120This study shows a general trend thatmore basic anions lead
to higherENR (less polar) values for the ionic liquids. This trend is the
opposite to that observed for the εr values reported above (more basic
anion = more polar ionic liquid). This shows the complexity of these
polarity phenomena and warns against attempts to use any one scale
to be a general polaritymeasure. It is known that inmolecular solvents
ENR is affected by the ability of the solvent to hydrogen bond donate
to the Nile red.132 It can be surmised that this dominates any
contribution from direct interaction of the ionic liquid (cation) with
Figure 4. Nile red.

K dx.doi.org/10.1021/cr1003248 |Chem. Rev. XXXX, XXX, 000–000
Chemical Reviews
REVIEW
electron pair donor and is traditionally derived from the 31P NMR
chemical shift of solutions of triethylphosphine oxide in the
solvent.115,144 These have not been measured for a wide range of
ionic liquids. However, the Raman spectra of diphenylcyclopropane,
particularly the {CdC þ CdO} mode, have been used to estimate
the AN of several ionic liquids.145 The values were found to vary in
the order [C4C1im][PF6] (27.7) > [C4C1C1
2im][BF4] (27.5) >
[C4C1im][BF4] (26.9) = [C4C1im]Cl (26.9) > [C4C1im]
[NTf2] (25.2) > [(C1OC2)(C2)2C1N][BF4] (23.9) > [(C1OC2)-
(C2)2C1N][NTf2] (20.3). These values are of the same order as that
for CHCl3 (23.1) and are considerably less than that for methanol
(41.5) and the values that the authors estimated usingmeasurements
basedonReichardt’s dye (ca. 36).While theAN is not the same as the
hydrogenbonddonorproperty of the liquid, onemight have expected
greater agreement for these values (see below). However, the general
trend that imidazolium-based ionic liquids have a greater AN than
quaternary ammonium salts is in agreement with expectations.
The νCdO IR stretching frequencies of acetone have been shown
to correlate with the AN of a range of solvents.146 The range of
values found in a variety of imidazolium-based ionic liquids was very
narrow (νCdO = 1706-1711 cm
-1), evenwhen the alkyl chain was
substitutedwith-OHand-CNgroups, and similar to that of polar
solvents (both protic and aprotic).147 Similar results were found for
theνCdO IR stretching frequencies ofDMF.When the authors used
Fe(CO)5 as the probe, they noted that the polarity of the ionic liquid
decreased with increasing alkyl chain length on the imidazolium
cation and that there was also an anion effect, similar to that seen for
Reichardt’s dye and Nile red.
5.2.3.3. Electron Paramagnetic Resonance. The 14N hyper-
fine coupling constant of the electron paramagnetic resonance
(EPR) spectrum of 4-amino-2,2,6,6-tetramethylpiperidine-1-oxyl, a
commercially available stable free radical, has been used to measure
the polarity of a small number of ionic liquids.148 The values correlate
very well with the ET(30) scale and so arise from a combination of
hydrogen bonddonation from the solvent andCoulombic, dipolarity,
and polarizability effects. EPR also provides information about the
tumbling of the probe solute in a solvent by its effect on the spin
relaxation times (τr). τr times in the ionic liquids are longer than those
in the molecular solvents studied. This was attributed partly to the
higher viscosities of the ionic liquids andpartly to a hydrogen-bonding
interaction between the cation and the nitroxide functionality of the
probe molecule. This technique has the advantage that it does not
require that the ionic liquid is transparent to UV and visible
wavelengths and may be used for colored ionic liquids.
The R-hydroxydiphenylmethyl radical bears a -OH group, and
consequently, its EPR spectra in ionic liquids highlight the impor-
tance of the anion in its solvation.149 The hyperfine coupling
constants were found to increase with increasing Gutmann donor
number (DN) of the anions, implying the formation of a hydrogen
bond between the probe and the ionic liquid. The values ranged
from those similar to that of 1,2-dichloroethane (aOH = 2.08) for
[C4C1im][PF6] (aOH = 2.03) to values closer to that of pyridine
(aOH = 3.25) for [C4C1im]2[WO4] (aOH = 3.21); [C4C1im]
[NTf2] (aOH = 2.28) had a value similar to that of benzonitrile
(aOH = 2.27). As expected, no correlation of the values was found
with the Reichardt’s dye spectra in the ionic liquids.
5.2.3.4. Fluorescence Spectra. Fluorescence spectra of a
number of probe dyes have been used to investigate solvation
phenomena in ionic liquids. The spectra of pyrene have generally
placed the ionic liquids in the polarity range of moderately polar
solvents.115,150,151 Although some contribution from solvent
hydrogen bond acidity has been noted,152 the mechanisms by
which spectral changes arise are poorly understood, and it is
difficult to draw conclusions in anything other than the most
general terms from thesemeasurements. The spectrum of pyrene
has been use to compare the microenvironment provided by a
supported ionic liquid phase (SILP) to the bulk ionic liquid.153
Salts with cations similar to those that give rise to ionic liquids
were covalently bound to polymeric resins via the cation. The
authors concluded that the functional surfaces of the SILP
maintained the same polarity as the equivalent bulk ionic liquids.
The Stokes shift of the fluorescence spectra of PRODAN has
been used to estimate the polarity of [C4C1im][PF6] as being
between those of chloroform and acetone.115 Correlations of the
fluorescence maxima of PRODAN and coumarin 153 (C153)
with ET(30) values for molecular solvents have been applied
to the results found in a number of ionic liquids {[C2C1im]
[BF4], [C4C1im][BF4], [C2C1im][NTf2], [C4C1im][NTf2],
and [C4C1pyrr][NTf2]}.
154 However, these values were signifi-
cantly below the values obtained with Reichardt’s dye itself.135
Time-resolved fluorescence spectra allow the dynamics of solva-
tion in ionic liquids to be probed. The literature reports an ongoing
debate. The questions that are at issue include whether the cations
and anions of the ionic liquid solvate on different time scales, the
relative importance of ion translation as opposed to dielectric
relaxation, and what effects the correlated motion of ions has on
solvation. On the basis of the fluorescence spectra of PRODAN,
C153, and Nile red,155 Samanta et al. attributed an initial fast
component of the solvation to the anion and a longer component
to collective motion of both anions and cations. Armstrong et al.156
have proposed that solvation of C153 is by the cation, which mainly
contributes to initial solvationbehavior via dipole relaxation. Although
neutral, these three dyes all carry electron pair donor sites such as
carbonyl groups, and none have any hydrogen-bond-donating sites,
which leads one to expect that solvation by the cation is most likely.
The observed cation effect on the electronic absorption spectra of
Nile red (see above) strongly supports this proposal. Also using
C153, Maroncelli et al.157 noted the similarity between data for some
ionic liquids and those for supercooled liquids close to their glass
transition. This is a very reasonable description of many ionic liquids
at room temperature. The dissimilarity that they observed between
imidazolium-based ionic liquids and phosphonium ionic liquids again
suggests preferential solvation by the cation and that the polarizability
of the delocalized electrons of the imidazolium ring makes an
important contribution to initial solvation by these ionic liquids.
On the basis of viscosity correlations, they also consider the rotation
of the cation to give a significant contribution to the solvation
dynamics, which is then followed on a longer time scale by ion
translation.
120-Apo-β-carotenoic 120-acid (120CA) is structurally very differ-
ent from those dyes used above, bearing a carboxylic acid functionality
and an extended polyene chain. The lifetime of its intramolecular
charge transfer state (τ1) has been used to investigate a wide range of
ionic liquids.158 In molecular solvents τ1 has been correlated with a
function derived from the static dielectric constant and the refractive
index of the solvent, but this did not hold for the ionic liquids. The
results are in the range 40-110 ps, which are in the range found for
short chain alcohols (e.g., ethanol, 109 ps, and methanol, 49 ps).
Comparing the results for the ionic liquids [Cat][NTf2], τ1 varies in
the order [C4C1C1
2im]þ (96 ps) > [C4(C1)3N]
þ (85 ps) ≈
[C4C1pyrr]
þ (83 ps) > [C4C1im]
þ (65 ps) ≈ [(C2)3S]
þ (64
ps). This strongly suggests that interactions between the cation and
the 120CA are an important contributor to the spectral changes. A
general, but imperfect, trend of increasing lifetime with increasing

N dx.doi.org/10.1021/cr1003248 |Chem. Rev. XXXX, XXX, 000–000
Chemical Reviews
REVIEW
bond acidity and basicity of the ionic liquids in comparison to
those with simple alkyl chains.174 A similar effect on the hydrogen
bond basicity was found by the inclusion of an ether linkage in a
phosphonium ionic liquid.175
5.2.5. Antagonistic Behavior in Hydrogen Bonding. It
has been demonstrated crystallographically46,176 and by neutron
diffraction,177 NMR spectroscopy,178 vibrational spectroscopy,179
mass spectrometry,180 and theory181 that hydrogen bonding is an
important structure-forming factor in many pure ionic liquids182 and
that the degree of interionic interaction is dependent upon the ions of
the ionic liquid. It was first proposed that these interionic hydrogen
bonds could have a significant impact upon the ability of ionic liquids
to hydrogen bond to solutes to explain the anion effects upon the
selectivities observed for some Diels-Alder reactions.183 It was
proposed that the ability of the ionic liquid to form a hydrogen
bond with a solute molecule appears to come from one of the ions,
but that competition between the solute and the counterion of the
ionic liquid weakens this interaction. This has subsequently been
reaffirmed by several of the studies detailed above and by molecular
dynamics simulations.184
Let us first consider hydrogen bond donation by the cation.
This can be described in terms of two competing equilibria. The
cation may hydrogen bond to the ionic liquid anion:
Cþ þ A- h Cþ---A- K0 ¼
½Cþ---A-
½Cþ½A-
ð4Þ
The cation may also hydrogen bond to the solute:
Cþ þ solute- h Cþ---solute K00 ¼
½Cþ---solute
½Cþ½solute
ð5Þ
It can be easily seen that because of the commonality of [Cþ] to
both equations, [Cþ---solute] arises as a consequence of bothK00
and K0. For any given solute, an ionic liquid with a stronger
hydrogen-bond-donating cation will lead to a higher value for K00
and a greater [Cþ---solute]. For the same solute and ionic liquids
with the same cation, a greater hydrogen-bond-accepting ability
of the ionic liquid anion leads to a largerK0 and hence a reduction
in [Cþ---solute]. To summarize, the overall ability of an ionic
liquid to donate a hydrogen bond to a solute comes from the
ability of the cation to act as a hydrogen bond donor, reduced to
some degree by the ability of the ionic liquid anion to act as a
hydrogen bond acceptor. In a similar way, the overall ability of an
ionic liquid to accept a hydrogen bond from a solute comes from
the ability of the anion to act as a hydrogen bond acceptor,
reduced to some degree by the ability of the ionic liquid cation to
act as a hydrogen bond donor.
5.2.6. Apparent Discrepancies between Polarity Mea-
surements in Ionic Liquids. The use of the dielectric constant
estimated bymicrowave dielectric spectroscopy as a descriptor of
the solvent-solute relationship in ionic liquids has received
some criticism.185 This rests on three arguments. The first is
that ionic liquid structures are heterogeneous on the nano-
scale186 and that no continuum-based description can capture
this level of detail. The second is that the high frequency of
the measurements fails to capture the defining feature of ionic
liquids—the translation of ions—giving rise to an inadequate
description of the ionic liquid.185,187 Finally, the dielectric
constants for most ionic liquids describe them as considerably
less polar than several of the polarity scales described above and
fail to explain the outcome of quantitative investigations of
reactions in ionic liquids (see below), yet these dielectric
constants do agree with some other measurements described
above.35,101,119,120
We suggest that, to explain these apparent contradictions, it is
necessary to consider the time scale of the various processes. The
highly structured nature of the ionic liquids has important implica-
tions for how they interact with solute species, particularly with
respect to any solvent reorientation around a solute.185,186 For
example, 2-amino-7-nitrofluorene exhibits excitation-wavelength-
dependent fluorescence spectra in [C4C1im][BF4], [C4C1im]
[PF6], and [C2C1im][BF4], which has been interpreted as the dye
molecules occupying distinct nonexchanging (on the time scale of the
measurements) environments in these ionic liquids.188 Undoubtedly,
solvent reorganization around a solute in ionic liquids is slower than in
most commonplacemolecular solvents. Amodel of solvation in ionic
liquids that does not encapsulate ion translation is likely to fit better to
measurements that occur on a time scale that is faster than that of ion
translation, such as in fluorescence35,120 or vibrational spectros-
copy,101 than those whose time scale is slower than that of ion
translation, such as chemical reactions involving atom or group
transfers. Although electronic absorption spectra arise from electronic
transitions in molecules that are rapid, their solvatochromism most
often arises from interactions of the solvent with the ground state of
the probe, which is in equilibriumwith its surroundings. Hence, these
polarity scales should not be expected to agree with dielectric
constants derived from microwave dielectric spectroscopy.
Another area inwhich there has been somedisagreement has been
in the estimation of the abilities of ionic liquids to act as hydrogen
bond donors. Measurements of the hydrogen bond donor ability of
ionic liquids that have used Reichardt’s dye have described some
ionic liquids as strong hydrogen bond donors. These results agree
well with those found for the solvatochromism of merocyanine dyes
and with the analysis of many of the reactions described below.136
Several other studies of hydrogen bonddonor abilities of ionic liquids
that have compared their results to these results, including those
based upon GC measurements,141,168 the UV-vis spectrum of
Fe(phen)2(CN)2,
167 the fluorescence spectra of PRODAN and
coumarin 153 (C153),154 and the Raman spectrum of diphenyl-
cyclopropane,146 describe these ionic liquids as considerably poorer
hydrogen bond donors. For several of these ionic liquids, not only
were the derived values lower, but the relative effects of changing
cations or anions were also different.What is immediately clear when
comparing across the whole range of measurements is that when the
assessment of the hydrogen-bond-donating ability of an ionic liquid is
based upon a measurement of the property of a charged probe, the
value derived is greater than when a neutral probe is used. This
observation can now be used to explain these differences.
The first assumption behind the use of empirical polarity scales to
study ionic liquids is that if the response of a particular probe solute,
or solutes, is the same as that in some knownmolecular solvent, then
it can be said that the polarities of the ionic liquid and the molecular
solvent are the same and that the appropriate value of the parameter
can then be assigned to the ionic liquid. To then compare the results
of these different scales for ionic liquids carries the second, implicit
assumption that the effect of transferring fromamolecular solvent to
an ionic liquid is the same for all probes. It is unlikely that this would
be perfectly true for any potential probe, but there may be some for
which this difference is great enough to significantly change the
results. Hence, it is important to consider the nature of the solute as
well as the solvent.
There is some debate among theoreticians as to whether the
interactions between cations and anions in ionic liquids should be
considered as Coulombic interactions with overlying hydrogen

P dx.doi.org/10.1021/cr1003248 |Chem. Rev. XXXX, XXX, 000–000
Chemical Reviews
REVIEW
protons are not hydrogenbonding strongly to the [NTf2]
- anionnor
is the imidazolium cation hydrogen bonding to the lone pair on the
amine. High-boiling, highly functionalized solutes [methoxybenzene,
(hydroxymethyl)benzene, 1,2-ethanediol, and 1,4-butanediol] dis-
solved in [C2C1im][NTf2] also showed somewhat higher activity
coefficients than their less functionalized counterparts.199 By contrast,
the same group also explored solutions of n-alcohols and benzene in
[C4C1im][C8SO3] and [C8C1im][BF4] to reveal extensive hydro-
gen-bonding interactions between methanol and the [C8SO3]
-
anion.200
Secondary and tertiary alcohols were discovered to interact more
favorably with the ionic liquid than the corresponding primary
alcohol containing the same number of carbons.201 This is likely to
be due once again to hydrogen bonds donated by the hydroxyl
proton to the anion of the ionic liquid, as one would expect on the
basis of their Kamlet-Taft values (see above). By increasing the
alcohol chain length {methanol, 1-butanol, and 1-hexanol dissolved
in [C6C1im][NTf2]}, a small increase in γ
¥ was found to occur.202
This suggests that for this ionic liquid the overall effect arises as a
combination of changing hydrogen bonding and interactions
between the alkyl chain of the alcohol and that of the imidazolium
cation.While there was little difference in γ¥ values for CH2Cl2 and
CHCl3 dissolved in [C2C1im][NTf2] and [C2C1C1
2im][NTf2],
196
Hi
E,¥ for CHCl3 was substantially more negative. The authors
attributed this to the high polarizability of the chlorine atom and the
“special” strength of ion-induced dipole interactions. However, it
seems to us that strong hydrogen bonding between CHCl3 and the
anion of the ionic liquid is more likely to be responsible.
γ
¥ values for alcohols, ketones, and aromatics dissolved in two
significantly different ionic liquids, [C4C1
4pyr][BF4] and
[C4C1im][PF6], yielded several interesting trends.
203 In addition
to the observation that anion effects were stronger than cation
effects, the authors noted some interesting trends with regard to
isomers. They found that branched alcohols interacted more
favorably with the ionic liquids than straight-chain alcohols of the
same carbon number. Additionally, for the xylenes, they found
that m-xylene interacted the least favorably with the ionic liquid.
It does appear that temperature plays a very strong role in the
selectivity of the various xylenes, with the order and ratios of
activity coefficients varying as the temperature is increased from
316 to 335 K. This temperature-dependent selectivity could be
an important design parameter for separation processes, with
xylene separation being an important industrial process. γ¥
values for 20 solutes dissolved in 3 [NTf2]
- ionic liquids plus
[C2C1im][C2SO4] revealed substantial deviations from Raoult’s
law for aliphatic hydrocarbons as opposed to aromatics, and this
was proposed as a potential means of separating solutes from an
ionic liquid via distillation.204 Selectivity increases of up to 168%
were reported for the use of ionic liquids as entrainers relative to
the standard 1-methylpyrrolidine for aqueous separations of
hydrocarbons. This study also revealed very strong interactions
between [C2C1im][NTf2] and various ketones, again an indica-
tion of the importance of hydrogen bonding in these systems.
Anthony et al.205 measured a variety of thermodynamic properties
for the dissolution of water in three ionic liquids: [C8C1im][BF4],
[C8C1im][PF6], and [C4C1im][PF6]. The reported measurements
included both Henry’s law coefficients and γ¥ values with a
comparison of the ionic liquid/water systems to organic solvents/
water. They found that theγ¥ value for water in [C8C1im][BF4] was
substantially lower than that in [C8C1im][PF6], which was slightly
higher than that in [C4C1im][PF6]. Not only did this demonstrate a
strong anion component to these interactions, but the relative
comparisons to organic solvents were quite interesting. While
[C8C1im][BF4] seemed to interact with water similarly to methanol,
[C4C1im][PF6] and [C8C1im][PF6] more closely resembled 2-pro-
panol or acetone. However, all of these ionic liquids are immiscible
with water, despite the γ¥ values {1.76, 6.51, and 5.36 for
[C8C1im][BF4], [C8C1im][PF6], and [C4C1im][PF6], respectively,
at 25 C} being similar to those of molecules that are miscible with
water. This clearly indicates that, despite very favorable interactions
between water and ionic liquids with “hydrophobic” anions {such as
[PF6]
-
}, other factors, most likely entropic ones, dominate the
solution thermodynamics.
An early review of the thermodynamic properties of solutes in
ionic liquids has been provided byHeintz,206 and a treatise on the
use of γ¥ values for selection of ionic liquids as entrainers has
been provided by Chiappe.207
5.3.1.2. Effect of Changing the Anion. Anion effects on γ¥
differ fundamentally from cation effects, mainly due to the different
interactions involved. As a simple example, the interactions of
hydrocarbon solutes with [NTf2]
- ionic liquids195 are distinctly less
favorable than those involving the [BF4]
- anion.192 This difference
was originally attributed to the [BF4]
- anion’s smaller size, but more
recent studies suggest that the intermolecular interactions are much
more dependent on the polarizability of the anion than initially
thought. In a diverse look at anion effects, Dobryakov et al.208
measured γ¥ values for eight alcohols in three ionic liquids with
varying properties: [C4C1im][PF6], [C4C1im][C1SO3], and
[C6C1im][NTf2]. They found that favorable hydrogen-bonding
interactions increased from [PF6]
- to [NTf2]
- to [C1SO3]
-, as
one would expect. The differences between the two hydrophobic
anions [PF6]
- and [NTf2]
- were relatively small, but the hydro-
philic [C1SO3]
- anion showed amarked lowering ofγ¥ values. This
provides strong evidence that specific hydrogen-bonding interactions
are much more powerful than any size-based effects.
Letcher et al.209 determined γ¥ values for organic solutes in
the ionic liquid [C6C1im][PF6]. Interestingly, they found that
benzene interacts more favorably with this ionic liquid than does
methanol. This suggests that hydrogen bonding to the [PF6]
-
anion is less important than the “cation-delocalized electron”
interactions, due largely to the weakly interacting nature of the
anion (large, symmetric, less polarizable, poor hydrogen bond
acceptor). A follow-up study explored many of the same solutes
(plus He) in the ionic liquid [C6C1im][BF4].
210 For alkanes,
alkenes, alkynes, and benzene, the γ¥ values were only 2-5% lower
when [BF4]
- was the anion as compared to [PF6]
-, indicating
essentially similar solute-solvent interactions. However, theγ¥ value
for methanol was more than 60% lower in [C6C1im][BF4] (0.75 vs
1.94 at 298.15 K), which has a much greater ability to act as a
hydrogen bond acceptor. This large difference in values suggests that
hydrogen bonding from the solute to the anion represents the largest
solute-solvent interaction between alcohols and this ionic liquid.
As part of that same series of studies, three very dissimilar ionic
liquids, [C8C1im]Cl,
211 [C6C1im][BF4],
212 and [C4C1im][(C1OC2-
OC2O)SO3],
213 were reported to have very similar γ¥ values. Also,
values for [C6C1im][NTf2]
214 were found to be very similar to those
for [C6C1im][PF6],
210 including the result that benzene interacts
more favorably with the ionic liquid than does methanol. This shows
that hydrogen bonding to the [NTf2]
- anion or [PF6]
- anion is
considerably weaker than for more hydrophilic ionic liquids.
For mixtures involving the ionic liquid [C4C1im][OTf], the γ
¥
values decrease with increasing temperature for alkanes and alkenes
and increase with increasing temperature for polar solutes and
water,214 further supporting the dominance of hydrogen-bonding

Q dx.doi.org/10.1021/cr1003248 |Chem. Rev. XXXX, XXX, 000–000
Chemical Reviews
REVIEW
interactions in the latter systems as hydrogen bonds weaken with
increasing temperature. These results were borne out later for polar
solutes dissolved in [C4C1im][OTf]. Hereγ
¥ values increased along
the series ethyl acetate > tetrahydrofuran >1,4-dioxane > acetone >
acetonitrile.215 This trend clearly indicates that aprotic solutes are
capable of favorably interacting with the ionic liquid. It is unclear
whether this is predominantly through hydrogen bond donation by
the cation to the solute or simply ion-dipole interactions. However,
the much lower values obtained for alcohols demonstrate that
solute-anion hydrogen bonding was more significant than either
solute-cation or ion-dipole interactions. The similarity in value
obtained for CHCl3 and CH2Cl2 (a trade-off of much polarity for a
little hydrogen bonding) clearly demonstrates the complexity of
balancing these interactions. A related study on [C2C1im][BF4]
revealed γ¥ values that were uniformly higher than for the same
solutes dissolved in [C4C1im][OTf].
216 The trends were, however,
identical with γ¥ values increasing along the series ethyl acetate >
tetrahydrofuran >1,4-dioxane > acetone > acetonitrile. Therefore,
there would seem to be a relatively strong anion-related effect.
This indicates that hydrogen bonding from the solute to the anion
is very important to the activity coefficients. Measurements con-
ducted using [C3C1C1
2im][BF4] were also performed.
217 In this
instance γ¥ values were again uniformly higher than for the same
solutes dissolved in [C4C1im][OTf] and slightly higher than for
[C2C1im][BF4]. The trends were again identical with γ
¥ values
increasing along the series ethyl acetate > tetrahydrofuran >1,4-
dioxane > acetone > acetonitrile {as was the case for [C2C1im]-
[BF4]}. In this case the cation of the ionic liquid is capped by amethyl
group at the C2 position of the imidazolium ring, greatly reducing its
potential to donate hydrogen bonds. Since the trend in solvation
remains unchanged, this strongly suggests that favorable molecular
solute-solvent interaction in ionic liquids can generally be classified
as solute-anion hydrogen bonding being stronger than ion-dipole
interactions being stronger than cation-solute hydrogen bonding.
This trend is general for the classes of ionic liquids that do not contain
specific protic functionality. Measurements in [C6C1im][OTf]
yielded values that were substantially higher (1-2 orders of
magnitude) than for the polar solutes studied previously,218 while
results obtained in [C4C1im][OTf] demonstrated that this ionic
liquid behaved with polarity between those of [C4C1im][NTf2] and
[C4C1im][BF4].
219
Measurements for hydrocarbons and methanol dissolved in the
ionic liquid [C4C1im][C8SO3]
220 yielded some interesting trends
in γ¥ values. Even with the long alkyl chain on the anion, hydrogen-
bonding interactions between methanol and the anion of the ionic
liquid were found to be more favorable than interactions between
the π electrons on benzene and the ionic liquid, and both of these
were found to be much more favorable than van der Waals
interactions between the hydrocarbons and the ionic liquid. How-
ever, the octyl chain did reduce the selectivity of the ionic liquid-
solute interactions, reducing its potential as an entrainer for separa-
tions. Similar results were reported for [C4C1im][C8SO3],
221where
γ
¥ values are significantly lower than in themore common [BF4]
--
or [NTf2]
--based ionic liquids, particularly for alkanes. Meanwhile,
solutes dissolved in [C2C1im][OTs] exhibit significantly higher γ
¥
values at the low-polarity end (alkanes), but γ¥ is 2 orders of
magnitude lower for methanol. Therefore, while hydrogen bond
donation by solutes to either anion appears to be strong and
dominant, the interaction of the octyl chain with alkanes was
significantly more favorable than for the tosylate anion. This is
interesting, as it shows that a wide range of interactions are possible
with nonpolar solutes (from favorable to very unfavorable) even
among solvents that display strong hydrogen-bonding interactions,
giving rise to some degree of amphiphilicity.
[C2C1im][CF3CO2] is an extremely hydrophilic ionic liquid.
As such, the activity coefficients for alkanes222 were found to be
extremely high {roughly twice that of [C2C1im][BF4]} and for
polar solutes (especially water) extremely low {roughly half that
of [C2C1im][BF4]}. This presents an interesting option for
separations as this polar solvent should behave very differently
with solvents of varying functionality, particularly for aliphatic/
aromatic separations and solutes with differences in hydrogen-
bonding interactions.
Measurements for various organic solutes (benzene, 1-buta-
nol, 2-pentanone, 1-nitropropane, and dioxane) dissolved in
[SCN]-- and [NO2]
--based ionic liquids revealed, as expected,
that strong interactions between the polar groups on the solute
and the ionic liquid anions dominated all other effects.171
Domanska et al.223 studied γ¥ values for organic solutes and
water dissolved in the very hydrophilic (and even nucleophilic)
ionic liquid [C2C1im][SCN]. Due to the extreme nature of the
anion employed, the interactions with nonpolar solutes (such as
alkanes) were markedly more unfavorable than for other ionic
liquids, while interactions with polar solutes (alcohols) and water
were extremely favorable. Arising from these effects, the separa-
tion efficiencies for both aromatic/aliphatic systems and alcohol
systems were very high, yielding a promising solvent for con-
ducting such separations. However, it bears noting that [SCN]-
is potentially highly reactive in the presence of electrophiles.
Deenadayalu et al.224 measured γ¥ values for organic solutes
dissolved in the unusual ionic liquid [C8C1im][(C1OC2-
OC2)OSO3]. While the values for solutes of low functionality were
in line with those of other ionic liquids, the γ¥ values for methanol
and benzene were very low, with that of methanol being much lower
than unity. This suggests that this anion has very favorable ion-
dipole interactions (stronger than usual) and that this anion is
extremely adept at accepting hydrogen bonds from methanol. This
anion would be of interest for dissolving solutes that have a highly
acidic component.
5.3.1.3. Effect of Changing the n-Alkyl Group on an
Imidazolium Cation. The relatively weak anion-solute interac-
tions available in [NTf2]
- ionic liquids helped to reveal the effect of
changing the alkyl group of some imidazolium ionic liquids.225
Although the trends were the same, γ¥ values for organic solutes
in [C6C1im][NTf2] were universally lower than those measured for
ionic liquids with shorter alkyl chains, even when more polar solutes
(acetone, acetonitrile) were involved.Hence, theweak cation-solute
interactions becomemore favorable with increasing cation alkyl chain
length. These data also agree well with those obtained previously by
Letcher using the same technique (gas-liquid chromatography).213
Similarly, measurements made in [C8C1im][BF4] for linear alkanes
were lower than in [C4C1im][BF4],
226 indicating an increased
lipophilicity for the ionic liquid with the longer alkyl chain that can
mitigate other unfavorable solvent-solute interactions to some
degree.
Foco et al.227 conducted a single study containing γ¥ values
for several classes of organic solutes (alkanes, alcohols, ketones,
ethers, aromatic hydrocarbons, halogenated compounds) in
[CnC1im][BF4] (n = 2, 4, 6, or 8). This provides a convenient
means for examining the effect of lengthening the alkyl chain
length on a substituted imidazolium cation. For most classes of
solutes, the value of γ¥ decreases as the alkyl chain length
increases. However, alcohols are an exception to this trend.
The γ¥ value of each alcohol dissolved in the ionic liquids

R dx.doi.org/10.1021/cr1003248 |Chem. Rev. XXXX, XXX, 000–000
Chemical Reviews
REVIEW
increased when changing from the ethyl to the butyl side chain on
the imidazolium ring. After that, the γ¥ values decreased with
increasing chain length. Therefore, there is a maximum value
obtained at n = 4, corresponding to [C4C1im][BF4] having the
least favorable interactions with alcohol solutes. This is likely to
be due to the [C2C1im]
þ cation having particularly favorable
interactions with the polar hydroxyl group on the alcohols. This
effect is obscured somewhat by the increase in chain length to n =
4, before the normal behavior reasserts itself. The behavior for
chlorinated solutes is similar, though less pronounced, ruling out
a strong hydrogen-bonding component to this effect. The results
of Inoue et al.228 for n-alkanes in [C4C1
4pyr][BF4] help to
confirm this hypothesis for a less acidic cation. γ¥ values for
polar solutes dissolved in [C4C1im][BF4]
194 also support the
notion that ionic liquids can act with dual nature in chromatog-
raphy columns—as nonpolar stationary phases for nonpolar
solutes and as polar stationary phases for more polar solutes.
This amphiphilic behavior is an interesting indication of the
balance between side chain interactions with nonpolar solutes
and ion-dipole interactions with more polar solutes.
Mutelet and Jaubert determined γ¥ values for organic solutes
in an ionic liquid with a very long side chain, [C16C1im][BF4].
229
In this extreme example, all experimentally determined γ¥ values
(for 31 organic solutes) were between 0.5 and 2.7, indicating that
ion-dipole interactions and hydrogen-bonding interactions are
being muted by van der Waals interactions. Interestingly, despite
the long side chain, the alkanes still had the least favorable
interactions, universally increasing with increasing chain length.
Alcohols (methanol, ethanol) were still very favorable in their
interactions, though less so than with other [BF4]
--based ionic
liquids. Aromatics and especially functionalized small organics
(choloroform, dichloromethane, 1,4-dioxane, acetonitrile) now
exhibited the most favorable interactions with the ionic liquid.
These results suggest that lengthening the alkyl chain length of
the cation can decrease the attraction between polar solutes and
the ionic liquid while making interactions with nonpolar solutes
more favorable. However, for imidazolium ionic liquids made
with the [BF4]
- anion, it does not appear that interactions with
alkanes will ever be more favorable than with solutes capable of
hydrogen bonding.
5.3.1.4. Effect of Changing the Cation Type. γ¥ values for
solutes dissolved in [C2C1im][NTf2] were only slightly lower
than in [C2C1C1
2im][NTf2],
195 indicating that there was no
large cation effect stemming from substitution of the acidic C2
proton, with the important exception that the value for acetoni-
trile was higher in [C2C1im][NTf2]. Acetonitrile is a strong
hydrogen bond acceptor; therefore, it would be expected to
interact strongly with the C2 proton. Hence, some other con-
sequence of the presence/absence of the C2 proton on the ionic
liquid cation must be operating to oppose the direct hydrogen-
bonding effect. The fact that similar trends were not observed for
the n-alcohols again indicates that the importance of hydrogen
bond donation from the solute to the ionic liquid anion is
greater than that of the hydrogen bond donation by the im-
idazolium ring to the solute. This also agrees with the observation
of water solubility being controlled by the anion much more so
than the cation. Substitution of the C2 proton to produce [C3C1-
C1
2im][BF4] appears to vastly increase γ
¥ values for low-polarity
solutes,230 an indication of much less favorable solvent-solute
interactions. This cation effect was further illustrated by compar-
ison with γ¥ values for organic solutes in [C3(C1)3N][NTf2].
231
Interactions with the tetraalkylammonium ionic liquid were
found to be significantly less favorable than those found with
any [NTf2]
- ionic liquid with imidazolium or pyrrolidinium
cations.
The phosphonium ionic liquid [(C6)3C12P][(C2F5)3PF3]
revealed γ¥ values that were all less than unity, indicating very
strong favorable interactions with the hydrocarbon solutes.232
The authors attribute this to the long C12 alkyl chain and cation
interactions with the delocalized π-electron systems. Also, these
ionic liquids had substantially more favorable interactions with
hydrocarbons (both aliphatic and aromatic) than with alcohols,
indicating that hydrogen-bonding interactions are less important
for this particular ionic liquid. Bannerjee and Khanna measured
γ
¥ values for solutes in three ionic liquids based on the
[C14(C6)3P]
þ cation.233 The anions were Cl-, [BF4]
-, and
[NTf2]
-. They found strong solute-solvent interactions (all γ¥
values were less than unity) similar to those found in previous
studies involving tetraalkylphosphonium ionic liquids.232 These
values are generally 1 order of magnitude lower than for the
corresponding imidazolium-based ionic liquids, indicating sub-
stantially more favorable interactions of a very strong nature. The
generally longer alkyl chains used for the tetraalkylphosphonium
cations that give rise to ionic liquids, coupled with the relatively
obscured charge of the cation, could explain this more lipophilic
nature. The only unfavorable interactions observed involved
alcohols dissolved in the [BF4]
--based ionic liquids. Since these
interactions were far less favorable than those for the less
hydrophilic Cl- and [NTf2]
- anions, it is likely that specific
interactions, such as hydrogen bonding, are unfavorable in these
systems. Measurements based on [(C4)3C1P][C1OSO3],
234 a
relatively more symmetric cation (compared to other phospho-
nium-based cations studied by the same group), yielded interac-
tions with alkanes that were substantially less favorable. The
geometry of these cations does appear to be relevant, as the less
delocalized charge combines with the bulk and symmetry of
the alkyl chains to exclude these solutes from the interior of
the cation. Interactions with more polar solutes (ion-dipole)
were less affected by the cation change than by the anion. Indeed,
despite the use of a hydrophilic anion, the ionic liquid dis-
played behavior very similar to that of [C6C1im][PF6]—a
hydrophobic ionic liquid. For the very hydrophobic ionic liquid
[C14(C6)3P]
þ[NTf2] the most favorable interactions were with
alkynes and benzene.235 The γ¥ values for four-carbon to six-
carbon alkanes and alkenes were very similar to that of methanol.
Therefore, in this extreme example, the hydrogen-bonding
effects are still present, though extremely muted, and favorable
van der Waals interactions are becoming more prevalent. It
would seem that generic ion-dipole interactions (between the
ionic liquid and the π-electrons of the solutes) are dominant in
these systems. Finally, measurements on [C14(C6)3P]
þ[(2,4,4-
trimethylpentyl)phosphinate] yielded quite extreme results.236
The very lipophilic nature of this ionic liquid leads to all solutes,
including straight-chain alkanes, having γ¥ values below unity.
The contrast between phosphonium-based and ammonium-
based cations is startling. Such a seemingly simple change yields
qualitatively different behaviors; both cations yield hydrophobic
ionic liquids, but tetraalkylammonium cations generate unfavorable
interactions with all substrates while tetraalkylphosphonium cations
yield ionic liquids displaying extraordinarily strong favorable inter-
actions with all substrates—from alkanes to alcohols. The relatively
charge diffuse nature of the larger phosphorus atomperhaps leads to
some sort of charge-shielding effect, while the alkyl groups are
relatively free to engage in van der Waals interactions with alkane

S dx.doi.org/10.1021/cr1003248 |Chem. Rev. XXXX, XXX, 000–000
Chemical Reviews
REVIEW
substrates. By contrast, the ammonium cations appear to behave
similarly to point charges, as almost inorganic ions in nature, leading
to very unfavorable interactions with all organic substrates.
Mutelet et al.237 reported γ¥ values for organic solutes
dissolved in the analogous ionic liquids [C3Cnim]Br and
[(CdC2)Cnim]Br (n = 3, 8, 10, or 12). Very strong hydrogen-
bonding interactions were inferred for alcohols dissolved in this
latter ionic liquid. The resulting γ¥ values were successfully
correlated using an LSER to characterize the solute-solvent
interactions. Measurements were also made in two more unusual
bromides, [1-[[(meth)acryloyloxy]hexyl]-3-methylimidazolium]-
Br and [1-[(acryloyloxy)propyl]-3-methylimidazolium]Br, with
cations that contain very polar side groups.238 These ionic liquids
yielded very unfavorable interactions with nonpolar alkanes and
very favorable interactions with polar compounds such as
alcohols, even compared to other ionic liquids (γ¥ values as
low as 0.15 for methanol). This indicates that polar substituents
on the cation can interact favorably with both acidic and basic
sites on a solute (such as the hydroxyl in an alcohol) to yield very
favorable solutions and thus great potential for reactions and
separations.
Both the chain length and cation type (imidazolium vs
pyrrolidinium) were examined in the same study of several
classes of organic solutes and water dissolved in a range of
[NTf2]
- ionic liquids.239 No noticeable effect of cation type on
the γ¥ values of alkanes dissolved in ionic liquids was observed.
However, there was a noticeable decrease in γ¥ as the alkyl chain
length on [CnC1im][NTf2] (n = 2, 4, 6, or 8) increased,
indicating that interactions become more preferable as the chain
length increases. This suggests that favorable interactions be-
tween the alkane and the alkyl chain on the imidazolium cation
are important. Similar trends for ethanol (despite the polarity
change) indicate that cation hydrogen bond donation to the
solute follows the same trend with respect to alkyl chain length.
As expected, the γ¥ values for water in these ionic liquids
demonstrated the opposite trend; γ¥ increased with increasing
alkyl chain length as the interactions became less favorable. The
authors also measured γ¥ values for solutes in a 50:50 mixture of
[C2C1im][NTf2] and [C4C1im][NTf2] and observed a close to
ideal mixing effect.
5.3.2. Modeling of γ¥ Values. In an early attempt to model
activity coefficients in ionic liquids, Diedenhofen et al.240 com-
pared the prediction of γ¥ values for 38 organic compounds in
several ionic liquids {[C2C1im][NTf2], [C1C1im][NTf2] and
[C4C1
4pyr][BF4]} at 2 different temperatures (314 and 344 K)
using the COSMO-RS model (conductor-like screening model
for real solvents).241 The data used were taken from several
literature sources. The general trends that they reported were
limited to an analysis of the model’s predictive capabilities, and
these results are discussed in more detail below.
Kato andGmehlingmeasured and correlated γ¥ values for various
organic solutes andwater in several imidazoliumandpyridinium ionic
liquids.242 The interactions between ionic liquids and these solutes
were generally found tobecome less favorable as the alkyl chain length
on the ionic liquid was increased. Also, the anion choice had a
significant effect on the γ¥ values for cyclohexane and benzene, with
interactions becoming more favorable along the series [C1SO4]
-,
[(C1OC2)OSO3]
-, [(C1)2PO4]
-, Cl-, [BF4]
-, [NTf2]
-, and
[PF6]
-. This strongly indicates the importance of delocalization of
electrons on the anion for determining solute-solvent interactions in
the absence of significant hydrogen bonding (ion-induced dipole
interactions).
Eike and Brennecke performed a large study on the prediction of
γ
¥ values for 38 organic solutes in 3 ionic liquids, [C4C1
3pyr][BF4],
[C2C1im][NTf2], and [C2C1C1
2im][NTf2].
243 They revealed that
QSPR techniques are quite accurate for the correlation of γ¥ values
for a wide range of solute/ionic liquid combinations. They also had
some limited success at predicting γ¥ values for “unknown” organic
solutes. The power of this technique, however, comes from the
physical understanding of solute-solvent interactions that can be
gleaned from the correlation parameters. Their key finding was that
the γ¥ values were most strongly dependent on the hydrophilicity
parameter log KOW, which is strongly dependent on anion selection.
Later QSPR analysis for γ¥ values of 38 solutes in 3 ionic liquids
confirmed this result.244
Verevkin et al. modeled solutions of aldehydes and ketones
dissolved in [C2C1im][NTf2] using several standard activity
coefficientmodels.245 In these systems, containing functionalities
with interesting potential specific interactions, the nonrandom
two-liquid (NRTL) model could accurately capture vapor-
liquid equilibrium data in these binary systems by extrapolating
from pure substance parameters. Similar conclusions were drawn
formixtures of [C4C1im][NTf2] with n-alcohols and benzene.
246
Despite the differences in specific interactions available in these
two systems (e.g., hydrogen bonding), the authors found the
NRTL equation superior for modeling the activity coefficients to
the UNIQUAC (universal quasichemical) or Margules equation.
Regular solution theory has also been used to model ionic
liquid solutions by estimating both solubility parameters and
Hansen parameters for the three ionic liquids [C4C1im][PF6],
[C8C1im]Cl, and [C2C1im][NTf2] from other physical property
measurements.247 The modeling was fairly successful, and the
authors also found that increasing the alkyl chain length of n-
alcohols decreased the favorable ionic liquid-alcohol interac-
tions considerably, further strengthening the premise that these
systems are dominated by hydrogen bonding from the alcohol
solute to the anion of the ionic liquid. The poor suitability of
regular solution theory for solutions containing polar solvents
would seem to limit the application of this particular model.
UNIFAC (universal functional activity coefficient) is a group
contribution model that has proved highly successful at correlating
γ
¥ values for a wide range of solutes in several [NTf2]-based ionic
liquids.248 These results compare favorably withCOSMO-RS-based
predictions and offer a possible route for predicting interactionswith
new solutes within a class of ionic liquids.
5.3.3. Summary. Activity coefficients at infinite dilution are an
excellent means of isolating solvent-solute interactions. They add
to the overall statement that most of the common ionic liquids
behave similarly to short-chain alcohols, in accordancewith spectros-
copic quantification of their solvent potential. However, the wide
availability of different cation and anion functionalities has led to a
varied set of results for solutes in ionic liquids.
In general, interactions with small alcohols (such as methanol)
and water are quite favorable; even water-immiscible ionic liquids
yield activity coefficients of less than unity at infinite dilution.
Those of functionalized organics can be either side of unity,
depending on the nature of the ionic liquid employed and the
functional group(s) of the solute. Unfunctionalized alkanes
nearly always yield the least favorable interactions, though the
activity coefficients can vary by up to 4 orders of magnitude
across the range of ionic liquids reported in the literature.
Functionalization to alkenes and then alkynes and finally aro-
matics has universally led to a lowering of activity coefficient and
thus ever more favorable solvent-solute interactions.

U dx.doi.org/10.1021/cr1003248 |Chem. Rev. XXXX, XXX, 000–000
Chemical Reviews
REVIEW
more viscous ionic liquid would lead to the photogenerated radicals
being kept together in the solvent cage longer and so givemore of the
rapid recombination, it would also lead to slower diffusion-controlled
recombination of the radicals once they had escaped the solvent cage.
Hence, it was not possible to derive a simple relationship between the
overall rate of recombination and the ionic liquids’ viscosities. The
authors did, however, note a distinct effect of changing the anion
from [NTf2]
- to [BF4]
- for the [C4C1im]
þ ionic liquids. They
considered that this might arise from direct interactions between the
ionic liquid ions and the activated complex for the reaction.However,
it should be noted that these types of solvent effects on the rates of
reactions are not usually large when there is no significant charge
rearrangement in the activation process.115 Their closer analysis of
the activation parameters for the reactions in these two ionic liquids
revealed a qualitative difference in the activation entropy for the
reactions, with the expected negative value (ΔSq = -107 J mol-1
K-1) in[C4C1im][BF4],butapositivevalue(ΔS
q=18Jmol-1K-1)
in [C4C1im][NTf2]. They concluded that this must arise due to
solvent reorganization in the transition process. This would imply
that the observed anion effect does not arise from the direct
interaction of the anions with either the starting materials or the
activated complex, but from the strength of the interaction between
the ionic liquids’ cations and anions, with the weaker of these,
[C4C1im]
þ
-[NTf2]
-, allowing at least some solvent reorganiza-
tion in the activation process and hence compensating for the large
negative entropy change associated with the bimolecular
recombination.
It can be seen from the above that the relationships between solute
kinetic behaviors and the ionic liquid viscosities of ionic andmolecular
starting materials for these reactions are quite different, and we now
propose a hypothesis to explain these apparently contradictory results.
For the ionic solutes the rates of the reactions described above do
appear to be classically diffusion controlled, whereas for neutral
solutes the rates of the reactions are considerably faster than expected.
In a study of N,N,N0N0-tetramethyl-p-phenylenediamine, TMPD,
and its mono- and dications in acetonitrile and a variety of ionic
liquids, [CnC1im][NTf2] (n = 4 or 10), [C4pyr][NTf2], or
[C14(C6)3P][NTf2], the diffusion rates were found to vary in the
order TMPD>TMPDþ >TMPD2þ.261 In acetonitrile the ratios of
these compounds were 1, 0.89, and 0.51, whereas in the ionic liquids
they were on average 1, 0.53, and 0.33. Clearly there is greater
difference in the diffusion rates of molecules and ions in ionic liquids
than is seen in acetonitrile.
We propose that when a salt is dissolved in an ionic liquid, its ions
interact directly with the charge-bearing part of the ionic liquid’s ions
and break the solvent-solvent ion interactions, replacing them with
solvent-solute ion interactions.Hence, the frictional forces and range
of motion experienced by these solute ions are similar in nature to
those experienced by the ions of the pure ionic liquid that give rise to
its bulk viscosity. Consequently, bulk viscosity measurements would
be expected, and are observed, to give a reasonable guide to the
diffusion of solute ions in ionic liquids and “normal” diffusion-
controlled behavior. However, molecular solutes such as those used
above are not expected to have such strong interactions with the ions
of the ionic liquids, unless they have the ability to form hydrogen
bonds with them.262 In the absence of these interactions, the solute-
solute Coulombic attraction of the ionic liquid ions is not (fully)
disrupted by these neutral startingmaterials. This has several potential
consequences.Thefirst is that it leaves neutral soluteswith a restricted
volume in which to move. With a reduced volume to occupy, the
encounter rate of the solute will be higher than expected for the
concentration when considering the whole volume of the ionic liquid
solution. The importance of this effect will depend upon the relative
strengths of the solvent-solvent ionic interactions as well as the
potential solvent-solute interactions and would be expected to be
sensitive to the types of cations and anions composing the ionic liquid.
In addition to this, the properties of this available volume will be
dictated more by the nature of the nonpolar parts of the ionic liquid,
the alkyl chains of the cations in the cases above. Hence, the frictional
forces experiencedbymoleculesmoving through this reduced volume
would be expected to be very different from those that give rise to the
bulk viscosities of the ionic liquids, particularly lacking the ion-ion
interactions. Consequently, bulk viscosity measurements do not give
a guide to the diffusion of these solutemolecules in ionic liquids. This
effect would be expected to be sensitive to the length of the alkyl
chains. Finally, if the solvent-solvent ionic interactions are disrupted
during the activation process for the reaction, then this will contribute
to a more favorable activation entropy for the reaction. This is
somewhat similar to the model used to interpret of the rates of the
SN2 reaction of tributylaminewithmethyl p-nitrobenzenesulfonate in
[C4C1im][OTf]
263 and would be expected to be sensitive to the
types of ions composing the ionic liquid. All of these sensitivities are
observed in the reactions described above.
When pyridinium ionic liquids have been used as solvents for
electron transfer reactions, a further effect has been observed.
When pulsed radiolysis is used on a neat ionic liquid, it is usual for
the solvated electron formed to not react rapidly with the ionic
liquid’s cations.264 However, when pyridinium ionic liquids are
used, solvated electrons react with the cation to give pyridinyl
radicals.265 These can, in turn, transfer an electron to solute
acceptors to give a solvent-mediated electron transfer
mechanism.266
6.3. Acid-Base Reactions
The well-known intramolecular proton transfer called “keto-
enol tautomerism” was one of the first chemical phenomena found
to be solvent dependent.267 The equilibrium of pentane-2,4-dione is
one of the most widely studied of these proton transfers
(Scheme 5). 1H NMR has been used to investigate these proton
transfers in the ionic liquids [C4C1im]X {X = [PF6], [BF4], [OTf],
or [NTf2]}.
268 The percentage of the enol present at fixed overall
concentrations of solute was used to compare the ionic liquids to
molecular solvents. Inmolecular solvents this was found to correlate
with the dielectric constant. However, the results in the ionic liquids
would lead to predictions of between 40 and 55 for their dielectric
constants, which are considerably higher than those found by
physical measurements (see above). It is likely that this equilibrium
is also very sensitive to hydrogen-bonding effects, which were not
captured in this analysis.
The tautomeric equilibrium of 2-nitrocyclohexanone has
also been studied in several molecular solvents and the ionic
liquids [C4C1im][PF6], [CnC1im][NTf2] (n = 2, 4, or 6), and
Scheme 4

V dx.doi.org/10.1021/cr1003248 |Chem. Rev. XXXX, XXX, 000–000
Chemical Reviews
REVIEW
[C4C1C1
2im][NTf2].
269 The equilibrium constants for the ionic
liquids were of the same order as those for polar aprotic solvents,
such as acetonitrile. By this measure, the most polar of the ionic
liquids was found to be [C4C1C1
2im][NTf2]. The relationship
found betweenΔG for this reaction and theHildebrand parameter,
δH, was used to predict the δH parameters for the ionic liquids,
which were found to be in the range 23.6-26.9 (J cm-3)1/2.
The discovery that acidic compositions of the chloroaluminate-
(III) ionic liquids produce superacidic protons, with Hammett
acidities as strong as -18, together with the potential of ionic
liquids to act as catalysts/solvents (see below) has stimulated a
considerable amount of interest in their acidity and in solutions of
acids in ionic liquids.270
The relative Lewis acidities of a series of chlorometallate ionic
liquids, [C4C1im]Cl-MClx [X(MClx) = 0.67;MClx =AlCl3, FeCl3,
ZnCl2, or CuCl], and [C4C1im][BF4] have beenmeasured using the
infrared spectra of CH3CN dissolved in the ionic liquids.
271 The
spectrum in [C4C1im][BF4] was found to be essentially unchanged
in comparison to that in neatCH3CN, suggesting that this ionic liquid
can be considered to be Lewis neutral. However, the spectra in the
other ionic liquids revealed a new peak that indicated that the ionic
liquids were Lewis acidic in the order [C4C1im]Cl-AlCl3
(2338 cm-1) > [C4C1im]Cl-ZnCl2 (2318 cm
-1) > [C4C1-
im]Cl-FeCl3 (2310 cm
-1) > [C4C1im]Cl-CuCl (2292 cm
-1).
The effect of changing the composition of [C4C1im]Cl-AlCl3 on
the IR spectrumofCH3CNwas also investigated and shown to range
from 2338 cm-1 for X(AlCl3) = 0.67 to 2333 cm
-1 for X(AlCl3) =
0.55; for compositions with X(AlCl3) e 0.50, this peak was not
observed, indicating a lack of Lewis acidity. In this study similar
conclusions were drawn using the IR spectrum of pyridine, which
additionally showed an indication of Brønsted acidity in [C4C1-
im]Cl-AlCl3 arising from adventitious water. When the halogen-
ometallate complexes contained both bromide and chloride, the IR
spectra of CH3CN indicated the same ordering of acidity, [C4pyr]-
Br-AlCl3 > [C4pyr]Br-FeCl3 > [C4pyr]Br-CuCl2, but with no
evidence for Lewis acidity of this copper(II) system.272 When the
mixed halide systems [C4C1im]X-AlCl3 [X(AlCl3) = 0.67; X = Cl,
Br, or I] were compared, the Lewis acidity increased in the order Cl-
< Br- < I-, as one would expect.273
The Hammett acidity function gives a method for the mea-
surement of the acidity of nonaqueous Brønsted acid systems.274
This has led a number of groups to measure the Hammett
acidities of ionic liquids. Ideally, the measurements of absorp-
tions of both protonated and unprotonated forms of a basic dye
in the solution of interest are used to calculate the equilibrium
constant between. In this scale, themore negative the value ofH0,
the stronger the acid. Neat sulfuric acid has a value of -12, and
any system more acidic than this is generally considered to be a
superacid. When the acid is dissolved in a solvent, the value ofH0
becomes more negative with concentration. Also the value of H0
for the same acid at a fixed concentration in different solvents will
vary.115 It is generally difficult to measure the acidity of the pure
ionic liquids at room temperature, so measurements have been
made either at higher temperature or in a dilute solution in a
molecular solvent, either CH2Cl2 or water.
The overwhelming message from the results in Table 3 is,
perhaps unsurprisingly, that they depend upon the conditions of
the measurement. This makes it difficult to compare the results
across different reports by different groups. Ionic liquids with anions
that are known to react with water to produce acid, e.g., [BF4]
- and
[PF6]
-, have been excluded, yet several discrepancies remain in the
table. However, some general conclusions can be drawn.
None of these ionic liquids are in the superacid range. When the
Brønsted acidity arises from the anion of the ionic liquid, as in the
Table 3. Hammett Acidities of Ionic Liquidsa
ionic liquid measurement conditions H0
[(HSO3)
4C4C1im][OTf]
275 CH2Cl2 solution (10 mM) 0.01
[(HSO3)
3C3C1im][HSO4]
276 CH2Cl2 solution (5 mM) -0.01
[(HSO3)
4C4pyr][HSO4]
277 CH2Cl2 solution (10 mM) 1.21
[(HSO3)
3C4pyr][pTSA]
277 CH2Cl2 solution (10 mM) 3.98
[(HSO3)
4C4pyr][HSO4]
278 H2O solution (32 mM) 1.42
[(HSO3)
4C4pyr][HSO4]
279 neat ionic liquid, 110 C -2.1
[(HSO3)
4C4pyr][pTSA]
278 H2O solution (32 mM) 1.6
[(HSO3)
4C4pyr][pTSA]
279 neat ionic liquid, 110 C -1.2
[(HSO3)
4C4pyr][OTf]
277 CH2Cl2 solution (10 mM) 1.06
[(HSO3)
4C4pyr][H2PO4]
279 neat ionic liquid, 110 C -3.3
[(HSO3)
3C3pyr][pTSA]
278 H2O solution (32 mM) 1.59
[(HSO3)
3C3C2bim][BF4]
280 CH2Cl2 solution (5 mM) 0.19
[(HSO3)
3C3C2bim][HSO4]
280 CH2Cl2 solution (5 mM) 0.05
[(HSO3)
3C4(C1)3N][HSO4]
281 H2O solution (32 mM) 1.4
[(HSO3)
4C4(C2)3N][HSO4]
281 H2O solution (32 mM) 1.4
[(HSO3)
4C4(C1)3N][pTSA]
281 H2O solution (32 mM) 1.58
[C1Him][pTSA]
276 CH2Cl2 solution (5 mM) 2.45
[C1Him][OTf]
282 CH2ClCH2Cl solution (4.38 mM) 2.76
[C4Him][OTf]
282 1.38
[HC1
2pyr][OTf]282 CH2ClCH2Cl solution (2.89 mM) 2.19
[HC1
2pyr][OTF]283 1.53
[HC1
2pyr][CH3SO3]
283 CH2ClCH2Cl solution (4.38 mM) 3.13
[Hpyrrone][CF3CO2]
284 3.48
[cplm][H2PO4]
284 CH2Cl2 solution (5 mM) 2.08
[cplm][CF3CO2]
284 CH2Cl2 solution (5 mM) 3.35
[cplm][ClCH2CO2]
284 CH2Cl2 solution (80 mM) 5.04
[cplm][PhCO2]
284 CH2Cl2 solution (80 mM) 6.07
CH2Cl2 solution (80 mM)
[C4C1im][HSO4]
275 CH2Cl2 solution (80 mM) 0.73
[C4C1im][HSO4]
276 CH2Cl2 solution (80 mM) 0.08
[C4C1im][H2PO4]
275 2.55
[C4C1im][H2PO4]
276 CH2Cl2 solution (10 mM) 1.75
[C5C1im][HSO4]
276 CH2Cl2 solution (5 mM) 1.09
[C5C1im][H2PO4]
276 CH2Cl2 solution (10 mM) 1.92
[C4C2bim][HSO4]
280 CH2Cl2 solution (5 mM) 1.26
[C5C2bim][HSO4]
280 CH2Cl2 solution (5 mM) 1.5
[C4C2bim][H2PO4]
280 CH2Cl2 solution (5 mM) 1.83
[C5C2bim][H2PO4]
280 CH2Cl2 solution (5 mM) 2.28
CH2Cl2 solution (5 mM)
CH2Cl2 solution (5 mM)
CH2Cl2 solution (5 mM)
a [CnCmbim]
þ = dialkylbenzimidazolium, [Hpyrrone]þ = 2-pyrrolido-
nium, [cplm] = caprolactamium, and [pTSA] = p-toluenesulfonate.
Scheme 5

W dx.doi.org/10.1021/cr1003248 |Chem. Rev. XXXX, XXX, 000–000
Chemical Reviews
REVIEW
[HSO4]
- and [H2PO4]
- ionic liquids, the cation appears to
generally have little effect upon the overall acidity of the ionic liquid.
However, when the cation bears the proton, the anion does appear
to influence the overall acidity of the ionic liquid. Using a Brønsted
acidic anion appears to maximize the acidity of the ionic liquid for
any given cation. This suggests that even when a Brønsted acid, the
anion is the strongest base in the pure ionic liquid.
The gas-phase ion pair [(HSO3)
3C3C1im][HSO4] has been
investigated with density functional theory (DFT).285 In spite of
the limitations of being restricted to a single ion pair, this study
did give some useful insights. The first was that there is a low
activation energy for the proton transfer between the -SO3H
group and the [HSO4]
-. It was also noted that hydrogen-
bonding interactions between the protons of the imidazolium
ring could act to stabilize the deprotonated -SO3
- group, but
these bonds broke when the group was protonated.
The acidities of solutions of the strongBrønsted acidH[NTf2] or
H[OTf] dissolved in a range of ionic liquids, [C4C1im]X {X =
[BF4], [PF6], [SbF6], [NTf2], or [OTf]}, [C4Him][NTf2], or
[(C2)3NH][NTf2], have also been investigated using theHammett
method.286 In these studies, no major differences were observed for
these acid solutions for ionic liquids with a common anion, even
when the cation itself is protic as in [C4Him]
þ. However, large
differences in the acidities were found when the anion was changed,
such that [PF6]
- > [BF4]
-
. [NTf2]
- > [OTf]-. This suggested
solvent leveling by the ionic liquid, with the anion playing the role of
base. Except in [C4C1im][OTf], the acidities of the H[NTf2]
solutions were consistently higher than those of the H[OTf]
solutions, but these differences were not as great as those between
the ionic liquids with different anions.
The acidities of solutions of weak carboxylic acids in
[C4C1C1
2im][NTf2] and [C4C1pyrr][NTf2] have also been
investigated using sodium 4-nitrophenolate {Na[ArO]} as an
indicator.287 The acidities of these solutions vary such that H2O
> [C4C1C1
2im][NTf2] > [C4C1pyrr][NTf2]. For the range of
acids used there was a linear correlation of the acidities of the
resultant solutions between water and [C4C1C1
2im][NTf2], but
not with [C4C1pyrr][NTf2], in which a different order of
acidities was found. Given the observed insensitivity of other
acidity measurements of solutions of acids in ionic liquids to the
nature of the cation,286 this is an unexpected result. The authors
interpret their results in terms of an equilibrium between the
dissolved sodium 4-nitrophenolate and molecular acid (HA) on
one hand and 4-nitrophenol (ArOH) and the sodium salt of
the conjugate base of the acid (NaArO) on the other (eq 9).
However, this probably does not well represent the species
actually present in the solution (see below).
HA þNa½ArO h ArOH þ NaArO ð9Þ
It is now possible to make some general comments about the
speciation of protons in ionic liquids. First, it should be noted
that protons are small and highly mobile species and acidic
solutions should be regarded as highly dynamic. Also, protons are
never found in a naked state in a condensed phase and will always
seek out the best available base. The data above suggest, and by
analogy with the formation of [HCl2]
- in [C1C1im]Cl
288 and in
chloride-rich compositions of chloroaluminate ionic liquids,289
that in neutral ionic liquids the simplest possible Brønsted acidic
species when the acid is dilute is [HY2]
- (where Y- = any
anion). The anion Y- may arise from the conjugate base of the
acid added or from the ionic liquid. Thus, the addition of an acid
to an ionic liquid in dilute solution is perhaps best viewed as
yielding two species in equilibrium:
HA þ X- h ½HXA- ð10Þ
½HXA- þ X- h ½HX2
-
þ A- ð11Þ
where X- is the ionic liquid anion. The concentrations of each
species depend upon the relative basicities and concentrations of
the ions X- and A-. It is also possible that the following
equilibrium may need to be considered:
2½HXA- h ½HA2
-
þ ½HX2
-
ð12Þ
At higher concentrations ions of the type [(HA)xX]
- can also be
formed, as is found in [Hpyr][(HF)21F].
290
While solvent leveling by the ionic liquids’ anions described
above is to have been expected,286 the ordering found {[PF6]
- >
[BF4]
-
. [NTf2]
- > [OTf]-} is not. Whether based upon calcu-
lated gas-phase acidities291 or measured Kamlet-Taft β values,141
the expected order would be {[PF6]
- > [NTf2]
- > [BF4]
- >
[OTf]-}. It is not immediately clear as to how this difference arises.
However, the chemical similarities within and differences between
the two sets of ions, [PF6]
- and [BF4]
-, and [NTf2]
- and [OTf]-,
are well recognized. One possible reason for the unexpected acidities
is that [PF6]
- and [BF4]
- are more reactive than [NTf2]
- and
[OTf]- and that this results in the formation of new species thatmay
contribute to the acidity of the solution (e.g., HF, BF3, or PF5).
Another possibility is that there may be subtle differences in the
nature of the solution species, [HX2]
-. These ions can take either a
symmetrical form, [Cl-H-Cl]- as found in [(C2)4N][HCl2], or
an unsymmetrical form, [Cl-H
3 3 3
Cl]- as found in [(C1)4N]-
[HCl2].
292 Computational studies on [HCl2]
- in solution in
[C1C1im]Cl indicate that this species is more stable (i.e., less acidic)
when it is in the symmetrical form.288 Hence, if [NTf2]
- and
[OTf]- are capable of forming the symmetric form of [HX2]
- and
[PF6]
- and [BF4]
- are not, this could, at least in part, explain the
results. However, at this stage this remains a conjecture.
Less attention has been given to the effects of ionic liquids on the
strengths of dissolved bases. This is in great part due to known
sensitivity of imidazolium ionic liquids to evenmild bases, which can
lead to the formation of the deprotonated N-heterocyclic carbenes
(NHCs) and seriously affect these types of reactions. Not only are
these NHCs reactive species,293 they can also act as catalysts.294,295
Hence, NHCs can be significant contributors to the chemistry
observed even when present in low concentrations, so relatively
weak bases that can only partially deprotonate the imidazolium
cation may be sufficient to produce these effects. Attempts have
been made to prevent the formation of these carbenes by substitut-
ing the imidazolium ring at the 2-position, usually with a methyl
group. However, it has been noted that these methyl protons are
also sufficiently acidic to potentially lead to undesired side
reactions.296 Other substitutions of the imidazolium have also been
used,297 but it is usually simpler to avoid imidazolium cations when
studying base solutes.However, even in the absence of this particular
problem with imidazolium ionic liquids, the more general problem
of the well-established instability of ammonium salts in the presence
of bases/nucleophiles can further restrict the use of bases.6,83,298
Alkylphosphonium ionic liquids are more stable to bases than their
ammonium analogues,39 but even they are not immune to reaction
with bases.
A study of the equilibrium constants of the reactions of
p-nitrophenol (ArOH) with butylamine (BuNH3), piperidine
(pip), and triethylamine (Et3N) (eq 13) was conducted in

Y dx.doi.org/10.1021/cr1003248 |Chem. Rev. XXXX, XXX, 000–000
Chemical Reviews
REVIEW
the order of those of the polar aprotic solvents and short-chain
alcohols.However, no simple correlationwas foundwith the dielectric
constant, even for the molecular solvents investigated, and closer
inspection revealed the importance of specific solvent-solute inter-
actions in determining the reaction rates. The range of the solvent
dependency of the reaction varied dramatically between the different
nucleophiles; for [CH3CO2]
- k2 values varied greatly between the
different solvents, whereas for [SCN]- k2 values showed very little
change. The solvents providing the fastest reaction also changed with
the different nucleophiles. Similarly, the trend of the relative nucleo-
philicities of the nucleophiles was different in different solvents.
To understand these complex results, a Kamlet-Taft LSER
approach was taken. Acceptable correlations, using data sets from
both ionic liquids andmolecular solvents, were achieved for all of the
nucleophiles, clearly demonstrating that in this reaction there is no
special ionic liquid effect and that all significant interactions between
the ionic liquids and these solutes are adequately described by an
appropriate combination of their Kamlet-Taft parameters. R
appeared in all of the LSERs generated and always had a negative
value. Two of the nucleophiles, I- and [Ac]-, had best fit LSERs
with β in the correlations, but with a much lower contribution that
was attributed to the antagonistic hydrogen-bonding effects de-
scribed above. π* appeared in the best fit LSERs for three of the
nucleophiles, [CN]-, [CH3CO2]
-, and [CF3CO2]
-.
The negative R values can be attributed to slowing of the reaction
by the solvent hydrogen bonding to the nucleophile, stabilizing it with
respect to the activated complex,which is a less strong hydrogenbond
acceptor than the nucleophile itself. Since the hydrogen bond donor
property of the ionic liquids is largely a result of the cation structure,
this is generally seen as a strong cation effect for the ionic liquids. It
can be seen that the stronger the hydrogenbond acceptor character of
the nucleophile (represented by the nucleophile β values listed in
Table 5), the larger this effect. This is particularly emphasized when
anions of similar structure are compared {i.e., Cl- > Br- > I-,
[CH3CO2]
- > [CF3CO2]
-, and [CN]- > [SCN]-}. The effect of
changing theR value of the solvent on the reaction broadly splits into
threegroups:Cl-≈ [CH3CO2]
->[CN]-≈ [CF3CO2]
-
≈Br->
I-≈ [SCN]-. This again fits with the qualitative order that would be
predicted from theβ values of the anionic nucleophiles alone,with the
exception of [CN]-. This is an excellent demonstration of the
intimacy of the solvent-solute relationship. While a solvent may
have the potential to behave as a hydrogen bond donor to a solute
species, this potential is only realized when the solute is a good
hydrogen bond acceptor. The apparent anomaly of the [CN]- result
was explainedby invoking ahard/soft classificationof the interactions,
with the ionic liquids being considered to be hard hydrogen bond
donors. Subsequent studies in [ClO4]
- and [PF6]
- ionic liquids
showed similar effects and also demonstrated the influence that
dissolved water can have on the rates of these reactions.304
When studying the activation parameters for these reactions in ionic
liquids, it was noted thatΔSq for the reactionwasmore similar to that
of the reaction of free solvated ions in molecular solvents, but ΔHq
was more similar to that of the ion-paired nucleophiles in molecular
solvents. It was concluded that this apparent anomaly is a direct result
of the nature of ionic association in ionic liquids. In an ionic liquid,
there are no molecules available to separate ions and the ions of the
solute will be repelled by the ions of the ionic liquid bearing the same
charge (cation-cation or anion-anion) and attracted by anions of
opposite charge. Hence, in an ionic liquid, a solute ion will always be
closely associated with ions of opposite charge. Studies of the
compound [C1C1im]Cl by neutron diffraction showed that the
anion is coordinated by six cations within 6.5 Å.177 This is the same
thing as neither a free ion nor an ion pair. This was used to propose a
mechanism for the reaction which required the removal of a cation
from the nucleophile to liberate an active site for reaction before the
activation process itself occurred, which also involved a reduction
hydrogen bonding. However, a subsequent simulation study has
shown that theMe-p-NBS prefers to occupy a position similar to that
of the cation in [C4C1im][PF6].
305 The chloride ion occupies the
anion position in [C4C1im][PF6], with stronghydrogenbonds to the
[C4C1im]
þ. This suggests that while the two reactants must come
together for the reaction to occur, there is no significant overall
breaking of hydrogen bonds in this step, and its energetic demand is
not likely to have a major effect on the rates of these reactions.
Consequently, the negativeR effect for this reaction can be attributed
to the weakening hydrogen bonds between the solvent and the
chlorine atom as its charge density decreases as the activated complex
is formed (Figure 8). The simulations also show that hydrogen bonds
are also formed to the nucleofuge, which is developing charge during
the activation process, but this is obviously not sufficient to compen-
sate for the loss of hydrogen bonding to the Cl-.
The nucleofugacity of the leaving group is an important contri-
butor to the rate of SN2 reactions that can also be subject to solvent
effects. The reactions of NaN3 with a range of primary, secondary,
and tertiary halides and tosylates have been investigated in the ionic
liquids [C4C1im][PF6], [C4C1im][NTf2], and [C6pyr][NTf2].
306
While the substrates were fully soluble in the ionic liquids, there was
always some solid NaN3, which had been added in 3-fold excess,
present in the reaction flask. The reactions proceeded cleanly without
Table 5. LSER Correlations for ln k2 Obtained for Some Anionic Nucleophiles in [C4C1im][NTf2], [C4C1im][OTf],
[C4C1pyrr][NTf2], DMSO, DCM, and MeOH303
nucleophile R2 LSER
Cl- (β = 1.00) 0.99 ln k2 = 0.21 - 7.56R
Br- (β = 0.67) 0.97 ln k2 = -0.87 - 5.38R
I- (β = 0.30) 0.95 ln k2 = -2.57 - 3.05 R þ 1.16 β
[CH3CO2]
- (β = 1.49) 1.00 ln k2 = -2.37 - 7.60R þ 2.65β þ 1.83π*
[CF3CO2]
- (no β available) 0.97 ln k2 = -9.18 - 4.94R þ 5.76π*
SCN- (β = 0.33) 0.89 ln k2 = -3.87 - 2.46R
CN- (β = 1.37) 0.99 ln k2 = -3.16 - 5.07R þ 5.78π*
Figure 8. Activated complex for the reaction of Me-p-NBS and anionic
nucleophiles.

AB dx.doi.org/10.1021/cr1003248 |Chem. Rev. XXXX, XXX, 000–000
Chemical Reviews
REVIEW
different dependency (recalculated here, eq 16), having a positive
coefficient for R.321 This suggests that any inhibition of the
nucleophile through hydrogen bonding is outweighed by stabi-
lization of the developing nucleofuge in the activated complex.
ln k2 ¼ - 7:7 þ 2:5Rþ 4:2π ð16Þ
In an interesting application of the increased nucleophilicity of
amines in ionic liquids, the amine-functionalized [(H2N)
2-
C2C1im][PF6] has been used as a scavenger for excess electro-
phile in a toluene/[C4C1im][PF6] biphasic system.
322 Here the
neutral electrophile (e.g., PhCH2Cl) reacts with the amino group
of the [(H2N)
2C2C1im]
þ ion to give ionic products {e.g.,
[(PhCH2NH2)
2C2C1im]
2þ
þ Cl-) which are then trapped in
the ionic liquid layer. This was found to be a more effective
approach than others based upon solid-supported scavengers.
The ligand substitution reactions of [Rh(CO)(PPh3)2(NO3)]
with substituted pyridines have been investigated in
[C6pyr][NTf2] and [C4C1im][PF6] and compared to those in
dichloromethane.323 It was found that conversions to the pro-
duct [Rh(CO)(PPh3)2py][NO3] were poor in CH2Cl2 for all
but the most basic substituted pyridines used, whereas in the
ionic liquids complete conversions were achieved in all cases
studied. This was attributed to the ability of the ionic liquids to
preferentially stabilize the product ions over the molecular
starting materials, while the reverse is true in dichloromethane.
Many catalytic cycles contain a step inwhich a coordinated solvent
molecule is displaced by an incoming reagent. Consequently, the
ease of solvent displacement is a concern when selecting the best
solvent for catalytic reactions. Ionic liquids have a reputation for
being “noncoordinating” solvents. The photoinduced decarbonyla-
tion of [(C6H6)Cr(CO)3] leads to the rapid formation of
[(C6H6)Cr(CO)2(solvent)]. The rates of solvent displacements
from these complexes, which in an ionic liquid are most likely to be
[(C6H6)Cr(CO)2X] {X = (PF6) or (NTf2)}, have been investi-
gated in [C4C1im][PF6] and [C4C1im][NTf2].
324 In [C4C1-
im][PF6] rapid displacement of [PF6]
- by water was observed,
even in ionic liquids that hadbeen through a drying procedure.When
CH3CN was included in the reaction mixture, [(C6H6)-
Cr(CO)2(CH3CN)] was formed with a bimolecular rate constant
that was 2 orders of magnitude greater than that of the reaction of
[(C6H6)Cr(CO)2(C2H4Cl2)] with CH3CN in dichloroethane,
showing the ease of this displacement in the ionic liquid. However,
in [C4C1im][NTf2] no reaction occurred with CH3CN, indicating
that [NTf2]
- is a considerably more coordinating ligand.
6.4.3. Type c SN2 Y þ [R—X]
þ Reaction. Only one
substantial study of a nucleophilic substitution of a neutral nucleo-
phile with a charged electrophile is available in the literature. The
reactivities of n-butylamine, di-n-butylamine, and tri-n-butylamine
with dimethyl(4-nitrophenyl)sulfonium salts {[p-NO2PhS(CH3)2]
X, Scheme 8} have been investigated using UV-vis spectroscopy in
[C4C1im]X and [C4C1pyrr]X {X = [NTf2] or [OTf]} and a range
of molecular solvents (toluene, dichloromethane, tetrahydrofuran,
acetonitrile, methanol).325
Attempts to correlate the rates of these reactions with dielectric
constants of the liquids proved unsuccessful, and a Kamlet-Taft
LSER approach (Table 7) was taken to understand the effects of the
solvents on these reactions. The strongest effect for all of the amines
is an overwhelmingly negative β effect. This arises from interactions
between the solvent's hydrogen-bond-accepting site and acidic
protons on [p-NO2PhS(CH3)2]
þ which reduce the electrophilicity
of the reacting carbon center. This is ameliorated to some extent by
favorable interactions with the N-H protons of BuNH2 and
Bu2NH. A less intense negative solvent R effect was observed for
BuNH2 and Bu2NH, which was not observed for Bu3N, probably
being masked by the very large β effect for this amine. Finally, a
positive dependence on π* was observed for all of the reactions.
These were effects again rationalized by consideration of the
different strengths of solute-solvent interactions for the starting
materials and activated complex of the reaction (see Figure 10).
6.4.4. Type d SN2 Y
-
þ [R—X]þ Reaction. The Hughes-
Ingold rules predict that the rates of the reactions of charged
nucleophiles with charged electrophiles will be slower in ionic
liquids than in molecular solvents. A study of the reactions of QCl
salts (Q = quaternary ammonium or imidazolium cation) with [p-
NO2PhS(CH3)2]
þ has confirmed this, but also showed a change in
mechanism between the ionic liquids and molecular solvents.326 In
the molecular solvents (THF, acetone, 1-butanol, propylene carbo-
nate, CH2Cl2, CH3CN, or DMSO) the reactions were shown to
proceed through a mechanism that first forms a [p-NO2PhS-
(CH3)2]
þCl- ion pair which then goes on to react with a further
QþCl- ion pair to yield the products, whereas in the ionic liquids
{[C4C1im]X, [C4C1pyrr]X, [C4C1C1
2im][NTf2]; X = [NTf2] or
[OTf]} there was no evidence for kinetically significant ion pairs
and the reaction was that expected between the [p-NO2PhS-
(CH3)2]
þ ion and Cl- (Scheme 9).
The ion-pairing effects in molecular solvents prevented direct
comparison of the rates of the reactions in these solvents and the
ionic liquids with an LSER.326 The LSER for the ionic liquids
alone showed a strong negative R correlation, showing that
hydrogen bonding to the Cl- ion slows the reaction.
Where the charged electrophile has been a transition-metal
complex, this change of behavior has not been seen, in that no
ion-pairing effects have been seen in the molecular solvents.327,328
However, in these reactions significant quantities of unreactive salts
are added tomaintain a high ionic strength throughout the reactions
in molecular solvents. This appears to be sufficient to prevent the
formation of kinetically active ion pairs. The reaction of [SCN]-
ions with [Pt(terpyridine)Cl]þ in a series of [C2C1im]X
Scheme 8
Table 7. Results of LSER Fits for Reactions of Butylamines
with [p-NO2PhS(CH3)2]þ
R3N LSER R
2
BuNH2 ln k2 = -2.38 - 3.59R - 4.16β þ 2.10π* 0.96
Bu2NH ln k2 = -2.66 - 2.79R - 5.01β þ 2.89π* 0.99
Bu3N ln k2 = -5.62 - 6.46β þ 4.26π* 0.87
Figure 10. Activated complex for the reaction of [p-NO2PhS(CH3)2]þ
with butylamine.

AC dx.doi.org/10.1021/cr1003248 |Chem. Rev. XXXX, XXX, 000–000
Chemical Reviews
REVIEW
{X = [NTf2], [N(CN)2], [OTf], or [C2OSO3]} ionic liquids and
CH3OH showed that the rates of the reactions in the ionic liquids
were considerably slower than that of the reaction in methanol and
varied with the anion. This was interpreted as the anion interacting
with the [Pt(terpyridine)Cl]þ to slow the rate of the reaction.
A study of the reactions of [SCN]- ions with a similar
complex, [Pt(di(2-picolyl)amine)Cl]þ, in [C4C1im][NTf2],
[(C1OC2)C1im][NTf2], and [C4C1pyrr]X {X = [NTf2] or
[OTf]}, CH3OH, DMSO, and H2O revealed LSERs with a large
positive π* effect and a lesser positive R effect.263 This observa-
tion was used to suggest that hydrogen bonding to the emerging,
hard Cl- nucleofuge was having a greater positive effect than any
negative effect caused by hydrogen bonding to the soft [SCN]-
nucleophile.
6.4.5. Type e SN1 R—X Reaction.Unimolecular substitution
reactions in ionic liquids have been far less studied than bimolecular
substitutions. The hydrolysis reactions of a number of organic
triflates in [C4C1im][NTf2] have been shown to lead to byproducts
that were the result of rearranged carbocations.329 A more detailed
examination of the hydrolysis of 1-adamantyl mesylate was found to
yield a mixture of 1-adamantanol and diadamantyl ether, suggesting
that the 1-adamantanol reacts with the adamantyl carbocation
intermediate. On the basis of the activation parameters for the
reaction, the authors proposed that the reaction to the 1-adamanta-
nol followed an SN1 mechanism, although a dependency of the
reaction rate on the concentration of water in the ionic liquid was
noted.The rate constants for the ionic liquid reactionswere between
those found for protic molecular solvents (CF3CH2OH, CH3OH,
or CD3CO2D) and DMSO-d6. The authors concluded that carbo-
cation intermediates did form in these reactions, but that the rates of
ionization were not extraordinary and that they depended upon the
nature of the leaving group. This suggests that specific interactions
between the ionic liquid and the solute species are more important
in determining the ability of an ionic liquid to promote ionization
(Scheme 10) than the ionic nature of the solvent itself.
A study of the rates of the methanolysis of (R)-3-chloro-
3,7-dimethyloctane to (S)-3-methoxy-3,7-dimethyloctane in
[C4C1im][NTf2]/MeOH mixtures showed that the ionic liquid
did not significantly affect the rate of formation of the carbocation
intermediate.330 However, analysis of the stereochemical outcome
showed that there was greater racemization in the ionic liquid,
suggesting that both faces of the carbocation were available for the
nucleophile to attack. This in turn suggests that the intermediate is
more ion paired in the methanol and more dissociated and
symmetrically solvated in the ionic liquid. Analysis of the activation
parameters for this reaction showed that addition of a small amount
of ionic liquid to the reaction mixture led to stabilization of the
activated complex and a decrease in the activation enthalpy.331
However, at higher concentrations of ionic liquid an increased
ordering of the system becomes a significant contributor to the
overall activation energy as the charged activated complex forms
from the neutral substrate. In contrast to the stereochemical out-
comes, the activation parameters suggest a significant degree of ion
pairing in the activated complex. As noted in previous studies of
SN2 reactions,
263 again this apparent contradiction arises because,
while the association of ions in ionic liquids shares some features
with ion pairing in molecular solvents, there are also significant
differences.
The related SN1 ionizations of 2-chloro-2-methylpropane in
[C2C1im][PF6] and water have been modeled using molecular
dynamics simulation.332 It is interesting see that in water the
reaction proceeded through a contact ion pair to dissociated ions.
In the ionic liquid there was no evidence for contact ion pairing;
rather the intermediate was a solvent-separated ion pair with, on
average, one ionic liquid cation and one ionic liquid anion
separating the solute ions. This arrangement has been shown
to have the solute ions fully screened from each other.333 These
observations, together with the fact that the calculated activation
energy for the reaction in water was much lower than for the
reaction in [C2C1im][PF6], suggest that the ionic liquid is not an
unusually ionizing solvent.
6.4.6. Type f SN1 [R—X]
þ Reaction. Arenediazonium
{[ArN2]X} salts are highly reactive intermediates that often under-
go substitution via an SN1 pathway. Dissolving [ArN2]X in
[CnC1im]X {n = 2 or 4; X = [PF6] or [BF4]} and heating led to
fluorodediazonization, forming ArF.334 Other ionic liquids,
[C2C1im]X {X = [CF3CO2], [OTf], or [OTs]} gave ArX as the
product. The authors suggested that this was the result of rapid
metathesis of the solute and solvent anions, leading to trapping of
the ionic liquid anion nucleophile.
While this might have been expected for these ionic liquids,
arenediazonium ions have been shown to react with the supposedly
non-nucleophilic [NTf2]
- anion.335,336 Chiappe et al. found that
when [PhN2][BF4] was dissolved in [C4C1im]Br, PhBr was formed.
However, when [PhN2][BF4] was dissolved in [C4C1im]Br/-
[C4C1im][NTf2] (1:2, 1:1, or 3:1), only products of the reaction
with [NTf2]
- were observed, yet when the same reaction was
conducted in water with a 1:1 [C4C1im]Br/[C4C1im][NTf2]
mixture, the products were bromobenzene and phenol. This suggests
that there is a radical change in the relativenucleophilicities ofBr- and
[NTf2]
- on going from the aqueous to the ionic liquid system. On
the basis of mass spectra of the [C4C1im]Br/[C4C1im][NTf2]
mixture, the authors proposed that this arises from the nucleophilicity
of the bromide ion being sufficiently suppressed by its strong
interaction with the [C4C1im]
þ cation of the ionic liquids for it to
become less nucleophilic than [NTf2]
-. Interestingly, it was only for
Br- that no halogenation was observed; in [C4C1im]Cl/[C4C1im]-
[NTf2] 20% chlorination was observed, and in [C2C1im]I/-
[C2C1im][NTf2] iodobenzene was the only product.
6.4.7. Ionizing vs Dissociating Solvents.The formation of
free solvated ions from a hypothetical molecular species can be
represented as occurring in two steps: ionization and dissociation
(Scheme 10).115 Early in the development of ionic liquids as
solvents for synthesis, it was widely supposed that they would be
especially highly ionizing solvents and that they would promote
the formation of charged species. The survey of nucleophilic
substitutions above suggests that while it is generally true that
they are ionizing solvents, this behavior is in no way extreme and
arises as a result of the specific solvent-solute interactions that
are occurring in solution, rather than as a result of their ionic
nature. However, these results do suggest that once ions have
Scheme 9 Scheme 10

AE dx.doi.org/10.1021/cr1003248 |Chem. Rev. XXXX, XXX, 000–000
Chemical Reviews
REVIEW
trifluoroacetates yielded the appropriately substituted R-methyl-
styrene in dry [C4C1im][NTf2], but significant amounts of
cumyl alcohols were formed when water was present. Kinetic
investigations confirmed a carbocation-forming E1 mechanism
for these reactions. Further confirmation of carbocation inter-
mediates could be obtained through the formation of products
that derive from rearranged carbocations, for instance, in the
reactions of R-keto triflates.
Dyson and Daguenet344 investigated the decomposition of the
chloride-templated metallocage [(Ni(atu)2)4Ni2Cl][BF4]3 (Hatu =
aminidothiorea). This decomposition leads to the release of chloride
ions into solution, and in molecular solvents it had previously been
noted that the equilibriumwas strongly influenced by the solvation of
the chloride ion.345 This equilibrium was studied in the ionic liquids
[cat][NTf2] {where [cat] = [C4C1im], [C4C1C1
2im], [C4C1pyrr],
[C4pyr], or [C5(C2)3N]}. Although significant differences in the
enthalpy of the reaction were seen, ranging from 4.32 kJ mol-1 for
[C4pyr][NTf2] to 16.8 kJ mol
-1 for [C5(C2)3N][NTf2], no simple
relationship between this value and any available solvent polarity
scale could be found. However, when [(HO)2C2C1im][NTf2] was
used, the decomposition proceeded to completion, suggesting the
importance of hydrogen bond donation from the ionic liquid to the
chloride ion, and the trend of increasing chloride ion solvation,
[C5(C2)3N]
þ < [C4C1pyrr]
þ < [C4C1im]
þ e [C4C1C1
2im]þ <
[C4pyr]
þ
, [(HO)2C2C1im]
þ, was proposed. The study of the
entropic effects in the ionic liquids suggested that not only do they
arise from rearrangements of the first solvation sphere around the
solutes, but that correlatedmotion of the ions of the bulk ionic liquids
is also involved.
6.6. Additions
The reaction of 4,4-bis(dimethylamino)benzophenone
(Michler’s ketone, MK) and tetracyanoethene (TCNE) has very
interesting solvent-dependent behavior. In nonprotic molecular
solvents an electron donor-acceptor complex, MK-TCNE, is
formed, but in protic solvents an addition reaction occurs to give a
zwitterionic compound.346 In the rigorously dried ionic liquids
[CnC1im][NTf2] (n = 2, 4, or 6), [C4C1im][PF6], and
[C4pyr][NTf2], theMK-TCNE complex was observed in the first
instance.347 The νmax in the visible spectrum of these solutions,
together with those in aprotic molecular solvents, could be corre-
lated with an LSER including contributions fromR, β, andπ* (R2 =
0.859, eq 17). This was noticeably different from the results for only
the molecular solvents, for which a better correlation could be
achieved using β and π* alone.336 The MK-TCNE species was
unstable in the ionic liquids on the time scale of a few hours, yielding
the radical ions [MK]þ and [TCNE]-. In molecular solvents this
is only seen in highly ionizing solvents, such as 1,1,1,3,3,3-hexa-
fluoropropanol. Even in trifluoroethanol these ions are only seen in
a mixture with the zwitterions. This strongly suggests that the ionic
liquids are acting as highly ionizing solvents in this case. In
[C4C1im][N(CN)2] an unknown product was formed that ap-
peared to be the same as that formed in other very high β solvents,
such as DMSO. When the ionic liquids were wet or contaminated
with other nucleophilic impurities, the zwitterion was formed.
νmax  10
-3
¼ 16:605 þ 1:043Rþ 3:018β- 2:351π ð17Þ
6.6.1. Electrophilic Addition. Electrophilic addition across a
double or triple carbon-carbon bond is the characteristic reaction
of alkenes and alkynes. In an attempt to find replacements for the
commonly used halogenated organic solvents for these reactions,
the bromination of alkenes was investigated in [C4C1im]X {X = Br,
[PF6], or [BF4]}.
348 In [C4C1im][PF6] and [C4C1im][BF4], the
reactions gave mixtures of syn and anti products and the likely
electrophile was identified as Br2 itself, as in molecular solvents.
When the rates of ICl and IBr additions were compared to those in
CH2Cl2, they were found to be appreciably faster. Applying the
Hughes-Ingold rules to these reactions in ionic liquids suggests a
mechanismwith an ionic intermediate. In [C4C1im]Br with Br2, the
actual brominating agent was found to be Br3
- and the bromina-
tions were anti-stereospecific. When Cl2 was used as a chlorinating
agent in [C4C1im]Br, 1,2-bromochlorides were produced, which
was again taken as an indication of the presence of a trihalde
electrophile, BrCl2
- in this case. The regioselectivities of these
reactions were found to vary with the particular trihalide halogenat-
ing agent and the structure of the alkene or alkyne. Using the same
preprepared trihalide salt as a solute in [C4C1im][PF6] hadno effect
on the regioselectivity of most of the reactions investigated, indicat-
ing that there was no change of intermediate in the different ionic
liquids. This intermediate was identified as being similar to that
previously proposed to explain the kinetics of the bromination of
alkynes in ionic liquids,349 in which early attachment by a halide
nucleophile occurs on a 1:1 unsaturated alkene-halogen, or
alkyne-halogen, π-complex (Scheme 13).
It might, at first glance, appear odd that these reactions do not
pass through an ionic bromonium-type intermediate in the ionic
liquids. However, that mechanism begins with Br2 as the
brominating agent, as in [C4C1im][PF6] and [C4C1im][BF4].
However, in [C4C1im]Br it is likely that Br3
- and Br- are the
only bromine-containing species present in kinetically significant
concentrations. Consequently, the absence of Br2 means that
any reaction mechanism that requires its presence will be
unavailable.
A kinetic study of bromination of alkenes and alkynes in the
ionic liquids [C2C1im][NTf2], [C4C1im]X {X = Br, [NTf2], or
[PF6]}, [C6C1im][NTf2], [C2C1C1
2im][NTf2], and [C4pyr]
[NTf2], using [C4C1im][Br3] as the brominating agent, showed
little variation and no systematic trend in the rates of these
reactions in the ionic liquids, all of which were faster than the
same reactions in 1,2-dichloroethane.350 However, when
[(C4)4N][ICl2] was used as the halogenating agent, definite
effects of changing the ionic liquid were seen. While it was
recognized that there were many effects on the changing rates of
these reactions in the ionic liquids, the hydrogen bond donor
Scheme 13

AF dx.doi.org/10.1021/cr1003248 |Chem. Rev. XXXX, XXX, 000–000
Chemical Reviews
REVIEW
ability of the ionic liquid cations was indicated to be an important
and readily identifiable one, probably due to the stabilization of
the emerging Cl- ion during the activation process. Finally,
another hydrogen-bonding effect was noted. These ICl2
- reac-
tions lead in the first place to the mixed iodo-chloro addition
product and a Cl- ion. This Cl- ion can act as a nucleophile
toward the initial addition product to give the dichloride. Ionic
liquids with strong hydrogen bond donor cations were found to
suppress the nucleophilicity of the Cl- ion and hence prevent
this reaction. Tetraalkylphosphonium trihalides have also been
used as halogenation agents.351
6.6.2. Nucleophilic Addition. In addition to the solvatochro-
mism of the longest wavelength absorption of merocyanine probes,
the changing absorbances of this band can be used to estimate the
equilibrium constant for the spiropyran-merocyanine equilibrium
(Scheme 3). This is again well correlated with Reichardt’s ET(30)
scale in molecular solvents.352 Also, by irradiating the sample, the
reaction can be shifted to the right, after which the kinetics of the
thermal re-equilibration, which can be thought of as a nucleophilic
addition, can be measured. The relationship of these values to the
nature of the ionic liquid has been shown to be more complex than
that of the solvatochromism. An initial study of 2,3-dihydro-10,30,30-
trimethyl-6-nitrospiro[1-benzopyran-2,20-1H-indole] (seeScheme3,
R=CH3) in a range of [NTf2]
- ionic liquids showed that therewere
no longer simple linear relationships with the ET(30) scale, but both
of these values are affected by the nature of the cation of the ionic
liquids, with stronger hydrogen-bond-donating ions leading to
slower re-equilibrium and a higher equilibrium concentration of
the merocyanine form.139 In subsequent investigations by the same
group, the lack of a linear correlation between the equilibrium
constants and the ET(30) scale was again noted for reactions in
ionic liquids.353 Analysis of the activation parameters for these
reactions showed that the changes in the activation entropies for
the reactions in the phosphonium ionic liquids studied were
significant in determining the observed solvent dependencies. This
suggests that solvent reorganization is occurring on the same time
scale as the reaction in this case. Entropic effects are poorly described
by polarity scales based on solvatochromism, which probably
explains the lack of linear correlation between these results and the
ET(30) scale. Similar results have been reported for a range of
imidazolium ionic liquids by Deng et al.138
6.6.3. 1,4-Conjugate Additions. 1,4-Conjugate additions
are conceptually simple reactions in which a nucleophile adds to a
so-called Michael acceptor, which is activated by the presence of
an electron-withdrawing group (EWG; Scheme 14). However,
they usually require the use of catalysts and unfavored solvents,
such as DMSO, to achieve the reaction. The first reported
conjugate addition in an ionic liquid used [C4N]Br as the solvent
for the reaction of β-oxosulfides of benzothiazole with various
Michael acceptors.354 This allowed the stereoselective synthesis
of spirocyclopropanes with a weaker base, sodium bicarbonate,
than usually required.
Conjugate additions of a range of amine and thiol nucleophiles
to various Michael acceptors in ethylammonium acetate have
recently been reported.355 The interaction of the cation with the
EWG of the Michael acceptor was proposed to explain the
enhanced reactivities observed. However, it should also be noted
that the ionic liquid increased the rates of the reactions of the
thiols and the primary and secondary amines used, but not of the
imidazoles, suggesting that activation of the nucleophile is also an
important contributor. Thus, the cation effect and the anion
effect on these reactions are acting synergistically.
At the same time, conjugate additions of malonodinitrile, and
other reagents, to chalcone were shown to proceed well in pure
ionic liquids, without the addition of any catalyst.356 The authors
explored several different possible explanations for this behavior
and concluded that the most likely reason was the different
dissociation constants of C-H acids in ionic liquids relative to
molecular solvents. However, they subsequently demonstrated
that residual methylimidazole, remaining in the ionic liquid from
its preparation, can act as a base catalyst for these reactions.357
The prolinate ionic liquid [C2C1im][Pro] has also been used for
conjugate additions of cyclohexanone to chalcones.358 The use of
this chiral ionic liquid led to both an increased reactivity
compared to that found in the ionic liquids [C4C1im][PF6]
and [C4C1im][BF4] and significantly improved enantioselectiv-
ity. L-Proline has been used as a catalyst for Michael addition of
ketones to nitrostyrene,359 so it is possible that unreacted L-
proline from the synthesis of the ionic liquid is acting as the
catalyst for the reaction. Another conjugate addition in which
impurities in the ionic liquid were important, this time giving a
detrimental effect, was in the metal {Ni(acac)2
3
2H2O, FeCl3
3
6
H2O, or Yb(OTf)3} catalyzed addition of acetylacetone to
methyl vinyl ketone in [C4C1im][PF6] and [C4C1im][BF4].
360
It was found that Cl- ions in the ionic liquids led to the formation
of stable metal halide complexes that did not catalyze the
reaction. In the same paper it was noted that water in the ionic
liquids did not have any detrimental effect on the reaction.
Another route to enantioselective conjugate additions was
realized by the use of (R,R)-trans-1,2-diaminocyclohexane as a
chiral auxiliary in ionic liquids.361 For the asymmetric addition of
ethyl cyclohexanone-2-carboxylate to methyl vinyl ketone, en-
antiomeric excesses of up to 91% were achieved. It was specu-
lated that the role of the ionic liquid might be to favor enamine
formation, the first step of the reaction. Interestingly, the
addition of a variety of metal complexes only had detrimental
effects on the reaction, presumably because this led to the
diamine preferentially acting as a ligand for the metal center
rather than a catalyst for the reaction.
The ionic liquids [(C8)3C1P][C1CO3] and [(C8)3C1P]
[HCO3] were prepared for use as a combined solvent-catalyst
for 1,4-conjugate additions.362 However, they were such potent
catalysts for the reaction of cyclohexanone and nitroethane that
they were used in a ratio of IL to cyclohexanone of 0.004, with the
reactant being the solvent. The Brønsted acidic ionic liquid
[(HO3S)
4C4pyr][p-CH3PhSO3] has also been used as a catalyst
in acetonitrile (10mol %) for the reaction of indole and a range of
1,4-unsaturated ketones.363
6.6.4. Diels-Alder Reactions. Diels-Alder cycloaddition
reactions were among the first to be studied in ionic liquids.364
This initial study explored the possibility of using the ionic liquid
[C2NH3][NO3] as a substitute for water, which has become a
popular solvent for these reactions. The reactions of cyclopenta-
diene with methyl acrylate and methyl vinyl ketone in this IL
showed a strong preference for the endo product (see
Scheme 15) and an acceleration of the reactions in comparison
to those in nonpolar organic solvents, although the increased rate
and selectivities were not as great as those seen in water.
Scheme 14

AG dx.doi.org/10.1021/cr1003248 |Chem. Rev. XXXX, XXX, 000–000
Chemical Reviews
REVIEW
Nevertheless, the ionic liquid has the advantage that moisture-
sensitive reagents can be used.
The Diels-Alder reaction of cyclopentadiene with methyl
acrylate is well-known to have strong solvent dependencies for both
its selectivity and its rate.115 It has even been used as the basis for a
solvent polarity scale, Ω25C {log(endo/exo)}.
365 The influence of
hydrogen bonding on the endo/exo ratio and rate of this reaction in
ionic liquids was rapidly recognized.183 Hydrogen bond donation
from the cation of the ionic liquid to the dienophile increased the
endo selectivity of the reaction, whilemore hydrogen bond accepting
anions decreased the endo selectivity of the reaction. This led to the
use of a range of protic ionic liquids as solvents for reactions of
cyclopentadiene with either methyl acrylate or dimethyl maleate.366
The endo/exo ratios of these reactions were particularly high when
[C1Him][NTf2] was used as the solvent. A semiempirical compu-
tational model was used to further confirm the hypothesis that
hydrogen bonding of protic imidazolium cations to the dienophile
gives rise to the effects observed in the ionic liquids. Subsequent
kinetic investigations of the rates of the reactions of cyclopentadiene
with alkyl (methyl, ethyl, or butyl) acrylates in a range of imidazo-
lium ionic liquids {[C2C1im][BF4], [C4C1im]I, [C4C1im][BF4],
[C4C1im][PF6], or [C8C1im][PF6]} again confirmed the influence
of hydrogen bonding on the rate of the reaction, but also showed
that the rate slowed significantly with the increased viscosity of the
ionic liquid.367 This had the effect that when the non-hydrogen-
bonding solvent CH2Cl2 (55mol %) was added to [C4C1im][BF4]
(45 mol %) and the mixture used as a solvent for the reaction, the
rate of the reaction increased by 20%.
The effect of the viscosity on the intramolecular Diels-Alder
reaction of (E)-1-phenyl-4-[2-[(3-methyl-2-butenyl)oxy]benzy-
lidene]-5-pyrazolone has also been investigated in a variety of
pyridinium tetrafluoroborate and bis(trifluorosulfonyl)imide ionic
liquids.368 Two similar, but separate, trends were found in the
correlation of the ln k1 with the viscosity of the ionic liquids for each
of the two sets. The authors attributed this difference in the behaviors
of the two sets of ionic liquids to microscopic frictional effects in
which the microviscosity is different from the macroscopically
observed value.
Dyson et al.369 conducted an extensive study of the selectivities and
rates of the reaction of cyclopentadiene with methyl acrylate in 30
different [NTf2]
- ionic liquids. They analyzed the results for the
endo/exo selectivity by reference to the IR spectra of the ionic liquids
and using the polarity scales derived frommeasurements of theNMR
spectra of a series of probe molecules to derive a predictive equation
for the selectivity of this reaction. Their kinetic data suggested that the
selectivity was kinetically controlled, with the rate of formation of the
exo product being unaffected by the presence of the ionic liquid,
whereas higher overall selectivities were accompanied by more rapid
formation of the endo product. This arose from a lowering of the
relative energy of the transition states leading to the endo products,
while the transition states yielding the exo products are not
significantly influenced by the choice of cation, at least in the cases
studied. They concluded that hydrogen bonding, while important,
cannot alone satisfactorily account for the observed selectivities and
that any property of the ionic liquid that favored interaction of its
cation with the activated complex leading to the endo product favored
greater endo selectivity and vice versa. Important cation properties
identified included hydrogen-bond-donating ability, steric bulk, with
long substituents on the cation leading to lower selectivities, pre-
sumably due to unfavorable steric interaction between the TSq and
the cation, chemical hardness, with strong electrostatic association
between the ionic liquid ions leading to less interaction between the
ionic liquid and the TSq, the presence of a low-energy lowest
unoccupied molecular orbital on the cation, and the fact that
polarizable cations can lower the energy of the endo TSq more than
that of the exo TSq. In a subsequent study they also showed that the
application of pressure further affects the rates of the Diels-Alder
reaction.370
In an attempt to derive LSERs that could be used to design ionic
liquids to be solvents for a wide range of Diels-Alder reactions, Bini
et al.371 investigated the selectivities and rates of the reactions
between cyclopentadiene and three dienophiles (acrolein, methyl
acrylate, and acrylonitrile) in ionic liquids and molecular solvents.
They succeeded in deriving an LSER for both properties of all three
reactions. That these LSERs could be achieved including both
molecular and ionic liquids shows that there are no qualitative
differences in the ways in which these solvents affect the reactions.
However, no two were precisely the same. For instance, when
investigating both the selectivities and rates of these reactions, it was
found that the solvent hydrogenbonddonation abilitywas important
in the reactions of acrolein and methyl acrylate, which both have
strong hydrogen bond acceptor sites, but not of acrylonitrile, which
does not. This demonstrates that the solvent effects that are seen in
theseDiels-Alder reactions are a functionof both the solvent and the
solute and that the properties of both should be considered when
making predictions. Also, the R2 values for the correlations of the
LSERs derived from the kinetic data were much poorer than those
found for the selectivities of the same reactions, suggesting that the
combination of factors leading to the changes in the rates of the
reactions are more complex than those leading to changes in
selectivity and are more sensitive to effects not covered in the
analysis, such as subtle fluctuations in the reaction conditions. This
also partly arises from the selectivities being the result of the ratio of
the reaction rates for the endo and exo products whereas the overall
rate of the reaction is the sum of the rates for the two products. The
consequence of this is that some effects are canceled out in the
selectivities of the reactions. For instance, viscosity is present in
the correlation for ln k2 for acrolein, whereas it is absent in the
explanation for its selectivities. This experimental study was part-
nered with a theoretical investigation.372 This focused on the two
limiting cases with the highest endo/exo selectivities, acrolein and
methyl acrylate in [C4Him][NTf2], compared to the lowest selec-
tivity case, acrylonitrile in [C4C1C1
2im][NTf2]. Again both the
abilities of the dienophile and the ability of the cation to enter into
hydrogen bonding were found to be the important determinants of
the strengths of the interactions between these compounds and
hence the selectivities of the reactions.
Contrary to statements found in standard texts that the solvent
effects on Diels-Alder reactions are well understood and
relatively straightforward, perhaps the most overwhelming con-
clusion to be drawn from the studies presented above is that
multiple factors influence differently both the rates and product
selectivities of the Diels-Alder reactions and that these factors
Scheme 15

AH dx.doi.org/10.1021/cr1003248 |Chem. Rev. XXXX, XXX, 000–000
Chemical Reviews
REVIEW
differ depending upon the reagents.370,371 Conclusions drawn on
the basis of experiments taking a multiparameter approach are
more reliable than those with a single variable, particularly when
that variable may be acting as a proxy for others that have not
beenmeasured. It is possible tomake rationalizations of observed
reactions using quantitative measurements and comparison with
multiple physicochemical parameters, but these are poor pre-
dictors for the effects of solvents on even closely related reac-
tions. A case in point is the intramolecular Diels-Alder reaction
of (E)-hexa-3,5-dienyl acrylate, which shows no strong rate effect
from the hydrogen bond donor property of the ionic liquid.373
The interactions between the cation and both the diene and
dienophile have been exploited in the asymmetric aza-Diels-
Alder reaction of Danishefsky’s diene with chiral imines.374
When ionic liquids prepared with cations derived from N-
methylephedrine were used as the solvent for these reactions,
moderate diastereoselectivities could be achieved. It was neces-
sary to have a cation that could interact with both reagents to
achieve the alignment in the activated complex that leads to the
preferred diastereomer (Figure 11).
The organocatalyst (5S)-5-benzyl-2,2,3-trimethylimidazolidin-4-
one has been used as the catalyst for the Diels-Alder reaction
between cyclohexadiene and acrolein in the ionic liquids
[C4C1im][PF6], [C4C1im][SbF6], [C4C1im][BF4], and [C4C1-
im][OTf].375 Both the endo/exo selectivity and the enantioselec-
tivity of the endo adduct were reported. The endo/exo ratio was 17:1
for all of the ionic liquids used (cf. 14:1 inCH3CN), but while the ee
of the endo adduct was approximately 90% in the hydrophobic ionic
liquids (cf. 94% inCH3CN), it was zero in the two hydrophilic ionic
liquids. Also the yield for the reaction in these two ionic liquids was
only 5% or 7%, compared to 70-76% in the hydrophobic ionic
liquids (cf. 82% in CH3CN). Clearly, water plays an important role
in these reactions. When water was deliberately added to the
[C4C1im][OTf] ionic liquid up to 10%, no effect on either outcome
of the reaction was seen. However, when this was increased to 20%,
both the yield and the enantioselectivity of the endo adduct
increased. It was concluded that water is needed for the hydrolysis
of the iminium ion formed during the catalytic cycle and that the
stronger interactions between the [OTf]- and [BF4]
- ions and
water prevented this fromoccurring. Similar behavior, althoughwith
a positive outcome for giving the required product, has been seen in
perruthenate-catalyzed reactions in ionic liquids.94
6.6.5. Oxidative Addition. The oxidative addition reaction
is encountered almost universally in mechanistic interpretations
of synthetically important, metal-mediated processes. The oxi-
dative addition of methyl iodide to Vaska’s complex {[(PPh3)2-
(CO)ClIr]; Scheme 16} has been studied in [C4C1im][OTf],
[C4C1im][NTf2], and [C6pyr][NTf2].
376 In molecular solvents
this reaction is generally agreed to proceed by an initial rate-limiting
formation of the pentacoordinate [(PPh3)2(CO)Cl(CH3)Ir]
þ.
Consequently, it was thought that it might be accelerated in
the ionic liquids. We found the observed rate constants under
pseudo-first-order conditions with a large excess of CH3I to vary
in the order toluene ≈ [C6pyr][NTf2] < [C4C1im][OTf] ≈
[C4C1im][NTf2] < DMF. When the concentration of CH3I was
changed, complex changes in the observed rates were seen in the
ionic liquids. This was interpreted as arising from the involve-
ment of another ligand species, possibly Cl- or free-base impuri-
ties left from the ionic liquids’ syntheses.
6.6.6. Metal Alkyl and Aryl Addition Reactions. Organo-
metallic compounds are often used in stoichiometric quantities as
activated reagents to facilitate the addition of one group to a
compound. Probably the most well-known of these is the Grignard
reaction, which is the addition of an alkyl- or arylmagnesium halide
(most often bromide) to an aldehyde or ketone to form a secondary
or tertiary alcohol, respectively (Scheme 17). Grignard reagents are
strong bases and as such are unsuitable for use with unprotected
imidazolium salts. The commercially available Grignard reagent
PhMgBr in THFwas found to give clear solutions in [C14(C6)3P]X
{X = Cl, Br, or [NTf2]} which were stable for over one month.
377
Reaction with acetone or benzaldehyde gave the expected alcohol
products after hydrolysis. Unfortunately, attempts to form the
PhMgBr directly in the ionic liquids were unsuccessful.
When iodoethane was reacted with magnesium metal in
[C4pyr][BF4], it led to the formation of the Grignard reagent.
378
However, when it was reacted with benzaldehyde, the product was a
mixture of 1-iodo-1-phenylpropane and 1,2-diphenylpinacol. The
former can be imagined to arise from a nucleophilic attack of iodide
on the expected Grignard product 1-phenyl-1-propanol, which was
confirmed by its addition to the reaction mixture, leading to the
iodide. For the latter a ketyl radical intermediate was proposed.
Addition of pyridine or triethylamine led to the formation of
1-phenyl-1-propanol. When expanding the scope of the reaction,
the authors found that ketones were unreactive, which is an order of
reactivity different from that found inmolecular solvents, which they
proposed resulted from the Grignard reagent having a polymeric
structure in the ionic liquid. It is interesting to note that not only was
the reaction unsuccessful when bromoethane was used, but when
Br- was present in the ionic liquid, this prevented the reaction of
iodoethane. While the lack of success of the reaction of bro-
moethane could arise from a failure of the initial oxidative addition
of the bromoalkane to the magnesiummetal, the suppression of the
reaction by free bromide suggests that interference with the Schlenk
equilibrium may be occurring.
The first of these organometallic additions to be demonstrated
in ionic liquids was the addition of the allyl group of tetraallyl-
stannane to the CdO group of benzaldehyde in good yields in
Figure 11. Proposed interactions between the chiral cation of an ionic
liquid and the reagents in the activated complex of a diastereoselective
aza-Diels-Alder reaction.374
Scheme 16
Scheme 17

AI dx.doi.org/10.1021/cr1003248 |Chem. Rev. XXXX, XXX, 000–000
Chemical Reviews
REVIEW
[C4C1im]X {X = [PF6] or [BF4]}.
379 No specific investigations
into the possibility of forming NHC-derived products were
made, but it was noted that in some cases there was greater loss
of benzaldehyde than tetraallylstannane. Allyldiisopropoxybor-
ane has been shown to produce the same products in similar
yields in [C4C1im]X {X = Br or [BF4]}, [C3C1im][BF4], and
[C2NH3][NO3].
380
The need to preform the allyl metal reagent was circumvented
by using allyl bromide as the reagent together with indium metal
in [C4C1im]X {X = [NTf2], [PF6], or [BF4]}.
381 The yields in
[C4C1im][NTf2] were particularly poor, suggesting that the HF
that would have been generated in these reactions, which were
conducted open to the air, may have been assisting the activation
of the metal surface. Unidentified impurities were noted in the
product mixtures that could indicate the effects of NHC forma-
tion in these systems. Tin and zinc have been used in a similar
way in [CnC1im][BF4] (n = 2 or 4).
382
Attempts to extend these methodologies beyond the allyl group
havemetwith limited success. The reaction of benzaldehydewithR-
bromoesters (the Reformatsky reaction) in ionic liquids have given
only moderate yields, except for the most highly activated
reagents.383 Reactions using diethylzinc led to the formation of
zinc-NHC complexes in [C4C1im]Br and [C4C1im][BF4] with
Br- contamination, and reasonable yields of addition products
could only be achieved in very pure [C4C1im][BF4].
384 The same
reaction in [C4C1im][PF6] gave reduction of the aldehyde starting
material to its alcohol. Clearly, a great deal more work is required for
these reactions to become synthetically useful.
6.7. Acid-Catalyzed Reactions
Acid catalysis is one of the most important technologies used
in the chemical industry, and it is applied throughout its value
chain. However, it presents a number of environmental problems
associated with the strong acids used and the waste that they
generate. HF, for instance, is volatile, corrosive, and highly toxic;
AlCl3 is also toxic and corrosive and produces large amounts of
HCl on contact with water. All strong acids produce large
amounts of waste, causing significant treatment problems. This
has led several groups to explore the possibility of using ionic
liquids as more environmentally sustainable alternatives.
6.7.1. Electrophilic Aromatic Substitutions. Clean elec-
trophilic substitution reactions remain one of the most important
targets of the synthetic chemical industry. Conventional methods
for these usually require strong acids, such HF or AlCl3, as catalysts.
Friedel-Crafts alkylations of benzene using halogenoalkanes were
among the earliest reactions to be investigated in chloroaluminate
ionic liquids.385 It was clearly demonstrated that acidic [X(AlCl3) >
0.5] ionic liquids were required to act as a combination of solvent
and catalyst for the reaction and that the reaction proceeded via the
dissociated carbonium ions. This was followed by interest in arene
alkylation with alkenes, which are more desirable starting materials
for industrial-scale reactions. These reactions are rarely completely
selective to a single product. Since the alkylated product is more
reactive than the parent arene, overalkylation is a common problem.
This can be minimized by the use of a large excess of the arene. It is
less easy to control the position along the alkyl chain at which it
attaches to the arene ring. This arises because the carbocations
formed during the reaction rearrange to more substituted forms,
which are more stable, leading to a number of isomeric products for
all but the simplest alkenes.
For the alkylation of anthracene with 2-chloropropane, both the
overall yield and the selectivity to 2-isopropylanthracene reached a
maximum when the composition of the [C2C1im]Cl-AlCl3 ionic
liquid reached X(AlCl3) = 0.67. In compositions of X(AlCl3)e 0.5
no reaction occurred,386 nor was any reaction observed in
[C2C1im]Cl-ZnCl2 [X(ZnCl2) = 0.67] and only the slightest
amount in [C2C1im]Cl-FeCl2 [X(FeCl3) = 0.67]. It was also
shown that the greatest yield and selectivity were found for reactions
at 30 Cand under these conditions for 4 h, suggesting that there is a
balance to be drawn between conditions that lead to insufficient
reaction and those that lead to overreaction.
For the alkylation of benzene with 1-dodecene in
[C4C1im]Cl-AlCl3 with a fixed acidity [X(AlCl3) = 0.67], the
primary factor influencing the conversion of the 1-dodecene was
the temperature of the reaction, with quantitative conversions
being achieved above 50 C.387 Attempts to repeat reactions in a
batch reactor were severely hampered by the exposure of the
ionic liquid tomoisture and its consequent deactivation. This was
somewhat ameliorated by the use of a continuous-flow reactor.
It has been demonstrated that the addition of HCl to a
[C2C1im]Cl-AlCl3 [X(AlCl3) = 0.67] ionic liquid leads to an
increased selectivity to the 2-substituted isomer when the reac-
tion is carried out to completion and under conditions that avoid
polyalkylation (i.e., large excess of benzene).388 Similar results
have been found when using either 1- or 2-methylnaphthalene as
the arene.389 This clearly suggests that the reaction is proceeding
via a free carbocation, which is rearranging to the more stable
highly substituted form, for example, Scheme 18. It is well
established that, unless specially treated, chloroaluminate ionic
liquids will contain protons derived from hydrolysis with adven-
titious water.1 Indeed, these reactions require the presence of
some protons to generate the necessary carbocation for the
reaction. Consequently, small amounts of adventitious water lead
to the reaction being possible, whereas larger amounts lead to the
deactivation of the ionic liquid. This makes it very difficult to
compare results obtained from ionic liquids that have been
handled differently because of the varying degree of contact
between the ionic liquids and water.
Detailed kinetic investigations of the reaction of cumene
(isopropylbenzene) withpropene in [C2C1im]Cl-AlCl3 [X(AlCl3) =
0.67] confirmed that the various products (di-, tri, and tetra-
isopropylbenzene) result from a series of successive alkylation
reactions.390 These also showed that it was necessary to take into
account the solubilities of the products to be able to explain the
observed selectivity for the reaction and to fit the kinetic model to
the data. To do this, and because of the difficulties faced in their
experimental determination, the authors used the COSMO-RS
model391 to predict the relative solubilities of the products. This
showed that the more alkylated products were less soluble in the
reactive ionic liquid phase, leading to an improved selectivity for
the monoalkylated product.
It is well established that varying the composition [X(AlCl3)] leads
to chloroaluminate ionic liquids of different acidities.1 This was
demonstrated to affect the yield and selectivity in the alkylation of
diphenyl oxidewith1-dodecene.392 Ionic liquidwithX(AlCl3)e0.55
did not show any catalytic activity in the alkylation, those with
X(AlCl3) > 0.60 gave increasing amounts of byproduct (e.g.,
Scheme 18

AK dx.doi.org/10.1021/cr1003248 |Chem. Rev. XXXX, XXX, 000–000
Chemical Reviews
REVIEW
Triflic acid itself has also been used as the promoter for the
adamantylation of toluene by 1-adamantyl halides (Cl or Br) or
1-adamantyl alcohol in [C4C1im][OTf].
401 In the ionic liquid
excellent yields of, almost exclusively, p-tolyladamantane could be
achieved, while in the comparable solvent 1,2-dichloroethane
m-tolyladamantane was the major product (ca. 40%), with p-
tolyladamantane and adamantane being produced in roughly equal
amounts. These data were rationalized by proposing that a more
complete formation of the adamantyl cation occurred in the ionic
liquid and that the transition state for the reaction was later in the
reaction coordinate and with more benzenium ion character.
Arene acylation is also an industrially important reaction. It is
in many ways a simpler reaction than alkylation. The reduced
reactivity of the product means that overacylation is not a
concern. The reaction requires the use of an activated acylating
agent, usually an acid chloride or acid anhydride. A Lewis acid,
commonly AlCl3, is required to generate the reactive acylium ion,
which then attacks the arene, leading to the substitution of a
proton on the arene. However, this reaction is not without its
problems. The AlCl3 is required in large excess to achieve useful
conversions, and it forms an adduct with the carbonyl oxygen of
the product. To liberate the product from this adduct, water is
added to hydrolyze the AlCl3, consuming it and preventing its
reuse. Early investigations of this reaction in chloroaluminate
ionic liquids showed that they were effective catalysts/solvents,
but that these problems also remained.1,389 Consequently, most
of the focus for research in this area has been to find alternative
systems that can overcome this need to destroy the catalyst to
isolate the reaction products.
It has been shown that [C4C1im]Cl-FeCl3 [X(FeCl3) = 0.67] is
a more effective catalyst/solvent for the reaction of benzene with
benzoyl chloride (Scheme 19) than [C4C1im]Cl-AlCl3 [X(AlCl3)
= 0.67], giving higher yields of benzophenone in shorter times.402
The same study showed that [C4C1im]Cl-ZnCl2 [X(ZnCl2) =
0.67] was almost ineffective as a catalyst/solvent for the same
reaction. The authors note that there is no correlation with the
acidity of the ionic liquids, which are in the order [C4C1im]Cl-
AlCl3 > [C4C1im]Cl-ZnCl2 > [C4C1im]Cl-FeCl3.
271 The ben-
zophenone product was isolated from the [C4C1im]Cl-FeCl3 ionic
liquid by simple extractionwith cyclohexane, and the ionic liquidwas
reused. However, on the fourth run the yield collapsed from 85-
95% to less than 30%. This was attributed to the accumulated loss of
FeCl3 into the product stream as a benzophenone-FeCl3 adduct.
Detailed spectroscopic measurements confirm the formation of a
similar acetophenone-FeCl3 adduct in [C4C1im]Cl-FeCl3 ionic
liquids.403 This study also confirmed that the acetylium ion
{[CH3CO]
þ
} is the key intermediate in the reaction in this ionic
liquid, as for the related chloroaluminate ionic liquids.
[C4C1im]Cl-InCl3 [X(InCl3) = 0.67] has been used in an
attempt to circumvent the problem of loss of catalyst into the
product stream.404 Since these ionic liquids are not readily
hydrolyzed, an aqueous workup is possible for the reactions of
benzoic anhydride with various arenes in the ionic liquids, which
are then recovered and reused. However, subsequent detailed
evaluation of the benzoylation of anisole in [C4C1im]Cl-InCl3
showed that even in this system the product was contaminated
with an average 0.36 ppm indium after each reaction, which
would require further separations in any commercial process and
lead to the eventual deactivation of the catalyst/solvent.405
Several groups have investigated the use of catalysts dissolved in
ionic liquids in an attempt to solve the problem of catalyst loss in
Friedel-Crafts acylations.Metal triflate salts dissolved in a variety of
ionic liquids have shown promise as catalysts for the benzoylation of
arenes with benzoyl chloride.406,407 Activities, as expressed through
conversion at fixed reaction time, varied with both the particular
metal salt used and the nature of the ionic liquid. As one might
expect of an acid-catalyzed process, ionic liquids with more basic
anions {e.g., [BF4]
- or [OTf]-} provided lower conversions than
those with less basic anions {e.g., [PF6]
- or [NTf2]
-
}. The cation
was also shown to impact the observed conversions, possibly due
to solubility effects. However, in all cases where recycling of the
catalyst solutions was attempted, diminishing conversions were
seen. When Bi2O3 was used as the catalyst in [C2C1im][NTf2],
the formation of [C2C1im][BiCl4] during the reactions was clearly
demonstrated. This occurs by the HCl byproduct of the acylation
reacting with the bismuth complex to give the unreactive
[BiCl4]
-.411 An analogous mechanism is likely to be occurring in
the other systems.
Since acylation with acid anhydrides does not lead to the
formation of HCl, the formation of inactive metal chloride
complexes cannot be a mechanism for catalyst deactivation.
However, detailed kinetic studies have demonstrated that cata-
lyst deactivation also occurs in these reactions.408 This elegant
study showed that the catalysis was by a Brønsted acid, generated
from a series of ligand exchange reactions, which was subse-
quently deactivated by complexation with the product. The use
of a biphasic ionic liquid/scCO2 (sc = supercritical) system
extended the active lifetime of the catalyst solution {In(OTf)3 in
[C4C1
4pyr][NTf2]}, probably by more effective separation of
the product from the catalyst.409 However, deactivation was not
totally prevented. An attempt at using zeolite catalysts in ionic
liquids for this reaction also led to catalysis by a Brønsted acid
liberated from the zeolite by ion exchange with the ionic liquid
and subsequent deactivation of the catalyst solution.410
An in situ generated 2,20-dihydroxy-1,10-binaphthyl (BINOL)
titanium(IV) complex in [C2pyr][BF4] or [C2pyr][CF3CO2] ionic
liquids has been used for the reaction of ethyl glyoxylate with
aromatic amines.411 Good yields and enantioselectivities were
achieved for this reaction, which only diminished a little after
successive runs of the reaction. These results suggest that the metal
complex is indeed the catalyst for the reaction and that its deactiva-
tion is lessened by the presence of this ligand environment.
The electrophilic nitration of aromatics is also industrially impor-
tant. The earliest studies in ionic liquids used acidic compositions of
[C2C1im]Cl-AlCl3withKNO3.
412 Itwasproposed that thenitrating
agent was in situ generated [NO2]
þ. However, this reaction was
accompanied by the product isolation difficulties that arise with the
use of chloroaluminate ionic liquids. [NO2]
þ can be delivered
directly as the salt [NO2][BF4]. When it was used in [C2C1im]Cl
for the nitration of toluene, NO2Cl was formed, which evaporated
from the reaction mixture when the apparatus was not sealed.413
Rapid (compared to nitration) ion metathesis was also seen in
[C2C1im]Cl-AlCl3, [C2C1im][PF6], and [C2C1im][OTf]. Of
course, the metathesis of [NO2][BF4] with [C2C1im][BF4] is a nil
reaction. No useful yields of nitrotoluenes could be achieved in any of
these ionic liquids. This was at least partly due to the formation of
nitrated imidazolium cations in the ionic liquids.
Scheme 19

AL dx.doi.org/10.1021/cr1003248 |Chem. Rev. XXXX, XXX, 000–000
Chemical Reviews
REVIEW
Silver, copper, iron, or ammonium nitrates combined with either
acetic anhydride or trifluoroacetic acid anhydride have also been
shown to act as arene nitration agents.413-415 It was proposed that
the ortho/para selectivity suggested that thenitration occurred via the
in situ formation of acyl nitrates. The formation of nitrated imida-
zolium cations was seen again, and it was recommended that
nonimidazolium ionic liquids, such as [C4C1pyrr]
þ or [(iC3)2-
C2NH]
þ salts, are preferred.413,415 HNO3/acetic anhydride has
been used to investigate the effect of the ionic liquid anion on the
nitration of phenols.416 Although little difference was seen in the
reactivities of activated substrates, greater differences were seen for
less reactive substrates (e.g., halobenzenes) in [C4C1pyrr][OTf] or
[C4C1pyrr][NTf2], for which significantly greater yields at a fixed
time were achieved in [C4C1pyrr][NTf2] than in [C4C1pyrr]-
[OTf]. No reaction was observed under the same conditions in
CH2Cl2. The results in the former ionic liquid were similar for these
deactivated areneswhen [NO2][BF4] was used as the nitrating agent
(with no addedAc2O). These results were used together as evidence
of nitration by [NO2]
þ itself, although its reactivity is moderated
through solvation by the ionic liquid anion. Similar results were seen
in reactions in [C4C1im][PF6], [C4C1im][BF4], and [C4C1C1
2im]-
[BF4] that agree with this interpretation, although these authors
speculated that the differences arose from solubility differences of
Ac2O in the ionic liquids.
417
Since theonly reactionbyproduct iswater, the use of aqueous nitric
acid for nitration has the potential to lead to an environmentally
benign synthesis of nitroaromatic compounds.However, strong acids,
such as concentratedH2SO4, are required in commercial processes to
give acceptable reactivities. These in turn lead to acidwastes that need
to be disposed of, reducing the environmental efficiency of the
processes. Consequently, a number of authors have attempted these
reactions in ionic liquids.WhenHNO3alonewas used as thenitrating
agent, truly dramatic ionic liquid anion effects were seen
(Scheme 20).418 In ionic liquids with triflate or hydrogen sulfate
anions, the expected nitrated arenes were produced. However, when
halide ionic liquids were used, halogenated areneswere produced, the
acid acting to oxidize the halide to the hypohalous acid halogenating
agent. The same was seen when HCl was the acid used in
[C4C1im][NO3]. Finally, inmethanesulfonate ionic liquids the nitric
acid acted as an oxidizing agent for toluene to yield benzoic acid.
Usingmethanesulfonic acid dissolved in [C4C1im][NO3] yielded the
same result. These results show that the origin of various ions in the
reaction mixture has no impact on the reactivity of that mixture. This
suggests that complete dissociation of the dissolved acid occurs and
the product mixture has a statistical distribution of ions.
[(HO3S)
nCnC1im][OTf] (n = 3 or 4) have been used to
supply catalysts for nitrations with 62% nitric acid.419 At 5%
loadings (arene:nitric acid = 2:1), these ionic liquids provided
good conversions for activated and moderately deactivated
arenes, but not for the strongly deactivated nitrobenzene. The
biphasic system gave easy separation of the product and recycling
of the aqueous layer, which contained both the ionic liquid and
unreacted nitric acid. Unfortunately, conversions rapidly de-
creased with successive cycles. However, this may simply be
due to the decreasing concentration of nitric acid in the solution.
The, by then well-known, propensity of imidazolium-based ionic
liquids to be nitrated led Liu et al. to prepare a range of
[(HO3S)
nCn(Cm)3N][HSO4] (n = 3 or 4; m = 1, 2, or 4) ionic
liquids for use in these reactions.420 TheHammett acidities of the
ionic liquids in CH2Cl2 were compared and their solubilities in
several nonpolar molecular solvents measured. Significant differ-
ences in reactivities were seen for moderately deactivated arenes,
but all of the ionic liquids gave increased yields and para-
selectivities compared to those seen for the same reaction in
their absence. The order of the yields was consistent with the
miscibility of the ionic liquid with the arene, rather than the
acidity of the ionic liquid. The authors proposed that this was
because the ionic liquids were behaving as phase transfer catalysts
and allowing the arene to come into contact with the charged
[NO2]
þ nitrating agent.
6.7.2. Isobutane Alkylation. The acid-catalyzed alkylation of
isobutane with 2-butene is conducted by the petrochemical industry
on amassive scale.421 Although at first sight simple, this reaction has
a mechanism that passes through a succession of carbocations that
readily rearrange to give amixture of octanes (Scheme 21), someC8
olefins, and heavier alkanes (>C8, typically 15-20%) and cracking
products of these alkanes (C5-C7 alkanes, typically ca. 20%). The
products of the reaction are added to automobile fuels to increase
their octane rating. The performance of the mixed alkylate product
depends upon its precise composition. Better performance is given
by mixtures with a higher overall proportion of octanes and within
that a greater trimethylpentane (TMP) to dimethylhexane (DMH)
ratio. The conventional acids used for this process are HF andmore
recently H2SO4.
In acidic compositions the ionic liquids [CnC1im]X-AlCl3 {n =
4, 6, or 8; X = Cl, Br, or I) are highly effective catalysts/solvents for
this reaction.422 The highest olefin conversions were found for the
[C8C1im]Br-AlCl3 ionic liquid. Comparison of the best ionic
liquid with H2SO4 showed that higher olefin conversions were
achieved with the ionic liquid, which also maintained its perfor-
mance for longer. It was particularly notable that the formation of
heavy products, which are commonly catalyst deactivators, was less
in the ionic liquids, largely due to the biphasic character of the ionic
liquid system. However, the TMP:HMP ratio was less in the ionic
liquid than in H2SO4. The catalytic performance of the ionic liquid
systems varied with both the cation and the anion of the ionic
liquids. This was found to arise from a combination of solubility
(increasing the cation alkyl chain length increases the hydrocarbon
solubility in the ionic liquid) and acidity effects. It was also found
that the TMP:DMH ratio was greater for the [CnC1im]Br-AlCl3
ionic liquids, with variation in the cation having little effect,
suggesting that acidity affects this result, but that reactant solubilities
do not. However, the overall selectivity to octanes was lower for
these ionic liquids, with a greater proportion of light cracking
products being formed. Variation of the composition of the
[C8C1im]Br-AlCl3 ionic liquid showed that as the ionic liquid
was made more acidic the amount of both cracking products and
oligomerization products (olefins and heavy products) increased
Scheme 20

AM dx.doi.org/10.1021/cr1003248 |Chem. Rev. XXXX, XXX, 000–000
Chemical Reviews
REVIEW
alongside the increased conversions. It was shown that the moisture
sensitivity of these ionic liquids led to gradual deactivation of the
catalyst/solvent. It has also been shown that the deliberate addition
ofHCl to a [(C2)3NH]Cl-AlCl3 [X(AlCl3) = 0.67] ionic liquid led
to lower selectivity to octanes, with a concomitant increase in the
proportion of cracking products.423 This arises because of the well-
established superacidity of these systems.1 Hence, as with many of
the electrophilic aromatic substitutions described above, extreme
care is required in controlling the conditions of these reactions.
Some Hþ, often from contact with adventitious water, is required
for the reaction to occur at all; too much leads to poor selectivity to
the required products, and consequently, prolonged exposure to
water leads to deactivation of the catalyst/solvent.
The addition of benzene (1%, w/w) to [(C2)3NH]Cl-AlCl3
[X(AlCl3) = 0.67] was found to decrease the formation of
byproduct, increasing the selectivity to octanes and the TMP:
DMH ratio, even though some alkylbenzenes were formed.424
This was attributed to the benzene moderating the ionic liquid’s
acidity, as indicated by the IR spectrum of CH3CN. It was
reasoned that the addition of benzene to the ionic liquid was
expected to have a limited impact upon the number of acidic
species in the solution and that the effect might be through its
coordination to the chloroaluminate species. However, it was also
noted that arenium cations, which might be the source of the
Brønsted acid in these reactions, are formed under these condi-
tions and that these cations could prevent the formation of
cracking products.
The addition of metal halide salts to [(C2)3NH]Cl-AlCl3
[X(AlCl3) = 0.67] was found to have only a marginal impact on
products for SnCl4, ZnCl2, or NiCl2, but led to significantly
improved overall selectivity to octanes and TMP:DMH ratio when
CuCl was used.423 In subsequent studies, this was attributed to the
formation of a mixed-metal [AlCl4CuCl]
- anion, which was
proposed to moderate the acidity of the ionic liquid.425 PetroChina
has introduced a commercial alkylation process using these “com-
posite ionic liquids” to tune the acidity of the ionic liquid for process
optimization.426 In 2006 this process was retrofitted into an existing
65 000 t/year sulfuric acid alkylation unit in China.427 The process
operates at ambient temperature and moderate pressure, with
increased yields, greater process unit capacity (40% greater), and
less reactor corrosion than the H2SO4 process. These properties
together give a more economic process.
HF/amine (up to 22:1 molar ratio) mixtures have also been used
for isobutane alkylation.290 These mixtures have formed the basis for
UOP’s commercial Alkad process.428 These compositions give
equilibrium mixtures (e.g., for pyridine, Scheme 22) that are both
far less volatile and less corrosive than pure anhydrous HF. The
volatility of the mixtures (which is proportional to the concentration
of free HF) increased with increasing HF:amine ratio, although no
simple relationship between the volatility at any given composition of
different amine mixtures and the basicity (pKa) of the amine was
found. It was also reported that the alkylate quality varied with the
composition of the mixture, although no details are available in the
public literature of the precise compositionof the commercial system.
However, since the more commonly used 9:1 HF/pyridine compo-
sition was said to be ineffective in this reaction, it can be speculated
that the free HF that is present in the more acidic compositions is
required for the reaction to occur. In this case, the catalysis is best
described as being by a solution of HF in the ionic liquid
[Hpyr][(HF)xF] (Scheme 22).
Protic ionic liquids have also been used as catalysts/solvents
for isobutane alkylation.429 Whether the acidic site was on the
cation {e.g., [(HO3S)
4C4C1im][NTf2]} or the anion {e.g.,
[C4C1im][HSO4]}, lower conversions and selectivities, both
selectivity to C8 products and TMP:DMH ratio, were found in
these ionic liquids than were found for sulfuric acid itself (97.3%
conversion, 61.7% C8 selectivity, and 6.1 TMP:DMH ratio).
[(HO3S)
4C4C1im][HSO4] gave a better TMP:DMH ratio, but
again poorer selectivity to C8 products and lower overall con-
version. Adding H2SO4 (75 wt %) to this ionic liquid gave
improved selectivities (62.3% C8 selectivity and 6.4 TMP:DMH
ratio) and hence a better product at only a slight cost in terms of
conversion (95.6%) compared to using sulfuric acid alone.
6.7.3. Esterification. Esters are widely used compounds in a
wide range of consumer products, from pharmaceuticals to fabrics.
Consequently, esterifications are important industrial processes, and
they have also attracted much fundamental academic interest. The
simple direct stoichiometric condensation reaction of an acid with an
alcohol should be a highly environmentally efficient process, with
water as its only byproduct (Scheme23).However, due to the usually
unfavorable position of the esterification equilibrium, reactive acid
derivatives or coupling agents are often required to achieve high
yields. Since ionic liquids often mix well with carboxylic acids and
short-chain alcohols but are only poorly miscible with simple esters,
most of the reactions discussed below begin asmonophasic solutions,
Scheme 21
Scheme 22

AN dx.doi.org/10.1021/cr1003248 |Chem. Rev. XXXX, XXX, 000–000
Chemical Reviews
REVIEW
but end as biphasicmixtures. These separations of the esters from the
ionic liquid containing aqueous layer in biphasic systems can be used
to drive the reaction to completion without requiring energetically
costly water removal. Hence, ionic liquids have been used on a
number of occasions in an attempt to conduct esterifications with
unactivated starting materials.
The first reported Fischer esterification in ionic liquids used acidic
chloroaluminate ionic liquids.430 However, the moisture sensitivity of
these ionic liquids means that they could never become practical
catalysts/solvents for these reactions. The Fischer esterification of
ethanol and acetic acid was also one of the first reactions to be
conducted in acid-functionalized ionic liquids.47 [(HO3S)
3C3-
(Ph)3P][pTSA] was liquid at the temperature of the reaction, but
solid at room temperature, giving a very simple method for its
isolation and recycle at the endof the reaction. It was found that yields
of the ethyl acetate productwere actually improved by the presence of
some water. Similar effects are known for solid polymeric sulfonic
acids. Optimum conditions for the reaction were found to be 1:1:7
reactants/ionic liquid/water.431 It is likely that the water is required
for proton transfer from the-SO3H group to the acetic acid to occur
efficiently.
Protic ionic liquids with a labile proton on the nitrogen of an
ammonium ion can also act as Brønsted acids in esterifications.432,433
When methanesulfonic acid was added to the [C1Hpyrr][C1SO3]
ionic liquid, its ester was found in the productmixture, which was not
seenwhen the reactionwas conductedwith the neat ionic liquid. This
clearly indicates that the catalysis was by the protic ionic liquid and
that there are negligible amounts of free methanesulfonic acid in the
pure ionic liquid itself.
Brønstedaciditycanalsobeimpartedtotheionicliquidviatheanion.
The ionic liquids [CnC1im][HSO4] (n=4or6), [C4C1im][H2PO4],
[((HO)2C2OC2)C1im][HSO4], and [(HO3S)
nCnC1im][PF6]
(n = 1 or 2) have been compared as catalysts/solvents for the Fischer
esterification of a range of acids and alcohols.434 The ionic liquids had
the advantage over concentratedH2SO4 that the reactionsweremore
selective to the desired esters. Comparison of the various imidazolium
ionic liquids showed a distinct anion effect on the rate of the
esterification, with the more acidic [HSO4]
- anion giving more
effective catalysis than [H2PO4]
-. For these ionic liquids the effect of
the cation was much less marked and probably associated with
changes in the solubilities of the reactants andproducts. Similar results
were observed with some ethylammonium-based ionic liquids.435
However, the sulfonic acid-bearing [(HO3S)
nCnC1im][PF6] (n = 1
or 2) ionic liquids were more effective catalysts than the [HSO4]
-
ionic liquids. Subsequently, it was demonstrated that the rates of
esterifications in [HSO4]
- ionic liquids could be greatly improved
when 5% H2SO4 was added to the ionic liquid.
436 It was noted that
when this reaction system was recycled, there was no loss of activity,
suggesting that the H2SO4 was not extracted during the workup.
Of course, it is possible to prepare ionic liquidswith both a sulfonic
acid-bearing cation and a Brønsted acidic anion.437 The ionic liquids
[(HO3S)
3C3pyr]X {X = [BF4]
-, [H2PO4]
-, [HSO4]
-, or
[pTSA]} have been compared in the esterification of benzoic acid
with simple alcohols. The yield of esters in [(HO3S)
3C3pyr][HSO4]
was considerably higher than that in the [BF4]
- or [pTSA]- salts.
[(HO3S)
3C3pyr][HSO4] was the least miscible of the ionic liquids
with the products, which led to the greatest shift of the esterification
reaction equilibrium to the product side. It was later confirmed by
Hammett acidity measurements that the difference in the rates of
the reactions arose from the different acidities of the ionic liquids.279
This has been reconfirmed in studies of other esterifications with
closely related ionic liquids,438,439 including the esterification of
acetonitrile.440
Two recent studies sought to directly compare the catalysis of the
Fischer esterification by ionic liquids with different types of potential
proton sources. In one, the catalytic activity of the ionic liquids
[(HO3S)
4C4C1im]X {X = [pTSA], [OTf], [CF3CO2], [HSO4], or
[H2PO4]}, [(HO3S)
4C4pyr][OTf], [C4C1im][PF6], and [HC4-
im][HSO4] in the esterification of butyric acid and methanol was
studied.441High yields were only found for the-SO3H-bearing ionic
liquids, with no significant difference between the pyridinium
and imidazolium ionic liquids with a common cation. For the
[(HO3S)
4C4C1im]X ionic liquids the yields varied with the anion
such that [H2PO4]
- < [CF3CO2]
-
, [pTSA]- < [HSO4]
- <
[OTf]-. In the second study the ionic liquids [C4Him][BF4],
[C4Him][OTf], [(C2)3NH][HSO4], [(HO3S)
4C4C1im][pTSA],
and [(HO3S)
4C4C1im][HSO4] were compared.
442 This found
that for ethanol with acetic acid the conversion of the acid varied
such that [C4Him][BF4] < [C4Him][OTf] , [(C2)3NH][HSO4]
< [(HO3S)
4C4C1im][pTSA] < [(HO3S)
4C4C1im][HSO4], in line
with the expected acidity of the ionic liquids. In this study, it was also
noted that when the reactants were water miscible, ionic liquids with
more hydrophilic anions gave better conversions, but when theywere
water immiscible, the reverse was true.
Esterification catalyzed by a solution of acid in an ionic liquid has
also been investigated. The effects of [C4C1im][BF4], [C4C1im]-
[PF6], and [C2C1im][PF6] on the equilibria of pTSA-catalyzed
Fisher esterifications of ethanol and acetic acid showed distinct
differences for the different ionic liquids.443 In the absence of ionic
liquid the equilibrium alcohol conversion reached was 63% in 6 h. In
the hydrophilic [C4C1im][BF4] only 38% conversion was achieved,
whereas in the hydrophobic [C4C1im][PF6] and [C2C1im][PF6]
73% and 83% conversions, respectively, were seen in 24 h under
otherwise identical conditions. This, almost certainly, arises from the
productwater separating from the [PF6]
- ionic liquids, displacing the
reaction equilibrium toward the products. A detailed kinetic study
compared the rates of the pTSA-catalyzed reaction of benzyl alcohol
and methoxyacetic acid in ionic liquids {[C4C1im][NTf2], [C4C1-
im][OTf], [C4C1C1
2im][NTf2], or [C4C1pyrr][NTf2]} and mo-
lecular solvents (toluene, CH3CN, or THF) using the LSER
method.444 An LSER that included both the ionic liquids and
molecular solvents was found (eq 18), indicating that there was no
change in mechanism in moving from molecular solvents to ionic
liquids. It also overwhelmingly demonstrated that the only factor that
was important was the basicity of the solvent. This was interpreted as
arising from the solvent reducing the acidity of the proton in solution
and hence its ability to catalyze the reaction. These two studies
together demonstrate that for both optimal equilibrium conversions
and rates the bacisity of the ionic liquids should be minimized when
they are being used as solvents for esterifications.
ln k2 ¼ - 7:25β- 3:59 ð18Þ
Binary metal isopropoxide complexes {Zr(IV), Fe(III),
Ga(III), and Sn(IV)} have also been used as catalysts for the
direct esterification of 4-phenylbutyric acid with benzyl
alcohol.445 When [C4C1im][PF6] or [C4C1im][BF4] was used
with Zr(OiPr)4, no reaction was seen, whereas the catalyst was
Scheme 23

AO dx.doi.org/10.1021/cr1003248 |Chem. Rev. XXXX, XXX, 000–000
Chemical Reviews
REVIEW
active in [C4C1im][OTf] and [C4C1im][NTf2]. This suggests
that HF, generated from the reaction of water with the [PF6]
-
and [BF4]
- ions, is suppressing the catalyst’s activity. However,
with all of the catalysts the use of large amounts of ionic liquid led
to decreased reactivity, probably because of direct interactions
between the ionic liquid ions and the catalysts.
Esterifications can also be effected by biotransformations in
ionic liquids.8 Due to their tolerance of organic solvents, the use
of lipases in nonaqueous solvents has been widely explored for
some time. Hence, they were considered to offer a good chance
of showing good activities in ionic liquids. However, early studies
of direct enzyme-catalyzed esterifications in ionic liquids have
been unpromising.446-449
Chemical catalysis of transesterifications has been far less
widely investigated than direct esterification. Early studies
showed that transesterifications of β-ketoesters with sulfamic
acid as the catalyst were more selective in [CnC1im]Cl (n = 3 or
5) than the same reactions in CH2Cl2 or hexane.
450 Similar
results have been found when [HSO4]
- ionic liquids have been
used as catalysts for these transesterifications.451 However, in
neither case was any explanation for how the ionic liquid prevents
attack by the alcohol at the β-carbon of the ketoester offered.
More recently, the internal transesterification of 2-hydroxypro-
pyl p-nitrophenyl phosphate in a number of ionic liquids showed
a distinct increase in the rate of the reaction in ionic liquids when
compared to water.452 For the ionic liquids [C4C1im]X {X =
[CH3CO2], [OTf], [InCl4], or [BF4]} the rate of the reaction
increased dramatically with the change of anion in the order
[CH3CO2]
- < [OTf]- < [InCl4]
- < [BF4]
-. However, no
simple rationalization for this observation could be found. For
the transesterification of cottonseed oil with methanol in a range
of ionic liquids with sulfonic acid-containing cations, the chang-
ing reactivities were attributed primarily to the Brønsted acidities
of the ionic liquids.453
The ring-opening polymerizations of lactones can be consid-
ered to be transesterifications. ε-Caprolactam has been polym-
erized by a range of rare-earth-metal triflate complexes in ionic
liquids.454 When the ionic liquid was [C4C1im][BF4], the
reactions were slower, and only afforded oligomers with Mn e
600, than when the ionic liquids were [C4C1im][PF6] or
[C4C1im][SbF6], in which Mn values of up to 4400 could be
achieved. When Sc(OTf)3 or Eu(OTf)3 was used as the catalyst
in [C4C1im][BF4], no reaction was observed.
6.7.4. Diels-Alder Reactions. A number of ionic liquids
with Lewis acidic anions have been used as a combination of
solvent and catalyst for Diels-Alder reactions. Lee noted that
using [C2C1im]Cl-AlCl3 ionic liquids with an excess of AlCl3
(51 mol %) as the solvent/catalyst led to dramatically increased
endo/exo selectivities for the reaction of cyclopentadiene and
methyl acrylate, when compared to those with an excess of
[C2C1im]Cl.
455 The increased reactivity and selectivity arise due
to the presence of [Al2Cl7]
- in the acidic compositions of the
ionic liquid. This was then subsequently used to invert the
selectivity of the reaction of cyclopentadiene and methyl trans-
crotonate from a slight exo preference to an endo preference.456
In an attempt to get away from the extreme moisture sensitivity
and corrosiveness of these chloroaluminate ionic liquids, the use
of [(HO)2C2NH3]Cl-MCl2 {M=Zn or Sn) with an excess (67
mol %) of the metal halide has been explored.457 Although the
sensitivity of these reactions to the precise reaction conditions
can affect the results, making direct comparisons difficult, the
endo/exo ratio achieved with the [(HO)2C2NH3]Cl-ZnCl2
ionic liquid was noticeably higher (24:1) than had been achieved
in the previously reported chloroaluminate ionic liquid (19:1).
Given that chloroaluminate species are generally regarded as
more potent Lewis acids than chlorozincate species, this differ-
ence may be the result of the composition of the [(HO)2-
C2NH3]Cl-ZnCl2 ionic liquid giving rise to a higher concentra-
tion of active Lewis acid catalyst.
The reaction of myrcene with acrolein yields two regioi-
somers, “p”-myrac aldehyde and “m”-myrac aldehyde. In acidic
[C4C1im]Cl-ZnCl2 these reactions are considerably faster and
marginally more selective to the “para” isomer than when
conducted in a solution of ZnCl2 in CH2Cl2.
458 This was
attributed to the increased concentration of the active Lewis
acid catalyst, which activates the dienophile via interaction with
the carbonyl oxygen. The reaction with myrcene was repeated
withmethyl acrylate, methyl methacrylate, andmesityl oxide, and
the relative reactivities of these compounds {acrolein > methyl
acrylate > methyl methacrylate > mesityl oxide} were used to
confirm this explanation.
Chiral C2-symmetric bis(oxazoline)-copper {BOX-Cu} com-
plexes have been used as catalysts for Diels-Alder reactions in ionic
liquids. When the IndaBOX-Cu-catalyzed reaction of cyclopenta-
diene with 3-acryloyloxazolidin-2-one (Scheme 24) was conducted
in [C4C1im]X {X = [BF4], [OTf], [PF6], or [SbF6]}, it was faster
and more selective when the less basic anions were used.459 It was
proposed that this was the result of the elevated water levels in the
hydrophilic ionic liquids, which were not predried before the
reactions. Recycling experiments with the product extracted by
diethyl ether demonstrated a steady drop in selectivity, although the
reactivity remained high. This suggests that the ligand was being
extracted from the solution, but that the copper remained in some
other form. The use of imidazolium-tagged BOX-Cu complexes
prevents this problem.460 Here it was found that marked improve-
ments in both rates and selectivities could be demonstrated in ionic
liquids with [NTf2]
- anions, when compared to using the same
catalyst in dichloromethane. It was also found that the time taken for
the catalytically active species to form was far shorter in the ionic
liquids. However, it was also found that ionic liquids with bromide
anions were substantially less effective solvents for these reactions.
In these experiments considerable care was taken to exclude
moisture, and its presence cannot be the explanation for the
difference. It is more likely to be that the Br- coordinates the
copper directly, preventing the formation of theBOX-Cu complex.
The same reaction of cyclopentadiene with 3-acryloyloxazolidin-
2-one (Scheme 24) has been studied using platinum diphosphine
complexes in [C2C1im][NTf2].
461 As with the copper-based Lewis
acid catalysts above, the rates of the reactions were appreciably
greater in the ionic liquid than in dichloromethane. Although good
selectivities could be obtained in the dichloromethane solutions,
these required that the reaction be performed at-20 C,whereas in
the ionic liquid excellent selectivities were achieved at room
temperature. Here the effect of the ionic liquid was attributed to
Scheme 24

AP dx.doi.org/10.1021/cr1003248 |Chem. Rev. XXXX, XXX, 000–000
Chemical Reviews
REVIEW
its stabilization of the catalytically active species, which formed
unreactive complexes in dichloromethane.
In an attempt to gain insight into the effects of Lewis acids in ionic
liquids, the reactions of acrolein,methyl acrylate, and acrylonitrilewith
cyclopentadiene have been studied in some detail (Scheme 15).462 In
[C4C1im][NTf2] the activity of the catalysts followed the trend ZnI2
<AlCl3 < InCl3 <Sc(OTf)3.The fact that the effect of theLewis acids
was much less on the reactions of acrylonitrile was used as evidence
that the catalysts interacted directly with the polar groups of the
dienophiles. It was necessary to keep the catalyst concentrations
below concentrations at which the polymerization of the starting
materials began to compete with the cycloaddition. When the ionic
liquids [C4C1im]Cl-MCl3 (M = Al or In) were used for the
reactions of acrolein, reactivities were similar to and selectivities were
less than those achieved using solutions of the respective metal halide
in [C4C1im][NTf2]. Other studies in which several Lewis acids have
been comparedhave used the ionic liquid [C6C1im][BF4].
463 For the
reaction of cyclopentadiene with methyl vinyl ketone the reactivies
followed the trend Li[OTf] ≈ Li[NTf2] < BF3 ≈ ZnI2 ≈ AlCl3 <
HOTf < HNTf2 < Y(OTf)3 < Sc(OTf)3 ≈ Sc(NTf2)3 < Ce-
(OTf)4
3
5H2O. It is interesting to note that both of the pairs
Li[OTf]/Li[NTf2] and Sc(OTf)3/Sc(NTf2)3 showed no significant
change in reactivity or selectivity with the change of anion. It was
suggested that the low reactivities of the lithium, zinc, and aluminum
catalysts were due to partial deactivation by water in the ionic liquid.
Kumar and Pawar464 combined the use of the dissolved
catalyst [Et3NSi(toluene)][B(C6F5)4] with an acidic chloroalu-
minate ionic liquid. The rates of reaction of the combined system
were greater than those of either component alone for the same
catalyst dissolved in a [BF4]
--, [PF6]
--, or [SbF6]
--based ionic
liquid or of a pure [C4pyr]Cl-AlCl3 ionic liquid. This use of the
ionic liquid as an active catalyst support allowed for a reduction in
the amount of catalyst used.
6.7.5. Dehydration. Renewed interest in the chemistry of
biomass-derived compounds has led to attempts to dehydrate
alcohols and particularly sugars in ionic liquids. 5-(Hydroxy-
methyl)furfural (HMF) can be prepared by the dehydration of
fructose (Scheme 25) by heating with solid [Hpyr]X, {X = Cl,
[pTSA], or [CF3CO2]} or [(C2)4N]Br in 65-75% yields without
the need for added acids.465 Only very poor yields were obtained in
this way from glucose. However, dehydration of the glucose
derivative sorbitol and other polyols could be achieved under these
conditions.466 This led Lansalot-Matras and Moreau to attempt the
reaction in [C4C1im][BF4] and [C4C1im][PF6] using Amberlyst-
15 as a solid acid catalyst.467 The best results were achieved when
DMSO was added as a cosolvent due to the poor solubility of
fructose in these ionic liquids. The ionic liquids 3-allyl-1-(4-sulfobu-
tyl)imidazolium triflate, [(HO3S)
4C4(CdC2)im][OTf], and 3-al-
lyl-1-(4-sulfuryl chloride butyl)imidazolium triflate, [(ClO2S)
4-
C4(CdC2)im][OTf], have both been used as catalysts for this
reaction, with the latter Lewis acidic form being more active than the
Brønsted acidic version.468
When [C1Him]Cl was used as the solvent at 90 C, no
cosolvent nor added acid was required to achieve reasonable
yields and rates.469 It was also found that selectivity to the
5-(hydroxymethyl)furfural was excellent, with no decomposition
products (e.g., levulinic acid) seen. If sucrose was used as the
starting material, it was rapidly cleaved by water into fructose,
which reacted to form HMF and glucose, which did not react
under these conditions. Subsequently, it was found that
[C2C1im]Cl could also act as a catalyst/solvent for the dehydra-
tion of fructose, showing that a protic ionic liquid is not required
for dehydration to occur.470 Dehydration of benzyl alcohols has
also been achieved under acid-free conditions in the ionic liquids
[C4C1im]X {X = Cl, Br, [BF4], or [PF6]} or [C6C1im]Br using
microwave heating.471 It is particularly interesting to note that
the rates as demonstrated by yields at fixed time varied in the
order [PF6]
- < [BF4]
- < Cl- < Br-, i.e., with increasing hydro-
philicity of the anion, but greater yield could be attained in less
time in [C6C1im]Br, with the more hydrophobic cation, than in
[C4C1im]Br.
The rate of the dehydration of fructose in [C2C1im]Cl could be
improved by the addition of several different metal chloride salts
(CrCl2, CrCl3, FeCl2, FeCl3, CuCl, CuCl2, VCl3, MoCl3, PdCl2,
PtCl2, PtCl4, RuCl3, or RhCl3).
477 In contrast to this, the dehydration
of glucose could only be achieved with CrCl2, for which good
selectivity to 5-(hydroxymethyl)furfural was found. It was thought
that this was because the solute metal complex {[CrCl3]
-
} was
coordinating to the glucose and facilitating its mutarotation into a
more reactive form. SnCl4 has also been reported to dehydrate
glucose to HMF in [C2C1im][BF4].
472 For a range of [C4C1im]X
{X = [BF4], [PF6], [NTf2], [OTf], or [saccharin]} ionic liquids, it
was found that more coordinating anions lead to lower HMF yields,
presumably by forming anunreactive tin complex.Theonly exception
to this trend was the very poor reactivity found in [C4C1im][PF6],
but in this case the glucose was observed to be insoluble. Having
settled on the [BF4]
- anion, the authors found that the yields of
HMF increased in the order [C4pyr][BF4] < [C4C1im][BF4] <
[C2C1im][BF4], which also gave a higher yield than DMSO.
1H
NMR evidence was presented for the formation of a similar
chelate complex between the tin and the glucose, which enabled
the reaction to proceed. Although the formation of HMF from
glucose has only been reported with the use of a metal catalyst, a
partial dehydration has been reported in some [NTf2]
- ionic liquids
at 200 C.473
6.7.6. Beckmann Rearrangement. ε-Caprolactam is a
monomer for the production of nylon-6. As such it is produced
industrially on a huge scale by several producers. The first step in its
synthesis is from cyclohexanone with hydroxylammonium sulfate,
whichmelts with decomposition at 120 C, to give the intermediate
cyclohexanone oxime, which is then treated with concentrated
sulfuric acid or oleum to yield ε-caprolactam. It has recently been
pointed out that in the second step of this process the product ε-
caprolactam reacts with the H2SO4 to give the salt ε-caprolacta-
mium hydrogen sulfate (Scheme 26), which melts at 60 C.474 This
ionic liquid then acts as a solvent for the reaction. The synthesis of ε-
caprolactam can therefore be considered to be a well-established
large-scale industrial ionic liquid process. The lower than expected
vapor pressures of solutions of SO3 in this ionic liquid suggest that
[HS2O7]
- is formed in solution.
Scheme 26Scheme 25

AQ dx.doi.org/10.1021/cr1003248 |Chem. Rev. XXXX, XXX, 000–000
Chemical Reviews
REVIEW
The industrial production of ε-caprolactam also yields ammo-
nium sulfate as a byproduct. Even though ammonium sulfate can be
used as a fertilizer, attempts have been made to find an ionic liquid
process that avoids this.475,476 It has been found that the ionic liquid
ε-caprolactamium tetrafluoroborate can act as a catalyst/solvent for
the reaction without the need to add acid, with product separation
being achieved using molecular solvents.477
The Beckman rearrangement of a number of ketoximes into
amides, catalyzed by Lewis acids (AlCl3, TiCl4, SnCl4, or BF3) have
been conducted in 17 different ionic liquids.478 AlCl3 and TiCl4
were found in all cases be more effective catalysts than SnCl4 and
BF3. Generally, changing the ionic liquid had a greater effect on the
reactions of these latter two catalysts than on the more reactive
AlCl3 and TiCl4. A small decrease in reactivity with increasing alkyl
chain lengths of the cation was observed, and the rate of the
reactions increased with C2 substitution of the imidazolium ring.
The impact of the anion of the ionic liquid on the rates of the
reactions is less clear. It appears to be a combination of the basicity
(possibly also hard/soft character) of the anion, hence its ability to
interact with Lewis acids, and the degree of hydration of the ionic
liquid, but no definitive conclusion can be drawn.
6.7.7. Ring Closing of Isonitrosoacetanilides. The ring
closing of isonitrosoacetanilides to give isatin 3-oximes can be
catalyzed by Lewis or Br€onsted acids in ionic liquids.479 No reaction
was observed in neutral [C4C1im]Cl-InCl3, with acidic composi-
tions being required to achieve modest reactivities. More useful
yields were obtained with Brønsted acids in [C4C1im][NTf2]. The
yields obtained with different dissolved acids increased in the order
pTSA (31%) < CF3CO2H (61%) < CH3SO3H (88%) < HBF4
(96%). However, HBF4 in [C4C1im][BF4] only yielded 25% of the
product. The authors attributed this difference in the behavior of the
two ionic liquids to the cationic charged intermediate being more
stabilized by the “low coordinating” anions.
6.7.8. Mannich Reaction. The Mannich reaction is a method
for preparing β-amino acids from an aldehyde, ketone, and mono-
alkylamine. It is usually catalyzed by either Brønsted or Lewis acids.
The reaction in the ionic liquids [C4C1im][pTSA], [C4C1im]-
[CF3CO2], [C4C1im][H2PO4], and [C4C1im][HSO4] displayed
an interesting reversal of activities in comparison to the reactions
described above where reactivities in these ionic liquids, or closely
related ones, were compared.480 No reaction at all was seen in
[C4C1im][HSO4], and the greatest yields were observed in
[C4C1im][pTSA]. This was rationalized by noting that although
acidwas required to catalyze the reaction, if the acidwere too strong, it
would react directly with the amine, preventing the formation of the
Mannich base.
6.7.9. Glycosylation. The synthesis of carbohydrate deriva-
tives remains an extremely demanding area of research. Many
syntheses require protection/deprotection protocols and sophis-
ticated coupling reagents to achieve a single anomeric product.
The use of Lewis acids as promoters for alcohol glycosylation with
trichloroacetinidates in ionic liquids has shown an inversion of
stereochemical outcome depending on the ionic liquid used.481 A
detailed NMR experiment showed that the triflate ion of
[C2C1im][OTf] formed a transientR-glycosyl triflate intermediate,
which guides the subsequent attack of the glycosyl acceptor from the
β-face. No such intermediate was formed in [C4C1im][PF6], which
behaved similarly to noncoordinating molecular solvents and
provided products of the opposite configuration.
6.7.10. Summary. Ionic liquids have shown themselves to
have great potential, both as acids themselves and as solvents for
other acids. They can influence the reactions of solute species in
many ways: as tunable acid systems, by being better solvents for
starting materials and/or intermediates than aqueous acid sys-
tems, by giving effective separation of products, etc. Often more
than one of these effects is operating in any one system, and
oversimplified interpretations of results based on just one of
these can fail to give adequate explanations for the observed
behaviors. Ionic liquids can also give rise to problems, such as
hydrolytic instability and the formation of acid wastes. Good
understandings of these various effects and careful selection of
the ionic liquid can lead to highly effective catalysts/solvents.
6.8. Base/Nucleophile-Catalyzed Reactions
Catalysis by bases in ionic liquids has received less attention
than acid catalysis. This is in great part due to the known
sensitivity of ionic liquids to even mild bases, which can seriously
affect these types of reactions.482 The formation of NHCs from
imidazolium ionic liquids is a particular problem. For example,
the disparity between the high benzaldehyde conversions (as
determined by GC) and low product yields in the base-catalyzed
Baylis-Hillman reaction of methyl acrylate and benzaldehyde
catalyzed by 3-hydroxyquinuclidine or 1,4-diazabicyclo[2.2.2]-
octane (DABCO) in the presence of [C4C1im]Cl can be
explained by a reaction of the benzaldehyde with [C4C1im]
þ,
via the deprotonated NHC.483 Consequently, any interpreta-
tions of the results for base-catalyzed reactions in imidazolium
ionic liquids must be viewed as potentially compromised unless a
specific investigation into the involvement of NHCs, whether for
good or ill, has also been reported and will not be reviewed here.
6.8.1. Aldol Reaction. The aldol addition is one in which a
carbon of one aldehyde or ketone adds to the carbonyl carbon of
another to yield a β-hydroxycarbonyl compound (Scheme 27),
which may then dehydrate to an R,β-unsaturated carbonyl com-
pound, at which point the reaction becomes known as the aldol
condensation (Scheme 27).Not only is the aldol reaction significant
in its own right, it also occurs as a step inmechanistic interpretations
of other widely used reactions. It can be acid catalyzed, but is most
often catalyzed by bases, traditionally metal hydroxides or alkoxides.
The presence of these bases leads to the possibility of the formation
of NHCs and the products of their reactions with the aldehyde
starting materials, as described above.483
The reaction of p-(trifluoromethyl)benzaldehyde, which is a
highly activated aldehyde for these reactions, with acetone in
[C4C1im][PF6] with no added catalyst showed no evidence of
reaction.484 When NaOH was added to [C4C1im][PF6] as an
ethanolic solution, the reaction of benzaldehyde with acetophe-
none (the Claisen-Schmidt variation of the aldol reaction) was
found to proceed with good conversions.485 However, even
when the ethanol had been removed by vacuum, ethyl benzoate
Scheme 27

AT dx.doi.org/10.1021/cr1003248 |Chem. Rev. XXXX, XXX, 000–000
Chemical Reviews
REVIEW
benzaldehyde and substituted benzaldehydes, with no additional
solvent.520 [(C1OC3)NH3][CH3CO2] has been used similarly for
the reactions of a number of aldehydes with a number of active
methylene compounds.521 A cyclic trimethylguanadinium lactate
ionic liquid has been used as a catalyst/solvent for the reactions of
benzaldehyde or a range of substituted benzaldehydes with maloni-
trile, ethyl cyanoacetate, or diethyl malonate, with reaction times and
yields reflecting the relative reactivities of the reactants.522 Ethylam-
monium nitrate has also been found to lead to high yields for the
reaction of p-nitrobenzaldehyde with malonitrile in 3 h at room
temperature517 and to act as a catalyst/solvent for other Knoevenagel
condensations more effectively than [C4C1im][PF6], [C4C1im]-
[BF4], water, or ethanol.
523 Finally, for this group of ionic liquids,
[C1Him][CF3CO2] has been used as the catalyst/solvent for the
reaction of substituted benzaldehydes with 2,2-dimethyl-1,3-dioxane-
4,6-dione,with reactions reaching 94%yields in 30min, as opposed to
54% {[C4C1im][PF6]}, 76% {[C4C1im][BF4]}, or 77% {[C6C1-
im][BF4]}.
524
This collection of results indicates that the protic ionic liquids with
basic anions are acting as effective catalysts/solvents for Knoevenagel
condensations and together allow for some speculation about how
this is achieved. The mechanism for the base-catalyzed Knoevenagel
condensation has the activated methylene compound being depro-
tonated to give a carbanion intermediate (Scheme 30). The basic
anions of the more effective ionic liquids above may be able to act as
the base in the formation of the carbanion.521 This could be further
enhanced by the carbanion being stabilized by ionic liquids in
comparison to molecular solvents. This is similar to the catalysis of
Michael reactions by 1-methylimidazole in ionic liquids.357 The basic
anions of the ionic liquid could also facilitate the dehydration step
(see above).
However, there is another potentialmechanism for these reactions.
The basicities of these anions are sufficient such that there will be
a significant concentration of the free amine present when the
cation carries a N-H proton.72 This has been demonstrated by
the formation of ethylamine-derived products in nucleophilic
aromatic substitution reactions in [C2NH3][NO3] or [C2NH3]-
[CH3CO2].
525 These free bases could also act as base catalysts for the
reaction. However, the same study showed that substituted benzal-
dehydes gave imines in [C2NH3][NO3], presumably via an iminium
intermediate (Scheme 31). This gives rise to the second possible
explanation of the results above. The free base, or perhaps the
ammonium ion directly, reacts with the aldehyde to form an
iminium ion, leading to nucleophilic catalysis of the reaction
(Scheme 32).526 This has implications beyond the Knoevenagel
reaction to any reaction for which this type of catalysis is possible
when conducted in protic ionic liquids487 and as demonstrated
in other aldol reactions.497
6.8.4. Keto-Enol Tautomerization. The keto-enol tauto-
merization of 2-nitrocyclohexanone can be base catalyzed
(Scheme 33). The equilibrium constant for the uncatalyzed reaction
has been used to predict Hildebrand parameters, δH, for ionic
liquids.269 The kinetics of the reaction catalyzed by excess pyridine
in [CnC1im][NTf2] and [CnC1C1
2im][NTf2] (n = 2, 4, or 6),
[C4pyr][NTf2], and [C4C1im][PF6] have been compared to those
in a range of both pure molecular solvents and solvent mixtures.527
The second-order rate constant correlated well with εr for the
molecular solvents, but it was not possible to include the ionic
Scheme 31
Scheme 32

AU dx.doi.org/10.1021/cr1003248 |Chem. Rev. XXXX, XXX, 000–000
Chemical Reviews
REVIEW
liquids in this analysis. However, a Kamlet-Taft type of analysis did
provide a good LSER (R2 = 0.991, eq 19), with the increased rates of
reactions found in the ionic liquids being largely due to their ability
to stabilize a more enolate-like transition state for the reaction than
found in molecular solvents.
ln k2 ¼ - 11:3 þ 1:86π þ 4:73Rþ 8:18FðεÞ þ 0:021δ
2
ð19Þ
6.8.5. Nucleophilic Substitutions. Indole is an ambident
nucleophile that can be alkylated at either the nitrogen or the C3
position of its pyrrole ring or both (Scheme 34), with a distinct
solvent effect on this regioselectivity.528 The reactions of indole with
various halogenoalkanes in [C4C1im][PF6] with KOH as the
catalyst gave N-alkylindoles in almost quantitative yields, with only
iodomethane and bromobenzene giving some dialkylation, as is
typical for polar aprotic solvents.529 This methodology has been
extended to other nitrogen heterocycles.530 Similar results were
found for the reactions of 1-bromo-3-phenylpropane with indole
catalyzed by potassium carbonate in acetonitrile/[C4C1im]X {X =
[PF6], [BF4], [NTf2], or [OTf]} mixtures, although with some
formation of 1-(3-phenylpropyl)-1H-indole as a byproduct.531 The
selectivity to the desired 1-(3-phenylpropyl)-1H-indole varied with
the anion in theorder [PF6]
-
≈ [BF4]
-> [NTf2]
-
≈ [OTf]-, with
the reactions in the latter ionic liquids also being slower. Pyrrole itself
has been found to almost exclusively undergo C-alkylation under
similar conditions in [C4C1im][SbF6], with some dialkylation.
319 In
a study using a wider range of ionic liquids, it was shown that the
selectivities found in alkylations with the halogenobutanes fell in the
order found for polar protic solvents for [C4C1im]Br, but in the
order found in polar aprotic solvents in [C7pyr]Br and
[C1C1im][OTs].
532
Base-promoted alkylation of 2-hydroxypyridine in the ionic
liquids [CnC1im][OTs] (n = 1 or 4), [C4C1im]Br, and
[Cnpyr]Br (n = 4 or 7) with a range of bases {MX; M = K, Na,
Li, or Ag; X = H, C4H9, OH, or OCH3} has been found to favor
N-alkylation over O-alkylation in comparison to molecular
solvents.533 The selectivities and rates of these reactions in the
[Cnpyr]Br (n = 4 or 7) ionic liquids were found to be greater than
for the imidazolium ionic liquids. It was suggested that this was
because hydrogen bonding between the imidazolium cations and
the anionic nucleophile was suppressing its reactivity.
Base (CsOH
3
H2O) catalyzed N-alkylations of primary amines in
[C4C1im]X {X = [PF6] or [NTf2]}, [C4C1C1
2im][PF6], or
[C6pyr][NTf2] have been shown to have excellent selectivity to
secondary amines, with amarkedly reduced overalkylation to tertiary
amines.534 The results rivaled those possible in the best available
alternative solvent, DMF. It is likely that this arises from the ability of
the ionic liquids, and indeed DMF, to activate the primary amines
more strongly than the secondary amines through hydrogen bond-
ing. Dialkylation was only found for the activated halides allyl
bromide and benzyl bromide. For the uncatalyzed reactions of
C8H17X the order of reactivity in [C4C1im][PF6] was [OTs]
- >
I-≈ Br- . Cl-, with the addition of CsOH (i.e., the activation of
the amine) reducing the differences.
The base-catalyzed rearrangement of the (Z)-phenylhydrazone
of 3-benzoyl-5-phenyl-1,2,4-oxadiazole into the relevant 4-
(benzoylamino)-2,5-diphenyl-1,2,3-triazole (Scheme 35) has been
studied in both ionic liquids, [C4C1im]X {X = [BF4], [PF6], or
[NTf2]}, [C4C1C1
2im][NTf2], or [C4C1pyrr][NTf2], and molec-
ular solvents (benzene, 1,4-dioxane, ethyl acetate, MeOH, or
MeCN).535 In the ionic liquids, the reaction was found to be first
order in both reactant and base, but in the molecular solvents, more
complex rate laws were found, which depended upon the solvent.
This prevents a quantitative analysis of the bimolecular rate con-
stants for the reactions in different solvents. The authors attributed
the kinetic behaviors in the ionic liquids to their heterogeneous
structures. However, the authors did report the pseudo-first-order
rate constants for the reactions at a fixed concentration of base for
the reactions with piperidine (kA,R), which we have analyzed using
the Kamlet-Taft approach and derived an acceptable LSER (eq 20,
R2 = 0.85). The positive β and π* effects are in line with the LSERs
found for other nucleophilic substitutions of amines (see above).
ln kA, R ¼ - 30:9þ 16:5βþ 20:4π ð20Þ
6.8.6. Esterification/Acetylation of Alcohols. Base-cata-
lyzed esterification has not received asmuch attention in ionic liquids
as its acid-catalyzed counterparts. The transesterification of soybean
oil with alcohols to give biodiesel and glycerol with ionic liquids has
been achieved with both basic (K2CO3, Cs2CO3, or K3PO4) and
H2SO4 catalysts.
536 With ethanol the transesterification in
[C4C1im]X {X = [NTf2], [BF4], or [PF6]}, [C4C1C1
2im][NTf2],
or [C2C1im][C2SO4] proceeded to almost quantitative yields with
no obvious effect of changing either the cation or anion, although
both the [BF4]
- and [PF6]
- ionic liquids decomposed under the
reaction conditions via hydrolysis with adventitious water. However,
the reactivities in [C4C1im][OTf] were appreciably lower. A
13C
NMR spectrum of [C4C1im][NTf2] with Cs2CO3 saw an almost
complete collapse of the signal due to the C2 carbon of the
imidazolium ring, but with no new signal for the NHC appearing.
Upon addition of ethanol this signal reappeared, suggesting that
Scheme 33
Scheme 34
Scheme 35

AX dx.doi.org/10.1021/cr1003248 |Chem. Rev. XXXX, XXX, 000–000
Chemical Reviews
REVIEW
This ruled out the possibility of a catalytic cycle that required chloride
dissociation in the rate-limiting step, and the authors proposed that a
differentmechanismwas operating in these ionic liquids compared to
water. After careful kinetic investigations they proposed that it was the
dissociation of the neutral p-cymene ligand that was occurring in the
formation of the active catalyst, which was likely to be dimeric and in
the presence of excess Cl- ions was in equilibrium with more heavily
chlorinated species (Scheme 38). This is one of very few examples in
the literature of a clearly demonstrated mechanistic change upon
using ionic liquids.
Differences in the solubilities of reactants and products in the
ionic liquid can also have a pronounced effect on the catalytic
activity. For example, Dupont et al.563 found that the Pd nanopar-
ticle catalyzed partial hydrogenation of 1,4-butadiene to butene was
far more selective in [C4C1im][BF4] compared to [C4C1im][PF6].
The much higher solubility of alkenes in the [PF6]
- ionic liquid
likely led to the reported 3-fold higher production of the over-
hydrogenated product butane. Curiously, in this system, the addi-
tion of water had no effect, despite the likely reduction in alkene
solubility that would result. The same authors reported higher
activity for the partial hydrogenation using Ru nanoparticles pre-
pared by controlled decomposition of [Ru(cod)(cot)] (cod = 1,5-
cyclooctadiene, cot = 1,3,5-cyclooctatriene) in [C4C1im][PF6] over
[C4C1im][BF4];
564 however, results with other substrates were
somewhat mixed, and high conversions rendered interpretation of
relative ionic liquid effects difficult.
Similarly, a number of ionic liquids were screened by Dyson
et al.565 for the hydrogenation of styrene using the ionic cluster
[H3Os4(CO)12]
-. While no clear solvation effect appears in
these results, the more viscous ionic liquids, based on the
[C4C1C1
2im]þ cation, yielded lower TOFs, while the most
active system used [C8C1pyrr][BF4], which was the only solvent
found to completely dissolve the substrate.
The importance of enantioselective catalysis to pharmaceutical
synthesis has led to many attempts to use ionic liquids for these
reactions.12,549 Feng et al. screened a variety of ionic liquids for a Rh-
catalyzed enantioselective hydrogenation of enamides.566 The results
show that adding alcohols or water to the ionic liquid increases both
the reaction rate and enantiomeric excess (ee) over those found in the
ionic liquid alone. However, there is no clear trend among the ionic
liquid anions with regard to hydrophobicity or basicity, with
[C4C1im][BF4] outperforming [C4C1im][PF6] substantially, while
little difference was seen between wet [C4C1im][BF4] and wet
[C4C1im][NTf2]. The longer chain [C8C1im]
þ cations also yielded
superior reaction media. Selectivity differences were more marginal,
with [C4C1im][BF4]/iPOH or [C8C1im][BF4]/iPOH mixtures
generally providing the same ee as pure MeOH, though
[C4C1im][PF6] did show a substantial drop in selectivity. The results
across a series of catalyst complexeswere evenmore confused, leading
the authors to draw no conclusions based on solvent effects alone. It
would appear that the primary effect of adding the water or alcohol is
to prevent the decomposition of the catalyst into a nonchiral colloidal
form. Furthermore, specific interactions between the protic (alcohol
or water) cosolvent and the basic enamide substrate were competing
with alcohol/anion hydrogen bonding, with the result thatmore basic
anions yielded higher activity and (marginally) higher selectivity. It
remains unclear how much of an effect the changing catalyst had on
this competition.
Hydrogen solubilitymay also be responsible for these differences.
Jessop et al. compared the hydrogenation of tiglic acid (which gives
higher ee's at lower H2 concentrations) with that of atropic acid
(which gives higher ee's at higher H2 concentrations) using Ru-
(O2CMe)2(R-tolBINAP) in several ILs.
567 At various H2 pressures
the ionic liquids were found to be generally inferior to MeOH for
the hydrogenation of atropic acid, with the differences in the ee
becoming smaller as the H2 pressure was increased. For the tiglic
acid system, the ionic liquids were generally superior at lower H2
pressures, with the differences again disappearing at higher pres-
sures. This indicates that while the H2 solubilities are lower in the
ionic liquids than in MeOH, the changing solubility with changing
pressure is less substantial. This result was also obtained for acid
substrates by DuPont, which found no difference between alcohol
and ionic liquid solvents over a pressure of 100 bar.568 The effect of
changing the ionic liquid on these systems is less clear, however, with
the [C4C1im][PF6] and [NTf2]
- ionic liquids yielding higher ee's
for tiglic acid hydrogenation than [C4C1im][BF4], despite the lower
solubility of H2 in the latter solvent.
556 The same solvent ordering
was found for the atropic acid reactions, perhaps indicating that
solvent effects on the catalytic reaction itself were at least as
important as H2 solubility, with more basic solvents strongly
hydrogen bonding to the carboxylic acid substrates and lowering
the ee, regardless of the preference of the substrate for a higher or
lower H2 concentration.
Yinghuai et al. used several different ionic liquids for the asym-
metric hydrogenation of ketones.569 Their results indicate that more
hydrophobic (or less basic) anions yielded higher ee's. However, the
authors used a rhodacarborane catalyst, and the most active ionic
liquid employed a very special decaborate anion, so specific catalyst
stabilization may be responsible. Li et al. deployed TPPTS-based Ru
complexes in a variety of ionic liquids.570 Their results are quite
complicated; while ionic liquids with [pTSA]- anions were slightly
more active than [C4C1im][BF4], [C4C1im][PF6] barely gave any
conversion at all. The highest ee was obtained using [C4C1im]
þ as
the cation instead of [C2C1im]
þ, [C8C1im]
þ, or [C12C1im]
þ,
indicating that perhaps substrate interactions with the hydrophobic
part of the ionic liquid (the cation side chain) have an effect on the ee,
possibly through diffusion control. It is interesting that for this system
the more hydrophilic (more basic anion) ionic liquids were generally
superior, in contrast to the results obtained by Yinghuai.569
Lin et al. used 2-propanol/ionic liquid mixtures and a Ru-
(BINAP)(DPEN)Cl2 catalyst for the asymmetric hydrogenation
of ketones.571 They found that [C4C1im][BF4] and [C4C1-
im][PF6] performed poorly when compared to [C3C1C1
2im]-
[NTf2]. They proposed two possible explanations for this discre-
pancy. The first was that the relatively more nucleophilic [PF6]
-
and [BF4]
- ions could be reducing the catalyst’s efficiency.
However, while [BF4]
- has been shown to be considerably
more coordinating, [PF6]
- and [NTf2]
- have much more similar
coordinating abilities.141 Their second explanation was that the
KOH cocatalyst could be deprotonating the C2 position of the
imidazolium ring, leading to the formation of NHCs, which could
then poison the catalyst. Their finding that the use of [C4C1C1
2-
im][BF4] gives results similar to those of [C3C1C1
2im][NTf2]
supports this latter argument. Indeed, it is possible that the
poisoning of homogeneous catalyst complexes by NHCs is respon-
sible for many irregular catalytic findings in ionic liquids.
Scheme 38

AZ dx.doi.org/10.1021/cr1003248 |Chem. Rev. XXXX, XXX, 000–000
Chemical Reviews
REVIEW
presence of ethanol. It is very interesting to note that the ee for
the reaction at first increased with the concentration of the ionic
liquid in the ethanol, but reached a maximum and then declined.
This was interpreted as a consequence of an increasing concen-
tration of ionic liquid ion pairs forming in solution, but that at
higher concentrations higher order clusters were forming, which
were less selective chiral transfer agents.
Ionic liquids have been shown to have striking effects on hydro-
genations, even when used as additives. For the Ru-catalyzed
asymmetric hydrogenation of β-keto esters both the rate and ee of
the reaction inmethanol were greatly affected by the addition of ionic
liquids, with the ee dropping considerably in the presence of
[C2C1im][NTf2] and increasing slightly when [(HO)
2C2(C1)3N]-
[NTf2] was added.
590 This stabilization effect also led to enhanced
recyclability of the catalyst system. This stabilization of the catalyst
was also seen in the Ru-BINAP-catalyzed hydrogenation of β-keto
esters in [(C2)3CnN][NTf2] ionic liquids.
591 In this study, the length
of the alkyl chain did not affect the ee of the reaction on the first run,
but the use of longer alkyl chains led to dramatically improved
performance on subsequent runs.
While the formation of metal colloids from organometallic metal
complexes can be seen as a problem for catalysis in ionic liquids, metal
nanoparticle catalysts have been extensively investigated as catalysts in
ionic liquids. The Rh nanoparticle catalyzed hydrogenation of styrene
was examined in a variety of ionic liquids.592,593 The selectivity
(ethylbenzene/ethylcyclohexane) was found to vary widely with both
the cation and anion. For [C4C1im]X the selectivitywas found to vary
with X in the order [N(CN)2] (100%) > [NTf2] (70%) > [PF6]
(40%) > [BF4] (8%), while for [cat][NTf2] the relative ordering was
[C4C1
4pyr] (85%) > [C4C1pyrr] (70%) = [C4C1im] (70%) >
[(OH)2C2C12(C1)2N] (40%). Although H2 solubility data are not
available for all of these ionic liquids, it appears that overhydrogenation
is most prevalent in the ionic liquids with the highest H2 solubility.
Incorporation of a hydroxyl functionality into the side chain of a
variety of ionic liquid cationswas found to enhance the catalytic rate of
the same reaction catalyzed by Rh nanoparticles with selectivities
approaching 100% in all cases.594This particular result could be due to
the likely increased diffusivity of H2 in these ionic liquids, as the
selectivity for acetylene hydrogenation was lower for Pd nanoparicles
in the same solvents compared to [C4C1im][PF6] due to
overhydrogenation.595 In addition, for the Ru nanoparticle catalyzed
hydrogenation of toluene in [C4C1im]X and [C10C1im]X {X =
[NTf2] and [BF4]}, there was a significant drop in activity when
[C4C1im][BF4] was employed.
596 The [BF4]
- ionic liquids per-
formedpoorly in all of these nanoparticle studies, and it is possible that
breakdown of the [BF4]
- anion, or the poisoning of the catalyst
surface by this anion, is the root cause for all of these findings.
In a heterogeneously catalyzed reactionmass transport is often
an issue. This was highlighted in a study byWasserschied wherein
the activity of a ruthenium on carbon catalyst suspended in a
single-phase ionic liquid systemwas twice as high as for a biphasic
ionic liquid/organic solvent system, but the selectivity for the
biphasic system was near 80% compared to 40-50% for the
single-phase approach.597 Removing diffusion limits can elim-
inate some ionic liquid effects, as for the hydrogenation of citral
using Pd nanoparticles in a thin layer of ionic liquid coated on
Si.598 In these studies, the activity in various ionic liquids was
nearly the same, with a slight drop for [C4C1im][PF6] compared
to [C4C1im][N(CN)2]. The same authors did report a small
cation effect on the selectivity of this reaction using Ru nano-
particles in [CnC1im][NTf2] (n = 4 or 6) ionic liquids.
599
Gholami et al. performed an LSER analysis of the solvent effects
on the synchronous hydrogenation of cyclohexene and acetone
by Pt/Al2O3 in mixtures of the ionic liquid (2-hydroxyethyl)
ammonium formate with ethanol and 2-propanol at various mole
fractions.600 The authors found that the rates of both hydrogenation
reactions increased with increasing π* value, and the selectivity
favors cyclohexene hydrogenation at higher π*.
7.3. Oxidations
Aswell as being included in the general reviews of catalysis in ionic
liquids,12 a more specific review of oxidations in ionic liquids is
available.601 Ionic liquids have a number of advantages as solvents for
oxidations, particularly when either Br- anions or pyridinium cations
are employed, as these are capable of oxygen transfer.More generally,
ionic liquids have been found to stabilize the free radicals necessary
for many oxidation mechanisms.602 Also, a wide variety of ionic or
peroxide oxidants can be used in ionic liquids, and many exhibit
increased stability. Many ionic liquids are stable under oxidizing
conditions and therefore constitute ideal solvents for this class of
reactions.603 However, it should be noted that oxidation of imida-
zolium cations has been observed,94 and unless this possibility has
been explicitly considered, the formation of these oxidation products
cannot be excluded. One of the main attractions of ionic liquids as
solvents for this class of reactions is the selectivities that they can give;
e.g., they have been found to promote oxidations of alcohols to
aldehydes, but suppress the overoxidation of these aldehydes to
acids.94,95 Ionic liquids can also be specially tailored to support ionic
liquid modified oxidants; e.g., periodinanes have been supported by
ionic liquids for the oxidation of alcohols in [C4C1im][BF4] and
[C4C1im][PF6].
604 Peracids605 and iodoacids606 have also been
supported for use in ionic liquids.
Both Ley607 andWelton94 performed early oxidation of alcohols in
ionic liquids catalyzed by tetra-N-propylammonium perruthenate
(TPAP) with the oxidant N-methylmorpholine N-oxide (NMO).
While the reaction proceeded in all cases to high conversion with
catalyst recycling, loss of catalytic activity was generally observed after
2-3 recycles. This was particularly evident when [(C4)4N]Br was
used as the ionic liquid (and oxidation promoter).607 Since the
systems were biphasic, this could be an indication of phase transfer
activity, which ionic liquids are known to promote.302Oxidationswith
molecular oxygen as the oxidant could be carried out using this
catalyst system, but it was found that an organic cosolvent and
vigorous agitation were necessary to maintain the reaction. This is a
direct consequence of the extremely low solubility of O2 in ionic
liquids608 coupled with their generally poor mass transport
properties.609
As with other forms of catalysis, oxidations can be affected by the
presence of impurities in ionic liquids. For example, Zhang et al.
reported that while benzyl alcohol is oxidized to benzaldehyde with
Scheme 39

BA dx.doi.org/10.1021/cr1003248 |Chem. Rev. XXXX, XXX, 000–000
Chemical Reviews
REVIEW
NaOCl in [C4C1im][PF6] at room temperature (even without
TEMPO), the reaction proceeded to much higher conversion and
selectivity using a cyclic guanidinium ionic liquid, [C6(C1)3cy-gu]-
[PF6].
610While it is possible that the guanidium ion is able to stabilize
free radicals much more easily than imidazolium salts, it seems more
likely that residual guanidine from the synthesis of the ionic liquid is
promoting the reaction. The base 4-(dimethylamino)pyridine
(DMAP) has been clearly demonstrated to be a promoter for the
Cu(II)-catalyzed oxidation of alcohols usingTEMPO in [C4C1pyrr]-
[PF6].
611 A mechanism was proposed by which the base deproto-
nates the alcohol to promote its oxidation. Hence, residual base in
ionic liquids is likely to be active in oxidation reactions.
The effect of varying the ionic liquid solvent on alcohol oxidations
has received some attention. The nature of the anion does seem to
play a major role in some oxidation catalyses in ionic liquids. Wolfson
noted thatCl-promotedoxidationof aliphatic alcohols using aRu(II)
catalyst more than OH-, indicating that this is not a simple matter of
basicity.612 The destructive nature of OH- at high concentrations
could also have been responsible for some catalyst degradation. This
conclusion is supported by Seddon,95 who found that the oxidation of
benzyl alcohol withO2 catalyzed by Pd(OAc)2 proceeded in a variety
of imidazolium chlorides and [BF4]
- salts, but not in [C4C1im]Br,
despite the oxidation-promoting properties of Br-.613 In each of these
cases, the removal of byproducts improved the yield. [(C4)4N]Brwas
used as a solvent for the PdCl2-catalyzed oxidation of indan-1-ol;
614
however, side reactions limited this reaction to low selectivity. It is
noteworthy that this reaction did not proceed if [(C4)4N]Clwas used
as a solvent; clearly the Br- ion is playing a key role as a catalyst
promoter. It is unclear why the contradiction exists between these
results. The role of Br- as a promoter in ionic liquids has been
established in other systems as well, such as for the oxidation of
primary and secondary aliphatic or benzylic alcohols in
[C2C1im][BF4] using an aryliodine diacetate bound to an imidazo-
lium ion. In this case, catalysis took place with or without a metal
catalyst, indicating catalysis by bromide impurities remaining in the
ionic liquid.615 Li et al.616 examined the oxidation of 2,3,6-trimethyl-
phenol with molecular oxygen using hydrated CuCl2 in a wide variety
of ionic liquids. The best results were obtained when using
[C4C1im][BF4] (82%), with the highly coordinating [HSO3]
-,
[NO3]
-, Cl-, and Br- all yielding selectivities of less than 35%.
The ionic liquids had drastic effects, even when used as a cosolvent in
1-hexanol.617 Under the conditions of these reactions, these latter
anions will lead to the formation of different new complex ions, and it
is probably these ions that lead to the different selectivities. It is
particularly notable that [NO3]
-, Cl-, and Br-, all of which are likely
to form [CuX4]
2- ions, also required longer reaction times to achieve
complete reaction. This effect will be entirely independent of the
observed cation effect,which the authors suggestmayhave arisen from
2,3,6-trimethylphenol solubilities.
In contrast, methyltrioxorhenium/H2O2 oxidation of aromatic
aldehydes and ketones was found to exhibit little discernible anion
effect between [C4C1im][BF4] and [C4C1im][PF6].
618 In a series of
detailed studiesAbu-Omar examined the kinetics and thermodynamics
of methyltrioxorhenium/H2O2 oxidations in a variety of ionic liquids
and found that the reaction rate was independent of the ionic liquid
type when dry, but instead depended linearly on the concentration of
water in the ionic liquid.619 This observation that an impurity (water)
can be solely responsible for effects on a reaction rate, when coupled
with the observation that the concentration of that impurity in various
ionic liquids is likely to vary with some other parameter, such as anion
type in the case ofwater,26 can lead to an apparent solvent effect (anion
basicity) that is actually aproxy for the true causeof theobserved results
(water concentration). The same authors found little solvent effect
between various dry ionic liquids or CH3CN for olefin epoxidation
using the same catalyst;620 however, the addition of an equal volume of
water increased the reaction rate 10-fold. This water effect could be
particularly important when aqueous H2O2 is used as the oxidant.
In an attempt to facilitate its use as an environmentally benign
oxidant, H2O2 has been used for several catalytic oxidations in ionic
liquids.621 Aqueous (30%) H2O2 was used with [C4C1im]4-
[W10O23] in [C4C1im][BF4] to oxidize both aliphatic and benzylic
alcohols.622 The authors reported that only secondary alcohols were
oxidized under these reaction conditions and achieved the selective
transformation of 1-phenyl-1,2-diethanol into R-hydroxyacetophe-
none. This selectivity was not apparent when [C4C1im]3[PO4-
(W(O)(O2)2)4] was used as the catalyst withH2O2.
623 All selectivity
and most activity were lost when either Na2WO4 or Na2MoO4 was
instead used as the catalyst; it can be conjectured that low catalyst
solubility could be a problem for these sodium salts.
Water slightly enhanced both the rate and selectivity of OsO4-
based catalyst systems for the asymmetric dihydroxylation of olefins
in [C4C1im][PF6],
624 either alone or when tBuOH was used as a
third solvent component. Interestingly, switching to a more hydro-
phobic cation, [C8C1im]
þ, decreased the efficiency, as did using a
more hydrophilic anion, such as [BF4]
-.625 It should be noted that
although the concentration of water is often higher in these
hydrophilic ionic liquids, they have been demonstrated to interact
sufficiently strongly with water to be dehydrating agents for solutes
when dry (see above).97 Hence, these results are not contradictory.
The Mo(VI)/H2O2-catalyzed oxybromination of phenylace-
tylene proceeded faster in ionic liquids than in the organic
solvent CH2Cl2.
626 There was also a difference in selectivity
among CH2Cl2, [C4C1im][PF6], and [C4C1im][NTf2]. This is
consistent with results using vanadium(V),627 where the oxybro-
mination selectivity increased with the hydrophobicity of the
anion, from 58% in [C4C1im][BF4] to >99% in [C4C1im][PF6].
The authors attributed this result to the structure of the anion
affecting the hydration of key intermediate species.
With aqueous H2O2 oxidations observed anion effects may also
arise from changes in phase behavior. This was proposed to be the
cause of differences in the rates of the oxidation of thiols to sulfones
by acetic acid-H2O2 in [C4C1im][BF4] and [C4C1im][PF6].
628
The much slower rate of the reaction when [C4C1im][BF4] was
used was explained by the formation of a biphasic aqueous ionic
liquid/tetradecane system with [C4C1im][BF4] rather than the
triphasic [C4C1im][PF6]/tetradecane/water system. The authors
concluded that this exclusion ofwater from the ionic liquid phase led
to the large difference in reactivity.
Many aqueous H2O2 oxidations reported in the literature are in
conjunction with water-sensitive ionic liquids. The decomposition of
wet ionic liquids to release byproducts, such asHF,which can catalyze
side reactions has been pointed out for olefin epoxidations,629 where
[BF4]
- or [PF6]
- ionic liquids underperformed compared to more
hydrolytically stable ionic liquids. Problems in catalysis associatedwith
this were first noted by Dupont,630 and its effect on catalytic
oxidations has been reviewed by Muzart.601
The Baeyer-Villiger oxidation of cyclohexanone with aqueous
H2O2 can be performed in biphasicwater/ionic liquid systems using a
range of ionic liquids of various hydrophilicities.631 The catalyst in this
study was [Pt(dppb)(μ-OH)]2[BF4]2 [dppb = 1,4-bis(diphenyl-
phosphino)butane]. Interestingly, the hydrophilic [C4C1im]X {X =
[BF4] or [OTf]} produced lower yields (ca. 5% in 5 h) than the
hydrophobic organic solvents investigated (CH2Cl2 or CHCl3), but
this could be increased to an 8% yield in 1 h in [C4C1im][PF6] and

BD dx.doi.org/10.1021/cr1003248 |Chem. Rev. XXXX, XXX, 000–000
Chemical Reviews
REVIEW
increased the reaction rate considerably, with conversions after 20 h
increasing from 46.7% for [C4C1im][NTf2] to 86.5% for
[C8C1im][NTf2] or [C10C1im][PF6]. For the [C4C1im]
þ ionic
liquids the anion effectwas quite strong,with the 46.7%conversion for
[C4C1im][NTf2] dropping to6.4% in [C4C1im][PF6].This couldbe
due to some hydrolytic decomposition of the [PF6]
- or simply a
matter of the greatly lower viscosity of [C4C1im][NTf2], which
should vastly increase gas/liquid transport for this reaction. Another
possibility is that the solubility of the catalyst complex is being affected
by the choice of the ionic liquid. For the water-soluble Rh-BISBIS
complex {BISBIS = sodium salt of sulfonated 2,20-bis[(dip-
henylphosphino)methyl]-1,10-biphenyl) ligand}, the reaction selec-
tivity (for both aldehyde and linear:branched ratio) was greatly
affected by the choice of the ionic liquid. The catalyst solubility in
[C4C1im][PF6] and [C4C1im][BF4] was limited, and these solvents
yielded aldehyde selectivities of 20% and 16%, respectively, and linear:
branched ratios of 1.5 and 8.2.664 Meanwhile, the catalyst was highly
soluble in various [CnC1im][OTs] ionic liquids, and the selectivities
jumped to 82-85% and the linear:branched ratios to 34-43. A
similar anion effect was observed by the same authors when Rh-
TPPTS was employed as the catalyst.665 In this case a clear cation
effect on the reaction rate, but not on the selectivity, was also observed,
with longer alkyl chains on the cation reducing the reaction rate.
Several [NTf2]
- ionic liquids were compared to toluene for
the hydroformylation of vinyl acetate.666 The ILs outperformed
toluene in both rate and selectivity, with the linear:branched ratio
surpassing 1:99 in the case of [C8(C1)3N][NTf2]. The authors
attribute the rate and selectivity enhancements to the IL affecting
the CO insertion equilibrium step of the catalytic cycle via ionic
stabilization of the 18e- complex.
Hydroformylation of alkenes using CO2 has also been reported in
biphasic ionic liquid/organic solvent systems.667 The conversion of
1-hexene was reported to increase as the organic solvent was changed
from cyclohexane to diethyl ether to tetrahydrofuran to toluene, in
line with increased cosolvent miscibility with the IL [C4C1im]Cl.
Thus, diffusion across the phase interface (and substrate solubility) is
likely controlling the reaction efficiency, though the miscible tolue-
ne/[C8C1im]Cl mixture featured a drop in selectivity.
7.5. Dimerizations
The dimerization of propene to hexenes with nickel(II) catalysts in
acidic chloroaluminate(III) ionic liquids was one of the earliest uses of
ionic liquids668 and was described in the original review.1 This work
was further developed and led eventually to the Institut Franc-ais du
Petrole’sDifasol process.549,669Other ionic liquids have also been used
as solvents for the biphasic oligomerization of ethane, with some
success. For example, the Ni-catalyzed oligomerization has been
performed in a wide variety of [CnC1im][PF6] (n = 4, 6, 8, or 10)
ionic liquids and organic solvents, with the effectiveness of the catalytic
system being found to decrease as the alkyl chain length increased,670
which has been attributed to the higher solubility of the oligomeriza-
tion products increasing the concentration of catalyst-poisoning
internal higher olefins.12b Also, the presence of impurities (most
notably Cl- and water) had a marked effect on product selectivity.
However, the oligomerization showed higher reactivity and selectivity
in all of the ionic liquids than in conventional solvents (ionic liquids >
CH2Cl2 > n-pentane > toluene > acetone > THF = 1,4-butanediol),
and the reaction rate (TOF) correlated well with the olefin solubility
and solvatochromic polarity data.670 Similarly, the dimerization of 1,3-
butadiene by PdCl2 proceeded faster in several [C4C1im]
þ ionic
liquids than in THF,671 although there was no clear effect of changing
the anion. The issue with halide impurities can be avoided by using
Brønsted acidic ionic liquids formedbydirect protonation,672,673which
can raise the reaction rate, though whether this is a solubility
phenomenon or catalyst activation is unclear. There was also a strong
anion effect in these systems, as observed through the higher activity in
[C4C1im][OTf] compared to [C4C1im][BF4], though this is likely
due to anion complexation with the catalyst as the same effect was
observed when the catalyst anion was changed to [OTf].672 Similar
anion effects were observed for the Fe-catalyzed dimerization of 1,3-
butadiene,674 where the more strongly coordinating [C4C1im][BF4]
greatly outperformed [C4C1im][PF6], and also for the selectivity of
the hydrodimerization of 1,3-butadiene by Pd(II), which favored
[C4C1im][BF4] over [C4C1im][PF6].
675 In this instance, the selec-
tivity difference was attributed to water miscibility, which is consistent
with the findings for the cyclooligomerization of arylethynes by Ru-
porphyrins that showed a strong preference for [C8C1im][PF6] over
[C4C1im][PF6], which the authors attributed to the water content of
the ionic liquid.676
7.6. Palladium-Catalyzed Coupling Reactions
Palladium-catalyzed coupling reactions are among the most
extensively studied catalytic processes in ionic liquids.12,584
Herrmann reported that many 1,3-dialkylimidazolium ionic
liquids formed Pd-NHC species that led to a large drop in
the yield of Heck reactions when compared to tetraalkylammo-
nium cation systems.677 This negative effect has also been
observed for the Sonogashira coupling in ionic liquids,678 which
proceeded to much higher yields when a tetraalkylphosphonium
cation replaced [C4C1im]
þ. The necessary presence of a strong
base to promote these reactions can lead to other unwanted side
reactions with the ionic liquid such as NHC formation for 1,3-
dialkylimidazolium ionic liquids, on many tetraalkylammonium
species being susceptible to the Hoffman elimination. The use of
the anion as the base promoter provides an excellent opportunity
to reduce the complexity of these systems,679 but the presence of
a large amount of basic anions (such as [OAc]-) can lead to
changes in selectivity.680
Anion effects have been noted as a key parameter in determin-
ing the efficiency of ionic liquids as reaction media for the Heck
reaction. It has been noted that halide ions (frequent synthetic
impurities) can promote the reaction and improve catalyst
stability,12b with I- > Br- > Cl-.681 The Heck coupling of
electron-rich olefins with aryl halides using Pd(OAc) proceeded
to very high selectivity in [C4C1im][BF4],
682 but switching to
[C4C1im]Br increased the reaction rate, which has been attrib-
uted to the incorporation of Br- into the catalytic Pd complex.12b
This result does not appear to translate to molecular solvents,
where the addition of [BF4]
- ionic liquids to DMSO and DMF
was found to improve the selectivity of the Heck reaction, while
the addition of Br- ionic liquids did not.683
The regioselective Heck coupling of aryl bromides with allylic
alcohols has been performed in a range of ionic liquids with Pd
complexes by Hardacre.684,685 Although the activity in ionic
liquids was higher than in organic solvents, a detailed kinetic
analysis of the reaction revealed that there was an induction
period when trans-bis(2,3-dihydro-3-methylbenzothiazol-2-
ylidene)diiodopalladium(II) was used as the catalyst, leading
the authors to conclude that Pd nanoparticles were forming
during the reaction and that they were the active catalysts. The
harsh conditions required by many Heck couplings can lead to
some side reactions occurring, and in this case the relatively poor
results when [BF4]
- 683 or Br- 684,685 was used as the anion
rather than the much less basic [NTf2]
- support the possible

BE dx.doi.org/10.1021/cr1003248 |Chem. Rev. XXXX, XXX, 000–000
Chemical Reviews
REVIEW
formation of nanoparticles. It is worth noting12b that the
[NTf2]
- ionic liquids have much lower viscosities, and if the
reaction is occurring at a nanoparticle/liquid interface, then
diffusion considerations may be dominant.
Ionic liquids can also have an effect on the regioselectivity of
Heck reactions. While linear selectivity was observed for Br--
based ionic liquids,685-687 branched products predominated in
[C4C1im][BF4].
685 Molecular solvents have been reported to
yield a mixture of regioisomers.12b
Asymmetric Heck couplings of alkynes and imines have also been
carried out using ionic liquids as solvents.688The ee’swere reported to
be slightly higher in [C4C1im][NTf2] (94%) than [C4C1im][PF6]
(86%). However, the choice of cation proved crucial, as the ee
dropped to 27%when [C8C1im][PF6] was the solvent, showing that
longer alkyl chains markedly decrease selectivity, perhaps through
interaction with the hydrophobic substrates. However, this does not
appear to be a simple hydrophilicity effect, as the ee was nearly
identical when [C4C1C1
2im][BF4] (78%), [C4C1C1
2im][NTf2]
(79%), or [C4C1C1
2im][PF6] (76%) was the solvent, so specific
interactions with the cation are likely responsible. The anion had a
more marked effect on yield, as once again the more basic [BF4]
-
ionic liquid caused a significant reduction.
A cation effect has been noted in palladium-catalyzed allylic
alkylations in [C4C1im][BF4].
689 In these reactions the base for
the reaction is generated during the catalysis via [MeOCO2]
-.
However, Xiao and Ross noted that hydrogen bonding from
[C4C1im]
þ prevented this by suppressing the nucleophilicity of
the base by hydrogen bonding to it.
Suzuki-Miyaura cross-coupling reactions of halogenoarenes with
arylboronic acids have also been investigated in ionic liquids. There
has been somediscussion of the role of Pd-NHCcarbine complexes.
Welton et al.690 investigated these reactions in a range of ionic liquids
in the presence of PPh3. They found that under carefully controlled
conditions catalytically active solutions could be prepared that were
sufficiently stable to allow recycling in [CnC1im]
þ (n = 2, 4, or 6)
ionic liquids, but not when [C4C1C1
2im]þ or [C4C1pyrr]
þ ionic
liquids were used. Subsequent investigation of the reaction conditions
and analysis of the catalytically active solutions revealed the presence
of a mixed phosphine-imidazolylidene palladium complex
(Figure 13 A). They found similar results when amine ligands were
used in place of the PPh3.
691 More recently, Dyson studied the cation
effects for this reaction in detail,692 using ether-functionalized cations
to prepare ionic liquids. In these experiments catalytically active,
recyclable solutions could be formed for ionic liquids that are not
capable of forming NHCs, e.g., [(C1OC2)C1
2pyr][NTf2]. Further-
more, when [(C1OC2)C1imy]2PdCl2 (Figure 13 B) was used as the
catalyst source, the reaction yields were very low, and the authors
concluded that Pd-NHC complex formation in these ionic liquids
terminated the catalytic cycle.
Although at first sight these studies may appear to contradict each
other, the details of the experiments are quite different. In theWelton
studies, strongly coordinating PPh3 or amine coligands are added to
the solutions and a meticulous initiation procedure to form the
[(C4C1imy)(PPh3)2PdX]
þ complex is conducted.690,691 The ether
functionalities of Dyson’s ionic liquids would be expected to be less
well coordinating to palladium than these functionalities.692 Combin-
ing these results shows that ether-functionalized ionic liquids are
capable of providing palladium complexes of sufficient stability to be
able to form catalytically active and recyclable solutions. When the
[(C1OC2)C1imy]2PdCl2 complexes are added to these ionic liquids,
they are so stable as to be poor catalysts for these reactions. However,
when appropriate coordinating ligands are added into the system, the
mono-NHC complexes [(C4C1imy)(PPh3)2PdX]
þ are formed and
the solutions are again both catalytically active and recyclable.
7.7. Olefin Metathesis
The Ru-catalyzedmetathesis of methyl oleate was studied in detail
in a variety of [C4C1im]X {[NTf2], [PF6], or [BF4]} ILs. The choice
of anion showed a clear effect on the reaction rate, with the activity
decreasing slightly in the anion order [BF4]
- > [NTf2]
-
≈ [PF6]
-
for the first-generation Grubbs catalyst and [NTf2]
-
≈ [BF4]
- >
[PF6]
- for the second-generation catalyst.693 Ligand selection had a
much more marked effect on the rate. Interestingly, this reaction has
been demonstrated to show a clear solvent polarity trend inmolecular
solvents which is not present in these ionic liquids.694
7.8. Summary
The effects of ionic liquids on transition-metal-catalyzed reactions
are much less well understood than those on stoichiometric or acid-
catalyzed reactions. This is true of solvent effects in general. Much of
the difficulty in this area is in its complexity. First, the selection of
catalyst is a usually the dominating factor in determining the outcome
of a transition-metal-catalyzed reaction, and even small changes to a
ligand can lead to large changes in catalyst performance.Consequently,
unless a reaction is studied using identical catalysts, it is not possible to
even confirm whether solvent effects are occurring. Also, impurities in
the system may also have a dominating effect on the catalysis, further
confounding analysis. Even when solvent effects have been identified,
the large number of possible sources of these effects, e.g., interaction
with the catalyst, interactionwith reactants, and solubilities of reactants
and catalysts, makes it difficult to identify a sole cause, if one exists. All
of these causes have been observed in ionic liquids.
8. OVERVIEW
In the past decade there have been extraordinary advances in
the understanding of how ionic liquids can affect chemical
synthesis. This astounding class of solvents has challenged both
our experimental and our intellectual abilities. Detailed descrip-
tions of how ionic liquids interact with solute species to change
their reactivities have begun to emerge, yet much remains to be
discovered. What could be better?
ASSOCIATED CONTENT
bS Supporting Information
List of abbreviations. This material is available free of charge
via the Internet at http://pubs.acs.org.
AUTHOR INFORMATION
Corresponding Author
*E-mail: t.welton@imperial.ac.uk.
Figure 13. Phosphine-imidazolylidene palladium complex [(C4C1imy)-
(PPh3)2PdX]
þ.

BG dx.doi.org/10.1021/cr1003248 |Chem. Rev. XXXX, XXX, 000–000
Chemical Reviews
REVIEW
C. M.; Plechkova, N. V.; Seddon, K. R.; Welton, T. Anal. Chem. 2007,
79, 4247.
(30) Clare, B. R.; Bayley, P. M.; Best, A. S.; Forsyth, M.; MacFarlane,
D. R. Chem. Commun. 2008, 2689.
(31) Kagimoto, J.; Noguchi, K.; Murata, K.; Fukumoto, K.; Nakamura,
N.; Ohno, H. Chem. Lett. 2008, 37, 1026.
(32) (a) Fukumoto, K.; Yoshizawa, M.; Ohno, H. J. Am. Chem. Soc.
2005, 127, 2398. (b) Ogihara, W; Yoshizawa, M.; Ohno, H. Chem. Lett.
2004, 33, 1022.
(33) Himmler, S.; K€onig, A.;Wasserscheid, P.Green Chem. 2007, 9, 935.
(34) The synthesis and isolation of [C4C1im]OH by metathesis
from the bromide salt and KOH in dry dichloromethane has been
reported: Ranu, B. C.; Banerjee, S.Org. Lett. 2005, 7, 3049. However, we
have been unable to reproduce this preparation and to isolate pure, dry
[C4C1im]OH by any other method.
(35) Bonho^te, P.; Dias, A.-P.; Papageorgiou, N.; Kalyanasundaram,
K.; Gr€atzel, M. Inorg. Chem. 1996, 35, 1168.
(36) Cassol, C. C.; Ebling, G.; Ferrera, B.; Dupont, J. Adv. Synth.
Catal. 2006, 348, 243.
(37) Wasserscheid, P.; van Hal, R.; B€osmann, A. Green Chem. 2002,
4, 400.
(38) Holbrey, J. D.; Reichert, W. M.; Swatloski, R. P.; Broker, G. A.;
Pitner, W. R.; Seddon, K. R.; Rogers, R. D. Green Chem. 2002, 4, 407.
(39) Bradaric, C. J.; Downard, A.; Kennedy, C.; Robertson, A. J.;
Zhou, Y. Green. Chem. 2003, 5, 143.
(40) Bowing, A. G.; Jess, A.; Wasserscheid, P.Chem.-Ing.-Tech. 2005,
77, 1430.
(41) Himmler, S.;H€ormann, S.; vanHal, R.; Schulz, P. S.;Wasserscheid,
P.Green Chem. 2006, 8, 887.
(42) Holbrey, J. D.; Reichert, W. M.; Tkatchenko, I.; Bouajila, E.;
Walter, O.; Tommasi, I.; Rogers, R. D. Chem. Commun. 2003, 28.
(43) Smiglak, M; Holbrey, J. D.; Griffin, S. T.; Reichert, W. M.;
Swatloski, R. P.; Katritzky, A. R.; Yang, H.; Zhang, D.; Kirichenko, K.;
Rogers, R. D. Green Chem. 2007, 9, 90.
(44) Bridges, N. J.; Hines, C. C.; Smiglak, M.; Rogers, R. D.Chem.—
Eur. J. 2007, 13, 5207.
(45) Maase, M.; Massonne, K. U.S. Patent US 2006/0149074 A1,
2006.
(46) Linder, T.; Sundermeyer, J. Chem. Commun. 2009, 2914.
(47) Cole, A. C.; Jensen, J. L.; Ntai, I; Tran, K. L. T.; Weaver, K. J.;
Forbes, D. C.; Davis, J. H. J. Am. Chem. Soc. 2002, 124, 5962.
(48) Yoshizawa, M.; Hirao, M.; Ito-Akita, K.; Ohno, H. J. Mater.
Chem. 2001, 11, 1057.
(49) Ohno, H.; Yoshizawa, M. Solid State Ionics 2002, 154, 303.
(50) Yoshizawa, M.; Xu, W.; Angell, C. A. J. Am. Chem. Soc. 2003,
125, 15411.
(51) Maase, M.; Massonne, K.; Halbritter, K.; Noe, R.; Bartsch, M.;
Siegel, W.; Stegmann, V.; Flores, M.; Huttenloch O.; Becker, M. World
Patent WO 2003 062171, 2003.
(52) (a) Swatloski, R. P.; Holbrey, J. D.; Rogers, R. D. Green Chem.
2003, 5, 361. (b) Cho, C.-W.; Pham, T. P. T.; Jeon, Y.-C.; Yun, Y.-S.
Green Chem. 2008, 10, 67.
(53) Freire, M. G.; Santos, L. M. N. B. F.; Fernandes, A. M.;
Coutinho, J. A. P.; Marrucho, I. M. Fluid Phase Equilib. 2007, 261, 449.
(54) Freire, M. G.; Carvalho, P. J.; Gardas, R. L.; Marrucho, I. M.;
Santos, L.M. N. B. F.; Coutinho, J. A. P. J. Phys. Chem. B 2008, 112, 1604.
(55) Seddon, K. R.; Stark, A.; Torres, M. J. Pure Appl. Chem. 2000,
72, 2275.
(56) (a) Shah, J. K.; Maginn, E. J. J. Phys. Chem. B 2005, 109, 10395.
(b) Anthony, J. L.; Anderson, J. L.; Maginn, E. J.; Brennecke, J. F. J. Phys.
Chem. B 2005, 109, 6366. (c) Buzzeo, M. C.; Klymenko, O. V.;
Wadhawan, J. D.; Hardacre, C.; Seddon, K. R.; Compton, R. G.
J. Phys. Chem. A 2003, 107, 8872. (d) Husson-Borg, P.; Majer, V.;
Gomes, M. F. C. J. Chem. Eng. Data 2003, 48, 480.
(57) (a) Pham, T. P. T.; Cho, C.-W.; Yun, Y.-S. Water Res. 2010,
44, 352. (b) Zhao, D. B.; Liao, Y. C.; Zhang, Z. D. Clean 2007, 35, 42.
(58) Luis, P.; Ortiz, I.; Aldaco, R.; Irabien, A. Ecotoxicol. Environ. Saf.
2007, 67, 423.
(59) Couling, D. J.; Bernot, R. J.; Docherty, K. M.; Dixon, J. K.;
Maginn, E. J. Green Chem. 2006, 8, 82.
(60) García-Lorenzo, A.; Tojo, E.; Teijeira, M.; Rodríguez-Berrocal,
F. J.; Gonzalez,M. P.;Martínez-Zorano, V. S.Green Chem. 2008, 10, 508.
(61) Schaffran, T.; Justus, E.; Elfert, M.; Chen, T.; Gabel, D. Green
Chem. 2009, 11, 1458.
(62) Ranke, J.; M€uller, A.; Bottin-Weber, U.; Stock, F.; Stolte, S.;
Arning, J.; St€ormann, R.; Jastorff, B. Ecotoxicol. Environ. Saf. 2007, 67, 430.
(63) Landry, T. D.; Brooks, K.; Poche, D.; Woolhiser, M. Bull.
Environ. Contam. Toxicol. 2005, 74, 559.
(64) Hough, W. L.; Smiglak, M.; Rodriguez, H.; Swatloski, R. P.;
Spear, S. K.; Daly, D. T.; Pernak, J.; Grisel, J. E.; Carliss, R. D.; Soutullo,
M. D.; Davis, J. H.; Rogers, R. D. New J. Chem. 2007, 31, 1429.
(65) (a) Zhang, S.; Sun, N.; He, X.; Lu, X.; Zhang, X. J. Phys. Chem.
Ref. Data 2006, 35, 1475. (b) Galinski, M; Lewandowski, A; Stepniak, I.
Electrochim. Acta 2006, 51, 5567.
(66) Freemantle, M. Chem. Eng. News 1998, 76 (13), 32.
(67) Eike, D. M.; Brennecke, J. F.; Maginn, E. J. Green Chem. 2003,
5, 323.
(68) Varnek, A.; Kireeva, N.; Tetko, I. V.; Baskin, I. I.; Solov’ev, V.
J. Chem. Inf. Model. 2007, 47, 1111.
(69) Holbrey, J. D.; Reichert, W.M.; Nieuwenhuyzen, M.; Johnston,
S.; Seddon, K. R.; Rogers, R. D. Chem. Commun. 2003, 1636.
(70) Krossing, I.; Slattery, J. M.; Daguenet, C.; Dyson, P. J.; Oleinikova,
A.; Weing€artner, H. J. Am. Chem. Soc. 2006, 128, 13427.
(71) Markusson, H.; Belieres, J.-P.; Johansson, P.; Angell, C. A.;
Jacobsson, P. J. Phys. Chem. A 2007, 111, 8717.
(72) Belieres, J.-P.; Angell, C. A. J. Phys. Chem. B 2007, 111, 4926.
(73) Luo, H.; Baker, G. A.; Lee, J. S.; Pagni, R. M.; Dai, S. J. Phys.
Chem. B 2009, 113, 4181.
(74) Kroon, M. C.; Buijs, W.; Peters, C. J.; Witkamp, G.-J. Thermo-
chim. Acta 2007, 465, 40.
(75) Baranyai, K. J.; Deacon, G. B.; MacFarlane, D. R.; Pringle, J. M.;
Scott, J. L. Aust. J. Chem. 2004, 57, 145.
(76) Deetlefs, M.; Seddon, K. R.; Shara, M. Phys. Chem. Chem. Phys.
2005, 8, 642.
(77) Slattery, J. M.; Daguenet, C.; Dyson, P. J.; Schubert, T. J. S.;
Krossing, I. Angew. Chem., Int. Ed. 2007, 46, 5384.
(78) Jacquemin, J.; Ge, R.; Nancarrow, P.; Rooney, D. W.; Gomes,
M. F. C.; Padua, A. A. H.; Hardachre, C. J. Chem. Eng. Data 2008,
53, 716.
(79) Gardas, R. L.; Coutinho, J. A. P. Fluid Phase Equilib. 2008,
265, 57.
(80) Palomar, J.; Ferro, V. R.; Torrecilla, J. S.; Rodriguez, F. Ind. Eng.
Chem. Res. 2007, 46, 6041.
(81) Schr€oder, U.; Wadhawn, J. S.; Compton, R. G.; Marken, F.;
Suarez, P. A. Z.; Consorti, C. S.; de Souza, R. F.; Dupont, J. New J. Chem.
2000, 24, 1009.
(82) Freire, M. G.; Carvalho, P. J.; Fernandes, A. M.; Marrucho,
I. M.; Queimada, A. J.; Coutinho, J. A. P. J. Colloid Interface Sci. 2007,
314, 621.
(83) Fletcher, K. A.; Pandey, S. J. Phys. Chem. B 2003, 107, 13532.
(84) Huddleston, J. G.; Visser, A. E.; Reichardt, W. M.; Willauer,
H. D.; Broker, G. A.; Rogers, R. D. Green Chem. 2001, 3, 156.
(85) Dyson, P. J.; Ellis, D. J.; Welton, T. Can. J. Chem. 2001, 79, 705.
(86) Schrekker, H. S.; Stracke, M. P.; Schrekker, C. M. L.; Dupont, J.
Ind. Eng. Chem. Res. 2007, 46, 7389.
(87) Yasuda, T.; Ogawa, A.; Kanno, M.; Mori, K.; Sakakibara, K.;
Watanabe, M. Chem. Lett. 2009, 37, 692.
(88) (a) Miskolczy, Z.; Seb€oknagy, K.; Biczok, L.; G€okt€urk, S. Chem.
Phys. Lett. 2004, 400, 296. (b) Goodchild, I.; Collier, L.; Millar, S. L.;
Prokes, I.; Lord, J. C. D.; Butts, C. P.; Bowers, J.; Webster, J. R. P.;
Heenan, R. K. J. Colloid Interface Sci. 2007, 307, 455. (c) Modaressi, A.;
Sifaoui, H.; Mielcarz, M.; Domanska, U.; Rogalski, M. Colloids Surf., A
2007, 302, 181.
(89) (a) Firestone, M. A.; Rickert, P. G.; Seifert, S.; Dietz, M. L.
Inorg. Chim. Acta 2004, 357, 3991. (b) Gaillon, L.; Sirieix-Plenet, J.;
Letellier, P. J. Solution Chem. 2004, 33, 1333.

BH dx.doi.org/10.1021/cr1003248 |Chem. Rev. XXXX, XXX, 000–000
Chemical Reviews
REVIEW
(90) (a) Kaar, J. L.; Jesionowski, A. M.; Berberich, J. A.; Moulton, R.;
Russell, A. J. J. Am. Chem. Soc. 2003, 125, 4125. (b) Ropel, L.; Belveze,
L. S.; Aki, S. N. V. K.; Stadtherr, M. A.; Brennecke, J. F. Green Chem.
2005, 7, 83. (c) Chapeaux, A.; Simoni, L. D.; Stadtherr, M. A.;
Brennecke, J. F. J. Chem. Eng. Data 2007, 52, 2462. (d) Zhao, H.; Baker,
G. A.; Song, Z.; Olubajo, O.; Zanders, L.; Campbell, S. M. J. Mol. Catal. B
2009, 57, 149.
(91) Kato, H; Miki, K.; Mukai, T.; Nishikawa, K.; Koga, Y. J. Phys.
Chem. B 2009, 113, 14754.
(92) (a) Weng, J.; Wang, C.; Li, H.; Wang, Y. Green Chem. 2006,
8, 96. (b) Li, C.; Zhao, Z. K.Adv. Synth. Catal. 2007, 349, 1847. (c) Li, C.;
Wang, Q.; Zhao, Z. K. Green Chem. 2008, 10, 177.
(93) Kim, D. W.; Hong, D. J.; Seo, J. W.; Kim, H. S.; Kim, H. K.;
Song, C. E.; Chi, D. Y. J. Org. Chem. 2004, 69, 3186.
(94) Farmer, V.; Welton, T. Green Chem. 2002, 4, 97.
(95) Seddon, K. R.; Stark, A. Green Chem. 2002, 4, 119.
(96) Ansari, I. A.; Gree, R. Org. Lett. 2002, 4, 1507.
(97) Amigues, E.; Hardacre, C.; Keane, G.; Migaud, M.; O’Neill, M.
Chem. Commun. 2006, 72.
(98) Sanders, M. W.; Wright, L.; Tate, L.; Fairless, G.; Chrowhurst,
L.; Bruce, N. C.;Walker, A. J.; Hembury, G. A.; Shimizu, S. J. Phys. Chem.
A 2009, 113, 10143.
(99) Jork, C.; Seiler, M.; Beste, Y.-A.; Arlt, W. J. Chem. Eng. Data
2004, 49, 852.
(100) Tran, C.d.; De Paoli Larcerda, S. H.; Oliveira, D. Appl.
Spectrosc. 2003, 57, 152.
(101) K€oddermann, T.; Wertz, C.; Heintz, A.; Ludwig, R. Angew.
Chem., Int. Ed. 2006, 45, 3697.
(102) (a) Lopez-Pastor, M.; Ayora-Ca~nada, M. J.; Valcarcel, M.;
Lendl, B. J. Phys. Chem. B 2006, 110, 10896. (b) Dominguez-Vidal, A.;
Kaun, N.; Ayora-Ca~nada, M. J.; Lendl, B. J. Phys. Chem. B 2007,
111, 4446. (c) Jeon, Y.; Sung, J.; Kim, D.; Seo, C.; Cheong, H.; Ouchi,
Y.; Ozawa, R.; Hamaguchi, H. J. Phys. Chem. B 2008, 112, 923. (d) Fazio,
B.; Triolo, A.; Di Marco, G. J. Raman Spectrosc. 2008, 39, 233.
(103) (a) Chang, H.-C.; Jiang, J.-C.; Liou, Y.-C.; Hung, C.-H.; Lai,
T.-Y.; Lin, S. H. J. Chem. Phys. 2008, 129, 44506. (b) Chang, H.-C.; Jiang,
J.-C.; Liou, Y.-C.; Hung, C.-H.; Lai, T.-Y.; Lin, S. H. Anal. Sci. 2008,
24, 1305.
(104) Zhang, L.; Xu, Z.; Wang, Y.; Li, H. J. Phys. Chem. B 2008,
112, 6411.
(105) Sun, B.; Jin, Q.; Tan, L.; Wu, P.; Yan, F. J. Phys. Chem. B 2008,
112, 14251.
(106) Mele, A.; Tran, C. D.; De Paoli Lacerda, S. H. Angew. Chem.,
Int. Ed. 2003, 42, 4364.
(107) Taga, T.; Machida, K.; Kimura, N.; Hayashi, S.; Umemura, J.;
Takenaka, T. Acta Crystallogr. 1987, C43, 1204.
(108) Taga, T.; Machida, K.; Kimura, N.; Hayashi, S.; Umemura, J.;
Takenaka, T. Acta Crystallogr. 1986, C42, 608.
(109) Downard, A.; Earle, M. J.; Hardacre, C.; McMath, S. E. J.;
Nieuwenhuyzen, M.; Teat, S. J. Chem. Mater. 2004, 16, 43.
(110) Hanke, C. G.; Lynden-Bell, R. M. J. Phys. Chem. B 2003,
107, 10873.
(111) (a) Schr€oder, C.; Rudas, T.; Neumayr, G.; Benkner, S.;
Steinhauser, O. J. Chem. Phys. 2007, 127, 234503. (b) Porter, A. R.;
Liem, S. Y.; Popelier, P. L. A. Phys. Chem. Chem. Phys. 2008, 10, 4240. (c)
Schr€oder, C.; Neumayr, G.; Steinhauser, O. J. Chem. Phys. 2009,
130, 194503.
(112) Moreno, M.; Castiglone, F.; Mele, A.; Pasqui, C.; Raos, G.
J. Phys. Chem. B 2008, 112, 7826.
(113) Wang, Y.; Li, H. R.; Han, S. J. Phys. Chem. B 2006, 110, 24646.
(114) Spickermann, C.; Thar, J.; Lehmann, S . B. C; Zahn, S.;
Hunger, J.; Buchner, R; Hunt, P. A.; Welton, T.; Kirchner, B. J. Chem.
Phys. 2008, 129, 104505.
(115) Reichardt, C. Solvents and Solvent Effects in Organic Chemistry,
3rd ed.; Wiley-VCH: Weinheim, Germany, 2003.
(116) M€uller, P. Pure Appl. Chem. 1994, 66, 1077.
(117) Katritzky, A. R.; Fara, D. C.; Yang, H.; T€amm, K. Chem. Rev.
2004, 104, 175.
(118) (a) Weinga1rtner, H.; Knocks, A.; Schrader, W.; Kaatze, U.
J. Phys. Chem. A 2001, 105, 8646. (b) Wakai, C; Oleinikova, A; Ott, M;
Weing€artner, H. J. Phys. Chem. B 2005, 109, 17028. (c) Daguenet, C.;
Dyson, P. J.; Krossling, I.; Oleinkova, A.; Slattery, J.; Wakai, C.;
Weing€artner, H. J. Phys. Chem. B 2006, 110, 12682. (d) Weing€artner,
H. Z. Phys. Chem. (Muenchen) 2006, 220, 1395. (e) Weing€artner, H.
Angew. Chem., Int. Ed. 2008, 47, 654. (f) Huang, M.-M.; Weing€artner, H.
ChemPhysChem 2008, 9, 2172.
(119) Singh, T.; Kumar, A. J. Phys. Chem. B 2008, 112, 12968.
(120) Baker, S. N.; Baker, G. A.; Kane, M. A.; Bright, F. V. J. Phys.
Chem. B 2001, 105, 9663.
(121) Earle, M. J.; Esperanc-a, J. M. S. S.; Gilea, M. A.; Canongia
Lopes, J. N.; Rebelo, L. P. N.; Magee, J. W.; Seddon, K. R.; Widegren,
J. A. Nature 2006, 439, 831.
(122) Luo, H.; Baker, G. A.; Dai, S. J. Phys. Chem. B 2008,
112, 10077.
(123) Jin, H.; O’Hare, B.; Dong, J.; Arzhantsev, S.; Baker, G. A.;
Wishart, J. F.; Benesi, A. J.; Maroncelli, M. J. Phys. Chem. B 2008, 112, 81.
(124) Lee, S. H.; Lee, S. B. Chem. Commun. 2005, 3469.
(125) Kilaru, P. K.; Scovazzo, P. Ind. Eng. Chem. Res. 2008, 47, 910.
(126) Swiderski, K.; McLean, A.; Gordon, C. M.; Vaughan, D. H.
Chem. Commun. 2004, 2178.
(127) DaSilveiro Neto, B. A.; Santos, L. S.; Nachtigall, F. M.; Eberlin,
M. N.; Dupont, J. Angew. Chem., Int. Ed. 2006, 45, 7251.
(128) Carmichael, A. J.; Seddon, K. R. J. Phys. Org. Chem. 2000,
13, 591.
(129) Ogihara, W.; Aoyama, T.; Ohno, H. Chem. Lett. 2004, 33, 1414.
(130) Chen, H.; Kwait, D. C.; G€onen, Z. S.; Weslowski, B. T.;
Abdallah, D. J.; Weiss, R.; G. Chem. Mater. 2002, 14, 4063.
(131) Dzyuba, S. V.; Bartsch, R. A. Tetrahedron Lett. 2002, 43,
4657.
(132) Moog, R. S.; Kim, D. D.; Oberle, J. J.; Ostrowski, S. G. J. Phys.
Chem. 2004, 108, 9294.
(133) (a) Reichardt, C. Chem. Soc. Rev. 1992, 147. (b) Reichardt, C.
Pure Appl. Chem. 2004, 76, 1903.
(134) Taft, R. W.; Kamlet, M. J. J. Am. Chem. Soc. 1976, 98, 2886.
(135) Reichardt, C. Green Chem. 2005, 7, 339.
(136) Wu, Y.; Sasaki, T.; Kazushi, K.; Seo, T.; Sakurai, K. J. Phys.
Chem. B 2008, 112, 7530.
(137) Drummond, C. J.; Furlong, D. N. J. Chem. Soc., Faraday Trans.
1990, 86, 3613.
(138) Zhang, S.; Zhang, Q.; Ye, B. X.; Zhang, X.; Deng, Y. J. Phys.
Chem. 2009, 113, 6012.
(139) Byrne, R.; Fraser, K. J.; Izgorodina, E.; MacFarlane, D. R.;
Forsyth, M.; Diamond, D. Phys. Chem. Chem. Phys. 2008, 10, 5919.
(140) Muldoon, M. J.; Gordon, C. M.; Dunkin, I. R. J. Chem. Soc.,
Perkin Trans. 2 2001, 433.
(141) Crowhurst, L.; Mawdsley, P. R.; Perez-Arlandis, J. M.; Salter,
P. A.; Welton, T. Phys. Chem. Chem. Phys. 2003, 5, 2790.
(142) Linert, W.; Jameson, R. F.; Taha, A. J. Chem. Soc., Dalton Trans.
1993, 3181.
(143) Ueno, K.; Imaizumi, S.; Hata, K.; Watanabe, M. Langmuir
2009, 25, 825.
(144) Gutmann, V. The Donor-Acceptor Approach to Molecular
Interactions; Plenum Press: New York, 1978.
(145) (a) Kimura, Y.; Fukuda, M.; Fujisawa, T.; Terazima, M. Chem.
Lett. 2005, 34, 338. (b) Fujisawa, T.; Fukuda, M.; Terazima, M.; Kimura,
Y. J. Phys. Chem. A 2006, 110, 6164.
(146) Cha, D. K.; Kloss, A. A.; Tikanen, A. C.; Fawcett, W. R. Phys.
Chem. Chem. Phys. 1999, 1, 4785.
(147) Tao, G.-H.; Zou, M.; Wang, X.-h.; Chen, Z.-y.; Evans, D. G.;
Kou., Y. Aust. J. Chem. 2005, 58, 327.
(148) Kawai, A.; Hidemori, T.; Shibuya, K. Chem. Lett. 2004,
33, 1464.
(149) Kawai, A.; Hidemori, T.; Shibuya, K. Chem. Phys. Lett. 2005,
414, 378.
(150) Street, K. W., Jr.; Acree, W. E., Jr.; Fetzer, J. C.; Shetty, P. H.;
Poole, C. F. Appl. Spectrosc. 1989, 43, 1149.

BJ dx.doi.org/10.1021/cr1003248 |Chem. Rev. XXXX, XXX, 000–000
Chemical Reviews
REVIEW
(200) Safarov, J.; Verevkin, S. P.; Bich, E.; Heintz, A. J. Chem. Eng.
Data 2006, 51, 518.
(201) Heintz, A.; Casas, L. M.; Nesterov, I. A.; Emel’yanenko, V. N.;
Verevkin, S. P. J. Chem. Eng. Data 2005, 50, 1510.
(202) Heintz, A.; Verevkin, S. P.; Lehmann, J. K.; Vasiltsova, T. V.;
Ondo, D. J. Chem. Thermodyn. 2007, 39, 268.
(203) Shimoyama, Y.; Hirayama, T.; Iwai, Y. J. Chem. Eng. Data
2008, 53, 2106.
(204) Krummen, M.; Wasserscheid, P.; Gmehling, J. J. Chem. Eng.
Data 2002, 47, 1411.
(205) Anthony, J. L.; Maginn, E. J.; Brennecke, J. F. J. Phys. Chem. B
2001, 105, 10942.
(206) Heintz, A. J. Chem. Thermodyn. 2005, 37, 525.
(207) Jork, C.; Kristen, C.; Pieraccini, D.; Stark, A.; Chiappe, C.;
Beste, Y. A.; Arlt, W. J. Chem. Thermodyn. 2005, 37, 537.
(208) Dobryakov, Y. G.; Tuma, D.; Maurer, G. J. Chem. Eng. Data
2008, 53, 2154.
(209) Letcher, T. M.; Soko, B.; Ramjugernath, D.; Deenadayalu, N.;
Nevines, A.; Naicker, P. K. J. Chem. Eng. Data 2003, 48, 708.
(210) Letcher, T. M.; Soko, B.; Reddy, P.; Deenadayalu, N. J. Chem.
Eng. Data 2003, 48, 1587.
(211) David, W.; Letcher, T. M.; Ramjurgernath, D.; Raal, J. D.
J. Chem. Thermodyn. 2003, 35, 1335.
(212) Letcher, T. M.; Domanska, U.; Marciniak, M.; Marciniak, A.
J. Chem. Thermodyn. 2005, 37, 587.
(213) Letcher, T. M.; Marciniak, A.; Marciniak, M.; Domanska, U.
J. Chem. Thermodyn. 2005, 37, 1327.
(214) Domanska, U.; Marciniak, A. J. Phys. Chem. B 2008,
112, 11100.
(215) Ge, M.-L.; Wang, L.-S. J. Chem. Eng. Data 2008, 53, 846.
(216) Ge, M.-L.; Wang, L.-S.; Wu, J.-S.; Zhou, Q. J. Chem. Eng. Data
2008, 53, 1970.
(217) Ge, M.-L.; Wu, J.-S.; Wang, M.-H.; Wang, L.-S. J. Chem. Eng.
Data 2008, 53, 871.
(218) Yang, X.-J.; Wu, J.-S.; Ge,M.-L.;Wang, L.-S.; Li,M.-Y. J. Chem.
Eng. Data 2008, 53, 1220.
(219) Ge, M.-L.; Wang, L.-S.; Li, M.-Y.; Wu, J.-S. J. Chem. Eng. Data
2007, 52, 2257.
(220) Letcher, T. M.; Marciniak, A.; Marciniak, M.; Domanska, U.
J. Chem. Eng. Data 2005, 50, 1294.
(221) Mutelet, F.; Jaubert, J.-N. J. Chromatogr., A 2006, 1102
256.
(222) Domanska, U.; Marciniak, A. J. Phys. Chem. B 2007,
111, 11984.
(223) Domanska, U.; Marciniak, A. J. Chem. Thermodyn. 2008,
40, 860.
(224) Deenadayalu, N.; Thango, S. H.; Letcher, T. M.; Ramjugernath,
D. J. Chem. Thermodyn. 2006, 38, 542.
(225) Heintz, A.; Verevkin, S. P.; Ondo, D. J. Chem. Eng. Data 2006,
51, 434.
(226) Heintz, A.; Verevkin, S. P. J. Chem. Eng. Data 2005, 50, 1515.
(227) Foco, G. M.; Bottini, S. B.; Quezada, N.; de la Fuente, J. C.;
Peters, C. J. J. Chem. Eng. Data 2006, 51, 1088.
(228) Inoue, G.; Iwai, Y.; Yasutake, M.; Honda, K.; Arai, Y. Fluid
Phase Equilib. 2007, 251, 17.
(229) Mutelet, F.; Jaubert, J.-N. J. Chem. Thermodyn. 2007, 39, 1144.
(230) Wang, M.-H.; Wu, J.-S.; Wang, L.-S.; Li, M.-Y. J. Chem. Eng.
Data 2007, 52, 1488.
(231) Heintz, A.; Vasiltsova, T. V.; Safarov, J.; Bich, E.; Verevkin,
S. P. J. Chem. Eng. Data 2006, 51, 648.
(232) Letcher, T. M.; Reddy, P. Fluid Phase Equilib. 2005, 235, 11.
(233) Banerjee, T.; Khanna, A. J. Chem. Eng. Data 2006, 51, 2170.
(234) Letcher, T. M.; Reddy, P. Fluid Phase Equilib. 2007, 260, 23.
(235) Letcher,T.M.;Ramjugernath,D.;Laskowska,M.;Krolikowski,
M.; Naidoo, P.; Domanska, U. J. Chem. Eng. Data 2008, 53, 2044.
(236) Letcher, T. M.; Ramjugernath, D.; Laskowska, M.;
Krolikowski, M.; Naidoo, P.; Domanska, U. J. Chem. Thermodyn. 2008,
40,1243.
(237) Mutelet, F.; Jaubert, J.-N.; Rogalski, M.; Boukherissa, M.;
Dicko, A. J. Chem. Eng. Data 2006, 51, 1274.
(238) Mutelet, F.; Jaubert, J.-N.; Rogalski, M.; Harmand, J.; Sindt,
M.; Mieloszynski, J.-L. J. Phys. Chem. B 2008, 112, 3773.
(239) Kato, R.; Gmehling, J. J. Chem. Thermodyn. 2005, 37, 603.
(240) Diedenhofen, M.; Eckert, F.; Klamt, A. J. Chem. Eng. Data
2003, 48, 475.
(241) Klamt, A. J. Phys. Chem. 1995, 99, 2224.
(242) Kato, R.; Gmehling, J. Fluid Phase Equilib. 2004, 226, 37.
(243) Eike, D. M.; Brennecke, J. F.; Maginn, E. J. Ind. Eng. Chem. Res.
2004, 43, 1039.
(244) T€amm, K.; Burk, P. J. Mol. Model. 2006, 12, 417.
(245) Verevkin, S. P.; Vasiltsova, T. V.; Bich, E.; Heintz, A. Fluid
Phase Equilib. 2004, 218, 165.
(246) Verevkin, S. P.; Safarov, J.; Bich, E.; Hasel, E.; Heintz, A. Fluid
Phase Equilib. 2005, 236, 222.
(247) Mutelet, F.; Butet, V.; Jaubert, J.-N. Ind. Eng. Chem. Res. 2005,
44, 4120.
(248) Wang, J.; Sun, W.; Li, C.; Wang, Z. Fluid Phase Equilib. 2008,
264, 235.
(249) Alvaro, M.; Ferrer, B.; García, H.; Narayana, M. Chem. Phys.
Lett. 2002, 362, 435.
(250) Grodkowski, J.; Neta, P. J. Phys. Chem. A 2002, 106, 5468.
(251) Gordon, C. M.; McLean, A. J. Chem. Commun. 2000, 1395.
(252) Alvaro, M.; Carbonell, E.; Ferrer, B.; Garcia, H.; Herance, J. R.
Photochem. Photobiol. 2006, 82, 185.
(253) Grodkowski, J.; Neta, P. J. Phys. Chem. A 2002, 106, 11130.
(254) Takahashi, K.; Sakai, S; Tezuka, H.; Hiejima, Y.; Katsumura,
Y.; Watanabe, M. J. Phys. Chem. B 2007, 111, 4807.
(255) McLean, A. J.; Muldoon, M. J.; Gordon, C. M.; Dunkin, I. R.
Chem. Commun. 2002, 1880.
(256) Skrzypczak, A.; Neta, P. J. Phys. Chem. A 2003, 107, 7800.
(257) Vieira, R. C.; Falvey, D. E. J. Phys. Chem. B 2007, 111, 5023.
(258) Paul, A.; Samanta, A. J. Phys. Chem. B 2007, 111, 1957.
(259) Sarkar, S.; Pramanik, R.; Seth, D.; Setua, P.; Sarkar, N. Chem.
Phys. Lett. 2009, 477, 102.
(260) Strehmel, V.; Wishart, J. F.; Polyansky, D. E.; Strehmel, B.
ChemPhysChem 2009, 10, 3112.
(261) Evans, R. G.; Klymenko, O. V.; Price, P. D.; Davies, S. G.;
Hardacre, C.; Compton, R. G. ChemPhysChem 2005, 6, 526.
(262) Padua, A. A. H.; Costa Gomes, M. F.; Canongia Lopes, J. N. A.
Acc. Chem. Res. 2007, 40, 1087 and references therein.
(263) Crowhurst, L.; Lancaster, N. L.; Perez Arlandis, J. M.; Welton,
T. J. Am. Chem. Soc. 2004, 126, 11549.
(264) Wishart, J. F.; Neta, P. J. Phys. Chem. B 2003, 107, 7261.
(265) Behar, D.; Neta, P.; Schulteisz, C. J. Phys. Chem. A 2002,
106, 3139.
(266) Vieira, R. C.; Falvey, D. E. J. Am. Chem. Soc. 2008, 130, 1552.
(267) Wislicenus, W. Liebigs Ann. Chem. 1896, 291, 147.
(268) Earle, M. J.; Engel, B. S.; Seddon, K. R. Aust. J. Chem. 2004,
57, 149.
(269) Angelini, G.; Chiappe, C.; De Maria, P.; Fontana, A.; Gasparrini,
F.; Pieraccini, D.; Pierini, M.; Siani, G. J. Org. Chem. 2005, 70, 8193.
(270) Johnson, K. E.; Pagni, R. M.; Bartmess, J. Monatsh. Chem.
2007, 138, 1077.
(271) Yang, Y.-L.; Kou, Y. Chem. Commun. 2004, 226.
(272) Zhao, Z.; Yuan, B.; Qiao, W.; Li, Z.; Wang, G.; Cheng, L.
J. Mol. Catal. A 2005, 235, 74.
(273) Xin, H. L.; Wu, Q.; Han, M.; Wang, D.; Jin, Y. Appl. Catal., A
2005, 292, 354.
(274) (a) Hammett, L. P.; Deyrup, A. J. J. Am. Chem. Soc. 1932,
54, 2721. (b) Hammett, L. P.; Dingwall, A.; Flexser, L. J. Am. Chem. Soc.
1934, 56, 2010.
(275) Gu, Y.; Zhang, J.; Duan, Z.; Deng, Y. Adv. Synth. Catal. 2005,
347, 512.
(276) Wang, Y.-Y.; Li, W.; Dai, L.-Y. Chin. J. Chem. 2008, 26, 1390.
(277) Cheng, G.; Duan, X.; Qi, X.; Lu, C. Catal. Commun. 2008,
10, 201.

BK dx.doi.org/10.1021/cr1003248 |Chem. Rev. XXXX, XXX, 000–000
Chemical Reviews
REVIEW
(278) Liu, X.; Zhou, J.; Guo, X.; Liu, M.; Ma, X.; Song, C.; Wang, C.
Ind. Eng. Chem. Res. 2008, 47, 5298.
(279) Xing, H.; Wang, T.; Zhou, Z.; Dai, Y. J. Mol. Catal. A 2007,
264, 53.
(280) Wang, Y.; Jiang, D; Dai, L. Catal. Commun. 2008, 9, 2475.
(281) Liu, X.; Liu, M.; Guo, X.; Zhou, J. Catal. Commun. 2008, 9, 1.
(282) Du, Y.; Tian, F. J. Chem. Res. 2006, 486.
(283) Duan, Z.; Gu, Y.; Zhang, J.; Zhu, L.; Deng, Y. J. Mol. Catal. A
2006, 250, 163.
(284) Du, Z.; Li, Z.; Guo, S.; Zhang, J.; Zhu, L.; Deng, Y. J. Phys.
Chem. B 2005, 109, 19542.
(285) Liu, X.-M.; Song, Z.-X.; Wang, H.-J. Struct. Chem. 2009,
20, 509.
(286) (a) Thomazeau, C.; Olivier-Bourbigou, H.; Magna, L.; Luts,
S.; Gilbert, B. J. Am. Chem. Soc. 2003, 125, 5264. (b) Robert, T.; Magna,
L.; Olivier-Bourbigou, H.; Gilbert, B. J. Electrochem. Soc. 2009,
156, F115.
(287) D’Anna, F.; La Marca, S.; Noto, R. J. Org. Chem. 2009,
74, 1952.
(288) Del Popolo, M. G.; Kohanoff, J.; Lynden-Bell, R. M. J. Phys.
Chem. B 2006, 110, 8798.
(289) Trulove, P. C; Osteryoung, R. A. Inorg. Chem. 1992, 31, 3980.
(290) Olah, G. A.; Mathew, T.; Goeppert, A.; T€or€ok, B.; Bucsi, I.; Li,
X.-Y.; Wang, Q.; Marinez, E. R.; Batamack, P.; Aniszfeld, R.; Prakash,
G. K. S. J. Am. Chem. Soc. 2005, 127, 5964.
(291) Koppel, I. J. Am. Chem. Soc. 2000, 122, 5114.
(292) Emsley, J. Chem. Soc. Rev. 1980, 9, 91.
(293) Nair, V.; Bindu, S.; Sreekumar, V. Angew. Chem., Int. Ed. 2004,
43, 5130.
(294) Enders, D.; Niemeier, O.; Hensler, A. Chem. Rev. 2007,
107, 5606.
(295) (a) Connor, E. F.; Nyce, G. W.; Mock, A.; Hedrick, J. L. J. Am.
Chem. Soc. 2002, 124, 914. (b) Nyce, G. W.; Lamboy, J. A.; Connor,
E. F.;Waymouth, R.M.; Hedrick, J. L.Org. Lett. 2002, 4, 3587. (c) Grasa,
G. A.; Kissling, R. M.; Nolan, S. P. Org. Lett. 2002, 4, 3583. (d) Grasa,
G. A.; Guveli, T.; Singh, R.; Nolan, S. P. J. Org. Chem. 2003, 68, 2812.
(296) Handy, S. T.; Okello, M. J. Org. Chem. 2005, 70, 1915.
(297) (a) Handy, S. C. J. Org. Chem. 2006, 71, 4659. (b) Jurcik, V.;
Wilhelm, R. Green Chem. 2005, 7, 844.
(298) (a) Murray, C. B.; Sandford, G.; Korn, S. R. J. Fluorine Chem.
2003, 123, 81. (b) Glenn, A. G.; Jones, P. B. Tetrahedron Lett. 2004,
45, 6967. (c) Ohtani, H.; Ishimura, S.; Kumai, M. Anal. Sci. 2008,
24, 1335.
(299) D’Anna, F.; Noto, R. Tetrahedron 2007, 63, 11681.
(300) D’Anna, F.; Vitale, P.; Noto, R. J. Org. Chem. 2009, 74, 6224.
(301) (a) Ingold, C. K. Structure and Mechanism in Organic Chem-
istry, 2nd ed.; Bell: London, 1969. (b) Hughes, E. D.; Ingold, C. K.
J. Chem. Soc. 1935, 244. (c) Hughes, E. D. Trans. Faraday Soc. 1941,
37, 603. (d) Hughes, E. D.; Ingold, C. K. Trans. Faraday Soc. 1941,
37, 657. (e) Cooper, K. A.; Dhar, M. L.; Hughes, E. D.; Ingold, C. K.;
MacNulty, B. J.; Woolf, L. I. J. Chem. Soc. 1948, 2043.
(302) Wheeler, C.; West, K. N.; Liotta, C. L.; Eckert, C. A. Chem.
Commun. 2001, 887.
(303) (a) Lancaster, N. L.;Welton, T.; Young,G. B. J. Chem. Soc., Perkin
Trans. 2 2001, 2267. (b) Lancaster, N. L.; Salter, P. A.; Welton, T.; Young,
G. B. J. Org. Chem. 2002, 67, 8855. (c) Lancaster, N. L.; Welton, T. J. Org.
Chem. 2004, 69, 5986. (d) Crowhurst, L.; Falcone, R.; Lancaster, N. L.;
Llopis-Mestre, V.; Welton, T. J. Org. Chem. 2006, 71, 8847.
(304) (a) Landini, D.; Maia, A. Tetrahedron Lett. 2005, 46, 3961. (b)
Betti, C.; Landini, D.; Maia, A. Tetrahedron 2008, 64, 1689.
(305) Arantes, G. M.; Ribeiro, M. C. C. J. Chem. Phys. 2008,
128, 114503.
(306) Chiappe, C.; Pieraccini, D.; Saullo, P. J. Org. Chem. 2003,
68, 6710.
(307) McNulty, J.; Nair, J. J.; Cheekoori, S.; Larichev, V.; Caretta, A.;
Robertson, A. J. Chem.—Eur. J. 2006, 12, 9314.
(308) (a) Kim, D. W.; Song, C. E.; Chi, D. Y. J. Am. Chem. Soc. 2002,
124, 10278. (b) Kim, D. W.; Song, C. E.; Chi, D. Y. J. Org. Chem. 2003,
68, 4281. (c) Kim, D. W.; Choe, Y. S.; Chi, D. Y. Nucl. Med. Biol. 2003,
30, 345.
(309) (a) Kim, D.W.; Chi, D. Y.Angew. Chem., Int. Ed. 2004, 43, 483.
(b) Kim, D. W.; Hong, D. J.; Jang, K. S.; Chi, D. Y. Adv. Synth. Catal.
2006, 348, 1719.
(310) (a) Shinde, S. S.; Lee, B. S.; Chi, D. Y. Org. Lett. 2008, 10, 733.
(b) Shinde, S. S.; Chi, H.M.; Lee, B. S.; Chi, D. Y.Tetrahedron Lett. 2009,
50, 6654.
(311) Murray, C. B.; Sanford, G.; Korn, S. R. J. Fluorine Chem. 2003,
123, 81.
(312) Stegmann, V.;Massonne, K. .World PatentWO2005 026089,
2005.
(313) Ren, R. X.; Wu, J. X. Org. Lett. 2001, 3, 3727.
(314) Nguyen, H.-P.; Kirilov, P.; Matondo, H.; Baboulene,M. J. Mol.
Catal. A 2004, 218, 41.
(315) Gupta, N.; Kad, G. L.; Singh, J. J. Mol. Catal. A 2009, 302, 11.
(316) (a) Royer, R.; Buisson, J.-P.; Demerseman, P.; Lechartier, J.-P.
Bull. Chim. Soc. Fr. 1969, 2792. (b) Royer, R.; Buisson, J.-P.; Demerse-
man, P. Bull. Chim. Soc. Fr. 1971, 4362. (c) Driver, G.; Johnson, K. E.
Green Chem. 2003, 5, 163. (d) Boonanahalli, S. K.; Kim, D.W.; Chi, D. Y.
J. Org. Chem. 2004, 69, 3340. (e) Schmid, C. R.; Beck, C. A.; Cronin, J. S.;
Staszak, M. A. Org. Process Res. Dev. 2004, 8, 670.
(317) Skrzypczak, A.; Neta, P. Int. J. Chem. Kinet. 2004, 36, 253.
(318) Bini, R.; Chiappe, C.; Pomelli, C. S.; Parisi, B. J. Org. Chem.
2009, 74, 8522.
(319) Jorapur, Y. R.; Lee, C.-H.; Chi, D. Y. Org. Lett. 2005, 7,
1231.
(320) Yau, H. M.; Howe, A. G.; Hook, J. M.; Croft, A. K.; Harper,
J. B. Org. Biomol. Chem. 2009, 7, 3572.
(321) Adam, C.; García-Rio, L.; Godoy, A.; Leis, J. R. Green Chem.
2006, 8, 596.
(322) Song., G.; Cai, Y.; Peng, Y. J. Comb. Chem. 2005, 7, 561.
(323) Sliger, M. D.; P’Pool, S. J.; Traylor, R. K.; McNeill, J., III;
Young, S. H.; Hoffman, N. W.; Klingshirn, M. A.; Rogers, R. D.;
Shaughnessy J. Organomet. Chem. 2005, 690, 3540.
(324) Swiderski, K.; McLean, A.; Gordon, C. M.; Vaughan, D. H.
Chem. Commun. 2004, 590.
(325) Ranieri, G.; Hallett, J. P.; Welton, T. Ind. Eng. Chem. Res. 2008,
47, 638.
(326) Hallett, J. P.; Liotta, C. L.; Ranieri, G.;Welton, T. J. Org. Chem.
2009, 74, 1864.
(327) Illner, P.; Begel, S.; Kern, S.; Puchta., R.; van Eldik, R. Inorg.
Chem. 2009, 48, 588.
(328) Correia, I; Welton, T. Dalton Trans. 2009, 4115.
(329) Creary, X.; Willis, E.; Gagnon, M. J. Am. Chem. Soc. 2005,
127, 18114.
(330) Man, B. Y. W.; Hook, J. M.; Harper, J. B. Tetrahedron Lett.
2005, 46, 7641.
(331) Yau, H. M.; Barnes, S. A.; Hook, J. M.; Youngs, T. G. A.; Croft,
A. K.; Harper, J. B. Chem. Commun. 2008, 3576.
(332) Shim, Y.; Kim, H. J. J. Phys. Chem. B 2008, 112, 2637.
(333) Lynden-Bell, R. M. Phys. Chem. Chem. Phys. 2010, 12, 1733.
(334) Laali, K. K.; Gettwert, V. J. J. Fluorine Chem. 2001, 107, 31.
(335) Bini, R.; Chiappe, C.; Marmugi, E.; Pieraccini, D. Chem.
Commun. 2006, 897.
(336) Laali, K. K.; Okazaki, T.; Bunge, S. D. J. Org. Chem. 2007,
72, 6758.
(337) Yadav, J. S.; Reddy, B. V. S.; Basak, A. K.; Narsaiah, A. V.
Tetrahedron Lett. 2003, 44, 2217.
(338) D’Anna, F.; Frenna, V.; Noto, R.; Pace, V.; Spinelli, D. J. Org.
Chem. 2006, 71, 5144.
(339) Newington, I.; Perez-Arlandis, J. M.; Welton, T. Org. Lett.
2007, 9, 5247.
(340) D’Anna, F.; Marullo, S.; Noto, R. J. Org. Chem. 2008, 73, 6224.
(341) D’Anna, F.; Frenna, V.; Pace, V.; Noto, R. Tetrahedron 2006,
62, 1690.
(342) D’Anna, F.; La Marca, S.; Noto, R. J. Org. Chem. 2008,
73, 3397.

BL dx.doi.org/10.1021/cr1003248 |Chem. Rev. XXXX, XXX, 000–000
Chemical Reviews
REVIEW
(343) D’Anna, F.; La Marca, S.; Lo Meo, P.; Noto, R. Chem.—Eur. J.
2009, 15, 7896.
(344) Daguenet, C.; Dyson, P. J. Inorg. Chem. 2007, 46, 403.
(345) Diaz, P.; Mingos, D. M. P.; Vilar, R.; White, A. J. P.; Williams,
D. J. Inorg. Chem. 2004, 43, 7597.
(346) Spange, S.; Vilsmeier, E.; Adolph, S.; F€ahrmann, A. J. Phys.
Org. Chem. 1999, 12, 547.
(347) Chiappe, C.; Pieraccini, D. J. Phys. Chem. A 2006, 112, 4937.
(348) (a) Chiappe, C.; Capraro, D.; Conte, V.; Pieraccini, D. Org.
Lett. 2001, 3, 1061. (b) Bortolini, O.; Bottai, M.; Chiappe, C.; Conte, V.;
Pieraccini, D. Green Chem. 2002, 4, 621.
(349) Chiappe, C.; Conte, V.; Pieraccini, D. Eur. J. Org. Chem.
2002, 2831.
(350) Chiappe, C.; Pieraccini, D. J. Org. Chem. 2004, 69, 6059.
(351) Cristiano, R.; Ma, K.; Pottanat, G.; Weiss, G. J. Org. Chem.
2009, 74, 9027.
(352) Song, X.; Zhou, J.; Li, Y.; Tang, Y. J. Photochem. Photobiol., A
1995, 92, 99.
(353) Byrne, R.; Coleman, S.; Fraser, K. J.; Raduta, A.; MacFarlane,
D. R.; Diamond, D. Phys. Chem. Chem. Phys. 2009, 11, 7286.
(354) Calo, V.; Nacci, A.; Lopez, L.; Lerario, V. L. Tetrahedron Lett.
2000, 41, 8977.
(355) Verma, A. K.; Attri, P.; Chopra, V.; Tiwari, R. K.; Chandra, R.
Monatsh. Chem. 2008, 139, 1041.
(356) Meciarova, M.; Toma, S. Chem.—Eur. J. 2007, 13, 1268.
(357) Meciarova, M.; Cigan, M.; Toma, S.; Gaplovsky, A. Eur. J. Org.
Chem. 2008, 4408.
(358) Qian, Y.; Xiao, S.; Liu, L.; Wang, Y. Tetrahedron: Asymmetry
2008, 19, 1515.
(359) Rasalkar, M. S.; Potdar, M. K.; Mohile, S. S.; Salunkhe, M. M.
J. Mol. Catal. A 2005, 235, 267.
(360) Gallo, V.; Mastrorilli, P.; Nobile, C. F.; Romanazzi, G.;
Suranna, G. P. J. Chem. Soc., Dalton Trans. 2002, 4339.
(361) Gallo, V.; Giardina-Papa, D.; Mastrorilli, P.; Nobile, C. F.;
Suranna, G. P.; Wang., Y. J. Organomet. Chem. 2005, 690, 3535.
(362) Fabris, M.; Lucchini, V.; Noe, M.; Perosa, A.; Selva, M. Chem.
—Eur. J. 2009, 15, 12273.
(363) Yu, C.-J.; Liu, C.-J. Molecules 2009, 14, 3222.
(364) Jaeger, D. A.; Tucker, C. E. Tetrahedron Lett. 1989, 30, 1785.
(365) Berson, J. A.; Hamlet, Z.; Mueller, W. J. Am. Chem. Soc. 1962,
84, 297.
(366) Janus, E.; Goc-Maciejewska, I.; yo_zynski, M.; Pernak, J.
Tetrahedron Lett. 2006, 47, 4079.
(367) Tiwari, S.; Kumar, A. Angew. Chem., Int. Ed. 2006, 45, 4824.
(368) Tiwari, S.; Khupse, N.; Kumar, A. J. Org. Chem. 2008, 73
9075.
(369) Vidis, A.; Ohlin, C. A.; Laurenczy, G.; K€uster, E.; Sedelmeier,
G.; Dyson, P. J. Adv. Synth. Catal. 2005, 347, 266.
(370) Vidis, A.; Laurenczy, G.; K€usters, E.; Sedelmeier, G.; Dyson,
P. J. J. Phys. Org. Chem. 2007, 20, 109.
(371) Bini, R.; Chiappe, C.; Llopis Mestre, V.; Pomelli, C. S.;
Welton, T. Org. Biomol. Chem. 2008, 6, 2522.
(372) Bini, R.; Chiappe, C.; Llopis Mestre, V.; Pomelli, C. S.;
Welton, T. Theor. Chem. Acc. 2009, 123, 347.
(373) Yanai, H.; Ogura, H.; Taguchi, T. Org. Biomol. Chem. 2009,
7, 3657.
(374) Pegot, B.; Van Buu, O. N.; Gori, D.; Vo-Thanh, G. Beilstein J.
Org. Chem. 2006, 2, 18.
(375) Park, J. K.; Sreekanth, P.; Kim, B. M. Adv. Synth. Catal. 2004,
346, 49.
(376) P’Pool, S. J.; Klingshirn, M. A.; Rogers, R. D.; Shaughnessy,
K. H. J. Organomet. Chem. 2005, 690, 3522.
(377) (a) Ramnial, T.; Ino, D. D.; Clyburne, J. A. C. Chem. Commun.
2005, 325. (b) Ramnial, T.; Taylor, S. A.; Bender, M. L.; Gorodetsky, B.;
Lee, P. T. K.; Dickie, D. A.; McCollum, B. M.; Pye, C. C.; Walsby, C. J.;
Clyburne, J. A. C. J. Org. Chem. 2008, 73, 801.
(378) Law, M. C.; Wong, K.-Y.; Chan, T. H. Chem. Commun.
2006, 2457.
(379) Gordon, C. M.; McClusky, A. Chem. Commun. 1999
1431.
(380) Kabalka, G. W.; Venkataiah, B.; Das, B. C. Green Chem. 2002,
4, 472.
(381) Gordon, C. M.; Ritchie, C. Green Chem. 2002, 4, 124.
(382) Law, M. C.; Wong, K.-Y.; Chan, T. H. Green Chem. 2002,
4, 161.
(383) Kitazume, T.; Kasai, K. Green Chem. 2001, 3, 30.
(384) Law, M. C.; Wong, K.-Y.; Chan, T. H.Green Chem. 2004, 6, 241.
(385) Boon, J. A.; Levisky, J. A.; Pflug, J. L.; Wilkes, J. S. J. Org. Chem.
1986, 51, 480.
(386) Chen, M.; Luo, Y.; Li, G.; He, M.; Xie, J.; Li, H.; Yuan, X.
Korean J. Chem. Eng. 2009, 26, 1563.
(387) Qiao, C.-Z.; Zhang, Y.-F.; Zhang, J.-C.; Li, C.-Y.Appl. Catal., A
2004, 276, 61.
(388) Qiao, K.; Deng, Y. J. Mol. Catal. A 2001, 171, 81.
(389) (a) Zhao, Z.; Qiao, W.; Li, Z.; Wang, G.; Cheng, L. J. Mol.
Catal. A 2004, 222, 207. (b) Zhao, Z.; Li, Z.; Wang, G.; Qiao,W.; Cheng,
L. Appl. Catal., A 2004, 262, 69.
(390) Joni, J.; Schmitt, D.; Schulz, P. S.; Lotz, T. J.; Wasserscheid, P.
J. Catal. 2008, 258, 401.
(391) Klamt, A.; Eckert, F. Fluid Phase Equilib. 2000, 172, 43.
(392) Piao, L. Y.; Fu, Y. L.; Tao, G. H.; Kou, Y. Catal. Today 2004,
93-95, 301.
(393) (a) Sun, X.; Zhao, S.; Wang, R. Chin. J. Chem. Eng. 2004,
12, 658. (b) Sun, X.; Zhao, S. Chin. J. Chem. Eng. 2006, 14, 289.
(394) Zhao, Z.; Qiao, W.; Wang, X.; Wang, G.; Li, Z.; Cheng, L.
Appl. Catal., A 2005, 290, 133.
(395) Jia, L.-J.; Wang, Y.-Y.; Chen, H.; Shan, Y. K.; Dai, L.-Y. React.
Kinet. Catal. Lett. 2005, 86, 267.
(396) Wasserscheid, P.; Sesing, M.; Korth, W. Green Chem. 2002,
4, 134.
(397) Shen, H.-Y; Judeh, Z. M. A.; Ching, C. B.; Xia, Q.-H. J. Mol.
Catal. A 2004, 212, 301.
(398) Liu, X.; Liu, M.; Guo, X.; Zhou, J. Catal. Commun. 2008, 9, 1.
(399) Song, C. E.; Shim, W. H.; Roh, E. J.; Choi, J. H. Chem.
Commun. 2000, 1695.
(400) Sarca, V. D.; Laali, K. K. Green Chem. 2006, 8, 615.
(401) Laali, K. K.; Sarca, V. D.; Okazaki, T.; Brock, A; Der, P. Org.
Biomol. Chem. 2005, 3, 1034.
(402) Li, C.; Liu, W.; Zhao, Z. Catal. Commun. 2007, 8, 1834.
(403) (a) Csihony, S.; Mehdi, H.; Horvath, I. T. Green Chem. 2001,
3, 307. (b) Csihony, S.;Mehdi, H.; Homonnay, Z.; Vertes, A.; Farkas, O.;
Horvath, I. T. J. Chem. Soc., Dalton Trans. 2002, 680.
(404) Earle, M. J.; Hakala, U.; Hardacre, C.; Karkkainen, J.; McAuley,
B. J.; Rooney, D. W.; Seddon, K. R.; Thompson, J. M.; W€ah€al€a, K. Chem.
Commun. 2005, 903.
(405) Hardacre, C.; Nancarrow, P.; Rooney, D. W.; Thompson,
J. M. Org. Process Res. Dev. 2008, 12, 1156.
(406) Ross, J.; Xiao, J. Green Chem. 2002, 4, 129.
(407) Gmouh, S.; Yang, H.; Vaultier, M. Org. Lett. 2003, 5, 2219.
(408) Goodrich, P.; Hardacre, C.; Mehdi, H.; Nancarrow, P.;
Rooney, D. W.; Thompson, J. M. Ind. Eng. Chem. Res. 2006, 45, 6640.
(409) Zayed, F.; Greiner, L.; Schulz, P. S.; Lapkin, A.; Leitner, W.
Chem. Commun. 2008, 79.
(410) Hardacre, C.; Katdare, S. P.; Milroy, D.; Nancarrow, P.;
Rooney, D. W.; Thompson, J. M. J. Catal. 2004, 227, 44.
(411) Malhotra, S. V.; Xiao, Y. Aust. J. Chem. 2006, 59, 468.
(412) Boon, J. A.; Lander, S. W., Jr.; Levisky, J. A.; Pfug, J. L.;
Skryznecki-Cooke, L. M.; Wilkes, J. S. Proc. Int. Symp. Molten Salts 1987,
6, 979.
(413) Laali, K. K.; Gettwert, V. J. J. Org. Chem. 2001, 66, 35.
(414) Rajagopal, R.; Srinivasan, K. V. Synth. Commun. 2003, 33, 961.
(415) Lancaster, N. L.; Llopis-Mestre, V. Chem. Commun.
2003, 2812.
(416) Dal, E.; Lancaster, N. L. Org. Biol. Chem. 2005, 3, 2812.
(417) Smith, K.; Liu, S.; El-Hiti, G. A. Ind. Eng. Chem. Res. 2005,
44, 8611.

BM dx.doi.org/10.1021/cr1003248 |Chem. Rev. XXXX, XXX, 000–000
Chemical Reviews
REVIEW
(418) Earle, M. J.; Katdare, S. P.; Seddon, K. R. Org. Lett. 2004,
6, 707.
(419) Qiao, K.; Yokoyama, C. Chem. Lett. 2004, 33, 808.
(420) Fang, D.; Shi, Q.-R.; Cheng, J.; Gong, K.; Liu, Z.-L. Appl.
Catal., A 2008, 345, 158.
(421) Corma, A.; Martinez, A. Catal. Rev. 1993, 35, 483.
(422) Ya, K. S.; Namboodiri, V. V.; Varma, R. S.; Smirniotis, P. G.
J. Catal. 2004, 222, 511.
(423) Huang, C.-P.; Liu, Z.-C.; Xu, C.-M.; Chen, B.-H.; Liu, Y.-F.
Appl. Catal., A 2004, 277, 41.
(424) Zhang, J.; Huang, C.; Chen, B.; Ren, P.; Pu, M. J. Catal. 2007,
249, 261.
(425) Liu, Y.; Hu, R.; Xu, C.; Su, H. Appl. Catal., A 2008, 346, 189.
(426) Liu, Z.; Xu, C.; Huang, C. U.S. Patent 0133056, 2004.
(427) Liu, Z.; Zhang, R.; Xu, C.; Xia, R. Oil Gas J. 2006, 104, 52.
(428) Olah, G. A. U.S. Patent 5073674, 1991.
(429) Tang, S.; Scurto, A. M.; Subramaniam, B. J. Catal. 2009,
268, 243.
(430) Deng, Y.; Shi, F.; Beng, J.; Qiao, K. J. Mol. Catal. A 2001,
165, 33.
(431) Forbes, D. C.; Weaver, K. J. J. Mol. Catal. A 2004, 214, 129.
(432) Zhu, H.-P.; Yang, F.; Tang, J.; He, M.-Y. Green Chem. 2003,
5, 38.
(433) Zhang, H.; Xu, F.; Zhou, X.; Zhang, G.; Wang, C. Green Chem.
2007, 9, 1208.
(434) Fraga-Dubreuil, J.; Bourahla, K.; Rahmouni, M; Bazureau,
J. P.; Hamelin, J. Catal. Commun. 2002, 3, 185.
(435) Ganeshpure, P. A.; George, G.; Das, J. J. Mol. Catal. A 2008,
279, 182.
(436) Arfan, A.; Bazureau, J. P. Org. Process Res. Dev. 2005, 9
743.
(437) Xing, H.; Wang, T.; Zhou, Z.; Dai, Y. Ind. Eng. Chem. Res.
2005, 44, 4147.
(438) Fang, D.; Zhou, X.-L.; Ye, Z.-W.; Liu, Z.-L. Ind. Eng. Chem. Res.
2006, 45, 7982.
(439) Liu, S.; Xie, C.; Yu, S.; Liu, F.; Ji, K. Catal. Commun. 2008,
9, 1634.
(440) Jiang, D.; Wang, Y. Y.; Tu, M.; Dai, L. Y. React. Kinet. Catal.
Lett. 2008, 95, 265.
(441) Zhao, Y.; Long, J.; Deng, F.; Liu, X.; Li, Z.; Xia, C.; Peng, J.
Catal. Commun. 2009, 10, 732.
(442) Wei, Z.; Li, F.; Xing, H.; Deng, S.; Ren, Q.Korean J. Chem. Eng.
2009, 26, 666.
(443) Jiang, T.; Chang, Y.; Zhao, G.; Han, B.; Yang, G. Synth.
Commun. 2004, 34, 225.
(444) Wells, T. P.; Hallett, J. P.; Williams, C. K.; Welton, T. J. Org.
Chem. 2008, 73, 5585.
(445) Sato, A.; Nakamura, Y.; Maki, T.; Ishihara, K.; Yamamoto, H.
Adv. Synth. Catal. 2005, 347, 1337.
(446) Park, S.; Viklund, F.; Hult, K.; Kazlauskas, R. J. Green Chem.
2003, 5, 715.
(447) Ulbert, O.; Frater, T.; Belafi-Bako, K.; Gubicza, L. J. Mol. Catal.
B 2004, 31, 39.
(448) Barahona, D.; Pfromm, P. H.; Rezac, M. C. Biotechnol. Bioeng.
2006, 93, 318.
(449) Lee, S. H.; Dang, D. T.; Ha, S. H.; Chang, W.-J.; Koo, Y.-M.
Biotechnol. Bioeng. 2008, 99, 1.
(450) Bo, W.; Ming, Y. L.; Shaun, S. J. Tetrahedron Lett. 2003,
44, 5037.
(451) Qureshi, Z. S.; Deshmukh, K. M.; Bhor, M. D.; Bhanage, B. M.
Catal. Commun. 2009, 10, 833.
(452) Domingos, J. B.; Dupont, J. Catal. Commun. 2007, 8, 1383.
(453) Wu, Q.; Chen, H.; Han, M.; Wang, D.; Wang., J. Ind. Eng.
Chem. Res. 2007, 46, 7955.
(454) Nomura, N.; Taira, A.; Nakase, A.; Tomioka, T.; Okada, M.
Tetrahedron 2007, 63, 8478.
(455) Lee, C. W. Tetrahedron Lett. 1999, 40, 2461.
(456) Kumar, A.; Pawar, S. S. J. Org. Chem. 2004, 69, 1419.
(457) Abbott, A. P.; Capper, G.; Davies, D. L.; Rasheed, R. K.;
Tambyrajah, V. Green Chem. 2002, 4, 24.
(458) Yin, D.; Li, C.; Li, B.; Tao, L.; Yin, D. Adv. Synth. Catal. 2005,
347, 137.
(459) Yeom, C.-E.; Kim, H. W.; Shin, Y. J.; Kim, B. M. Tetrahedron
Lett. 2007, 48, 9035.
(460) Doherty, S.; Goodrich, P.; Hardacre, C.; Knight, J. G.;
Nguyen, M. T.; Pa^rvulescu, V. I.; Paun, C. Adv. Synth. Catal. 2007,
349, 951.
(461) Doherty, S.; Goodrich, P.; Hardacre, C.; Luo, H.-K.; Rooney,
D. W.; Seddon, K. R.; Styring, P. Green Chem. 2004, 6, 63.
(462) Vidis, A.; K€usters, E.; Sedelmeier, G.; Dyson, P. J. J. Phys. Org.
Chem. 2008, 21, 264.
(463) (a) Silvero, G.; Arevalo,M. J.; Bravo, J. L.; Avalos,M.; Jimenez,
J. L.; Lopez, I. Tetrahedron 2005, 61, 7105. (b) Lopez, I.; Silvero, G.;
Arevalo, M. J.; Babiano, R.; Palacios, J. C.; Bravo, J. L. Tetrahedron 2007,
63, 2901.
(464) Kumar, A.; Pawar, S. S. J. Org. Chem. 2007, 72, 8111.
(465) Fayet, C.; Gelas, J. Carbohydr. Res. 1983, 122, 59.
(466) Duclos, A.; Fayet, C.; Gelas, J. Synthesis 1994, 1087.
(467) Lansalot-Matras, C.; Moreau, C.Catal. Commun. 2003, 4, 517.
(468) Bao, Q.; Qiao, K.; Tomida, D.; Yokoyama, C. Catal. Commun.
2008, 9, 1383.
(469) Moreau, C.; Finiels, A.; Vanoje, L. J. Mol. Catal. A 2006,
253, 165.
(470) Zhao, H.; Holladay, J. E.; Brown, H.; Zhang, Z. C. Science
2007, 316, 1597.
(471) Kumar, R.; Sharma, A.; Sharma, N.; Kumar, V.; Sinha, A. K.
Eur. J. Org. Chem. 2008, 5577.
(472) Hu, S.; Zhang, Z.; Song, J.; Zhou, Y.; Han, B. Green Chem.
2009, 11, 1743.
(473) Satoh, H.; Kaga, H.; Kakuchi, T.; Satoh, T.; Takahashi, K.
Chem. Lett. 2009, 38, 1178.
(474) Fabos, V.; Lantos, D.; Bodor, A.; Balint, A.-M.; Mika, L. T.;
Sielcken, O. E.; Cuiper, A.; Horvath, I. T. ChemSusChem 2008, 1, 189.
(475) (a) Peng, J.; Deng, Y.Tetrahedron Lett. 2001, 42, 403. (b) Ren,
R. X.; Zueva, L.; Ou, W. Tetrahedron Lett. 2001, 42, 8441.
(476) Du, Z.; Li, Z.; Gu, Y.; Zhang, J.; Deng, Y. J. Mol. Catal. A 2005,
237, 80.
(477) Guo, S.; Du, Z.; Zhang, S.; Li, D.; Li, Z.; Deng, Y. Green Chem.
2006, 8, 296.
(478) Zicmanis, A.; Katevica, S.; Mekss, P. Catal. Commun. 2009,
10, 614.
(479) Pinto, A. C.; Moreira Lapis, A. A.; da Silva, B. V.; Bastos, R. S.;
Dupont, J.; Neto, B. A. D. Tetrahedron Lett. 2008, 49, 5639.
(480) Zhao, G.; Jiang, T.; Gao, H.; Han, B.; Huang, J.; Sun, D. Green
Chem. 2004, 6, 75.
(481) (a) Poletti, L.; Rencurosi, A.; Lay, L.; Russo, G. Synlett 2003,
15, 2297. (b) Rencurosi, A.; Lay, L.; Russo, G.; Caneva, E.; Poletti, L.
J. Org. Chem. 2005, 70, 7765. (c) Rencurosi, A.; Lay, L.; Russo, G.;
Caneva, E.; Poletti, L. Carbohydr. Res. 2006, 341, 903.
(482) Sowmiah, S.; Srinivasadesikan, V.; Tseng, M.-C.; Chu, Y.-H.
Molecules 2009, 14, 3780.
(483) Aggarwal, V. K.; Emme, I.; Mereu, A. Chem. Commun.
2002, 1612.
(484) Kotrusz, P.; Kmentova, I.; Gotov, B.; Toma, S.; Solcaniova, E.
Chem. Commun. 2002, 2510.
(485) Formentín, P.; García, H.; Leyva, A. J. Mol. Catal. A 2004,
214, 137.
(486) Mehnert, C. P.; Dispenziere, N. C.; Cook, R. A. Chem.
Commun. 2002, 1610.
(487) Cota, I.; Goknzalez-Olmos, R.; Iglesias, M.; Medina, F. J. Phys.
Chem. B 2007, 111, 12468.
(488) Cordova, A. Tetrahedron Lett. 2004, 45, 3949.
(489) (a) Gruttadauria, M.; Riela, S.; Lo Meo, P.; D’Anna, F.; Noto,
R. Tetrahedron Lett. 2004, 45, 6113. (b) Gruttadauria, M.; Riela, S.;
Aprile, C.; Lo Meo, P.; D’Anna, F.; Noto, R. Adv. Synth. Catal. 2006,
348, 82.

BN dx.doi.org/10.1021/cr1003248 |Chem. Rev. XXXX, XXX, 000–000
Chemical Reviews
REVIEW
(490) Zhou, W.; Xu, L.-W.; Qiu, H.-Y.; Lai, G.-Q.; Xia, C.-G.; Jiang,
J.-X. Helv. Chim. Acta 2008, 91, 53.
(491) (a) Kotrusz, P.; Alemayehu, S.; Toma, S.; Schmalz, H.-G.;
Adler, A. Eur. J. Org. Chem. 2005, 4904. (b) Guo, H.-M.; Cun, L.-F.;
Gong, L.-Z.; Mi, A.-Q.; Jiang, Y.-Z. Chem. Commun. 2005, 1450.
(492) Davey, P. N.; Forsyth, S. A.; Gunaratne, N.; Hardacre, C.;
McKeown, A.; McMath, S. E. J.; Rooney, D. W.; Seddon, K. R. Green
Chem. 2005, 7, 224.
(493) Lombardo, M.; Easwar, S.; Pasi, F.; Trombini, C.; Dhavale,
D. D. Tetrahedron 2008, 64, 9203.
(494) (a) Miao, W.; Chan, T. H. Adv. Synth. Catal. 2006, 348, 1711.
(b) Luo, S.; Mi, X.; Zhang, L.; Liu, S.; Xu, H.; Cheng, J.-P. Tetrahedron
2007, 63, 1923. (c) Yang, S.-D.; Wu, L.-Y.; Yan, Z.-Y.; Pan, Z.-L.; Liang,
Y.-M. J. Mol. Catal. A 2007, 268, 107. (d) Zhou, L.; Wang, L. Chem. Lett.
2007, 36, 628.
(495) Lombardo, M.; Pasi, F.; Easwar, S.; Trombini, C. Adv. Synth.
Catal. 2007, 349, 2061.
(496) Hu, S.; Jiang, T.; Zhang, Z.; Zhu, A.; Han, B.; Song, J.; Xie, Y.;
Li, W. Tetrahedron Lett. 2007, 48, 5613.
(497) (a) Zhu, A.; Jiang, T.; Wang, D.; Han, B.; Liu, L.; Huang, J.;
Zhang, J.; Sun, D. Green Chem. 2005, 7, 514. (b) Zhu, A.; Jiang, T.; Han,
B.; Huang, J.; Zhang, J.; Ma, X. New J. Chem. 2006, 30, 736.
(498) Jiang, T.; Gao, H.; Han, B.; Zhao, G.; Chang, Y.; Wu, W.; Gao,
L.; Yang, G. Tetrahedron Lett. 2004, 45, 2699.
(499) Rosa, J. N.; Alfonso, C. A. M.; Santos, A. G. Tetrahedron 2001,
57, 4189.
(500) (a) Kim, E. J.; Ko, S. Y. Helv. Chim. Acta 2003, 86, 894. (b)
Kumar, A.; Pawar, S. S. J. Mol. Catal. A 2004, 21, 43.
(501) Hsu, J.-C.; Yen, Y.-H.; Chu, Y.-H. Tetrahedron Lett. 2004,
45, 4673.
(502) Lin, Y.-S.; Lin, C.-Y.; Liu, C.-W.; Tsai, T. Y. R. Tetrahedron
2006, 62, 872.
(503) Johnson, C. L.; Donkor, R. E.; Nawaz, W.; Karodia, N.
Tetrahedron Lett. 2004, 45, 7359.
(504) (a) Gong, H.; Cai, C.-q.; Yang, N.-f.; Yang, L.-w.; Zhang, J.;
Fan, Q.-h. J. Mol. Catal. A 2006, 249, 236. (b) Zhao, S.-H.; Zhang, H.-R.;
Feng, L.-H.; Chen, Z.-B. J. Mol. Catal. A 2006, 258, 251. (c) Zhao, S.;
Zhao, E.; Shen, P.; Zhao, M.; Sun, J. Ultrason. Sonochem. 2008
15, 955.
(505) Mi, X.; Luo, S.; Cheng, J.-P. J. Org. Chem. 2005, 70, 2338.
(506) Mi, X.; Luo, S.; Xu, H.; Zhang, L.; Cheng, J.-P. Tetrahedron
2006, 62, 2537.
(507) Lenard~ao, E. J.; de Oliveira Feijo, J.; Thurow, S.; Perin, G.;
Jacob, R. G. Tetrahedron Lett. 2009, 50, 5215.
(508) (a) Pegot, B.; Vo-Thanh, G.; Gori, D.; Loupy, A. Tetrahedron
Lett. 2004, 45, 6425. (b) Garre, S.; Parker, E.; Ni, B.; Headley, A. D. Org.
Biomol. Chem. 2008, 6, 3041.
(509) Morrison, D. W.; Forbes, D. C.; Davis, J. H., Jr. Tetrahedron
Lett. 2001, 42, 6053.
(510) Forbes, D. C,; Law, A. M.; Morrison, D. W. Tetrahedron Lett.
2006, 47, 1699.
(511) Xu, X. M.; Li, Y. Q.; Zhou, M. Y. Chin. J. Org. Chem. 2004,
24, 1253.
(512) Cai, Y.; Peng, Y.; Song, G. Catal. Lett. 2006, 109, 61.
(513) Paun, C.; Barklie, J.; Goodrich, P.; Gunaratne, H. Q. N.;
McKeown, A.; Pa^rvulascu, V. I.; Hardacre, C. J. Mol. Catal. A 2007,
269, 64.
(514) Fan, X.; Hu, X.; Zhang, X.; Wang, J. Aust. J. Chem. 2004,
57, 1067.
(515) Khan, F. A.; Dash, J.; Satapathy, R.; Upadhyay, S. K. Tetra-
hedron Lett. 2004, 45, 3055.
(516) Wang, Y.; Shang, Z.-c.; Wu, T.-x.; Fan, J.-c.; Chen, X. J. Mol.
Catal. A 2006, 253, 212.
(517) Hu, Y.; Chen, J.; Le, Z.-G.; Zheng, Q.-G. Synth. Commun.
2005, 35, 739.
(518) Chen, Z.-C.; Zheng, Q.-G. Synthesis 2003, 555.
(519) Li, Y. Q.; Xu, X. M.; Zhou, M. Y. Chin. Chem. Lett. 2003,
14, 448.
(520) Yue, C.; Mao, A.; Wei, Y.; L€u, M. Catal. Commun. 2008, 9, 1571.
(521) Wang, W.-J.; Cheng, W.-P.; Shao, L.-l.; Liu, C. H.; Yang, J.-G.
Kinet. Catal. 2009, 50, 186.
(522) Xin, X.; Guo, X.; Duan, H.; Lin, Y.; Sun, H. Catal. Commun.
2007, 8, 115.
(523) (a) Hangarge, R. V.; Jarikote, D. V.; Shingare, M. S. Green Chem.
2002, 4, 266. (b) Hu, Y.; Wei, P.; Huang, H.; Le, Z.-G.; Chen, Z.-C. Synth.
Commun. 2005, 35, 2955.
(524) Darvatkar, N. B.; Deorukhkar, A. R.; Bhilare, S. V.; Salunkhe,
M. M. Synth. Commun. 2006, 36, 3043.
(525) Adam, C. G.; Fortunato, G. G.; Mancini, P. M. J. Phys. Org.
Chem. 2009, 22, 460.
(526) Erkkil€a, A.; Majander, I.; Pihko, P. M. Chem. Rev. 2007,
107, 5416.
(527) Angelini, G.; DeMaria, P.; Chiappe, C.; Fontana, A.; Gasbarri,
C.; Siani, G. J. Org. Chem. 2009, 74, 6572.
(528) Cardillo, B.; Casnati, G.; Pochini, A.; Ricca, A. Tetrahedron
1967, 23, 3771.
(529) Earle, M. J.; McCormac, P. B.; Seddon, K. R. Chem. Commun.
1998, 2245.
(530) Le, Z.-G.; Chen, Z.-C.; Hu, Y.; Zheng, Q.-G. Synthesis
2004, 208.
(531) Jorapur, Y. R.; Jeong, J. M.; Chi, D. Y. Tetrahedron Lett. 2006,
47, 2435.
(532) Vavilina, G.; Zicmanis, A.; Mekss, P.; KlavinsChem. Heterocycl.
Compd. (N. Y.) 2008, 44, 530.
(533) Vavilina, G.; Zicmanis, A.; Mekss, P.; Klavins Chem. Hetero-
cyclyic Compd. (N. Y.) 2008, 44, 549.
(534) Chiappe, C.; Pieraccini, D. Green Chem. 2003, 5, 193.
(535) D’Anna, F.; Frenna, V.; Noto, R.; Pace, V.; Spinelli, D. J. Org.
Chem. 2006, 71, 9637.
(536) Lapis, A. A. M.; de Oliveira, L. F.; Neto, B. A. D.; Dupont, J.
ChemSusChem 2008, 1, 759.
(537) Ju, H.-Y.; Manju, M. D.; Park, D.-W.; Choe, Y.; Park, S.-W.
React. Kinet. Catal. Lett. 2007, 90, 3.
(538) Forsyth, S. A.; MacFarlane, D. R.; Thompson, R. J.; von
Ltzstein, M. Chem. Commun. 2002, 714.
(539) Gholap, A. R.; Venkatesan, K.; Daniel, T.; Lahoti, R. J.;
Srinivasan, K. V. Green Chem. 2003, 5, 693.
(540) Duan, Z.; Gu, Y.; Deng, Y. J. Mol. Catal. A 2006, 246, 70.
(541) Lui, Y.; Liu, L.; Lu, Y.; Cai, Y.-Q. Monatsh. Chem. 2008,
139, 633.
(542) Swatloski, R. P.; Spear, S. K.; Holbrey, J. D.; Rogers, R. D.
J. Am. Chem. Soc. 2002, 124, 4974.
(543) (a) Wu, J.; Zhang, J.; Zhang, H.; He, J.; Ren, Q.; Guo, M.
Biomacromolecules 2004, 5, 266. (b) Barthel, S.; Heinze, T. Green Chem.
2006, 8, 301. (c) Liu, C. F.; Sun, R. C.; Zhang, A. P.; Ren, J. L.; Geng,
Z. C. Polym. Degrad. Stab. 2006, 91, 3040. (d) Liu, C. F.; Sun, R. C.;
Zhang, A. P.; Ren, J. L.; Wang, X. A.; Qin, M. H.; Chao, Z. N.; Luo, W.
Carbohydr. Res. 2007, 342, 919. (e) Liu, C. F.; Sun, R. C.; Zhang, A. P.;
Qin, M. H.; Ren, J. L.; Wang, X. A. J. Agric. Food Chem. 2007, 55, 2399.
(544) Haibo, X.; King, A.; Kilpelainen; Granstrom, M.; Argyropoulos,
D. S. Biomacromolecules 2007, 8, 3740.
(545) (a) Nyce, G. W.; Glauser, T.; Connor, E. F.; M€ock, A.;
Waymouth, R. M.; Hedrick, J. L. J. Am. Chem. Soc. 2003, 125, 3046.
(b) Dove, A. P.; Pratt, R. C.; Lohmeijer, B. G. G.; Culkin, D. A.; Hagberg,
E. C.; Nyce, G. W.; Waymouth, R. M.; Hedrick, J. L. Polymer 2006,
47, 4018.
(546) (a) Volland, M.; Seitz, V.; Maase, M.; Flores, M.; Papp, R.;
Massonne, K.; Stegmann, V.; Halbritter, K.; Noe, R.; Bartsch, M.; Siegel,
W.; Becker, M.; Huttenloch, O., BASF SE.World PatentWO 03062251,
2003. (b) Maase, M.; Huttenloch, O., BASF SE. World Patent WO
061416, 2005.
(547) Chojnowski, J.; Cypryk, M.; Fortuniak, W. Heteroat. Chem.
1991, 2, 63.
(548) http://www.basf.com/group/corporate/en/function/conversions:
/publish/content/sustainability/eco-efficiency-analysis/images/BASF_Eco-
Efficiency_Label_Basil_2005.pdf.

BO dx.doi.org/10.1021/cr1003248 |Chem. Rev. XXXX, XXX, 000–000
Chemical Reviews
REVIEW
(549) Cornils, B., Herrmann, W. A., Horvath, I. T., Leitner, W.,
Meckingm, S., Olivier-Bourbigou, Vogt, D., Eds. Multiphase Homoge-
neous Catalysis; Wiley-VCH: Weinheim, Germany, 2005.
(550) Chauvin, Y.; Mussmann, L.; Olivier, H. Angew. Chem., Int. Ed.
Engl. 1995, 34, 2698.
(551) Wasserscheid, P.; Waffenschmidt, H.; Machnitzki, P.; Kottsieper,
K. W.; Stelzer, O. Chem. Commun. 2001, 451.
(552) (a) Kottsieper, K. W.; Stelzer, O.; Wasserscheid, P. J. Mol.
Catal. A 2001, 175, 285. (b) Brauer, D. J.; Kottsieper, K. W.; Liek, C.;
Stelzer, O.; Waffenschmidt, H.; Wasserscheid, P. J. Organomet. Chem.
2001, 630, 177. (c) Sirieix, J.; Ossberger, M.; Betzemeier, B.; Knochel, P.
Synlett 2000, 11, 1613.
(553) Favre, F.; Olivier-Bourbigou, H.; Commereuc, D.; Saussine
Chem. Commun. 2001, 1360.
(554) Brasse, C. C.; Englert, U.; Salzer, A.; Waffenschmidt, H.;
Wasserscheid, P. Organometallics 2000, 19, 3818.
(555) Riisager, A.; Wasserscheid, P.; van Hal, R.; Fehrmann, R.
J. Catal. 2003, 219, 452.
(556) Dyson, P. J.; Laurenczy, G.; Ohlin, C. A.; Vallance, J.; Welton,
T. Chem. Commun. 2003, 2418.
(557) Widegren, J. A.; Finke, R. G. J. Mol. Catal. A 2003, 198
317.
(558) Dyson, P. J.; Ellis, D. J.; Parker, D. G.; Welton, T. J. Mol. Catal.
A 1999, 150, 71.
(559) Dyson, P. J.; Russell, K.; Welton, T. Inorg. Chem. Commun.
2001, 4, 571.
(560) Rossi, L. M.; Machado, G.; Fichtner, P. F. P.; Teixeira, S. R.;
Dupont, J. Catal. Lett. 2004, 92, 149.
(561) Daguenet, C.; Dyson, P. J. Organometallics 2004, 23, 6080.
(562) Daguenet, C.; Dyson, P. J. Organometallics 2006, 25, 5811.
(563) Umpierre, A. P.; Machado, G; Fecher, G. H.; Morais, J.;
Dupont, J. Adv. Synth. Catal. 2005, 347, 1404.
(564) Silveira, E. T.; Umpierre, A. P.; Rossi, L. M.; Machado, G.;
Morais, J.; Soares, G. V.; Baumvol, I. J. R.; Teixeira, S. R.; Fichtner,
P. F. P.; Dupont, J. Chem.—Eur. J. 2004, 10, 3734.
(565) Zhao, D.; Dyson, P. J.; Laurenczy, G.; McIndoe, J. S. J. Mol.
Catal. A 2004, 214, 19.
(566) Pugin, B.; Studer, M.; Kuesters, E.; Sedelmeier, G.; Feng, X.
Adv. Synth. Catal. 2004, 346, 1481.
(567) Jessop, P. G.; Stanley, R. R.; Brown, R. A.; Eckert, C. A.; Liotta,
C. L.; Ngo, T. T.; Pollet, P. Green Chem. 2003, 5, 123.
(568) Berger, A.; de Souza, R. F.; Delgado, M. R.; Dupont, J.
Tetrahedron: Asymmetry 2001, 12, 1825.
(569) Yinghuai, Z.; Carpenter, K.; Bun, C.-C.; Bahnmueller, S.; Ke,
C.-P.; Shanmugham Srid, V.; Kee, L.-W.; Hawthorne, M. F. Angew.
Chem., Int. Ed. 2003, 42, 3792.
(570) Xiong, W.; Lin, Q.; Ma, H.; Zheng, H.; Chena, H.; Li, X.
Tetrahedron: Asymmetry 2005, 16, 1959.
(571) Ngo, H. L.; Hu, A.; Lin, W. Tetrahedron Lett. 2005, 46, 595.
(572) Berthod, M.; Joerger, J.-M.; Mignani, G.; Vaultier, M.; Lemaire,
M. Tetrahedron: Asymmetry 2004, 15, 2219.
(573) Dyson, P. J.; Ellis, D. J.; Henderson, W.; Laurenczy, G. Adv.
Synth. Catal. 2003, 345, 216.
(574) Hu, A.; Ngo, H. L.; Lin,W.Angew. Chem., Int. Ed. 2004, 43, 2501.
(575) Lam, K. H.; Xu, L.; Feng, L.; Ruan, J.; Fan, Q.; Chan, A. S. C.
Can. J. Chem. 2005, 83, 903.
(576) Anderson, K.; Goodrich, P.; Hardacre, C.; Rooney, D. W.
Green Chem. 2003, 5, 448.
(577) Solinas, M.; Pfaltz, A.; Cozzi, P. G.; Leitner, W. J. Am. Chem.
Soc. 2004, 126, 16142.
(578) Giernoth, R.; Krumm, M. S. Adv. Synth. Catal. 2004, 346, 989.
(579) Xu, D.-Q.; Hu, Z.-Y.; Li, W.-W.; Luo, S.-P.; Xu, Z.-Y. J. Mol.
Catal. A 2005, 235, 137.
(580) Ruta, M.; Yuranov, I.; Dyson, P. J.; Laurenczy, G.; Kiwi-
Minsker, L. J. Catal. 2007, 247, 269.
(581) (a) Virtanen, P.; Karhu, H.; Kordas, K.; Mikkola, J.-P. Chem.
Eng. Sci. 2007, 62, 3660. (b) Virtanen, P.; Mikkola, J.-P.; Salmi, T. Ind.
Eng. Chem. Res. 2007, 46, 9022.
(582) Kume, Y.; Qiao, K.; Tomida, D.; Yokoyama, C. Catal. Com-
mun. 2008, 9, 369.
(583) Virtanen, P.; Salmi, T.; Mikkola, J.-P. Ind. Eng. Chem. Res.
2009, 48, 10335.
(584) Welton, T.; Smith, P. J. Adv. Organomet. Chem. 2004, 51, 251.
(585) Hu, Y.; Yu, Y.; Hou, Z.; Li, H.; Zhao, X.; Feng, B. Adv. Synth.
Catal. 2008, 350, 2077.
(586) Wang, J.; Qin, R.; Fu, H.; Chen, J.; Feng, J.; Chen, H.; Li, X.
Tetrahedron: Asymmetry 2007, 18, 847.
(587) Wang, J.; Feng, J.; Qin, R.; Fu, H.; Yuan, M.; Chen, H.; Li, X.
Tetrahedron: Asymmetry 2007, 18, 1643.
(588) Baan, Z.; Finta, Z.; Keglevich, G.; Hermecz, I. Green Chem.
2009, 11, 1937.
(589) (a) Schulz, P. S.; M€uller, N.; B€osman, A.; Wasserscheid, P.
Angew. Chem., Int. Ed. 2007, 46, 1293. (b) Schneiders, K.; B€osman, A.;
Schulz, P. S.; Wasserscheid, P. Adv. Synth. Catal. 2009, 351, 432.
(590) €Ochsner, E.; Schneiders, K.; Junge, K.; Beller, M.; Wasserscheid,
P. Appl. Catal., A 2009, 364, 8.
(591) Floris, T.; Kluson, P.; Bartek, L.; Pelantova, H. Appl. Catal., A
2009, 366, 160.
(592) Leger, B.;Denicourt-Nowicki, A.;Roucoux, A.;Olivier-Bourbigou,
H. Adv. Synth. Catal. 2008, 350, 153.
(593) Leger, B.; Denicourt-Nowicki, A.; Olivier-Bourbigou, H.;
Roucoux, A. Inorg. Chem. 2008, 47, 9090.
(594) Yang, X.; Yan, N.; Fei, Z.; Crespo-Quesada, R. M.; Laurenczy,
G.; Kiwi-Minsker, L.; Kou, Y.; Li, Y.; Dyson, P. J. Inorg. Chem. 2008,
47, 7444.
(595) Ruta, M.; Laurenczy, G.; Dyson, P. J.; Kiwi-Minsker, L. J. Phys.
Chem. C 2008, 112, 17814.
(596) Prechtl, M. H. G.; Scariot, M.; Scholten, J. D.; Machado, G.;
Teixeira, S. R.; Dupont, J. Inorg. Chem. 2008, 47, 8995.
(597) Obert, K.; Roth, D.; Ehrig, M.; Sch€onweiz, A.; Assenbaum, D.;
Lange, H.; Wasserscheid, P. Appl. Catal., A 2009, 356, 43.
(598) (a) Arras, J.; Steffan, M.; Shayeghi, Y.; Ruppert, D.; Claus, P.
Green Chem. 2009, 11, 716. (b) Arras, J.; Steffan, M.; Shayeghi, Y.;
Claus, P. Chem. Commun. 2008, 4058. (c) Steffan, M.; Lucas, M.;
Brandner, A.; Wollny, M.; Oldenburg, N.; Claus, P. Chem. Eng. Technol.
2007, 30, 481.
(599) Arras, J.; Ruppert, D.; Claus, P. Appl. Catal., A 2009, 371, 73.
(600) Khodadadi-Moghaddam, M.; Habibi-Yangjehb, A.; Gholami,
M. R. J. Mol. Catal. A 2009, 306, 11.
(601) Muzart, J. Adv. Synth. Catal. 2006, 348, 275.
(602) Marcinek, A.; Zielonka, J.; Ge-bicki, J.; Gordon, C.M.; Dunkin,
I. R. J. Phys. Chem. A 2001, 105, 9305.
(603) (a) Peng, J.; Deng, Y. New J. Chem. 2001, 25, 639. (b) Zhao,
D.; Wu, M.; Kou, Y.; Min, E. Catal. Today 2002, 74, 157. (c) Sun, J.;
Fujita, S.-i.; Bhanage, B. M.; Arai, M. Catal. Commun. 2004, 5, 83. (d)
Panchgalle, S. P.; Choudhary, S. M.; Chavan, S. P.; Kalkote, U. R.
J. Chem. Res. 2004, 550. (e) Kim, D. W.; Hong, D. J.; Seo, J. W.; Kim,
H. S.; Kim, H. K.; Song, C. E.; Chi, D. Y. J. Org. Chem. 2004, 69, 3186.
(604) Yadav, J. S.; Reddy, B. V. S.; Basak, A. K.; Narsaiah, A. V.
Tetrahedron 2004, 60, 2131.
(605) Yadav, J. S.; Reddy, B. V. S.; Basak, A. K.; Narsaiah, A. V.Chem.
Lett. 2004, 33, 248.
(606) (a) Liu, Z.; Chen, Z.-C.; Zheng, Q.-G.Org. Lett. 2003, 5, 3321.
(b) Karthikeyan, G.; Perumal, P. T. Synlett 2003, 2249. (c) Qian,W.; Jin,
E.; Bao,W.; Zhang, Y.Angew. Chem., Int. Ed. 2005, 44, 952. (d)Qian,W.;
Jin, E.; Bao, W.; Zhang, Y. J. Chem. Res. 2005, 613.
(607) Ley, S. V.; Ramarao, C.; Smith, M. D. Chem. Commun.
2001, 2278.
(608) Anthony, J. L.; Maginn, E. J.; Brennecke, J. F. J. Phys. Chem. B
2002, 106, 7315.
(609) Hardacre, C; Mullan, E. A.; Rooney, D. W.; Thompson, J. M.
J. Catal. 2005, 232, 355.
(610) Xie, H.; Zhang, S.; Duan, H. Tetrahedron Lett. 2004, 45, 2013.
(611) Jiang, N.; Ragauskas, A. J. Org. Lett. 2005, 7, 3689.
(612) Wolfson, A.; Wuyts, S.; De Vos, D. E.; Vankelecom, I. F. J.;
Jacobs, P. A. Tetrahedron Lett. 2002, 43, 8107.

BP dx.doi.org/10.1021/cr1003248 |Chem. Rev. XXXX, XXX, 000–000
Chemical Reviews
REVIEW
(613) Balandina, T. A.; Larina, T. V.; Trebushat, D. V.; Adonin,
N. Y.; Kuznetsova, N. I. React. Kinet. Catal. Lett. 2009, 97, 191.
(614) Ganchegui, B.; Bouquillon, S.; Henin, F.; Muzart, J. Tetrahe-
dron Lett. 2002, 43, 6641.
(615) Qian, W.; Jin, E.; Bao, W.; Zhang, Y. Angew. Chem., Int. Ed.
2005, 44, 952.
(616) Guan, W.; Wang, C.; Yun, X.; Hu, X.; Wang, Y.; Li, H. Catal.
Commun. 2008, 9, 1979.
(617) Wang, C.; Guan, W.; Xie, P.; Yun, X.; Li, H.; Hu, X.; Wang, Y.
Catal. Commun. 2009, 10, 725.
(618) Bernini, R.; Coratti, A.; Provenzano, G.; Fabrizi, G.; Tofani, D.
Tetrahedron 2005, 61, 1821.
(619) Owens., G. S.; Abu-Omar, M. M. J. Mol. Catal. A 2002,
187, 215.
(620) Owens, G. S.; Durazo, A.; Abu-Omar, M. M. Chem.—Eur. J.
2002, 8, 3053.
(621) (a) Bianchini, G.; Crucianelli, M.; de Angelis, F.; Neri, V.;
Saladino, R. Tetrahedron Lett. 2005, 46, 2427. (b) Chhikara, B. S.;
Tehlan, S.; Kumar, A. Synlett 2005, 63. (c) Chhikara, B. S.; Chandra, R.;
Tandon, V. J. Catal. 2005, 230, 436.
(622) Chhikara, B. S.; Tehlan, S.; Kumar, A. Synlett 2005, 63.
(623) Chhikara, B. S.; Chandra, R.; Tandon, V. J. Catal. 2005,
230, 436.
(624) (a) Branco, L. C.; Afonso, C. A. M. Chem. Commun.
2002, 3036. (b) Song, C. E.; Jung, D.-U.; Roh, E. J.; Lee, S.-G.; Chi,
D. Y. Chem. Commun. 2002, 3038.
(625) Branco, L. C.; Afonso, C. A. M. J. Org. Chem. 2004, 69, 4381.
(626) Conte, V.; Floris, B.; Galloni, P.; Silvagni, A. Adv. Synth. Catal.
2005, 347, 1341.
(627) Conte, V.; Floris, B.; Galloni, P.; Silvagni, A. Pure Appl. Chem.
2005, 77, 1575.
(628) Lo, W.-H.; Yang, H.-Y.; Wei, G.-T. Green Chem. 2003, 5, 639.
(629) K€uhn, F. E.; Zhao, J.; Abrantes, M.; Sun, W.; Afonso, C. A. M.;
Branco, L. C.; Gonc-alves, I. S.; Pillinger, M.; Rom~ao, C. C. Tetrahedron
Lett. 2005, 46, 47.
(630) Dupont, J.; Spencer, J. Angew. Chem., Int. Ed. 2004, 43, 5296.
(631) Conte, V.; Floris, B.; Galloni, P.; Mirruzzo, V.; Scarso, A.;
Sordi, D.; Strukul, G. Green Chem. 2005, 7, 262.
(632) Wang, J.-R.; Liu, L.; Wang, Y.-F.; Zhang, Y.; Deng, W.; Guo,
Q.-X. Tetrahedron Lett. 2005, 46, 4647.
(633) Namboodiri, V. V.; Varma, R. S.; Sahle-Demessie, E.; Pillai,
U. R. Green Chem. 2002, 4, 170.
(634) Li, X.; Geng,W.; Zhou, J.; Luo,W.;Wang, F.;Wang, L.; Tsang,
S. C. New J. Chem. 2007, 31, 2088.
(635) Sun, H.; Harms, K.; Sundermeyer, J. J. Am. Chem. Soc. 2004,
126, 9550.
(636) Sun, H.; Li., X.; Sundermeyer, J. J. Mol. Catal. A 2005,
240, 119.
(637) Jiang, N.; Ragauskas, A. J. J. Org. Chem. 2007, 72, 7030.
(638) Ansari, I. A.; Gree, R. Org. Lett. 2002, 4, 1507.
(639) Wu, X.-E.; Ma, L.; Ding, M.-X.; Gao, L.-X. Chem. Lett. 2005,
34, 312.
(640) Jiang, N.; Ragauskas, A. J. Tetrahedron Lett. 2005, 46, 3323.
(641) Wu, X.-E.; Ma, L.; Ding, M.-X.; Gao, L.-X. Synlett 2005, 607.
(642) Hou, Z.; Han, B.; Gao, L.; Jiang, T.; Liu, Z.; Chang, Y.; Zhang,
X.; He, J. New J. Chem. 2002, 26, 1246.
(643) Blanchard, L. A.; Hancu, D.; Beckman, E. J.; Brennecke, J. F.
Nature 1999, 399, 28.
(644) Peng, J.; Shi, F.; Gu, Y.; Deng, Y. Green Chem. 2003, 5, 224.
(645) Bianchini, G.; Crucianelli, M.; de Angelis, F.; Neri, V.;
Saladino, R. Tetrahedron Lett. 2005, 46, 2427.
(646) Li, Z.; Xia, C.-G. J. Mol. Catal. A 2004, 214, 95.
(647) (a) Angueira, J. E.; White, G. M. J. Mol. Catal. A 2005, 227, 51.
(b) Angueira, J. E.; White, G. M. J. Mol. Catal. A 2005, 238, 163.
(648) Brausch, N.; Metlen, A.; Wasserscheid, P. Chem. Commun.
2004, 1552.
(649) Fukuyama, T.; Yamaura, R.; Ryu, I. Can. J. Chem. 2005,
83, 711.
(650) Fukuyama, T.; Inouye, T.; Ryu, I. J. Organomet. Chem. 2007,
692, 685.
(651) Calo, V.; Giannoccaro, P.; Nacci, A.; Monopoli, A.
J. Organomet. Chem. 2002, 645, 152.
(652) (a)Mizushima, E.; Hayashi, T.; Tanaka,M.Green Chem. 2001,
3, 76. (b) Mizushima, E.; Hayashi, T.; Tanaka, M. Top. Catal. 2004,
29, 163.
(653) Lin, Q.; Yang, C.; Jiang, W.; Chen, H.; Li, X. J. Mol. Catal. A
2007, 264, 17.
(654) Dong, W.-S.; Zhou, X.; Xin, C.; Liu, C.; Liu, Z. Appl. Catal., A
2008, 334, 100.
(655) Shi, F.; Deng, Y.; Ma, T. S.; Peng, J.; Gu, Y.; Qiao, B. Angew.
Chem., Int. Ed. 2003, 42, 3257.
(656) Shi, F.; Zhang, Q.; Li, D.; Deng, Y. Chem.—Eur. J. 2005,
11, 5279.
(657) (a) Zhang, Q.; Shi, F.; Gu, Y.; Yang, J.; Deng, Y. Tetrahedron
Lett. 2005, 46, 5907. (b) Shi, F.; Zhang, Q.; Gu, Y.; Deng, Y. Adv. Synth.
Catal. 2005, 347, 225.
(658) Shi, F.; Peng, J.; Deng, Y. J. Catal. 2003, 219, 372.
(659) Shi, F.; He, Y.; Li, D.; Ma, Y.; Zhang, Q.; Deng, Y. J. Mol. Catal.
A 2006, 244, 64.
(660) Kim, H. S.; Kim, Y. J.; Bae, J. Y.; Kim, S. J.; Lah, M. S.; Chin,
C. S. Organometallics 2003, 22, 2498.
(661) McNulty, J.; Nair, J. J.; Robertson, A. Org. Lett. 2007, 9, 4575.
(662) Haumann, M.; Riisager, A. Chem. Rev. 2008, 108, 1474.
(663) Webb, P. B.; Sellin, M. F.; Kunene, T. E.; Williamson, S.;
Slawin, A. M. Z.; Cole-Hamilton, D. J. J. Am. Chem. Soc. 2003,
125, 15577.
(664) Lin, Q.; Fu, H.; Jiang, W.; Chen, H.; Li, X. J. Chem. Res.
2007, 216.
(665) Lin, Q.; Jiang,W.; Fu, H.; Chen, H.; Li, X.Appl. Catal., A 2007,
328, 83.
(666) Williams, D. B. G.; Ajam, M.; Ranwell, A. Organometallics
2007, 26, 4692.
(667) Tominaga, K.-I.; Sasaki, Y. Chem. Lett. 2004, 33, 14.
(668) Chauvin, Y.; Gilbert, B.; Guibard, I. J. Chem. Soc., Chem.
Commun. 1990, 1715.
(669) Olivier-Bourbigou, H.; Travers, P.; Chodorge, J. A. Pet.
Technol. Q. 1999, Autumn, 141.
(670) (a) Wasserscheid, P.; Gordon, C. M.; Hilgers, C.; Muldoon,
M. J.; Dunkin, I. R. Chem. Commun. 2001, 1186. (b) Wasserscheid, P.;
Hilgers, C.; Keim, W. J. Mol. Catal. A 2004, 214, 83.
(671) Silvana, S. M.; Suarez, P. A. Z.; de Souza, R. F.; Dupont
J. Polym. Bull. 1998, 40, 401.
(672) Picquet, M.; Tkatchenko, I.; Tommasi, I.; Wasserscheid, P.;
Zimmermann, J. Adv. Synth. Catal. 2003, 345, 959.
(673) Picquet, M.; Poinsot, D.; Stutzmann, S.; Tkatchenko, I.;
Tommasi, I.; Wasserscheid, P.; Zimmermann J. Top. Catal. 2004, 29, 139.
(674) Ligabue, R. A.; Dupont, J.; de Souza, R. F. J. Mol. Catal. A
2001, 169, 11.
(675) Dullius, J. E. L.; Suarez, P. A. Z.; Einloft, S.; de Souza, R. F.;
Dupont, J.; Fischer, J.; De Cian, A. Organometallics 1998, 17, 815.
(676) Conte, V.; Elakkari, E.; Floris, B.; Mirruzzo, V.; Tagliatesta, P.
Chem. Commun. 2005, 1587.
(677) B€ohm, V. P. W.; Herrmann, W. A. Chem.—Eur. J. 2000,
6, 1017.
(678) Plechkova, N. V.; Seddon, K. R. Chem. Soc. Rev. 2008, 37, 123.
(679) B€ohm, V. P. W.; Herrmann, W. A. Chem.—Eur. J. 2000,
6, 1017.
(680) Calo, V.; Nacci, A.; Monopoli, A.; Ferola, V. J. Org. Chem.
2007, 72, 2596.
(681) (a) Handy, S. T.; Okello, M. Tetrahedron Lett. 2003, 44, 8395.
(b) Jeffery, T. Tetrahedron 1996, 52, 10113.
(682) (a) Xu, L.; Chen, W.; Ross, J.; Xiao, J. Org. Lett. 2001, 3, 295.
(b) Mo, J.; Xu, L.; Xiao, J. J. Am. Chem. Soc. 2005, 127, 751.
(683) Mo, J.; Liu, S.; Xiao, J. Tetrahedron 2005, 61, 9902.
(684) Forsyth, S. A.; Gunaratne, H. Q. N.; Hardacre, C.; McKeown,
A.; Rooney, D. W. Org. Process Res. Dev. 2006, 10, 94.

BQ dx.doi.org/10.1021/cr1003248 |Chem. Rev. XXXX, XXX, 000–000
Chemical Reviews
REVIEW
(685) Forsyth, S. A.; Gunaratne, H. Q. N.; Hardacre, C.; McKeown,
A.; Rooney, D. W.; Seddon, K. R. J. Mol. Catal. A 2005, 231, 61.
(686) Calo, V.; Nacci, A.; Lopez, L.; Mannarini, N. Tetrahedron Lett.
2000, 41, 8973.
(687) Calo, V.; Nacci, A.; Monopoli, A.; Spinelli, M. Eur. J. Org.
Chem. 2003, 1382.
(688) Rosa, J. N.; Santos, A. G.; Afonso, C. A. M. J. Mol. Catal. A
2004, 214, 161.
(689) Ross, J.; Xiao, J. Chem.—Eur. J. 2003, 9, 4900.
(690) (a) Mathews, C. J.; Smith, P. J.; Welton, T. Chem. Commun.
2000, 1249. (b) McLachlan, F.; Mathews, C. J.; Smith, P. J.; Welton, T.
Organometallics 2003, 22, 5350.
(691) (a) Mathews, C. J.; Smith, P. J.; Welton, T. J. Mol. Catal. A
2003, 206, 77. (b)Mathews, C. J.; Smith, P. J.; Welton, T. J. Mol. Catal. A
2004, 214, 27.
(692) Yang, X.; Fei, Z.; Geldbach, T. J.; Phillips, A. D.; Hartinger,
C. G.; Li, Y.; Dyson, P. J. Organometallics 2008, 27, 3971.
(693) Thomas, P. A.; Marvey, B. B. Int. J. Mol. Sci. 2009, 10, 5020.
(694) Marvey, B. B.; Segakweng, C. K.; Vosloo, M. H. Int. J. Mol. Sci.
2008, 9, 615.