Fundamentals of Ionic Liquids E-Book

Douglas R. MacFarlane

• Herausgeber: Wiley-VCH Verlag GmbH & Co. KGaA
• Kategorie: Wissenschaft und neue Technologien
• Sprache: Englisch

Seitenzahl: 545

Veröffentlichungsjahr: 2017

Beschreibung

Written by experts who have been part of this field since its beginnings in both research and academia, this textbook introduces readers to this evolving topic and the broad range of applications that are being explored.
The book begins by examining what it is that defines ionic liquids and what sets them apart from other materials. Chapters describe the various types of ionic liquids and the different techniques used to synthesize them, as well as their properties and some of the methods used in their measurement. Further chapters delve into synthetic and electrochemical applications and their broad use as "Green" solvents. Final chapters examine important applications in a wide variety of contexts, including such devices as solar cells and batteries, electrochemistry, and biotechnology.
The result is a must-have resource for any researcher beginning to work in this growing field, including senior undergraduates and postgraduates.

Leseprobe

Table of Contents

Cover

Title Page

Copyright

Chapter 1: An Introduction to Ionic Liquids

1.1 Prologue

1.2 The Definition of an Ionic Liquid

1.3 A Brief Perspective

1.4 Aprotic Versus Protic ILs

1.5 An Overview of IL Applications

1.6 Key Properties and Techniques for Understanding ILs

1.7 New Materials Based on ILs

1.8 Nomenclature and Abbreviations

References

Chapter 2: The Structure of Ions that Form Ionic Liquids

2.1 Introduction

2.2 Ionic Interactions and the Melting Point

2.3 Effect of Ion Size and Crystal Packing

2.4 Charge Delocalization and Shielding

2.5 Ion Asymmetry

2.6 Influence of Cation Substituents

2.7 Degrees of Freedom and Structural Disorder

2.8 Short-Range Interactions – Hydrogen Bonding

2.9 Dications and Dianions

2.10

T

m

Trends in Other IL Families

2.11 Concluding Remarks

References

Chapter 3: Structuring of Ionic Liquids

3.1 Introduction

3.2 Ionicity, Ion Pairing and Ion Association

3.3 Short-Range Structuring

3.4 Structural Heterogeneity and Domain Formation

3.5 Hydrogen Bonding and Structure

3.6 Experimental Probes of Structure

3.7 Simulation Approaches to Understanding Structure

3.8 Structuring at Solid Interfaces

3.9 Ionic Liquid Structure in Confined Spaces

3.10 Impact of Structure on Reactivity and Application

3.11 Concluding Remarks

References

Chapter 4: Synthesis of Ionic Liquids

4.1 Introduction

4.2 Synthesis of ILs

4.3 Characterization and Analysis of ILs

4.4 Concluding Remarks

References

Chapter 5: Physical and Thermal Properties

5.1 Introduction

5.2 Phase Transitions and Thermal Properties

5.3 Surface and Tribological Properties

5.4 Transport Properties and their Inter-relationships

5.5 Properties of Ionic Liquid Mixtures

5.6 Protic ILs, Proton Transfer, and Mixtures

5.7 Deep Eutectic Solvents and Solvate ILs

5.8 Concluding Remarks

References

Chapter 6: Solvent Properties of Ionic Liquids: Applications in Synthesis and Separations

6.1 Introduction – Solvency and Intermolecular Forces

6.2 Liquid–Liquid Phase Equilibrium

6.3 Gas Solubility and Applications

6.4 Synthetic Chemistry in ILs – Selected Examples

6.5 Inorganic Materials Synthesis

6.6 Biomass Dissolution

6.7 Concluding Remarks

References

Chapter 7: Electrochemistry of and in Ionic Liquids

7.1 Basic Principles of Electrochemistry in Nonaqueous Media

7.2 The Electrochemical Window of Ionic Liquids

7.3 Redox Processes in ILs

7.4 Electrodeposition and Cycling of Metals in ILs

7.5 Electrosynthesis in Ionic Liquids

7.6 Concluding Remarks

References

Chapter 8: Electrochemical Device Applications

8.1 Introduction

8.2 Batteries

8.3 Fuel Cells

8.4 Dye-Sensitized Solar Cells and Thermoelectrochemical Cells

8.5 Supercapacitors

8.6 Actuators

8.7 Concluding Remarks

References

Chapter 9: Biocompatibility and Biotechnology Applications of Ionic Liquids

9.1 Biocompatibility of Ionic Liquids

9.2 Ionic Liquids from Active Pharmaceutical Ingredients

9.3 Biomolecule Stabilization in IL Media

9.4 Concluding Remarks

References

Index

End User License Agreement

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Guide

Cover

Table of Contents

Begin Reading

List of Illustrations

Chapter 1: An Introduction to Ionic Liquids

Figure 1.1 Illustration of the difference between the volatilization of PILs, where neutral species evaporate into the gas phase and aprotic ILs that exist as tightly bound ion pairs in the gas phase.

Chapter 2: The Structure of Ions that Form Ionic Liquids

Figure 2.1 Interaction energies for (a) uncharged atoms and (b) ions.

Figure 2.2 Free-energy change with temperature for a solid and a liquid.

Figure 2.3 Within a sea of ions, there will be a combination of attractive forces (e.g. between cation and anion) and repulsive forces (e.g. between cation and cation).

Figure 2.4 Effect of alkyl chain length,

n

, on the temperature at which different [C

n

mim]

+

salts become liquid. This transition temperature represents either the

T

g

(mostly likely those data points below about −50 °C, marked with hollow symbols) or

T

m

. (Data from López-Martin

et al

. 2007 [7].)

Figure 2.5 Packing diagrams for [NMe

4

][BF

4

] viewed down the

c

-axis showing (a) full unit cell contents (the [BF

4

]

−

ions are shown with their measured disorder) and (b) charged atoms only (N

+

shown in blue, charge center of [BF

4

]

−

shown in orange) [4]. Packing diagrams for [Prop]Br as viewed down the

c

-axis showing (c) full unit cell contents and (d) charged atoms only (N

+

shown in blue, Br

−

shown in red).

Figure 2.6 (a) Charge density on the dicyanamide (left) and [NTf

2

]

−

(right) anions. The blue and red colors represent negative and positive charge density, respectively. (Izgorodina 2011 [18]. Reproduced with permission of Royal Society of Chemistry.) (b) Space-filled structure showing the steric shielding of the N

+

(shown in blue) in the

N

-methyl-

N

-butylpyrrolidinium cation.

Figure 2.7 Examples of anions with delocalized charges.

Figure 2.8 The structure of [C

3

mpyr]Cl shown with 50% thermal ellipsoids (hydrogen atoms are shown by spheres of set size, as these are placed in geometrically fixed idealized positions) [22].

Figure 2.9 Examples of asymmetric charge-delocalized anions.

Figure 2.10 Cation structure and

T

m

values for the salts when paired with the [PF

6

]

−

anion.

Figure 2.11 Two conformers of the [NTf

2

]

−

anion.

Figure 2.12 Strength of a hydrogen bond has both distance and angle criteria.

Figure 2.13 Structures of example dications, and the

T

m

of the salts using [NTf

2

]

−

anions.

Chapter 3: Structuring of Ionic Liquids

Figure 3.1 Gas-phase geometry of [P

4,4,4,4

][Cl] by molecular orbital theory calculations.

Figure 3.2 Probability distributions of (a) the anions and (b) the imidazolium cations around an imidazolium cation derived from the EPSR model for liquid [C

1

mim]Cl, [C

1

mim][PF

6

], and [C

1

mim][NTf

2

]. For anion distributions in the top panel (a), the contour level was chosen to enclose the top 5% of the molecules, while for the cation distributions in lower panel (b), the contours enclose the top 20% of molecules each in the distance range 0–9 Å.

Figure 3.3 Snapshot presenting the fitted ethylammonium nitrate bulk structure at thermal equilibrium (298 K). From left to right, the front face of the 3D simulation box of (a) 500 [EtNH

3

]

+

cations and 500 [NO

3

]

−

anions, (b) apolar –C–C– domains only (–NH

3

and NO

3

omitted), and (c) anionic 500 [NO

3

]

−

only (the 500 [EtNH

3

]

+

omitted). Atom coloring is C (gray), H (white), N (blue), O (red).

Figure 3.4 Dependence of

D

spacing on the length

n

of the alkyl chain on the cation [21].

Figure 3.5 Raman spectra of [ethylimidazolium][NTf

2

] at 20 °C and 140 °C. The fitting (Lorentzian) components corresponding to the cis and trans contributions are shown in red (solid line) and blue (dashed line), respectively, and the structures of the conformers are shown on the left. The spectra show that the relative population of trans conformer decreases (and cis increases) with temperature.

Figure 3.6 Example of a radial distribution function plot, redrawn from a study on [C

6

mim][NTf

2

] [43]. RDFs are taken between the central nitrogen of the anion and the center of mass of the imidazolium cation. Black line: cation–anion, red line: anion–anion, blue line: cation–cation. It is evident that the polar network extends over several nanometers without losing the correlation between shells of ions of opposing sign.

Figure 3.7 Snapshots of simulation boxes containing 700 ions of [C

4

mim][PF

6

] (left, length of box, l = 49.8 Å) and [C

6

mim][PF

6

] (right, length of box, l = 52.8 Å). Red versus green color shows charged (polar) versus nonpolar domains, respectively, and aggregation of the alkyl chains in nonpolar domains when the alkyl side chain of the IL is longer than or equal to four.

Figure 3.8 Side view of the first adsorbed layer of ionic liquids [C

2

mim][NTf

2

] (left) and [C

8

mim][NTf

2

] (right) on mica. Atoms of the cation are largely in blue or black. The atoms of the mica substrate are red, yellow, and purple. This is the result from one set of simulations – two sets are compared in Ref. [49].

Figure 3.9 Representative snapshot extracted from the simulation of [C

4

mim][PF

6

] and carbide-derived carbon nanoporous electrodes. Green spheres: [PF

6

]

−

; red spheres: [C

4

mim]

+

; blue rods: carbon–carbon bonds.

Chapter 4: Synthesis of Ionic Liquids

Scheme 4.1 (a) General mechanism of quaternization. (b) Synthesis of 1-ethyl-3-methylimidazolium bromide [C

2

mim]Br.

Scheme 4.2 General mechanism for the alkylation of 1-methylimidazole with methyltriflate to produce 1,3-dimethylimidazolium trifluoromethanesulfonate [C

1

mim][CF

3

SO

3

].

Scheme 4.3 Generalized metathesis reaction for exchanging anions.

Scheme 4.4 Synthesis of [C

2

mim][BF

4

] via a two-step reaction: (a) quaternization and (b) metathesis.

Scheme 4.5 Resin-based ion-exchange reaction.

Scheme 4.6 Synthesis of [Ch][Ac] via a two-step ion-exchange reaction.

Scheme 4.7 Synthesis of [C

2

mim][BF

4

] via a methylcarbonate salt.

Scheme 4.8 (a) Schematic of a microreactor used to synthesize

N

-3-(3-trimethoxysilylpropyl)-1-methyl imidazolium chloride via a flow process. (Löwe

et al.

(2010) [14]. Reproduced with permission of Elsevier.) (b) Conventional method of synthesizing

N

-3-(3-trimethoxysilylpropyl)-1-methyl imidazolium chloride via reflux.

Scheme 4.9 Alkylation of 1-methyl imidazole with methyl-trifluoromethanesulfonate to produce 1-methyl-3-methyl imidazolium trifluoromethanesulfonate [C

1

mim][CF

3

SO

3

].

Scheme 4.10 Formation of the solvate ionic liquid [Li(G4)][NTf

2

] from equivalent mixtures of [G4] and Li[NTf

2

].

Scheme 4.11 Synthesis of an alkoxy-ammonium salt.

Scheme 4.12 Synthesis of the zwitterionic liquid 1-methylimidazolium-3-(propanesulfonate) via the reaction between

N

-methyl imidazole and propanesultone.

Scheme 4.13 (a) Synthesis of 3-butyl-1-(butyl-4-sulfonyl)imidazolium zwitterion and (b) reaction between the zwitterion and trifluoromethanesulfonic acid to form the Bronsted acid IL 3-butyl-1-(butyl-4-sulfonyl)imidazolium trifluoromethanesulfonate.

Scheme 4.14 One-pot synthesis of imidazolium-based mixture ILs with various cations and a common anion.

Figure 4.1 Synthesis of a polymer IL from an IL. (Tomé and Marrucho 2016 [36]. Reproduced with permission of Royal Society of Chemistry.)

Scheme 4.15 Synthesis of a polymerized methacryloyl-functionalized imidazolium IL. (Yuan and Antonietti 2011 [39]. Reproduced with permission of Elsevier.)

Scheme 4.16 Synthesis of

n

-butylammonium acetate.

Scheme 4.17 Synthesis of an amino acid-based IL.

Scheme 4.18 Preparation of a chiral oxazolinium salt from (

S

)-valine methyl ester. (Baudequin

et al.

2005 [45]. Reproduced with permission of Elsevier.)

Chapter 5: Physical and Thermal Properties

Figure 5.1 DSC traces showing examples of typical thermal behavior of ILs (a)

T

g

only, (b)

T

g

, crystallization, and melt, and (c) solid–solid phase transition(s) before the melt, as seen in organic ionic plastic crystals.

Figure 5.2 Entropy as a function of temperature illustrating the Kauzman Paradox.

Figure 5.3 Effect of ion pair volume on the glass transition.

Figure 5.4 A disc of a typical organic ionic plastic crystal (OIPC), in this case [C

2

mpyr][NTf

2

] being removed with a scalpel blade.

Figure 5.5 Illustration of the difference between the volatilization of PILs, where neutral species evaporate into the gas phase, and aprotic ILs that can exist as tightly bound ion pairs in the gas phase. (Adapted from Earle

et al

. 2006 [10] with permission from Nature Publishing Group.)

Figure 5.6 Dynamic equilibrium of DIMCARB.

Figure 5.7 (a) Typical rising temperature thermogravimetric analysis (TGA) used to determine the onset of decomposition. (Armel

et al.

2011 [23]. Reproduced with permission of Royal Society of Chemistry.), and (b) example plot showing the time taken for 1% degradation of an IL, determined by a series of isothermal TGA runs at different temperatures. (Baranyai

et al.

(2004) [24]. Reproduced with permission from CSIRO Publishing.)

Figure 5.8 (a) Proposed decomposition of the NTf

2

anion followed by (b) possible mechanisms of attack on the [C

2

mim] cation, in [C

2

mim][NTf

2

] [35] and (c) possible decomposition pathways for [C

4

mim][Cl] and [C

4

mim][PF

6

] [34].

Figure 5.9 Potential energy profile calculated by

ab initio

computational methods during the motion of an iodide ion from one face (behind) to the other (in front) of the imidazolium ring. The maxima at around 160 and 240 ° represent the highest barriers (∼16 kJ mol

−1

) to the motion. The equivalent barrier in dimethylethylimidazolium iodide is >40 kJ mol

−1

.

Figure 5.10 Simulated impedance plane plot for a 4000 Ω electrolyte resistance (

R

el

).

Figure 5.11 Walden plot of molar conductivity versus inverse viscosity for a range of ionic liquids.

Figure 5.12 Correlations between (a) molecular volume and viscosity and (b) molecular volume and conductivity; filled squares: [C

n

(CN)mim] salts of various anions; triangles: [C

n

mim][BF

4

] or [PF

6

] salts; circles: [C

n

mim][N(CN)

2

] salts; diamonds [C

n

mim][NTf

2

] salts.

Figure 5.13 Binary phase diagram for two solids.

Figure 5.14 Eutectic formation in a binary mixture of salts.

Figure 5.15 Structure of [Me-DBU][NTf

2

].

Figure 5.16 Viscosity data for mixtures of [C

4

mim][Me

2

PO

4

] and [C

4

mim][NTf

2

] fitted to Eq. (5.16) (dashed line) and Eq. (5.18) (solid line).

Figure 5.17 Molar conductivity data for mixtures of [C

4

mim][Me

2

PO

4

] and [C

4

mim][NTf

2

], using both impedance spectroscopy (Λ

m

, Imp) and the self-diffusion coefficients (Λ

m

, NMR), fitted to Eq. (5.19).

Figure 5.18 Walden plot data for addition of various ILs to [C

3

mpyr][NTf

2

] [65]. The common starting point in each mixture series is [C

3

mpyr][NTf

2

], shown by the arrow.

Figure 5.19 Deviations Δ

W

of the data from the ideal line in the Walden plot with increasing concentrations of [NH

4

][SCN] in [C

2

mim][OAc] and [C

2

mim][EtOSO

3

].

Figure 5.20 Four possible outcomes in protic ionic liquids (a) mostly un-ionized acid and base; (b) hydrogen-bonded acid–base pairs (no proton transfer); (c) mostly ionized PIL (proton transfer has occurred, independent ions); (d) associated PIL (proton transfer has occurred, but ions associated through hydrogen bonding).

Figure 5.21 (a) Walden plot for binary PIL [dema][HSO

4

] + [dema][NTf

2

] mixtures and (b) ionicity of binary PIL [dema][HSO

4

] and [dema][NTf

2

] mixtures at 30 °C.

Chapter 6: Solvent Properties of Ionic Liquids: Applications in Synthesis and Separations

Figure 6.1 Appearance and position of a UCST and LCST near a region of single-phase liquid mixtures.

Figure 6.2 DFT calculations of the interactions between N

2

and a variety of IL anions. In the case of the [fap]

−

anion, there are two modes of interaction in the N

2

–anion complex [17].

Scheme 6.3 Formation of three different products from toluene in the presence of different ILs.

Scheme 6.2 Formation of the Wheland intermediate [24].

Scheme 6.3 Heck reaction in the presence of palladium acetate catalyst and IL [25].

Figure 6.3 Imidazolium-based ionic liquid tagged with a palladium complex.

Scheme 6.4 Reaction of cyclopentadiene with dimethyl maleate [31].

Figure 6.4 H-bonding interaction of an imidazolium cation with the carbonyl oxygen of methyl acrylate in the activated complex of a Diels–Alder reaction.

Scheme 6.5 (a) Reaction of acetyl chloride with benzene in the presence of an acidic chloroaluminate IL, and (b) the proposed mechanism of the acylation reaction [36].

Scheme 6.6 (a) Reaction between benzaldehyde and methyl acrylate catalyzed by DABCO, and (b) the proposed general mechanism of the role of the nucleophilic amine in the Baylis–Hilman reaction [38].

Scheme 6.7 Methanolysis of linalool chloride in [C

4

mim][NTf

2

] [42].

Scheme 6.8 Reaction of diethylchlorophosphate with ethanol-

d

6

in the presence of various ILs [44].

Figure 6.5 Structure of [C

2

mim]Br-zeolite template material consisting of hexagonal prismatic units with the formula Al

8

(PO

4

)

10

H

3

·3C

6

H

11

N

2

.

Chapter 7: Electrochemistry of and in Ionic Liquids

Figure 7.1 Schematic of a three-electrode electrochemical cell. (CE = counter-electrode; RE = reference electrode; WE = working electrode).

Figure 7.2 Scanning potential techniques: current–voltage profile during a cyclic voltammetry experiment.

Figure 7.3 Comparison of an aqueous reference electrode and a nonaqueous reference electrode, both composed of a metal (Ag) electrode separated from the bulk electrolyte by a fritted tube.

Figure 7.4 Electrochemical windows of [C

2

mim][NTf

2

] (red) and [C

2

C

1

mim][NTf

2

] (blue) versus I

−

/I

−

3

.

Figure 7.5 (a) Electrochemical window of [C

4

mim][BF

4

]+16 wt% Cl

−

at (1) 25 °C, (2) 40 °C, (3) 60 °C, (4) 70 °C, and (5) 80 °C; (Li

et al.

2006 [11]. Reproduced with permission of Elsevier.) (b) Electrochemical window of (1) vacuum-dried (24 h, 60 °C) [C

4

mim][BF

4

], (2) atmospheric [C

4

mim][BF

4

], and (3) wet [C

4

mim][BF

4

] at 25 °C. (O'Mahony

et al.

2008 [12]. Reproduced with permission of American Chemical Society.)

Figure 7.6 Electroactive species that can be used as internal calibrants in ILs.

Figure 7.7 Examples of redox-active species used in DSSCs.

Figure 7.8 CVs of a 10 mmol

−1

solution of [Co(bpy)

3

]

2+

in [C

2

mim][B(CN)

4

] at different scan rates (20, 50, 100, 200, 500, 800, and 1000 mV s

−1

) on a glassy carbon electrode. Each CV shows the two redox reactions: one (at positive potentials) attributed to the [Co(bpy)

3

]

2+/3+

redox couple, and the other (at negative potentials) to [Co(bpy)

3

]

1+/2+

.

Figure 7.9 (a) Reduction of benzophenone in a dry IL [27]. (b) Reversible one-electron electrochemical oxidation of TEMPO.

Figure 7.10 Cyclic voltammograms obtained for the reduction of [Bu

4

N]

4

[α-S

2

W

18

O

62

] at a 3-mm diameter glass carbon electrode, at the scan rate 0.1 V s

−1

. (a) Solution (5 mM) in [C

4

mim][BF

4

] and (b) adhered solid in contact with [C

4

mim][BF

4

].

Figure 7.11 Examples of redox-active ILs, where the electroactive species are incorporated into the cation or anion or both.

Figure 7.12 Cyclic voltammograms of 0.1 M (black line) or 0.2 M (red line) ZnX

2

in an [NTf

2

]

−

IL and also 0.2 M·ZnX

2

with 2.5 wt% water added (blue line) versus Zn/Zn

2+

at 25 °C (where X = [NTf

2

]

−

) on a glassy carbon electrode.

Figure 7.13 Overview of possible electrosynthesis reactions, showing reactive intermediates that are formed electrochemically and a range of possible subsequent reactions to form the final product.

Figure 7.14 Electrochemical window of fluoride-containing ionic liquids.

Figure 7.15 Electrochemical fluorination of pthalides in an ionic liquid.

Figure 7.16 Electrochemical synthesis of organic carbamates in an ionic liquid.

Figure 7.17 Electrochemical reduction of benzoyl chloride to benzil in an ionic liquid.

Figure 7.18 Electroreductive coupling of organic halides in an ionic liquid.

Chapter 8: Electrochemical Device Applications

Figure 8.1 Schematic of lithium-ion battery (arrows indicate direction of Li

+

motion during discharge and the direction of ‘normal’ current in the circuit (electron flow is in the opposite direction).

Figure 8.2 Cyclic voltammogram of Li deposition and stripping in an IL electrolyte, comparing the neat [P

1,1,1,i4]

[fsi] and solutions of 0.5, 3.2, and 3.8 mol kg

−1

Li[fsi] in [P

1,1,1,i4]

[fsi] at a Ni working electrode with a potential sweep rate of 20 mV s

−1

at 25 °C.

Figure 8.3 Dependence of total ion conductivity and viscosity as a function of Li

+

concentration in ‘inorganic–organic’ ILs based on mixtures of [P

1,1,1,i4

][fsi] and Li[fsi] at 20, 50 and 80 °C. In this mixture, a concentration of Li[fsi] = 3.2 mol kg

−1

corresponds to an equimolar mixture of Li[fsi] and [P

1,1,1,i4

][fsi]. Beyond this point, the IL contains more Li

+

ions than [P

1,1,1,i4

] cations.

Figure 8.4 Charging of a Li|inorganic–organic IL|LiCoO

2

cell compared to one containing a standard organic carbonate electrolyte.

Figure 8.5 Li(glyme) solvate structure in a solvate IL.

Figure 8.6 Schematic of the operation of a polymer electrolyte membrane fuel cell. The PEM is sandwiched between two porous, electrically conductive, and catalytically active electrodes.

Figure 8.7 (a) Schematic of H

+

transport in a PIL of type [BH

+

][A

−

]. Here, the protonated cations [BH

+

] of the PIL deprotonate at the cathode and generate free base species [B], which migrate toward the anode to serve as a vehicle for the H

+

species formed by the hydrogen oxidation reaction. The anions [A

−

] of the PIL, due to their low basicity, either serve as the proton defects (hopping sites) or can also serve as carriers (vehicles) if sufficiently basic. (Rana

et al.

2012 [32]. Reproduced with permission of Elsevier.) (b) Examples of proton-conducting protic ILs and OIPCs, where dema = diethylmethylammoniun and dmeda =

N

,

N

-dimethylethylenediamine.

Figure 8.8 Redox shuttling in a dye-sensitized solar cell or a thermoelectrochemical cell, with the Co

II/III

(bpy)

3

redox couple as an example.

Figure 8.9 (a) Configuration of a simple electrochemical double-layer capacitor and (b) the ion distribution and potential profile at the positive electrode, where φ

M

is the potential at the electrode surface and φ

s

is the potential of the Stern layer. Further from the electrode is the Gouy–Chapman diffuse layer.

Chapter 9: Biocompatibility and Biotechnology Applications of Ionic Liquids

Figure 9.1 Toxicity data (EC

50

values) as a function of cation side-chain length

n

in halide and [BF

4

]

−

salts; Black squares:

Chlorella vugaris

in [C

n

mim]Cl; Red circles: Leukaemia cells IPC-81 in [C

n

mim]Cl; Blue triangles:

V. fischeric

in [C

n

mim][BF

4

]; Green triangles:

S. obliquus

in [C

n

mim]Br.

Figure 9.2 Biodegradable nicotinic and pyridinic cation derivatives.

Figure 9.3 Preparation of propantheline cyclamate as a dual-active pharmaceutical ionic liquid.

Figure 9.4 Confocal microscopy images of (i) blue stained nucleic acids and (ii) a yellow fluorescent protein expressing pDNA transfected into HEK 293T cells after storage in phosphate buffered saline solution (left), 20% [Ch][dhp] (middle), and 50% [Ch][dhp] (right) after 28 days of storage at 37 °C. The much higher rate of transfection on the right indicates much greater stability of the pDNA during storage in the 50% [Ch][dhp].

Figure 9.5 Stabilization of small interfering RNA in hydrated [Ch][dhp].

Figure 9.6 Action of added acid on a choline H

2

PO

4

−

/HPO

4

2−

([Ch][dhp/mhp]) buffer ionic liquid. The spectra are of phenol red indicator present in the IL. The addition of up to 50 mol% acid produces little change in the spectrum of the indicator. Adding 250 mol% acid, which swamps the buffer, produces a distinct color change and shift in the spectrum. (MacFarlane

et al.

2010 [20]. Reproduced with permission of Royal Society of Chemistry.)

Figure 9.7 Fluorescence micrographs of L-929 cells growing actively on [Ch]tartrate cross-linked collagen materials. Scale bars represent 50 mm.

List of Tables

Chapter 1: An Introduction to Ionic Liquids

Table 1.1 Glossary of structures and nomenclature abbreviations used in this book

Chapter 5: Physical and Thermal Properties

Table 5.1 Surface tension (γ) of ionic liquids at 20 °C (in mN m

−1

)

Table 5.2 Viscosity for various ILs in mPa.s at

T

= 25 °C

Table 5.3 Ionic conductivities of various ILs in mS/cm at

T

= 25 °C

Chapter 6: Solvent Properties of Ionic Liquids: Applications in Synthesis and Separations

Table 6.1 Distribution coefficient

D

for Cd

2+

and Cu

2+

between IL and aqueous solution at pH 7 ± 5%

Table 6.2 Henry's Law constants for a variety of gases in different ionic liquids at 298 K

Table 6.3

Chapter 9: Biocompatibility and Biotechnology Applications of Ionic Liquids

Table 9.1 Protic pharmaceutical ILs and modulation of their solubility properties

Domínguez de María, P. (ed.)

Ionic Liquids in Biotransformations and Organocatalysis

Solvents and Beyond

2012

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Ohno, H. (ed.)

Electrochemical Aspects of Ionic Liquids, 2nd edition

2 Edition

2011

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Wasserscheid, P., Welton, T. (eds.)

Ionic Liquids in Synthesis

2 Edition

2008

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ISBN: 978-3-527-62119-4

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Endres, F., MacFarlane, D., Abbott, A. (eds.)

Electrodeposition from Ionic Liquids

2008

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Wasserscheid, P., Stark, A. (eds.)

Handbook of Green

Chemistry - Green Solvents

Volume 6 - Ionic Liquids

2010

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Fundamentals of Ionic Liquids

From Chemistry to Applications

Douglas R. MacFarlane, Mega Kar and Jennifer M. Pringle

Authors

Prof. Douglas R. MacFarlane

Monash University

School of Chemistry

Wellington Rd

Clayton

3800 Melbourne

Australia

Dr. Mega Kar

Monash University

School of Chemistry

Wellington Rd

Clayton

3800 Melbourne

Australia

Dr. Jennifer M. Pringle

Deakin University

Institute for Frontier Materials

221 Burwood Hwy

Burwood

3125 Melbourne

Australia

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Cover Design Grafik-Design Schulz

Chapter 1An Introduction to Ionic Liquids

1.1 Prologue

The benefits of using salts in their liquid form as electrolytes or reaction media have long been recognized. For example, Faraday developed his laws of electrolysis in the 1830s using molten metal halide salts. However, most researchers would rather avoid using solvents that require heating to hundreds of degrees. Therefore, while ‘high-temperature molten salts’ are extremely useful for certain applications, they are not in widespread use in research laboratories or industry. In contrast, ‘ionic liquids’ – and we will discuss the definition of these below – are proving to be very exciting for a very wide range of applications at more moderate temperatures. Particularly over the last decade, scientists working in many different areas of research have started to realize the special properties of ionic liquids (ILs) and to embrace the promise of these new materials. As more applications are discovered, and more families of ILs are developed, so the field continues to grow. Thousands of papers on ILs are now published every year, and, more importantly, increasing numbers of researchers are experimenting with ILs and discovering firsthand how their unique properties can help their work.

So here we are, just over 100 years since the first ‘room-temperature ionic liquid’ was discovered, at a point where these materials are now proving so promising and widely applicable that it is important for a broad range of scientists and engineers to have an appreciation of the basics of ILs. Thus, the purpose of this book is to serve as an introduction to the key concepts and applications of ILs for those venturing into this field for the first time. Hopefully, the book will also inspire further curiosity and enthusiasm for exploring these exciting and very unique materials.

Our goal in this book is to provide a thorough introduction to the field appropriate to the level of a finishing undergraduate science student or a beginning postgraduate student. Our emphasis is on illustrative examples and the background chemistry sufficient to understand the fundamentals of ILs and their applications. For further reading, we have referenced more extensive reviews where they exist. To provide background on fundamental concepts and methods that may not be readily accessible in standard textbooks, we have included Concept Toolbox items as breakout text boxes in various places throughout the chapters. This first chapter provides a broad overview of the field, the materials involved, their properties, and their applications, of which more details can be found later in this book.

1.2 The Definition of an Ionic Liquid

The phrase ‘ionic liquid’ was coined only relatively recently to refer to ambient-temperature liquid salts, and the definition has since been the subject of much discussion and some evolution. The most useful practical definition of an IL is

‘A liquid comprised entirely of ions.’

We can delve into this a little deeper. By this definition, is an IL different from a molten salt? The answer is: ‘No’ – the term ‘molten salt’ refers to the liquid phase of a crystalline salt, for example, NaCl. ‘Ionic Liquid’ covers that, but also covers a broader range of possibilities. Imagine a mixture of the two salts Na[fsi] and [C3mpyr][NTf2] (see Table 1.1 for an explanation of abbreviations). This mixture of salts is a liquid at room temperature and is properly called an IL by our definition. In fact an IL could contain a very large number of different ions. Note that, in common usage, the term ‘molten salt’ has also come to mean mixtures of salts, although the term itself clearly indicates a single compound.

Table 1.1 Glossary of structures and nomenclature abbreviations used in this book

Quaternary cations

Abbreviation

Structure

Alkylmethylimidazolium

[C

n

mim]

+

Alkyldimethylimidazolium

[C

n

dmim]

+

Alkylmethylpyrrolidinium

[C

n

mpyr]

+

Alkylmethylpiperidinium

[C

n

mpip]

+

Ammonium

[N

n

,

n

,

n

,

n

]

+

Phosphonium

[P

n

,

n

,

n

,

n

]

+

Ether-functionalized

[N

R,R,R,2O1

]

+

,

etc.

Cholinium

[Ch]

+

Sulfonium

[S

R,R′,R″

]

+

Guanidinium

[Gdm]

+

Tetrafluoroborate

[BF

4

]

−

Hexafluorophosphate

[PF

6

]

−

Trifluoroacetate

[tfa]

−

Triflate

[OTf]

−

or [CF

3

SO

3

]

−

Bis(fluorosulfonyl)imide

a

[fsi]

−

or [N(SO

2

F)

2

]

−

Bis(trifluoromethanesulfonyl)imide

a

[NTf

2

]

−

or [N(SO

2

CF

3

)

2

]

−

Dicyanamide

[N(CN)

2

]

−

or [dca]

−

Tetracyanoborate

[B(CN)

4

]

−

Fluoroalkylphosphates

[fap]

−

, [efap]

−

, etc.

Dihydrogenphosphate

[H

2

PO

4

]

−

or [dhp]

−

Hydrogen sulfate

[HSO

4

]

−

p

-Toluenesulfonate or tosylate

[Tos]

−

Tetrahydroborate or borohydride

[BH

4

]

−

a Also referred to as ‘amides’.

In principle (in fact an important thermodynamic principle), the IL obtained by mixing Na[fsi] and [C3mpyr][NTf2] is exactly the same as that obtained by mixing the appropriate quantities of Na[NTf2] and [C3mpyr][fsi]. The points of origin are irrelevant in defining the IL; only the quantities of the individual ions present are important. In fact, such ILs with very high concentrations of Na or Li salts are proving to be highly effective as electrolytes for Na and Li batteries [1].

Some definitions of IL add a temperature range, such as ‘below 100 °C’, to the definition but this is not necessary. In fact, it is limiting to do so, since it can blinker our perspective on which compounds or mixtures may be useful for certain applications. Indeed, there are many, quite valuable, applications of ILs at temperatures above 100 °C, for example, the preparation of MnOx water oxidation catalysts by electrodeposition at 130 °C [2]. The key requirement for this is that the IL be a liquid at 130 °C. It is convenient if it is also liquid at room temperature, but it is not necessary for this to be the case and one should certainly not exclude from consideration compounds having melting points >100 °C for an application such as this. Similarly, a definition that includes ‘a salt having a melting point below 100 °C’ (or some other temperature such as room temperature) is also an unnecessary restriction because in some cases the melting point may be practically difficult to find and measure. The supercooling of liquids below their equilibrium melting points is a well-known phenomenon, and in some cases the liquid becomes so viscous that the crystalline phase never forms on a practical timescale. This is particularly true of mixtures of salts, which we have agreed are perfectly good ILs, because the melting points of individual compounds is often sharply depressed in mixtures. We will discuss the melting points of ILs in Chapter 2, and multicomponent phase diagrams and behavior further in Chapter 5. With all of this in mind, referring to a melting point in our definition of an IL becomes unhelpful.

The meaning of the word ‘ion’ in this definition also needs some discussion. Species such as Cl− are obvious, as are simple molecular ions such as [NO3]−. Things become more subtle when we consider metal coordination complex ionic species such as [AlCl4]−, which were used extensively during the early work on ILs, as reviewed briefly in the following. These are certainly ionic as written, and as long as they continue to stay bound in the real liquid, for long times, then they fit our definition. However, there is always an equilibrium process by which such complex species are formed, and therefore we must always recognize the presence, in equilibrium, of some amount of the component species. In the case of [AlCl4]−, this might be Al3+ and Cl−; as long as these components involved are also ionic, we still have an IL, although with more complex ‘speciation’ than is at first apparent. This speciation, and how it responds to variables such as temperature, can influence the IL's properties significantly.

In the case of Co(H2O)62+, which sits in equilibrium with H2O and Co2+, the situation is more complicated; [Co(H2O)6][NO3]2 is a simple hydrate salt that melts into a stable liquid at 55 °C. Where the other species are neutral (e.g., H2O), we have non-ionic components intruding into our IL and at some stage we must begin to think of the liquid as a mixture rather than a pure IL (recall that our definition includes the word ‘entirely’). It becomes important to consider where the dividing line lies between a ‘pure’ IL, meaning containing only the ions we think are present, and a ‘mixture’ of ions and other species. MacFarlane and Seddon proposed some time ago that a practical approach to this problem be adopted based on the fact that very few of any of our laboratory chemicals are absolutely pure – in fact, 100% purity is a practical impossibility. So, in reality, we may accept something as ‘pure’ if it is at least 95% of the desired compound – or perhaps 99% for high-quality/sensitivity work. Another perspective, which underlies much of our laboratory chemistry, is that we consider something to be pure if the level of contaminants (often starting with water and the ever-present dissolved N2 and O2) is not sufficient to significantly affect its properties; this requires thought, judgment, and discussion of what is ‘significant’. This, quite reasonable, practical approach avoids pedantic arguments about whether something does or does not fit our definition of an IL. This also allows our definition to cover a host of very interesting and important complex-ion species.

Therefore, to summarize, throughout this book we will simply refer to an IL as being any ‘liquid that is comprised entirely of ions’, with the word ‘entirely’ being used in a practical rather than absolute sense. Although our definition now includes the classical field of molten salts, we will focus mostly on the ambient-temperature ILs that have become of immense interest in the last 20 years.

Other types of liquid systems that may be considered related to ILs but fall outside of our definition, and are not covered here, are liquid mixtures of ions with molecular species or ions with water. These complex liquid systems have been discussed in a number of recent reviews and books [3–7].

1.3 A Brief Perspective

In this book we will discuss many of the different types of ILs now available and the applications being developed. However, it is also important to appreciate the origins of this field and the path through different ion chemistries that has led us this point. As our definition above highlights, it is a mistake to put ILs and high-temperature molten salts into two separate classifications and treat them as different fields. While this book primarily concentrates on those species that are air- and water-stable and liquid at room temperature, because those are the focus of the majority of present research reports, other salt families also have much to teach us [8, 9].

The first ‘discovery’ of a room-temperature IL is often attributed to Paul Walden who made ethyl ammonium nitrate, which is a protic ionic liquid (a class of IL discussed further below) [10]. He recognized that the low melting point (13–14 °C) was a result of the larger organic cation that decreases the degree of ion association compared to an inorganic salt. It is also likely that liquid organic salts were recognized by organic chemists with annoyance long before Walden, when they probably described them as ‘intractable oils’ and disposed of them.

Remarkably, the ability of ammonium salts to dissolve cellulose was first recognized in 1934 [8, 11]; the use of ILs for processing biomass (which commonly includes cellulose) is now an important and extensive area of investigation, discussed in more detail in Chapter 6.

Angell and his group began to investigate ambient-temperature systems in the 1970s [12], and then in the 1980s chloroaluminate-based salts were developed that combined pyridinium or imidazolium cations with the tetrachloroaluminate anion ([AlCl4]−), again demonstrating the value of a large organic cation in reducing the melting point [9, 13, 14]. Depending on the composition, some of these salts are liquid at room temperature and primarily found application in batteries and metal electrodeposition. However, the composition of these materials is quite complex as, depending on the relative concentrations of the two components, they can also form multivalent species such as [Al2Cl7]− and have either (Lewis) acidic, neutral, or basic properties. These ILs (and their starting materials) are also very moisture sensitive, with [AlCl4]− hydrolyzing to form corrosive HCl, which severely limits their application.

One of the most significant advances in the evolution of ILs came with the discovery of air- and water-stable species in 1992 [15] by Wilkes and his group. This report combined the 1-ethyl-3-methylimidazolium cation ([C2mim]+, which is still one of the most widely used cations today) with [BF4]−, [NO3]−, [NO2]−, and acetate anions. We now know that some fluorinated anions (e.g., [BF4]− and [PF6]−) are in fact susceptible to hydrolysis [16], particularly at increased temperatures, and produce HF. Nonetheless, the significantly improved stability that these anions impart compared to the chloroaluminates is highly beneficial.

From this point in time, the field of ILs expanded rapidly, both in terms of the different ions used and in the range of applications being investigated. These applications include those on a research lab-scale and also an industrial scale; the commercial use of ILs has been under development since the late 1990s [8]. Table 1.1 shows the structures and common abbreviations of the most widely used IL ions. It can be seen that there are a variety of different fluorinated anions in use because these commonly produce low melting points, relatively low viscosities, and good electrochemical stability. This is primarily a result of the charge delocalization over the anion (because F is so electronegative). The relationship between the chemical structure of ILs and the different physical properties (viscosity, melting point, etc.) is discussed throughout this book. The [NTf2]− anion is a classic example of a fluorinated anion [17, 18], and for the same reasons as this is attractive for IL electrolytes, it has also been widely used in the battery community for many years (originally just as the lithium salt, Li[NTf2]).

The cations used to make ILs are still predominantly nitrogen-based, although phosphonium cations are becoming increasingly common and can impart better stability [19, 20]. Methods for synthesizing common nitrogen-based ILs are discussed in Chapter 4. Another emerging family of ILs comprises ‘solvate ionic liquids’, which are concentrated mixtures of salts and a solvent – most commonly lithium salts and oligoether (also known as ‘glyme’) solvents [21, 22]. The multiple oxygens on the glyme strongly solvate the lithium ion, essentially forming a [Li(glyme)]+ cation that is balanced by the counter-ion from the lithium salt (most commonly [NTf2]−). Thus, they have a relatively ionic nature and similar behavior to ILs (e.g., low vapor pressure) despite the presence of a molecular component. Not all concentrated mixtures of lithium salt and solvent form ‘solvated ILs’, that is, if the coordination of the Li+ is not strong enough. In that case, the physical properties of the mixture would not be consistent with a pure cation/anion combination and would therefore fall outside our definition of an IL. However, many concentrated mixtures do form ‘solvated ILs’ and these show great promise in various battery applications [21, 22].

While much of the early work on ILs/molten salts was driven by interest in their use as electrolytes, the potential benefit of these materials to ‘green chemistry’ – on a lab scale and in industry – has also been an important driving force. However, ILs should not be claimed as ‘green’ solvents simply because they have negligible vapor pressure (and thus do not emit the toxic vapors of many volatile organic solvents). Factors such as biodegradation, toxicity, and recycling of the IL must also be considered. The ‘12 Principles of Green Chemistry’ [23] also include parameters such as reducing the number of reaction steps, minimizing overall energy consumption, and decreasing the amount of materials used, which are as important for designing green engineering processes; all of these must be considered before classifying the use of an IL as ‘green’. The toxicity and biodegradation of ILs is discussed further in Chapter 9.

Finally, one other frequently proclaimed feature of ILs is their ‘tunability’. Once we understand that we can use large charge-diffuse ions to decrease the melting point of high-temperature molten salts, then we realize that there are an extraordinarily large number of ions that we can choose from to make our IL. Therefore, if we know which properties the different ions will produce, we can ‘tune’ the IL to suit our application (e.g., choosing a fluorinated anion to increase the electrochemical stability or decrease miscibility with water). The phrase ‘task-specific’ ionic liquid is sometimes used, which originally referred to the use of functionalized alkyl groups to improve the partitioning of specific metal ions into the IL phase from water [24], but this terminology has since expanded in scope. The use of IL mixtures to achieve the desired properties allows a further dimension in tunability, and this is discussed further in Chapter 5. In general, however, we are still some way from being able to design a cation/anion combination from first principles to optimize a specific physical property. The reasons for this will likely become clearer as we progress through this book and the complexity of the relationships between the chemical and the physical properties of ILs is discussed further. In the future, computational techniques for modeling ILs, which are introduced briefly in the following and in Chapter 5, are likely to bring significant benefit toward ‘designing ILs to suit the task’.

1.4 Aprotic Versus Protic ILs

Thus far we have discussed ions that do not contain any labile (or transferable) protons; these are called ‘aprotic’ ILs meaning ‘not protic’. On the other hand, ILs can actually be formed by proton transfer from a Bronsted acid to a Bronsted base, and these are termed ‘protic ionic liquids’ (PILs). This family has some unique characteristics that need additional consideration with respect to how they fit within our definition of an IL.

The physical properties of PILs can depend very strongly on the relative strengths of the acid and base, as this determines the extent of proton transfer to form the salt:

1.1

The free-energy change associated with the proton transfer process, as well as the relationship between this and ΔpKa (the difference between the pKa values of the acid and base), has a direct influence on properties of the liquid such as its vapor pressure, which in this case is nonzero at most temperatures, and its ‘ionicity’ (as discussed further in Chapter 5). This is because the free-energy change determines the extent of proton transfer and thus the amount of free acid and base present in the PIL. It was this problem that prompted Seddon and MacFarlane to propose the 95% (or 99%) guideline discussed above as a practical definition of when these could be called ILs. The presence of these neutral species affects the melting point (one could think of these as impurities within the PIL). Also, it can be difficult to remove all traces of water from PILs, as extended drying under vacuum will remove the volatile acid or base components, shifting the above equilibrium and changing the composition.

It is important to note that the acid/base stoichiometry in a PIL does not have to be 1 : 1; in fact, it is quite hard to achieve genuinely 1 : 1 stoichiometry. The actual composition can impact significantly the chemical and thermal properties of the liquid [25, 26]. Such off-stoichiometry compositions must also be considered to be mixtures in most cases – the melting behavior of mixtures is also discussed further in Chapter 5.

1.5 An Overview of IL Applications

When considering all the different applications of ILs, of which there is an ever-growing number, it is useful to relate these back to their unique properties. We often introduce ILs as having properties that can include ‘low melting point, negligible vapor pressure, good electrochemical and thermal stability, and tunable structures’, and so forth, although it is very important in such an introduction to note that not all ILs have all of these properties. So how do these characteristics translate to application, and which properties still need the most investigation and improvement? The following provides some examples of particular properties attracting certain applications.

The low melting point of ILs drives interest in their use as pharmaceutical salts, where the cation or anion is an active pharmaceutical ingredient [27–29]. The low melting point removes the concern of a salt crystallizing into an alternative polymorph (crystal structure) from that which has been trialed and patented, as the formation of polymorphs has significant medical and legal implications. Having the drug in a liquid form may also make it easier to be administered to patients. This application of ILs is discussed further in Chapter 9.

The ionic nature of ILs also means that they provide quite unique solvation environments compared to conventional molecular solvents, and this is exploited in a variety of different synthetic reactions, materials processing/extraction, and gas separation (as discussed further in Chapter 6). There is also extensive interest in their use for biomass processing. Biomass is a complex mixture of materials that can include cellulose, hemicellulose, keratin, lignin, chitin, and so on, depending on its source, and it is a valuable sustainable resource. However, these material mixtures are notoriously insoluble in molecular solvents, which makes their chemical processing difficult. ILs are being investigated for both the dissolution of a variety of different biomaterials and for their processing into higher value products [29, 30]. For biotechnological applications, the ability of ILs to dissolve and stabilize enzymes and proteins, DNA, and RNA is also extremely valuable. Note, however, that in these applications often the IL must contain a small amount of water, and therefore they are sometimes referred to as ‘hydrated ILs’. These applications are discussed further in Chapter 9.

Unique solubilizing properties, coupled with good electrochemical stability, also underlie the use of ILs for rare-earth processing and recycling [31]. Rare-earth metals are used in significant quantities globally as they have unique magnetic, luminescent, and electrochemical properties, but supplies are increasingly becoming limited. ILs are promising for two techniques – as an ionic extractant for the separation of rare-earth salts, and as the medium for the subsequent electrodeposition of the pure rare-earth metal. The electrodeposition of rare-earth metals requires very electrochemically stable, aprotic media, and until recently this process has primarily utilized high-temperature molten salts. The use of ILs for the electrodeposition of a variety of metals is discussed in Chapter 7.

Excellent electrochemical stability is arguably one of the most important characteristics of some ILs, as evidenced by their extensive use in electrochemical devices, for electrowinning, water splitting, and so on [32–34]. Water splitting describes the electrolysis of water into hydrogen and oxygen. This is of great interest for the large-scale production of hydrogen as a fuel, particularly if the energy for the electrolysis can be provided by the sun (i.e., photoelectrochemical water splitting). Central to this process is an efficient electrocatalyst for the water oxidation reaction, as it is the kinetics of this reaction that is the more limiting factor. ILs have been used as a medium for the synthesis of good water-splitting catalysts such as MnOx, where x is typically between 1.5 and 2 [2]. Although these deposition and oxidation processes are still under investigation, one of the key properties of the IL here appears to be the structure that it imparts to the liquid phase (and the liquid/electrode interphase) during the reaction, which influences the thermodynamics of the process. The use of ILs to enhance the water oxidation reaction (i.e., decrease the overpotential) is also central to their use in metal–air batteries. The use of ILs in electrochemical processes and devices is discussed in Chapters 7 and 8.

In addition to the synthesis of water-splitting catalysts mentioned previously, ILs have, of course, found application in a wide range of other synthetic reactions – organic, inorganic, biological, and so on. Again, related to the unique properties of ILs, the benefits of their use are multifold. Most simply, the ability to dissolve materials that are insoluble in common organic solvents is highly advantageous, while their ionic nature helps in stabilizing nanoparticle dispersions during synthesis. At a more sophisticated level, the term ‘ionothermal synthesis’ describes the use of ILs as both the solvent and as a structure-directing tool [29, 35]. This has been widely utilized for the synthesis of a range of valuable materials including molecular sieves, metal–organic frameworks (MOFs), reduced graphene oxides, and polyoxometalates, during which the nature of the IL used can impact the resultant material structure. In this application, the low vapor pressure and good thermal stability of ILs also provide an important advantage: such reactions use temperatures up to 200 °C, and therefore if water is used (i.e., in ‘hydrothermal synthesis’), this requires the reaction to be performed in an autoclave. In contrast, the ionothermal reaction can often be carried out in ordinary equipment. However, while the templating effects of ILs for the synthesis of these materials are intriguing, full understanding or control of the self-assembly processes that underlie such synthesis requires further investigation. The nature of the IL and their structure (in the liquid phase and at solid/liquid interfaces) also influence the synthesis of a variety of nanoparticles [36]. The structural ordering of ILs, as well as its impact on some synthetic processes, is discussed in Chapter 3.

Other synthetic applications of ILs include the use of chiral ILs in stereoselective synthesis, where they can be used as the catalyst or as the synthetic medium [37]. Here, the most important and unique property of the IL is the ability to incorporate a chirally pure moiety into the ionic structure, the source of which may be natural (e.g., amino acid-based ILs) or synthetic. Chiral ILs also have application in chromatography and spectroscopy, for example, as NMR chiral shift reagents.

In addition to chiral catalysis, ILs can also be used for heterogeneous, homogeneous, and biocatalysis (whole-cell and enzyme-catalyzed). This is a very wide-ranging application of ILs and is the focus of multiple books and reviews [38–40], and is discussed further in Chapter 6. In this application there are two important material classes: supported ILs, that is, the use of a porous solid support to immobilize an IL containing a dissolved homogeneous metal catalyst, and ‘solid catalyst with IL layer’, where a heterogeneous catalyst is coated with an IL to improve the properties [40]. In terms of the application of ILs in the liquid form, these are commonly used in a biphasic system with another solvent. In this application, the solubilizing properties of the IL are the key.

Although it is not yet a large field, it is interesting to note the use of ILs as heat transfer fluids, as this is an example of an application that benefits from the relatively high heat capacity and good thermal stability of ILs. For this application, low vapor pressures, viscosity, and corrosivity are also important requirements. Nanoparticles can also be added to the IL to improve the heat capacity and thermal conductivity [41, 42].

In the field of energetic materials [43], the huge structural variability of ILs is a great advantage, as are their low vapor pressure, wide liquid range, and good thermal stability. Energetic materials are those with a large amount of stored chemical energy that can be released by, for example, shock, heating, or applying friction, and are thus used as explosives, propellants, and so on. As one class of such materials, ‘hypergolic ILs’ refer to ILs designed to ignite when in contact with a suitable oxidizer. These ILs are being designed as a replacement for hydrazine, which is highly toxic and difficult to handle, for use in propellants. Nitrogen-containing heterocyclic cations (e.g., substituted alkyl ammonium, imidazolium, or triazolium cations) are commonly combined with anions such as nitrate, dicyanamide, nitrocyanamide, cyanoborate, azide, or aluminum borohydride (which contain energetic groups –NO2, –N3