Correlation of Structures and Solvent Properties of Ionic Liquids

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‘Thank the Lord’ ?

That sounded like a prayer.

A prayer in a public school.

God has no place within these walls,

just like facts don’t have a place

within an organized religion.

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Abstract.

The aim of the thesis was to find out a correlation between the structural properties of ionic liquids and their influence on organic reactivity.

Kinetic studies were performed on nucleophilic substitutions (Menschutkin reaction and dediazonation, Chapter 1) and Diels-Alder reactions (Chap-ter 2). Data obtained, together with ab-initio calculations, confirm an ac-tive role of the hydrogen bond interaction between ionic liquid and sub-strates/activated complex in the influence of the reactivity.

In the dediazonation reactions, competition between the scarce nucle-ophilic [Tf₂N]− _{anion and bromide, surprisingly bears to the formation of}

the incorporating products with [Tf₂N]− _{only (section 1.2.2). ESI-MS}

exper-iments on pure and mixed ionic liquids were performed in order to understand the inner interactions responsible for this peculiar behaviour (section 1.2.3). In order to better elucidate the structural features responsible of the physical properties of ionic liquids, two computer aided approaches were used in this thesis: prediction of the melting point of a class of pyridinium substituted ionic liquids using a recursive neural network model (Chapter 3) and calculation of the physical properties of [bmim][PF₆] and [bmim][BF₄] through molecular dynamic simulations (Chapter 4).

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List of Figures

1 Structure of some typical cations of ionic liquids. . . xvi 2 Synthesis path for the preparation of ionic liquids examplified

for an imidazolium salt. . . xix 3 Comparison of the partial radial distribution functions for (a)

the cation-cation distribution and (b) the cation-anion distribu-tion for the 1,3-dimethylimidazolium chloride, hexafluorophos-phate, and bis{(trifluoromethyl)sulfonyl}imide salts. Each ra-dial distribution function is calculated from the center of the imidazolium ring, from the phosphorus atom in the case of [PF6],

and from the nitrogen atom in the case of [Tf2N]−. . . xxviii

4 Probability distributions of (a) the anions and (b) the imida-zolium cations around an imidaimida-zolium cation derived from the EPSR model for liquid [mmim]Cl, [mmim][PF6], and [mmim][Tf2N]. xxx

5 Reichardt’s betaine dye. . . xxxvi 6 Solvent polarity scale according to EN

T values. . . xxxviii

7 Competition between anion and solute for hydrogen bonding with cation. . . xliii

1.1 ILs cations investigated with ESI-MS technique. . . 7

1.2 Menschutkin reaction between N-methylimidazole and various substituted benzyl halides. . . 9 1.3 Kinetic plot of the reaction between N-methyl-imidazole and

p-methyl-benzyl chloride in [bm2im][Tf2N] at 333 K. 265 nm,

pathlenght=0.1 cm. . . 11

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iv LIST OF FIGURES

1.4 Linear plot kobsagainst various concentration of N-methyl-imidazole.

Reactions with p-methoxy-benzyl chloride in [bm2im][Tf2N] at

333 K. . . 12 1.5 SN2 mechanism. . . 18

1.6 Hammett plots for the reaction between N-methylimidazole and various substituted benzyl chlorides in organic solvents and ionic liquids at 333 K and 1 atm. Data shown in Tables 1.3, 1.4, 1.5, 1.6, 1.7, 1.8, 1.9 . . . 19 1.7 Hammett plot for the reaction between N-methylimidazole and

various substituted benzyl chlorides in [bmim][Tf2N] at 333 K

and 1 atm. Data shown in Tables 1.3, 1.4, 1.5, 1.6, 1.7, 1.8, 1.9 20 1.8 Mechanism schemes for dediazonation.. . . 24 1.9 Kinetic plot of the reaction between PhN2BF4 (8 × 10

−4 _{M) and}

[bmim][Br] (8 × 10−3 _{M) in [bmim][Tf}

2N] at 298 K, 262 nm,

path-lenght=0.1 cm . . . 26 1.10 Two compounds obtained from the incorporation of [Tf2N]− by

PhN2BF4 at room temperature. . . 26

1.11 Kinetic plot of the reaction between PhN2BF4 (8 × 10

−4 _{M) and}

[bmim][Tf2N] at 298 K, 262 nm, pathlenght=0.1 cm. . . 27

1.12 Peak distribution in positive-ion ESI mass spectra of [emimX] in acetonitrile, X = Br−_{, OTs}−_{, I}−_{, PF}−

6, Tf2N−. . . 30

1.13 Relative abundances of the fragment ions in the MS/MS spec-tra of [emim· · ·Br· · ·hmim]+

and [bupy· · ·Br· · ·picol]+

supramolecu-lar aggregates (m/z values refer to 79_Br). _{. . . 31}

1.14 Calculated (B3LYP/6-311+G(d,p)//HF/6-31G(d)) geometries of the homologous [(PF6)· · ·bmim· · ·(PF6)]− complex. . . 36

1.15 Calculated (B3LYP/6-311+G(d,p)//HF/6-31G(d)) geometries of the two isomers of [(PF6)· · ·bmim· · ·(TsO)]− complex. . . 36

1.16 Calculated (B3LYP/6-311+G(d,p)//HF/6-31G(d)) geometries of the homologous [(BF4)· · ·bmim· · ·(BF4)]

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Index v

1.17 Calculated (B3LYP/6-311+G(d,p)//HF/6-31G(d)) geometries of the two isomers of [(PF6)· · ·bmim· · ·(BF4)]

− _complex._{. . . 38}

1.18 Kinetic scheme of the competition between [Tf2N]

− _{and Br}−

to react with PhN2BF4 . . . 39

2.1 The energy diagram for concerted and stepwise mechanisms of the Diels-Alder reaction between butadiene and ethene. . . 49 2.2 Woodward-Hoffmann (WH), Salem, Houk, and Alston (SAH)

and Singleton’s [4+3] (S43) secondary orbital interactions in a Diel-Alder reaction between butadiene and acrolein. . . 50 2.3 Diels-alder reaction scheme for cyclopentadiene and methyl

acry-late (R=−COOMe), acrolein (R=−CHO) and acrylonitrile (R=−CN). 54 2.4 Comparison of the endo-selectivities of the Diels-Alder reaction

of cyclopentadiene with different dienophiles. . . 59 2.5 The relationship between the log(endo/exo) of the reaction

be-tween cyclopentadiene and acrylonitrile against the Berson’s empirical solvent parameter at 298 K, Ω298. . . 61

2.6 The relationship between the log(endo/exo) of the reaction be-tween cyclopentadiene and acrolein against the Berson’s empir-ical solvent parameter at 298 K, Ω298. . . 62

2.7 Linear plot for ln(endo/exo) against the hydrogen bond donor ability of the solvent α for the reaction with acrolein at 298 K (R2_{= 0.919).} _{. . . 65}

2.8 Calculated versus observed selectivities of the Diels-alder be-tween acrolein and cyclopentadiene for several solvents at 298 K (R2_{= 0.983). ln(endo/exo) was calculated from equation 2.4.}_{. . 67}

2.9 Interaction between the permanent dipoles of cyclopentadiene and methyl acrylate in the transition state for the formation of the endo- and exo-adduct. . . 69 2.10 Second order plot obtained using GC technique of the reaction

between acrylonitrile and cyclopentadiene in [Hbim][Tf2N] at 298

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vi LIST OF FIGURES

2.11 Pseudo-first order plot obtained using UV technique of the reac-tion between acrylonitrile and cyclopentadiene in [bm2im][Tf2N]

at 298 K. The rate constant kobswas obtained fitting the

exper-imental curve with the formula reported in equation 1.2. . . 72 2.12 Linear plot obtained using UV technique (kobsvs various dienophile

concentrations) of the reaction between acrylonitrile and cy-clopentadiene in [bm2im][Tf2N] at 298 K. . . 74

2.13 Calculated versus observed second order rate constants of the Diels-Alder reaction between acrolein and cyclopentadiene for several solvents at 298 K. Calculated data are obtained from equation 2.9. . . 78

2.14 Calculated versus observed second order rate constants of the Diels-Alder reaction between methyl acrylate and cyclopentadi-ene for several solvents at 298 K. Calculated data are obtained from equation 2.9 . . . 80 2.15 Isokinetic relationship between ∆H‡ _{and ∆S}‡ _{observed for the}

Diels-Alder reaction between cyclopentadiene and acrolein in different solvents. . . 86

2.16 Isokinetic relationship between ∆H‡ _{and ∆S}‡ _{observed for the}

Diels-Alder reaction between cyclopentadiene and methyl acry-late in different solvents. . . 87 2.17 Isokinetic relationship between ∆H‡ _{and ∆S}‡ _{observed for the}

Diels-Alder reaction between cyclopentadiene and acrylonitrile in different solvents. . . 88

2.18 Transition state geometries. [Hbim]+

and acrolein. . . 95

2.19 Transition state geometries. [Hbim]+

and methyl acrylate. . . . 95 2.20 Transition state geometries. [bm2im]

+ _{and acrylonitrile.} _{. . . 95}

2.21 Energetic estimation of the effect of the cation coordination on the Diels-Alder reaction. . . 97 2.22 HOMO orbitals for TS with exo (bottom) and endo (top)

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Index vii

2.23 The HOMO and LUMO energies for cyclopentadiene compared to the energies of the dieneophiles with and without cation coordination.. . . 100

3.1 Outline of the traditional QSPR/QSAR approach. Structural features of the molecule are represented by a proper set of nu-merical descriptors obtained by using different encoding ap-proaches (f function). The target property is then obtained applying a regression procedure (g function). . . 109

3.2 Outline of the RecNN approach to QSPR/QSAR. RecNN takes as input the molecular graph generated by fR encoding it by

the adaptive function fE. The g function uses this internal

representation to produce the final predicted value. . . 111

3.3 Representation of the N-(2-fluoro-5-hydroxybenzyl)-3-methylpyridinium cation as a tree. The numbers indicate the priority of the nodes

on cyclic moieties. . . 113

4.1 Schematic representation of [bmim]+_{and [PF}

6]−and [BF4]−united

atom molecular topology based on the GROMOS FF. Atom types are described in Table 4.1. Charge groups are enclosed by dashed lines. . . 151 4.2 Plots of the pressure fluctuactions against time for [bmim][PF6]

at 298 K and 1 atm. M = _R_t+t₀ t0 Pij(t ′_)dt′2 t0 . . . 158 4.3 Plots of the pressure fluctuactions against time for [bmim][BF4]

at 298 K and 1 atm. M = _R_t+t₀ t0 Pij(t ′_)dt′2 t0 . . . 159 4.4 Plot of velocity vs time for [bmim][BF4] at 298 K and 1 atm.

Periodic perturbation method. . . 160

4.5 Center of mass radial distribution functions of [bmim][PF6] at

298 K and 1 atm. For the cation only the imidazolium ring is considered. . . 161

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viii LIST OF FIGURES

4.6 Center of mass radial distribution functions of [bmim][BF4] at

298 K and 1 atm. For the cation only the imidazolium ring is considered. . . 162 4.7 Spatial distribution probability density of anions around cations

for [bmim][PF6] at 298 K and 1 atm. The contours enclose regions

with a probability density above 0.000025. . . 164

4.8 Spatial distribution probability density of anions around cations for [bmim][BF4] at 298 K and 1 atm. The contours enclose regions

with a probability density above 0.000025. . . 165 4.9 DFT geometry optimization for [(PF6)· · ·bmim· · ·(PF6)]− at the

b3lyp/6-311++g(d,p) level. . . 165

4.10 DFT geometry optimization for [(BF4)· · ·bmim· · ·(BF4)]

− _{at the}

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List of Tables

1 Effect of cation and anion structure on the melting point of imidazolium based ionic liquids . . . xxiv 2 ET(30) and ETN values for several ionic liquids and moleculare

solvents. . . xxxvii

3 EN

T and Kamlet-Taft values for a selection of ionic liquids and

molecular solvents. . . xli

1.1 Permittivity (ε), solvatochromic parameters (α, β, π∗_{) and}

co-hesive pressure (δ2_{) for various solvents.} _{. . . 10}

1.2 Second order rate constant for the reaction between N-methylimidazole and benzyl halide in various solvents at 298 K. . . 11

1.3 Second order rate constant for the reaction between N-methylimidazole and benzyl chloride in various solvents at 333 K. . . 15 1.4 Second order rate constant for the reaction between N-methylimidazole

and p-metoxy-benzyl chloride in various solvents at 333 K. . . . 15 1.5 Second order rate constant for the reaction between N-methylimidazole

and m-metoxy-benzyl chloride in various solvents at 333 K. . . 16

1.6 Second order rate constant for the reaction between N-methylimidazole and p-methyl-benzyl chloride in various solvents at 333 K. . . . 16 1.7 Second order rate constant for the reaction between N-methylimidazole

and m-methyl-benzyl chloride in various solvents at 333 K.. . . 17 1.8 Second order rate constant for the reaction between N-methylimidazole

and p-chloro-benzyl chloride in various solvents at 333 K. . . . 17

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x LIST OF TABLES

1.9 Second order rate constant for the reaction between N-methylimidazole

and p-bromo-benzyl chloride in various solvents at 333 K. . . . 18

1.10 Slope ρ (with RMSE in brackets) of the Hammett correlations for electron donating substituents (EDS) and electron with-drawing substituents (EWS). . . 21

1.11 Relative quantities of compounds 1 and 2 (cfr. Figure 1.10) obtained from the incorporation of [Tf2N]− by PhN2BF4at room temperature and variable quantities of [bmim][Tf2N] and [bmim][Br]. 28 1.12 Relative abundances of fragment ions (I) originating from the MS/MS decompositions of [C1_{· · ·Br· · ·C}2_]+ _species. _{. . . 32}

1.13 Qualitative order of intrinsic bond strength of various anions to [Hmim]+_{, [emim]}+_{, [bmim]}+_{, and [bm} 2im] +_{. The ESI-MS elution} solvent is reported in brackets. . . 33

1.14 Energy values required to reduce to one-half the starting inten-sity of [X1 · · ·bmim· · ·X1 ]− _{species in CID mass spectra.} _{. . . 35}

1.15 Retention time of the nucleophilic substitution products. . . 42

1.16 Relative yields of the nucleophilic substitution products. . . 43

1.17 Fragmentation pattern of compounds 1 and 2 (GC-MS). . . 43

1.18 IR, 1_H-NMR, 13_{C-NMR data for compound 1 and 2 mixture.} _{. 44}

2.1 solvatochromic parameters for organic solvents and ionic liq-uids: EN T is Reichardt’s solvent polarity index, α and β are a measure of the solvent hydrogen bond donor acidity and ac-ceptor basicity of the solvent respectively and π∗ _{is an index of} solvent dipolarity/polarizability. . . 56

2.2 parameters for organic solvents and ionic liquids: VM the mo-lar volume, ∆U is the internal energy of the solvent, δ2 _{is the} cohesive pressure and η is the viscosity. . . 57

2.3 endo/exo selectivities observed for the reaction between cy-clopentadiene and three dienophiles (acrolein, methyl acrilate and acrylonitrile) at 298 K. . . 58

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Index xi

2.4 data of the log(endo/exo) of the Diels-Alder reaction of cyclopen-tadiene with methyl acrylate, acrylonitrile and acrolein at 298 K. . . 60 2.5 fitting parameters and correlation coefficients for the Berson’s

plot log(endo/exo)298 = a + bΩ298 for reactions with acrolein and

acrylonitrile at 298 K. . . 63 2.6 Different dual parameter correlations of ln(endo/exo) with α, π∗

and VM for reaction between cyclopentadiene and acrolein at

298 K. . . 66 2.7 Comparison of the significance of the proposed solvent

parame-ters that affect the endo- selectivity of the Diels-Alder reaction between cyclopentadiene and methyl acrylate at 298 K. . . 68 2.8 Kamlet-Taft descriptors that characterize the dienophiles

stud-ied in this work. . . 70 2.9 Second order rate constants of the Diels-Alder reaction between

cyclopentadiene and three dienophiles at 298 K. . . 73 2.10 activation parameters for the Diels-Alder reaction between acrolein

and cyclopentadiene in various solvents in the temperature range 288-298 K. ∆G‡ _{was calculated from ∆G}‡ _{= ∆H}‡_{− T ∆S}‡ _where

T = 298 K. . . 83 2.11 activation parameters for the Diels-Alder reaction between methyl

acrylate and cyclopentadiene in various solvents in the tem-perature range 288-298 K. ∆G‡ _{was calculated from ∆G}‡ ₌

∆H‡_{− T ∆S}‡ _{where T = 298 K.}_{. . . 84}

2.12 activation parameters for the Diels-Alder reaction between acry-lonitrile and cyclopentadiene in various solvents in the tem-perature range 288-298 K. ∆G‡ _{was calculated from ∆G}‡ ₌

∆H‡_{− T ∆S}‡ _{where T = 298 K.}_{. . . 85}

2.13 Relevant quantities from the theoretical calculations with [Hbim]+

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xii LIST OF TABLES

2.14 Relevant quantities from the theoretical calculations with [bmim]+

cation. . . 91 2.15 Relevant quantities from the theoretical calculations with [bm2im]

+

cation. . . 92

2.16 Relevant quantities from the theoretical calculations without cation. . . 93 3.1 Name and melting point of the investigated substituted

pyri-dinium bromides taking from the Beilstein database. . . 123 3.2 Cross validation. Training set B+C (84 molecules). Results

of the 16 trials. (threshold=50 K). For the meaning of the id number refer to Table 3.1. . . 126 3.3 Cross validation. Test set A (42 molecules). Results of the 16

trials. (threshold=50 K). For the meaning of the id number refer to Table 3.1. . . 128

3.4 Cross validation. Training set A+C (84 molecules). Results of the 16 trials. (threshold=50 K). For the meaning of the id number refer to Table 3.1. . . 131 3.5 Cross validation. Test set B (42 molecules). Results of the 16

trials. (threshold=50 K). For the meaning of the id number refer to Table 3.1. . . 133 3.6 Cross validation. Training set A+B (84 molecules). Results

of the 16 trials. (threshold=50 K). For the meaning of the id number refer to Table 3.1. . . 136 3.7 Cross validation. Test set C (42 molecules). Results of the 16

trials. (threshold=50 K). For the meaning of the id number refer to Table 3.1. . . 138 3.8 Single validation. Training set (100 molecules). Results of the

16 trials. (threshold=50 K). For the meaning of the id number refer to Table 3.1. . . 142

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Index xiii

3.9 Single validation. Test set (26 molecules). Results of the 16 trials. (threshold=50 K). For the meaning of the id number refer to Table 3.1. . . 143 3.10 Comparison of mean absolute errors (MAE), correlation

coef-ficients (R) and standard deviations (RMS) recorded in our experiments with literature data for the training sets. The cor-risponding results for the test sets are reported in Table 3.11. . 145 3.11 Comparison of mean absolute errors (MAE), correlation

coef-ficients (R) and standard deviations (RMS) recorded in our experiments with literature data for the test sets.. . . 146 4.1 Atom Types and Partial Atom Charges of [bmim]+

, [PF6]−, and

[BF4]

−_{. See Figure 4.1 for the atom name correspondence.} _{. . . 153}

4.2 [PF6]

−_{and [BF} 4]

−_{vdW Parameters. Interactions between unlike}

atom types were calculated using the geometric combination rule.154 4.3 Simulated and experimental values of density (ρ), viscosity (η),

isothermal compressibility (κ) for [bmim][PF6] at 298.15 K and 1

atm. . . 156

4.4 Simulated and experimental values of density (ρ), viscosity (η), isothermal compressibility (κ) for [bmim][BF4] at 298.15 K and 1

atm. . . 156 4.5 Cation-anion radial distribution function (rdf ) maxima and

min-ima for [bmim][PF6] (Figure 4.5). . . 163

4.6 Cation-anion radial distribution function maxima and minima for [bmim][BF4] (Figure 4.6). . . 163

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Introduction

Ionic liquids (ILs) have been known since 1914, when ethylammonium ni-trate, [EtNH₃][NO₃], with a melting point of 12◦_{C was discovered by Walden}

[1]. This seems to have created little interest at the time. Indeed, for the next forty years, there was no activity in this field, until other low melting ionic compounds were discovered by chemists looking for an alternative method to electrodeposition of aluminium. On mixing and heating various alkylpyri-dinium chlorides with aluminium chloride, the white powder reacted together to form a colourless liquids which resembled water [2]. However ionic liquids remained something of a curiosity until only recently, when they were redis-covered as alternatives to common organic solvents for synthetic applications. A major breakthrough in the chemistry of ionic liquids was realized in 1992, with the report from Wilke and Zaworotko of a series of air-stable imida-zolium salts with anions such as tetrafluoroborate and hexafluorophosphate [3]. In contrast to the chloroaluminate salts, these ionic liquids were tolerant to a variety of polar functional groups and opened the door to a wider range of chemical and electrochemical applications.

Nowadays, the term ionic liquids has replaced the older phrase molten salts (or melts), which suggests that they are high-temperature, corrosive, viscous media (like molten minerals). The reality is that ionic liquids can be liquid at temperatures as low as -96◦_{C. Furthermore, room-temperature}

ionic liquids are frequently colourless, fluid, and easy to handle. Therefore, in the patent and academic literature, the term ionic liquids now refers to liquids composed entirely of ions that are fluid around or below 100◦_{C. At}

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xvi Introduction

temperature above their melting points, useful organic salts afford stable liquids suitable for a variety of chemical applications where the ionic liquid acts as a solvent, an electrolyte or a reagent for chemical transformations. There are now many ILs based on organic cations that can be coupled to a wide range of anions to provide solvents with specific chemical and physical properties. The cations are generally bulk organic structures with low sym-metry. The most common employed until now are based on ammonium, sul-fonium, phosphonium, imidazolium, pyridinium, picolinium, pyrrolidinium, thiazolium, oxazolium and pyrazolium cations differently substituted (Figure 1).

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Introduction xvii

According to the identity of the anions, ILs may be divided into 4 groups: - systems based on AlCl3and organic salts like 1-butyl-3-methylimidazolium

chloride, [bmim][Cl], whose Lewis acidity can be varied by changing the relative amounts of organic salt/AlCl₃. With a molar excess of AlCl₃ these ILs are Lewis acids while with an excess of organic salt they are Lewis bases. Lewis neutral liquids contain equimolar amounts of organic salt and AlCl3. These ILs are extremely hygroscopic and

han-dling is only possible under dry atmosphere. They are often referred as first generation ILs.

- systems based on anions like PF−6, BF −

4 and SbF −

6. They are nearly

neutral and air stable although a slow hydrolysis of PF−6, leading to

detectable amounts of HF, can be observed. - systems based on anions like CF₃SO−

3 and (CF3SO2)2N− ([Tf2N]−)

are much more stable towards such reactions and generally are char-acterized by the lowest melting points, the lowest viscosities and the highest conductivities of currently available ILs. However, these com-pounds present some drawbacks, generally associated with a high price (in particular those having [Tf₂N]− _{as anion) and the presence of}

flu-orine, which makes more complicated the disposal of spent of these ILs. In addition they may contain traces of halides (chlorides and bro-mides) arising from the preparation procedure. For these reasons, the research of new ILs bearing inert low coordinating and no-fluorinated anions represents a field of intense investigation in the chemistry of ILs. Among the possible alternatives recently proposed must be cited the ILs having as counter anions carboranes ([CB₁₁H₁₂]−_{, [CB}

11H6Cl6]−,

[CB₁₁H₆Br₆]−_{) [4] and orthoborates [5], and dicyanoamide ([N(CN)} 2]−).

- systems based on anions like alkylsulfates [6] and alkylsulfonates [7], which are relatively cheap, do not contain fluorine atoms and could be prepared in a high pure form and are characterized by a wide electro-chemical window and air stability.

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xviii Introduction

Ionic liquids based on 1,3-dialkylimidazolium cations have been investi-gated in detail, partly due to their wide use as solvents in synthesis as well as in catalysis [8]. The attraction of the imidazolium cation in synthetic ap-plications is mainly because the two substituents can be varied to modify the properties of the solvent. 1-Butyl-3-methyl and 1-ethyl-3-methyl imidazolium cations ([bmim] and [emim]) are probably the more investigated structures of this class, although the alkyl chain may be properly functionalized

Ionic liquids are readily prepared from commercial available reagents [9]. The initial step in the synthesis of ionic liquids is the quaternization of an amine or phosphane to form the cation. The alkylation process shows the advantages that a wide range of cheap haloalkanes are available and substitutions generally occur smoothly at reasonably temperature. Any halide but fluoride can be formed by this way. Quaternizations between 1-alkylimidazoles, trialkylamines or triphenilphosphines with methyl triflate, methyl tosilates, octyl tosylate or dimethylsulphate have also been used for the direct preparation of ionic liquids: in principle any alkyl compound con-taining a good leaving group may be used in this way. The excellent leaving group ability of the triflate and tosylate anions means that the direct quat-ernization can generally be carried out at ambient temperature. Of course, this approach has the major advantage of generating the desired ionic liquid with no side products. In all cases, after reaction is complete it is necessary to remove all excess of solvent and starting material by heating the salt under vacuum. Care should be taken during this step as overheating can result in a reversed Menschutkin, especially when halide salts are involved.

In the case in which it is not possible to form the desired anion directly by the quaternization reaction, a further step follows. For example, starting from an imidazolium halide, two different paths to obtain the anion are possible. In the first one, the halide can be treated with a Lewis acid MXy

to form an ionic liquid of the type [R′

Rim][MX_y+1] (Figure 2). Alternatively it is possible to exchange the halide ion X− _{with a proper one by three}

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Introduction xix

Figure 2: Synthesis path for the preparation of ionic liquids examplified for an imidazolium salt.

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xx Introduction

alternative way:

- by the addition of a metal salt M+_[A]− _{followed by removing of the}

by-product M+_X−_{. In this case silver salts are generally used to precipitate}

the halide salt. Cheaper synthetic steps involve the use of sodium salts in dichloromethane or acetonitrile and take advantage of the poor solubility of M+_X− _{in the reaction medium as the driving force of the}

overall process.

- by the use of an ion exchanger resin, which provide a highly pure IL, if the exchange reaction has proceeded to completion.

- by displacement of the halide with a strong acid. This step is generally used to obtain hydrophobic ionic liquids and needs to remove the by-product HX with water extraction.

One of the primary driving forces behind research into ionic liquids is the perceived benefit of substituting traditional industrial solvents, most of which are volatile organic compounds (VOCs), with nonvolatile ionic liq-uids. Replacement of conventional solvents by ionic liquids would prevent the emission of VOCs, a major source of environmental pollution. Really, ionic liquids are not intrinsically green (some are extremely toxic), but they can be designed to be environmentally benign, with large potential bene-fits for sustainable chemistry [10]. Moreover, ionic liquids have a number of properties that make them suitable media for conducting chemical synthesis and catalysis [11, 12]:

- They dissolve many catalysts, polar organic compounds and gases and even support biocatalysis

- They have favourable thermal stabilities and operate over large ranges. Most melt below room temperature and only start to decompose above 300◦_{C, which gives a temperature range in which to conduct synthesis}

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Introduction xxi

- They may be designed to be immiscible with many organic solvents and water, and both cations and anions can be modified to give specific solubility properties, as well as other physical or chemical properties. This has been used to affect total catalyst recovery in a number of transition metal catalysed reactions.

- They are good solvents for a wide range of both inorganic and organic materials and unusual combinations of reagents can be brought into the same phase.

- Their liquid range allows for large kinetic control, which, coupled with their good solvent properties, allows small reactor volumes to be utilised. - They are often composed of poorly coordinating ions, so they have the

potential to be highly polar non coordinating solvents

- They have no vapour pressure and therefore do not evaporate. This means they do not escape into the environment like volatile organic solvents and it is also easy to remove VOCs from the ionic medium under vacuum or by distillation

Essentially there is no limit to the number of different ionic liquids that can be engineered with specific properties for chemical applications. However a number of problems still need to be overcome before their use becomes widespread:

- Many are difficult to prepare in high pure form and current methods that provide pure ionic liquids are generally expensive. Moreover, a re-cent study by Seddon et al. [13] has shown that the physical properties of ionic liquids (melting points, density, and viscosity) strongly depend on the purity of the ionic liquid. Scale-up could be a problem in some cases.

- The viscosity of ionic liquids is often quite high. In addition to this, impurities can have a marked influence and may increase the viscosity as well.

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xxii Introduction

- Catalysts immobilized in ionic liquids are sometimes leached into the product phase.

- Product recovery is not always easy

Ohno et al. [14] have proposed a new class of ionic liquids obtained by di-rect neutralization of five different imidazoles with nine organic acid, includ-ing tetrafluoroboric, hexafluorofosphoric and bis(trifluoromethanesulfonyl)imide acid, in order to obtain highly pure solvents. The synthesis is very simple and provides salts which have low melting point, low glass transition temperature and high conductivity.

Regarding product recovery from ionic liquids, it is important to note that extraction with water works only for hydrophilic products and with hydrophobic ionic liquids, distillation is not suitable for poorly volatile or thermally labile products, and liquid-liquid extraction using organic solvents results in cross-contamination. Brennecke et al. [15] find that non-volatile organic compounds like naphtalene can be extracted from ionic liquids using supercritical carbon dioxide.

A detailed list, although not exhaustive (due to the increasingly growth of new ILs), of 1,3-dialkylimidazolium salts has been presented in various re-views [16, 17, 18]. In several cases, all that is known about these salts is the method of synthesis and the fact that they are liquid at room temperature; in a few cases, a full range of physical properties are available. The unique set of properties of ionic liquids is attributed to a dual behaviour resulting from their quasi-molecular structure formed by three-dimensional sopramolecular polymeric networks of anions and cations linked mainly by C−H hydrogen bonds through the proton in the 2-position. Knowledge of the intimate na-ture, type and strength of hydrogen-bonding in these sopramolecular assem-blies is therefore fundamental to understand the very unique properties of ionic liquids. A number of X-ray diffractions studies have been conducted to probe the solid-state structure of various 1,3-dialkylimidazolium salts, re-vealing the presence of an extended H-bonded network, although the shortest

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Introduction xxiii

interaction does not always involve the proton in the 2-position [19, 20]. In solution, such assemblies have been demonstrated mainly by NMR [21], elec-trochemical studies [22] and reactivity data [23]. In the gas phase the relative strength of hydrogen-bonding and the structure of supramolecular assemblies in ionic liquids have been determined by ESI-MS experiments [24]. For the ESI-MS experiments conducted during this study, see section 1.2.3.

On the basis of literature data, low symmetry [25], weak intermolecular interactions [26] the presence of H-bonding and a large charge distribution [27] are believed to be responsible for the low melting points of ionic liquids (due to the reduction of the lattice energy of the crystalline form of the salt) [28].

Indeed, the strength of electrostatic attraction is determined by the amount of charge on the ions and how far away from each other they are. It is possible to make ions that have their charge on a central atom and have large groups bonded to them so that nothing else can get close: this greatly enhances the distance between the charges. Moreover, replacing the anion with a bulkier one and delocalizing the charge over many atoms, it is even more difficult for the charges to come together. All these features could be used to ratio-nalize the low melting point of ILs, even if some problems may arise from the uncertainty correlated with literature data (many ionic liquids undergo supercooling and the temperature of the phase transition can considerably vary depending on the rate of cooling or warming). Nevertheless, some gen-eral criteria may be drawn. Indeed, according to data reported in Table 1 the cation with the highest symmetry, [mmim]+_{, has a higher melting point}

than the other less symmetric cations. In addition, as the chain length in-creases beyond an ethyl group, there is little change in the melting point until it reaches 10 carbon atoms and the melting point starts to increase again, presumably because of van der Waals interactions between the alkyl chains. Besides the cation, the nature of the anion also has a major influence on both the physical and the chemical properties of the ionic liquid: this effect

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xxiv Introduction

ionic liquids melting point

◦_C [mmim][BF₄] 103 [emim][BF₄] 6 [bmim][BF4] -81 [hmim][BF4] -82 [omim][BF₄] -78 [C₁₀mim][BF₄] -4 [emim][Cl] 81 [bmim][Cl] 65 [bm₂im][Cl] 187 [bmim][I] -72 [bmim][PF6] -61 [emim][OTf] -9 [bmim][OTf] 16 [em₂im][OTf] 6 [mmim][Tf₂N] 22 [emim][Tf2N] -3 [bmim][Tf2N] -4 [emim][N(CN)₂] -3 [bmim][N(CN)₂] -90 [bmim][TSAC] -1.5 [bmim][BOB] -29.2

Table 1: Effect of cation and anion structure on the melting point of imida-zolium based ionic liquids

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Introduction xxv

is however not easy to predict. The effect that charge delocalization and an irregular shape of the anion has on the melting point is reflected in Table 1 by the high proportion of anions such as halides, dicyanamide ([N(CN)₂]−_),

hex-afluorophosphate and bis{(trifluoromethyl)sulfonyl}imide ([Tf₂N]−_{) among}

the room temperature ionic liquids. Delocalization of the negative charge within the S−N−S backbone of the bis{(trifluoromethyl)sulfonyl}imide an-ion, for example, combined with steric protection provided by the sulfonyl oxygen atoms of the trifluoromethylsulfonyl groups presumably decrease ion-ion interaction-ions within the solid thus facilitating ion-ion mobility [29]. In general the melting points provided by smaller anions are greater that those ob-tained with larger more bulky anions. It has been pointed out that the anion fluorination plays a key role for the determination of ionic liquids physical properties [30]. Indeed, the lack of fluorine atoms in NMes2-containing ILs

with respect to Tf2N-based analogues determines lowering of the melting

points. The significantly lower mass of the anion in the non fluorinated case overcome the effect of an increased hydrogen bonding. Moreover, the differ-ent basicities of the anions can also be expected to play a role in determining the melting point of the species. The HNMes₂ is a weaker acid (pKa = 2.85) [31] than Tf2NH (pKa = 1.7) [32] and thus forms a stronger conjugate

base. This is expected to result in a greater degree of association with the cation. As stated above, the proton in the 2-position of imidazolium based ionic liquids is quite acidic as evidenced both spectroscopically and from its reactivity. Such acidity would suggest that the cation can form hydrogen bonds with the anions as well as with compounds dissolved in ionic liquids.

The presence of specific interactions clearly influences the melting point of an ionic liquid. What is unusual, based on the acidity of the 2-proton, is that when it is replaced by a methyl group, the melting point of the ionic liquid increases rather than decreases. Considering the number of proper-ties that determine the melting point of ionic liquids, finding a correlation between this property and chemical composition is not so easy: only mod-est success was obtained for imidazolium bromide, using computer generated

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xxvi Introduction

molecular descriptors [33]. In order to investigate these aspects, a recent neural networks method was use in this work and is presented in Chapter 3. Strictly related with the melting points are the crystal structures of the ILs and, although the structural organization is surely much less in a liquid than in a crystal, the structural organization of a crystal lattice provides a reasonable starting point for understanding structural features in the liquid phase. The crystal structures of several imidazolium ILs have been reported, including the largely used [PF6]− salts [34]. In [emim][PF6] each anion is

surrounded by six imidazolium cations and each imidazolium by six anions. The close points of contact between the imidazolium cation and anion are the aromatic protons and the nitrogen atoms. The methyl and ethyl groups are oriented in alternating direction to give the most efficient packing while beginning to separate the charged and the neutral portions of the cation. In [C12mim][PF6] this separation is more dramatic and a lamellar bi-layer type

organization can be evidenced. The bulk structural organization consists therefore of two alternating non-polar regions, the salt-like packing of anions and cations heads, and the hydrocarbon-like regions of the alkyl tails.

It is noteworthy that this salt-type packing of ions will result in a non-polar, but polarizable medium. The presence of polar (or polarizable) and non polar domain, has been recognized [35] also in the double-layer crystal structures of 1-butyl-2,3-dimethylimidazolium tetrafluoroborate and hexaflu-orophosphate ([bm2im][BF4] and [bm2im][PF6]), whereas no such domain has

been evidenced in the structure of [bm₂im][SbF₆]. In these three non hydro-gen bonded ILs (no evidences of hydrohydro-gen bonding between counter anions have been obtained for [bm₂im][BF₄], [bm₂im][PF₆] and [bm₂im][SbF₆]) pack-ing is determined by anion size. It has been therefore hypothesized [31] that, when a strong anion cation hydrogen bonding is possible (for example in imidazolium chlorides) this becomes the dominant force determining pack-ing of the ions in the crystal lattice and by far exceeds the contribution of cation-cation repulsive interactions. These turn to be important in the case of the ILs containing the [bm₂im]+ _{cation and small anions ([BF}

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Introduction xxvii

and [PF6]−) which are not hydrogen bonded. Finally, in [bm2im][SbF6] also

the cation-cation repulsive interactions are no longer influential owing the lattice expansion. It is evident from this brief summary that the reasons determining the low melting points of ILs, as well as the nature and entity of the interactions present at the liquid state, are far to be completely under-stood and much more work, both experiment and theoretical, is necessary to obtain a more complete picture. It is important to note that the physical properties of ILs not only depend on their structure, but also on their purity grade, therefore, a correlation between chemical composition and physical properties is not always straightforward.

A variety of experimental techniques have been used to investigate liquid structure including neutron diffraction, X-ray scattering and extended X-ray absorption fine structure (EXAFS) [36, 37, 38, 39], and NMR spectroscopy [21]. The importance of the elucidation of liquid structures is clear because this gives an indication as to which interactions are important within the phase and, therefore, which dominate many chemical and physical properties of the liquid, for example, solvation, density, viscosity, and polarity. Enderby et al. [40] were the pioneers in the study of molten salt/ionic liquid structure and clearly demonstrated that the structure of molten NaCl was dominated by alternating anion and cation interactions. In this case, the local order extended out to three to four anion-cation pairs and the molten salt is a highly structured liquid. Since this time, experimental determination of ionic media has expanded significantly with the examination of a wide range of single-component and two-component salts.

A range of 1,3-dimethylimidazolium ([mmim]+_{) salts have been examined}

using neutron diffraction as analogues for the longer chain length derivatives [41, 42, 43]. These salts are symmetric and were used in the experimental studies in order to simplify the analysis. Although they have higher melting points than the ionic liquids based on asymmetric alkyl chain substitution on the ring nitrogen atoms, they still provide useful generic information about ionic liquids. Figure 3 shows the radial distribution of chloride around a

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cen-xxviii Introduction

tral imidazolium cation in [mmim]Cl [44]. Strong charge ordering was found to be present in this ionic liquid with the anions and cations alternating in the radial distribution function in agreement with that previously reported for alkaline halide molten salts.

Figure 3: Comparison of the partial radial distribution functions for (a) the cation-cation distribution and (b) the cation-anion distribu-tion for the 1,3-dimethylimidazolium chloride, hexafluorophosphate, and bis{(trifluoromethyl)sulfonyl}imide salts. Each radial distribution function is calculated from the center of the imidazolium ring, from the phosphorus atom in the case of [PF6], and from the nitrogen atom in the case of [Tf2N]−.

Similarly, probability radial distribution functions have also been deter-mined for the corresponding ionic liquids with hexafluorophosphate([PF₆]) [42] and bis(trifluoromethyl)sulfonylimide ([Tf₂N]−_{) [43] anions as shown in}

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Introduction xxix

Figure 3. Differences are observed between each of the distributions as a re-sult of the differing anions present. An examination of the cation-cation con-tacts shows that in [mmim]Cl the cations are separated by 0.55 nm, while for the hexafluorophosphate and bis{(trifluoromethyl)sulfonyl}imide analogues, the spacings are 0.63 and 0.70 nm, respectively, that is the cation to cation contacts become larger as the size of the anion is increased, Cl−_{< [PF}

6]− <

[Tf2N]−. Similarly, the anion-cation distances are expanded with anion size:

0.42 nm (Cl−), 0.45 nm ([PF6]−), and 0.52 nm ([Tf2N]−). Although in each

case charge ordering was observed, the anion-cation-anion alternating pattern is less pronounced in the case of [mmim][Tf₂N] than for either the [mmim]Cl or [mmim][PF₆] liquids as shown by the almost coincident position of the second shells of the cations and anions in [mmim][Tf₂N] at ≈ 1.3 nm.

In Figure 4, the spatial probability distribution maps of the anions and cations are shown for each ionic liquid, where it is clear that a gradual change is observed in the space that the anions and cations occupy.

The main observations are that progressing from chloride to hexafluo-rophosphate to bis{(trifluoromethyl)sulfonyl}imide, the anions interact less with the ring hydrogens and the cations and anions start to occupy different positions. These effects are the result of the size and charge distribution on the anion changing. The point-charge-like behavior decreases as the anion size increases and the charge becomes more delocalized. This has the effect of reducing the hydrogen bonding accepting ability of the anion, and thus the interaction with the ring hydrogens reduces. A further consequence of this delocalization of the charge is that ionic bonding in the liquid becomes softer and results in increased overlap of the anions and cations in the radial distribution of [mmim][Tf2N]. Moreover, as the anion size increases, the

an-ions and catan-ions must occupy mutually exclusive positan-ions in order to pack most efficiently, and thus it is only in the chloride liquid that an onion-skin structure may be achieved.

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struc-xxx Introduction

Figure 4: Probability distributions of (a) the anions and (b) the imidazolium cations around an imidazolium cation derived from the EPSR model for liquid [mmim]Cl, [mmim][PF6], and [mmim][Tf2N].

ture. In the case of [mmim]Cl and [mmim][PF₆], a remarkable similarity between the reported crystal structure data and the interactions character-izing the liquid state is found. For example, in the solid-state structure of [mmim]Cl, hydrogen-anion contacts dominate the interactions, each cation interacting with six anions; this structure has been also found in the liquid structure. In the crystal structure, the closest distance between cations is the van der Waals contact distance between two methyl hydrogens in adjacent cation dimers (0.25 nm). This is not associated with an attractive interaction, and the anion-cation interactions control the structure, as expected. How-ever, it is important to stress that this is the shortest cation-cation distance in the liquid structure. In contrast, a correspondence between the crystal structure and liquid structure is not found for [mmim][Tf2N]. This could be

explained by the differences in the conformation of the anion between the solid and liquid states. The bis{(trifluoromethyl)sulfonyl}imide anion can adopt either cis or trans conformation but only the cis form is present in the

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Introduction xxxi

crystal [45].

Computer simulations of ionic liquids.

Although experimental data on ionic liquid structures is limited, a wide range of simulations have been applied to this media originally pioneered by the seminal work of Lynden-Bell and co-workers [46]. Comparison of simulation data with models produced from the neutron and X-ray diffraction techniques presented above show good general agreement.

Theoretical studies, especially with the aid of computer simulations, have also been carried out to help understand the details of ILs properties. Molec-ular dynamics (MD) simulations have been utilized to study the diffusion and viscosity of ILs [47, 48, 49], solubility of ILs [50, 51, 52, 53, 54], confined ILs [55], and thermodynamic and transport properties [56], dynamics [57, 58], surface [59, ?, 60] and bulk [61, 62, 63, 64, 65, 66] structures of ILs.

The large amount of experimental and theoretical effort has led researchers to an era when general theories need be developed to systematically explain and predict the physical properties of various species of ILs, rather than to just study individual ILs separately. A unified understanding will allow the systematic design of ILs, targeting customized applications with designated requirements.

A very interesting feature in ionic liquids is the spatial heterogeneity due to tail aggregation [65, 66] which is independent of the specific types of ions. It has been found both in simulations [48, 61, 62, 46] and in experiments [41, 42] that the local environment around the IL ions is highly anisotropic. The nonpolar tail groups were found to aggregate and form separated tail domains, while the charged head groups of the cations and the anions formed a continuous charged network by retaining their local structures with various lengths of the cationic side chain.

A mechanism was suggested [65, 66] to explain the aggregation phe-nomenon. The electrostatic and van der Waals (vdW) interactions are two

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xxxii Introduction

competitive interactions in the ILs. For the charged groups, the electrostatic interactions dominate the local behavior of the charged groups, while the short-range VDW interactions are only significant and repulsive when these groups come very close. These interactions are so strong that the charged groups roughly retain their local structures and form continuous charged domains in different IL systems with various side-chain lengths. By con-trast, the electrostatic interactions on the nonpolar tail groups are negligible compared with their collective vdW interactions, with the latter driving the tail groups to aggregate. Because the head groups and the tail groups of the cations are connected by chemical bonds, the competition of these two groups leads to an inhomogeneous spatial distribution of ionic liquids, in which the tail groups of the cations tend to aggregate to form isolated tail domains, while the charged groups adjust their global behavior to retain their local structures.

The spatial heterogeneity can be related to the diffusive behavior of ILs. When the alkyl chain is short, the amphiphilic character of the cations is not so apparent. Consequently, the cations distribute almost uniformly in space, with the diffusive behavior of the ions being close to simple isotropic liquids. With a longer alkyl chain, the heterogeneously distributed tail domains result in effective binding energies between cations, so that the ions move in a more hopping-like fashion rather than the free diffusion as in a simple liquid, resulting in a slower diffusion. This explains the experimentally observed diffusion decrease, or equivalently, viscosity increase, with longer alkyl chains [67, 68].

One interesting experimental observation is that the C₂ to C₉ IL sys-tems have a very strong tendency to form glasses. By contrast, the C10

and longer side-chain systems are easier to crystallize. This might be inter-preted by the mechanism described as follows: when the side chain is not very long, the collective long-range electrostatic interactions dominating the charged groups and the collective short-range vdW interactions dominating the nonpolar groups are comparable. The competition between these two

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Introduction xxxiii

interactions having different length scales leads to the self-generated glassy state independent of the cooling rate. When the side chain is long enough, the charged groups are greatly localized, so the global behavior of the system is dominated by the collective shortrange vdW interactions. Therefore, the tendency of selfgeneration of glassy state is attenuated due to the weaker competition between those two interactions having different length scales.

Polarity of ionic liquids.

Most chemical reactions are carried out in solution, therefore, the selec-tion of a proper solvent, suitable for the process under investigaselec-tion, is of paramount importance for the success of this reaction. The key characteris-tics of a solvent are those that determine how it will interact with potential solutes. Although for molecular liquids this is most commonly recorded as the polarity of the pure liquid, expressed through its dielectric constant, the inhability of this scale to provide adequate correlations with many ex-perimental data has lead to the conclusion that solvent polarity cannot be measured by a single macroscopic physical bulk solvent parameter, consider-ing solvents as a nonstructured homogeneus continuum [69, 70]. While the concept of solvent polarity is easily grasped in a qualitative sense, the pro-vision of an operational definition remains elusive. IUPAC defines solvent polarity as the sum of all possible, non-specific and specific, intermolecular interactions between the solute ions or molecules and solvent molecules, ex-cluding such interactions leading to definite chemical alterations of the ions or molecules of the solute. Solvent polarity is therefore much better de-scribed by molecular microscopic empirical solvent parameters derived from suitable solvent-dependent reference processes. It has long been known that UV/Vis absorption spectra of chemical compounds may be influenced by the surrounding medium and that solvents can bring about changes in the posi-tion, intensity and shape of absorpton bands. This phenomenon, generally termed as solvatochromism, is caused by differential solvation of the ground

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xxxiv Introduction

and first excited state of the molecule used. For example, if with increasing solvent polarity, the ground state is better stabilized by solvent than the excited state, negative solvatochromism will results, that is a hypsochromic (or blue) shift will be observed. In applying solvatochromisms, however, one has to remember that no single probe molecule is capable of providing a suitable scale solvent polarity, due to the fact that solute-solvent interactions involve a number of distinct and different intermolecular forces. Moreover, in applying such single parameter solvent scales to rationalize solvent effect on organic reactivity, it is tacitly assumed that the combination of solute-solvent interactions between the reference solute and the solvent is almost the same as with the particular substrate under consideration. Obviously, this is an oversimplification which causes serious limitations of the single-parameter approach to medium effect. On the other hand, studies of solute-solvent interactions by means of solvatochromic probes are both simple and conve-nient, if only the interpretation is carefully considered. Solvatochromic and fluorescent probe compounds have been extensively utilized to determine the polarity of solvents [69]. Altough attemps to measure the dielectric constant were made [71, 72] the solvatochromic approach have been applied also for ionic liquids [16].

Reichardt’s betaine dye.

Probably the most widely used empirical scale of solvent polarity is the ET(30) scale, which is based on the negatively solvatochromic

2,6-diphenyl-4-(2,4,6-triphenyl-N-pyridino)phenolate as probe molecule [73] (Figure 5). ET(30) values are simply defined as the molar electronic transition energies

(in kcal/mol) of the solute, measured in solvents of different polarity at room temperature and normal pressure according to equation 1.

ET(30) = hcνmaxNA (1)

where νmaxis the frequency of the maximum absorbtion at lowest energy, and

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Introduction xxxv

respectively.

This scale has been subsequently revised and normalized by Reichardt and Harbusch-Gornert [74, 75], (equation 2) using water and tetramethylsilane as extreme polar and non polar reference solvents, respectively.

EN T =

ET(solvent) − ET(TMS)

ET(water) − ET(TMS)

(2) When choosing a probe molecule able to characterize ionic liquids, our attention focused on Reichardt’s betaine not only because it is widely ap-plied, and therefore ET(30) values are known for a wide range of molecular

solvents, but also because of its inherent properties which make this zwit-terionic dye particularly appropriate to measure simultaneously a variety of solute/solvent interactions (Figure 5). Indeed, it exhibits:

- a large permanent dipole moment, suitable for the registration of non specific dipole/dipole and dipole/induced dipole interactions between solute and solvent

- a large polarizable aromatic π system, consisting of 44 electrons, suit-able for the registration of non specific dispersion interactions

- with the phenolate oxygen, a strong electron-pair donor (EPD) or bond acceptor (HBA) centre, suitable for specific hydrogen-bond interaction with hydrogen hydrogen-bond donor (HBD) solvents

- a positive delocalised and shielded charge of the pyridinium moiety so that EPD of solvents are practically not registered by this betaine dye. Reichardt’s betaine exhibits one of the largest solvatochromic effect of any known compound: its negatively solvatochromic intramolecular absorp-tion band is hypsochromically shifted by 9730 cm−1 _{(537 nm) on going from}

diphenyl ether (λmax = 810 nm) to water (λmax = 453 nm). Differential

solvation of the highly dipolar zwitterionic electronic ground state and the less dipolar first excited state with increasing solvent polarity, leads to the large negative solvatochromism observed.

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xxxvi Introduction

Figure 5: Reichardt’s betaine dye.

The ET(30) and ETN scales of solvent polarity for ionic liquids are

sum-marized in Table 2, whereas a graphical representation of the EN

T solvent

polarity is given in Figure 6

Solvents like ethylammonium nitrate, dimethylammonium chloride and dimethylammonium methylcarbammate are found in the polarity range of other dipolar HBD solvents such as alchools, whereas fully alkylated ammo-nium salts like tetrahexylammoammo-nium benzoate or [bmpyrr][Tf2N] exhibit a

non HBD behaviour. Simple [bmim]-based ionic liquids show ET(30) values

which fall into a narrow range, comprises between 48.32 and 53.22. Com-parison with molecular solvents suggests that these salts are similar to short chain alcohols, like ethanol and 1-butanol. Alteration of the anion has very little effect on the EN

T values, with the exception of [Tf2N]− and [N(CN)2]−

which seem to be less polar than the others. These values are in agreement with those reported by Charmicael and Seddon using Nile Red as probe [76], even if an enhanced sensitiveness towards anion identity could be envisaged with the latter dye. Indeed, even if the range of values is narrow also when Nile Red is applied, the small variations in polarity seems to be determined

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Introduction xxxvii solvents ET(30) ETN [Et3NH][NO3] 61.61 0.954 [bmim][PF6] 52.52 0.675 [emim][Tf₂N] 52.17 0.662 [bmim][Tf₂N] 51.61 0.645 [hmim][Tf2N] 51.86 0.653 [omim][TF2N] 51.1 0.629 [bm₂im][Tf₂N] 48.39 0.546 [om₂im][Tf₂N] 47.7 0.525 [bmim][BF4] 52.30 0.667 [bm2im][BF4] 49.36 0.576 [C₃CNmim][BF₄] 53.22 0.695 [C₃CNmim][N(CN)₂] 51.88 0.653 [C3CNmim][Tf2N] 52.65 0.677 [HO(CH2)2mim][Tf2N] 60.80 0.954 [CH3O(CH2)2mim][Tf2N] 54.09 0.722 [bmim][MesNAc] 51.75 0.650 [emim][N(CN)₂] 51.72 0.649 [bmim][N(CN)2] 51.10 0.629 [bpy][Tf2N] 51.65 0.646 [bmpyrr][Tf₂N] 48.32 0.544 [hex₄N][C₆H₅COO] 43.98 0.41 [bmim][OTf] 51.95 0.656 water 63.1 1.000 methanol 55.4 0.762 ethanol 51.90 0.650 1-butanol 49.7 0.586 acetonitrile 46.0 0.47 dimethylsulfoxide 45 0.44 acetone 42.2 0.355 dichloromethane 40.7 0.309

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xxxviii Introduction

Figure 6: Solvent polarity scale according to EN

T values.

by the anion in the case of ILs containing short 1-alkyl groups, and by the cation for those containing long 1-alkyl groups. Despite in both cases ionic liquids result similar to low molecular alcohols, a different order of polarity could be detected between the two scales:

- [PF₆]− _{> [BF}

4]− > [OTf]− > [MesNAc]− > [Tf2N]− > [N(CN)2]−

following ET(30)

- [NO2]−> [NO3]− > [BF4]− > [Tf2N]− > [PF6]− following Nile Red

In 1-alkyl-3-methylimidazolium [Tf2N] ionic liquids, despite an

appre-ciable decrease in ET(30) is detected on passing from [emim] to [bmim], a

further increase in the alkyl chain length has virtually no influence on the EN

T values. The introduction of a hydroxyl group, ether group or a nitrile

increase the solvent polarity as determined by the EN

T scale,while a methyl

group at C(2) causes a significant drop in the ET(30). In keeping with the

presumption that changes in ET(30) are dominated by hydrogen-bond of the

solvent, a reduction of ET(30) in [bm2im][Tf2N] reflects modification of the

hydrogen bond acidity of the cation, thus confirming the more time proposed ability of H(2) to form hydrogen bond with a solute [24, 77]. Because of its zwitterionic structure, the solvatochromic properties of Reichardt’s dye is

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Introduction xxxix

indeed strongly affected by the hydrogen-bond acidity of the solvent, since hydrogen-bond-donating solvents stabilize the ground state more than the excited state. Taft and Kamlet calculated that about 2/3 of the shift in the absorbance maximum of Reichardt’s dye could be assigned directly to specific interactions involving the phenoxide oxygen [78]. In effect, the ET(30) scale

is largely, but not exclusively, a measure of the hydrogen-bond acidity of the solvent system. It is therefore not surprising that ET(30) for [bmpyrr][Tf2N]

is practically the same as [bm2im][Tf2N], which in turn is not very far from

the value reported for acetonitrile.

Kamlet-Taft solvatochromic parameters.

The scales presented in the precedent paragraphs are based on the spec-tral data of a single standard probe molecule. They are, therefore, of some-what limited value in the correlation analysis of other solvent-dependent processes, because they respond to a combination of nonspecific and spe-cific solute-solvent interactions, which are typical for the chemical structure of the probe molecule, i.e. its ability to register dispersion, dipole-dipole, hydrogen-bond and other possible intermolecular interactions. The reason for the introduction of multiparameter equations is the observation that solute-solvent interactions, responsible for the solute-solvent influence on equilibria, rates and absorptions are caused by a multitude of nonspecific and specific inter-molecular forces between solute solvent molecules. The chance of developing individual empirical parameters for each of these interactions and combine them into a multiparameter equation depends on the possibility of finding solvatochromic probe molecule able to interact with solvents by only one of the existing solute-solvent interaction mechanisms. Kamlet et al. [79, 80, 81] have developed one of the most ambitious and very successful linear solvation energy relationships (LSER), which in the simplified forms could be written as reported in equation 3.

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xl Introduction

where XY Z is a solvent-dependent solute property and (XY Z)₀is the regres-sion value for a reference solvent system, π∗_{is the solvent dipolarity/polarizability,}

α is the solvent hydrogen bond donating acidity (HBD), β is the solvent hy-drogen bond accepting basicity (HBA) and δ is a polarizability correction term that is 0.0 for non halogenated aliphatic solvents, 0.5 for polyhalo-genated aliphatic solvents and 1.0 for aromatics. The three scales are nor-malized so that π∗ _{value for dimethyl sulfoxide, α value for methanol and β}

value for hexamethylphosphoramide is one.

The solvatochromic parameter π∗ _{measures the ability of a solvent to}

stabilize a neighbouring charge or dipole by virtue of non specific dielectric interactios. Therefore, π∗ _{values represents a blend of dipolarity and}

polariz-ability of the solvent. For selected solvents, i.e. nonpolychlorinated aliphatic solvents, with a single dominant bond dipole moment, π∗ _{values are very}

nearly proportional to the solvents molecular permanent dipole moment. The π∗ _{scale is so named because it is derived from solvent effects on the}

π-π∗ _{absorptions of seven nitroaromatics used as primary probe molecules.}

Within the scope of this thesis, the π∗ _{parameters were estimated by using}

N,N-diethyl-4-nitroaniline, a non-hydrogen bond donor solute. As with this probe hydrogen bond interactions are practically excluded, with equation is possible to calculate the π∗ _value.

νmax = ν0+ sπ∗ (4)

where νmax is the frequency of the absorbance maximum, ν0 = 27.52 kK and

s = -3.128.

The π∗ _{values for the ionic liquids are higher than that for non-aqueous}

molecular solvents and are reported in Table 3.

Although differences between the ILs are small, both the cation and the anion affect the values. In molecular solvents, π∗ _{reports the effect of the}

dipolarity and polarizability of the solvent. However, it should be noted that what is being measured is a property of the solute, namely the differential

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Introduction xli solvents EN T π∗ α β [Et₃NH][NO₃] 0.954 1.12 1.10 0.46 [bmim][PF₆] 0.675 1.015 0.654 0.246 [emim][Tf2N] 0.662 0.977 0.664 0.237 [bmim][Tf2N] 0.645 0.971 0.635 0.248 [hmim][Tf₂N] 0.653 0.971 0.650 0.259 [bmim][SbF₆] 0.629 1.039 0.639 0.146 [bm2im][Tf2N] 0.546 1.002 0.413 0.243 [bmim][BF4] 0.667 0.984 0.665 0.451 [bm₂im][BF₄] 0.576 1.083 0.402 0.363 [C₃CNmim][BF₄] 0.695 1.078 0.638 0.337 [C3CNmim][N(CN)2] 0.653 1.124 0.516 0.508 [C3CNmim][Tf2N] 0.677 1.029 0.646 0.219 [bmim][MesNAc] 0.650 1.027 0.594 0.778 [bmim][N(CN)₂] 0.629 1.129 0.464 0.708 [bmpyrr][Tf2N] 0.544 0.954 0.427 0.252 [bmim][OTf] 0.656 1.006 0.625 0.464 water 1.000 1.33 1.12 0.14 methanol 0.762 0.73 1.05 0.61 acetonitrile 0.47 0.799 0.190 0.370 toluene 0.100 0.532 -0.213 0.077 acetone 0.355 0.704 0.202 0.539 hexane 0.009 -0.12 0.07 0.04 Table 3: EN

T and Kamlet-Taft values for a selection of ionic liquids and

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xlii Introduction

stabilization of the more polar excited state with respect to the ground state of the dye. Hence, π∗ _{is derived from the change in energy of the absorption}

maximum of the dye that is induced by the local electric field generated by the solvent. Therefore, it is no surprise that π∗ _{has been affected by the}

ion-dye interactions now possible in the ionic liquid. All of the [Tf₂N]− _{based ILs}

lie at the low end of the range of values observed. At first sight this could be explained considering the reduction in the strength of coulombic interactions as the charge on the anion becomes delocalised. However, despite this rela-tionship is able to justify the decreasing in π∗ _{on passing from [PF}

6]−, [BF4]−

and [OTf]− _{to [Tf}

2N]−, it is not followed by [N(CN)2]− and [MesNAc]−.

The [bmim]+ _{and [hmim]}+ _{based ionic liquids have the same π}∗_{, but}

lower values than the [emim]+ _{and [bm}

2im]+ with a common anion and the

[bmpyrr][Tf2N] ionic liquid has the lowest value of all. The introduction of a

nitrile group in the alkyl chain determines an increase in π∗ _{for [Tf} 2N]

− _and

[BF₄]− _{derivatives, whereas it has practically no effect on [N(CN)}

2]− salts.

The solvatochromic parameter α is a quantitative, empirical measure of the ability of a bulk solvent to act as a hydrogen-bond donor (HBD) toward a solute. The determination of this value arises from the comparison of solvent induced shifts of the longest wavelength π-π∗ _{absorption band of two similar}

probe molecule, one of which cannot act as hydrogen-bond acceptor toward HBD solvent (N,N-diethyl-4-nitroaniline), whereas the other can (Reichardts betaine), and can be obtained by equation 5

α = ET(30) − 14.6(π

∗_{− 0.23δ) − 30.31}

16.5 (5)

As can be seen from the data reported in Table 3, the values are largely determined by the nature of the cation, but there is also a smaller anion effect. It has long been known that all three imidazolium ring protons are acidic, and it was predicted that hydrogen bonding to solutes would be significant in the absence of hydrogen bond accepting anions. In general [bm₂im]+_salts

have lower a values, with respect to the [bmim]+ _{analogues, reflecting the}

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Introduction xliii

based ILs, [bmpyrr][Tf2N] has a slightly highly value than [bm2im][Tf2N].

Although one might expect this latter relationship to be reverse, it probably takes in account the number of protons that can be involved in the hydrogen bonding to the probe molecule. The value reported for [EtNH₃][NO₃] is also in agreement with this hypothesis. It is noteworthy that the α value is not greatly influenced by the presence of a nitrile group in the alkyl chain: in this case a slight increase could be observed for [N(CN)2]−derivatives, whereas an

opposite trend could be envisaged in [BF4]−-containing salts. On the other

hand, no appreciable variation could be detected in [Tf₂N]−_{-based ILs.}

Focusing on the [bmim]+ _{cation salts, there is a clear anion effect: as}

the anion becomes more basic, the HBD ability of the ionic liquid decrease. Considering that the anion is not able to directly interact with the positive charge on the nitrogen atom of Reichardts dye, due to the steric hindrance exerted by the phenyl rings, another explanation needs to be called for. Wel-ton et al. [82] proposed a competition between the anion and the betaine dye solute for the proton to justify the anion effect on the HBD ability of the ionic liquid. According to his study, the a value of the ionic liquids are controlled by the ability of the liquid to act as a HBD moderated by its HBA ability as expressed in terms of competing equilibria, as reported in Figure 7.

Figure 7: Competition between anion and solute for hydrogen bonding with cation.