May 2, 2015 - 102 IONIC LIQUIDS IN BULK AND AT AN INTERFACE .... Data as a Function of Temperature for Three PyrrolidiniumâBased Ionic Liquids.
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Chapter 5
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Ionic Liquids in Bulk and at an Interface: Self‐Aggregation, Interfacial Tension, and Adsorption MOHAMMAD TARIQ
Instituto de Tecnologia Química e Biológica, Universidade Nova de Lisboa, Oeiras, Portugal
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Department of Chemical Engineering, College of Engineering, Qatar University, Doha, Qatar
KARINA SHIMIZU and JOSÉ N. CANONGIA LOPES
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Instituto de Tecnologia Química e Biológica, Universidade Nova de Lisboa, Oeiras, Portugal Centro de Química Estrutural, Instituto Superior Técnico, Universidade de Lisboa, Lisboa, Portugal
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BENILDE SARAMAGO
Centro de Química Estrutural, Instituto Superior Técnico, Universidade de Lisboa, Lisboa, Portugal
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LUÍS PAULO N. REBELO
Instituto de Tecnologia Química e Biológica, Universidade Nova de Lisboa, Oeiras, Portugal
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5.1 INTRODUCTION
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Ionic liquids (ILs) are a new class of salts with melting point below 100°C, so that they are liquid at or near room temperature. Due to their remarkable properties, like negligible vapor pressure, high thermal stability, nonflammability, wide electrochemical window, enhanced solvent quality, and, above all, easy recycling procedures, they are considered to be potential substitutes for many traditional organic solvents. There has been an explosion of academic research, and their introduction in industrial applications is well underway (Plechkova and Seddon [1]). Much of this interest has centered on their designation as “designer solvents”; that is, they can be fine‐tuned by the independent selection of cations and anions to be the optimum solvent with a desired set of properties for a given application. Despite the recent widespread interest in room‐temperature ILs for numerous applications, especially of interfacial relevance, they have not yet been fully subjected to the molecular‐level scrutiny that other liquid surfaces and interfaces have. For instance, ILs are finding use as gas‐capture media. This process involves collision between the Ionic Liquid-Based Surfactant Science: Formulation, Characterization, and Applications, First Edition. Edited by Bidyut K. Paul and Satya P. Moulik. © 2015 John Wiley & Sons, Inc. Published 2015 by John Wiley & Sons, Inc.
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102 IONIC LIQUIDS IN BULK AND AT AN INTERFACE
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gas‐phase molecules and the liquid surface, which is the first step in establishing equilibrium. Upon collision, gas‐phase molecules can be scattered instantaneously or can be accommodated by the liquid‐phase molecules present at the surface. This elementary process is highly dependent on the chemical nature of the liquid surface (Lovelock [2]). The arrangement of ions at the interface is a topic of fundamental importance. Identification of the ions potentially present at a variety of interfaces and of their arrangement there is often not obvious. At present, there are several descriptions available for the gas–liquid interface. In general, there are three possibilities (Fig. 5.1): the anions are able to approach the gas interface closer, on average, than the cations; cations are closer to the interface than anions; and both ions are evenly distributed at the interface. These limiting conditions will have an impact on the uptake and accommodation of the gas‐phase molecules into the IL (Santos and Baldelli [3]). The variety of ions that compose ILs and their inbuilt complexity will also have an obvious impact on such superficial structure. In a classical system like an electrolyte solution, ions generally avoid the gas– liquid boundary mainly due to the existence of weaker interactions at the surface. Therefore, ions prefer to remain in the bulk. This is the traditional interpretation based on the Gibbs equation and surface tension (ST) measurements. It should be noted that some recent molecular dynamics (MD) simulations and spectroscopy results challenge this interpretation (Jungwirth and Tobias [4] and Ghosal et al. [5]). Most electrolyte anions typically approach closer to the gas boundary than cations. Again, this is interpreted in two ways: many of those (molecular) anions are less strongly solvated than cations by molecular solvents, and/or those anions are more polarizable and thus experience less self‐charge repulsions. In essence, the overall structure and energetics of the ions at the surface are intimately influenced by the solvent, often water (Santos and Baldelli [3]). ILs are salts—substances composed essentially of cations and anions. This fact differentiates them from simple ionic solutions, where ions are dissolved in a molecular medium. The lack of a molecular solvent has a profound impact in the structure of the IL near a surface boundary. This chapter encompasses our most recent efforts in the investigation of such behavior.
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Figure 5.1 Three limiting cases for ions at the gas–liquid boundary: (a) anions closer than cations, (b) cations closer than anions, or (c) mixed cations and anions.
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ILs IN BULK: AGGREGATE FORMATION 103
5.2 ILs IN BULK: AGGREGATE FORMATION
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Diluted aqueous solutions of ILs containing ions with long alkyl side chains ( typically larger than octyl chains) can exhibit self‐aggregation phenomena (Blesic et al. [6]) similar to that of conventional surfactants (Tariq et al. [7], Greaves and Drummond [8], Lee et al. [9], and Ali et al. [10]). In these cases, some of the unique properties of ILs, such as their low melting point or their rather sophisticated solvation behavior, become minor or of no consequence (Angell et al. [11]). Nevertheless, aqueous solutions of ILs have some unique and important characteristics: on the one hand, demixing leads to liquid–liquid equilibria instead of solid precipitation (like in the case of traditional inorganic salts); on the other hand, hydrophilic ILs that are totally miscible with water can form electrolyte solutions ranging in concentration from pure IL to pure water. The mixture of an IL with a molecular compound such as water can exhibit different kinds of ionic association or aggregation over a wide range of compositions (Bernardes et al. [12]). Within this framework, we have selected a few IL systems in order to enhance the present understanding about the aggregation behavior of ILs in aqueous solutions and to compare their behavior with that of traditional surfactants. The studies started with diluted regimes that lead to the existence of either isolated ionic species, ion pairs, or micellar states; later, they were progressively extended to more concentrated IL solutions, where the appearance of other aggregation patterns are expected. The first step was to use isothermal titration calorimetry (ITC), a powerful technique for the quantitative investigation of the aggregation properties of surfactants, to study the temperature‐dependent aggregation behavior of three dialkylpyrrolidinium bromide ILs in aqueous solutions, namely, N‐dodecyl‐N‐methylpyrrolidinium bromide, [C1C12Pyrr]Br; N‐butyl‐N‐dodecyl‐pyrrolidinium bromide, [C4C12Pyrr]Br; and N‐butyl‐ N‐octylpyrrolidinium bromide, [C4C8Pyrr]Br. These systems are structurally very close to tetraalkylammonium salts that are used as traditional ionic surfactants, and the similarities or differences to be encountered can contribute to elucidate the mechanisms and the structures responsible for the aggregation of ILs in aqueous solutions. The values of critical micelle concentration (CMC) and micellization enthalpy, ΔHmic, determined using ITC for the aqueous solutions of [C1C12Pyrr]Br, [C4C12Pyrr]Br, and [C4C8Pyrr]Br at three different temperatures are listed in Table 5.1. For all systems and temperatures, except for [C4C8Pyrr]Br at 323.15 K,
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TABLE 5.1 Critical Micellar Concentration (CMC) and Micellization Enthalpy, ΔHmic, Data as a Function of Temperature for Three Pyrrolidinium‐Based Ionic Liquids Ionic Liquid
T (K)
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ΔHmic (kJ.mol−1)
[C1C12Pyrr]Br
288.15 298.15 323.15 288.15 298.15 323.15 288.15 298.15 323.15
13.6 ± 0.3 15.3 ± 0.5 16.1 ± 0.5 7.2 ± 0.5 6.1 ± 0.4 8.7 ± 0.3 169.3 ± 0.6 144.7 ± 0.8 148 ± 3
2.5 ± 0.3 −1.8 ± 0.4 −11.4 ± 0.6 7.4 ± 0.4 2.6 ± 0.1 −8.8 ± 0.2 10.20 ± 0.06 7.2 ± 0.2 −0.3 ± 0.7
[C 4C12Pyrr]Br [C 4C 8Pyrr]Br
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104 IONIC LIQUIDS IN BULK AND AT AN INTERFACE
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Figure 5.2 Enthalpy change as a function of concentration for titration of the aqueous solution of [C4C8Pyrr]Br (1.6 M) into water at 323.15 K. The data (squares) were fitted in a sigmoidal curve (filled curve), and the CMC was obtained as the zero of the second derivative (dotted curve, multiplied by 10).
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the determination of CMC and ΔHmic from titration curves was not difficult because the aggregation process is complete and heat flow differences before and after the CMC are sharp. With [C4C8Pyrr]Br at 323.15 K, the shape of the titration curve is more complex due to a less abrupt transition in the heat flow trends, ΔHobs, before and after CMC (cf. Fig. 5.2). Here, contrary to [C1C12Pyrr]Br and [C4C12Pyrr]Br, the dilution effect in the premicellar region is much more important than the effect due to demicellization. The slope after passing the CMC reflects the equilibrium between smaller, less structured aggregates and a solution containing a higher concentration of free monomers. Both effects contribute to gradual, rather than steep, changes in the enthalpy in the CMC region. This behavior was attributed to shorter alkyl chains being less hydrophobic, leading to smaller aggregation numbers and to lower cooperativity in the micellization process (Olofsson and Loh [13] and Bijma et al. [14]). By analyzing the calorimetric data, further insights could be gained into the energetics of the self‐organization phenomena. The CMC values for [C1C12Pyrr]Br and [C4C12Pyrr]Br are one order of magnitude lower than those for [C4C8Pyrr]Br, an observation compatible with easier micellization processes—and thus lower CMCs—when longer alkyl chains are present. For instance, it is known that in aqueous solutions of ILs composed of 1‐alkyl‐3‐methylimidazolium cations with chloride or bromide ions, the lower limit of alkyl side chain length leading to the formation of micelles is around C6 or C7 (Blesic et al. [6, 15, 16]). The smaller CMC values of [C4C12Pyrr]Br compared to those of [C1C12Pyrr]Br also indicate a more favorable micellization when two longer chains are present, C4 and C12, because of more favorable interactions between the hydrophobic groups. Table 5.1 also shows typical nonmonotonous dependences of CMC on temperature for [C4C12Pyrr]Br and [C4C8Pyrr]Br, which can be interpreted as the interplay between the two driving forces that concur to micellization in aqueous solutions: enthalpy‐ versus
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ILs IN BULK: AGGREGATE FORMATION 105
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entropy‐driven processes (Attwood and Fluorence [17]). The gradual breakdown of the hydrogen‐bonded network of liquid water as temperature increases (and the resulting increase of the entropy of liquid water) leads to enthalpy‐driven micellization processes at higher temperatures, whereas entropy‐driven processes usually dominate at lower temperatures. The CMC of [C1C12Pyrr]Br increases monotonically with temperature in the range studied, indicating that any CMC minima are outside this range, probably below 288.15 K. This is the IL with lower ΔHmic values, which means that a large contribution from the entropy term is not necessary to achieve micellization (negative ΔGmic values), a situation that occurs at relatively low temperatures. In the temperature range studied, the enthalpy of micellization of [C1C12Pyrr]Br and [C4C12Pyrr]Br is positive at lower temperatures, becoming negative at higher temperatures (Table 5.1), which is consistent with the behavior observed for tetraalkylammonium bromide surfactants (Bashford and Woolley [18]). Higher values of ΔHmic for [C4C12Pyrr]Br in comparison with [C1C12Pyrr]Br indicate stronger interactions between the alkyl chains in the former case, compatible with the lower values of CMC obtained. On the other hand, for [C4C8Pyrr]Br, the enthalpy of micellization decreases and approaches zero around 323.15 K, meaning that both contributions to the enthalpy are equally important: these are the disruption of the structural organization of water molecules around hydrophilic and hydrophobic domains of the IL surfactant and the restoring of the hydrogen‐bond structure of water when the micelle is formed. This can also explain the peculiar shape of the calorimetric titration curve observed in that case. In order to prove that diluted aqueous solutions of IL‐based surfactants have similar characteristics to those of traditional ionic surfactants, we have conducted a set of comparisons and empirical correlations between the properties of solutions of the two classes of surfactants (Tariq et al. [7]). Data concerning the aggregation process are compared for tetraalkylammonium bromide and N,N‐pyrrolidinium bromide surfactants with either one or two long alkyl chains attached to the cationic head group. The four series—tetraalkylammonium or N,N‐pyrrolidinium head groups with either one or two alkyl chains—yield an interesting and broad base for comparison and correlation. We note that both head groups contain a nitrogen atom connected to four aliphatic chains, two of which can “grow” to any desired length, while the two others retain their size (either two methyl groups or a pyrrolidinium ring). The main difference between the tetraalkylammonium bromide and N,N‐pyrrolidinium bromide surfactants is that the former are traditional ionic surfactants (when pure, they melt only above 200°C, near their decomposition temperature), while the latter belong to the IL class, with melting points below or around 100°C when pure. The most assessable data, obtained experimentally for all compounds, are the values of CMC. These are, of course, directly related to the tendency for self‐ aggregation—the lower the CMC, the higher the tendency—and many authors have correlated CMC values (or their logarithms) with the size of the alkyl chains of different surfactants (Zana [19]). In the four series under discussion, such comparisons are made in Figure 5.3, plots a and b, where ln(CMC) are plotted as a function of n, the number of carbons in the longest chain of the surfactant cation (plot a), or n + m, the total number of carbon atoms in the two “variable” chains of the cations (plot b). Both plots show distinct trends for each series, although both tetraalkylammonium and N,N‐dialkylpyrrolidinium cations with a single long chain show almost superimposed trends.
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106 IONIC LIQUIDS IN BULK AND AT AN INTERFACE (a)
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Figure 5.3 Critical micelle concentration (CMC) at 298 K as a function of the number of carbons in the first (n) and second (m) chains of CnTAB (cross), CnCmDAB (x‐cross), [CnC1Pyrr] Br (rhomb), and [CnCmPyrr]Br (square). (a) As a function of n only; (b) as a function of n plus m; (c) as a function of n plus m/2.
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In plot 3a, the surfactants belonging to series with two long side chains show, for a given size of the longest chain, lower CMC values than the analogous compound with just one chain, for instance, CMC (C12C4DAB)
