Articles

DOI: 10.1002/cphc.201500148

Aqueous Brønsted–Lowry Chemistry of Ionic Liquid Ions Gordon W. Driver*[a] Ionic liquids have become commonplace materials found in research laboratories the world over, and are increasingly utilised in studies featuring water as co-solvent. It is reported herein that proton activities, aH + , originating from auto-protolysis of H2O molecules, are significantly altered in mixtures with common ionic liquids comprised of Cl¢ , [HSO4]¢ , [CH3SO4]¢ , [CH3COO]¢ , [BF4]¢ , relative to pure water. paH + values, recorded in partially aqueous media as ¢log(aH + ), are observed over a wide range (~ 0–13) as a result of hydrolysis (or acid dissocia-

tion) of liquid salt ions to their associated parent molecules (or conjugate bases). Brønsted–Lowry acid–base character of ionic liquid ions observed is rooted in equilibria known to govern the highly developed aqueous chemistry of classical organic and inorganic salts, as their well-known aqueous pKs dictate. Classical salt behaviour observed for both protic and aprotic ions in the presence of water suggests appropriate attention need be given to relevant chemical systems in order to exploit, or avoid, the nature of the medium formed.

1. Introduction Liquid salts, known in the chemical sciences since the 1800s, have become increasingly popular media for all manner of scientific investigation due to lowered liquid temperatures available with modern materials.[1] This feature has enabled a broader spectrum of scientist to explore their utility through subsequent ease of specialist training formerly required for their preparation, handling and use. The term “ionic liquid”, in use already ~ 80 years,[2] is synonymous with the term “liquid salt”, regardless of temperature, though there is a tendency in the current literature to restrict possible fusion temperatures to those salts which melt “below the boiling point of water”. These interesting materials are frequently investigated for performance enhancement (relative to molecular systems) in their pure state, towards improvement of any given chemically driven event, from space exploration[3] and energy storage,[4] to embalming fluids[5] and as solvents for synthesis of specialty chemicals,[6] but are increasingly found in mixtures with molecular co-solvents, with water regularly being employed.[7] It has long been known that many ions and molecules possess unavoidable, solvent-dependent Brønsted–Lowry (BL) and/or Lewis (L) pKs that are defined via production of corresponding conjugate acids and bases, when favourable thermodynamics authorise a stabilising event. While ionic liquids’ chemistry has seen rapid advancement over an ever increasing broadness of scope, BL acid–base chemistry, highly developed and exploited in the organic and inorganic chemistry communities, has received little attention, and data are largely unavailable for applications employing salt mixtures with molecular liquids as co-solvent. The apparent underdevelopment is surpris[a] Dr. G. W. Driver Department of Chemistry Ume” University KBC-huset, Linnaeus v•g 10, 90187 Ume” (Sweden) E-mail: [email protected] Supporting Information for this article is available on the WWW under http://dx.doi.org/10.1002/cphc.201500148.

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ing as BL acid–base chemistry, in aqueous, partially aqueous and non-aqueous systems, has a ubiquitous presence in most branches of the chemical sciences; it plays a central role in daily life, in scientific education, in chemical research and in countless industrial processes. It is, however, understandable in recognition of the fact that pH is scaled to, and strictly reliable in, dilute aqueous solutions (e.g. to acid–base concentrations of < 1 m). For higher acid–base concentrations, in aqueous, partially aqueous and non-aqueous systems, one may alternatively qualify acidity through measurement of paH + , equivalent to pH through Equation (1):[8] paH þ ¼ ¢logðaH þ Þ ¼ pH¢d

ð1Þ

When the correction factor, œ d, defined in terms of liquid junction potentials and associated activity coefficients, is known to a high degree of accuracy, paH + becomes quantitative in nature. Although recognition of BL acid–base character available in aqueous ionic liquid solutions is critical for greater understanding and control of existing and developing chemistries, to date few investigations have been undertaken with regard to the presence of water and its influence on the chemistry of interest. Protic salts may exhibit the expected native BL acid character in the pure salt, although few literature studies are available that feature investigations of their potency in terms of known acidity functions (e.g. Hammett), due to challenges presented by fully non-aqueous systems.[9] Addition of known BL acids and bases (or their structurally appended moieties) to ionic liquid media have been investigated,[10] but the possibility for BL character arising directly as a result of ion reactivity (aprotic or protic) with water, continues to go unrecognised en masse by the ionic liquids community, except perhaps where systems containing [BF4]¢ and [PF6]¢ are concerned. It is the general case, however, that liquid salts comprised entirely of conjugate BL acids and bases, often with accurately character-

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Articles ised aqueous pK values, possess latent BL acid–base character that one may expect to become active in chemical systems that produce, via chemical transformation, or require, various quantities of water. In this work, the qualitative BL acid–base behaviour of various liquid salt solutions containing water from ~ 35–95 mol % was investigated. The varieties of commonly available hydrolysis and acid dissociation equilibria, which form the very foundation upon which liquid state BL acidity has been developed, as defined by reactions between H2O molecules and ions of opposite charge, are presented in Table 1 below. These reactions, already well-known in classical aqueous salt chemistry,[11] are unavoidably of equal significance for the assorted liquid types.

Table 1. Representative BL and L acid–base equilibria encountered in water, liquid salts and their solutions. C1im Ž 1-methyl-1H-imidazole, [HC1im] + Ž 3-methyl-1H-imidazolium cation, [C4C1im] + Ž 3-butyl-1-methyl1H-imidazolium cation, and [C2C1im] + Ž 3-ethyl-1-methyl-1H-imidazolium cation. C4Cl Ž n-butyl chloride.

water (protic and aprotic) 1) CCl3COOH + H2OÐ[H3O] + + [CCl3COO]¢ (BL acid dissociation) 2) (CH3)3N + H2OÐ[(CH3)3NH] + + HO¢ (BL base hydrolysis) 3) [CO3]2¢ + H2OÐ[HCO3]¢ + HO¢ (conjugate BL base hydrolysis) parent molecules to conjugates (protic and aprotic) 4) HBr + C1imÐ[HC1im] + + Br¢ (BL acid–base neutralisation) 5) C4Cl + C1imÐ[C4C1im] + + Cl¢ (L acid–base neutralisation) salt water solutions (protic and aprotic) 6) [C4C1im] + + [BF4]¢ + 2 H2OÐ[C4C1im] + + H2OH…F + [BF3OH]¢ (anion hydrolysis) 7) [HC1im] + + Cl¢ + H2OÐC1im + Cl¢ + [H3O] + (cation acid dissociation) 8) [HC1im] + + [HBr2]¢ + H2OÐ[HC1im] + + 2Br¢ + [H3O] + (anion acid dissociation) 9) [C2C1im] + + [CH3COO]¢ + H2OÐ[C2C1im] + + CH3COOH + HO¢ (anion hydrolysis) 10) [HC1im] + + [HSO4]¢ + H2OÐ[HC1im] + + SO42¢ + [H3O] + (anion acid dissociation) 11) [C4C1im] + + [CH3OSO3]¢ + H2OÐ[C4C1im] + + [HOSO3]¢ + CH3OH (anion hydrolysis)

2. Results and Discussion We find partially aqueous paH + values, reported as ¢log(aH + ), span more than 13 decades across the various salt solutions investigated, where aH + characterises H2O proton activity, or the effective concentration of protons (i.e. aH + = [H3O + ]fH + ). paH + versus H2O mol % plots (Figure 1) demonstrate, to varying degrees, the induction of BL acid–base chemistry due to ion hydrolysis or acid dissociation reactions that occur when any of [HC1im]Cl, [C4C1im][BF4], [C4C1im][CH3SO4], [C4C1im][HSO4], [C4C1im][CH3COO] and [H2C=C2C1im]Cl are mixed with quantities of water. Such reactivity must be expected generally, in much the same way as for typical organic and inorganic salts, since aqueous pKs for a given salt ion persist regardless of the identity of the oppositely charged counter ion. These findings introduce interesting and unavoidable questions concerning ChemPhysChem 2015, 16, 2432 – 2439

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Figure 1. Plot of ¢log (aH + ) versus mol % water for various liquid-salt solutions. Legend: * [C4C1im][CH3COO], ! [C4C1im][CH3SO4], N [H2C=C2C1im]Cl, & [C4C1im][BF4], ! [C4C1im][BF4]*, * [C4C1im][HSO4], & [HC1im]Cl. Data for [C4C1im][BF4]* from Ref. [14].

possible chemical effects of acidity in wet salts that are as yet not fully addressed in the ionic liquids’ literature. Affected applications of interest are included for discussion below. With partially aqueous compositions of > 50 mol % H2O, expected liquid junction potentials (LJPs), arising at the ceramic junction of the pH electrode, were observed to be minimal and reproducible across multiple measurements in the paH + range ~ 3–11, as is typical for measurements in aqueous solutions.[12] Salt-rich sample solutions (i.e. < 50 mol % H2O) however, which represent increasingly non-aqueous compositions were also observed to include minimal and reproducible LJPs, although it was within these compositions that such effects were expected to become most pronounced. That is to say, with increasingly non-aqueous compositions, the aqueousbased glass pH reference electrode was expected to continuously reduce its ability to function as a true reference due to an increasing mismatch of chemical potential between the sample medium and that of the reference electrolyte. Non-constant LJP development at the interface between the aqueous electrolyte and the partially aqueous electrolyte is known to contribute unpredictably to recorded cell electromotive forces (EMFs), and manifest through an expected onset of non-linear electrode response (on the logarithmic scale) across various proton concentrations. Such effects can produce errors in measured cell EMFs of > 40 mV,[13] a sizeable contribution considering that the pH unit spans 59.2 mV for a Nernstian circuit. Other sources of error, including those due to strongly alkaline and acid samples are detailed later in the text. Results of the paH + measurements, which are found to be reproducible and therefore interpretable (i.e. semi-quantitative), in various aqueous solutions of six liquid salts are given in Figure 1, and further detailed in Table 2. Figure 1 and Table 2 present clear, semi-quantitative, evidence ionic liquids investigated exhibit water concentration dependent BL acid–base character. Variation of water concentration was found to induce changes of up to 3 paH + units for aqueous solutions of [HC1im]Cl, [C4C1im][HSO4], [C4C1im] [CH3COO] and [H2C=C2C1im]Cl, across water fractions investigated. Overall, the BL behaviour is typical of aqueous solutions

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Articles Acidities of [C4C1im][HSO4] and [H2C=C2C1im]Cl increase with IL[a] (H2O initial)[b] R[c] mol % H2O[d] [IL][e] [m] paH œ SEM[f] N[g] R2[h] aH water content although the accompanying chemistries of the [C4C1im][HSO4] 15.8 94.0 2.13 0.79 œ 0.01 8 0.162 protic anion [HSO4]¢ and that (0.227 w % H2O) 7.24 87.8 3.17 0.88 œ 0.004 8 0.132 1.69 62.8 4.64 1.22 œ 0.01 8 0.98 0.060 observed for Cl¢ are vastly differ0.92 38.7 4.96 1.39 œ 0.03 8 0.041 ent. The former anion is known 0.54 34.9 5.14 1.49 œ 0.06 6 0.032 for its native non-aqueous [C4C1im][CH3COO] 7.81 88.7 3.05 9.83 œ 0.02 5 1.48 Õ 10¢10 Brønsted acidity[9a] which certain(0.469 w % H2O) 1.77 63.9 4.57 12.09 œ 0.01 6 0.96 8.13 Õ 10¢13 0.57 36.1 5.07 13.27 œ 0.02 9 5.37 Õ 10¢14 ly facilitates a relative ease of [C4C1im][CH3SO4] 8.19 89.2 2.83 8.88 œ 0.02 7 1.32 Õ 10¢9 proton transfer to H2O mole¢10 (0.303 w % H2O) 1.72 63.2 4.23 9.02 œ 0.01 7 0.99 9.55 Õ 10 cules, to yield [SO4]2¢ and 0.54 35.3 4.64 9.15 œ 0.03 7 7.08 Õ 10¢10 ¢6 [H3O] + . The mild decrease of [H2C=C2C1im]Cl 7.06 87.5 3.98 5.07 œ 0.03 8 8.51 Õ 10 (0.493 w % H2O) 1.94 65.9 6.29 5.76 œ 0.01 8 1.74 Õ 10¢6 paH + observed as the system be0.99 0.80 44.5 7.22 6.31 œ 0.03 6 4.90 Õ 10¢7 comes H2O-rich suggests [HSO4]¢ 0.76 43.2 7.26 6.38 œ 0.03 8 4.17 Õ 10¢7 is a stronger acid in the aqueous [HC1im]Cl 8.01 88.9 4.14 0.84 œ 0.01 7 0.145 system, with pKa(aq) < (na)[i] 1.76 68.7 7.76 0.33 œ 0.01 7 0.468 0.94 0.75 42.8 9.04[j] 0.22 œ 0.01 7 0.603 pK (partially aq), quite possibly a 0.55 35.4 9.34[j] 0.06 œ 0.02 7 0.871 a result of increased activity of [k] ¢5 [C4C1im][BF4] 6.02 85.8 3.38 4.78 1 1.66 Õ 10 H2O molecules in H2O-rich com(0.307 w % H2O) 1.86 65.0 4.53 4.74 œ 0.01 5 0.96 1.82 Õ 10¢5 positions, where the ratio of 0.61 38.0 5.04 4.62 œ 0.02 5 2.40 Õ 10¢5 “free” to solvating water mole+ + [a] [C4C1im] Ž 1-butyl-3-methyl-1H-imidazolium cation, [H2C=C2C1im] Ž 1-allyl-3-methyl-1H-imidazolium cules would be expected to incation, [HC1im] + Ž 3-methyl-1H-imidazolium cation, [CH3COO]¢ Ž acetate anion, [CH3SO4]¢ Ž methyl sulfate anion. [b] Initial water content of sample in weight % as the average value of three measurements by KF coulocrease. Acidity levels observed metric titration. [c] R = mole H2O/mole IL. [d] Final water content after correction for initial salt water content. for the two most water-rich solu+ [e] [IL] in water. [f] Mean paH values; pH meter calibrated using 3 standard aqueous buffer solutions at tions are similar to 1 m aqueous pH 4.00, 7.00 and 9.00; S.E.M. Ž standard error of the mean. [g] N Ž number of measurements taken per solutions of NaHSO4, with pH sample. [h] R2 Ž coefficient of determination obtained by linear regression of data points for each salt. [i] na Ž not applicable; this sample was a solid at this temperature. [j] Maximum values as the solutions were saturated. … 1, demonstrating the similarity [k] paH + was initially > 5 but stabilised to this value after ~ 10 min. between aqueous liquid salt and classical salt solution chemistries. In essence, aqueous NaHSO4 of classical salts, where increases in conjugate ion concentraachieves the same level of acidity at 1/2 to 1/3 the concentrations generally serve to increase the acidity or basicity obtion of the analogous liquid salt solutions. Such effects could served. Exceptions were noted for salts where a decrease in also include contributions from accompanying hydration numthe pKa value prevailed on through to water-rich compositions. bers of the [C4C1im] + salt which are expected to greatly Two possible exceptions were noted with the [C4C1im][BF4] exceed those of the Na + analogue, based on volumetric argu+ and [C4C1im][CH3SO4] solutions, with [d ¢log(aH )/d H2O ments.[8] The paH + similarity with pH values for [C4C1im][HSO4] ¢3 mol %] ‹ 10 over the range of ~ 35–90 mol % water, suggessuggests œ d is negligible, as might be expected from the imtive of a buffering capacity. The former salt produces solutions, plied reduction of fH + . in the range ~ 3–5 m [IL], with a ~ threefold reduction in acidity Mild acidity found in the case of [H2C=C2C1im]Cl was surprisrelative to a 1 m aqueous solution of H[BF4], with pH 1.58 (see ing although similar behaviour was reported for [C4C1im]Cl in equilibrium 6 of Table 1). Literature pH values for [C4C1im][BF4] various anhydrous alcohol solutions.[16] In that study, however, have been included in Figure 1 to emphasise the waning influ“apparent” pH values measured were reported as “absolute”, ence of the solute pKa on approach to the limit of 100 mol % without necessary attention directed towards use of the fully H2O, where proton activities tend towards the expected autoaqueous reference electrode (calibrated using aqueous pH protolytic limit of pure water.[14] paH + measurements attemptstandards) with strictly anhydrous samples. This in fact prevents meaningful comparisons between recorded fully noned in [C4C1im][PF6] failed due to this salt’s immiscibility with water, which resulted in an observable miscibility gap. Given aqueous alcoholic pH values and reference values obtained in enough time, however, [C4C1im][PF6] would certainly induce the aqueous system, where errors of up to 3 pH units are posacidity in the water phase due to the well-documented hydrolsible.[12] Additionally, any expectation for complete salt dissoci¢ [15] ysis of [PF6] , yielding HF(aq) and PF5(g). ation, in the low polarity alcohols, would require highly deProtic [HC1im]Cl was by far the most acidic salt observed of tailed assumptions, in contrast to those salts dissolved in those investigated via the expected conjugate acid dissociawater.[17] tion, and provided the highest ionic strength solutions. The With aqueous [H2C=C2C1im]Cl solutions, acidity is unlikely 35.4 mol % H2O solution appears to be more acidic than at a function of impurities directly introduced during salt synthe42.8 mol % H2O, due to differences in water activity; dilution efsis as all starting materials were distilled and dried. In any case, fects prevail at the higher mol % H2O contents, as expected. allyl chloride commonly contains olefins, chloropropanes and Table 2. Brønsted–Lowry character of various liquid salts for [cation][anion]·RH2O at 21 8C. þ

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þ

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Articles chloropropenes;[18] acidic impurities are still less probable with 1-methylimidazole. While the primary active chemical event occurring to provide such quantities of H + remains unclear, an explanation may find roots in the localised hydrolysis theory proposed long ago by Robinson and Harned.[19] Therein it was proposed that lone pairs of H2O interact strongly with the cation, forming a stable hydrogen bonded complex through first hydration shell polarisation effects; positive charges (i.e. protons) are thought to be directed away from the cation, inducing electronic distortion of the participating water molecule. Within such a configuration, a single H2O proton is proposed to interact strongly with a neighbouring anion, with concomitant formation of a cation···H¢O¢H···anion solventshared ion-pair complex, resulting in elongation and subsequent weakening of its bond to oxygen, without occurrence of complete proton displacement. By analogy, acidity recorded for [H2C=C2C1im]Cl-based solutions is rationalised as originating from strong hydrogen bonding between H2O and the proton located on the carbon atom (C2) situated between nitrogen atoms of the amidinium linkage comprising the imidazolium ring. The chloride ion, residing on the periphery, would then interact strongly with the polarised proton of the water molecule to invoke the recorded acidity of a solvent shared ion-pair complex. One structure resulting from this proposed localised hydrolysis between H2O and [H2C=C2C1im]Cl is presented in Figure 2.

Figure 2. Geometry of the proposed cation···H¢O¢H···anion solvent-shared ion-pair complex, optimised at the BP86/KTZVP level with electronic energies calculated at the MP2 level using the 6-311 + G(d,p) basis set. The calculations were performed using GAMESS-US v. 1 May 2013 (R1) employing the COSMO model of Klamt and Eckert (distances in units of ængstrçm).[20] Labels A and B identify protons of H2O, with distances O¢HA = 1.001 æ and O¢HB = 0.974 æ; HA defines the active proton. Vibrational analyses were performed in all cases to ensure optimised geometries were at their true energy minimum.

In a notable study focusing on the state of water in liquidsalt-rich solutions, water molecules were found to exist in a free state, with protons interacting individually with different anions through hydrogen bonding interactions to form symmetric anion···H¢O¢H···anion-type complexes.[21] Therein experimental hydrogen bond enthalpies were reported to increase with the expected relative anion basicity, from [PF6]¢ with DH = ¢7.5 kJ mol¢1 to [NO3]¢ with DH = ¢20.1 kJ mol¢1. The MP2 electronic energy of [H2C=C2C1im] +¢H2O¢Cl¢ calculated in this work is found to be lower than the sum of the paired ChemPhysChem 2015, 16, 2432 – 2439

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ions and individual H2O species by DEMP2 = ¢34.2 kJ mol¢1. Additionally, we obtain DHMP2 = ¢6.5 kJ mol¢1, for the proposed cation···H¢O¢H···anion complex, corrected with MP2 electronic energies, which gives a relatively good agreement with experimental data considering a single geometry was used. In combination, these results suggest stability is gained via the complexation reaction, in both salt-rich and -poor compositions, through hydrogen bonding interactions. This finding, however, conflicts with a previous investigation where the possibility for cation–H2O–anion complex formation was excluded, although the [C2C1im] + salts of interest in that work possessed flexible, large-volume, anions (e.g. bis{(trifluoromethane)-sulfonyl}imide and ethylsulfate) that carry vastly reduced relative charge-tovolume ratios.[22] Another published investigation concluded [C2C1im]Cl, at infinite dilution in water, prefers to form the contact ion pair in a highly idealised Car–Parrinello simulation employing a single cation and anion amongst 60 H2O host molecules.[23] The chloride ion, however, is a known structure-destroying ion; the solvent-shared ion pair may therefore prevail at higher concentrations.[24] Both [C4C1im][CH3COO] and [C4C1im][CH3SO4] become less basic with increased water content with the behaviour of the former salt indicating dilution effects are invoked on the approach to 100 mol % water. It is noted the acetate salt yields paH + = 9.83 in its ~ 3 m solution, quite close to that of a 1 m aqueous Na[CH3COO] solution, with pH 9.37. While the latter salt yields a solution pH similar to the ionic liquid solution, it does so at 1/3 of its ionic strength. Again, this suggests reduced fH + values for protons in the ionic liquid solution, arising from differences in effective water–ion contact interactions. Interestingly, [C2C1im][CH3COO] is reported to absorb large quantities of moisture from the ambient environment, up to 27 w % water, or 78 mol %, indicating the importance of careful salt handling prior to use.[25] Chemical reactivity of aqueous [CH3SO4]¢ introduces the possibility of more complex behaviour, relative to [CH3COO]¢ , where the parent compound, (CH3)2SO4, is known to hydrolyse above 18 8C, with stepwise formation of CH3OSO3H, H2SO4, and CH3OH.[26] Similar hydrolytic instability of the methyl-ester bond of liquid [CH3SO4]¢ salts is also known, with anion hydrolysis yielding equivalents of [HSO4]¢ and CH3OH.[27] paH + values recorded in this work, from 21 measurements over the range 35–90 mol % H2O, however, indicate [C4C1im][CH3SO4] salt solutions yield basic media. This in turn suggests stability of the anion where the ester-acid form has a reported pKa = ¢3.54 endowing its conjugate, [CH3SO4]¢ , with pKb = 17.54.[28] The chemistry of this salt differs considerably in the aqueous environment from its sodium analogue, where a saturated mixture of the latter yields paH + = ¢0.10 œ 0.01 (N = 7), with [Na[CH3SO4]] = 7.46 m, and paH + = 1.33 œ 0.02 (N = 7) at 1.44 m (not saturated). Acidity observed is rationalised as the result of anion decomposition driven by impurity acid remaining after synthesis, being either H2SO4, [HSO4]¢ or a mixture thereof, known to catalyse ester cleavage through the oxygen¢sulfur bond, to produce CH3OH and [HSO4]¢ via an SN2 pathway. The liquid salt, therefore, has a potential to produce acidic aqueous media as well.

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Articles We turn now to the pH measurement itself, which exploits a combination glass membrane electrode that exhibits an electrical potential difference in response to development of a proton concentration differential between the sample, on the outer surface, and the internally enclosed aqueous HCl–KCl buffered reference electrolyte. The potential difference imposed across the glass membrane is expressed between the two electrolytic half cells on either side, and gives rise to an activity-dependent cell EMF, characterised by the Nernst equation, Ecell = E0¢2.303RT/F log(aH + ), to yield aH + of [H3O] + ions available in the sample electrolyte. The cell is constructed to operate using the following arrangement: AgðsÞ jAgClðsÞ jsat: KClðaqÞ ,HClðaqÞ jj1 samplejj2 sat: KClðaqÞ jAgClðsÞ jAgðsÞ

where the barrier labelled j j 1 is characterised by a glass membrane and that labelled j j 2 by a ceramic frit junction that allows contact between the sample and the outer reference electrolyte solution (i.e. KCl(aq) with constant aAg + ), thereby closing the circuit. Activity coefficients fH + and fHO¢ , which are defined by deviation from ideal system behaviour (e.g. fH + = fHO¢ = 1 for pure deionised water, with ionic strength, m = 0) are required for thermodynamic treatment of the real system, to account for repulsive and attractive ion–solute interactions that alter their activities. In an aqueous salt solution, m ¼ 6 0, fH + ¼ 6 fHO¢ ¼ 6 1 and + [H ] therefore becomes an increasingly poor approximation of aH + . While effects of ionic strength for the salt solutions investigated (e.g. m … 2–9 m) are captured by, and included with, aH + , the magnitude of fH + and fHO¢ are not known. Mean salt activity coefficients, f œ = (f + f-)1/2, of the pure liquids, also required for thermodynamic descriptions of the real system, are identified here from a selection of effective ion concentrations available in the literature.[29] Magnitudes are typically in the range of 0.75 > f œ > 0.52 for common ionic liquids comprised of imidazolium cations, in combination with any of [NTf2]¢ , [PF6]¢ , [CF3COO]¢ , and [CF3SO3]¢ . Specifically, [C4C1im][BF4] exhibits f œ = 0.64, with a œ = 3.4 m, in its pure state. Pure salt f œ values are, however, of partial utility only, since these variables are certain to differ greatly in the corresponding aqueous solution. These facts, therefore, preclude estimation of the thermodynamic equilibrium constant, Kreal, which requires knowledge of effective ion concentrations, as given for the following example of [CH3COO]¢(aq) [Eq. (2)]: w½CH3 COO¤¢ þ xH2 OG

K real

¢ HyCH3 COOH þ zHO

ð2Þ

with [Eq. (3)]: Kreal ¼

y z ½CH3 COOH¤y fCH ½HO¢ ¤z fHO ¢ 3 COOH ¢ w w x x ½½CH3 COO¤ ¤ f½CH3 COO¤¢ ½H2 O¤ fH2 O

ð3Þ

where f½CH3 COO¤¢ Ž f œ = (f + f-)1/2, is implied since individual f + and f- values are undefined. ChemPhysChem 2015, 16, 2432 – 2439

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The linear electrode response observed across the wide water content differential for each of the six salts, shown in Figure 1, is intriguing to say the least, especially considering each datum presented is the mean of multiple measurements of each individual, partially aqueous solution, where complete electrode withdrawal was followed by a rinse-and-dry routine between each subsequent measurement. If present, significant error in the recorded cell EMFs, resultant from LJPs (linearly proportional in concentration), manifest in terms of distinguished departures from linearity in plots of ¢log(aH + ) versus mol % H2O (Figure 1), as one feature of linear functions plotted in the semi-logarithmic scale is their appearance as pronounced curves. Large contributions from such “linear” errors would serve to “straighten out” the exponentiality of the data, and thereby induce curvature in the resulting semi-logarithmic plot. We emphasise this was not observed for any of the salt– water systems investigated in this study. Alkaline and acid errors, also linear in concentration, are likewise known to invoke changes to the recorded cell EMF values of up to 0.5 pH units (œ 30 mV) at extreme pH values. In combination, such errors for salt solutions investigated in this work appear to be minimal, estimated here to be of the order of < œ 3.6 mV (a 6 % standard error of the estimate, or ~ < 0.1 paH + unit). The error is based generously on the largest coefficient of non-determination, (1¢R2), obtained from the salts investigated (Table 2) and is typical of analytical work where extreme accuracy is not required.[30] The apparent proportionality shared between ¢log(aH + ) and mol % H2O suggests that acidity–basicity observed in H2O-rich solutions is continuous on approach to the salt-rich ones; acidity–basicity cross-over was not observed. paH + values obtained therefore indicate, in the minimum case, the qualitative BL acid–base character expected for each salt solution. At the same time, this implies that ionic liquid systems containing water content below limits expected for reliable aqueous pH electrometer quantification are in no way exempt from the possibility of latent BL acid–base character expression, as their well-known aqueous pK values would demand. It should be noted that paH + values obtained across different samples of the same salt, synthesised in different laboratories, may vary according to the method of synthesis, the degree of exposure to ambient conditions (e.g. CO2(g) uptake, initial H2O content), sample history and the water quality used for sample preparation, such that specific paH + values reported here are not intended to serve as reference values for any particular salt. Qualitatively, however, these results serve to indicate whether the given aqueous/partially aqueous liquid salt solutions will yield basic or acidic media. Overall these results clearly signify reactivity of typical liquid-salt ions long assumed to be inert in the presence of water. While H2O molecules are in general more highly ionised in the partially aqueous salt systems than in their native pure liquid state, proton activities appear to be reduced relative to those of the corresponding classical aqueous salt solutions. This observation is consistent with an earlier study showing indicator acids to be less dissociated in liquid-salt solvents than in water,[31] due to reduced dielectrics of the pure salt.[32]

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Articles Though classical salt behaviour may be expected for aqueous and partially aqueous liquid-salt solutions, it is worthwhile to consider there are also notable departures from commonly expected BL character of known acids and bases in the pure liquid salt. Ions of anhydrous protic liquid salts may possess drastically shifted pK values, and further differ from the analogous aqueous system in terms of speciation, stability and accompanying chemical equilibria.[33] 3-methyl-1H-imidazolium bromide, [HC1im]Br, produced from equivalents of dry hydrogen bromide (HBr) admitted directly to anhydrous liquid 1-methylimidazole (C1im) (Equilibrium 4, Table 1) is a notable example. Both component molecules possess known BL acid– base character, well exploited in aqueous solution chemistry. In the non-aqueous mixture, however, the pure anhydrous salt that results possesses a non-acidic nitrogenic proton. [HC1im] + so strongly retains its “acidic” proton in the liquid salt, that the high-temperature 1H–15N doublet (1 H) is easily observed, in the 1H NMR spectrum, giving rise to a coupling constant J(1H–15N) = 102 Hz; the associated 15N–1H doublet (1 H) is observed with equal ease in the 15N spectrum. The spectra, shown in Figure 3 following, clearly indicate the non-exchanging nature of the nitrogenic proton, at 105 8C, and therefore attest to the stability of [HC1im] + . This is undoubtedly the result of a larger operative DpK gap between [HC1im] + relative to HBr, in the non-aqueous environment, relative to the analogous aqueous system. It is possible to maintain conditions that preserve cation integrity, although such stability is sensitive to water’s competitive, basic nature, and easily destroyed by it. Aqueous-induced acid dissociation produces [H3O] + and molecular 1-methylimidazole according to the latter’s well-known aqueous pKa ~ 7, even with extremely low water content, according to Equilibrium 7 of Table 1. At the same time, it is clear that liquid [HC1im]Br would serve as a stable solvent for chemistries employing reagents, or yielding products, that possess lower nonaqueous BL pKa values. Such a result in the anhydrous system is to be expected, since the molecular base 1-methylimidazole (C1im) is inherently strong, asserted by its large gas-phase proton affinity (PA), C1im(PA) = 958 kJ mol¢1 (compared with H2O(PA) = 693 kJ mol¢1).[33] Cation formation results in an N¢H + bond with 37 % s character, with 50 % being the maximum value for a pure sp bond. Unfortunately, few reports are available that detail the significance of high temperature, non-aqueous stability of the N¢H + bond observed with high purity salts, which is easily captured in the typical variable temperature 1H NMR experiment through observance of non-binomial coupling between 1H¢14N producing the 1:1:1 proton triplet resonance.[10c] While the possibility for this type of non-aqueous, non-exchanging N¢H + behaviour was observed some time ago, it is, unfortunately, poorly understood and continues to receive little attention in the literature.[34] One unifying aspect common to the many disciplines that now comprise the ionic liquids community is that much of the research effort focuses on their performance as replacement solvents/chemical components in existing and developing chemistries. Mixtures of liquid salts with molecular co-solvents ChemPhysChem 2015, 16, 2432 – 2439

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Figure 3. 1H and 15N NMR spectra of the N¢H + functionality of [HC1im]Br at 105 8C. a) 1H spectrum at 105 8C showing a doublet (1 H) centred at 12.51 ppm, due to spin–spin splitting of the 1H (I = 1=2 ) resonance by 15N (I = 1=2 ) with J(1H-15N) = 102 Hz. The weaker central resonance is due to the magnetic quantum number mI = 0 arising from 14N (I = 1) coupling with 1 H. b) 15N spectrum at 105 8C showing a doublet (1 H) centred at ¢176.65 ppm, due to spin splitting of the 15N resonance by 1H with J(15N¢1H) ~ 100 Hz. Note: The NMR sample tube was filled with neat liquid salt and contained a sealed capillary of [D6]DMSO for field lock.

are now commonplace in varieties of investigations that focus on such key areas as (but not limited to): chemical analysis, polymerisation chemistry, thermodynamics, selective solvent design, electrochemistry, enzyme catalysis, and as reagent/ product for assortments of organic and inorganic bond forming-breaking chemical transformations.[35] In general, recognition of liquid salt ion reactivity in the aqueous milieu is of great value for countless chemical processes. Its absence in the literature suggests a need for a broad-scale, sweeping re-analysis of the relevant chemical systems. Rather than taking a case-by-case approach, we instead provide representative examples that describe the general effect of aqueous induced BL acid–base character in biotransformations and other commonly encountered chemical systems, in Table 3. In cases where BL acid–base participation of the solvent would serve to complement or enhance the chemistry intended, such knowledge can be accounted for, and applied in the design of experiments being considered. In cases where

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Articles Table 3. Implications of aqueous induced BL character of IL ions of various applications employing H2O co-solvent.

Aqueous application

Implications summary

enzyme activation[36]

Enzyme activity becomes a function of IL hydrolysis with buffer capacity greatly exceeded leading to lack of pH control. Assessment of IL performance based on incorrect pH adjustment that may approach pH regime of known enzyme deactivation. High density, non-derivatised cellulose solutions desired for fibre development require water at various stages of the process. Biomaterials are sensitive to acidic media through well-known carbohydrate chemistry. Acidic media formed by ILs in water are to be avoided when derivatisation reactions are undesired. Toxicity of ILs becomes more important as breadth of application increases, especially with water co-solvent. Many bioprobes used in toxicological assessments are pH sensitive. Phase behaviour as a function of polarity and BL character goes unnoticed. Water + salt mixtures are assumed to be stable with effects of additional species due to BL character ignored. MD studies require foreknowledge of species involved in the system of interest. This requires understanding and recognition of solvent (i.e. IL) reactivity in order to probe details of mechanisms involved in the process. Acid catalysed reaction that produces water during the course of the reaction. Salts giving rise to basic aqueous media would be undesirable. Salts providing acidic aqueous media would potentially enhance the reaction and increase product yields.

biomass processing[37]

toxicity assessment thermodynamics[38] MD studies[39] Fischer esterification[40]

such behaviour would serve to impede or complicate the target chemistry, the material can be avoided. The overall question arises, “does acid dissociation or ion hydrolysis in the wet liquid salt enhance or interfere with the intended chemistry”? The answer of course depends on the application, and both situations are possible. In the first instance, one merely need consider classical aqueous system behaviour of a given conjugate ion (e.g. the pKa), as a guide for gauging whether or not BL character would interfere with the chemistry being investigated, since these highly characterised systems express and typify the relevant traits of BL character expected.

ate salt. Most, if not all, ionic liquid ions commonly employed have parent molecules that possess highly developed aqueous Brønsted–Lowry acid–base chemistry through their well-characterised pK values. These liquid salts, comprised solely of oppositely charged conjugate ions, are therefore susceptible to hydrolysis, or acid dissociation (proton transfer) reactions, in chemistries where water is required or becomes available. That is, fundamental changes to the chemical nature of ionic liquid ions, in the presence of water, are to be expected for systems endowed with a natural predisposition for aqueous solution Brønsted–Lowry acid–base chemistry.

3. Conclusions

Experimental Section

The aqueous Brønsted–Lowry acid–base chemistry of six protic or aprotic liquid salt solvents has been reported. A water-concentration-dependent Brønsted–Lowry acid–base character for these common ionic liquids, was observed, with paH + values spanning a range of ~ 0–13. aH + values, determined from paH + measurements, which record ¢log(aH + ) in partially aqueous media, differ from the neutral equilibrium value expected for pure water (i.e. aH + = [H3O + ]fH + ¼ 6 1 Õ 10¢7), by orders of magni+ tude in most cases (e.g. aH … 10¢1 to 10¢14), over a range ~ 35–90 mol % water. Overall, the BL acid–base character found for ionic liquids is concordant with that found for classical organic and inorganic salts, known for their participation in BL acid–base equilibria in the aqueous environment. This result was expected since liquid salt ions possess well-known aqueous pK values. These results additionally indicate water molecules possess a higher degree of effective ionicity in the liquid salt environment relative to that of pure water (i.e. H2O molecules can be highly ionised in ILs). Various implications are realised for applications investigated in such systems, where analysis of the operative chemistry must allow the possibility for natural BL chemistries of ionic liquid ions to manifest with water co-solvent, as their wellknown aqueous pK values demand. This knowledge in turn provides an additional solvent pre-selection criterion, where a poor system may be identified before too much scientific effort has been invested, and replaced with a more appropri-

[C4C1im][BF4] and [C4C1im][PF6], purchased from Acros with stated purities of Š 98 %, were used as received. All other ionic liquids investigated were prepared according to known literature procedures.[10c, 41] Aqueous ionic liquid electrolyte samples were individually prepared by weighing predetermined volumes of fresh 18 MW ultra-pure water into preweighed quantities of the liquid salts, to the desired H2O mol % compositions. Initial trace quantities of water contained in each salt were quantified using the Karl-Fischer coulometric titration (except for [HC1im]Cl, which was used as a powdered solid) and included in the final sample water content. Partially aqueous proton activities, aH + , were recorded using a common glass membrane working electrode (cat. #14002–850 VWR, containing an internal Ag/AgCl reference) combined with a pH electrometer (sympHony SB70P VWR). The pH electrode was rinsed with 18 MW ultra-pure water and wiped dry between measurements. Additionally, aqueous calibration buffers (pH 4.00, 7.00, 9.00) were periodically re-checked to ensure the correct, calibrated pH value was consistently reproduced over the course of measurements. Accordingly, large liquid junction potentials, exposed via observance of buffer pH off-set, were not observed. Uptake of atmospheric CO2(g), available at a partial pressure of ~ 0.4 mbar, by the ionic liquids prior to paH + measurement was of concern due to the possibility of H2CO3 formation with water quantities initially present (~ 0.3–0.5 w % water, c.f. Table 2). To this end, care was taken to minimise sample (and 18 MW ultra-pure water) exposure to ambient CO2(g), although the possibility for significant quantities of the gas to absorb in the samples prior to measurement was negligible since even with a CO2(g) pressure of ~ 1 bar, solubility is limited to mole fractions of < 10¢3 (i.e. CO2 if present, is so at infinitely dilute concentrations).[42]

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Received: February 23, 2015 Published online on June 10, 2015

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