Validation of pH Standards and Estimation of the Activity Coefficients

Article Cite This: J. Phys. Chem. B 2018, 122, 10593−10599

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Validation of pH Standards and Estimation of the Activity Coefficients of Hydrogen and Chloride Ions in an Ionic Liquid, Ethylammonium Nitrate Ryo Kanzaki,* Hikaru Daiba, Hitoshi Kodamatani, and Takashi Tomiyasu

J. Phys. Chem. B 2018.122:10593-10599. Downloaded from pubs.acs.org by YORK UNIV on 12/10/18. For personal use only.

Department of Earth and Environmental Sciences, Graduate School of Science and Engineering, Kagoshima University, Korimoto, Kagoshima 890-0065, Japan ABSTRACT: We selected and validated the pH values of three standard materials that function in the protic ionic liquid, ethylammonium nitrate (EAN). The pH values of 0.05 mol kg−1 phthalate, oxalate, and phosphate buffers were 4.93 (0.04), 2.12 (0.04), and 7.13 (0.06), respectively (the values in the parentheses denote the standard deviation). Because the pH of EAN ranges from 0 to 10, with a neutral pH of 5, these materials are usable as acidic, basic, or neutral standards. The standard electrode potential of silver−silver chloride in EAN was 127.2 (1.7) mV. The activity coefficients of hydrogen and chloride ions remain equal to unity in EAN of a wide concentration range, which indicates that the effective ionic strength is independent of the solute ion concentration. In addition, the estimated value of the transfer activity coefficient of chloride ion suggests a weaker solvation in EAN compared with water in spite of a ubiquitous cation (C2H5NH3+). These behaviors of ions in EAN can be explained by the unique solvation in the ionic liquid through direct ion−ion electrostatic interactions.

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scale is addressed.42−49 Analogous investigations were performed previously for a number of conventional nonaqueous solvents;50−56 however, these types of studies on ionic liquids are few and provide inadequate information. According to IUPAC, pH is a notational definition on the basis of the activity of hydrogen ion (aH), and only the practical procedure to determine pH is provided by potentiometric measurements. That is, the pH value of the sample solution is determined from the potential gap of the pH indicator electrode, often using a glass electrode, from in the pH-reference solution.57 For use as the pH-reference, the lists of pH values of several standard aqueous solutions57 and nonaqueous solutions of potassium hydrogen phthalate58 are provided. On the other hand, pH (or aH) can be expressed by the specific molality and the activity coefficient of hydrogen ions. These two pH values, so-called electrochemically determined and concentration-based, are confirmed to be consistent.59−62 However, appropriate pH standards have not been established yet in ionic liquids. Therefore, selection of pH standard materials, validation of their pH value, and comparison with the concentration-based value must first be achieved in ionic liquids.

INTRODUCTION Ionic liquids are electrolytes that have melting temperature below 100 °C or around ambient temperature and provide a special reaction field that is solely filled with ions. Its subclass of onium salt, namely, protic ionic liquid (PIL),1−5 is primarily characterized by the cation’s dissociable hydrogen, which makes PILs highly applicable to acid catalyzing reactions,6−9 fuel cells,10,11 and so on. Because it is linked to the acid−base properties of PILs, a number of studies concerned with the hydrogen behavior have been performed. In particular, a parameter ΔpKa, based on the difference in the pKa values between the constituent anion and cation in water, exhibits empirically a possible relationship with the ionic character of PIL.1,12−16 However, ions in neat PIL experience serious perturbations in their acidity and basicity. In addition, the pKa values of the conjugate acids of the anions are often unreliable because they tend to behave as strong acids in water. Therefore, a current ongoing debate discusses the actual definition of ΔpKa and what exactly it expresses.17 On the other hand, PILs made from a very weak acid were found to have a surprising ionic property, where hydrogen ions are expected to rarely dissociate according to their ΔpK a values.18−23 From this point of view, the direct observation based on solvatochromism,24−31 NMR,32,33 titrimetric pKa determination,34−41 and so on, of the actual hydrogen ion activity is a significant target of current research. Recently, scaling of pH in ionic liquids using a commonly applicable pH © 2018 American Chemical Society

Received: September 10, 2018 Revised: October 21, 2018 Published: October 23, 2018 10593

DOI: 10.1021/acs.jpcb.8b08870 J. Phys. Chem. B 2018, 122, 10593−10599

The Journal of Physical Chemistry B

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Ethylammonium nitrate (EAN) is the first described liquid salt in 1914,63 and nowadays is the most archetypical PIL. Its characteristics, such as macroscopic physicochemistry,64−72 liquid structure,73−79 dynamics,80,81 and so on, are the most widely known out of all PILs. In our previous study, we have demonstrated that EAN is capable of maintaining ordinary acid−base equilibria where the ionization profile is completely described as a function of pH using the dissociation constants.43 In addition, we have also demonstrated that the ion-sensitive field-effect transistor (ISFET) electrode acts as a pH indicator in EAN, in which a glass electrode does not work.71 Therefore, developing the pH standard materials will promote the investigation and application of acid−base reactions in EAN. In this study, we selected pH standard materials and validated their pH values in accordance with the orthodox procedure that was applied to the conventional molecular solvents. During potentiometric measurements, we observed that the silver−silver chloride electrode acts as a stable reference electrode, and we observed unique behaviors for the activity coefficients of ions in EAN.

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RESULTS AND DISCUSSION Standard Potential of the Silver Chloride. The silver− silver chloride electrode is one of the best choices as a reference electrode in water. Its standard electrode potential in water is determined in a hydrochloric acid solution using a Harned cell.57,82 HCl dissociates completely in water to form an equal amount of lyonium H3O+ and solvated Cl−. Instead of HCl,82,83 we used the EAN solution containing equimolar TfOH and C2H5NH3Cl because TfOH is supposed to dissociate completely in EAN to form an equal amount of HNO3,71 and C2H5NH3Cl is the solvate of Cl−. Consequently, this solution is considered to be equivalent to HCl solution during electrochemical measurements. The following represents the Harned cell that we used in these experiments Pt(H 2)|EAN(H+, Cl−)|Ag/AgCl

(1)

Figure 1 shows the emf of the cell, E, while titration with TfOH + C2H5NH3Cl solution, which corresponds to 0.1 or

EXPERIMENTAL SECTION

Reagents. EAN was synthesized following previously reported procedures43,69,71 from aqueous solutions of nitric acid and ethylamine (both 70%; Kanto Kagaku, Japan). The water content in the synthesized EAN, determined by the KarlFischer method, was 80−200 ppm. Because a completely neutral EAN is not possible, the excess amount of HNO3 or C2H5NH2 of the synthesized EAN was quantified by separate neutralization titrations, then the same amount of distilled C3H7NH2 or TfOH was added to neutralize. Again, the acid− base balance of this stock EAN was determined by the second neutralization titrations. The excess concentration was less than 1 mM (M = mol dm−3), which was taken into account during the subsequent measurements and analyses. Molarity and molality of EAN solutions were mutually converted using the specific density (1.212), which was determined by a pycnometer. TfOH (>99%, provided in ampoule; Kanto Kagaku), C2H5NH3Cl (>98%; Merck, Germany), sodium chloride (99.98%; Kanto Kagaku), potassium hydrogen phthalate (>99.95%; Kanto Kagaku), sodium dihydrogen phosphate (99%; Kanto Kagaku), disodium hydrogen phosphate (>99%; Nacalai Tesque, Japan), oxalic acid anhydrate (99%; Merck), and disodium oxalate (99.5%; Kanto Kagaku) were used without further purification. Potentiometric Titration. Potentiometric titrations were performed at 298.15 K in a glass vessel equipped with a thermostatic water jacket. The electrolyzed silver on a platinum wire was partially oxidized to silver chloride for use as a silver−silver chloride electrode. We confirmed in separate experiments that the silver−silver chloride electrode performed as expected in aqueous solution. After titrant was added to the solution with an autoburette, equilibrium was achieved within 5−10 min, then electromotive force (emf) was recorded. Hydrogen gas was bubbled in the sample to fill the vessel throughout the titration and discharged into the atmosphere. The atmospheric pressure was given by an aneroid barometer and assumed to be equivalent to the partial pressure of hydrogen, pH2, in the vessel.

Figure 1. emf of the Harned cell. Neat EAN was titrated with 0.1 mol kg−1 (filled symbols) or 0.01 mol kg−1 (blank symbols) TfOH + C2H5NH3Cl solution of EAN. A least-squares regression provides (127.2 ± 1.7) mV − 59.16 log(mHmCl/m°2).

0.01 mol kg−1 HCl solution of EAN. The Nernst equation of this cell is expressed as follows pH RT RT ° (EAN) − E = EAgCl ln(aHaCl) − ln 2 F 2F p° (2) where EAgCl ° (EAN) is the standard silver chloride redox potential in EAN, R = 8.314 J/mol·K, T = 298.15 K, F = 96485 C/mol, and the standard pressure, p°, is 1013 hPa. The effect of the atmospheric pressure on E is negligible, and is hereafter ignored. Using the molalities of H+ and Cl− (mH and mCl, respectively), E is expressed by the following equation m m RT ln 10 2RT ° (EAN) − E = EAgCl log H 2Cl − ln γ± F F m° (3)

where γ± is the mean activity coefficient of H+ and Cl−, given by γ± = γHγCl , and m° = 1 mol kg−1. Unlike the HCl aqueous solution, mH and mCl are not strictly identical because of two experimental reasons, that is, (1) imbalance of HNO3 and C2H5NH2 in synthesized EAN and (2) not completely equimolar TfOH and C2H5NH3Cl in the titrant. As shown in Figure 1, E of log(mHmCl/m°2) falls along a straight line with a slope of 59.73 (0.60) mV (the values in the parentheses denote the standard deviation). This value, which is consistent with RT ln 10/F = 59.16 mV at 298.15 K, indicates that there is an ideal Nernst response of the Harned cell in EAN. In 10594


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The Journal of Physical Chemistry B addition, excellent linearity suggests that the γ± value remains constant throughout the molality range. Therefore, E versus log(mHmCl/m°2) was reanalyzed with the ideal slope (59.16 mV), as shown by the solid line in Figure 1. Finally, we obtained E°AgCl(EAN) of 127.2 (1.7) mV from the intercept. The experimental points shown in Figure 1 span a region that is significantly diluted. In particular, several data points at the right of the graph plot along the line only when autoprotolysis of EAN is taken into account in mH. One can consider that this is an infinitely dilute condition, where γ± should be unity. Consequently, γ± = 1 is retained throughout the concentration range observed. We discuss the activity coefficients for each of the respective ions later. Validation of the pH of Standards. According to the IUPAC recommendations,58 0.05 mol kg−1 of potassium hydrogen phthalate was chosen as the primary pH standard for several nonaqueous solutions. We thus adopted the identical condition for the first pH standard. The emf of the Harned cell (1) was recorded in an EAN solution of 0.05 mol kg−1 potassium hydrogen phthalate with changing the chloride ion molality by titration. The titrant solution (0.1 mol kg−1 C2H5NH3Cl along with 0.05 mol kg−1 potassium hydrogen phthalate) was prepared from the same solution of the titrand in order to keep the analytical concentration of potassium hydrogen phthalate constant throughout the titrations. The emf value, E, in this solution is described using the chloride molality mCl by the following equation. m RT ln 10 RT ln 10 ° (EAN) + E = EAgCl pH − log Cl F F m° RT − ln γCl (4) F As shown by the filled circles in Figure 2, E as a function of −log mCl/m° yields a straight line with a slope of 57.63 (0.38)

intercept at log(mCl/m°) = 0 of 418.9 mV, which yields a pH of 4.931 (0.043). Note that the neutral pH of EAN is around five according to the autoprotolysis constant of EAN.64,71 Therefore, potassium hydrogen phthalate is useful as a pH standard for neutral EAN. We expect that a pair of NaH2PO4 and Na2HPO4 will produce a basic pH buffer according to the pKa value (pKa,2 = 7.6643). In Figure 2, filled squares show E for cell (1) depending on Cl− molality in the EAN solution containing 0.05 mol kg−1 of 1:1 NaH2PO4 and Na2HPO4. The cell yields a linear relationship between E and −log mCl/m° with a slope of 56.16 mV. Following the same procedure mentioned above, we finally observed a pH of 7.135 (0.056). For an acidic buffer, although its pKa was not known at this stage, we selected oxalic acid. In Figure 2, triangles display E for cell (1) depending on Cl− molality in the EAN solution containing 0.05 mol kg−1 of 3:1 H2(ox) and Na2(ox) (ox2− = oxalate). The slope was 57.17, and finally a pH of 2.122 (0.038) was given. These pH values of three buffer solutions are summarized in Table 1. They cover the acidic, neutral, and basic regions of EAN. Table 1. Comparison between Electrochemically Determined and Concentration-Based pHsa

a

Values in parentheses denote the standard deviation.

We here compare these pH values with those based on the hydrogen ion molality. Taking into account the acid dissociation constants of phthalic acid in EAN, that is, pKa,1 = 3.73 and pKa,2 = 5.88,43 a 0.05 mol kg−1 potassium hydrogen phthalate solution yields an mH of 1.19 × 10−5 mol/kg. Because acid dissociation constants are molarity-based values, the molality and molarity of all compounds involved in this calculation are simply converted using the specific density of EAN. Finally, this yields a pH = −log γHmH/m° of 4.92, where we apply γH = 1 as mentioned later. This value agrees with the pH determined from the electrochemical procedure mentioned above within the standard deviation. With regard to the phosphoric acid solution, using pKas,43 we found that −log γHmH/m° is equal to 7.13, which agrees with the pH determined electrochemically. With regard to oxalic acid, the acid dissociation constants in EAN were determined in this work by the separate titrations. The detailed procedure has been described elsewhere.43 Typically, 5 cm3 EAN solution of 10 mM sodium oxalate was titrated by EAN solution of 0.25 M TfOH. In Figure 3, three series of titration data are plotted. By means of least-squares fitting, we obtained pKa,1 and pKa,2 of 1.96 (0.06) and 5.28 (0.02), respectively (for only pKas, values in parentheses denote three standard deviations). The calculation curve using the obtained pKa values is drawn by the solid line in Figure 3, which reproduces excellently the experimental points. Using these pKas, −log γHmH/m° is equal

Figure 2. emf of the Harned cell. EAN solutions of 0.05 mol kg−1 phthalate (circles), phosphate (squares), and oxalate (triangles) buffers were titrated with a C2H5NH3Cl + 0.05 mol kg−1 buffer solution of EAN.

mV. Although the difference from the ideal Nernst response (59.16 mV) was slightly larger than the standard deviation, we could not find out a thermodynamically proper reason to interpret simultaneously the smaller slope and the linearity. Therefore, we assume the ideal Nernst response in this solution. Besides, γCl remained constant throughout the concentration range observed. As a result, similar to the way for obtaining EAgCl ° (EAN), we reanalyzed Figure 2 with the ideal slope, as shown by the solid line. Finally, we obtain an 10595


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Figure 3. Potentiometric titration curve of oxalic acid in EAN. The Na2(ox) solution of EAN is titrated with the TfOH solution of EAN. H means the hydrogen ion in EAN (exists as HNO3). The solid line represents the theoretical curve calculated from finally obtained pKa,1 and pKa,2 values.

to 2.22. This value satisfactorily agrees with the electrochemically determined pH, while the difference is larger than the standard deviation. Overall, these pH buffer solutions associate well with the hydrogen ion concentration. Activity Coefficients of H+ and Cl− in EAN. As previously mentioned, the mean activity coefficient of H+ and Cl−, γ±, remains constant at unity in EAN, whereas γCl does not depend on mCl, as is shown in Figure 2. Therefore, γH = 1 and γCl = 1 are supposed in a wide concentration range in EAN. This is a strange behavior because, in conventional molecular liquids, the activity coefficient of ions depends strongly on their concentration. However, it can be reasonably explained by the nature of the ionic liquids. In molecular solvents, the electrostatic potential generated by an ion decays gradually by the ionic atmosphere of the coexisting ions. On the other hand, in ionic liquids, oppositely charged ions exist ubiquitously around the ion, so that the electrostatic potential is rapidly screened and possibly disappeared at the nearest neighbor by one counterion at a minimum. In other words, the effective ionic strength in the conventional meaning is saturated in ionic liquid and remains constant in spite of the addition of ions. This situation has also been found in C2mimTf2N.84 Hence, it may be a fundamental aspect of ionic liquids. In order to consider the solvation state of H+ and Cl− in EAN, their transfer activity coefficients from water to EAN are estimated from emf of the following cell equipped with the double junction bridge

Figure 4. Potential diagram in water and EAN. Values with an asterisk denote that AgCl was immersed in water.

SHEEAN than the SHE in water (SHEW), which yields the transfer activity coefficient of H+ from water to EAN, log γtrW,H→+EAN , of 0.96. Therefore, HNO3 in EAN is a stronger acid than H3O+ in water. Although a rough estimation, this value is excellently consistent with previous results, suggesting that EAN is an acidic solvent that is 1 pH unit more acidic than water.43 For Cl−, the redox potential of AgCl in EAN is 38 mV → EAN of 0.64. lower than that in water, resulting in a log γtrW,Cl − → EAN → EAN > 1) indicates that Cl− (or γtrW,Cl The positive log γtrW,Cl − − is destabilized in EAN, while the transfer Gibbs energy, → EAN → EAN − ΔGtrW,Cl = RT ln γtrW,Cl , is only less than 4 kJ/mol. The − positive transfer Gibbs energy was also observed in the transfer from water to most of organic solvents;55,85 however, it is significantly smaller than the typical values (13.2 kJ/mol for methanol to 40.3 kJ/mol for dimethylsulfoxide). The relatively → EAN − small ΔGtrW,Cl value might be surprising if Cl− in EAN was assumed to be fully solvated by the oppositely charged ions, C2H5NH3+, through electrostatic interactions. Even though the smaller dielectric constant of EAN80 than water should be taken into account, presumably, coexisting NO3− predominantly destabilize Cl− in EAN because of the repulsive interactions. This competitive solvation contributes negatively to the net solvation Gibbs energy. Such a significant effect of like ion is a distinguishable feature in ionic liquids because, in conventional solvents, like ions rarely make contact with each other and their electrostatic contribution has been almost ignored so far.

Pt(H 2)|sample EAN solution||EANint||NaCl aq|Ag/AgCl (5)

The double bars denote liquid junctions with a 4G glass filter. NaClaq denotes an aqueous solution of 0.1 M NaCl, in which silver−silver chloride is immersed, and EANint is the electrode internal solution, which prevents water from contaminating the EAN sample solution. As previously reported,43 the practical standard potential of the cell is −234 mV. Calculating the activity coefficient of Cl− in NaClaq using a simple assumption (log γ± = −0.5 0.1 ), we obtained the redox potential of AgCl to be 165 mV against the standard hydrogen electrode (SHE) in EAN (SHEEAN). If the liquid junction potential between the EAN internal solution and 0.1 M NaClaq is ignored, this potential is considered at the same level as the redox potential of AgCl in water. Finally, as indicated in Figure 4, this hypothesis results in 57 mV higher

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CONCLUSIONS In this study, we propose three pH standard buffers that function in EAN and validate their pH values. These values cover the pH range of EAN from acidic to basic. In addition, we confirmed that these pH values are consistent with the concentration of hydrogen ion. Hence, this set of pH standards are applicable for pH measurements in EAN solutions. Especially, it facilitates using an ISFET electrode because ISFET is a more convenient pH probe than the hydrogen electrode, although its standard potential is always not readily available. Because EAN is an easily available ionic liquid, the 10596


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(11) Miran, M. S.; Yasuda, T.; Susan, M. A. B. H.; Dokko, K.; Watanabe, M. Binary Protic Ionic Liquid Mixtures as a Proton Conductor: High Fuel Cell Reaction Activity and Facile Proton Transport. J. Phys. Chem. C 2014, 118, 27631−27639. (12) Miran, M. S.; Kinoshita, H.; Yasuda, T.; Susan, M. A. B. H.; Watanabe, M. Physicochemical properties determined by ΔpKa for protic ionic liquids based on an organic super-strong base with various Brønsted acids. Phys. Chem. Chem. Phys. 2012, 14, 5178−5186. (13) Stoimenovski, J.; Izgorodina, E. I.; MacFarlane, D. R. Ionicity and Proton Transfer in Protic Ionic Liquids. Phys. Chem. Chem. Phys. 2010, 12, 10341. (14) Angell, C. A.; Byrne, N.; Belieres, J.-P. Parallel Developments in Aprotic and Protic Ionic Liquids: Physical Chemistry and Applications. Acc. Chem. Res. 2007, 40, 1228−1236. (15) Chen, K.; Wang, Y.; Yao, J.; Li, H. Equilibrium in Protic Ionic Liquids: The Degree of Proton Transfer and Thermodynamic Properties. J. Phys. Chem. B 2018, 122, 309−315. (16) Shen, M.; Zhang, Y.; Chen, K.; Che, S.; Yao, J.; Li, H. Ionicity of Protic Ionic Liquid: Quantitative Measurement by Spectroscopic Methods. J. Phys. Chem. B 2017, 121, 1372−1376. (17) Reid, J. E. S. J.; Agapito, F.; Bernardes, C. E. S.; Martins, F.; Walker, A. J.; Shimizu, S.; da Piedade, M. E. M. Structure−Property Relationships in Protic Ionic Liquids: A Thermochemical Study. Phys. Chem. Chem. Phys. 2017, 19, 19928−19936. (18) Doi, H.; Song, X.; Minofar, B.; Kanzaki, R.; Takamuku, T.; Umebayashi, Y. A New Proton Conductive Liquid with No Ions: Pseudo-Protic Ionic Liquids. Chem.Eur. J. 2013, 19, 11522−11526. (19) Watanabe, H.; Doi, H.; Saito, S.; Sadakane, K.; Fujii, K.; Kanzaki, R.; Kameda, Y.; Umebayashi, Y. Raman Spectroscopic Speciation Analyses and Liquid Structures by High-Energy X-ray Total Scattering and Molecular Dynamics Simulations forNmethylimidazolium-Based Protic Ionic Liquids. Bull. Chem. Soc. Jpn. 2016, 89, 965−972. (20) Watanabe, H.; Doi, H.; Saito, S.; Matsugami, M.; Fujii, K.; Kanzaki, R.; Kameda, Y.; Umebayashi, Y. Hydrogen bond in Imidazolium Based Protic and Aprotic Ionic Liquids. J. Mol. Liq. 2016, 217, 35−42. (21) Orzechowski, K.; Pajdowska, M.; Czarnecki, M.; Kaatze, U. Complexation and Proton Transfer in the Binary System Propionic Acid−Triethylamine Evidence from the Composition Dependencies of Mixture Properties. J. Mol. Liq. 2007, 133, 11−16. (22) Kanzaki, R.; Doi, H.; Song, X.; Hara, S.; Ishiguro, S.-i.; Umebayashi, Y. Acid-Base Property of N-Methylimidazolium-Based Protic Ionic Liquids Depending on Anion. J. Phys. Chem. B 2012, 116, 14146−14152. (23) Vitorino, J.; Bernardes, C. E. S.; da Piedade, M. E. M. A general strategy for the experimental study of the thermochemistry of protic ionic liquids: enthalpy of formation and vaporisation of 1methylimidazolium ethanoate. Phys. Chem. Chem. Phys. 2012, 14, 4440−4446. (24) Thomazeau, C.; Olivier-Bourbigou, H.; Magna, L.; Luts, S.; Gilbert, B. Determination of an Acidic Scale in Room Temperature Ionic Liquids. J. Am. Chem. Soc. 2003, 125, 5264−5265. (25) Robert, T.; Magna, L.; Olivier-Bourbigou, H.; Gilbert, B. A Comparison of the Acidity Levels in Room-Temperature Ionic Liquids. J. Electrochem. Soc. 2009, 156, F115−F121. (26) Shukla, S. K.; Kumar, A. Do Protic Ionic Liquids and Water Display Analogous Behavior in terms of Hammett Acidity Function? Chem. Phys. Lett. 2013, 566, 12−16. (27) Reid, J. E. S. J.; Bernardes, C. E. S.; Agapito, F.; Martins, F.; Shimizu, S.; da Piedade, M. E. M.; Walker, A. J. Structure−property relationships in protic ionic liquids: a study of solvent-solvent and solvent−solute interactions. Phys. Chem. Chem. Phys. 2017, 19, 28133−28138. (28) del Valle, J. C.; Blanco, F. G.; Catalán, J. Empirical Parameters for Solvent Acidity, Basicity, Dipolarity, and Polarizability of the Ionic Liquids [BMIM][BF4] and [BMIM][PF6]. J. Phys. Chem. B 2015, 119, 4683−4692.

result of this study may promote the understanding and application of PILs. For thermodynamics, we observed that the activity coefficients of H+ and Cl− remain constant at unity in EAN. This suggests that the Debye−Hückel treatment is no longer relevant for ionic liquids. In addition, the transfer activity coefficients of single ions, H+ and Cl−, were estimated from the redox potential of silver chloride in EAN. These estimations reveal that EAN destabilizes Cl− to some extent in comparison with water, in spite of the ubiquitously existing C2H5NH3+. Although these trends seem strange in comparison with conventional solvents, they can be explained by the fact that ionic liquids are filled with ions. It is worth noting that there is possibly a negative contribution to the solvation energy caused by the contact of like ions, which can occur exceptionally in ionic liquids. This kind of competition should be taken into account when considering ion solvation in ionic liquids. Further knowledge of ion’s behavior in ionic liquids is required.

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AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. Phone: +81-99-2858106. ORCID

Ryo Kanzaki: 0000-0002-7638-6849 Hitoshi Kodamatani: 0000-0002-0136-2754 Author Contributions

The manuscript was written through contributions of all authors. All authors have given approval to the final version of the manuscript. Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS This work was supported by JSPS KAKENHI grant numbers JP26410157 and JP18H03392. REFERENCES

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