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Copper (II) Chloro Complex Formation Thermodynamics and Structure in Ionic Liquid, 1-Butyl-3-Methylimidazolium Trifluoromethanesulfonate Ryo Kanzaki, Shuma Uchida, Hitoshi Kodamatani, and Takashi Tomiyasu J. Phys. Chem. B, Just Accepted Manuscript • DOI: 10.1021/acs.jpcb.7b06287 • Publication Date (Web): 22 Sep 2017 Downloaded from http://pubs.acs.org on September 24, 2017

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The Journal of Physical Chemistry

Copper (II) Chloro Complex Formation Thermodynamics and Structure in Ionic Liquid, 1Butyl-3-Methylimidazolium Trifluoromethanesulfonate

Ryo Kanzaki,* Shuma Uchida, Hitoshi Kodamatani, Takashi Tomiyasu Department of Earth and Environmental Sciences, Graduate School of Science and Engineering, Kagoshima University, Korimoto, Kagoshima, 890-0065, Japan

ACS Paragon Plus Environment

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Abstract

Metal ions in ionic liquids are laid under an unprecedented reaction field. In order to assess the reaction thermodynamics of metal ions in such a situation, Cu2+-chloro complex formation was examined with spectroscopic and calorimetric titrations in an ionic liquid, 1-buthyl-3methylimidazolium trifluoromethanesulfonate (C4mimTfO).

In addition, the effect of the

structure of the solvated complexes on the complexation mechanism was investigated with the aid of DFT calculations. Chloro complexation successively proceeded and finally provided a [CuCl4]2‒ species, which is also the final product in conventional molecular solvents. Their stability constants were comparable to those in molecular solvents. Interestingly, in spite of the charged solvent in the ionic liquid, the entropy profile of the complexation resembled that in the conventional molecular liquids. This indicates that the entropy gain of the released solvent species from the complexes is the main driving force of the chloro complexation in the ionic liquid. In contrast, unlike the major molecular solvents, the total coordination number of Cu2+ is saturated to 4 in the ionic liquid, and the Cl‒ complexation tends to be accompanied with a 1:1 exchange of the solvent TfO‒ from the complex. In addition, this ligand exchange was almost athermal. This possibly indicates that the coordination number is dominated by the electrostatic hindrance among the ligands including the solvent ions in the primary coordination sphere.


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Introduction Ionic liquids are extraordinary solvents whose cohesive energy arises from electrostatic forces.1-3 Unlike conventional molecular solvents, ions in ionic liquids are laid in the reaction field that is governed by electrostatic interaction.

Metal ions play a key role in many

applications of ionic liquids, such as electrochemical devices, synthesis, extraction, and electrodeposition.4-10 In the use of metal ion reaction, the activity of the metal ions is affected by the solvation state. From this viewpoint, the solvation structure of metal ions in ionic liquids has been widely studied. In particular, lithium ion has attracted significant attention because its molecular-scopic information provides insights into the ionic conductivity of lithium in secondary batteries.11-16

Interestingly, bis(trifluoromethanesulfonyl) amide (Tf2N‒), a major

anion for this kind of ionic liquids, has been reported to undergo conformational isomerization in the solvation sphere of Li+.11-15 Sodium and magnesium are also promising active materials for secondary batteries.16,17 In addition, the solvation structure of lanthanides is important when considering the extraction of radioisotopes using ionic liquids.18-24 Although the information for first transition metals is scarce,25-34 as a whole, the geometric configuration of the primary neighbor atoms around the metal center exhibits minor differences as compared to molecular solvents. With regard to thermodynamics, however, an amazingly small number of research has been performed for experimentally assessing the reactivity of metal ions in ionic liquids.18,19,25,34,35 Lewandowski et al. investigated the transfer Gibbs energy of some metal ions from dimethylsulfoxide (DMSO) to ionic liquids by potentiometric measurements.35 The values range up to approximately ±100 kJ/mol depending on the anion of the ionic liquid, although the values are difficult to be assigned to a single ion.

Considering complex formation, some metal ion


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complexes with anions of low Lewis basicity such as Cl‒ and NO3‒ have been found to form in ionic liquids.25-27 Although chloro complexation occurs in conventional molecular solvents also,36-41 the higher-order chloro complex exists only at very high Cl‒ concentrations in aqueous solutions.42,43 In the past, the complexation thermodynamics of metal ions in the molecular solvents have revealed specific behaviors of the solvents that cannot be explained by continuum consideration for solvents, e.g., steric hindrance of the solvates and the effect of the solvent structure on complexation.44 However, experimental information is too poor to elucidate the metal ion complexation behavior in ionic liquids at the present stage. In ionic liquids, metal ions would be surrounded by the anions of the ionic liquid, like ionic crystals. In contrast, usual anions of ionic liquids such as trifluoromethanesulfonate (TfO‒) have low Lewis basicity;45 therefore, they are considered to be excluded from the metal coordination sphere in the molecular solvents.

Since this situation has not appeared in conventional solvents so far,

experimental information is a first step for understanding the characteristics of ionic liquids as a solvent. In this study, we investigated the formation of Cu2+ chloro complexes in a typical ionic liquid, 1-butyl-3-methylimidazolium trifluoromethanesulfonate, as the model reaction of a metal ion in ionic liquids.

Thermodynamic measurements and density functional theory (DFT)

calculations were carried out to compare the metal ion solvation and complexation mechanism in molecular and ionic solvents.

Experimental Section Ionic liquid 1-butyl-3-methylimidazolium trifluoromethanesulfonate (C4mimTfO) was synthesized via bromide salt (C4mimBr) following the procedure reported in the literature.46 The starting materials (i.e., 1-methylimidazole, ACROS, 99%; bromobutane, Kanto Chemical, 98%;


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and trifluoromethanesulfonyl acid, Kanto Chemical, >99%) were used after distillation once. The product was further washed with distilled water two or three times after passing the AgNO3 test, and was finally dried in a vacuum oven at 60 °C. The water content of the product was checked by Karl Fischer titrator (Kyoto Electronics Manufacturing, MKS-500) and was found to be lower than 200 ppm. Copper (II) trifluoromethanesulfonate, Cu(TfO)2, and 1-butyl-3-methylimidazolium chloride (C4mimCl) were used after drying over P2O5 in vacuo without further purification. Spectroscopic titration was performed using a double beam spectrophotometer (JASCO, V560). The spectra of C4mimTfO containing 0.3-0.8 mM (M = mol dm‒3) Cu(TfO)2 in a 1 cm quartz cell were recorded with distilled water in a Teflon-stoppered cell as a reference. The measurement cells were placed in a cell room under a nitrogen atmosphere and thermostated at 25.0 °C ± 0.1 °C with circulated water. Then, aliquots of a 3 mM C4mimCl solution of C4mimTfO were added up to 1 cm3 using an automatic titrator (Kyoto Electronics Manufacturing, APB-510) directly into the cell. Before recording the spectra, the solution was stirred for over 15 min until stable absorption for each titration point was achieved.

The

absorbance spectrum of pure C4mimTfO was separately recorded as the baseline. Titration was performed three times with different initial Cu(TfO)2 concentrations. Calorimetric titration was performed using an isoperibol-type calorimeter (Tokyo Riko, MPC11) equipped with a uniquely designed titration and measurement system. A weighed amount (typically 2.5 cm3) of C4mimTfO containing 2 mM of Cu(TfO)2 and pure C4mimTfO as a reference were placed in the twin-vessels. The Cu(TfO)2 solution, initially colorless by eyes, turned yellow after the titration. Each vessel, made of polypropylene with a stainless-steel bottom, was set on a thermo-module attached on an aluminum heat bath at a controlled


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temperature of 25.000 °C ± 0.007 °C. Then, aliquots of a 40 mM C4mimCl solution were added using a 1 cm3 micro-syringe with an on-line-controlled automatic syringe pump (HARVARD 11+) placed inside the calorimeter. The solution was stirred by the Tefron-coated propeller. From the voltage signal of the pair of thermo-modules, a theoretical adiabatic curve was drawn assuming Newton’s law of cooling in order to evaluate the heat of reaction at each titration point. The heat capacity of the cell was determined before and after titration from the Joule heat obtained using a co-immersed 100 Ω chip resistor. The heats of reactions ranged from 0.01 to 0.05 J, while the calibration curve showed good linearity (the relative standard deviations were less than 0.5%) up to 0.1 J. The heats in the blank titration were negligibly small, so the heat of dilution was ignored. For the sake of reproducibility, titration was performed four times. DFT calculations for [CuCln(TfO)m](2‒n‒m)+ complexes were performed using the Gaussian 09 program package.47 The computation was mainly carried out using the computer facilities at Research Institute for Information Technology, Kyushu University. Geometry optimization was performed for each complex in an isolated gaseous phase according to B3LYP using a basis set with diffuse and polarization functions, 6-31+G*, followed by a vibrational normal mode analysis. The binding enthalpy was determined by subtracting the formation enthalpies of Cu2+, nCl‒, and mTfO‒ from that of [CuCln(TfO)m](2‒n‒m)+. The basis set superposition error (BSSE) was ignored. In fact, we have confirmed it to be negligible by calculating some clusters using the counterpoise method.


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Results and Discussion Absorbance Spectra: Figure 1(a) shows the experimental absorption spectra of Cu(TfO)2 in C4mimTfO at varying C4mimCl concentrations up to a Cu2+:Cl‒ ratio of 1:6. In these spectra, the absorbance is divided by the Cu(TfO)2 analytical concentration. The addition of C4mimCl caused a monotonic increase in the intensities at 290 and 410 nm to generate absorption maxima, and the intensities at 260 and 370 nm increased up to a certain C4mimCl concentration and decreased thereafter. Observed data were well explained when the formation of four complexes, [CuCln](2‒n)+ (the number of coordinated Cl‒, nCl = 1, 2, 3, and 4), was assumed (see Figure S1) and the absence of nCl = 2 and/or nCl = 3 complex(es) was discarded. The stability constants were determined by optimizing spectroscopic and calorimetric data (described below) simultaneously. The detailed procedure is described in the supporting information (SI). The resulting stability constants, as well as the complexation enthalpies and entropies, are summarized in Table 1. Figure 2(a) shows the stepwise stability constants (log Kn) in C4mimTfO and in some molecular solvents (dimethylformamide; DMF, DMSO, acetonitrile; AN, and propylene carbonate; PC). Also, in these molecular solvents, [CuCl4]2‒ is successively formed as the final product. monotonically.

The log Kn value in C4mimTfO gradually decreases with nCl, but not

According to the values of the stability constant, the activity of Cu2+ in

C4mimTfO is comparable to that in conventional solvents, although it is less sensitive to nCl. The individual absorption spectra of [CuCln](2‒n)+ complexes are shown in Figure 1(b). The spectrum of the nCl = 4 complex was similar to that of [CuCl4]2‒ in DMF (indicated by a dotted line in Figure 1(b)).37 In DMF, [CuCl4]2– is the final product of the chloro complexation with no solvating DMF in the primary coordination sphere. Also, in C4mimTfO, the Cu2+ center of the nCl = 4 complex can be assumed to be placed under the same electronic conditions; that is, Cu2+


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is fully surrounded by four Cl‒ ions without the solvent TfO‒. According to the selection rules, the absorption spectra of nCl = 3 and 4 complexes reveal the presence of Cu2+ with a D2d symmetry.37,48 In the absence of Cl‒, Cu2+ showed very low molar absorptivity in the visible region, which is reminiscent of the Jahn-Teller-distorted Cu2+ in water. According to FT-IR and Raman spectroscopy, the solvation number of Zn2+ in TfO‒-based ionic liquids was estimated to be 3.8 or 4.5 depending on the counter-cation.29 Our DFT calculation revealed a square-planar symmetry of the coordinating oxygen atoms to the Cu2+ center in [Cu(TfO)4]2‒ (see below). Therefore, [Cu(TfO)4]2‒ will be the predominant solvate species of Cu2+ in C4mimTfO (nCl = 0). Interestingly, pseudo-isosbestic points were observed at 280, 320, and 380 nm from nCl = 2 to 4 complexes. This implies that the nCl = 3 complex may exist as an intermediate of nCl = 2 and 4. However, note that the absence of the nCl = 3 complex led to a very poor data fit and was therefore discarded.

Structures of the Complexes: The calorimetric titration curve is exhibited in Figure 3 by the apparent reaction enthalpy, i.e., the experimental heat of reaction at each titration point is divided by the added amount of C4mimCl. The solid line is the theoretical curve calculated by using stability constants and complexation enthalpies in Table 1, indicating that these values were reliably determined. Figure 2(b) shows the stepwise complexation Gibbs energy (∆Gn°, ∆Gn° = ‒RT ln Kn), enthalpy (∆Hn°) and entropy (∆Sn°, in T∆Sn°) as a function of nCl. The significantly large ∆H2° value indicates that [CuCl]+ + Cl‒ → [CuCl2] complexation is energetically unfavorable, whereas the large and positive ∆S2° value promotes this reaction. This trend was reversed in the next step; i.e., the negative ∆H3° and negligible ∆S3° suggest that [CuCl2] + Cl‒ → [CuCl3]‒ complexation is enthalpically driven. As shown in Figure 2(b), ∆Hn° and T∆Sn°


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varies nearly in parallel, thereby resulting in a minor variation in ∆Gn°. This so-called enthalpyentropy compensation effect implies that the coordination structure deeply affects the complexation mechanism of Cu2+. Both of (i) the desolvation of Cl– in the bulk and (ii) the release of the TfO‒ from the solvated chloro complex positively contribute to the complexation entropy. However, the former value would not depend on the complexation step. Therefore, the desolvation process of the complex is concluded to affect entropically the chloro complex formation mechanism in C4mimTfO. Figure 4(a) shows the overall complexation enthalpy (∆βHn°). As shown, the ∆βHn° value gradually decreased to a value of ∆βH4° = –12 kJ/mol, except for ∆βH2°. The ∆βH4° value corresponds to the ligand exchange of four solvating TfO‒ species with the same number of Cl‒, i.e., [Cu(TfO)4]2– + 4Cl‒ → [CuCl4]2– + 4TfO‒. Through a simple calculation, the ion-exchange process is found to be nearly athermal, or slightly exothermic (3 kJ/mol). Conversely, the exceptionally large and positive ∆βH2° value indicates the transient lowering of the total coordination number of the Cu2+ complex in the nCl = 2 complex. In fact, the different shape of the nCl = 2 spectrum compared with the nCl = 1 implies the difference in the coordination geometry between these complexes. Hence, a 3-fold [CuCl2(TfO)]‒ is likely the main form of the nCl = 2 complex. When the nCl =3 complex is formed, the total coordination number becomes 4 again: [CuCl2(TfO)]‒ + Cl‒ → [CuCl3(TfO)]2‒, and the ∆βH3° value returned to near zero. In this reaction, Cl‒ can enter the primary coordination sphere without enthalpic cost and entropic gain for the TfO‒ release. This consistently explains the large and negative ∆H3° and the very low ∆S3° values (Figure 2 (b)). To clarify the solvation structure of the Cu2+ complexes, DFT calculations were performed. The complexation enthalpies for nCl = 0 to 4 complexes and the corresponding geometries of


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their coordinating atoms are summarized in Table S1. Each complex was optimized from 2-8 initial configurations.

The primary neighboring atoms of the most plausible complex for

respective nCl, as discussed below, are extracted in Figure 5. In the optimized geometries of [Cu(TfO)4]2‒ and [CuCl(TfO)3]2‒, the four nearest neighbor atoms were in approximate squareplanar symmetry, where the latter complex showed a slight D2d character (the bond angles made by Cu2+ and the respective trans-position atoms were 157° and 158°). An optimized geometry of [Cu(TfO)3]‒ was also provided. Although the stationary point was not achieved for [Cu(TfO)5]2‒ , the complex did not disaggregate during the geometry optimization process. Conversely, one TfO‒ departed from the [Cu(TfO)6]4‒ complex. These results imply that the solvation number of 4 ± 1 can be allowed for Cu2+ in C4mimTfO. In the case of the nCl = 2 complex, both [CuCl2(TfO)]‒ and [CuCl2(TfO)2]2‒ provided the optimized geometries; however, the formation enthalpy of the former complex was approximately 200 kJ/mol larger than that of the latter. Although it is difficult to determine the real form of the complex based only on the DFT energy in the isolated gaseous phase, the [CuCl2(TfO)]‒ structure of the nCl = 2 complex is consistent with the aforementioned thermodynamic results. For the nCl = 3 and 4 complexes, geometry optimization did not converge for the complexes containing TfO‒. However, since the nCl = 3 spectrum showed a D2d character, the existence of [CuCl3(TfO)]2‒ is assumed. To explain these results, [CuCl3]‒ and [CuCl3(TfO)]2‒ are proposed to coexist as the nCl = 3 complex. In fact, the [CuCl3]‒ geometry appears to have sufficient spatial capability to approach for TfO‒ even though it is not explicitly coordinated to Cu2+ in the solution.

If the spectra of [CuCl3]‒ and

[CuCl3(TfO)]2‒ are supposed to be similar to those of [CuCl2(TfO)]‒ and [CuCl4]2‒, respectively, the spectrum assigned to the nCl = 3 complex becomes a combination of both those spectra, which plausibly explains the presence of isosbestic points. However, secondary and further


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solvation should be taken into account to obtain a more reliable insight. With respect to the gaseous phase calculation, a stable geometry for complexes is not given with a total coordination number of 5 and more. This implies that the total coordination number of the Cu2+ chloro complex in C4mimTfO is restricted to 4 by the electrostatic hindrance of the ligand anions in the coordination sphere, although it might be weakened by the cations in the outer sphere.

Comparison with Other Solvents: In Figure 4, the overall enthalpy and entropy (∆βSn°) values for the chloro complexation of Cu2+ in C4mimTfO are compared with those in some molecular solvents (DMF, DMSO and AN).36-38 The order of ∆βH° in the molecular solvents clearly depends on the Gutmann’s donor number. Namely, the larger the donor number of the solvent (DMSO: 29.8 ≥ DMF: 26.6 > AN: 14.1), the more endothermic is the substitution by Cl‒. As shown in Figure 4(a), the donor number of C4mimTfO likely lied between those of DMF and AN.

Hence, the value estimated using

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Na-NMR (20.5)49 is reasonable, although the

dependence on the solvation energy of Cl‒ should be taken into consideration. Interestingly, as shown in Figure 4 (b), the overall entropy of the complex in C4mimTfO showed a roughly similar behavior to these solvents until ∆βS3°. This implies that the same driving-force works in the chloro complexation. That is, in spite of the charged solvent species, the entropy gain associated with the desolvation of TfO‒ promotes the chloro complexation in the ionic liquid, while the direct Cu2+-TfO‒ interaction has a minor effect. On the other hand, the reaction mechanism shows a difference in the total coordination number. In these molecular solvents, Cu2+ is essentially 6-fold, although the solvation from the axial position is loosened owing to the Jahn-Teller effect.50-52 This causes a permanent decrease in the total coordination number during the formation of the [CuCl4]2‒ complex. Conversely, the total coordination number of the chloro


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complexes of Cu2+ does not exceed 4 in C4mimTfO. As a result, the chloro complexation tends to occur through a 1:1 ligand exchange. In such a case, (i) the solvent TfO‒ is released from the primary coordination sphere and subsequently (ii) inserted into the bulk, accompanied with (iii) the desolvation of Cl‒ in the bulk before (iv) entering the coordination sphere of Cu2+. Although the information of complexation dynamics of Cu2+ is very scarce,53 as far as the initial and final states are considered for each complexation step, the Cu2+ complex maintains an unchanged formal net charge (‒2) except for the nCl = 2 complex. As a result of the compensation for the electrostatic interactions associated with the abovementioned (i) to (iv) steps, the overall enthalpy of the complex is balanced. This could be the reason for an athermal ligand exchange of TfO‒ and Cl‒. The minor increase in ∆G2° in Figure 2 (b), or a drop of in log K2 in Figure 2(a), would be ascribed to the shrink of the net charge of the complex, or the partial compensation between the enthalpic cost and entropic gain of the desolvation. With regard to ionic liquids, the information of the complexation thermodynamics is quite rare. In Figure 2(b), the complexation enthalpies of Ni2+ with NO3‒ in C4mimTf2N, as reported by Melchior et al.,25 are shown. As the final product is [Ni(NO3)3]‒ (nNO3 = 3), the shifted scale along the coordination number is used. Although a direct comparison in this case is difficult, it is worth noting that the trend from ∆H2° to ∆H3° resembles that from ∆H3° to ∆H4° in our case. It is interesting if the framework of ion-exchange complexation can explain generally the reaction mechanism of metal ions in ionic liquids.

Conclusions The chloro complexation constant of Cu2+ in the ionic liquid (C4mimTfO) was found to be at the same level as that in molecular solvents. That is, the Cu2+ cation solvated by the TfO‒ anion


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was neither significantly stabilized by electrostatic interactions nor activated owing to the solvents’ low Lewis basicity. The reason is clarified; as is the case for molecular solvents, the main driving force of chloro complexation in C4mimTfO was the desolvation entropy of the complex, and its contribution to the stability constants is similar between ionic and molecular solvents.

In contrast, a difference in the reaction mechanism was observed.

The total

coordination number of Cu2+ is limited to 4 due to the electrostatic hindrance at the primary coordination sphere. As a result, the complexation proceeds with 1:1 ligand exchange of Cl‒ and TfO‒, except for the nCl = 2 complex. In addition, the ligand exchange was almost athermal. Hence, the charge of the ligands in the primary coordination sphere is indicated to dominate the total coordination number of metal ion in the ionic liquid. This is a characteristic feature in the ionic liquid because the coordination number tends to be rather dominated mainly by the steric hindrance in molecular solvents. In spite of such a difference, the complexation enthalpy value was not inconsistent with the estimated donor number of C4mimTfO. As shown by AlarconEsposito et al.,54 the donor number might be a reliable parameter. As a whole, our results highlighted the similarity between ionic and molecular solvents. However, it is still unclear whether this is a general rule for ionic liquids. As a proof to this query, further experimental data for various ionic liquids are required.


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–1 3 –1

5000 (a)

(b)

4000 4

2+

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Absorption Coefficient of Cu (mol dm cm )


3000

2–

[CuCl4] in DMF

3 1 2

2000

4 1000

3 0

0

2 250 300 350 400 450 250 300 350 400 450 500 wavelength (nm)

Figure 1. (a) Absorption spectra of Cu(TfO)2 in C4mimTfO ionic liquid with changing C4mimCl concentration. The absorption coefficient corresponds to the experimental absorbance divided by Cu(TfO)2 analytical concentration. (b) The finally obtained individual spectra of [CuCln](2‒n)+ complexes in C4mimTfO ionic liquid. The number refers to the coordination number of Cl‒ of the complex (nCl in the text). The dotted line denotes the [CuCl4]2‒ spectrum in dimethylformamide (DMF).37


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PC AN DMF DMSO

12 10 8 6 4 2 0

1

2

3

4

–1

(a)

log Kn

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∆Gn°, ∆Hn°, T∆Sn°, and ∆βGn° in kJ mol

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60 (b)

T∆S°

40 20 ∆H°

0 ∆Hβ°

-20 -40

∆G°

1

-60

3 NO –

2

1

2

3

3

4 Cl–

–

coordination number of Cl

Figure 2. (a) Stepwise stability constant of [CuCln](2‒n)+ in log Kn plotted against the coordination number of Cl‒ of the complex (nCl in the text). Filled squares linked by the dotted line (for eye-guide) correspond to C4mimTfO ionic liquid (present data). Blank symbols indicate conventional molecular liquids; propylene carbonate (PC, ◊ ),40 acetonitrile (AN, ▵ ),36 dimethylformamide (DMF, ◌ ),37 and dimethylsulfoxide (DMSO, ▿ ).38 (b) Stepwise thermodynamic quantities of complexation, ∆Gn°, ∆Hn° and T∆Sn° plotted against nCl (the dotted lines are for eye-guide).

The blank circle indicates stepwise complexation enthalpy of

[Ni(NO3)n](2‒n)+ in C4mimTf2N,25 which is plotted on the stacked axis with a notation of NO3‒ (see text).


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Apparent Reaction Enthalpy (kJ/mol)


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15 10 5 0 -5

0

1

2

3

4 –

5

2+

Concentration Ratio (Cl /Cu )

Figure 3. Apparent reaction enthalpy in the calorimetric titration of Cu(TfO)2 with C4mimCl in C4mimTfO ionic liquid (see text) plotted against concentration ratio of Cl‒ to Cu2+. Four series of data are shown by respective marks. The solid lines are theoretical curve calculated by using finally obtained stability constants and formation enthalpies of [CuCln](2‒n)+ complexes


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∆βHn°/kJ mol

(b) 400

300

200

–1 –1

40 (a) 30 20 10 0 -10 -20 -30 -40 -50 -60 -70 1

∆βSn°/J mol ·K

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


–1

Page 17 of 30

DMF AN DMSO

2

3

4

1

2

3

4

100

0

–

coordination number of Cl

Figure 4. Overall complexation enthalpies (a) and entropies (b) of [CuCln](2‒n)+ plotted against the coordination number of Cl‒ of the complex (nCl in the text). Filled squares linked by the dotted line (for eye-guide) correspond to the C4mimTfO ionic liquid. Blank symbols indicate conventional molecular liquids; dimethylsulfoxide (DMSO, ▿ ),38 dimethylformamide (DMF, ◌),37 and acetonitrile (AN, ▵).36


17


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Figure 5.

Page 18 of 30

Extracted geometries of the primary neighbor of cupper (II) in the respective

optimized complex structures by DFT calculation. The shadowed spheres express chlorines


18

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1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


Table 1. Stability constants and thermodynamic quantities (in kJ/mol) of chloro complexation of Cu2+ in C4mimTfO Stepwise

log Kn

∆Gn° a

∆Hn°

Τ∆Sn° b

1

6.7 (0.7)

‒38 (4)

‒5.1 (0.8)

33 (4)

2

5.8 (0.1)

‒33.2 (0.6)

30 (2)

63 (2)

3

6.3 (0.2)

‒36 (1)

‒31 (2)

5 (2)

4

4.0 (0.9)

‒23 (5)

‒6 (1)

17 (5)

Overall

log βn

∆βGn°

∆βHn°

Τ∆βSn°

2

12.5 (1.4)

‒71 (8)

25 (2)

96 (8)

3

18.8 (1.7)

‒107 (9)

‒6.1 (0.7)

101 (10)

4

22.8 (2.6)

‒130 (15)

‒12.0 (0.9)

118 (15)

Values in parentheses refer to the standard deviations. a ∆Gn° = ‒RT ln Kn (T = 298.15 K, R = 8.314 J/mol⋅K). b ∆Sn° = ∆Hn°/T + R ln K.


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Page 20 of 30

ASSOCIATED CONTENT Supporting Information. Detailed procedure of data analysis, dependence of absorption intensities of Cu2+ at given wavelengths on the Cu2+:Cl‒ concentration ratio, and the formation enthalpies of [CuCln(TfO)m](2‒n‒m)+ complexes estimated by DFT calculations (PDF)

AUTHOR INFORMATION Corresponding Author *Ryo KANZAKI, [email protected] Author Contributions The manuscript was written through contributions of all authors. All authors have given approval to the final version of the manuscript. Funding Sources ACKNOWLEDGMENT This work was supported by JSPS KAKENHI Grant Number 26410157.


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Absorption Spectroscopy and DFT/TDDFT Calculations. J. Phys. Chem. B 2015, 119, 8754-8763. 54. Alarcón-Espósito, J.; Contreras, R.; Tapia, R. A.; Campodónico, P. R. Gutmann’s Donor Numbers Correctly Assess the Effect of the Solvent on the Kinetics of SNAr Reactions in Ionic Liquids. Chem. Eur. J. 2016, 22, 13347-13351.


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TOC Graphic


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Copper(II) Chloro Complex Formation ... - ACS Publications

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