Imbalance between Anion and Cation Distribution at Ice Interface with

Article pubs.acs.org/JPCC

Imbalance between Anion and Cation Distribution at Ice Interface with Liquid Phase in Frozen Electrolyte As Evaluated by Fluorometric Measurements of pH Hiroki Watanabe, Takuhiro Otsuka, Makoto Harada, and Tetsuo Okada* Department of Chemistry, Tokyo Institute of Technology, Meguro-ku, Tokyo 152-8551, Japan S Supporting Information *

ABSTRACT: When an aqueous electrolyte is frozen, anions and cations are distributed between liquid and ice phases in different fashions. This partition imbalance is relaxed by the transfer of H+ and OH− to each phase, resulting in the acidification of the liquid phase when the cation is better distributed in the ice phase than the anion and in the basification in the opposite situation. In this work, a pH change in the liquid phase has been precisely evaluated by fluorescence ratiometry with pyranine as the pH probe. For frozen alkali chlorides (LiCl, NaCl, and KCl), the liquid phase is always basified by freezing due to the preferential partition of Cl− over the alkali metal cations. Changes in pH are quantitatively analyzed by a partition model, in which the distribution of an ion between the liquid and ice phases is determined by the partition coefficient. Since the concentration of a salt (i.e., ions) in the liquid phase in contact with ice becomes higher as freezing proceeds, the concentration of the ions in the ice phase is higher near the interface with the liquid phase and decreases toward the interior of ice. When the temperature of a frozen electrolyte increases, the ionic imbalance is relaxed to some extent by melting of ice near the interface.

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INTRODUCTION Ice plays important roles in the circulation of substances in the global environment; for example, some atmospheric reactions are accelerated on the surface or in the interior of ice.1−3 Since ice found in the natural environment is not pure but contains impurities such as the sea salt, their effects on the physicochemical nature of ice should be well taken into account to discuss the phenomena occurring in a frozen aqueous system.4 We can consider frozen water−NaCl as a model system of naturally occurring ice. Thermodynamics suggests that while only solid phases (ice and NaCl hydrates) are present at the temperature lower than the eutectic point (Teu = ca. 252 K for water/NaCl),5 a liquid phase coexists with solid ice at the temperature above Teu ; the liquid water phase coexistent with ice is hereinafter referred to WPI. The WPI in the water/NaCl system is a highly concentrated aqueous NaCl. Other minor impurities, if any, are also concentrated in the WPI unless their solubility limits are reached. This freeze concentration accelerates some reactions, and that, in some cases, modifies the reaction pathway.3 We have utilized the physicochemical features of ice containing a solute such as NaCl (hereinafter called “doped ice”) to develop analytical methods, such as chromatography, microreaction systems, and preconcentration system.6−12 The management of the WPI volume is essential for the successful developments of these analytical means utilizing physicochemical nature of doped ice. When the dopant concentration is low, the volume of the WPI is so small that the direct determination of the WPI volume is not possible. However, ice chromato© 2014 American Chemical Society

graphic experiments have revealed that the WPI volume is precisely determined from the relevant phase diagram as long as supercooling does not occur.10 The freeze concentration of minor components in the WPI has allowed us to design a novel sample enrichment protocol for capillary electrophoresis.11 The concentration of NaCl in the WPI at −9.3 °C is, for example, 2.5 M.5 If 1 mM NaCl solution is frozen at this temperature, all of the minor solutes as well as NaCl are enriched in the WPI by 2500-fold. Another important feature of the WPI is that some reactions are anomalously enhanced therein. Some reaction enhancements have been explained well by the concentration effect; obviously, second or higher order reactions are kinetically accelerated as the concentrations of reactants become higher.13 However, some other effects that cannot be explained only by the freeze concentration have been also reported. We have found that the complexation constants of some crown ethers in the WPI are 4 orders of magnitude larger than the corresponding ones in bulk water at the same temperature.11 In addition, it has been confirmed that the alkaline hydrolysis of fluorescein diacetate more rapidly occurs in this medium than in bulk water, albeit the mechanism of this acceleration has not been well elucidated.11 Some possible mechanisms have been suggested for the reactions in frozen solution. Klán et al. reported that poor molecular diffusions in the frozen solution, Received: March 31, 2014 Revised: June 25, 2014 Published: July 6, 2014 15723

dx.doi.org/10.1021/jp5031653 | J. Phys. Chem. C 2014, 118, 15723−15731

The Journal of Physical Chemistry C

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for example, can alter reaction selectivity.14 It was also indicated that the catalytic effect of the defects of the ice crystal facilitates the transfer of H+ and enhances the reaction kinetics in frozen aqueous solutions.15,16 In some cases, reactions are affected by a shift of liquid phase pH upon freezing. Takenaka et al.,17 for example, suggested that gallic acid oxidation is accelerated by freezing and revealed that this is mainly caused by an increase in pH of the WPI. The increase in pH is caused by the preferential incorporation of Cl− in the ice phase over Na+. Such an imbalance between anion partition and cation partition to ice has been studied from various viewpoints. Workman−Reynolds freezing potential has been known for a long time; measurable potential (up to some tens of volts) is generated between ice and liquid phases during ice growth.18−22 This potential depends on the salt concentration and the type of ions involved in a solution. A highly electronegative anion (e.g., F−) can replace a water molecule (or the oxygen atom in it) in the ice crystal, and similarly, a molecule having a similar structure to that of water, e.g., NH4+, is well incorporated in the ice crystal. Therefore, when ice is, for example, in contact with solution containing F−, negative freezing potential is yielded. In contrast, positive potential is generated when solution of an NH4+ salt is frozen.19 Such partition imbalance occurs for other salts as well, though the sign and magnitude of potential is varied. The freezing potential is thus caused by the transient charge imbalance in the WPI and the ice phase. This potential is relaxed by the transfer of H+ and OH− to each phase and becomes zero when the advancement of the ice front is stopped.23 The ionic partition to ice has been evaluated in a direct way. Gross and his co-workers24 determined the distribution coefficient of36 Cl− for various alkali salts between ice and WPI. According to their paper, the average value of the distribution coefficient for Cl− to ice was 2.7 × 10−3, which only slightly depended on the type of cation. The determination of ion concentrations in the ice phase should follow careful procedures because the ions incorporated in ice crystals should strictly be discriminated from those precipitated as salts in the interstitial space in polycrystalline ice. A similar issue is involved in the evaluation of the ionic partition based on freezing potential measurements. Since potential generation involves complex transient processes causing charge separation, the quantitative evaluation of the ionic partition from freezing potential is not straightforward. After the relaxation of the imbalances of an electric charge in ice and liquid phases, pH in the WPI is shifted due to the transfers of H+ and OH− as stated earlier. When an anion is better partitioned to ice, the WPI is basified to keep electroneutrality and to compensate for a positive charge excess in the WPI. In contrast, the preferable incorporation of a cation in ice results in the acidification of the WPI. This pH shift is an equilibrium phenomenon rather than a transient process, and, therefore reliable evaluation is possible if an appropriate probe is applied. Precise measurements of pH in the WPI eventually allow us to quantitatively evaluate the imbalance of the ionic partition to ice upon freezing. Recent developments of potentiometric pH sensors have allowed direct pH measurements in the WPI in frozen samples.25−28 However, for reliable potentiometric measurements of pH, a large volume of the WPI should be present in a frozen sample and, therefore, a relatively high salt concentration is required. The 19F NMR shift of fluorobenzoic acid has also been proposed as an effective probe for pH in frozen solutions.29 From a perspective of wide applicability, pH-

indicator dyes are useful probes to determine pH in frozen solutions. Absorption spectrometry has indicated that the protonation of cresol red is enhanced in the frozen state because of the higher concentration of an acid such as HCl.30 Fluorescence spectroscopy has higher applicability to a frozen system than absorption spectrometry. Wren and Donaldson31 employed reflective fluorescence spectroscopy to measure pH on the surface of ice and discussed its relation to the oceanic oxidation of bromide. Cheng et al.23 have recently reported confocal fluorescence microscopic visualization of pH changes of the liquid phase for a frozen aqueous electrolyte. In the present paper, we also employ a pH-sensitive fluorescent dye, pyranine, to determine pH in the WPI in a frozen aqueous alkali chloride. Because of the high salt concentration in the WPI, appropriate corrections should be made to the equilibrium constants for the rigorous determination of pH based on fluorometric measurements.

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EXPERIMENTAL SECTION Fluorescence spectra were measured with a Hitachi fluorescence spectrophotometer Model F-4500. The sample chamber of the spectrophotometer was purged by dried air to prevent the dew and frost formation on the sample surface. The temperature of a sample was maintained on a Peltier array controlled by a Cell System Peltier controller Model TDC2030R. The reversed side of the Peltier module was cooled by a chiller. The temperature of ice was monitored with a thermistor. Sample solutions were prepared with Milli Q water. Pyranine and alkali chlorides (MCl; LiCl, NaCl, and KCl) were of analytical grade and used as received. Solution pH was adjusted by adding aqueous MOH (LiOH, NaOH, or KOH) so that a single alkali cation was involved in the system. The concentration of pyranine was kept 1000th as low as that of an alkali salt in an initial solution before freezing; this allowed us to treat the solution as a simple binary mixture of water and an alkali chloride. A freshly prepared solution was poured into a homemade Cu cell after deaeration, and then a cover glass slip was put on the solution surface. After the solution was completely frozen in liquid nitrogen, the coverslip was removed to provide the flat ice surface for reflective fluorescence measurements. The Cu cell was then fitted to the Cu stage on the Peltier array installed in the sample chamber of the spectrophotometer. The angle of incidence of the excitation light was kept 45° by fixing the Peltier array in the chamber at an appropriate position. Spectra were measured 10 min after the constant temperature was reached. A sample solution directly put in the Cu cell was in direct contact with the Cu cell surface. Dissolved Cu could affect the fluorescence spectra of pyranine. A PEEK cell was used instead of the Cu cell to assess the effect of dissolved Cu on the pH measurements. No differences between these two cells were confirmed. The Cu cell was used throughout this work because of its high thermal conductivity. For the measurements of liquid samples, a usual 1 cm silica glass cuvette for fluorescence measurements was fixed on a homemade Cu cell holder fitted to the Cu stage. The fluorescence intensities at an emission wavelength (λem) of 510 nm were measured with an excitation wavelength (λex) of 455 nm and at λex = 405 nm. The ratio of the intensity at these different excitation wavelengths (I455/I405) provides the information on pH of a sample. The detailed procedure for pH 15724



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Figure 1. Room-temperature absorption (A and B) and fluorescence (C and D) spectra of pyranine in solutions adjusted at pH 3.0 and pH 9.3.

Figure 2. Excitation spectra of pyranine in frozen NaCl. Fluorescence measurements at 510 nm. (A) Effect of the initial concentration (c0) of NaCl at −5.4 °C. (B) Effect of the temperature (a: −7.2 °C; b: −5.4 °C; and c: −3.4 °C) for c0 =50 mM and 50 μM pyranine. pH in original solutions, 7.0.

the shift of its absorption wavelengths; the undissociated species (HPy) gives an absorption maximum at 405 nm, whereas it is shifted to 455 nm when the proton is dissociated (Py−). It is known that the proton dissociation of HPy at the excited states so rapidly occurs that excited Py− always emits fluorescence.34,35 This explains why HPy and Py− give identical fluorescence spectra despite different absorption spectra, and no wavelength shifts are observed in the fluorescence spectra of HPy and Py−. Solution-phase absorption and fluorescence spectra of pyranine with different concentrations are also given in Figure 1. Effects of concentration on the spectral features of pyranine

measurements in the frozen state is given in the Supporting Information.

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RESULTS AND DISCUSSION pH Evaluation with Pyranine. Pyranine has been utilized for determining pH in small spaces, to which potentiometric pH sensors are inapplicable in a direct fashion.32,33 Figure 1 depicts absorption and fluorescence spectra of pyranine in acidic and basic solutions. Since pKa for pyranine was determined to be 7.6 in an aqueous solution at 298 K, it is fully protonated at pH 3.0 but is almost dissociated at pH 9.3. Dissociation of the OH group in a pyranine molecule induces 15725



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are not seen in either absorption or fluorescence spectra. It is known that the aggregate formation of some dye molecules leads to spectral shifts.36,37 Highly accumulated negative charges in a pyranine molecule prevent such aggregate formation. In actuality, an increase in the pyranine concentration does not lead to either spectral shifts or the emergence of additional bands as shown in Figure 1. Figure S1 depicts a fluorescence micrograph of KCl-doped ice containing pyranine. Bright points represent the WPI, suggesting that fluorescence is emitted from pyranine dissolved in WPI and no signals come from the ice phase. Figure 2 shows the excitation spectra of pyranine in frozen NaCl. Figure 2A shows the effect of the concentration of NaCl and pyranine in the original solution. As discussed below in detail, the NaCl concentration in the WPI is constant at the constant temperature, and, therefore, the pyranine concentration therein is also constant. However, since the volume of the WPI almost linearly increases with increasing NaCl concentration, the amount of pyranine in the detection volume linearly increases during this change; thus, the fluorescence intensity also increases as shown in Figure 2A. Figure 2B shows an effect of temperature. The concentrations of NaCl and pyranine in the WPI increase as the temperature becomes lower. Essential spectral features are not affected by the temperature, i.e., by the pyranine concentration in the WPI, suggesting that pyranine does not form aggregates even in the frozen state. This is supported by the size of the WPI reported in our previous paper,6 which revealed that the size of a liquid phase in frozen electrolyte is in the micrometer range. The aggregation would be enhanced if the dye molecules were confined in much smaller spaces, e.g., at very low temperature.30 However, the present situations are not the case. The ratio of fluorescence intensity measured with the excitation at 455 nm to that at 405 nm gives information on pH. In an aqueous solution at 298 K, pH of a solution around 7.6 ± 1.5 can be determined by the ratiometric measurement of fluorescence intensities. Figure S2 shows a typical calibration graph for the pH determination in solution. This calibration graph could be applied to the determination of pH in solution under the ambient condition. As described in detail in the Supporting Information, the fluorescence intensity ratio (I455/ I405) was related to the dissociation degree of pyranine, and then this relation was applied to the determination of pH in the frozen state. In the present study, the solution always contains an alkali chloride as the main dopant and pyranine as the minor component. The system was considered as a simple water-alkali chloride binary mixture. Pyranine is basically condensed in the WPI when a solution is frozen and the temperature is kept between the freezing point and Teu of the binary system. Figure S3 shows the freezing point depression curves for water/LiCl, water/NaCl, and water/KCl.5 The present discussion is confined in the range of concentration lower than the eutectic concentration and in the range of temperature higher than Teu of the individual system. The concentration of LiCl, NaCl, or KCl in the WPI (cWPI) is, for example, 2.5 M at −13.1 °C, −9.3 °C, and −8.8 °C, respectively, regardless of the initial salt concentration (c0) in a solution before freezing. Figure 3 shows the dependence of pH of the WPI on c0 of LiCl, NaCl, and KCl when cWPI = 1.0, 1.5, 2.0, and 2.5 M. It should be noted that the temperature providing the given cWPI is different for LiCl, NaCl, and KCl. As shown in Figure S4, the fluorescence intensity ratio (I455/ I405) in frozen sample showed

Figure 3. Comparison of pH in the WPI determined by fluorescence ratiometry (plots) with that calculated assuming no ion incorporation in ice (curves) for LiCl, NaCl, and KCl. cWPI was 1.0, 1.5, 2.0, and 2.5 M at the temperature (in the order of decreasing temperature) labeled in the corresponding figures.

a time-change after the sample prepared in liquid nitrogen was transferred on the Peltier stage (even after the temperature became constant). Therefore, spectra were measured at least 10 min after the temperature became constant. The relative standard deviation of the measured fluorescence intensity ratio was smaller than 5% as shown in Figure S4. The uncertainty of pH is smaller than 0.05. In this study, pH of a solution before freezing was precisely adjusted to 7.0 before freezing; an HPy− Py− pair acts as a pH buffer. A simple conclusion from Figure 3 may be the acidification of the WPI by freezing because pH is lowered from the initial value of 7.0. However, in actuality, the WPI is basified as discussed below. In the present case, the acid−base equilibrium of pyranine occurs in the WPI, which is an aqueous solution of a high salt concentration. Solution equilibria therein can be treated in a similar way to those in bulk solutions. When an aqueous electrolyte is frozen, a salt is concentrated in the WPI. The salt concentration in the WPI is so high that the dissociation constant of pyranine should be appropriately corrected. Table 1 Table 1. Dependence of pKa on the Salt Concentration pKa

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concentration/M

LiCl

NaCl

KCl

2.5 2.0 1.5 1.0 0.5 0.25

6.53 6.62 6.72 6.77 6.88 7.03

6.74 6.79 6.84 6.88 6.98 7.05

6.88 6.92 6.98 7.00 7.04 7.12



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lists pKa of pyranine determined in the model WPI. A high salt concentration obviously enhances the dissociation of pyranine to a significant extent; pKa of pyranine in the WPI is no longer 7.6. The temperature also affects pKa of pyranine. In order to simultaneously take these effects into account, pKa values were determined in the model WPI, i.e., in 2.5 M LiCl at −13.1 °C, 2.5 M NaCl at −9.3 °C, and 2.5 M KCl at −8.8 °C etc. The pKa value was, for example, lowered to 6.53, 6.74, and 6.88 for the model WPI of LiCl, NaCl, and KCl, respectively. The curves in Figure 3 show pH calculated with corrected pKa assuming that ions are simply concentrated in the WPI and are not partitioned to the ice phase. Due to the shift of pKa of pyranine in the WPI, pH becomes lower upon freezing even though no ion partition to ice occurs. Of importance is that experimental points are deviated from the calculated curves in the basic direction, suggesting that the liquid phase is basified by freezing. Under any conditions (c0 and temperature), pH of the liquid phase is higher than the corresponding one calculated assuming that no ionic partition occurs. This means that Cl− is incorporated in ice better than any alkali cations studied here. As stated already, pH of a solution of MCl (M+ is an alkali cation) before freezing was adjusted at 7.0 by adding MOH. In the WPI, the following charge balance equation can be written:

[Cl−]ice − [M+]ice =

VWPI ([Py −]WPI − [M+]add ) Vice

(4′)

−

We can calculate [Py ]WPI from pH of the WPI, and pKa and the concentration of pyranine therein, and also [M+]add WPI from the concentration of MOH used to adjust pH of the initial solution at 7.0 and the freeze concentration factor, which is equal to cWPI/c0 for MCl. Figure 4 shows the alkali cation excess concentration in the WPI for LiCl, NaCl, and KCl under various conditions,

[M+]WPI + [H+]WPI + [M+]add WPI = [Cl−]WPI + [OH−]WPI + [Py −]WPI +

(1)

[M+]add WPI

where [M ]WPI and are the concentrations of a metal ion originating from MCl and from MOH added for pH adjustment. Distinction between [M+]WPI and [M+]WPIadd facilitates the estimation of ionic incorporation in ice. The alkali cation excess concentration over the Cl − concentration in the WPI is simply given by [M+]WPI − [Cl−]WPI = [OH−]WPI + [Py −]WPI − [H+]WPI − [M+]add

(2)

Similarly, the chloride excess concentration in the ice phase is represented by [Cl−]ice − [M+]ice =

VWPI ([M+]WPI − [Cl−]WPI ) Vice

(3)

Figure 4. Dependence of the alkali cation excess concentration in the WPI on c0 and temperature. Curves are results of fitting to eq 12 with (KCl − KM) as the fitting parameter.

where V is the volume. Substitution of eq 1 into eq 2 yields [Cl−]ice − [M+]ice V = WPI ([OH−]WPI + [Py ‐]WPI − [H+]WPI − [M+]add ) Vice

calculated on the basis of eq 2′. The cation excess concentration is at a level of 10−4 M and decreases with increasing c0 or temperature. Interpretation of pH shift by ionic partition. The pH shifts by freezing are interpreted on the basis of an ion-partition model, which assumes that an individual ion is distributed between ice and liquid phases. The partition coefficient of an ion species, i, is defined in a usual way

(4) −

The calculation of [OH ]WPI requires the self-dissociation constant of water (KW), which is also varied with the temperature and salt concentration. However, it has been reported that the dependence of Kw on the ionic strength is not very large,38 and, in addition, Kw is lowered with decreasing temperature. Thus, [OH−]WPI at pH of Na+ ≥ Li+. Wilson and Haymet have shown that freezing potential depends on the electrolyte concentration.7 Since the data reported by Cobb and Gross did not necessarily pay an attention to the concentration effect, the order of ionic partition revealed in the present work cannot be deduced directly from freezing potential data. However, the present results are qualitatively consistent with freezing potential data as well.

dVWPI

⎛ci ⎞ ⎟ ln⎜ WPI i ⎝ c0 ⎠

cWPI/M

(KCl − KM) was determined by curve-fitting of eq 12 to the data depicted in Figure 4.

Substitution of eq 8 into eq 6 yields m iice = −

⎛c ⎞ cV ρ Δmiice = (K Cl − KM) 0 0 WPI ln⎜ WPI ⎟ VWPI VWPI ρice ⎝ c0 ⎠

(10)

The amounts of ions entrapped in ice are so small that the concentration of an anion in the WPI is always equal to that of Cl M Cl a cation therein. Therefore, cM WPI = cWPI = cWPI and c0 = c0 = c0 are assumed in eq 10. 15728



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We have studied the local structures of Br− in frozen electrolyte by X-ray absorption fine structure (XAFS).40,41 Although Br− should also be partitioned in the ice phase, a main part of the ion in the frozen state is in the form of a salt or dissolved in the solution present in the interstitial space (grain boundaries) in polycrystalline ice. Therefore, the local structure of Br− given in our previous work does not reflect that of Br− partitioned to the ice phase. Molecular dynamics simulations have indicated that both Na+ and Cl− are entrapped in the approximately middle of the hexametric cavity of the ice Ih crystal.42,43 The coordination numbers of the ions accommodated in the ice crystal are smaller than the corresponding ones in bulk solutions, and the coordination distances in the ice phase are also different from those in bulk water. Thus, ions do not adopt the usual hydration structures when partitioned to ice. When the hydration structure of an ion is disturbed in a given medium, a poorly hydrated ion tends to be stabilized therein relative to its well-hydrated counterparts; i.e., the phase transfer of a poorly hydrated ion in general, is more favorable than that of the latter in terms of the Gibbs energy of transfer.44−46 The distribution of ions to an ion-exchange resin from an aqueous solution is a typical example; i.e., a poorly hydrated counterion has higher affinity to an ion-exchange resin than a well-hydrated counterion because of a smaller loss of the Gibbs energy of transfer from the solution phase to the resin phase. For alkali cations, cation-exchange selectivity in water basically follows the order of Cs+ > Rb+ > K+ > Na+ > Li+.47 In the present study, it has been revealed that the partition of ions becomes lower in the order of K+ > Na+ > Li+. This order of ionic partition to the ice phase is interpreted in a similar fashion to the distribution of countercations to an ion-exchange resin. Figure 6 shows the temperature dependence of the anion excess concentration in ice. The Cl− excess concentration

temperature near Teu. The ionic partition is possibly affected by a freezing rate. It is in general difficult to precisely control the temperature of a sample during a freezing process. Ice samples were prepared on the Peltier array to get qualitative information on the effect of freezing rate. Although a series of measurements resulted in slightly larger (KCl − KM) than those listed in Table 2, there were no essential differences. In addition, only a marginal hysteresis has been confirmed when the temperature increases and then again decreases as shown in Figure S5, strongly suggesting that the partition equilibrium at the ice/ WPI interface is basically reached at any temperature.

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CONCLUSION The WPI is basified when an alkali chloride solution is frozen. This has been predicted from freezing potential measurements and other previous work. In this paper, we have shown quantitative changes in pH in the WPI of a frozen alkali chloride. An advantage of pH measurements by fluorescence ratiometry is that a small change in the concentration (10−7 M level) can be evaluated as a clear change in pH. No other methods could sense such small concentration changes occurring in a frozen sample. It is another important advantage of this method that a difference in ion partition between an anion and cation is selectively probed; the phase transfer as a salt or ion-pair does not contribute to pH changes in the WPI, but ionic partition does. This approach is applicable to any other neutral salts, and the ionic partition to the ice phase is quantitatively evaluated if the phase diagram is known. Since the partition imbalance between a cation and an anion is not very large in most of the neutral salts, only a small change in pH in the WPI should take place upon freezing. The elucidation of the local structures of an ion accommodated in ice crystals is also of fundamental interest and will facilitate understanding of the mechanism of ionic partition to ice. The partition coefficient of an ion implies that the ion concentration in ice is as high as 1 mM or slightly higher. This concentration is possibly high enough to study the structural feature of ions entrapped in the ice crystal by XAFS. We believe that such an experimental study makes significant contributions to the advancements of ice chemistry and possibly to that of aquatic chemistry.

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ASSOCIATED CONTENT

* Supporting Information S

Details of the procedure for pH measurements, a fluorescence micrograph of ice doped with pyranine and KCl, a calibration graph for pH measurements, freezing depression curves for LiCl, NaCl, and KCl/water systems, time changes of the fluorescence intensity ratio (I455/ I405) with time, and hysteresis of the fluorescence intensity ratio for a frozen NaCl system. This material is available free of charge via the Internet at http://pubs.acs.org.

Figure 6. Temperature dependence of the chloride excess concentration in ice for LiCl, NaCl, and KCl when c0 = 10 mM.

decreases as the temperature increases, because the concentrations of ions in the WPI decrease with an increase in the temperature. This figure demonstrates that the ionic imbalance in ice is more marked near the ice/WPI interface. There is a gradient in an anion excess concentration in ice, which decreases from the liquid interface to the deep interior of the ice phase. A part of ice involving a large Cl− excess first melts to form the liquid phase as the temperature of MCl-doped ice increases past Teu. Since interfacial ice is acidified to greater extent than the residual part of ice, the neutralization of the WPI with increasing temperature is more marked at the lower

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AUTHOR INFORMATION

Corresponding Author

*Phone and Fax +81-3-5734-2612. E-mail [email protected]. ac.jp. Notes

The authors declare no competing financial interest. 15729



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(22) Wilson, P. W.; Haymet, A. D. J. Workman−Reynolds Freezing Potential Measurements between Ice and Dilute Salt Solutions for Single Ice Crystal Faces. J. Phys. Chem. B 2008, 112, 11750−11755. (23) Cheng, J.; Soetjipto, C.; Hoffmann, M. R.; Colussi, A. J. Confocal Fluorescence Microscopy of the Morphology and Composition of Interstitial Fluids in Freezing Electrolyte Solutions. J. Phys. Chem. Lett. 2010, 1, 374−378. (24) Gross, G. W.; Wong, P. M.; Humes, K. Concentration Dependent Solute Redistribution at the Ice−Water Phase Boundary. III. Spontaneous Convection. Chloride Solutions. J. Chem. Phys. 1977, 67, 5264−5273. (25) Sundaramurthi, P.; Shalaev, E.; Suryanarayanan, R. “pH Swing” in Frozen Solutions-Consequence of Sequential Crystallization of Buffer Components. J. Phys. Chem. Lett. 2010, 1, 265−268. (26) Sundaramurthi, P.; Suryanarayanan, R. Thermophysical Properties of Carboxylic and Amino Acid Buffers at Subzero Temperatures: Relevance to Frozen State Stabilization. J. Phys. Chem. B 2011, 115, 7154−7164. (27) Sundaramurthi, P.; Suryanarayanan, R. The Effect of Crystallizing and Non-Crystallizing Cosolutes on Succinate Buffer Crystallization and the Consequent pH Shift in Frozen Solutions. Pharm. Res. 2011, 28, 374−385. (28) Sundaramurthi, P.; Shalaev, E.; Suryanarayanan, R. Calorimetric and Diffractometric Evidence for the Sequential Crystallization of Buffer Components and the Consequential pH Swing in Frozen Solutions. J. Phys. Chem. B 2010, 114, 4915−4923. (29) Robinson, C.; Boxe, C. S.; Guzman, M. I.; Colussi, A. J.; Hoffmann, M. R. Acidity of Frozen Electrloyte Solutions. J. Phys. Chem. Lett. 2006, 110, 7613−7616. (30) Heger, D.; Klanova, J.; Klan, P. Enhanced Protonation of Cresol Red in Acidic Aqueous Solutions Caused by Freezing. J. Phys. Chem. B 2006, 110, 1277−1287. (31) Wren, S. N.; Donaldson, D. J. How Does Deposition of Gas Phase Species Affect pH at frozen Salty Interfaces? Atmos. Chem. Phys. 2012, 12, 10065−10073. (32) Yamaguchi, A.; Namekawa, M.; Kamijo, T.; Itoh, T.; Teramae, N. Acid-Base Equilibria inside Amine-Functionalized Mesoporous Silica. Anal. Chem. 2011, 83, 2939−2946. (33) Kano, K.; Fendler, J. H. Pyranine as a Sensitive pH Probe for Liposome Interiors and Surfaces. pH Gradients across Phospholipid Vesicles. Biochim. Biophys. Acta, Biomembr. 1978, 509, 289−299. (34) Spry, D. B.; Goun, A.; Bell, C. B., III; Fayera, M. D. Identification and Properties of the 1La And 1Lb States of Pyranine. J. Chem. Phys. 2006, 125, 144514. (35) Mohammed, O. F.; Dreyer, J.; Magnes, B.-Z.; Pines, E.; Nibbering, E. T. J. Solvent-Dependent Photoacidity State of Pyranine Monitored by Transient Mid-Infrared Spectroscopy. ChemPhysChem 2005, 6, 625−636. (36) Ageev, D. V.; Patsaeva, S. V.; Ryzhikov, B. D.; Sorokin, V. N.; Yuzhakov, V. I. Influence of Temperature and Ethanol Content on Aggregation of Rhodamine 6G Molceules in Aqueous Ethanol Solutions. J. Appl. Spectrosc. 2008, 75, 653−657. (37) Zehentbauer, F. M.; Moretto, C.; Stephen, R.; Thevar, T.; Gilchrist, J. R.; Pokrajac, D.; Richard, K. L.; Kiefer, J. Fluorescence Sectroscopy of Rhodamine 6G: Concentration and Solvent Efects. Spectrochim. Acta Part A 2014, 121, 147−151. (38) Kron, I.; Marshall, S. L.; May, P. M.; Hefter, G.; Konigsberger, E. The Ionic Product of Water in Highly Concentrated Aqueous Electrolyte Solutions. Monatsh. Chem. 1995, 126, 819−837. (39) Cobb, A. W.; Gross, G. W. Interfacial Electrical Effects Observed during the Freezing of Dilute Electrolytes in Water. J. Electrochem. Soc. 1969, 116, 796−804. (40) Harada, M.; Tasaki, Y.; Qu, H.; Okada, T. Hydration of Ions and Salt Crystallization in Liquid Phase Coexistent with Ice at Temperature below Eutectic Point. RSC Adv. 2012, 2, 461−466. (41) Tasaki, Y.; Harada, M.; Okada, T. Eutectic Transition of Local Structure for Bromide Ion in Bulk and on Surface of Doped Ice. J. Phys. Chem. C 2010, 114, 12573−12579.

ACKNOWLEDGMENTS This work has been supported by a Grant-in-Aid for the Scientific Research from the Japan Society of the Promotion of Science and by SENTAN from the Japan Science and Technology Agency.

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Imbalance between Anion and Cation Distribution at Ice Interface with

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