Effect of KCl and CsCl on the Electrical Conductivity of Molten LiF–KBr

Article pubs.acs.org/jced

Effect of KCl and CsCl on the Electrical Conductivity of Molten LiF−KBr at the Critical Composition Victor P. Stepanov,* Larisa M. Babushkina, Stanislav I. Dokashenko, and Dmitri S. Peshkin Institute of High-Temperature Electrochemistry, Russian Academy of Sciences, Ural Division, Academicheskaya str. 20, Yekaterinburg, 620990 Russia S Supporting Information *

ABSTRACT: The electrical conductivity was measured from the melting point to 1280 K for molten 0.7 LiF−0.3 KBr (its composition corresponds to the top of the miscibility gap) containing (2.3, 4.4, 6.5, 8.8, and 11.2) mol % KCl or (1.2, 2.5, 5.5, and 10.2) mol % CsCl to establish the influence of this solute on the stability of the two-phase system. These results indicate that the temperature dependences of the conductivity along the saturation lines for all of the mixtures studied herein are similar to one another. Hence, this demonstrates that small additions of KCl and CsCl to the dissolving melt of LiF-KBr do not exert a substantial influence on its type of criticality. In the vicinity of the critical point, the temperature dependence on conductivity differences for melts is investigated and is described by the equation Δκ ≈ (Tc − T)k, where k is the critical exponent (k = 0.98). The critical temperature changes as a function of the mixture composition and depends on the ion size of the salt added. The critical temperature increases continuously with the addition of CsCl to molten LiF-KBr, whereas it decreases as the fraction as KCl is added. This circumstance must occur during the organization process, as salts accumulate in the dissolving molten mixture, and they prevent the confluence of the phases at a given operating temperature. To interpret the experimental results, the charged hard sphere model for ionic melts in the Debye−Hückel approximation was used with an account of the excluded volume.

■

INTRODUCTION

The realization of this idea requires one to solve many problems. For example, does the solute influence the stability of the initial two-phase system? If this influence occurs, then what quantity of solute can be dissolved in the phases without leading to the confluence of the phases? In the proposed work, we report on the results by measuring the electrical conductivity of the immiscible reciprocal mixture of 0.7 LiF− 0.3 KBr with the addition of KCl or CsCl to establish the influence of the solute on the miscibility gap parameters. In the literature, this question has not been adequately addressed. The electrical conductivity of the 0.7 LiF−0.3 KBr melt whose composition corresponds to the top of the miscibility gap6 was investigated earlier.7

1

The effect of stratification in the molten salt mixtures suggests great possibilities for its use in different technological processes. One example is the accumulation and the emission of the heat that occurs from the phase transition of mixing and demixing in thermal devices.2 The application of immiscible melts for the concentration of any substance in one of the phases may be promising. This possibility is demonstrated for the salt systems KNO 3 −AgCl, K 2 S 2 O 7 −AgCl, KNO 3 −AgBr, and LiCl− KAlCl4.3−5 The distribution data obtained in these works indicate that useful separation can be achieved because of the different solutes of the components that coexist in the phases. For example, the molar fraction CsCl in the light phase of the system LiCl−KAlCl4 is enriched by KAlCl4, which proved to be 18 times more enriched than the concentration of CsCl in the bottom phase, which was enriched by LiCl.5 New prospects unfold for the dissolving alkali halide melts6 caused by unique technological parameters. The most valuable qualities of this system are good electrical conductivity7 and a high thermodynamic and radiation durability. Specifically, these properties of the salt melts can help to resolve the question of nuclear fuel regeneration and utilization of the fission products.8 Actually, the salt melts will be considered successful if they achieve a sufficiently concentrated solution of these products in one of the phases such that the solution can be subsequently extracted from the phase by means of electrolysis. © XXXX American Chemical Society

■

EXPERIMENTAL SECTION The electrical conductivity measurements were made by an impedance method using the experimental setup described previously.7 All of the parts of the cell contacting the melts were of corrosion stable materials. One of the major sections of the setup is a measuring unit made of a synthetic ruby that showed strong chemical stability in the melts investigated in this study (Figure 1). The ruby cylinder has two longitudinal channels for the platinum working electrodes made of 99.99 % pure Received: April 11, 2012 Accepted: July 4, 2012

A

dx.doi.org/10.1021/je300415f | J. Chem. Eng. Data XXXX, XXX, XXX−XXX

Journal of Chemical & Engineering Data

Article

ments was estimated to be 0.5 %, and the relative instrumental error was 0.06 %. Particular attention was paid to the elimination of any impurities that could arise from the reactions of the test salts with atmospheric oxygen and moisture. The salts (> 99.5 % purity) were dried under vacuum, and the temperature was gradually increased thereafter up to the melting point. The dried salts were subjected to zone refinement six times in pure flowing argon. The measurements were performed in an atmosphere of pure argon after waiting for the temperature to stabilize after each temperature change to achieve equilibrium. This approach led to a constant electrical conductivity in each phase at a certain point of the melt at each stationary temperature, which was evidence that the system reached thermodynamic equilibrium.

■

RESULTS AND DISCUSSION Figure 2 displays the electrical conductivity as a function of temperature for the LiF−KBr melt which contains (2.3, 4.4, 6.5,

Figure 1. Measuring unit for the investigation of electrical conductivity of immiscible melts: 1, Pt crucible with melt; 2, alumina cylinder; 3, Pt work electrodes; 4, Pt/Pt−Rh thermocouple.

platinum wire measuring 0.1 cm in diameter. The lower tips of the electrodes were in a plane parallel with the melt surface, which allowed us to probe the thicknesses of the melts at all of the steps. The two electrodes were immersed in the liquid sample, which was in a platinum crucible enclosed in a quartz container with a controlled gas space that was hermetically closed by a stopper. The temperature was measured by a Pt/ Pt−Rh thermocouple, calibrated against a reference thermocouple, where the temperature was recorded by an analogue-todigital converter with an accuracy of 0.5 K. The heating component of the experimental unit included a tube resistance furnace. The Pt/Pt−Rh thermocouple in the immediate vicinity of the heating section was used to control the strength of the current in the windings automatically. This geometry of the heating furnace allowed the temperature of the melts to be maintained within ± 1 K, while the zone of isothermal heating was 70 mm high. The impedance was measured with a potentiostat/galvanostat Parstat 2273 in the frequency range from 1 Hz < f < 1 MHz with a 50 mV amplitude sine-wave excitation. The measured impedance was extrapolated to infinite frequency to correct the data for the electrode polarization. The short-circuit cell was investigated to allow for the cell and the electrical wires impedances. Cell calibration was conducted at the same frequencies and temperatures as in the basic measurements using a melt of potassium chloride of known conductivity.9 The temperature dependence of the cell constant was taken into account. The absolute accuracy of the conductivity measure-

Figure 2. Electric conductivity κ as a function of temperature for molten LiF−KBr containing ⧫, 0;7 □, 2.3; ×, 4.4; ◊, 6.5; △, 8.8; ●, 11.2 mol % KCl and ▲, 1.2; ○, 2.5; ∗, 5.5; ■, 10.2 mol % CsCl.

8.8, and 11.2) mol % KCl or (1.2, 2.5, 5.5, and 10.2) mol % CsCl in the two-phase region along the saturation line. Figure 2 displays the conductivities of the molten mixtures that lie in the range from (2.2 to 3.7) S·cm−1, which is between the values characteristic for LiF on the one end of the interval and for the more heavy alkali halides on the other end. This observation is evidence of a mutual solubility of these salts. The temperature rise is accompanied by an increase in the conductivity of all of the low phases, as is expected in pure halide melts.9 Obviously, these dependences are determined by both the change in the particle interaction energy in the melts and the variations in the phase compositions as a result of the mutual solution of the components in each saturated phase. When enriched in the heavy halide phases, both factors act in the same direction because the conductivity increases with increasing temperature as a result of both the weakening of the ion−ion bonding and the increase in the concentration of the high conductivity lithium fluoride phase.9 It is worth mentioning that the temperature dependence of the conductivity of the melts of LiF−KBr with and without KCl and CsCl in the heavy phase is collinear and increasing monotoniB

dx.doi.org/10.1021/je300415f | J. Chem. Eng. Data XXXX, XXX, XXX−XXX

Journal of Chemical & Engineering Data

Article

dependent on both the quantity of added alkali chloride and on its nature. Figures 4 and 5 clearly illustrate this thesis. Figure 4

cally with rising temperature. It would appear that, in this case, no notable change of structure is occurring for the melts. However, the electrical conductivity of the lithium fluorideenriched phases does not change in a similar way. For all mixtures, the electrical conductivity decreases when the temperature increases. Apparently, the impact on electrical conductivity with increasing temperature in the light phase is due to the predominant influence from the increasing concentration of the less mobile ions (e.g., K+, Cs+, Br−, and Cl−), rather than the Li+ and F− ions. Given the low solubility in lithium fluoride of the alkali halides having large ions,6 the conductivity of the upper phase would appear to be an exceptionally strong function of composition. The temperature factor which acts in the opposite direction cannot compensate for this effect. It is possible that the anomalous behavior of the light phase is due to a rearrangement of the structure of the liquid lithium fluoride in the presence of a small amount of the alkali halides with their large ions. Of the alkali halide forming ions, Li+ and F− have the smallest radii, so lithium fluoride is characterized by a distribution of interionic energy which is approximately evenly distributed over the volume of the melt. The appearance of the alkali halide RX with a cation R+ and anion X− in this melt with a larger size causes a local nonuniformity of the structure in the mixture.7 Theoretically, the near-identical nature of the change in electrical conductivity for both the mixture LiF−KBr and for the fusion process with the addition of KCl and CsCl is an important finding. Figure 3 reveals the difference, Δκ = κ1 − κ2,

Figure 4. Difference in the electrical conductivities Δκ of upper and lower phases at 1200 K as a function of concentration of KCl, ○ and CsCl, ⧫.

Figure 5. Critical temperature of mixing for melts versus the content of KCl, ○ and CsCl, ⧫.

shows the difference in the electrical conductivity of the equilibrium phases (Δκ) at 1200 K, which depends on the concentration of both KCl and CsCl. It is evident that Δκ monotonically increases in proportion to an increase in the concentration of CsCl in the fusion. On the contrary, Δκ decreases when the quantity of KCl increases in the melt. The mixture containing 8.8 mol % KCl at 1200 K reaches the critical point of mixing, whereas the fusion with 11.2 molar % KCl is already single-phase at this temperature. Figure 5 shows the critical temperature of mixing (Tc) as a function of the content of the solutes in the melts. The critical temperature changes with the mixture composition, depending on the ion size in the salt addition. The critical temperature increases continuously with the addition of CsCl to molten LiF−KBr, whereas it decreases continuously with the addition of KCl. In both cases, the curves are slightly concave down with respect to the expected linear dependence upon composition. There are many factors which may be useful in understanding the mutual solubility of inorganic salts. The reciprocal alkali halide ternary mixtures investigated in this work represent a convenient model object in the development of a theory for critical phenomena. A pair potential in these melts can be

Figure 3. Difference in the electrical conductivities Δκ of upper and lower phases for melts LiF−KBr containing ⧫, 0;7 □, 2.3; ×, 4.4; ◊, 6.5; △, 8.8; ●, 11.2 mol % KCl and ▲, 1.2; ○, 2.5; ∗, 5.5; ■, 10.2 mol % CsCl as a function of the difference between critical Tc and current T temperatures.

in the electrical conductivities, κ1 and κ2, of the upper and lower phases, respectively, depends upon the difference between the critical temperature, Tc, and the oven temperature, T. It is evident that all experimental points lie on one curve. Therefore, small additions of KCl and CsCl to the dissolving LiF−KBr melt do not exert a substantial influence on its criticality type. In the vicinity of the critical point, the temperature dependences of conductivity differences for the melts investigated are described by the equation Δκ ≈ (Tc − T)k, where the critical exponent k = 0.98. Thus, KCl and CsCl did not appear to influence the form of the miscibility gap in our experiment, which is typical for the mixture LiF−KBr. However, Figure 2 shows that the critical temperature changes substantially relative to the temperature scale, which is C

dx.doi.org/10.1021/je300415f | J. Chem. Eng. Data XXXX, XXX, XXX−XXX

Journal of Chemical & Engineering Data

■

described with a long-range Coulomb interaction and the Born−Mayer repulsion at small distances.10 These two interaction types were considered in the charged hard sphere model for ionic melts to interpret the experimental phase diagrams with incomplete miscibility.11 In this study, the specifics of the melts were calculated as a mixture of primitive electrolytes in which the ions have identical charges but different sizes. The Coulombic character of the particle interaction in these systems renders it possible to use the Debye−Hückel approximation while accounting for the excluded volume to describe the melt exsolution. An analysis of the contributions from these components in the exchange chemical potential revealed the nature of the immiscibility of these melts. Demixing arises from the different shielding abilities of the ions: the smaller the radius of the ion, the better its ability to shield. Therefore, the cation and anion with the smallest sizes tend to segregate and separate the system into a pair of phases with different component concentrations. In these systems, the alignment of the constituent ion sizes in the melts provides a main factor in the determination of the liquid−liquid phase transition.11 A theoretical analysis11 shows that an increase in the ionic radii difference is accompanied by an increase in the critical mixing temperature according to the correlation Tc ≈ (dRX − dLiF)2. In this correlation, dRX is the sum of radii of the large ions of the mixtures, and dLiF is the sum of the radii of the small ions. Our experimental results concerning the effect of KCl or CsCl on the miscibility of LiF− KBr are in close agreement with this theoretical conclusion. Indeed, dKCl < dKBr < dCsCl.12 Therefore, one may predict that the addition of KCl to the mixture of LiF−KBr should decrease the critical temperature, and the addition of CsCl increases the critical temperature. This condition must be met during the organizational process, as the salts accumulate during the dissolution of the molten mixture to preserve the system in the two-phase state at any given operating temperature.

Article

REFERENCES

(1) Kendall, J.; Crittenden, E. D.; Miller, H. K. A study of the factors influencing compound formation and solubility in fused salt mixtures. J. Am. Chem. Soc. 1923, 45, 963−996. (2) Pacheco, J.; Showalter, S. K.; Kolb, W. J. Development of a molten-salt thermocline thermal storage system for parabolic trough plants. J. Sol. Energy Eng., Trans. ASME 2002, 124, 153−159. (3) Kennedy, J. H. Distribution of thallium chloride between KNO3 and AgCl and between K2S2O7 and AgCl. J. Phys. Chem. 1961, 65, 1030−1033. (4) Kennedy, J. H. Distribution of TlBr between KNO3 and AgBr. J. Chem. Eng. Data 1964, 9, 95−98. (5) Moore, R. H. Distribution coefficients for certain actinide and fission product chlorides in the immiscible salt system LiCl-KAlCl4. J. Chem. Eng. Data 1964, 9, 502−505. (6) Margheritis, C.; Flor, G.; Sinistri, C. Miscibility gaps in fused salts. Z. Naturforsch. 1973, 28A, 1329−1334. (7) Stepanov, V. P.; Babushkina, L. V.; Dokashenko, S. I. Liquid + liquid equilibrium of lithium fluoride with potassium and rubidium halides. J. Chem. Thermodyn. 2012, 51, 12−16. (8) Mecham, N. J.; Jonke, A. A. Commercial aspects of fuel processing. Reactor Fuel Proc. 1966, 9, 137−146. (9) Janz, G. J.; Dampier, F. W.; Lakshminarayan, G. R.; Lorenz, P. K.; Tomkins, R. P. T. Molten Salts; National Standard Reference Data, NBS; National Institute of Standards and Technology: Gaithersburg, MD, 1968; Vol. 15, pp 1−168. (10) Sundheim, B. R. Fused Salts; McGraw-Hill Book Co.: New York, 1964. (11) Tkachev, N. K. Miscibility of salt melts and ionic size mismatch. Melts 1999, 5, 90−94. (12) Fumi, F. G.; Tosi, M. P. Ionic sizes and Born repulsive parameters in the NaCl-type alkali halides. J. Phys. Chem. Solids 1964, 25, 31−43.

■

CONCLUSION The temperature and concentration dependences of the electrical conductivity are investigated for dissolving melts consisting of 0.7 LiF−0.3 KBr with the addition of KCl or CsCl. The addition of KCl increases the critical temperature of the LiF−KBr mixture, while the addition of CsCl decreases the critical temperature. When these shifts in the critical temperature are taken into account, the differences in conductivity measured at compositions on either end of an isothermal tie line at temperatures below critical were the same for both KCl and CsCl.

■

ASSOCIATED CONTENT

S Supporting Information *

Temperature dependence of the electrical conductivity of dissolving melts. This material is available free of charge via the Internet at http://pubs.acs.org.

■

AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. Funding

This work was financially supported by the Russian Foundation for Basic Research (Grant No. 11-03-00149). Notes

The authors declare no competing financial interest. D

dx.doi.org/10.1021/je300415f | J. Chem. Eng. Data XXXX, XXX, XXX−XXX