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ISSN : 1738-1894(Print)
ISSN : 2288-5471(Online)
Journal of Nuclear Fuel Cycle and Waste Technology Vol.18 No.4 pp.549-562
DOI : https://doi.org/10.7733/jnfcwt.2020.18.4.549

# Investigation on Dissolution and Removal of Adhered LiCl-KCl-UCl3 Salt From Electrodeposited Uranium Dendrites Using Deionized Water, Methanol, and Ethanol

Dimitris Payton Killinger, Supathorn Phongikaroon*
Virginia Commonwealth University, 401 W Main St, Richmond, VA 23284, USA
Corresponding Author. Supathorn Phongikaroon, Virginia Commonwealth University, E-mail: sphongikaroon@vcu.edu, Tel: +1-804-827-2278
August 1, 2020 ; October 21, 2020 ; December 22, 2020

## Abstract

Deionized water, methanol, and ethanol were investigated for their effectiveness at dissolving LiCl-KCl-UCl3 at 25, 35, and 50℃ using inductively coupled plasma mass spectrometry (ICP-MS) to study the concentration evolution of uranium and mass ratio evolutions of lithium and potassium in these solvents. A visualization experiment of the dissolution of the ternary salt in solvents was performed at 25℃ for 2 min to gain further understanding of the reactions. Aforementioned solvents were evaluated for their performance on removing the adhered ternary salt from uranium dendrites that were electrochemically separated in a molten LiCl-KCl-UCl3 electrolyte (500℃) using scanning electron microscopy with energy dispersive spectroscopy (SEM-EDS). Findings indicate that deionized water is best suited for dissolving the ternary salt and removing adhered salt from electrodeposits. The maximum uranium concentrations detected in deionized water, methanol, and ethanol for the different temperature conditions were 8.33, 5.67, 2.79 μg·L-1 for 25℃, 10.62, 5.73, 2.50 μg·L-1 for 35℃, and 11.55, 6.75, and 4.73 μg·L-1 for 50℃. ICP-MS analysis indicates that ethanol did not take up any KCl during dissolutions investigated. SEM-EDS analysis of ethanol washed uranium dendrites confirmed that KCl was still adhered to the surface. Saturation criteria is also proposed and utilized to approximate the state of saturation of the solvents used in the dissolution trials.

## 초록

Department of Energy Technology Research & Development

## 1. Introduction

The morphologies and compositions of electrodeposited uranium, as well as some lanthanides, have been studied to further understand and optimize the pyroprocessing separation method [1-10]. Each study utilized different solvents to remove residual salts from the surface of the electroreduced metal to allow for analyzation. Solvents reported for washing adhered salts from the metals include deionized water [1-4], a mixture of deionized water with either methanol [5] or ethanol [6], ethanol [7], and ethanol with partial additions of deionized water during ultrasonic cleaning [8]. Additionally, some studies have reported failed attempts at washing adhered salts from the metals without specifying which solvents were being used [9] or not specifying a washing technique at all [10]. Drying techniques, when reported, also varied, mostly with drying by air [1, 5-7] or with the rinsing of methanol before drying by air [2].

The differences in techniques used to remove the residual salt from the electrochemically separated metalsin addition to the differences in molten salt systems and analytes studied-make it difficult to conclude the most effective method to remove the chloride-based salts adhered to metals. This missing gap becomes the main motivation for this study to investigate and determine a solvent that can satisfactorily rinse and wash the electrodeposited metal from such separation techniques on the same platform. The solvents of interest considered are taken from the literature, being ethanol, methanol, a purified water. The experiments were conducted at three different temperatures (25, 35, and 50℃) to observe any effect on the salt dissolution process. A visualization of the ternary salt dissolving in the solvents of interests was conducted. Additionally, solvents were tested on their performance to actually remove adhered LiCl-KCl- UCl3 salt from uranium dendrites that have been electrochemically separated in molten LiCl-KCl-UCl3 (500℃, 1wt% U) at an overpotential of 50 mV vs. Ag(I)/Ag.

## 2. Experimental

### 2.1 Salt Preparation

LiCl-KCl-UCl3 ternary salt (55.8 mol% LiCl - 44.2 mol% KCl, 1wt% U) was prepared in an argon atmosphere glovebox. The lithium chloride (LiCl, 99.95%) and the potassium chloride (KCl, 99.99%) were purchased from Alfa Aesar (ampouled under argon). Uranium chloride (69.78wt% UCl3 in LiCl-KCl eutectic), procured from Idaho National Laboratory, was used to achieve 1wt% U in the ternary salt system (total mass of 35 g). The salts were mixed in an alumina crucible and dried to 250℃ for 6 hours in a modified benchtop muffle furnace. Following the drying step, the salts were set to melt at 500℃ for 24 hours. Cyclic voltammetry was performed and was in agreement with previous work done on a LiCl-KCl-UCl3 ternary system [4].

### 2.2 Salt Dissolution Study Procedures

A quartz tube (OD = 4 mm, ID = 2 mm) and syringe were used to draw the molten salt out from the alumina crucible and allowed to freeze in the tube. The frozen ternary salt was then removed from the tube and divided into several pieces, approximately 2.54 cm in length for storage, prior to the experiment (see Fig. 1). The entirety of the molten salt was used to produce these ingots. The light purple hue of the salt ingot is characteristic of the UCl3 in the ternary salt.

#### 2.2.1 Quantitative analysis of salt dissolution

A salt ingot was taken and divided further into approximately 32.0 mg tablets. Solvents of interest used in the experiment include deionized water (Direct-Q® 3 UV Water Purification System; 18.2 MΩcm @ 25℃) ethanol (94-96%, Alfa Aesar), and methanol (99%, Alfa Aesar). At the beginning of each trial, a test vial was filled with 50 mL of one of the solvents and a ternary salt tablet was carefully dropped in. Every 3 min for 30 min following, a 1 mL sample was drawn from the test vial and placed into an individual secondary vial. Test vials remained shut (cap on), except when collecting a sample, to minimize solution evaporation to the ambient. Prior to each sample collection, test vials were slowly flipped several times to gently mix the solution. Samples were taken half the distance from the surface of the solution to the bottom of the vial.

The 10 samples from each trial were then diluted with nitric acid (Fisher Scientific, 67-70%) and analyzed for elemental composition (uranium, lithium, and potassium) using inductively coupled mass spectrometry (ICP-MS). The dissolution experiment was conducted at 25, 35, and 50℃. Table 1 provides a summary of ternary salt tablet masses used in this portion of the study.

#### 2.2.2 Visual analysis of salt dissolution

Salt ingots were taken and suspended in a volume (~25 mL) of each solvent (purified water, methanol, and ethanol) to visualize the progression of salt dissolution at room temperature (~25℃). Using a friction anchor to secure the salt ingot to the lid of a cylindrical vial, the salt ingot was lowered into the solvent containing vial. The dissolution was recorded with an iPhone 8 for approximately 2.5 min. Frames of the video were extracted at different times (0, 30, 60, 90, and 120 seconds) and will be shown and discussed later.

### 2.3 Solvent Performance on Uranium Dendrites with Residual Salts

Uranium dendrites were electrochemically separated in 35 g of LiCl-KCl-UCl3 (1wt% U) molten salt electrolyte at 500℃ with an overpotential of 50 mV against a Ag(I)/ Ag reference electrode (molten LiCl-KCl-AgCl (1wt%) in a thinned-end pyrex tube) onto a tungsten rod (working electrode with diameter = 1.5 mm). A molybdenum basket (1 mm wire) housing a depleted uranium chip was used as the counter electrode. An alumina sheathed K-type thermocouple was used to monitor the temperature of the molten salt throughout the uranium electrolysis.

The dendrites were harvested by cutting the tungsten rod just above the electrodeposited uranium (it should be noted that these dendrites were carefully scraped off of the tungsten rod). Fig. 2 shows a harvested uranium dendrite from the described conditions with the adhered LiCl-KCl- UCl3 salt on the outer surface (purple). Based on the coulombs passed during the separation time of approximately 30 min, the theoretical mass of uranium deposited was estimated to be ~60 mg.

The harvested uranium dendrite was separated into four equally sized portions for the different solvent rinsing and drying techniques investigated. It is important to note that the mass of adhered salt was not measured. Table 2 provides a summary of all the combinations of washing and drying investigated in this study, as well as the allotted dissolution and drying times of each case. Each solvent was allowed 25 mL of volume and 12 hours to remove the adhered salt without mechanical or acoustic agitation. For the methanol and ethanol conditions, the dendrites were allowed to dry by natural evaporation of the methanol and ethanol from its surface. However, due to the nature of water, two methods of drying were investigated: (1) drying in the oven for water evaporation, and (2) by taking advantage of the miscible properties between methanol and water, the dendrite rinsed with water was also rinsed with methanol to remove the water from the dendrite surface to quickly evaporate at room temperature.

Upon completion of each rinse, the washed dendrites were secured onto a scanning electron microscopy (SEM) mount for microscopy analysis of the washing technique. Additionally, energy dispersive spectroscopy (EDS) was performed to determine the surface compositions of the washed samples.

## 3. Results and Discussion

### 3.1 Inductively Couple Plasma Mass Spectrometry (ICP-MS)

ICP-MS results for concentrations of uranium, lithium, and potassium results in the deionized water, methanol, and ethanol at different temperatures have been plotted and analyzed. The relationship between relative standard deviation (RSD; a percentage), standard deviation (σ), and the mean (μ) in the formula below (Eq. 1):

$R S D = σ μ ×100$
(1)

In each of the succeeding figures (Fig. 3-5), the average ternary salt mass used in the trials and the respective maximum relative deviations (RSDmax; the largest relative standard deviation observed in the trials) are provided in the legends. Note that extremely large RSDs (>100) found in Fig. 3-5 are likely due to inadequate rinsing between ICP-MS analysis steps causing cross-contamination with adjacent samples; the worst case of this is exhibited in Fig. 5(b) for methanol at the 24-minute mark with an RSDmax of 6,573.5.

The evolution of uranium concentration (CU) during dissolution LiCl-KCl-UCl3 (1wt% U) in the aforementioned solvents at 25, 35, and 50℃ is provided in Fig. 3(a), 3(b), and 3(c) respectively. Upon inspection, the highest achievable concentration of uranium (CU,max) with the allotted experimental time in these solvents was achievable with purified water with corresponding maximum concentrations of 8.33, 10.62 and 11.55 μg·L-1 respective to the 25, 35, and 50℃ conditions. Uranium appears to have achieved a solubility equilibrium limit in purified water for the given conditions as early as 9 min into the trials (Fig. 3(a) and 3(c)) or as late as 15 min into the trials when the change in concentration begins to decrease (Fig. 3(b)).

Unlike purified water, the evolution of CU in methanol and ethanol exhibits behavior indicating the solubility equilibrium has not been achieved; here, CU increases steadily in both solvents throughout the majority of the experiment. Of the two alcohols, methanol was able to achieve a higher CU,max. CU,max recorded for the 25, 35, and 50℃ conditions are 5.67, 5.73, and 6.75 μg·L-1 for methanol and 2.79, 2.50, and 4.73 μg·L-1 for ethanol.

The mass ratio (ξM) evolutions of lithium and potassium are provided in Fig. 4(a)-(c) and Fig. 5(a)-(c) respectively. Here, ξM is considered the mass (m [μg]) ratio of the metal of interest to the maximum uranium value of the corresponding solvent-temperature combination condition, that is:

$ξ M = C M C U , m a x = m M m U , m a x ; M = L i , K$
(2)

Note that the values in Eq. 2 are normalized to demonstrate the behavior of the concentration evolution and the change in analyte uptake in the different solvents in relation to the maximum uranium concentration recorded for that experiment.

The mass ratios of lithium in water, methanol, and ethanol behave similarly to that of the uranium concentration in the three solvents at all conditions. In this case, ξLi exhibits a plateau (indicating an apparent solubility equilibrium) as early as 9 min (Fig. 4(a) and 4(c)) or as late as 15 min (Fig. 4(b)), while in methanol and ethanol, the ratio gradually increases over time throughout the experiment. The maximum mass ratio of lithium (ξLi,max) in methanol and ethanol exceeds that of water for the 35℃ and 50℃ conditions, indicating that as the temperature increases, so does the uptake of lithium (per the maximum uranium taken up in the allotted experimental time) into the solvents and exceeding the rate of purified water. However, the absolute values of the lithium concentrations will indicate that water still dissolved more lithium chloride than methanol and ethanol at all temperatures. For example, at 50℃, the maximum ratio of lithium uptake to the maximum uranium uptake taken in water, methanol, and ethanol is 21, 27, and 23, respectively. The maximum concentrations of uranium for this condition have been provided above, and when multiplied with these ratios, it results in the maximum lithium concentrations of approximately 242.55 μg·L-1 (in water), 182.25 μg·L-1 (in methanol), and 108.79 μg·L-1 (in ethanol).

The mass ratio evolution of potassium (ξK,max) in water exhibits similar behavior to that of lithium and uranium: the indication of an apparent solubility equilibrium is reached at either the 9-minute mark (Fig. 5(a) and 5(c)) or the 15-minute mark (Fig. 5(b)). Yet, there is a difference in behavior when considering the evolution of ξK,max in methanol and ethanol for all three temperatures. At 25℃, ξK,max remains nearly zero for both solvents with a maximum mass ratio (ξK,max,max) of 0.7 and 0.5 for methanol and ethanol, respectively. For methanol at 35℃ and 50℃, ξK,max begins to increase around the 18-minute and 6-minute marks and recorded maximum values of 6 and 13, respectively. For ethanol, ξK,max remains zero for the final two temperature conditions.

The maximum concentrations of uranium (CU,max), mass ratios of lithium (ξLi,max) and potassium (ξK,max,max), and associated RSDs are provided in Table 3. For the allotted ternary salt dissolution time, there is a general increase in the CU,max uptake with temperature for the three solvents. For each temperature investigated, deionized water had the greatest CU,max of the three solvents, followed by methanol, and then ethanol with the least CU,max measured. When comparing the trends of ξLi,max and ξK,max,max, the three solvents exhibit different behaviors. With increasing temperature, ξLi,max and ξK,max,max in deionized water decrease, corresponding to a decreasing trend in absolute concentration of the two alkaline earth metals. Conversely, ξLi,max and ξK,max,max in methanol increase with increasing temperature, corresponding to an increase in absolute concentration. Ethanol does not seem to exhibit any definite trend in maximum mass ratios with increasing temperatures.

Although not a perfect comparison, the lanthanides (Lns) often behave chemically similar to that of the actinides in solution; therefore, their comparison to their chloride solubilities against UCl3 is suitable. In general, lanthanide chlorides studied in water have high solubilities and increase with temperature, even in mixtures with alkali or alkaline earth chlorides [11, 12]. Previous works have also investigated the solubilities of LiCl and KCl in aqueousorganic (e.g., water-methanol/water-ethanol) and organicorganic (e.g., methanol-ethanol) solvents [13-15]. It was found that the solubility of LiCl and KCl would (i) increase with temperature in methanol, (ii) decrease with temperature in ethanol, (iii) increase with water concentration in the aqueous-organic systems, (iv) increase with methanol concentration in organic-organic systems, and (v) decrease with ethanol concentration in organic-organic systems. In addition, the solubility of LiCl would have generally 2-3 orders of magnitude greater than that of KCl in identical conditions (solvent composition, temperature, etc.). The findings in this study support the first listed observations; however, there are discrepancies with (ii) and the observed magnitudes of the solubility. That is, the general trend of varying temperatures in ethanol is inconclusive and the solubility equilibrium of KCl is higher than LiCl in water for all temperatures. Also, the results in this work strongly support observations (iii), (iv), and (v); here, the solubilities (equilibriums) of LiCl and KCl are greatest in water, followed by methanol, and ethanol.

The saturation level of a solution for a given solute is useful information for all chemical processes. It is important to find modified criteria to approximate the degree of saturation of a liquid solution [16]. Shaltry and Phongikaroon [17] used alternative criteria to ensure their experimental dissolution conditions were near zero (i.e., near ideal conditions). Based on Refs. [16] and [17], a similar approach can be used to find the critical molality for when a solute-solvent combination is saturated (Ξ =1); that is,

$Ξ = m s o l u t e V s o l v e n t 10 ( C s a t − C 0 ) = { < 1 ( unsaturated ) 1 ( saturated ) > 1 ( supersaturated )$
(3)

from this aspect, the solute mass (msolute, kg) to be added to the solvent volume (Vsolvent, L) share a ratio (critical molality when Ξ =1) that has units of [kg·L-1]. From our current study, any initial concentration (C0, kg·L-1) of the solute in the solvent prior to a trial commencing can be assumed zero. The solubility limits (Csat, kg·L-1) of LiCl and KCl in methanol, ethanol, and water at 25℃ can be found from Refs. [13], [14] and [15] and converted to [kg·L-1]. It is important to note that we are unaware of any solubility limits reported for UCl3 in water, methanol, or ethanol; therefore, we have used published information of cerium in water as a surrogate for uranium in water (also, there are no reports on solubility limits for CeCl3 in methanol and ethanol) [11]. Table 4 provides the calculated theoretical mass of each salt present in the ternary salt mass tablets used, as well as the solubilities of the salts in the three solvents, and the approximated volume required for a saturated solution of the solute and solvent as a percentage of the solvent volume used in the experiments above (i.e., [Vsolvent /50 mL] × 100%). The approximated volume percentages were rounded to three significant figures.

The approximated volume percentages required for the water trials at 25℃ indicate that there is sufficient solvent to completely dissolve the components of the tablets without needing to consider the competing dissolution rates of the three salts. The amount of solvent required to dissolve the theoretical mass of LiCl is approximately 3- and 4-times that of water in methanol and ethanol, respectively. The amount of solvent required to dissolve the theoretical mass of KCl is approximately 44- and 1510-times that of water in methanol and ethanol, respectively. It should be noted that water (out of three solvent types) was able to completely dissolve the submerged ternary salt tablets used in these experiments; furthermore, there was some residual salt left over from the methanol and ethanol trials. The mass or composition of the remaining salts were not recorded or investigated. The above findings indicate that competing dissolution rates of the three solvents have a higher relative effect on each other in ethanol and methanol than they do in water. Considering the above discussion, further investigation is required to understand the discrepancy with the aforementioned observations as well as the solubility of UCl3 to understand the competing dissolutions of the UCl3, LiCl, and KCl in the ternary salt.

### 3.2 Visualization of LiCl-KCl-UCl3 (1wt% U) Dissolution in Water, Methanol, and Ethanol

The visual progression of the ternary salt dissolution in room temperature water, methanol, and ethanol are provided in Fig. 6, 7, and 8 respectively. In Fig. 6 (deionized water, 25℃), note the reduction in diameter of the salt ingot with time while still exhibiting relatively a similar hue as that at 0 seconds. In conjunction with the concentration evolution of uranium and the mass ratio evolutions of lithium and potassium in deionized water at 25℃ (Fig. 3(a), 4(a), and 5(a)), this indicates that the deionized water dissolves the UCl3, LiCl, and KCl at similar rates. However, there is a noticeable increasing discrepancy between the top and bottom diameters of the submerged portion of the salt ingot as time advances. This is likely due to the dissolution of ternary salt at the top of the ingot affecting the reaction rate of dissolution of the ternary salt at the bottom of the ingot. That is, at 0 seconds, it can be assumed that deionized water does not contain any of the dissolved ternary salt. As time increases, all portions of the ingot dissolve into the deionized water, however, the ternary salt dissolved at the top of the ingot sinks (due to gravity) towards the bottom of the ingot, increasing the ternary salt concentration toward the bottom. This causes a decrease in available water molecules and a decrease in the uranium concentration gradient near the solid-liquid interface, which ultimately affects the reaction rate of dissolution at the lower portion of the ingot. A previous study confirmed the existence of this natural convection phenomenon during the dissolution process of LiCl, SrCl2, CeCl3, LaCl3, PrCl3, and YCl3 salt beads in water [17]. A remedy for difference in dissolution rates would be to submerge the entire ingot into the deionized water horizontally. This would decrease the residency time of dissolved salt near the solid-liquid interface as it sinks, allowing for a more consistent dissolution reaction rate through the length of the ingot and ultimately obtaining a more uniform shape through the duration of the dissolution.

By inspection of Fig. 7 and 8, there is little to no change in the salt ingots diameter within the first 2 min. However, the change of the initial color to an off-white color for methanol and light pink for methanol is an indication that the UCl3 is being dissolved first at the solid-liquid interface, leaving behind the bulk of the ternary salt (LiCl and KCl). Furthermore, the off-white color of the salt ingot at 120 seconds in the methanol trial compared to the light pink color of the salt ingot in the ethanol trial is an indication that methanol takes up the UCl3 much quicker than ethanol. This claim is supported by examining and comparing the concentration evolution of uranium in methanol and ethanol at 25℃ (Fig. 3(a)).

### 3.3 Scanning Electron Microscopy with Energy Dispersive Spectroscopy (SEM-EDS)

Upon completion of washing and drying of samples per Table 2, the uranium dendrites were interrogated using SEM-EDS to obtain microscopic and elemental mapping visualizations on how well the adhered salt was removed using each washing technique and what was the composition of the remaining residual salt. Quantitative results (weight percent [wt%] of the three elements) for each case are provided in Table 5. It is important to note that samples have not been analyzed for lithium due to the inherent limitations of the SEM-EDS device. Therefore, lithium presence and elemental mapping were not discussed or presented in this section with the results. Also, samples were not analyzed for tungsten due to the lack of alloy formation between tungsten and uranium determined in a previous study [18].

In the ethanol washed specimen (Fig. 9), only a small portion of the uranium dendrite is visible, seen in Fig. 9(a) on the left-hand side. Further investigation concluded that the remainder of the structure found in Fig. 9a is mostly chlorine (Fig. 9(b)) and potassium (Fig. 9(c)) with miniscule amounts of uranium (Fig. 9(d)). By comparison, the methanol trial allowed for slightly more visualization of the uranium dendrites morphologies (Fig. 10(a)) and left behind less potassium (Fig. 10(b)) and chlorine (Fig. 10(c)). In both the methanol and ethanol cases, a larger volume of solvent would likely have dissolved the remaining salt on the dendrites per the criteria in Eq. 3 and the approximations in Table 4.

The water washing technique performed exceptionally well at allowing visualization of the uranium dendrite morphologies (Fig. 11(a) and Fig. 11(b)). However, the ovendried method did leave slightly more chlorine in some capacity on the surface of the dendrites than the alternative methanol-rinsing method.

The resulting values in Table 5 provide two key points (while also considering the discussions in this and previous sections). First, of the techniques investigated, the waterbased rinsing technique outperformed the ethanol- and methanol-based techniques when dissolving the adhered LiCl-KCl-UCl3 ternary salt off from the electrochemically separated uranium dendrites. The differences between the two drying techniques used with the water-based rinsing techniques are also minute.

Second, the quantitative EDS results of the ethanol washing techniques shows a ratio of approximately 1:1 for potassium to chlorine. Given that the ratio of potassium to chlorine in KCl is 1:1, the assumption can be made that the ethanol washing technique performed unsatisfactorily at dissolving and removing adhered KCl from the surface of the uranium dendrites.

## 4. Conclusion

Deionized water, methanol, and ethanol have been investigated as possible solvents for the removal of adhered chloride salts from the surface of electrodeposited metals to allow for satisfactory study of samples. LiCl-KCl-UCl3 ternary salt was dissolved in the three solvents at several temperatures (25, 35, and 50℃) over a period of 30 min with the absence of mechanical or acoustic agitation. Samples of the solution were taken periodically (every 3 min). ICPMS analysis of the samples indicated that water can achieve a greater apparent solubility equilibrium of LiCl, KCl, and UCl3 than methanol and ethanol. Saturation criteria used indicate that the water dissolution trials had sufficient solvent for dissolving the component salts of the ternary salt, whereas further investigation into the competing dissolutions of LiCl, KCl, and UCl3 in methanol and ethanol are needed. A photographic depiction of the dissolution the ternary salt in the solvents at 25℃ confirms that water has greater capability at dissolving the ingot in its entirety, whereas methanol and ethanol tend to dissolve the UCl3 at the solid-liquid interface much faster than the LiCl or KCl. Finally, electrochemically separated uranium dendrites were washed with the three solvents at 25℃. It was determined that a water wash in tandem with either a methanol rinse and air dry or an oven dry perform better at removing adhered salt from the surface of the dendrites over a methanol or ethanol rinse.

## Acknowledgement

This study is based upon work supported by the Department of Energy Technology Research & Development Program under award number DE-AC07-05ID14517. The authors would like to thank Dr. James King from Idaho National Laboratory for suggestions and discussions on uranium morphology. The authors would also like to thank Meredith B. Eaheart for her contributions to the experiments discussed in this report.

## Figures

LiCl-KCl-UCl3 salt ingot prepared for salt dissolution experiment.

Electrochemically deposited uranium dendrite separated in LiCl- KCl-UCl3 salt at 500℃ with a 50 mV (vs Ag(I)/Ag) overpotential.

Uranium concentration in the studied solvents at (a) 25℃, (b) 35℃, and (c) 50℃.

Mass ratio evolution of lithium (ξLi) in the studied solvents at (a) 25℃, (b) 35℃, and (c) 50℃.

Mass ratio evolution of potassium (ξK,max) in the studied solvents at (a) 25℃, (b) 35℃, and (c) 50℃.

Visualization of ternary salt in purified water over 120 seconds (25℃).

Visualization of ternary salt in methanol over 120 seconds (25℃).

Visualization of ternary salt in ethanol over 120 seconds (25℃).

Ethanol Wash Microscopy and EDS Elemental Mapping: (a) microscopy and EDS elemental mapping of (b) chlorine, (c) potassium, and (d) uranium.

Methanol Wash Microscopy and EDS Elemental Mapping: (a) microscopy and EDS elemental mapping of (b) chlorine, (c) potassium, and (d) uranium.

Water Wash (Oven-Dry) Microscopy and EDS Elemental Mapping: (a) microscopy and EDS elemental mapping of (b) chlorine, (c) potassium, and (d) uranium.

Water Wash (Methanol-Rinse) Microscopy and EDS Elemental Mapping: (a) microscopy and EDS elemental mapping of (b) chlorine, (c) potassium, and (d) uranium.

## Tables

Ternary salt tablet masses

Experimental conditions and procedures determining performance of rinsing technique on salt-adhered electrochemically separated uranium dendrites

Summary of maximum analyte concentration and analyte mass ratios measured with RSD

Volume percent of experimental volumes used required to completely dissolve theoretical masses of salt from ternary salt tablets at 25℃

Summary of chlorine, potassium, and uranium surface presence after dendrite washing

## References

1. C.H. Lee, T.J. Kim, S. Park, S.J. Lee, S.W. Paek, D.H. Ahn, and S. K. Cho, “Effect of Cathode Material on the Electrorefining of U in LiCl-KCl Molten Salts”, J. Nucl. Mater., 488, 210-214 (2017).
2. S.L. Marshall, L. Redey, G.F. Vandegrift, and D.R. Vissers, Electroformation of Uranium Hemispherical Shells, Argonne National Laboratory Report, ANL-89/26 (1990).
3. K.C. Marsden, B.R. Westphal, M.N. Patterson, and B. Pesic, “Purity of Uranium Product from Electrochemical Recycling of Used Metallic Fuel”, IPRC, Fontana, Wisconsin (2012).
4. K.C. Marsden, “Measurement and Analysis of Uranium and Cerium Depositions from LiCl-KCl Eutectic” Diss., University of Idaho (2015).
5. H. Tang and B. Pesic, “Electrochemistry and the Mechanisms of Nucleation and Growth of Neodymium during Electroreduction from LiCl-KCl Eutectic Salts on Mo Substrate”, J. Nucl. Mater., 458, 37-44 (2015).
6. K. Serrano, P. Taxil, O. Dugne, S. Bouvet, and E. Puech, “Preparation of Uranium by Electrolysis in Chloride Melt”, J. Nucl. Mater., 282(2-3), 137-145 (2000).
7. C.S. Wang, Y. Liu, H. He, F.X. Gao, L.S. Liu, S.W. Chang, J.H. Guo, L. Chang, R.X. Li, and Y.G. Ouyang, “Electrochemical Separation of Uranium and Cerium in Molten LiCl-KCl”, J. Radioanal. Nucl. Chem., 298(1), 581-586 (2013).
8. T.C. Totemeier and R.D. Mariani, “Morphologies of Uranium and Uranium-Zirconium Electrodeposits” J. Nucl. Mater., 250(2-3), 131-146 (1997).
9. H. Tang and B. Pesic, “Electrochemical behavior of LaCl3 and morphology of la deposit on molybdenum substrate in molten LiCl-KCl eutectic salt”, Electrochim. Acta, 119, 120-130 (2014).
10. H. Tang and B. Pesic, “Electrochemistry of ErCl3 and morphology of erbium electrodeposits produced on Mo Substrate in Early Stages of Electrocrystallization from LiCl-KCl Molten Salts”, Electrochim. Acta, 133, 224-232 (2014).
11. T. Mioduski, C. Gumiński, and D. Zeng, “IUPAC-NIST solubility data series. 87. Rare earth metal chlorides in water and aqueous systems. Part 2. Light Lanthanides (Ce-Eu)”, J. Phys. Chem. Ref. Data, 38(2), 441-62 (2009).
12. T. Mioduski, C. Gumiński, and D. Zeng, “IUPACNIST solubility data series. 87. rare earth metal chlorides in water and aqueous systems. Part 3. heavy lanthanides (Gd-Lu)”, J. Phys. Chem. Ref. Data, 38(2), 925-1011 (2009).
13. D. Lide, “Aqueous Solubility of Inorganic Compounds at Various Temperatures”, in: CRC Handb. Chem. Phys. 85th ed., CRC Press, New York (1968).
14. S.P. Pinho and E.A. Macedo, Solubility of NaCl, NaBr, and KCl in water, methanol, ethanol, and their mixed solvents, J. Chem. Eng. Data, 50(1), 29-32 (2005).
15. M. Li, D. Constantinescu, L. Wang, A. Mohs, and J. Gmehling, “Solubilities of NaCl, KCl, LiCl, and LiBr in methanol, ethanol, acetone, and mixed solvents and correlation using the liquac model”, Ind. Eng. Chem. Res., 49(10), 4981-4988 (2010).
16. J.P. Hsu and B.T. Liu, “Dissolution of solid particles in liquids: A reaction-diffusion model”, Colloids Surf., 69(4), 229-238 (1993).
17. M. Shaltry and S. Phongikaroon, “Experimental study of salt bead dissolutions in aqueous solvents”, Ind. Eng. Chem. Res., 53(34), 13550-13556 (2014).
18. C.H. Lee, T.J. Kim, S. Park, S.J. Lee, S.W. Paek, D.H. Ahn, and S.K. Cho, “Effect of cathode material on the electrorefining of U in LiCl-KCl molten salts”, J. Nucl. Mater., 488, 210-214 (2017).
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