Study on Electrolytic Recoveries of Carbonate Salt and Uranium from

Jan 16, 2009 - The concentrations of U were analyzed by a coulometric method using .... However, as the carbonate species were gone out of the anodic ...
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Ind. Eng. Chem. Res. 2009, 48, 2085–2092

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Study on Electrolytic Recoveries of Carbonate Salt and Uranium from a Uranyl Peroxo Carbonato Complex Solution Generated from a Carbonate-Leaching Process Kwang-Wook Kim,* Yeon-Hwa Kim, Se-Yoon Lee, Eil-Hee Lee, Kee-Chan Song, and Kyuseok Song Korea Atomic Energy Research Institute, 150 Deokjin, Yuseong, Daejeon, 305-600, Republic of Korea

This work studied an electrolytic decarbonation by using an electrolytic system consisting of a cation exchange membrane-equipped cell and a gas absorber for simultaneous recoveries of carbonate salt and uranium from a uranium-bearing carbonate solution generated from a uranium leaching process using carbonate and hydrogen peroxide. The system was effective for the recoveries of carbonate salt and uranium as a precipitate of uranium peroxide hydrate from the solution and for a treatment of the residual hydrogen peroxide in the solution. The existence of uranium and hydrogen peroxide in a carbonate solution did not affect the electrolytic recovery characteristics of the carbonate salt from the solution. The behavior of the voltage or current during the electrolysis in the system could be explained by the characteristics of change of the species generated depending on the pH in the anodic and cathodic chambers. In the system, the constant current operation was better than the constant voltage operation from the energy consumption point of view. The electrolytic method used in this work for the recovery of carbonate salt from a carbonate-leaching solution containing uranium was evaluated to be environmentally friendly because it could minimize the generation of secondary wastes. 1. Introduction Carbonate leaching of uranium has been used to extract uranium from its ore.1-3 Recently, a great deal of interest has been shown in the process using oxidative leaching of uranium from spent nuclear fuel in a high alkaline carbonate media instead of an acid media to enhance its safety and economic competitiveness and to minimize the generation of secondary waste streams.4-9 Such processes using carbonate salt such as M2CO3 (M ) Li+, K+, Na+, NH4+) consist of a few steps of dissolution and precipitation of uranium or other elements. If the carbonate salt is not recycled in the processes, an enormous amount of carbonate salt should be accumulated in the processes or be released to the environment, which weakens the advantages and the environmental friendliness of the processes. Therefore, in those processes, a step is required to recycle the carbonate salt used in the processes. To recover the carbonate salt used in those processes, it is necessary to understand the chemical reactions relevant to the processes. When uranium oxide is dissolved in a carbonate solution in such a process, an oxidant is usually used to facilitate the dissolution rate of uranium oxide in the solution. H2O2 is considered to be the most suitable among several candidate oxidants of H2O2, K2S2O8, NaOCl, and so on, because it is a salt-free oxidant and because it can make the uranium oxide have a sufficiently fast dissolution rate and a high uranium solubility in a carbonate solution.2,8,9 Uranium dioxide in a carbonate solution mixed with H2O2 is known to be oxidatively dissolved in the form of a complicated uranyl peroxo-carbonato complex ion as shown in eq 1.9-11 UO2 + xCO32- + yH2O2 + 2zOH- ) [UO2(O2)y(CO3)x]2-2z-2x + 2yH2O + 2e- (1) When the uranyl peroxo-carbonato complex solution is acidified, the carbonate species embedded in the complex ion can * To whom correspondence should be addressed. Tel.: +82 42 868 2044. Fax: +82 42 868 2351. E-mail: [email protected].

be released as CO2 gas according to eq 2, and the residual uranyl peroxo complex ions, UO2(O2)34-, can be precipitated as a uranium peroxide of UO2(O)2 · 4H2O at pH 3 to 4 as shown in eq 3. The solubility of the uranium peroxide is known to be very low.12-14 Therefore, the uranium dissolved in the carbonate solution with H2O2 can be almost completely recovered as a precipitate, when it is acidified. H2CO3(CO2 v ) T HCO3- T HCO3-2 +

(2)

UO2(O2)3 + 8H f UO2(O2) · 4H2O (3) The CO2 gas released from the solution by acidification can be recovered by using a gas absorber which an alkaline solution flows downward within. When acid and alkaline solutions are used directly for the acidification of a uranyl peroxo-carbonato complex solution and for the recovery of the released CO2 gas, large amounts of the acid and alkaline solutions should be used, which are excess to the stochiometric amount of the carbonate salt used in the uranium-leaching step. This eventually results in a generation of a large amount of secondary wastes. For such a carbonate-leaching process to have better environmental friendliness, the carbonate salt used in the process should be recycled as much as possible with a minimal usage of additional chemicals into the system. An electrolytic cell with a cation exchange membrane can generate an alkaline solution and an acidic solution through water split reactions in its respective cathodic and anodic chambers. Such an electrochemical device has attracted a great deal of attention recently for the pH control of a solution because of its advantages such as a minimal generation of secondary wastes, easy operation, remote control, and so on.15,16 When such an electrolytic cell is used, almost a zero-release of waste solution can be accomplished during the acidification of a uranyl peroxo-carbonato complex solution and the recovery of the CO2 gas as a carbonate salt solution without consuming large amounts of acid and alkaline solutions. In this work, an electrolytic system consisting of a cation exchange membrane-equipped cell and a gas absorber was 4-

10.1021/ie800990r CCC: $40.75  2009 American Chemical Society Published on Web 01/16/2009

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Figure 1. Schematic diagram of an apparatus used for electrolytic decarbonation and recycling of a sodium carbonate salt from a uranyl peroxo-carbonato complex solution.

Figure 2. Schematic diagram of chemical and electrochemical reactions occurring in a cell with a cation exchange membrane (CEM) during an electrolytic acidification of uranyl peroxo-carbonato complex solution.

devised to be used for recycling the carbonate salt in a process using carbonate media for separation of uranium or other useful materials from spent nuclear fuel or uranium ore. Its characteristics of electrolytic decarbonation of a carbonate solution and recovery of the released CO2 gas as a carbonate salt solution were investigated with several variables, and then the results were compared to those obtained by using a uranyl peroxo-carbonato complex solution. 2. Experimental A schematic diagram of an electrolytic system used in this work for the electrolytic decarbonation and recycling of a sodium carbonate salt from a uranium-containing carbonate solution is shown in Figure 1. A divided electrolytic cell, which consisted of a cation exchange membrane (Nafion 424), an anode of IrO2, and a cathode of Ti, was used. The size of each electrode was W 2 cm × H 4 cm × T 0.1 cm of a Madras type with a mesh of 6 mm × 12 mm. For the experiments of an electrolytic decarbonation and recycling of the carbonate salt without using uranium, a 0.1 M NaOH solution and a 0.5 M Na2CO3 solution mixed with 0.1 M Na2SO4 were used as an initial catholyte and anolyte, respectively. The Na2SO4 solution in the anolyte was used as a supporting electrolyte to prevent an operational cutoff due to a rapid increase of the cell voltage, which occurred because the conductivity of an anolyte bearing only a carbonate salt decreased rapidly as its electrolytic decarbonation was progressed (see Figure 3). For the experiments using uranium-containing carbonate solution, UO2 powder of 2.27 g was dissolved in a 0.5 M Na2CO3 solution with 1.0 M H2O2 of 100 mL. At that time, the uranium in the solution exists in the form of complicated uranyl peroxo-carbonato complex ions according to eq 1.9-11 The anolyte and catholyte

Figure 3. Changes of the pH and cell voltage (A) and carbonate concentrations (B) with electrolysis time in a system without using a supporting electrolyte in anolyte.

of 100 mL were circulated into their chambers and reservoirs, respectively. A CO2 gas-release line from the anodic chamber was connected to the bottom of a CO2-adsorber through which the catholyte solution was circulated in the direction of from the top to the bottom, as seen in Figure 1. The gas absorber was prepared by packing a Pyrex glass tube of 40 cm in length and 2.5 cm in diameter with silica beads of 1 mm in diameter. The void fraction of the packed column was about 40%. In our preliminary experiments, the column was confirmed to be able to completely absorb the CO2 gas released from the anodic chamber. The pH in each anolyte and catholyte reservoir was recorded online by a pH meter and a data logger controlled by a computer. The solutions in the anodic and cathodic chambers were sampled at regular intervals for the analyses. The carbonate concentration was analyzed by an inorganic carbon analyzer (Shimadzu TOC-V CSH/TNM-1). The concentrations of sulfate, sodium, and hydroxyl ions were analyzed by using an autotitrator and an ion chromatograph (Dionex ICS 90). The concentrations of U were analyzed by a coulometric method using Arsenazo III. H2O2 was analyzed with a reflectoquant (Merck RQflex plus 10). All the reagents used in this work were reagent grade, and they were dissolved, as received, in demineralized water of 18.2 MΩ prepared by a double distillation and one ion exchange (Milli-Q plus). All the experiments were carried out at 25 °C.

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3. Results and Discussion This work investigates the electrolytic characteristics and phenomena during decarbonation, recovery of carbonate salt, removal of residual hydrogen peroxide, and recovery of uranium from a carbonate-leaching solution containing uranium and H2O2 by using an electrolytic system of Figure 1. Therefore, it is necessary to understand the reactions relevant to the carbon species and the electrochemical reactions occurring at the anodic and cathodic chambers in the electrolytic system. Sodium carbonate is a salt of biprotic acid, H2CO3, and bicarbonate is amphoteric, so it acts as a weak acid and as a Brφnsted base, as expressed in eqs 4 to 8, which explain the interrelationship of eq 2. According to the mole fraction changes of carbonate species as a function of pH, which are calculated using the equations, only carbonate ions exist in the solution at pH of more than 12, and CO2 gas is completely released form the solution at pH of less than 4. In the range between pH 4 and pH 8, the carbonate ions, bicarbonate ions, and carbonic acid coexist in the solution in different mole fractions of their species, depending on the solution pH. Kb2

CO3-2 + H2O 98 HCO3 + OH

Ka2

-2 HCO+ H+ 3 98 CO3

Kb2 ) 2.1 × 10-4

Ka2 ) 4.76 × 10-11

Kb1

HCO3 + H2O 98 H2CO3 + OH

Ka1

+ H2CO3 98 HCO3 +H

H2O ) H+ + OH-

(4)

(5)

Kb1 ) 2.2 × 10-8 (6)

Ka1 ) 4.55 × 10-7

KW ) Ka1Kb1 ) Ka2Kb2 ) 10-14

(7) (8)

The water-splitting reactions at the anode and cathode are generally known to change, depending on the solution pH. The evolutions of oxygen and hydrogen are expressed by eqs 9 and 10 in an acidic condition and by eqs 11 and 12 in a basic condition.13 2H2O f O2 + 4H+ + 4e2H+ + 2e- f H2

at cathode in acid

-

-

4OH f O2 + 2H2O + 4e -

at anode in acid

-

2H2O + 2e f H2 + 2OH

(9) (10)

at anode in base

(11)

at cathode in base

(12)

Figure 2 shows the schematic electrochemical reactions occurring in the cell of Figure 1, based on eqs 1 to 12. When a sodium carbonate solution with uranyl peroxo-carbonato complex ions is put into the anodic chamber, proton ions generated in the anodic chamber through the oxygen evolution reaction of eq 11 convert the carbonate species embedded in the uranyl peroxo-carbonato complex ions to CO2 according to eqs 4 and 8. As the CO2 gas is released out of the solution, the remaining cations of Na+ in the anodic chamber should be migrated into the cathodic chamber through the cation exchange membrane to meet the electric neutrality of the anodic solution. On the other hand, in the cathodic chamber, a NaOH solution is generated by the OH- ions generated through the water split reaction of eq 12 and the Na+ ions from the anodic chamber. When the CO2 gas goes into the cathodic chamber with more concentrated NaOH, it is converted to carbonate ions again through eqs 4 to 8. The remaining uranium ions of UO2(O2)34-

Figure 4. Changes of the pH and cell voltage (A), of the carbonate concentrations (B), and of the sodium and hydroxyl ions in the cathode reservoir and of supporting electrolyte ion in the anodic reservoir (C) with electrolysis time.

in the anodic chamber are converted to UO2(O2) · 4H2O of a solid precipitate, as the anodic solution is acidified below pH 4.12-14 As all the carbonate salt ions are moved from the anodic chamber to the cathodic chamber, an osmotic phenomenon takes place between the two chambers so that some of the water in the anodic chamber moves into the cathodic chamber as well. Before investigating the electrolytic recycling of the carbonate salt from the uranyl peroxo-carbonato complex ion solution, which is prepared by dissolving UO2 in a Na2CO3 solution with H2O2, several electrolytic characteristics and phenomena occurring in the cell were investigated in advance by using a carbonate solution without uranium because the electrolytic behavior of the carbonate solution without uranium in the cell was found to be very similar to that by using the uranyl peroxo-carbonato complex ion solution (see Figures 4 and 10).

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Figure 5. Changes of the conductivities of anolyte and catholyte, and the calculated resistivity and voltage due to the electrolytes with electrolysis time.

Figure 3 shows the pH changes in the anodic and cathodic reservoirs and their cell voltage (A) and the carbonate concentrations in the anodic and cathodic reservoirs (B) in the system of Figure 1. In the experiment, the anolyte circulating through the anodic chamber and its reservoir was 0.5 M Na2CO3 without adding any supporting electrolyte into the solution, and the catholyte circulating through the cathodic chamber, its reservoir, and the gas absorber was 0.1 M NaOH. A constant current of 80 mA/cm2 (1.36 A) was applied to the cell. As the electrolysis progresses, the pH in the cathodic chamber increased steadily due to the water split reaction of eq 12 until the time when the first inflection point on the anodic pH appeared and decreased steadily again, then it got finally a steady state. The pH in the anodic reservoir decreased, revealing an inflection point at pH of about 8, by a titration of the carbonate solution with the proton ions produced through the water split reaction of eq 11. According to the calculation of mole fraction of carbonate species depending on pH with using eqs 4 to 8, the mole fraction of bicarbonate in the solution reaches a maximum at pH 8 where the carbonic acid begins to exist in the solution. Below pH 4, bicarbonate almost disappears, and only carbonic acid exists in the solution. Therefore, the total carbonate concentration in the anodic chamber did not apparently change until pH 8. Below pH 8, it decreased rapidly because of the generation of carbonic acid which is released as CO2. On the other hand, the carbonate concentration in the cathodic reservoir, of which solution was circulated within the CO2 absorber, increased inversely. However, as the carbonate species were gone out of the anodic solution, the anodic conductivity rapidly decreased so that the cell voltage increased exponentially and the electrolysis was finally cutoff because of a limit of the voltage compliance of the potentiostat. This result implies that a carbonate solution to be acidified has to contain a suitable supporting electrolyte within the solution, which should not interfere with the decarbonation and its recovery processes. Figure 4 shows the results of an experiment carried out in the same way as Figure 3 except for adding a supporting electrolyte of 0.1 M Na2SO4 into the carbonate solution of 0.5 M Na2CO3. Being different from the results of Figure 3, the cell voltage change became stable again without a rapid shooting only with a small peak at around 125 min. When the pH of the anodic chamber reached about 5, where the cell voltage peak was shown, it decreased rapidly down to less than 2. The whole pH behavior is just the same as that of a typical titration curve

of a carbonate solution by acid. When a 0.5 M Na2CO3 solution of 100 mL is titrated by an acid solution, the amount of the proton ions required for the solution pH to reach the second inflection point at pH 3.2 is 0.1 moles. When a current of 1.36 A is supplied to the anodic solution, the theoretical time taken to generate the same amount of proton ions in the solution is 119 min, which is a little shorter than the experimental time for the anodic solution pH to reach 5 in Figure 4(A) of about 125 min. This means that the total electric charge was required more than the theoretical one. In this case, the electrolytic efficiency to accomplish a decarbonation of the solution is about 95%. At pH of less than 4 in the anodic solution, the carbon concentration became almost zero, because only carbonic acid, which is so unstable that CO2 gas is easily released from the solution like in a soda, exists below pH 4. On the other hand, the carbonate concentration in the cathodic reservoir increased to about 0.42 M in a steady state. As mentioned above, the water in the anodic chamber moves into the cathodic chamber because of an osmotic phenomenon due to the migration of the sodium ions from the anodic chamber. The measured volume change of the cathodic solution during the electrolysis is represented in Figure 4(B). The water transfer starts from the beginning of the electrolysis, and then it stops at the time when the change of the carbonate concentration became a steady state. About 20% of the initial volume of the anodic solution was transferred into the cathodic chamber. In Figure 4, the closed symbols and the open symbols represent the actually measured concentrations and the water transfer-corrected concentrations of carbonate ion, Na+, and OH- in the system, respectively. Taking into consideration the water transfer, the corrected carbonate concentration in the cathodic reservoir at a steady state was evaluated to be the same as the initial carbonate concentration of the anodic reservoir. These results imply that the CO2 released from the anodic chamber can be recovered completely into the cathodic reservoir without any gas slip in the CO2 gas absorber. Figure 4(C) shows the changes of the Na+ and OH- concentrations in the cathodic reservoir and the SO4-2 concentration in the anodic reservoir with and without a correction of water transfer. The respective behaviors of Na+ and OH- were observed to be the same as in the electrolytic system, and their concentration increments are almost twice as much as the change of carbonate concentration in the cathodic reservoir during the electrolysis. These results imply that a 2-fold amount of Na+ to the carbonate transferred from the anodic

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Figure 6. Changes of the pH and cell voltage (A), of the carbonate concentrations (B), and of the sodium and hydroxyl ions in the cathode reservoir in the anodic reservoir (C) at different applied current densities with electrolysis time.

chamber is migrated to the cathodic chamber to maintain the electric neutrality in the solution. The SO4-2 ions were not observed in the cathodic reservoir during the electrolysis. This means that the SO4-2 ions, which were used as a supporting electrolyte in the anolyte, stayed in the anodic reservoir without any transfer to the cathodic reservoir through the cation exchange membrane. To look into the reason for the appearance of a peak at about 125 min during the change of cell voltage where a sharp drop of the anodic pH occurs, the changes of the conductivities of the anolyte and catholyte at 80 mA/cm2 in the same electrolytic system as that used in Figure 4 were measured independently, and the results are shown in Figure 5. The behaviors of the anolyte and catholyte in the system can be explained by the

characteristics of the changes of the species in each chamber of the cell during the electrolysis. The molar conductivities of CO3-2, HCO3-, Na+, H+, and OH- are 69.3, 44.5, 50.1, 315.8, and 198 Ω-1 cm-1 mol-1, respectively. During the acidification of the carbonate solution in the anodic reservoir, as the CO3-2 ions in the anodic reservoir are released out of the solution as CO2 and Na+ ions are migrated to the cathodic chamber, the conductivity of anolyte decreases until the completion of decarbonation, and then it increases again because the H+ ions of a high molar conductivity increase in the anodic reservoir. On the other hand, in the cathodic reservoir, the conductivity increases at the beginning because of the migration of Na+ ions and the generation of OH- ions, but it decreases again until the complete decarbonation in the anodic reservoir at about 125 min because the CO3-2 ions, of which the conductivity is relatively less than that of OH-, are generated. After no more CO2 gas is fed into the cathodic reservoir from the anodic reservoir, the conductivity of the catholyte increases again because OH- ions are steadily generated at the cathode by the water split reaction. The overall solution resistivity of both electrolytes was calculated as a reciprocal of the summation of the individually measured conductivities of the anolyte and catholyte, as shown by the right y-axis in Figure 5. In the experiment of Figure 4, the current applied to the system was 1.36 A, and the distance between the anode and cathode was 0.5 cm. By using such information, the voltage change due to the anolyte and catholyte could be calculated and was plotted as another right y-axis in Figure 5, as well. The calculated voltage pattern is very similar to that of the cell voltage measured directly for the system in Figure 4(A), showing a peak at 125 min. On the basis of these results, the ohmic voltage drop due to the cation exchange membrane was expected to be very low. Figure 6 shows the results of an experiment carried out in the same way as Figure 4 at several applied current densities of 60, 80, and 120 mA/cm2. The times required for the pH of the anolyte to reach 4, where the decarbonation of the solution was completed, became shorter, and the peaks of the cell voltage increased almost in proportion to the applied current. Also, the recovery of CO2 released from the anodic chamber became faster with the applied current, and so did the increase of the Na+ concentration in the cathodic chamber. Figure 7 shows the results of an experiment carried out in the same way as Figure 4 at several applied cell voltages of 4, 6, and 8 V. As the applied cell voltage increases, the times required for the pH of the anolyte to reach 4 become faster as well. The patterns of the measured current are reverse to those of the voltage changes in Figure 6 carried out at the constant currents, and they had minimal peaks around pH 4 as well. At an operation of 4 V, the times for the decarbonation in the anodic chamber and for the OH- concentration in the cathodic chamber to reach a steady state are considerably retarded, compared with the cases of the operations at 6 and 8 V. Figure 8 shows the energy consumed and the ratio of the actually supplied electric charge to the theoretical electric charge for the complete decarbonation in the anodic reservoir at the constant current or constant voltage operations, which are calculated from the results of Figure 6 and Figure 7. The results show that the constant voltage operation consumes more energy than the constant current operation for a decarbonation of the same amount of carbonate salt in the anodic chamber. The inset box shows that all the ratios of the actually supplied electric charge to the theoretical electric charge are over 1, and they increase with the applied current at a constant current operation or with the applied voltage

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Figure 8. Calculated energy consumption and ratio of total supplied electric charge to theoretical electric charge for the complete decarbonation in the anodic reservoir at constant current and constant voltage operations.

9 shows the results of an experiment carried out in the same way as Figure 4 except for using 1.0 M H2O2 in an initial carbonate solution of 0.5 M Na2CO3. All the results of Figure 9 are the same as those of Figure 4 except that the cell voltage only at the beginning until about 60 min was lower than that without using H2O2 in the carbonate solution. This is ascribed to a dissociation of H2O2 to perhydroxyl ions (HO2-) in alkaline condition as shown in eq 13 so that the conductivity of the carbonate solution with H2O2 is higher when compared with the same carbonates solution without H2O2. H2O2 ) HO2- + H+

(13)

H2O2 is decomposed into water and oxygen in alkaline conditions, but it is sable in acidic conditions. In Figure 9(C), the H2O2 in the anodic reservoir was observed to be completely decomposed to be zero concentration even in acidic conditions after 125 min. The decomposition of H2O2 is considered to be due to its oxidation at the anode as shown in eq 14. H2O2 ) O2 + 2H+ + 2e-

Figure 7. Changes of the pH and cell voltage (A), of the carbonate concentrations (B), and of the sodium and hydroxyl ions in the cathode reservoir in the anodic reservoir (C) at different applied cell voltages with electrolysis time.

at a constant voltage operation. This is considered to be because the concentration of the H+ ions in the anodic reservoir becomes higher at higher current or voltage operations so that more H+ ions are migrated more to the cathodic chamber through the cation exchange. The migration of ions is known to be proportional to the ion concentration and voltage difference between the anode and cathode.15,16 This means that the current efficiency of the electrolytic system becomes lower from 100% at the higher current or voltage operations. Because this work investigated the electrolytic characteristics during a recycling of carbonate salt as well as a recovery of uranium from a uranyl peroxo-carbonato complex ion solution, which is generated from an oxidative uranium-leaching by H2O2 in a carbonate solution as mentioned above, it is necessary to observe the effect of H2O2 on the electrolytic decarbonation and recovery of the carbonate salt in the system of Figure 1. Figure

(14)

Figure 10 shows a comparison of the experiments carried out at 80 mA/cm2 with a 0.5 M Na2CO3 solution and a uranyl peroxo-carbonato complex solution which was prepared by dissolving UO2 in a carbonate solution of 0.5 M Na2CO3 containing 1.0 M H2O2, of which the uranium concentration was 2 g/L. The changes of the cell voltage and pH in the anodic and cathodic reservoirs in both cases were almost similar until the pH in the anodic reservoir reached about 4 where a cell voltage peak appeared and fluctuations of pH in the anodic and cathodic reservoirs began. As the uranyl peroxo-carbonato complex solution was acidified in the anodic reservoir, the carbonate species of the complex ion were released as CO2 from the solution, and the residual uranyl peroxo complex ions in the solution, UO2(O2)34-, were observed to be almost completely precipitated as a uranium peroxide of UO2(O)2 · 4H2O below pH 3.12-14 The chemical form of the uranium precipitate was confirmed by an XRD analysis, but the result is not presented in this work. The uranium precipitate in the anodic chamber caused the fluctuations in the anodic and cathodic pH curves and a rapid increase of cell voltage during the precipitation. When the pH in the anodic reservoir reached 2, the electrolysis was turned off because the decarbonation was completed already, and most of the uranium in the solution was precipi-

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Figure 10. Comparison of the changes of pH in anodic and cathodic reservoirs and cell voltage in the cases of using a carbonate solution with and without bearing uranium (A), and change of the concentrations of carbonate, H2O2, uranium in the anodic reservoir (B) with electrolysis time.

Figure 9. Comparison of the changes of the pH and cell voltage (A), of carbonate concentrations (B) in anodic and cathodic reservoirs in the cases of using a carbonate solution with and without 1.0 M H2O2, and change of the concentration H2O2 in the anodic reservoir (C) with electrolysis time.

tated. The decarbonation in the anodic reservoir and the recovery of Na2CO3 in the cathodic reservoir by the CO2 absorber were accomplished very similarly to the case of using only the carbonate salt in Figure 4. After a filtration of the uranium precipitate in the anodic reservoir, the residual uranium concentration in the solution was a few parts per million, so the recovery yield of the uranium from the initial carbonate solution was more than 99%. Only a negligible amount of uranium was found in the cathodic reservoir, which means that the uranyl ions in the anodic chamber during the acidification were difficult to migrate into the cathodic side through the cation exchange membrane. In this work, no metal ions with a higher valance than +1 were confirmed to be transferred through the cation exchange membrane. However, these results are not presented

in this paper. The change of the H2O2 concentration in the anodic chamber was very similar to the result of the experiment using only Na2CO3 and H2O2 without uranium in Figure 9(C). On the basis of all the above results, the electrolytic system of Figure 1 can be said to be effectively used for the recoveries of carbonate salt and uranium from a uranium-containing carbonate solution coming from a uranium-leaching process by using a carbonate solution and H2O2 without putting any chemicals into the system, as the residual H2O2 in the solution was decomposed simultaneously, which results in a minimal generation of secondary wastes from the system. The residual Na+ and SO4-2 ions remained in the solution after the filtration of uranium precipitate, which was used as a supporting electrolyte for the anolyte, could be recovered by using an electrodialysis cell. 4. Conclusions An electrolytic system consisting of a cation exchange membrane-equipped cell and a gas absorber, which was suggested in this work, could be effectively used for the recovery of carbonate salt from the solution produced in a uranium-leaching process using carbonate and hydrogen peroxide, where the uranium could be simultaneously recovered as a uranium peroxide hydrate with a high yield and residual hydrogen peroxide in the solution was gotten rid of from the solution as well. The existence of uranium and hydrogen peroxide in a carbonate solution did not affect the electrolytic recovery characteristics of carbonate salt from the solution itself. The behavior of the voltage or current

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during the electrolysis in the system could be explained by the characteristics of change of carbonate species depending on the pH in the anodic and cathodic chambers. In the system, the constant current operation was better than the constant voltage operation from the energy consumption point of view. The electrolytic method used in this work for the recovery of carbonate salt from a carbonate-leaching solution containing uranium was evaluated to be environmentally friendly because it could minimize the generation of secondary wastes. Literature Cited (1) Mason, C. F. V.; Turney, W. R. J. R.; Thomson, B. M.; Lu, N.; Longmire, P. A.; Chisholm-Brause, C. J. Carbonate leaching uranium of uranium contaminated soils. EnViron. Sci. Technol. 1997, 31, 2707. (2) Hiskey, J. B. Hydrogen peroxide leaching of uranium in carbonate solutions. Trans. Instn. Min. Metall. 1980, 89, C145. (3) Merritt, R. C. The extractive metallurgy of uranium; Colorado School of Mines Research Institute, 1971. (4) Pepper, S. M.; Brodnax, L. F.; Field, S. E.; Zehnder, R. A.; Valdez, S. N.; Runde, W. H. Kinetic study of the oxidation dissolution of UO2 in aqueous carbonate media. Ind. Eng. Chem. Res. 2004, 43, 8188–8193. (5) Asanunma, N.; Harada, M.; Ikeda, Y.; Tomiyasu, H. New approach to nuclear reprocessing in a non-aqueous solution. Proc. Global 2001; Paris: France; September 9-13, American Nuclear Society, 2001. (6) Ohyama, K.; Nomura, K.; Washiya, T.; Tayama, T.; Yano, K.; Shibata, A.; Komaki, J.; Chikazawa, T.; Kikuchi, T. Development of Uranium Crystallization System in NEXT Reprocessing Process. Proc. Global 2007; Boise: Idaho: U.S.A; September 9-13, American Nuclear Society, 2007. (7) Goff, G. S.; Taw, F. L.; Peper, S. M.; Brodnax, L. F.; Runde, W. H. Uranium, Neptunium, and Plutonium complexes in aqueous hydrogen peroxide-carbonate solution; Stability and Speciation. ACS Northwest Regional Meeting; Boise: ID: U.S.A; June 17-20, American Chemical Society, 2007.

(8) Goff, S.; Taw, F. L.; Peper, S. M.; Brodnax, L. F.; Field, S. E.; Runde, W. H. Separation of uranium from fission product in spent nuclear fuel using aqueous H2O2-carbonate solutions. AIChE 2006 annual meeting; San Fransisco; November 12-17, 2006. (9) Kim, K.-W.; Chung, D.-Y.; Yang, H.-B.; Lim, J.-K.; Lee, E.-H.; Song, K.-C.; Song, K.-S. A conceptual process study for recovery of uranium alone from spent nuclear fuel by using high alkaline carbonate media. Nucl. Technol. 2008, in print. (10) Goldik, J. S.; Noel, J. J.; Shoesmith, D. W. Surface electrochemistry of UO2 in dilute alkaline hydrogen peroxide solutions; part II. Effect of carbonate ions. Electrochim. Acta 2006, 51, 3278. (11) Peper, S. M.; Brodnax, L. F.; Field, S. E.; Zehnder, R. A.; Valdez, S. N.; Runde, W. H. Kinetics study of the oxidative dissolution of UO2 in aqueous carbonate media. Ind. Eng. Chem. Res. 2004, 43, 8188. (12) Kubatko, K.-A. H.; Helean, K. B.; Navrotsky, A.; Burns, P. C. Stability of peroxide-containing uranyl material. Science 2003, 302, 1191. (13) Debets, P. C. X-ray diffraction data on hydrated uranium peroxide. J. Inorg. Nucl. Chem. 1963, 53, 727. (14) Djogic, R.; Cuculic, V.; Brainica, M. Precipitation of uranium (VI) peroxide (UO4) in sodium perchlorate solution. Croat. Chem. Acta 2005, 78 (4), 575. (15) Kim, K.-W.; Kim, I.-T.; Park, G.-I.; Lee, E.-H. Development of a continuous electrolytic system with discharging only one pH-controlled Stream and its Characteristics. Water Res. 2007, 41, 303. (16) Kim, K.-W.; Kim, I.-T.; Park, G.-I.; Park, H.-S.; Lee, E.-H.; Kim, E.-H. Electrochemical de-chlorination/transformation of metal chloride for the preparation of NZP structure product. J. Radioanal. Nucl. Chem. 2008, 275 (3), 595.

ReceiVed for reView June 25, 2008 ReVised manuscript receiVed November 4, 2008 Accepted November 9, 2008 IE800990R