1. Introduction

The oxide reduction (OR) process has been the subject of much research aimed at applying the oxide fuels discharged from pressurized water reactors in the metallic fuel pyroprocessing. This metallic fuel pyroprocessing was successfully demonstrated through the Experimental Breeder Reactor-II project in the USA [1]. Once the oxide fuels are converted into their metallic forms by OR operation, U is selectively recovered via an electro-refining process and transuranic (TRU) nuclides are then recovered along with U via an electro-winning process [2]. The OR process normally employs LiCl as a molten salt and is operated at around 650℃ [3-6]. Oxide fuels are loaded into a metallic cathode basket, with platinum normally employed as the anode owing to its high stability under the harsh condition of high temperatures and a corrosive environment (evolution of O₂ gas) [3-6]. However, a recent investigation claimed that repeated OR operations cause noticeable degradation of the platinum anode [7]. Previously, carbon was proposed as an alternative to platinum for the anode [8, 9]. In these reports, the operating voltage was maintained at less than 3.47 V to avoid decomposition of the LiCl salt and the consequent production of chlorine gas at the anode. The reactions employed in this carbon anode system can be summarized as follows:

<Cathode>

<Anode>

Although the carbon anode system showed some promise, the formation of carbonate ions through the following reaction (7) was pointed out as an issue [8, 9].

These carbonate ions are electrochemically decomposed to produce carbon dust at the cathode, according to the following reaction (8), leading to poor process efficiency.

In order to address the carbonate issue, our group recently proposed the use of a carbon anode at a high potential to decompose the LiCl salt itself. Applying higher voltage results in the production of chlorine gas at the anode via the following equation (9):

<Anode>

Here, the reactions mentioned above to produce O₂, CO, and CO₂ at the anode are still involved in the high-voltage operation due to oxygen ions which are liberated from oxide fuels. Hence, the carbonate formation reaction will proceed even during the high-voltage OR operation, while it was predicted that chlorine gas generated at the anode may react with Li₂CO₃ or Li₂O to produce LiCl [10, 11]. The previous work showed that the concentration of Li₂CO₃ remained below 0.3wt% after repeated experiments at high cell voltage levels (above 10 V) [10]. This result can be explained by various chemical reactions that can take place between Li₂CO₃ and Cl₂ along with other chemicals, such as Li, Li₂O, and C [10]. The direct reaction between Li₂CO₃ and Cl₂ (shown below as reaction (10)) was previously introduced by Nakamura et al. [12], in which the effects of the reaction temperature and particle size were investigated. The authors could achieve high proportions of LiCl in the temperature range of 350–500℃ with particle sizes smaller than 0.82 mm. However, detail information about the experimental conditions such as gas flow rate, the amount of reactant, and partial pressure of Cl₂ was not provided in that report [12].

In this work, the reaction above was investigated as functions of the reaction temperature, the total flow rate of Ar and Cl₂ (Q), and the chlorine partial pressure (pCl₂) to identify whether it is a reasonable explanation for the limited accumulation of Li₂CO₃.

2. Experimental

A quartz tube (4 cm diameter) equipped with an electrical furnace in the middle was employed as a reactor. Both ends of the tube were sealed with silicone stoppers connected to Teflon tubes to ensure the flow of gas. The flow rates of the Ar and Cl₂ gases were controlled independently using mass flow controllers (MFCs, Kofloc co., Japan). Li₂CO₃ powder as a reactant was weighed (0.25–1.00 g for each experiment) and then positions in the middle of the quartz tube using an alumina boat. The reactor was purged using Ar gas before the furnace was heated to remove air inside the reactor. After the furnace reached the setting temperature with a ramping rate of 10℃∙min^–1, Cl₂ gas was fed into it by controlling the MFC. The Cl₂ gas flow and the heater were turned off at the end of reaction. The pCl₂ value was derived from the inlet flow rate ratio of Cl₂ over sum of Ar and Cl₂.

The effect of the reaction temperature was investigated by repeating experiments at 300, 400, 500, and 650℃ with 100 mL∙min^–1 of Q and 5.07 kPa of pCl₂ for various durations. The experiments were conducted at 650℃ for 30 min with pCl₂ of 5.07 kPa and various Q (100–300 mL∙min^–1) and initial weight of sample (W _i, 0.25–1.00 g) to study the effect of Q/W_i ratio. The effect of pCl₂ was investigated for various pCl₂ conditions of 2.03, 3.04, 5.07, and 10.1 kPa at 650℃ with 100 mL∙min^–1 of Q and 1.00 g of W _i. The experiments were also conducted for 20 min with 200 mL∙min^–1 of Q and 0.25 g of W _i.

The progress of the chlorination reaction was calculated by measuing the weight change in the sample before and after the reaction. A structural analysis of the reaction products was conducted by means of the X-ray diffraction (XRD, Bruker D8 Advance).

3. Results and discussion

In order to quantify the progress of the reaction, the conversion ratio of Li₂CO₃ to LiCl, α, was defined as follows,

where W_f and W _i represent the final and initial weight of the samples, respectively, and M.W._LiCl and M.W._Li2CO3 correspondingly indicate the molecular weight of LiCl (= 42.39 g·mol^–1) and of Li₂CO₃ (= 73.89 g·mol^–1). Before getting started with chlorination reaction, the stability of Li₂CO₃ under the gas flow and reaction temperature was verified. This was done to confirm that 1) Li₂CO₃ is not blown away by the gas flow and 2) Li₂CO₃ does not react with alumina boat via the following reaction.

After heating Li₂CO₃ under a 300 mL∙min^–1 Ar flow at 650℃ for 2 h, no changes were found in the weight of Li₂CO₃ indicating that the above issues can be ruled out in this study. This outcome is supported by previous works, in which thermal decomposition temperature of Li₂CO₃ was reported to be 1,300℃ [13] and beginning of the decomposition reaction was found at 720℃ [14].

The effects of the reaction temperature on the reaction kinetics were investigated by repeating the experiments at 300, 400, 500, and 650℃, and these results are shown in Fig. 1. It is obvious in the figure that lowering the reaction temperature profoundly diminishes the reaction rate. One interesting outcome was that the reactions did not complete at low temperatures of 300 and 400℃. This result suggests that the diffusion of chlorine atoms within the Li₂CO₃ particle is too slow to achieve complete conversion at this temperature range. According to this result, it is beneficial to operate the OR process at 650℃ with regard to the removal of Li₂CO₃ in the salt. However, the trend was a little bit different in previous work done by Nakamura et al. [12]. The authors could achieve high conversion ratio even at 350℃ and proposed a temperature range 400–506℃ for high conversion ratio. It is interesting that the authors reported an abrupt decrease in the conversion ratio at 550℃. This discrepancy might have come from the different reaction conditions such as reactor design, gas flow rate, and chlorine partial pressure.

Fig. 1

Experimental results measured at reaction temperatures of 300, 400, 500, and 650℃ when the pCl₂ condition is 5.07 kPa and Q is 100 mL∙min⁻¹.

An XRD measurement was taken to confirm the reaction product, and the result is shown in Fig. 2. The peaks could be assigned using the three phases of Li₂CO₃ (JCPDF no. 01-087-0729), LiCl (JCPDF no. 01-074-1972), and LiCl-H₂O (JCPDF no. 01-070-9971). Here, it would be reasonable to assume that the LiCl-H₂O phase was formed during the handling and measurement of the reaction product in air. This result confirms that the reaction between Li₂CO₃ and Cl₂ proceeded without any unexpected by-products and that the assumptions employed during the calculation of α are reasonable. In addition, the results also propose that the direct chlorination reaction using Li₂CO₃ and Cl₂ is a viable means of synthesizing high-purity LiCl.

Fig. 2

XRD measurement results of Li₂CO₃ reacted at 500℃ for 2 h under a 5.07 kPa of pCl₂ and 100 mL∙min⁻¹ of Q condition. Peaks of the pertinent reference materials are also shown below.

The effects of Q/W _i ratio were investigated by repeating identical experiments for various W _i (0.25, 0.50, and 1.00 g) and Q (100, 200, and 300 mL∙min^–1) values, and the results are shown in Fig. 3. The experimental runs were conducted for 30 min while pCl₂ and the reaction temperature were held constant at 5.07 kPa and 650℃, respectively. It is clear in the figure that an increase in the Q/W _i ratio leads to an increase in α, which proves that the chlorination reaction is under control of Cl₂ mass transfer in the gas phase. A change in the Q/W _i ratio should not affect the reaction rate when the reaction is under control of chemical reaction. In order words, under the experimental conditions of this work, the consumption rate of Cl₂ via the reaction with Li₂CO₃ is faster than the transfer rate of Cl₂ at the surface of Li₂CO₃.

Fig. 3

Experimental results obtained after reaction for 30 min. at pCl₂ = 5.07 kPa and 650℃ for various W _i and Q values.

Fig. 4(a) shows the experimental results for various pCl₂ conditions with a constant Q of 100 mL∙min^–1 and W _i of 1.00 g (Q/W _i ratio 100 mL∙min^–1·g^–1). It is clear in the figure that a higher pCl₂ leads to a higher reaction rate as expected from the above results. A linear relationship between pCl₂ and α was found as shown in Fig. 4(b). As discussed above, these results are not available for analysis of reaction kinetics due to dominant influence of Cl₂ mass transfer rate. However, this data is very promising in terms of high-voltage OR operation, as it proves fast chemical reaction between Li₂CO₃ and Cl₂ at the OR operation temperature. In addition, an average utilization value of 37.9% was derived at α = 0.5 proving that Cl₂ is an efficient reactant for this reaction. The utilization ratio of Cl₂ was calculated by dividing the amount of Cl₂ consumed during the reaction by the supplied amount of Cl₂.

Fig. 4

(a) Experimental results obtained at 650℃ for various pCl₂ conditions with Q of 100 mL∙min⁻¹ and W _i of 1.00 g. (b) Linear fitting results derived from the 1 h reaction results of (a).

The effect of pCl₂ was also investigated with a constant Q of 200 mL∙min^–1 and W _i of 0.25 g (Q/W _i ratio 800 mL∙min^–1·g^–1). Fig. 5(a) shows the experimental results with a corresponding linear fitting result. This result indicates that the reaction rate is still under control of chlorine mass transfer. Here, it should be noted that the slope of the linear fitting increased by 27% when the Q/W _i ratio increased 8 times from 100 to 800 mL∙min^–1·g^–1. This outcome reveals that an increase in the Q/W _i ratio does not result in a proportional increase in α, and it means lower Cl₂ utilization at higher Q/W _i ratio. The reaction time–α relationship at Q of 200 mL∙min^–1 and W _i of 0.25 g (Q/W _i ratio 800 mL∙min^–1·g^–1) is shown in Fig. 5(b), which was conducted at 5.07 kPa of pCl₂ and 650℃. It is interesting to observe a linear relationship between the α value and the reaction time over the entire range, because the reaction rate decreased with increasing reaction time at Q/W _i ratio of 100 mL∙min^–1·g^–1 (in Fig. 4(a)). The effect of Q/W _i ratio on α is well displayed in Fig. 5(c), in which the results in 5(b) are shown along with the results of Fig. 4(a) with pCl₂ of 5.07 kPa. This result suggests that the slowing down of the reaction rate at Q/W _i ratio of 100 mL∙min^–1·g^–1 came from insufficient supply of Cl₂ owing to its slow mass transfer rate, whereas this phenomenon was eliminated by increasing the Q/W _i ratio to 800 mL∙min^–1·g^–1. The utilization ratio of Cl₂ at α = 0.5 was 20.0% with Q/W _i ratio of 800 mL∙min^–1·g^–1, which was significantly lower than 37.9% with Q/W _i ratio of 100 mL∙min^–1·g^–1.

Fig. 5

(a) Linear fitting results derived from the 20 min reaction results at 650℃ with Q of 200 mL∙min⁻¹ and W _i of 0.25 g.

(b) Experimental results obtained as a function of reaction time at 650℃ with Q of 200 mL∙min⁻¹ and W _i of 0.25 g.

(c) Comparison of (b) with the results in Fig. 4(a) at an identical pCl₂ of 5.07 kPa.

4. Conclusions

The chemical reaction between Li₂CO₃ and Cl₂ was investigated for various reaction temperatures, Q/W _i ratios, and pCl₂ values. The reaction temperature was identified as a key parameter that determines the maximum α value, and at least 500℃ was required for complete conversion of Li₂CO₃ into LiCl. The relationship between the Q/W _i ratio and α value revealed that the overall reaction rate is under control of Cl₂ mass transfer in the gas at 650℃. The key outcomes of this work show that the reaction between Li₂CO₃ and Cl₂ is fast at 650℃, and this reaction is strongly suggested as a mechanism which takes place at the anode during the high-voltage OR operation to limit the accumulation of Li₂CO₃ in the salt.