Gas–Liquid Reactive Crystallization of Lithium Carbonate by a Falling

Nov 19, 2013 - The gas–liquid reactive crystallization process of CO2 gas and LiOH solution was systematically studied by a falling film column to p...
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Gas−Liquid Reactive Crystallization of Lithium Carbonate by a Falling Film Column Yu-Zhu Sun,* Xing-Fu Song, Miao-Miao Jin, Wang Jin, and Jian-Guo Yu* State Key Laboratory of Chemical Engineering, East China University of Science and Technology, Shanghai 200237, China S Supporting Information *

ABSTRACT: The gas−liquid reactive crystallization process of CO2 gas and LiOH solution was systematically studied by a falling film column to prepare Li2CO3 crystals. Three important parameters concerning the fluid dynamic characteristics of the falling film column were obtained: the Reynolds number, the falling film thickness, and the exposure time. Li2CO3 is the intermediate product during the reactive absorption process, and the final pH of the solution should be controlled within the range of 9.0−9.5 to achieve high product yields. The effects of operating variables on the absorption rate, crystallization rate, and particle size distribution were experimentally explored and theoretically analyzed. Results showed that higher temperatures and LiOH concentrations significantly enhance reactants utilities. A novel method was adopted to quantitatively describe the interaction between absorption and crystallization, and results proved that the crystallization process promotes the transport of CO32− ions from the interfacial liquid film to the bulk solution and, thus, facilitates the overall absorption process. Product characterizations were performed by SEM, XRD, and ICP-MS.

1. INTRODUCTION Lithium carbonate (Li2CO3) is one of the most important lithium salts widely applied in batteries, materials, pharmaceuticals, information industry, and atomic industry.1,2 As such, the synthesis of high purity Li2CO3 has drawn increasing research attention in recent years. Several technologies for the preparation of high-purity Li2CO3, such as recrystallization, precipitation, carbonation, and so on,3 have been explored. Among these methods, the route of gas−liquid reactive crystallization of CO2 gas and LiOH solution is an important method. The structure of a gas−liquid reactor determines the mass transfer process, which, in turn, directly influences the reaction rate and product quality. The various gas−liquid reactors currently available can be divided into three types according to their gas−liquid contact characteristics: gas-bubble reactors, liquid-film reactors, and liquid-drop reactors. Gas−liquid reactive crystallization is frequently performed in a stirred tank or column reactor,4−8 in which the mass-transfer rate is generally slow and the utility of the gas reagent is low. To facilitate gas absorption or control the product morphology, some intensification techniques, such as the three-stage column,9 the Couette−Taylor reactor,10 the microbubble device,11 the membrane contactor,12 reverse micelles,13 the draft-tube reactor,14 HiGee-technology,15,16 and the microreactor,17 among others, have been attempted. Most of the publications relating to gas−liquid reactive crystallization focused on the CO2−Ca(OH)2 system; very few concerned CO2−LiOH system.18 For the gas−liquid reactive crystallization of CO2 gas and LiOH solution, enlarging the gas−liquid interfacial area accelerates the absorption rate and enhances the utility of the gas reagent. But intensifying absorption or controlling the product quality by reducing the diameter of the gas bubbles in a gas-bubble reactor and reducing the diameter of liquid drops in © 2013 American Chemical Society

a liquid-drop reactor is not advisable because Li2CO3 crystals tend to form a very thick fouling layer on the gas or liquid outlet holes.3 This fouling layer chokes the gas or liquid flow. Thus, compared with gas-bubble or liquid-drop reactors, liquidfilm reactors exhibit potential merits. The reaction between CO2 gas and LiOH solution is an obvious exothermic reaction. Rapid heat transfer is also necessary to maintain a stable reaction temperature. Considering these reasons, the falling film technique is more favorable for the present LiOH−CO2 system. The falling film column has been widely used to enhance mass and heat transfer rates during absorption,19,20 reaction,21,22 condensation,23 and evaporation.24,25 However, to the best of the authors’ knowledge, the falling film technique has not been applied in gas−liquid reactive crystallization. The characteristics of reactive absorption have been extensively investigated, focusing on the relationship between the gas−liquid mass transfer and the reaction in the liquid.26−29 Gas−liquid reactive crystallization is a far more complex process than a pure reactive absorption because of the occurrence of crystallization. Unfortunately, limited attention has been paid to the interaction between reactive absorption and reactive crystallization.30,31 Some significant and interesting issues still require further research. For example, the trilateral relationships among physical absorption, gas−liquid reaction, and crystallization have yet to be discussed, whether or not crystallization accelerates or impedes gas absorption is also unknown. In this research, the performance of a falling film column in the gas−liquid reactive crystallization of CO2 gas and LiOH solution was investigated. The impacts of operational variables Received: Revised: Accepted: Published: 17598

August 17, 2013 November 11, 2013 November 19, 2013 November 19, 2013 dx.doi.org/10.1021/ie402698v | Ind. Eng. Chem. Res. 2013, 52, 17598−17606

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(e.g., temperature, LiOH concentration, flow rate of CO2 gas, circulation rate of LiOH solution) on the absorption rate, particle size distribution (PSD), particle shape, and yielding rate were experimentally studied and theoretically analyzed. In particular, the interaction between absorption and crystallization was quantitatively described.

measurement was repeated 3−5 times to obtain an average, and the data scatter was less than 0.5%. 2.3. Procedures. The experiments were carried out in semibatch mode. First, 400 mL of LiOH solution was placed in the slurry tank, and the magnetic stirrer and the water bath were turned on. When the solution temperature had become stable, the CO2 valve was opened, and the gas flow rate was adjusted to a target level. After the air in the absorption column had been replaced by CO2 gas, the pump was started to convey the LiOH solution to the top of the falling column. Under the influence of gravity, the LiOH solution formed a thin liquid film on the wall surface and reacted with the absorbed CO2 to produce Li2CO3 precipitates. The LiOH solution flowed along the vertical wall to the outlet, through to the slurry tank, and was pumped to the top of the falling film column again. The slurry was circulated between the slurry vessel and the falling film column until completion of the carbonation reaction. During an experiment, a pH apparatus (Seven Multi, Mettler Toledo) was employed to monitor the pH value evolution. The slurry samples were withdrawn intermittently with a pipet and then rapidly filtered through a 0.22 μm membrane. The clear solutions were titrated with dilute HCl solution to determine the concentrations of the residual Li+, OH−, HCO3−, and CO32− according to the double-indicator method. After the carbonation reaction, the Li2CO3 slurry was filtered and washed by ethonal to collect the product. Recovered Li2CO3 particles were then dried in a vacuum oven at 70 °C for 6 h. The PSD of the product was analyzed by a Mastersizer 2000 instrument (Malvern). The particle morphology was observed by a scanning electron microscope (SEM; JSM-6360LV, JEOL), and crystals were characterized by X-ray diffraction (XRD; D/MAX 2550 VB/PC, Rigaku). Cation element analysis was conducted by inductively coupled plasma mass spectrometry (ICP-MS; Perkin-Elmer, NexION 300X), and the fluoride ion was determined by ion chromatography (DX 600, Separation column: Ionpac AS11-HC 5*250 mm).

2. EXPERIMENTAL SECTION 2.1. Materials. LiOH·H2O (AR grade) was bought from Sinopharm Chemical Reagent Co., Ltd., and the CO2 gas was purchased from Shanghai Central Industrial Gas Co., Ltd. LiOH·H2O powders were dissolved in deionized water and then filtered by a membrane to prepare the LiOH solutions. 2.2. Equipment. Figure 1 shows a schematic diagram of the experimental apparatus, which consists of a falling film column,

3. RESULTS AND DISCUSSION 3.1. Characteristics of the Fluid Dynamics of the Falling Film Column. Some important fluid dynamics parameters in falling film columns, like Reynolds number, falling film thickness, and exposure time of the falling film, have been studied and correlated by previous researchers (see eqs 1−3).28 The Reynolds number of the falling liquid film is defined in eq 1:

Figure 1. Schematic diagram of falling film adsorption: (1) CO2 gas cylinder; (2) gas flow meter; (3) falling film column; (4) slurry tank; (5) magnetic stirrer; (6) pump; (7) water bath.

a CO2 feeding system, a slurry circulation system, and a thermostat bath. The jacketed glass falling film column was 30 mm in internal diameter and 1000 mm in height. The gas and liquid phases featured paralleled flow. The CO2 gas was fed to the top side of the falling film column through a pipe, and the gas flow was controlled by a flow meter. The LiOH solution was pumped into a distributor located on the top of the falling film column, which helped achieve a homogeneous liquid film on the inner surface of the column. The carbonation of LiOH is an exothermic reaction, so a water bath (DC2006, Shanghai Hengping Apparatus Factory) was adopted to keep the reaction temperature constant with an accuracy of 0.1 °C. A homemade jacketed glass tank with an inner diameter of 80 mm was placed under the falling film column to collect the slurry. A magnetic stirrer was used to suspend the particles in the slurry tank. This vessel was also used in a comparative experiment, which investigated the carbonation rate of LiOH solution in a conventional stirred tank. The density and viscosity values of the LiOH solutions at different temperature were measured by pycnometers and an Ubbelohde viscometer, respectively. Each

Re =

Q lρ π dμ

(1)

where Ql is the liquid flow rate, ρ is the liquid density, d is the internal diameter of the falling film column, and μ is the liquid viscosity. The density and viscosity of the LiOH solutions at different temperatures can be seen in the Supporting Information Figures S1a-b. The Reynolds numbers, all of which fall below 250, are plotted in Figure 2a. This result indicates that all of the falling films exhibit laminar flow under the experimental conditions. The falling film thickness can be calculated according to eq 2: ⎛ 3μQ l ⎞1/3 δf = ⎜ ⎟ ⎝ πgd ⎠ 17599

(2)

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Figure 2. Fluid dynamics characteristics of the falling film column: (a) Reynolds number; (b) falling film thickness; (c) exposure time of the falling film.

Figure 2b describes the impacts of liquid flow rate and solution concentration on the falling film thickness. Increasing the liquid phase circulation rate or solution concentration increases the film thickness, whereas increasing the temperature reduces the falling film thickness. The exposure time of the falling film can be obtained by eq 3: 2/3 1/3 ⎛ 2L ⎞⎛ 3μ ⎞ ⎛ πd ⎞ ⎜ ⎟ te = ⎜ ⎟ ⎜ ⎟ ⎝ 3 ⎠⎝ gρ ⎠ ⎜⎝ Q ⎟⎠ l

(3)

where the L is the length of the falling film column. The dependencies of the exposure time on the experimental conditions are shown in Figure 2c. These results suggest that increases in temperature and liquid flow rate cause apparent decreases in the exposure time, whereas enhancement of the liquid concentration leads to a slight increase in the exposure time. 3.2. Carbonation of the LiOH Solution. The theoretical mechanism of the gas−liquid reactive crystallization between CO2 gas and LiOH solution can be interpreted briefly below on the basis of double film theory (Figure 3). The global process includes three steps: gas−liquid mass transfer, gas−liquid chemical reaction, and reactive crystallization. First, CO2 gas dissolves into the liquid phase. Second, gas−liquid reactions take place in the interfacial liquid film. For example, in the current system, the dissolved CO2 gas reacts with OH− ions in the interfacial liquid film, producing HCO3− ions (eq 4). These HCO3− ions then react with OH− ions and form CO32− ions (eq 5).32 Third, CO32− ions are transferred from the interfacial

Figure 3. Mechanism of gas−liquid reactive crystallization between CO2 gas and the LiOH solution.

liquid film to the bulk solution, where they react with Li+ ions and precipitate Li2CO3 crystals (eq 6). CO2 + OH− → HCO3−

(4)

HCO3− + OH− → CO32 − + H 2O

(5)

CO32 − + 2Li+ → Li 2CO3↓

(6)

Because eq 5 is a very rapid reaction, eqs 4−6 can be integrated with eq 7 to describe the entire gas−liquid reactive crystallization process: CO2 + 2OH− + 2Li+ → Li 2CO3 ↓ + H 2O

(7)

With increasing absorption of CO2, the Li2CO3 crystals in the slurry are gradually bicarbonated,33 as shown in eq 8: 17600

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Figure 4. Carbonation process of the LiOH solution (T = 20 °C, C0 = 2.0 mol/L, Ql = 18 L/h, Qg = 0.5 L/min): (a) ion concentration evolution; (b) PSD evolution.

Figure 5. Effects of operating variables on the carbonation of LiOH: (a) original concentration of LiOH solution (T = 20 °C, Ql = 18 L/h, Qg = 0.5 L/min); (b) circulation rate of LiOH solution (T = 20 °C, C0 = 2.0 mol/L, Qg = 0.5 L/min); (c) flow rate of CO2 (T = 20 °C, C0 = 2.0 mol/L, Ql = 18 L/h); (d) temperature (C0 = 2.0 mol/L, Ql = 18 L/h, Qg = 0.5 L/min).

CO2 + Li 2CO3(s) + H 2O → 2HCO3− + 2Li+

persistently absorbed into the solution and transformed into CO32−; Li2CO3 crystals do not appear until the concentration of Li2CO3 reaches supersolubility. This period is represented by pure reactive absorption without crystallization and can be called the precrystallization period. A drastic increase in solution turbidity indicates the occurrence of nucleation and the beginning of crystallization. The entire absorption process can be divided into the carbonation and bicarbonation stages,

(8)

The absorption process of CO2 into LiOH solution clearly undergoes consecutive reactions in which the objective product Li2CO3 is the intermediate product. As such, the reaction end point must be carefully controlled to achieve high yields, because both insufficient and excessive reactions may cause product loss. During the initial stage of the carbonation process, the solution basically remains clear because though CO2 gas is 17601

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Figure 6. Effects of operating variables on the PSD: (a) original concentration of LiOH solution (T = 20 °C, Ql = 18 L/h, Qg = 0.5 L/min); (b) circulation rate of LiOH solution (T = 20 °C, C0 = 2.0 mol/L, Qg = 0.5 L/min); (c) flow rate of CO2 (T = 20 °C, C0 = 2.0 mol/L, Ql = 18 L/h); (d) temperature (C0 = 2.0 mol/L, Ql = 18 L/h, Qg = 0.5 L/min).

pH = 10 and the pH corresponding to the maximum yield. According to Figure 5a, the final pH in the solution should be controlled within 9.0−9.5. Another experiment under the same conditions (T = 20 °C, C0 = 2.0 mol/L, Ql = 18 L/h, Qg = 0.5 L/min) was performed in a stirred tank for comparison. Results show that achieving the same reaction degree of pH = 9, the carbonation time in the falling film column (26 min) was less than half of that in the conventional stirred reactor (60 min),3 likely because the falling film column enlarges the gas−liquid interfacial area and the rapid flowing film provides a fresh liquid surface for reaction, which greatly enhances the absorption rate. Compared with stirred tanks or column reactors, falling film columns need additional pumps to circulate LiOH solution, but considering the remarkable acceleration in absorption rate and significant increase in reactants utilities, this technique presents its prominent advantages. Falling film columns are also suitable for theoretical researches on absorption processes because of their simple structures and definite gas−liquid contact areas. Figure 4b exhibits the evolution of PSD during the gas− liquid reactive crystallization process shown in Figure 4a. After nucleation, the particles quickly grew. For example, at t = 7.5 min, the volume mean particle size was only 5.72 μm. At t = 10 min, the particle size surged to 41.62 μm, and at t = 20 min, it reached 74.78 μm. Although crystal growth undoubtedly contributes to the increase in particle size, agglomeration plays a more significant role in this process. During the initial stage of crystallization, the solid content is so low that the

and the carbonation stage can be further divided into precrystallization and crystallization stages. Figure 4a shows the typical concentration evolutions of different ions in the solution (T = 20 °C, C0 = 2.0 mol/L, Ql = 18 L/h, Qg = 0.5 L/min). The [Li+] first decreases and then increases because the LiOH solution first reacts with CO2 gas to precipitate Li2CO3, after which Li2CO3 crystals are bicarbonated to soluble LiHCO3. The minimum [Li+] in the solution corresponds to the maximum yield of Li2CO3. Figure 4a also indicates that the pH is the direct reflection of the carbonation process. During the first stage, the pH is basically maintained at 12. At this stage, both the [Li+] and [OH−] obviously decline whereas the [CO32−] increases rapidly and remains at a relatively high level, which indicates a period of stable gas absorption and reactive crystallization. The pH then drops drastically to 10. The [Li+] and [HCO3−] slowly increase with the pH value, and [CO32−] consistently decreases, which indicates the bicarbonation of Li2CO3 crystals. At the end of the bicarbonation, the pH reaches about 7.8. When pH > 10, eq 7 is the dominant reaction and eq 8 can be ignored.28 When pH < 10, eq 2 rapidly becomes the dominant reaction, and the [HCO3−] increases persistently from pH = 10. So pH = 10 can be regarded as the turning point of eqs 7 and 8. It should be noted that the lowest [Li+] or the highest yield of Li2CO3 does not occur at pH = 10, because though bicarbonation begins at pH = 10, the supersaturation of Li2CO3 in the liquid phase is still maintained, so the Li2CO3 crystals will be further produced. Thus, there is a delay between 17602

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Figure 7. Analysis of the interaction between crystallization and absorption: (a) relationship between NA and the [OH−] (C0 = 2.0 mol/L, Ql = 18 L/h, Qg = 0.5 L/min); (b) relationship between NA and S (T = 30 °C, C0 = 2.0 mol/L, Ql = 18 L/h, Qg = 0.5 L/min); (c) comparisons between ra and rc (T = 20 °C, C0 = 2.0 mol/L, Ql = 18 L/h, Qg = 0.5 L/min); (d) comparisons between ra and rc (T = 30 °C, C0 = 2.0 mol/L, Ql = 18 L/h, Qg = 0.5 L/min).

supersaturation, which slows the crystal growth rate and lessens particle agglomeration. As is known, agglomeration can be described as a three-step process, collision of particles, adhesion by weak forces such as van der Waals forces, and solidification by crystal growth. So agglomeration increases with supersaturation. This conclusion has been discussed by a number of previous studies.35−38 3.3.2. Effect of the Circulation Rate of the LiOH Solution. As can be seen in Figure 5b, when the circulation rate of the LiOH solution was decreased, the absorbed CO2 and Li2CO3 product quantities also decreased. Reductions in the circulation rate of the LiOH solution cause corresponding decreases in the liquid surface renewal rate, which, in turn, causes a reduction in the absorption rate. Figure 6b shows that as the circulation rate of the LiOH solution decreased, the particle size increased slightly, because increasing agglomeration may take place under these conditions and the fragmentation of agglomerates becomes weak due to the decrease in turbulence of the solution.39,40 When the circulation rate of the LiOH solution is very slow, a few particles may adhere to the inner wall surface, because the solution flow cannot completely flush down the particles. Of course, when the circulation rate of the LiOH solution is high enough, e.g., 12 L/h, a continuous stable liquid film forms, and the fouling can be avoided in the experiments. 3.3.3. Effect of the Flow Rate of CO2. Figure 5c shows that as the flow rate of CO2 increased from 0.5 to 1.0 L/min, both the absorption and crystallization rates increased and the

collision probability between the particles is low and agglomeration is limited. As the carbonation reaction proceeds, both the solid content and particle size increase, which enhances the collision probability and, thus, facilitates the formation of agglomeration. As for the agglomeration mechanism, orthokinetic agglomeration is the dominant one in this system.34 The SEM images in section 3.5 also confirm this conclusion. 3.3. Effects of Operating Variables on the Carbonation of LiOH. 3.3.1. Effect of the Original Concentration of the LiOH Solution. Figure 5a shows that as the original concentration of the LiOH solution decreased from 2.0 to 1.5 mol/L, both the absorbed CO2 quantity and the Li2CO3 product quantity quickly decreased, because the low reagent concentration exerts a negative impacts on the absorption and crystallization rates. This result suggests that a high LiOH concentration is necessary to obtain a high yield rate. An extreme and interesting experimental phenomenon occurs when the LiOH concentration is adequately low enough. At 0.5 mol/L LiOH, almost no crystals appear during the entire carbonation process and the solution is directly bicarbonated without crystallization, because the supersaturation of Li2CO3 in the bulk solution may never exceed the supersolubility of Li2CO3 due to the low solubility product of [Li+] and [CO32−]. As can be seen in Figure 6a, with the decrease of LiOH concentration, the particle size becomes smaller. Because the low concentration of the LiOH solution leads to low 17603

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continuous decreases in the [OH−]. The subsequent reaction was the bicarbonation process, and the apparent absorption rate during the entire bicarbonation process became further slower according to the calculation results. Figure 7b provides a clue for the emergence of the plateau stage. The formation and existence of the supersaturation of Li2CO3 led to the crystallization process, which may facilitate the absorption rate. To clarify the interaction between absorption and crystallization, the average absorption rate of CO2 (ra) and the average crystallization rate of Li2CO3 (rc) obtained from the same period are plotted in Figure 7c,d. The definitions of ra and rc are shown as eqs 10 and 11, respectively.

carbonation time decreased from 20 to 15 min. This result indicates that higher flow rates of CO2 bring about the earlier emergence of the bicarbonation reaction and may be explained as follows. The interfacial reaction between CO2 and OH− occurs rapidly, and the gas−liquid mass-transfer process is the slower step of the global absorption process. Increasing the gas flow rate can elevate the CO2 supply rate and reduce the gas film thickness, which leads to increases in the mass-transfer and absorption rates. However, increases in the CO2 flow rate cause low reagent gas utility, as evidenced by the observation that despite doubling the CO2 flow rate, the total absorbed CO2 and Li2CO3 quantities remained nearly unchanged. Figure 6c illustrates that the influence of the flow rate of CO2 on the particle size is very slight. 3.3.4. Effect of Temperature. Low temperatures generally favor absorption processes. However, in this system, when the temperature varied from 10 to 40 °C, no difference in the absorbed CO2 quantity was observed (Figure 5d). Two explanations are believed to account for this phenomenon. First, as the temperature increases, the gas−liquid reaction rate increases also. Meanwhile, the liquid viscosity decreases markedly, which effectively reduces the mass-transfer resistance. The most significant benefit of a higher temperature is the remarkable rise in the output of Li2CO3. For example, the output of Li2CO3 at 40 °C is 0.71 mol/L, more than the triple of 0.22 mol/L at 10 °C. Li2CO3 is a special salt, and its solubility presents a trend opposite that of temperature. An increase in temperature reduces the solubility of Li2CO3, enhances the supersaturation level of Li2CO3, and accelerates the crystallization rate. The precrystallization period at T = 10 °C lasted nearly 10 min, whereas that at T = 40 °C lasted only 30 s. Figure 6d shows the impact of the temperature on the particle size. As the temperature varied from 10 to 40 °C, the volume mean particle size grew from 49.60 to 96.29 μm. Analysis indicates that increasing the temperature intensifies the crystallization process and the occurrence of agglomeration because of the high supersaturation and suspension density. The surface energy of the crystals also increases with increasing temperature, which means that crystals can form agglomerates more readily at higher temperature conditions than at lower ones. 3.4. Analysis of the Impact of Crystallization on the Absorption Process. The effect of crystallization on the absorption process is quantitatively described in this section. The apparent absorption rate can be calculated using eq 9:

Na =

ra = rc =

(10)

Δt ΔNLi 2CO3 Δt

(11)

In Figure 7c,d, the x-axis represents the reaction period, and the interval time between two adjacent pairs of bars is 2.5 min. For example, No. 1 represents the 0−2.5 min interval, whereas No. 8 represents the 17.5−20 min interval. The y-axis represents the increment rate. Parts c and d of Figure 7 allow us to draw four important conclusions: First, during the initial stage of carbonation, the ra is far higher than the rc, especially when the temperature is low, such as 20 °C, which indicates that enhancement in absorption rate originates from the gas− liquid reaction. Second, as the carbonation reaction proceeds, the [OH−] decreases and both the gas−liquid reaction and absorption rates drop down. In contrast, Li2CO3 supersaturation and Li 2CO 3 solid contents increase, which accelerates the crystallization rate of Li2CO3. These findings indicate that crystallization plays a significant role in promoting the absorption rate because crystallization facilitates the transport of CO32− ions from the interfacial liquid film to the Li2CO3 particles and speeds up the absorption rate. Third, under some conditions, the crystallization rate of Li2CO3 exceeds the absorption rate of CO2. In this case, the consumption rate of CO32− in the bulk solution is higher than the supply rate of CO2 in the interfacial phase, and crystallization takes the place of the gas−liquid reaction, becoming the fastest step in the global reactive absorption process as the decisive factor in enhancing the absorption rate. Fourth, when the crystallization rate of Li2CO3 surpasses the absorption rate of CO2, it generally causes an obvious increase in the absorption rate of CO2. 3.5. Product characterization. Figure 8 illustrates typical SEM images of the Li2CO3 products. The products show flower-like particles composed of plate-like primary crystals. The Li2CO3 crystals tend to cling together and form agglomerates in a suspension such that even ultrasound

ΔNCO2 πdLΔt

ΔNCO2

(9)

According to classic gas−liquid reaction theory and previous literature, when pH > 10, eq 7 is the dominant step, and the relationship between the apparent absorption rate and [OH−] should present a proportional tendency.28 However, Figure 7a shows that the apparent absorption rate of CO2 does not exhibit a simple linear tendency with the [OH−]. The relationship between the apparent absorption rate and the [OH−] could be divided into three stages: in the first stage, the apparent absorption rate went down drastically with decreasing [OH−]; in the second stage, whereas the [OH−] clearly continued to decline, the apparent absorption rate did not decrease accordingly and instead maintained at a plateau; in the third stage, the apparent absorption rate further decreased with

Figure 8. SEM images of the products (T = 20 °C, C0 = 2.0 mol/L, Ql = 18 L/h, Qg = 0.5 L/min). 17604

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waves still cannot prevent agglomeration.18 Supersaturation is the most important factor influencing the crystallization process and product morphology. The authors’ previous study found that when S is very low (e.g., 8.0), the Li2CO3 product is grape-like agglomerate.3 In this CO2−LiOH system, due to the limitation of the gas-transfer rate, all the S levels are moderate (1.5−3.5), so all the products are flower-like particles. The crystalline phases of the Li2CO3 products were identified using powder XRD. An example of the XRD patterns obtained is shown in Figure 9. All of the diffraction peaks are sharp and

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the bulk solution and promotes the overall absorption rate. 5. Product characterizations show that the products are flower-like particles composed of a number of plate-like primary crystals. XRD patterns present high Li2CO3 crystalline, and the purity of the products meets the requirements of battery grade Li2CO3.

ASSOCIATED CONTENT

S Supporting Information *

The density and viscosity of the LiOH solutions at different temperatures are shown in Figures S1a-b, the corresponding physical properties of pure water are also plotted as a reference (water data from: Wu, D. R. Chemical Process Design Hand Book, 4th ed.; Chemical Industry Press: Beijing, 2009). This information is available free of charge via the Internet at http:// pubs.acs.org.

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AUTHOR INFORMATION

Corresponding Authors

*Y.-Z. Sun: e-mail, [email protected]. *J.-G. Yu: e-mail, [email protected]. Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS This work was financially supported by the National High Technology Research and Development Program of China (863 Program, No. SS2012AA060701), Doctoral Fund of Ministry of Education (No. 20120074120013), Technology Innovation Action Plan of Shanghai (No. 11dz1205502), and the Open Project of State Key Laboratory of Chemical Engineering (No. SKL-ChE-11C03).

Figure 9. Typical XRD characterization of the products (T = 20 °C, C0 = 2.0 mol/L, Ql = 18 L/h, Qg = 0.5 L/min).

narrow, presenting very good agreement with the standard XRD pattern of Li2CO3 (JCPDS Card No. 22-1141). No impurity peak was observed, which indicates the product features very good crystalline Li2CO3. Purity analysis showed that the Li2CO3 contents of the products were above 99.8% and that impurities observed were within the required limits for battery-grade Li2CO3.

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NOMENCLATURE

Symbols

C = concentration, mol·L−1 C0 = original concentration of LiOH solution, mol·L−1 d = diameter of the falling film column, m g = gravitational acceleration, 9.81 Qg = flow rate of CO2 gas, L·min−1 Ql = flow rate of LiOH solution, L·h−1 N = amount of substance, mol·L−1 NA = absorption rate of CO2 gas, mol·m−2·s−1 PG = pressure of CO2 gas, Pa Re = reynold number ra = absorption rate of CO2, mol·L−1·s−1 rc = crystallization rate of Li2CO3, mol·L−1·s−1 S = supersaturation ratio T = temperature, °C t = time, s te = exposure time of the falling film, s

4. CONCLUSIONS The gas−liquid reactive crystallization process of CO2 gas and LiOH solution was systematically studied using a falling film column to prepare Li2CO3 crystals. The main conclusions are as follows: 1. Three important fluid dynamics parameters concerning the falling film column, including the Reynolds numbers, the falling film thickness, and the exposure time of the falling film, were obtained. 2. To obtain high Li2CO3 yields, the final pH of the carbonation reaction should be controlled within the range 9.0−9.5. 3. The effects of operating variables on the absorption rate of CO2, crystallization rate, and particle size were experimentally explored and theoretically explained. Results show that increases in temperature and LiOH concentration lead to significant increases in reactant utility and particle size. Increases in the CO2 gas flow rate facilitate absorption but exert nearly no impact on the particle size. As the circulation flow rate rises up, both the CO2 absorption rate and Li2CO3 yield increase. 4. Under some conditions, crystallization facilitates the transport of CO32− ions from the interfacial liquid film to

Greek Letters

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δf = falling film thickness, m ρ = density, kg·m−3 μ = liquid viscosity, m·Pa·s

REFERENCES

(1) Sun, Y. Z.; Song, X. F.; Wang, J.; Yu, J. G. Unseeded Supersolubility of Lithium Carbonate: Experimental Measurement and Simulation with Mathematical Models. J. Cryst. Growth 2009, 311, 4714. 17605

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dx.doi.org/10.1021/ie402698v | Ind. Eng. Chem. Res. 2013, 52, 17598−17606