The advent of lithium (Li)-ion technology in 1991 was marked by the commercialization of the first LiB, and this signified the beginning of a transformative epoch in technology [1,2]. The notable advantages of Li in batteries compared to other metals are its high electrochemical potential and energy density [3]. These compelling benefits have resulted in LiBs becoming indispensable constituents of portable electronic devices.

The introduction of LiB technology resulted in an unprecedented increase in its use, which was initially fueled by the adoption of portable electronics (such as mobile phones, computers, smartphones, and tablets) [4]; however, the scope of the technology broadened as it matured. By 2010, LiBs had permeated the power tool market and further expanded into the realm of electric vehicles. However, despite the diversity of their applications, portable electronics remained the dominant consumer of LiBs until the early-2010s [5], and the pattern of LiB consumption has since changed dramatically. Once predominantly consumed by portable electronics, the demand landscape has expanded and diversified. A significant amount of usage now relates to larger applications, such as electric vehicles and grid-scale energy storage systems [6]. This substantial shift indicates a transition towards large-scale energy storage solutions, and it underscores the influential role of LiBs in the evolving landscape of sustainable energy technologies.

From a mere few thousand units in 2010, the stock of electric vehicles (EVs) had increased dramatically to 11.3 million by 2020, and projections indicate a burgeoning fleet of 142 million EVs on roads by 2030 [7]. In parallel with this growth, the global production of Li tripled between 2010 and 2020 [8]. Based on these projections, the demand for Li is expected to reach 2 million tons per year by 2030 [9], and current projections estimate an 18- to 20-fold increase in the demand for Li by 2050, if existing extraction policies persist. However, the use of novel and more sustainable extraction strategies could potentially escalate this demand 40-fold by 2050 [10].

With the development of lithium secondary batteries, Li has become an essential resource for modern technology, especially in the form of Li carbonate (Li₂CO₃). This compound is the main source of Li, and it accounts for most of the world’s Li production (~60%) [11,12]. The demand for LiBs alone is expected to increase the production of Li₂CO₃ to a scale of 2 million metric tons annually by 2030 [13]. Li₂CO₃ is typically obtained from the hard rock mining of spodumene ore and from Li concentrated in the salt lake brine [14]. In hard rock mining, spodumene is chemically processed to liberate Li ions (Li⁺), which are then precipitated as Li₂CO₃, primarily through the addition of sodium carbonate (Na₂CO₃) [15–17]. Na₂CO₃ is also commonly employed as the carbonate source to recover Li from the salt lake brine via a series of steps involving the evaporation of brine, lithium concentration, and inducing Li₂CO₃ precipitation [18–20].

It is acknowledged that the accumulation of atmospheric carbon dioxide (CO₂) is a serious threat to the planet [21], and the Paris Agreement in 2015 established long-term temperature goals aimed to control average global warming at 2°C and to limit it to 1.5°C by the end of the 21st century [22]. In this respect, the rapidly growing use of the LiB is recognized as a carbon reduction technology that can ameliorate the problems associated with increases in atmospheric CO₂ caused by decades of industrialization and urbanization.

However, it is essential to consider the lifecycle of these batteries to fully understand their environmental impact. Notably, the production and recycling of Li, a crucial component of LiBs, can be a significant source of CO₂ emissions [23]. For example, an EV battery with an energy density of approximately 78 kWh/kg produces an estimated 172–196 kg CO₂ eq/kWh through raw material extraction, cell manufacturing, and assembly [24,25]. Therefore, although LiBs are currently promoted as tools to achieve carbon neutrality, their production processes indirectly contribute to CO₂ emissions.

To realize a circular economy and a sustainable Li production and recycling cycle, it is necessary to explore methods to mitigate these CO₂ emissions. One promising approach involves the use of CO₂ for Li precipitation. CO₂ is an abundant greenhouse gas, and utilizing it for Li precipitation could provide dual benefits where CO₂ emissions associated with Li production are mitigated and turned into a valuable asset [18]. This method has shown promise in several studies, and it represents a significant step toward a more sustainable and eco-friendly Li production process. CO₂ is typically regarded as a waste product, but it could be transformed into a valuable asset for Li production, thereby contributing to a sustainable and circular economy [26–27]. In this context, it has become increasingly vital to explore and understand Li₂CO₃ precipitation methods while also investigating innovative techniques that could potentially revolutionize the industry and result in a more sustainable and low-carbon society [28]. In addition to the CO₂ reduction benefits, the production of Li₂CO₃ using CO₂ is expected to be beneficial in terms of Li₂CO₃ purity [29].

In this review, we investigate and report the precipitation of Li₂CO₃ using CO₂ as the carbonate source. Until recently, Na₂CO₃ has been used as the carbonate source for Li precipitation; however, CO₂ could be a viable substitute, and carbonate mineralization is a promising carbon capture utilization technology that can be applied to the production of Li₂CO₃.

This study aims to overcome the existing limitations of Li precipitation methodologies by investigating the use of CO₂ gas for Li precipitation. Leveraging CO₂ for this purpose could potentially create a greener and more efficient approach for extracting Li, which aligns with the broader goals of advancing towards a low-carbon and sustainable future (Fig. 1). Despite its importance, there have been limited systematic reviews of Li₂CO₃ production using CO₂. To the best of my knowledge, there has yet to be a comprehensive review paper addressing this specific topic. By analyzing the mechanics, methods, and environmental implications of using CO₂ gas for Li precipitation, this review aims to determine its potential as a viable alternative to the traditional use of Na₂CO₃.

This section focuses on the Li₂CO₃ precipitation processes used to extract and recycle Li and the prevailing reliance on Na₂CO₃ as the primary agent.

An Li-rich solution and a carbonate source are essential for all Li₂CO₃ production processes [11]. The Li-rich solution is prepared using the following methods in these three processes:

First, in hard rock mining, extracted spodumene ore undergoes a series of processing stages. It is first crushed and heated in a kiln to produce beta-spodumene, which is more reactive and amenable to Li extraction. The beta-spodumene is then subjected to acid roast using sulfuric acid to generate an Li sulfate (Li₂SO₄) solution [15,16].

Second, to recover Li from brine, an Li-rich solution is first prepared by concentrating the brine and separating the Li. Li salt-rich brines are prepared by extracting Li from underground reservoirs and concentrating the Li content through solar evaporation in large ponds. Concentrated brine can be pretreated with an Li-selective precipitant, such as phosphate, to produce a primary precipitate, which is re-dissolved to prepare an Li-rich solution [18,19].

Lastly, when recycling Li from spent LiBs, an Li-rich solution is prepared by acid leaching after shredding the spent battery, and this is considered to be a hydrometallurgical process [30]. Pyrometallurgy is also used as a recycling process for spent batteries; however, this method has now been limited for use in Li extraction due to its high energy consumption [31] and the potential loss of Li during the process [32].

A carbonate source is then prepared using Na₂CO₃. In Li extraction and recycling methodologies, Na₂CO₃ plays a pivotal role in the precipitation of Li from Li-enriched solutions [11], and the chemical process is described by the following equation:

Through this reaction, Li⁺ in the solution combine with sodium carbonate (Na₂CO₃) to produce Li₂CO₃. The Li₂CO₃ precipitates owing to its low solubility in water, which enables its separation from the rest of the solution [33].

However, the precipitation of Li₂CO₃ using Na₂CO₃ has notable drawbacks. One significant disadvantage is the production of sodium- based by-products, which require further processing or disposal [34]. This can lead to additional refining steps that add to the process costs. Furthermore, the use of Na₂CO₃ does not contribute to carbon capture or reduction strategies; therefore, exploring alternative methods, such as the use of CO₂ as a carbonate source for Li precipitation, would provide both economic and environmental advantages [26, 27].

The precipitation of Li₂CO₃ using CO₂ gas is gaining attention from the Li production and recycling industry. The process essentially involves the reaction between the Li⁺ present in an aqueous solution and CO₂, which results in Li₂CO₃.

As shown in Fig. 2, the precipitation of Li₂CO₃ is completed in three steps, and these are outlined as follows:

First, CO₂ gas is dissolved in water to become carbonic acid (H₂CO₃), which lowers the pH of the solution:

However, if the pH is below 6.3, H₂CO₃ is the dominant carbonate species, which is theoretically unfavorable for Li₂CO₃ precipitation. When the pH exceeds 6.3, bicarbonate is the dominant carbonate species. When it reacts with Li, it exists as Li bicarbonate (LiHCO₃), which is highly soluble and ionizable.

Therefore, in most cases Li and bicarbonate ions exist in water systems.

Finally, when the pH exceeds 10.3, carbonates are the dominant carbonate species:

According to previous studies, in practice, the production of Li₂CO₃ precipitation occurs at pH 8 and above [34,35], depending on the concentrations of Li⁺ and carbonate.

Although a chemical understanding of the reaction is relatively straightforward, applying the process requires optimization of these reaction parameters. Various factors (such as temperature, pH, and the Li⁺ concentration of the solution) significantly influence the yield and efficiency of the reaction [36]; these factors are described in detail below:

First, precipitation is affected by temperature. the solubility of Li₂CO₃ shifts with increasing temperature and decreases with increasing temperature. Quantitatively, at room temperature (20°C), Li₂CO₃ has a solubility of approximately 0.18 mol/L, but this decreases to approximately 0.1 mol/L at higher temperatures of around 90°C. This decrease in solubility shows a linear correlation with temperature [37], which highlights the significance of temperature for the effective precipitation of Li₂CO₃ [38].

Second, pH strongly influences the carbonation process. For instance, the carbonate species in the solution may be bicarbonate or carbonate depending on the pH (refer to Eq. (5)). Considering that LiHCO₃ is approximately seven times more soluble in water than Li₂CO₃ [39], it can be appeared that the Li₂CO₃ precipitation is affected by the pH. On the other hand, the pH of a solution affects the solubility of CO₂, and the precipitation reaction occurs only when CO₂ dissolved in the liquid phase actively participates in the carbonation reaction. The reaction sequence includes the following steps:

A critical aspect of this process is maintaining a high pH, which is typically achieved by introducing sufficient amounts of hydroxide ion. Despite the theoretical considerations favoring a pH above 10.33 and 8.95 for the formation of CO₃²⁻ in freshwater and seawater, respectively [40], practical observations show that a pH above 8 is required for effective carbonate ion formation. This discrepancy may arise from various factors under real experimental conditions, including temperature and the presence of other ions.

Third, the process is influenced by the concentrations of the Li⁺ in the solution. Assuming that the supply of CO₂ is sufficiently large, the concentration of Li⁺ determines whether supersaturation occurs and affects the precipitation reaction. The supersaturation ratio, calculated as the solubility-to-concentration ratio, must exceed 1 to induce the formation of Li₂CO₃ precipitates [41]. The supersaturation ratio affects both the nucleation and growth rates of the precipitates. A high supersaturation ratio is preferred when a small particle size is desired because it leads to a higher nucleation rate relative to the growth rate, and consequently to a smaller particle size [42]. The purity of the solution plays an important role in the success of this process, and certain ions in the reaction mixture can interfere with the formation of Li₂CO₃. For example, the presence of sodium or potassium ions can lead to the formation of their bicarbonates, which may compete with Li for CO₂ and reduce the overall yield of Li₂CO₃. In addition, the presence of divalent ions, such as calcium or magnesium, may lead to the formation of insoluble carbonates, which can potentially cause scaling issues in industrial setups [43].

In addition, the reaction between Li and CO₂ can be promoted by increasing CO₂ partial pressure [44], and it can also be influenced by factors such as the stirring rpm and CO₂ feed rate. Therefore, it is important to optimize the reaction conditions to achieve the desired Li₂CO₃ precipitation. The following sections present information from various studies focusing on the production of Li₂CO₃ using CO₂.

In this section, previous studies comparing the precipitation of Li₂CO₃ with CO₂ and that of Li₂CO₃ with Na₂CO₃ are presented. Establishing this comparison provides a foundation for analyzing and assessing the innovations and potentials of using CO₂-based processes. The analysis aims not to resolve or identify the challenges with each method but to offer a comparative perspective of Li₂CO₃ precipitation methodologies.

Han et al. conducted experiments on the semi-batch precipitation of Li₂CO₃ using two distinct methods: precipitation using CO₂ (heterogeneous) and precipitation using an aqueous Na₂CO₃ solution (homogeneous) [45]. For both methods, the shape of the primary crystals varied slightly depending on the temperature: ellipsoid shapes were observed at 50°C, while elongated sheets were formed at 25°C. In precipitation using CO₂, the final pH value played an important role, and it was ascertained that a level of 8 was required to halt the CO₂ input. A higher rotation speed resulted in smaller particle sizes because the smaller bubbles accelerated the chemical reaction. Particles were larger at a higher temperature, and a doubled crystal yield was observed at 50°C. However, the gas feed rate did not significantly influence the particle size.

In the case of precipitation using an aqueous Na₂CO₃ solution, the particle size was also smaller at higher rotational speeds, but the pumping rate of the Na₂CO₃ solution made no notable impact. The crystal yield was higher than that obtained by precipitation using CO₂, because a higher concentration of CO₃²⁻ was provided by the Na₂CO₃ saturated solution. A high alkaline concentration was required to precipitate Li₂CO₃ from the Li₂SO₄ solution. The use of solid Na₂CO₃ improved the recovery of Li without increasing the overall solution volume. It was considered that the residues from both processes could potentially be used to precipitate other salts, such as Li₃PO₄. The study concluded that both methods were viable for recovering Li as Li₂CO₃ salt from a Li₂SO₄ solution; however, it was acknowledged that additional research is needed to enhance the yield and crystal purity.

In a comprehensive analysis conducted by Battaglia et al., the precipitation performance of Li₂CO₃ in terms of recovery and purity was examined using both CO₂ and Na₂CO₃, as indicated in Table 2 [34]. In their approach, an aqueous Li solution derived from dilute Li-rich brine was used for CO₂ precipitation. The pH of this solution was meticulously regulated at 8.5 or higher through the addition of NaOH. From their findings, the CO₂ precipitation method was found to perform well in terms of recovery and purity in all comparisons, although a lower recovery rate was observed in the reference case. In particular, the authors found that CO₂ precipitation was very effective in removing Mg, as shown in Table 2. Furthermore, by incorporating an ethanol washing step, they successfully obtained Li₂CO₃ with a near purity of 100%. These results underscore the potential of using CO₂ as a precipitant to obtain high-purity Li₂CO₃.

Although the CO₂ precipitation method for Li₂CO₃ extraction has not been as extensively studied as the Na₂CO₃ approach, it shows promise for optimization. Although many studies have reported higher recovery rates with Na₂CO₃, there is a potential to enhance the novel CO₂ method through better temperature and pH management, which suggests its potential to match, or even surpass the efficiency of the conventional method. The following sections focus on research efforts that have sought to enhance the performance of the CO₂ method by integrating various novel methods.

This section introduces the technique used to feed CO₂ gas by bubbling it directly into an aqueous Li solution to precipitate Li. This method provides a substantial advantage by preventing the introduction of Na-based impurities that are typically associated with traditional Na₂CO₃ precipitation, where implementing pure CO₂ gas circumvents extraneous contamination. However, this approach presents a significant challenge: the persistent decrease in pH inherent to the use of CO₂ gas necessitates consistent and meticulous adjustments to preserve optimal pH levels.

One study introduced the use of CO₂ as a substitute for Na₂CO₃ for the carbonation of Li hydroxide in a concentrated KOH solution to produce Li₂CO₃ [46]. This study revealed that the process can be divided into four distinct stages: Stage I, rapid consumption of free hydroxide ions; Stage II, burst nucleation; Stage III, gradual precipitation; and Stage IV, the formation of bicarbonate, as shown in Fig. 3a. In Stage I, the carbonation rate gradually increased as the concentration of hydroxide ions decreased rapidly, indicating that reaction in Eq. (8) was predominant. This stage prepares for the supersaturation necessary for the nucleation of Li₂CO₃. In Stage II, a direct correlation between the carbonation rate and time was observed, indicating the burst nucleation of Li₂CO₃ and the predominance of reaction in Eq. (9). In stage III, the carbonation rate began to level off, indicating that the precipitation of Li₂CO₃ was near equilibrium. Reactions in Eq. (8) and Eq. (9) occurred predominantly during this stage. Finally, in stage IV, the presence of bicarbonate indicated an excess injection of CO₂, and the reverse reaction of Eq. (9) was dominant during this period. As shown in Fig. 3b–c, the effect of temperature was dominant during the early stages (stages I and II). Temperature significantly affected the maximum carbonation rate and the particle size distribution (PSD) of Li₂CO₃. As the temperature increased from 30°C to 90°C, the maximum carbonation rate also rose from ~67% to ~83%, primarily due to the decrease in solubility of Li₂CO₃. Interestingly, the maximum volume average diameter of Li₂CO₃ particles was achieved at 70°C, illustrating the intricate influence of temperature on the particle size of Li₂CO₃. Furthermore, the study found a correlation between the absorption rate of CO₂ and the crystallization rate of Li₂CO₃, suggesting that crystallization notably enhances the CO₂ absorption efficiency, especially at its onset. Li₂CO₃ produced during the carbonation process was of high purity and crystallinity, validating the carbonation process as a promising approach for CO₂ utilization and Li recovery from potent alkaline solutions.

Menghua Tian’s research group [47] made several noteworthy observations and suggested that the presence of ammonium ions ([NH₄⁺]) significantly influences the precipitation of Li₂CO₃. Specifically, these ions can impede the recarbonation of Li₂CO₃ crystals (i.e., dissolution of Li₂CO₃) and prevent over-carbonation, thereby simplifying the gas–liquid reactive crystallization process. In the research, it was found that by increasing the NH₃ · H₂O concentration from 200 g/L to 400 g/L, the supersolubility of the solution increased from 0.38 mol/L to 0.55 mol/L, the reactive crystallization induction period was extended from 38 min to 50 min, [Li₊] conversion improved from 43.0% to 49.6% within 120 min, and the particle size decreased, from around 52 μm to 28 μm. This enhancement resulted in a longer induction period and a more stable metastable solution. They also found that the absorption process of CO₂ in an NH₃·H₂O solution was more temperature-sensitive than that in an LiOH–CO₂ system. This characteristic can be effectively utilized to regulate the quantity of absorbed CO₂, thereby providing a mechanism for controlling the crystallization process.

In their further research [48], Tian et al. demonstrated that the reaction temperature significantly affects the Li₂CO₃ particle size and distribution. Increasing the temperature from 25°C to 60°C doubled the mean particle size. They also found that higher temperatures decreased the secondary nucleation rate, while increasing the growth rate. These findings provide crucial insights into particle size regulation during Li₂CO₃ production. In their innovative approach, they introduced microbubbles into a weakly basic LiCl-NH₃·H₂O–CO₂ system to enhance gas–liquid mass transfer and CO₂ absorption, which resulted in improved local supersaturation and the creation of reactive micro-zones [35]. Using CO₂ as a precipitant for the production of Li₂CO₃ helped to avoid metal ion impurities and provided a pathway for beneficial CO₂ conversion. In addition, the introduction of microbubbles significantly reduced the particle size of Li₂CO₃ from a D90 of 32.5 μm to 14.6 μm. This reduction was attributed to increased local supersaturation, which promoted nucleation and the formation of reactive micro-zones that inhibited crystal growth. This method yielded a Li₂CO₃ product that met the battery-grade requirements in terms of both purity and particle size.

A coupled reaction and solvent extraction process for the production of Li₂CO₃ from LiCl and CO₂ was introduced in Zhou et al. [49] Considering that the reaction is not spontaneous because the pH is lowered by carbonation, this study emphasized the need to remove hydrogen chloride to allow the reaction to proceed uninterrupted. The resultant Li₂CO₃ crystals from the room temperature open reactor process were observed to have a dual PSD: micro-sized particles (0.1–1 μm) and larger ones (10–100 μm). Interestingly, the bulk crystals fell within a smaller size range. This study provided insights into the crystal growth conditions. Radial growth occurred in the free aqueous phase resulting in large flakes or bulk crystals, whereas growth within water-in-oil structures yielded smaller ellipsoidal crystals. This study also revealed that factors such as a large phase ratio, surfactant addition, low temperature, and a short reaction time can help increase the formation of small-particle crystals while curtailing the growth of larger-particle crystals.

Similar to the application of microbubbles and the formation of water-in-oil structures, which significantly improved the precipitation of Li₂CO₃ in the aforementioned research, similar systematic adjustments can also lead to enhanced sedimentation characteristics. In the next section, we introduce studies that have employed a range of systems to exploit large specific surface area and enhance the properties and production of Li₂CO₃.

In this section, our focus moves towards improving the CO₂ gas reaction rate by utilizing large specific surface areas. A fundamental principle of chemistry is that an increased reaction surface area accelerates the reaction rate. In the context of Li₂CO₃ production, systems that artificially create structures with large specific surface areas, such as films or small droplet shapes, can increase the interaction area between Li⁺ and CO₃ ²⁻, thereby enhancing the reaction efficiency. Innovative studies have expanded the reactive interface by incorporating techniques such as spraying the liquid onto CO₂ gas or passing the liquid through a microchannel, including microfilm. In this section, we explore such systems, how they operate, and their impact on the production of Li₂CO₃, thereby shedding light on their potential to improve traditional production methods.

An effective strategy used to increase the specific surface area involves the reaction being conducted in the form of a thin film, as shown in Fig. 4a [50]. In this study, a system for constructing a falling film was developed. In an associated study, the gas–liquid reactive crystallization of CO₂ gas and LiOH solution was systematically investigated using a falling-film column to produce Li₂CO₃ crystals. This study identified three critical fluid dynamic parameters of the falling-film column: the Reynolds number, falling-film thickness, and exposure time. The reaction time was 26 min at pH 9 in a falling-film column, whereas that in a conventional stirred reactor was 60 min. During crystallization, the mean particle size changed from 5.72 μm at t = 7.5 min to 41.62 μm at t = 10 min, and then to 74.78 μm at t = 20 min. When the CO₂ flow rate increased from 0.5 L/min to 1.0 L/min, the carbonation time decreased from 20 min to 15 min. At varying temperatures from 10°C to 40°C, there was no difference in the absorbed CO₂, but the output of Li₂CO₃ at 40°C was 0.71 mol/L, which was more than triple the amount of 0.22 mol/L at 10°C. In addition, as the temperature increased from 10°C to 40°C, the volume mean particle size increased from 49.60 μm to 96.29 μm. It was determined that to obtain high Li₂CO₃ yields, the final pH of the carbonation reaction should be maintained within the range of 9.0–9.5. Interestingly, crystallization expedited the transport of CO₃ ²⁻ from the interfacial liquid film to the bulk solution, thereby promoting the overall absorption rate. The resulting products exhibited high purity (99.8%), with the particles displaying a distinctive flower-like composition comprising numerous plate-like primary crystals.

Sun et al. [51] used a spinning disk reactor (SDR) to synthesize Li₂CO₃ through the liquid reactive crystallization of LiOH and CO₂ (Fig. 4b). The study showed that the particle size and yield rate of Li₂CO₃ were influenced by multiple factors. In this respect, ultrasound was found to be a primary determinant of particle size, as it prevented agglomeration and resulted in smaller particle sizes. Other factors, such as the reaction temperature, CO₂ gas flow rate, and the spinning disk rotation speed were also found to influence the particle size. For example, when the concentration of LiOH solution was 1.5 mol/L at a temperature at 20°C, the yielding rate was less than 30%, whereas with an LiOH solution with a concentration of 2.0 mol/L at a temperature at 40°C, the yielding rate rocketed to 73%. The authors also found that high temperature increased the average volume size of the Li₂CO₃ particles and increasing the temperature from 20°C to 40°C resulted in an average particle size of 37–90 μm. Several operational parameters were also investigated, including the rotational speed of the rotating disc, circulation speed of the LiOH slurry, CO₂ flow rate, and the influence of the ultrasonic field. The ultrasonic field, temperature, and CO₂ flow rate were considered to be the main factors affecting the particle size. In contrast, the yield was significantly affected by the temperature, LiOH solution concentration, and CO₂ flow rate, while the ultrasonic field and rotational speed of the rotating disc had limited effects. The scanning electron microscopy images revealed flower-shaped particles composed of plate-like primary crystals.

In another study by Sun et al. [52], spray pyrolysis of LiHCO₃ was employed to produce Li₂CO₃ hollow spheres (Fig. 4c). The X-ray diffraction pattern of the product closely matched the standard pattern, suggesting a pure Li₂CO₃ crystalline phase in the synthesized product. Scanning electron microscopy revealed that the self-assembled hollow spheres were composed of primary particles approximately 200 nm in size. Furthermore, a crystal size distribution analysis indicated that the macro-volume means crystal size ranged from 4 μm to 9 μm, and this varied based on experimental conditions. Notably, the Brunauer–Emmett–Teller (BET) surface area reached 7.24 m²/g, surpassing the highest value previously reported and suggesting its potential application in the Li-battery industry.

One method used to broaden the specific surface area within confined spaces involves the formation of microbubbles within small tubes (Fig. 4d) [42]. A study adopting this approach utilized a microfiltration membrane dispersion microreactor to execute the precipitation reaction between LiOH and CO₂, using a water–ethanol mixture as the solvent. This research underlines the crucial roles played by the interphase mass-transfer rate of CO₂ and the solubility of Li₂CO₃ in achieving successful nanoparticle synthesis. The solubility ratio of LiOH to Li₂CO₃ surged to nearly two orders of magnitude higher than that in water when the mass fraction of ethanol (ϕ) was approximately 0.9. Nanoparticles with a narrow size distribution (coefficient of variation (CV) < 0.3) were reliably obtained with a ϕ value of 0.870 or above. The particle size demonstrated sensitivity to ϕ and the flux of the liquid/slurry feed, with increases in either parameter causing a reduction in the particle size from 150 nm to 30 nm. The use of LiOH slurry up to a concentration of 10 wt% slightly increased the particle size, but it did not influence the uniformity of the nanoparticle morphology. Given the potential for solvent recycling through partial water removal, this method of Li₂CO₃ nanoparticle preparation is a green process in which water is the only byproduct.

The studies discussed in this section have advanced the technology to precipitate Li₂CO₃ by utilizing the large specific surface area of the gas–liquid interface. In the next section, we will explore the precipitation of Li₂CO₃ using the unique properties of solubility and temperature.

Extracting Li from various sources is a complex process due to the presence of multiple impurities, including Mg²⁺, Na⁺, Ca²⁺, K⁺, B, and SO₄²⁻. As over 99.9% pure Li₂CO₃ is required for LiB production [53], it is essential to remove these impurities. Of these impurities, the removal of Mg²⁺ (3.4–113.7 g/L) is of particular significance, due to their high prevalence in Li-bearing (0.06–1.21 g/L) sources [54].

Any technique employed for the separation of similar elements must exploit the minor differences between them. Mg²⁺ has a high charge density that is double that of Li⁺, but it has roughly the same ionic radius and is readily hydrated [55]. Therefore, the extraction efficiency of Li depends heavily on the Mg-to-Li mass ratio, and effective separation is crucial. This challenge was addressed by utilizing a strategy that capitalizes on the solubility differences between LiHCO₃ and Li₂CO₃. When CO₂ is dissolved in water at lower pH levels, it dissociates into HCO₃⁻. In the pH range where HCO₃⁻ dominates, there is potential for the formation of LiHCO₃. However, LiHCO₃ is considered to be metastable; therefore, it is generally believed that LiHCO₃ exists in an ionic state as separate Li⁺ and bicarbonate (HCO₃⁻) ions (Fig. 2).

The subsequent application of heat leads to the precipitation of Li₂CO₃, during which CO₂ is introduced into the preliminarily precipitated Li₂CO₃, converting it back into LiHCO₃. This solution is then dissolved in water to facilitate the removal of impurities including Mg²⁺. The final round of heat application precipitates the solution, yielding high-purity Li₂CO₃. Leveraging the differences in solubility and meticulous purification processes effectively eliminates Mg²⁺ and other impurities and significantly contributes towards the manufacture of high-quality LiBs.

Focusing on the influence of various factors that enhance precipitation via solubility, the research conducted by Tao on refining crude Li₂CO₃ procured from Qinghai’s salt lake brines in China is instructive [56]. This study comprehensively investigated the direct carbonation of Li₂CO₃ slurries using CO₂–water solutions, wherein Li₂CO₃ was converted into the more soluble LiHCO₃. Experiments were conducted in a three-phase mechanically agitated slurry-bed reactor under various conditions. The parameters analyzed included CO₂ pressure, temperature, the CO₂ flow rate, Li₂CO₃ particle size, slurry filling degree, and the agitation speed. The findings demonstrated that an increase in the CO₂ pressure, agitation speed, CO₂ flow rate, or a decrease in the Li₂CO₃ particle size, slurry filling degree, reaction temperature, or solid concentration resulted in an increased Li₂CO₃ dissolution rate. Interestingly, the immersion time before carbonation had a negligible effect on the carbonation ratio of Li₂CO₃. The process kinetics are well represented by the equation 1 – (1 – X)1/3 = kt, where X represents the conversion of Li₂CO₃, k is the reaction rate constant, and t is the reaction time up to 50 min. This equation describes the dynamic nature of Li₂CO₃ carbonation. Using multiple carbonation cycles, it was possible to convert all insoluble Li₂CO₃ into LiHCO₃. This study highlights the importance of maintaining room temperature, considering high CO₂ partial pressures (>0.6 MPa), performing multiple carbonation cycles, and ensuring a high agitation speed for optimal Li₂CO₃ carbonation. Moreover, this process enables potential CO₂ recovery by heating the LiHCO₃ solution under vacuum.

Xu developed a process for creating battery-grade Li₂CO₃ from Damxungcuo saline lake [14] that involved a two-stage Li₂CO₃ precipitation process in a hydrometallurgical system to remove impurities. Initially, industrial-grade Li₂CO₃ was obtained by removing Fe³⁺, Mg²⁺, and Ca²⁺ from Li-containing liquor. Subsequently, the industrial-grade Li₂CO₃ was treated with CO₂ to yield more soluble bicarbonates, while EDTA-Li was used to chelate Ca²⁺ and Mg²⁺. The decomposition of Li₂CO₃ at 85°C yielded insoluble Li₂CO₃, resulting in a high purity (99.6%) and homogeneous Li₂CO₃ in the final precipitate. The process effectiveness was influenced by certain factors; for example, an L/S mass ratio of 30:1 favored Li₂CO₃ slurry formation and hot water washing promoted ion removal. Furthermore, the cyclic use of the filtrate assisted in improving Li recovery, thereby paving the way for large-scale production.

Linneen focused on the transformation of industrial-grade Li chloride into high-purity battery-grade Li₂CO₃ and achieved an impressive purity of >99.95% [38] through a four-step precipitation process. The treatment solution was concentrated in the salt lake, and it contained contaminants such as K, Na, Mg, Ca, Cu, Ni, and iron chloride at concentrations ranging from 2.5 M to 10 M. In the first step, the heavy metals and the majority of Mg were eradicated by increasing the pH of the solution, and the remaining Mg and Ca were then removed through the addition of sodium oxalate, which successfully reduced the calcium levels to 5–6 ppm in the 10 M solution. A higher molarity and ionic strength of the LiCl solution facilitated more effective impurity removal. Finally, high-purity Li₂CO₃ was obtained through two stages of precipitation: initial precipitation from the brine solution, followed by a second purification step using pressurized CO₂. This final step allowed for the removal of entrapped Na and K from the first precipitate, ultimately yielding >99.95 wt% purity Li₂CO₃.

Zhou et al. successfully synthesized battery-grade Li₂CO₃ via a carbonation–decomposition method using crude Li₂CO₃ that contained various impurities [57]. The starting material was impure, defective, crude Li₂CO₃ with a purity of 98.56 wt%, and contained 5,873 ppm of SO₄²⁻, along with other impurities such as Na, K, Ca, Fe, and Mg. This study underlines the significance of using filter operations to enhance product quality. After filtration, the purity of Li₂CO₃ increased significantly to 99.65 wt%, and the amounts of SO₄²⁻ and Fe dropped to 222.0 ppm and 6.6 ppm, respectively. Interestingly, despite the use of the qualitative filter paper, the Ca content increased from 14.9 ppm in crude Li₂CO₃ to 249.6 ppm. The study also employed ethylenediaminetetraacetic acid disodium salt (EDTA-2Na) for further purification, and it was highly efficient at removing metal cations, particularly Ca and Fe ions. Post-EDTA-2Na treatment, the Ca and Fe contents plummeted to 2.9 ppm and 0.7 ppm, respectively, improving the whiteness of the sample to 97.8 Wb. The removal efficiencies obtained for Ca, Fe, and SO₄²⁻ were calculated as 80.5%, 99.4%, and 97.5%, respectively.

The research conducted by Meng Shi focused on a unique electrodialysis (ED) method for recovering Li and manganese from spent LiBs [58]. Unlike the hydrometallurgical routes traditionally used for Li recycling, this study used electrodialysis (ED), employing a CO₂-capture agent, N-methyldiethanolamine, as a regenerable catholyte. Both surrogate solutions and Li-rich leachates were tested to recover Li and Mn as Li₂CO₃ and MnCO₃, respectively. The process achieved battery-grade purity levels (99.6% purity) of recovered Li₂CO₃ from both the surrogate and leachate solutions. Notably, the energy consumption of the ED process was low, at 0.5 kW h/g Li, and Li₂CO₃ yields as high as 48.4% were reported. Despite the Li₂CO₃ yield being lower than what is typically needed for industrial applications, this study highlights several factors that can be optimized to increase the yield, such as the anolyte pH, Li₂CO₃ precipitation temperatures, and the choice of catholyte tertiary amine solutions. The ED method also suggests the possibility of a carbon-negative process with CO₂ as the only chemical consumed. This study presents a promising, environmentally friendly, and energy-efficient method for recovering Li₂CO₃ during LiB recycling.

The method proposed here leverages CO₂ as a leaching agent to recycle spent Li iron phosphate (LiFePO₄) at room temperature [29]. This strategy relies on the weak acidity of H₂CO₃, which restricts the leaching of Fe and P elements, thereby ensuring high extraction selectivity for Li. Notably, the role of CO₂ is two-fold; in addition to facilitating the leaching process, it also aids in the chemical precipitation of leached Li⁺ to yield value-added Li₂CO₃.

The CO₂-leaching strategy exhibits clear environmental and economic advantages. A life cycle analysis revealed that it consumes low amounts of energy and would reduce greenhouse gas emissions and provide substantial economic benefits. With an impressive recovery yield of high-purity Li₂CO₃ products and the ability to fix approximately 120 kg of CO₂ per ton of spent LiFePO₄, the use of this strategy demonstrates effective carbon reduction.

In addition, the elimination of strong acids/bases and additional precipitants hindered the creation of high-salt wastewater, high-lighting the environmental compatibility of this method. These processes also caused the lowest energy consumption, greatest reduction in greenhouse gas emissions, and sizable economic profits per LiFePO₄ cell unit. Specifically, the CO₂ leaching strategy, particularly the CO₂-O₂ recycling route, exhibited the lowest energy consumption at 2.29 MJ per kg of LiFePO₄ cells, the greatest reduction in greenhouse gas emissions at 194 g/kg of LiFePO₄ cells, and substantial economic profits at $4.04 per kg of LiFePO₄ cells.