Developing a feasible and low-cost strategy for the recovery of calcium fluoride efficiently from fluoride-containing wastewater is very essential for the recycle of fluoride resources. Herein, a modified lime precipitation method was employed to recover CaF2 from fluorinated wastewater using a special icy lime solution. Intriguingly, the highest F removal was greater than 95% under the optimal condition, leaving a fluoride concentration from 200 to 8.64 mg/L, while the lime dosage was much lower than that of industry. Importantly, spherical-shaped CaF2 particles with a 93.47% purity and size smaller than 600 nm were recovered, which has a high potential for the production of hydrofluoric acid. Besides, the precipitation was significantly affected by Ca/F molar ratio, stirring time, temperature, and solution pH. Furthermore, the thermodynamics and kinetics were investigated in detail to reveal the crystallization process. As a result, the defluorination reaction followed the pseudo-second order reaction kinetics model. Also, CO2 in the air adversely influenced the CaF2 purity. Based on this facile method, a high lime utilization efficiency was applied to defluorination, which contributed to protecting the environment and saving costs. This study, therefore, provides a feasible approach for the green recovery of fluorine resources and has significance for related research.

  • A facile icy lime precipitation method was employed to remove F.

  • Highly pure spherical particles with CaF2 of more than 93% were recovered.

  • Thermodynamics implied that CO2 had an inhibitory effect on the purity of CaF2.

  • Change in F fitted well with the pseudo-second order reaction kinetics model.

  • The co-precipitation of CaF2 and CaCO3 occurred.

In recent years, the semiconductor (Sinharoy et al. 2024) and photovoltaic industries (Hu et al. 2024) have witnessed rapid development, resulting in the discharge of a significant amount of wastewater containing fluorine with concentrations ranging from ten to thousands of mg/L during production processes (Li et al. 2021; Qiu et al. 2022). This not only pollutes the water system but also leads to wastage of valuable fluorine resources. Therefore, it is crucial to treat fluorine-containing wastewater before its discharge. In this case, the converting fluorine in wastewater into recyclable substances like fluorite (CaF2) would be an ideal approach for environment protection and resource recovery.

Fluorite (CaF2) is widely used in chemical, optics, national defense and fluorine chemical industries (Gao et al. 2021; Huang et al. 2022; Sun et al. 2022) due to its excellent characteristics. For example, fluorite serves as a memory material for manufacturing battery capacitors in the semiconductor industry (Park et al. 2023). Specifically, high-grade fluorite is the most important raw material for the production of hydrofluoric acid (Pelham 1985; Yen Chau Nguyen et al. 2021). Nevertheless, with the rapidly growing demand for fluorite globally, the reserve of fluorite decreases rapidly, e.g., only 230 million t (Yin et al. 2016). In addition, the fluorite mine production worldwide in 2019 was about 7,000 kt, heavily concentrated in three countries: China (4,000 kt), Mexico (1,200 kt), and Mongolia (670 kt) (Takaya et al. 2021). Inevitably, a steady supply of fluorite has become a major concern for numerous countries. In this regard, recovering calcium fluoride from existing fluorine resources becomes an important and promising strategy in the future.

Currently, common methods for treating fluorine-containing wastewater primarily include chemical precipitation, adsorption, membrane treatment, electric coagulation, and fluidized bed (Zeng et al. 2019; Shao et al. 2021; Wan et al. 2021; Qiu et al. 2022; Wang et al. 2022). Among these methods, chemical precipitation is widely recognized as a promising defluorination technology owing to its simplicity and cost-effectiveness. Due to the strong affinity of fluoride to calcium (Shao et al. 2021), calcium salts or calcium-based materials are commonly used as precipitants to remove fluoride from the wastewater. For example, Chen et al. (2022) recovered fluoride from groundwater and industrial wastewater by extracting calcium hydroxide from waste eggshells. Ezzeddine et al. (2014) achieved a reduction in fluoride concentration to approximately 8 mg/L using commercially available lime as a defluorination agent combined with reverse osmosis.

Nevertheless, the current precipitation operation often encounters challenges such as complex operation flowsheet, high costs and low utilization rates (He et al. 2020). For instance, Zhai et al. (2013) reported that direct addition of lime to fluorine-containing wastewater resulted in treated water with significantly higher concentrations of fluoride ions ranging from 30 to 50 mg/L than the required discharge limits. In addition, the low purity (only 20–40% (Zhou et al. 2023)) of recovered calcium fluoride needs to be further purified prior to use, thereby inevitably increasing the subsequent processing steps and costs. This indicates that the efficiency of lime treatment of fluorine-containing wastewater should be further improved. Therefore, developing a feasible chemical precipitation method to recover high-purity calcium fluoride from the wastewater is urgent and meaningful to environmental protection and high-efficient utilization of fluoride resource.

In this study, a special icy lime-induced precipitation method was proposed to recover calcium fluoride efficiently from fluorine-containing wastewater. Furthermore, the impact of various factors including Ca/F molar ratio, stirring time, reaction temperature, pH on fluoride removal efficiency was systematically investigated. Notably, highly pure spherical-shaped particles of calcium fluoride were successfully recovered using this approach. Besides, thermodynamics and kinetics were investigated to reveal the precipitation mechanisms. The composition and morphology of the obtained precipitates were characterized by X-ray diffraction and SEM-EDS. Henceforth, this study provides a straightforward and viable strategy for effectively recovering high-purity calcium fluoride from fluorine-containing wastewater.

Sample preparation

Sodium fluoride (98%, AR, Aladdin) and deionized water (IQ7000, MERCK, Germany) were used to simulate fluoride wastewater, with a fluoride concentration of about 400 mg/L. Concentrated hydrochloric acid (38%, AR, Sinopharm) and NaOH (96%, AR, Sinopharm) were used to prepare HCl and NaOH solutions to adjust the solution pH.

Experimental procedure

First, calcium hydroxide solution was prepared by using calcium oxide (96%, AR, Sinopharm) and deionized water in an icy water bath where calcium oxide can be dissolved more quickly and efficiently, thereby improving subsequent defluorination. The prepared calcium hydroxide solution was then mixed with the same volume of fluoride wastewater to prepare 200 mg/L F solution. The mixture was stirred at a speed of 200 rpm under constant temperature, which was further separated using the centrifuge at 10,000 rpm (TG18I, YINGTAI, China). Afterwards, the residual concentration of F in the supernatant was detected, while the precipitated substances were washed using deionized water and then dried in the electric blast drying oven (101-1AB, Taisite, China) to constant weight.

Analytical methods and characterization

F concentration was detected using a pH/Eh meter (pH 2100, EUTECH, America) with a selective electrode (CHN090, EUTECH, America) for fluorine ion. The composition and crystallization of obtained precipitates were detected by X-ray diffraction (Cu Kα, λ = 0.15406 nm, D8 Advance, Bruker, Germany) at 2θ ranging from 20 to 80°, with a scanning speed of 2°/min. In addition, the SEM-EDS (Gemini 300, ZEISS, Germany) was used to detect the morphology and elemental composition of the precipitates, with an acceleration voltage of 20 k eV. Thermodynamic analysis was conducted using HSC Chemistry. The removal efficiency of fluoride (η) was calculated using the following equation:
formula
(1)
where C1 (mg/L) was the initial concentration of F in the mixed solution, C2 (mg/L) was the residual F concentration in the supernatant after precipitation and centrifugation.

Fluoride removal efficiency

As shown in Equations (2) and (3), CaO in water can be rapidly converted to Ca(OH)2 which further dissociates into Ca2+. Considering that the solubility of Ca(OH)2 decreases at increasing temperature, the batch precipitation experiments in the icy water bath are expected to promote the dissolution of Ca(OH)2 and improve the fluoride removal efficiency:
formula
(2)
formula
(3)
We carried out five parallel tests to study the efficiency of defluorination with dissolved lime at different temperatures and performed error analysis (error bars expressed as standard deviation (SD)). As shown in Figure 1, the dissolved lime in the icy water bath (278 K) had a remarkable role on fluoride removal. For instance, the highest removal efficiency was 95.68% when the solution was controlled at 278 K. However, the removal efficiency was significantly reduced at 298 and 328 K, respectively. This indicated that dissolved lime in the icy water bath greatly improved the utilization rate of lime and then achieved much higher fluoride removal efficiency, compared with common ambient lime precipitation method.
Figure 1

Effect of temperature on the removal efficiency of fluoride.

Figure 1

Effect of temperature on the removal efficiency of fluoride.

Close modal

Effects of Ca/F molar ratio

The influence of Ca/F molar ratio on the F removal efficiency (η) was investigated. The Ca/F molar ratio was calculated by the amount of CaO before dissolution and the initial fluorine concentration of wastewater. According to Equation (4), the eventual formation of CaF2 can be achieved when the Ca/F molar ratio is 0.5. However, due to the influence of temperature, pH and other factors, excess precipitant is normally required to ensure F removal efficiency and produce more CaF2:
formula
(4)
formula
(5)
formula
(6)
As shown in Figure 2(a), η increased from 80.10 to 93.56% quickly with Ca/F molar ratio from 0.50 to 0.75. When Ca/F molar ratio increased to 1.00, η increased slightly to 94.06%, whereas when Ca/F molar ratio was 1.25, η decreased to 93.05%. Therefore, a Ca/F molar ratio of 0.75 was selected for further study. It should be noted that the total amount of precipitating agent was 1.5 times as to the theoretical value, significantly lower than that applied in the industry (Zhai et al. 2013). This greatly reduces the amount of lime and the treatment cost.
Figure 2

Effects of (a) Ca/F molar ratio, (b) stirring time, (c) initial pH, and (d) temperature on η (stirring time = 60 min, initial pH = 5, temperature = 298 K when the Ca/F mole ratio changed).

Figure 2

Effects of (a) Ca/F molar ratio, (b) stirring time, (c) initial pH, and (d) temperature on η (stirring time = 60 min, initial pH = 5, temperature = 298 K when the Ca/F mole ratio changed).

Close modal

It is speculated that high content of Ca(OH)2 in the reaction solution is subject to absorb CO2 from the air, resulting in less available Ca(OH)2 for defluorination (Equation (5)), and resulting in CaCO3 precipitation (Equation (6)). In addition, the higher Ca/F molar ratio indicated more alkaline reaction. The solubility of CaF2 in alkaline solution at 298 K is higher than that of CaCO3 (Lacson et al. 2021), increasing F concentration in the solution, which is not conducive to defluorination. This may be the reason why the efficiency of fluoride removal decreases when the amount of lime is increased to a certain value. As a result, the amount of precipitant should be controlled reasonably.

Effects of stirring time

Further precipitation experiments were performed to investigate the influence of reaction time on η. As shown in Figure 2(b), η increased apparently from 90.29 to 92.83% at a longer stirring time from 15 to 30 min. A higher η of 93.48% was found when the stirring time was further increased to 45 min. However, when the stirring time was increased to 60 min, η only slightly increased to 93.56%. It should be noted that the increased reaction time may increase the absorption of CO2 in the alkaline solution, negatively influencing the purity of calcium fluoride. The effect of fluoride removal was not significantly improved by increasing time. This suggested the optimal stirring time would be 45 min.

Effects of pH

The pH of the fluorine-containing wastewater was adjusted to 1, 2, 3, 4, and 5, respectively. Figure 2(c) illustrates that an extremely low η of 1.93 and 3.66%, respectively, was observed at pH 1 and 2 of the wastewater, indicating minimal precipitation occurred under strongly acidic conditions. With the increase of pH of the wastewater to 3, the η increased dramatically to 95.07%. When the pH of the fluorine-containing wastewater was 4, η increased slightly to 95.68%, leaving a residual F concentration of 8.64 mg/L. However, when the pH of fluorinated wastewater was 5, there was a decrease in η down to 93.48%.

To unveil the pH influence mechanism on η, the pH evolution of the mixed solution during the reaction was conducted, as shown in Figure 3. When the initial pH of the wastewater was 5, the highest pH of the mixed solution was 12.4 visibly after the addition of icy lime. Under this condition, more CO2 from the air can be absorbed to compete with F for Ca2+ or Ca(OH)2, reducing the removal efficiency of fluoride. Moreover, the generation of more CaCO3 resulted in decreased purity of calcium fluoride, further deteriorating the recovery and utilization processes.
Figure 3

pH evolution of fluorinated wastewater with different initial pH.

Figure 3

pH evolution of fluorinated wastewater with different initial pH.

Close modal
In order to better reveal the experimental results, the Eh–pH phase diagrams of Ca–F–H2O system and C–Ca–H2O system at 298 K were also plotted using the HSC Chemistry, as shown in Figure 4.
Figure 4

Eh–pH phase diagram of Ca–F–H2O and C–Ca–H2O systems at 298 K.

Figure 4

Eh–pH phase diagram of Ca–F–H2O and C–Ca–H2O systems at 298 K.

Close modal

From Figure 4, the left side of vertical line a was the region where Ca2+ and HF dominate, indicating that Ca2+ and HF exist stably in an acidic system. Thus, this region was not conducive to the reaction of F with Ca2+ to form CaF2. Lacson et al. (2021) also reported that fluorine mainly exists in the form of HF in a strongly acidic environment, consistent with the results of this experiment.

Furthermore, the middle of vertical lines a and b was the dominant region of F and CaF2, indicating that HF would dissociate F and CaF2 can exist stably in the pH range of 2.5–13.5. In addition, line b suggested that CaCO3 could exist in the region at pH above 4.6. In order to reduce the negative influence of CaCO3 on the purity of calcium fluoride, the pH of the wastewater should be as low as possible, but when the pH was below 4, the removal efficiency was too low. Therefore, the optimal initial pH of the fluorine-containing wastewater was determined at pH 4.

Thermodynamics

The effect of reaction temperature on the η was also studied as shown in Figure 2(d). When reaction temperature was controlled at 298 K, the highest η of 95.68% was achieved. However, with the increase of temperature, the η slightly decreased. When the reaction temperature was 328 K, the removal efficiency of fluoride reduced to 92.55%. The experimental results showed that increasing temperature was not conducive to defluorination.

In order to uncover the effect mechanism of reaction temperature, the thermodynamic parameters such as enthalpy change (ΔH), entropy change (ΔS) and Gibbs free energy change (ΔG) of relevant reactions during the precipitation process were calculated by using the Reaction Equations module in HSC Chemistry 6.0 software, as shown in Table 1 and Figure 5.
Table 1

Thermodynamics of precipitation reaction system at 298 K

Reaction equationsΔH (kJ)ΔS (J/K)ΔG (kJ)KaLog(Ka)
Main reaction (1) Ca2+(aq) + 2F(aq) = CaF2 −14.222 151.293 −59.330 2.484E + 010 10.395 
(2) 2NaF(aq) + Ca(OH)2 = CaF2 + 2NaOH(aq) −44.950 −25.828 −37.249 3.361E + 006 6.526 
Side reaction (3)  −107.512 −172.396 −56.112 6.784E + 009 9.831 
(4)  11.718 198.193 −47.373 1.996E + 008 8.300 
(5) CO2(g) + Ca(OH)2 = CaCO3 + H2−113.025 −135.509 −72.623 5.300E + 012 12.724 
Reaction equationsΔH (kJ)ΔS (J/K)ΔG (kJ)KaLog(Ka)
Main reaction (1) Ca2+(aq) + 2F(aq) = CaF2 −14.222 151.293 −59.330 2.484E + 010 10.395 
(2) 2NaF(aq) + Ca(OH)2 = CaF2 + 2NaOH(aq) −44.950 −25.828 −37.249 3.361E + 006 6.526 
Side reaction (3)  −107.512 −172.396 −56.112 6.784E + 009 9.831 
(4)  11.718 198.193 −47.373 1.996E + 008 8.300 
(5) CO2(g) + Ca(OH)2 = CaCO3 + H2−113.025 −135.509 −72.623 5.300E + 012 12.724 
Figure 5

(a) ΔH and (b) ΔG of reactions at different temperatures.

Figure 5

(a) ΔH and (b) ΔG of reactions at different temperatures.

Close modal

Table 1 shows the thermodynamics of precipitation reaction process at 298 K. It can be clearly seen that the product of reactions (1) and (2) was CaF2 while the product of side reactions (3)–(5) was CaCO3, due to the presence of CO2 in the air. Meanwhile, the Gibbs free energy change values (ΔG) of reactions (1)–(5) at 298 K were all less than 0, indicating that all reactions could occur spontaneously. So it should be noted that the occurrence of reactions (3)–(5) is unavoidable at 298 K, which would inhibit the formation of calcium fluoride due to the formation of calcium carbonate. Chaudhary & Maiti (2019) reported that co-precipitation of F and Ca2+ was the main mechanism for fluoride removal from industrial wastewater when using synthetic calcium hydroxide nanorods, and the existence of calcium carbonate was due to the absorption of CO2 in the air by the alkaline mixture. Moreno et al. (2021) also proved that Ca(OH)2 with high humidity captured CO2 from the air to produce CaCO3 at 298 K and atmospheric pressure.

The variations of ΔH and ΔG for five reactions in the range of 298–328 K are illustrated in Figure 5 to further study the reaction temperature correlated thermodynamics.

As shown in Table 1 and Figure 5(a), ΔH of reactions (1)–(3) and (5) was less than 0, indicating that the four reactions were exothermic reactions. In contrast, the ΔH of reaction (4) was greater than 0, indicating that this reaction was endothermic. Therefore, increasing temperature would promote the reaction (4) to the right, resulting in the generation of more CaCO3. This is contrary to our research objectives. The ΔH of reaction (1) and reaction (4) rose gradually with the increase of temperature, meaning that the two reactions were readily affected by temperature. The difference was that the ΔH of the other three reactions changed gently with temperature, implying that these reactions were weakly affected by temperature.

Notably, the ΔG values of reactions (1) and (4) decreased with increasing temperature from Figure 5(b), indicating that heating would promote the spontaneous enhancement of these two reactions and increase their likelihood of occurrence. In other words, the increase of reaction temperature may cause the increase of CaF2 and CaCO3 formed by ions in the system. In contrast, the ΔG values of the other three reactions elevated with the increase in temperature, suggesting that heating may weaken the spontaneity of these reactions and was not conducive for the formation of corresponding products. To sum up, the influence of temperature on the reactions in the system is complex. However, it is evident that raising the reaction temperature will promote calcium carbonate production.

The above thermodynamic analyses indicated that the decrease of defluorination was due to the competition between CO2 and F for Ca2+ or Ca(OH)2, and increasing temperature was conducive to the generation of CaCO3. Therefore, the reaction of defluorination and recovery of calcium fluoride in this study was more favorable at 298 K, which could reduce energy consumption and control treatment costs.

Kinetics

Precipitation kinetics are important to evaluate the defluorination process (Ruan et al. 2017). Therefore, the reaction kinetics model was simulated to reveal the evolution trend of fluoride concentration and the related precipitation mechanisms.

The evolution of fluoride concentration and defluorination efficiency is shown in Figure 6(a). The F concentration dramatically dropped to 35.04 mg/L at 300 s, indicating the highest reaction rate in the initial stage. Subsequently, at 3,000 s, the defluorination efficiency reached 95.62%, which only increased slightly compared to 300 s, indicating a decreased reaction rate. In the initial phase, the high concentration of reactants resulted in a rapid reaction, and the reaction rate was controlled by the chemical reaction. However, with the decrease of reactants concentration, the collision probability between reactant ions decreased, and the reaction rate changed to a controlled diffusion process (Farhoosh 2018). Therefore, the reaction rate decreased rapidly from the highest value in the process of defluorination.
Figure 6

Residual fluoride concentration (C) and defluorination efficiency (η) (a) and fitting graph of pseudo-first and pseudo-second order reactions for residual fluoride (b) at different reaction times.

Figure 6

Residual fluoride concentration (C) and defluorination efficiency (η) (a) and fitting graph of pseudo-first and pseudo-second order reactions for residual fluoride (b) at different reaction times.

Close modal
According to the above F removal data, the rates (Farhoosh 2018) of the pseudo-first order and pseudo-second order reaction can be calculated based on Equations (7) and (8):
formula
(7)
formula
(8)
where C (mg/L) was the concentration of fluoride in the reaction system at time t (s), and k1 (s−1) was the pseudo-first order reaction rate constant, k2 (mg−1 L s−1) was the pseudo-second order reaction rate constant.
The integration of Equations (7) and (8) followed Equations (9) and (10), respectively:
formula
(9)
formula
(10)
where a1 (mg/L) and a2 (mg/L) were integration constants.

Figure 6(b) illustrates the fitted pseudo-first order and the pseudo-second order reaction kinetics model. The relevant parameters are listed in Table 2.

Table 2

Fitted pseudo-first order and the pseudo-second order reaction kinetics

EquationkaR2
Pseudo-first order y = exp (kxa−0.0047 −5.2902 0.9281 
Pseudo-second order y = 1/(akx−5.4908 × 10−5 0.00501 0.9884 
EquationkaR2
Pseudo-first order y = exp (kxa−0.0047 −5.2902 0.9281 
Pseudo-second order y = 1/(akx−5.4908 × 10−5 0.00501 0.9884 

The fitted curve is shown in Figure 6(b) and the R2 value is shown in Table 2. K represents the speed of reaction concentration change with time and a can reflect the initial concentration of F in wastewater to a certain extent. According to Figure 6(b), the precipitation reaction was more consistent with the pseudo-second order reaction kinetics model.

The precipitate produced under optimal condition (Ca/F molar ratio = 0.75, stirring time = 45 min, temperature = 298 K) is shown in Figure 7(a). Obviously, the small white precipitate produced tended to aggregate, which appeared in powder form.
Figure 7

(a) Precipitates and (b) XRD patterns of precipitates (pH = 4, Ca/F molar ratio = 0.75, stirring time = 45 min, reaction temperature = 298 K) and standard spectrum diagram of CaF2, CaCO3.

Figure 7

(a) Precipitates and (b) XRD patterns of precipitates (pH = 4, Ca/F molar ratio = 0.75, stirring time = 45 min, reaction temperature = 298 K) and standard spectrum diagram of CaF2, CaCO3.

Close modal

The peaks shown in the XRD pattern (Figure 7(b)) were consistent with the peaks of standard calcium fluoride (PDF#35-0816) and calcium carbonate (PDF#05-0586). The plane index of calcium fluoride was 111, 220, 311, 400 and 331 in turn, consistent with the research results obtained by Liu et al. (2022) and Wan et al. (2020). Wang et al. (2021) also reported two types of peaks representing calcium fluoride and calcium carbonate in the precipitate. Moreover, the XRD pattern showed that the peak strength of calcium fluoride was significantly higher than that of calcium carbonate, indicating that the main component of the precipitate was CaF2 and only contained a small amount of CaCO3.

In order to further understand the surface morphology and composition of the precipitate, SEM-EDS was used to characterize the precipitate sample, and the results are shown in Figure 8.
Figure 8

SEM (a–c) and EDS (d–h) images and particle size distribution (i) of precipitated particles.

Figure 8

SEM (a–c) and EDS (d–h) images and particle size distribution (i) of precipitated particles.

Close modal

As shown in Figure 8(a) and (b), the smaller particles tended to be more quadrate while the larger particles tended to be more spherical in the precipitate. Furthermore, it can be clearly seen from Figure 8(b) that the surfaces of large spherical particles were formed by stacking layers of flake crystals, with ladder-like edges, similar to the hexahedron (Figure 8(c)). In contrast, the surfaces of the small quadrate particles were smooth, without layered crystals.

EDS images (Figure 8(d)–(h)) showed that the main elements in the sample were F and Ca, with a small amount of C and O. This indicated that there was co-precipitation of CaF2 and CaCO3 in the reaction process. According to the energy spectrum (Figure 8(h)), calculated by the F weight percentage, the main component of precipitate was CaF2 with a purity of 93.47%.

In order to obtain more effective information on the recovered calcium fluoride particles, ImageJ was used to measure the size of 206 particles as shown in Figure 8(a). The particle size distribution and Gaussian fitting results are shown in Figure 8(i).

Figure 8(i) shows that the particle size distribution range was between 50 and 500 nm, with an average size of 123.89 nm. Specifically, 47 particles were located in 50–100 nm, accounting for 22.8% of all the particles. The small particle size range of 500–550 nm only accounts for 0.5%. Wang et al. (2013) also reported similar particle size distribution when investigating the compaction flocculation process for treating fluorine-containing wastewater. These small particles of calcium fluoride are more advantageous in the production of hydrofluoric acid.

Therefore, based on the above discussion, it can be learned that the as-produced calcium fluoride from treating fluorine-containing wastewater by using unique icy lime solution endowed high purity and relatively smaller particle size, which was suitable and helpful for subsequent hydrofluoric acid production.

In this work, lime dissolved in icy deionized water was applied to produce calcium hydroxide as precipitant for F removal from fluorinated wastewater. At 298 K, when Ca/F molar ratio was 0.75, stirring time of 45 min, initial pH of 4, the efficiency of fluoride removal reached at 95.68%. Accordingly, the F concentration in the effluent water decreased ultimately to 8.64 mg/L. Comparatively, the treatment effect was significantly better than that of related studies (Table 3). For instance, the relevant studies of Hu et al. (2024) only reduced the fluorine-containing wastewater to 20–30 mg/L. This further highlighted the high efficiency of dissolved lime in an ice water bath for fluoride removal. In addition, we successfully obtained calcium fluoride sludge with CaF2 content of 93.47%.

Table 3

Comparison of the results of this study with those of other relevant studies

PrecipitantResidual F level (mg/L)η (%)
Our results Lime 8.64 95.68 
Sinharoy et al. (2024)  Calcium salt 80–40 82.50 
Hu et al. (2024)  Ca(OH)2/CaCl2 20–30 97.00 
PrecipitantResidual F level (mg/L)η (%)
Our results Lime 8.64 95.68 
Sinharoy et al. (2024)  Calcium salt 80–40 82.50 
Hu et al. (2024)  Ca(OH)2/CaCl2 20–30 97.00 

Crucially, our study had the following obvious advantages: (1) the removal efficiency of lime and purity of calcium fluoride sludge were improved creatively through an ice water bath; (2) the influence of CO2 in the air on the purity of calcium fluoride and fluoride removal were analyzed by thermodynamics for the first time; and (3) the treatment method was simple and efficient, and the treatment goal was achieved without the help of external forces or high-efficiency reaction devices. Therefore, the icy lime method is an economical and applicable method of removing fluoride, thereby providing a green way to produce high-purity calcium fluoride from fluorine-containing wastewater.

However, our research still has some shortcomings in some aspects, and the future direction of work is mainly as follows: (1) To explore how impurities affect the formation of calcium fluoride in the treatment process of industrial fluorine-containing wastewater and (2) to expand the study of the specific influence mechanism of CO2 concentration, flow rate and contact time in the air on the defluorination process.

This research was supported by the Key Research and Development Program of Hubei Province (2021BCA127).

All relevant data are included in the paper or its Supplementary Information.

The authors declare there is no conflict.

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