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Article

Removal of Low Concentrations of Er(III) from Water Using Heptadecyl-1,1-bisphosphonic Acid

by
Chunhua Bai
1,
Xiaoning Yang
1,2 and
Guanghui Li
1,*
1
College of Mining and Coal, Inner Mongolia Key Laboratory of Mining Engineering, Inner Mongolia University of Science and Technology, Baotou 014010, China
2
CNNC Tongliao Uranium Industry Limited Liability Company, Tongliao 014010, China
*
Author to whom correspondence should be addressed.
Minerals 2024, 14(6), 534; https://doi.org/10.3390/min14060534
Submission received: 6 April 2024 / Revised: 18 May 2024 / Accepted: 19 May 2024 / Published: 22 May 2024
(This article belongs to the Special Issue Green and Efficient Recovery/Extraction of Rare Earth Resources)
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Abstract

:
The removal of low concentrations of rare-earth ions (e.g., Er(III)) from water has stimulated interest in the field of mineral processing and water treatment. Here, an ion-exchange and complexation-assisted precipitation method for the removal of low concentrations of Er(III) from water using heptadecyl-1,1-bisphosphonic acid (HBPA) was investigated. The results showed that effective cation-exchange between Er(III) ions and the bisphosphonate headgroup was achieved, and the solution pH abruptly decreased from 6.5 to around 3.1 at the first stage, which further led to the formation of less soluble Er(III) heptadecyl-1,1-bisphosphonate complexes. While low concentrations of Er(III) ions in water are typically treated by the addition of HBPA, followed by the addition of sodium bicarbonate (adjusting the pH to 6–8) and activate carbon, Er(III) ions could be efficiently removed from aqueous solution after about 30 min based on the cation-exchange and complexation-assisted precipitation method. Additionally, the removal of ultra trace amounts of Er(III) ions was not significantly affected by coexisting trace amounts of alkaline-earth metal ions (Mg2+, Ca2+ and Sr2+). HBPA is an effective Er(III) chelator, which may be a potential and promising alternative technique to remove Er(III) ions from aqueous solutions.

1. Introduction

Due to their unique optical, electrical, magnetic, superconducting and catalytic properties [1], rare-earth elements (REEs) are increasingly recognized for the development of traditional and emerging strategic industries, such as energy, military, aerospace and electronic information, chemical, metallurgy and materials, biomedical engineering and other high-tech fields [2,3,4,5,6,7,8,9]. Therefore, the global availability, distribution and use of commercial products containing REEs, along with the impact of the mining of REEs (particularly for soluble rare-earth ions) on surrounding areas, have aroused wide concern among researchers and the public over the last few years due to the increasing mobility and bioavailability of REEs [10,11,12,13]. Additionally, REEs are biologically inessential elements, and they may conversely accumulate in animal or human bones through biochemical processes over their lifespan [14,15], which inevitably leads to adverse effects, such as bone toxicity, cytotoxicity and growth inhibition [16,17,18]. Moreover, end-of-life REEs containing products may induce severe environmental impacts and economic consequences, along with the security of the rare-earth supply chain [19]. Usually, the concentration of REEs in groundwater is less than 0.1 mg/L, but the concentration of REEs is higher in wastewater from the REE industry (ranging from 0.5 to 10 mg/L). To alleviate water pollution and control REE emissions from the REE industry in China, the permissible limits for wastewater discharges from the REE industry were set to total REEs ≤ 100 mg/L, in which any REE ≤ 10 mg/L except for Gd ≤ 5 mg/L. Hence, the effective removal and recovery of low concentrations of REEs from aqueous solution has become an important issue, especially for left-over residues during the production of REEs.
A variety of methods have been reported for the separation, enrichment and removal of low concentrations of REEs from aqueous solution, such as precipitation, ion- exchange, solvent extraction, adsorption and membrane processes [20,21,22,23,24,25,26,27,28]. On average, solvent extraction accounts for the production of the lightest REEs [29] due to its advantages of high capacity and efficient separation performance, which proves that organic compounds of phosphorus (2-ethylhexylphosphonic acid mono-2-ethylhexyl ester, HEHEHP, P507; di(2-ethylhexyl) phosphoric acid, HDEHP, P204) are effective and convenient extractants for REE recovery [30]. However, the recovery of rare-earth ions (e.g., Er(III)) from low-concentration solutions is extremely challenging and not economic using solvent extraction or chemical precipitation methods, so alternative extractants and adsorbents have been recently developed for REE recovery, such as di-(n-octyl) phosphinate salt, a Cyanex 572 and TBP mixture, various resins containing amino phosphonic and aminomethyl phosphonic functional groups, diglycolamic acid ligands, natural polymers, clay minerals, magnetic nanoparticles and a biomass adsorbent [31,32,33,34,35,36,37,38,39], in which the spirulina adsorbent and the bubble-supported organic liquid membrane method have been used to separate low concentrations of Er(III) ions, but it remains a challenge to completely remove Er(III) ions from aqueous solutions [40,41]. Among these methods, precipitation as a simple and inexpensive method is widely applied in chemical separation, hydrometallurgy and metal recovery processes, and complexation-assisted filtration has proven to be a very useful tool for the separation and removal of low concentrations of metals from wastewater [42,43,44,45], for this method may be utilized to separate metal cations selectively in an admixture by simultaneously using precipitation and complexation processes [46,47,48,49]. According to the HSAB theory [50], REEs have been classified as hard acids that readily form a complex with the P-O bond from organic compounds of phosphorus, and the formed complexes are very stable [51,52]. Correspondingly, several phosphonate-based polymers and N-heterocyclic donors have been recently studied for the separation of REEs and wastewater treatment [53,54,55]. However, the technique of complexation-assisted filtration has broad potential to efficiently remove very low concentrations of REEs (at g/L level) from water; that is, the residual concentration of total rare-earth ions is set equal to the maximum residue limit for toxic elements (e.g., maximum cadmium limit standards, 10 μg/L). Therefore, the removal of low concentrations of REEs from water is of great significance to improving drinking water quality. Bisphosphonates have a chemically stable P-C-P backbone, exhibit a strong affinity for alkaline-earth metal ions in aqueous solution [56,57] and may form multi-coordinated and highly stable metal complexes layered to hierarchically nanoporous structures [58,59], so bisphosphonates can display high affinity for REEs. A previous study proved that heptadecyl-1,1-bisphosphonic acid (HBPA, Figure 1c) with a long hydrocarbon tail is an excellent chelating agent to Cr(III), Fe(III), Ni(II) and Cu(II) ions due to the hydrophobic long alkyl chains leading to the direct separation of metal complexes without any additional precipitation steps [60]. Accordingly, based on the precipitation and complexation processes, HBPA was herein used to remove low concentrations of Er(III) from an aqueous solution, and the effects of pH, time, temperature, dosage and coexisting ions on the Er(III) removal rate and the Er(III) complexation capacities of HBPA were investigated.

2. Materials and Methods

2.1. Materials

Erbium (III) chloride hexahydrate (99.5%) was purchased from Aladdin Ltd. (Shanghai, China). HBPA was synthesized, and the estimated purity was >99% according to the previous 1H-NMR data. Purified water (Wahaha, Hangzhou China) was used throughout the experimental work. Tap water samples were obtained from different locations in the city of Baotou (Inner Mongolia, China), and water samples were analyzed for dissolved inorganic components using an Agilent ICPOES730 inductively coupled plasma optical emission spectrometer (ICP-OES). The hydrophobicity of Er(III) nonyl-1,1-bisphosphonate complexes was measured by water contact angle and evaluated through wettability tests.

2.2. Removal Experiments

Erbium (III) chloride aqueous solution was prepared by accurately weighing and transferring Erbium (III) chloride hexahydrate (0.1 g) into a 1000 mL volumetric flask and dissolving and diluting to volume with purified water, and the initial Er(III) concentration (c0, 44.147 mg/L) was determined by ICP-OES.
In a typical removal experiment, Erbium (III) chloride aqueous solution (100 mL, 0.0264 mmol) was added into a three-necked flask (250 mL) at a certain temperature (with a stirring speed of 450 rpm), then HBPA (12 mg, 0.03 mmol) was added, and the pH value of the aqueous solution was measured with a pH meter. After 30 min, the pH was maintained at a constant value, indicating that the complexation reaction approached near completion; then, activated carbon (~10 mg) was added, followed by stirring for approximately 10 min; and filtration was performed. Notably, the addition of activated carbon favored the adsorption of hydrophobic alkylbisphosphonate erbium ultra-nano particles and the subsequent solid–liquid separation, for the ultra-nano colloidal particles of Er(III) bisphosphonate complexes in aqueous solution can easily permeate through filter paper. Additionally, for comparison, after the addition of activated carbon, the aqueous solution was further treated to the pH range of 7.5–8.5 with 1 mol/L NaHCO3 solution, and the effect of pH on Er(III) removal was also investigated.
Removal procedure of REEs mixed in tap water: into a cleaned and dried three-necked flask (500 mL), tap water (400 mL), Pr(III) and Er(III) chloride aqueous solution (20 mL) were mixed (the analytical results obtained using ICP are shown in Table 1), and then HBPA (128 mg, 0.32 mmol) was added into the solution and mixed for 60 min. Finally, the tap water sample was treated with 1 mol/L NaHCO3 solution to obtain a pH range of 7.5–8.5, followed by the addition of activated carbon and filtration.
The Er(III) removal efficiency and complexation capacity at t time can be calculated by using Formulas (1) and (2), and the residual concentration of Er(III) ions in the aqueous solution is used to calculate the complexation capacity by using the mass balance equation [61,62]:
Removal efficiency (%) = [(c0 − ct)/c0] × 100%,
qt = [(c0 − ct)V]/m,
where c0, ct, qt, V and m denote the initial concentration of Er(III) ions (mg/L), the Er(III) concentration at corresponding time t (mg/g), the complexation capacity of the HBPA at corresponding time t (mmol/g), the volume of the solution (L) and the amount of HBPA (g), respectively.
Experimental pH titration of nonyl-1,1-bisphosphonic acid (0.288 g, 1.0 mmol) in water (100 mL) was performed using a NaOH solution (1.0 mol/L). The titration reaction was stirred at 25.0 ± 0.1 °C, and the pH was recorded using a pH meter.

3. Results

3.1. Effect of Temperature on Er(III) Removal and Complexation Capacities

Generally, the binding of Er(III) to bisphosphonate is strongly pH-dependent. Increases in pH can accelerate the reaction rate, and, in this case, the Er(III) removal efficiency also increased. Figure 1a shows a pH curve of the reaction between HBPA and erbium chloride aqueous solution at different temperatures (3.5–51 °C), and the curve reveals that the solution’s pH changed with time in an “L” shape on the whole, indicating that HBPA could be rapidly chelated with low concentrations of Er(III) in water due to the cation-exchange between Er(III) and H+ from the HBPA headgroup. At the initial stage of the reaction, the fastest removal process occurred at 30 °C, and it took ~230 s to decrease the pH to 3.1. The slowest removal process occurred at 3.5 °C, reaching pH ~3.4 after ~580 s. Interestingly, it showed a sharp decline followed by a gradual decline from 0 to 10 min, which is different from other pH curves from 0 to 10 min, indicating a two-step mechanism for Er(III) removal in water. A two-step mechanism is a process of cation-exchange and complexation, which is strongly affected by the nature of the solvent, temperature and pH. Firstly, the water molecules begin to form ice crystals at 3.5 °C [63], which means a certain number of compounds are trapped in a nice-like structure, further leading to a slower diffusion rate for Er(III) hydrate and HBPA and thereby resulting in a longer reaction time. Secondly, HBPA is less soluble in water and can self-assemble and form dynamically cross-linked hydrophobic domains, making it such that some number of ions (including Er3+ and H+ ions) are locked within the hydrophobic domains [64], but these locked ions may be released into water. The above feature was observed at 3.5 °C. At high reaction temperatures (≥22 °C), a dramatic decrease was observed in the pH curve (from ~6.5 to 3.2 ± 0.1). Hence, it might be concluded that several factors greatly influenced the removal process, including the water solubility of HBPA, the protonation degree and the rate of cation-exchange between H+ and Er(III). Water, as an associated, network liquid with a wide range of H-bond strengths, blends inter- and intramolecular coupling, and, therefore, the hydrogen bond strength appears to be a function of donor–acceptor interaction and of longer-range dipole–dipole alignment in the hydrogen-bonded networks [65]. At 3.5 °C, low-temperature water with limited molecular motion has a local hydrogen bond configuration toward a continuous random hydrogen bonding network for the amorphous and vitreous ice [66]. This led to a slower diffusion rate of Er(III) hydrate in aqueous solution, as well as a lower exchange rate between Er(III) and H+.
Higher temperatures increase the energy of the molecules and thus increase the rate of diffusion of Er(III) species. Additionally, the solubility and diffusion rate of HBPA molecules in water increased with the increasing temperature, which significantly increased the exchange rate between Er(III) and H+ ions, thus resulting in a shorter exchange reaction time at the initial stage. From Figure 1a, at temperatures of 3.5–24 °C, it can be experimentally observed that the final pH value decreased to ~3.3 ([H+] = 0.04–0.05 mmol/L). At higher temperatures (30–51 °C), diffusion is accelerated, hence the pH value of the reaction quickly decreased to 3.1 ([H+] = 0.08 mmol/L) within ~30 min, and the final pH of the solution is entirely derived from released H+ ions by HBPA, indicating that the released H+ ions generally increased with increasing temperatures during the exchange process of Er(III) and H+ ions. However, regardless of the temperature, the mixture solution reacted for ~60 min, and the solution’s pH hardly changed.
The Er(III) complexation capacities of HBPA over time at different temperatures are shown in Table 2 and Figure 1b, and the results showed the complexation capacity had a trend of increasing first and then decreasing, albeit with a slight fluctuation in the complexing capacity at 35 °C. At 3.5–24 °C, the complexation capacity of HBPA for Er(III) was around 265.3 mg/g, then the capacity of Er(III) increased with increasing temperature and reached 360.8 mg/g at 30 °C, and the removal rate of Er(III) was the highest at temperature of 40 °C and reached 99.9%. Subsequently, the complexation capability showed a slight downward trend as the reaction temperature increased to 51 °C, indicating that the acceleration of the hydrolysis reaction of Er(III) bisphosphonate complexes could be induced by a temperature increase.
In brief, within the temperature range from 3.5 °C up to 51 °C, the different degrees of deprotonation of HBPA, together with the presence of water molecules bonded to the Er(III) ion, led to the coexistence of different Er(III) bisphosphonate complexes. Nevertheless, the complexation reaction between HBPA and Er(III) was still completed within ~60 min, demonstrating that HBPA is valid for Er(III) removal over the temperature range of cold (4 °C) to hot (51 °C).
Minerals 14 00534 g001
Figure 1. The solution’s pH curve and complexation capacities over time at different temperatures. (a) pH curve of the solution over time at different temperatures (3.5–51 °C). (b) Complexation capacity of HBPA on Er(III) at different temperatures (3.5–51 °C). (c) Structure of HBPA.
Figure 1. The solution’s pH curve and complexation capacities over time at different temperatures. (a) pH curve of the solution over time at different temperatures (3.5–51 °C). (b) Complexation capacity of HBPA on Er(III) at different temperatures (3.5–51 °C). (c) Structure of HBPA.
Minerals 14 00534 g001
Table 1. Analytical results for Er(III) ions contaminated tap water sample before and after treatment with HBPA.
Table 1. Analytical results for Er(III) ions contaminated tap water sample before and after treatment with HBPA.
ElementCa2+Mg2+Sr2+Ba2+SO42−Er3+
Initial con. (mg/L)39.4827.460.460.03129.064.76
Final con. (mg/L)11.785.620.190.005128.560.008
Expulsion (%)71.6%80.5%60.9%83.3%-99.8%
- Final concentration is nearly the same as that after treatment.
Table 2. Complexation capacity of HBPA on Er(III) at different temperatures.
Table 2. Complexation capacity of HBPA on Er(III) at different temperatures.
Temperature (°C)Final Con. (mg/L) aComplexation Capacity (mg/g)
3.512.311265.3 (72.1%)
2212.632262.6 (71.4%)
2411.418272.7 (74.1%)
300.847360.8 (98.1%)
355.383323.0 (87.8%)
400.044367.5 (99.9%)
455.731320.1 (87.0%)
519.663287.4 (78.1%)
a The reaction time is 60 min.

3.2. Effect of the HBPA Dosage on Er(III) Removal

To study the effects of different amounts of HBPA on Er(III) removal, we calculated the complexation capacities based on the final pH of different HBPA/Er(III) molar ratios over time (Table 3, Figure 2a). The results showed that the complexation capacity of HBPA decreased from 360.8 mg Er(III) per gram to 230.1 mg Er(III) per gram with an increase in the proportion of HPBA, and there was a general downward trend, indicating that increasing the amount of HBPA led to a decrease in the complexation capacity of per unit dose. Specifically, as the amounts of HBPA increased from 11.5 mg to 12.0 mg, the complexation capacity showed an upward trend with the highest complexing capacity being obtained, and then the complexation capacity gradually declined with an increase in the dosage of HBPA from 12.5 mg to 18.0 mg, demonstrating that too much HBPA led to a negative effect and the overuse of HBPA. Therefore, the optimal molar ratio of HBPA/Er(III) (1.5/1) exhibited the highest Er(III) removal capacity.
The number of bound protons per HBPA may provide valuable information for exploring the structure of Er(III) bisphosphonate complexes (Figure 2b). It was hypothesized that free [H+] in aqueous solution might be equal to the final pH; thus, the final pH can be used to calculate a value of free [H+], which ranged from 4 to 8 per three HBPA molecules. H+ released was up to about 67 percent of the total protons of HBPA (2/3 mol/mol) at 30 °C, indicating that more or less [H+] ions were still bound in Er(III) bisphosphonate complexes, and various Er(III) bisphosphonate hydrated complexes may be formed as illustrated in Figure 3b, which is similar to the formation of different types of Eu(III)-oxalate species complexes in previous research [67]. To evaluate and compare the effects of reaction times on the residual concentration of Er(III) ions, reactions were performed using different times with different dosages. The residual concentrations of Er(III) ions were 6.082 mg/L and 0.437 mg/L as the solutions were treated for 60 min with different amounts of HBPA (11.5 mg and 12.5 mg, respectively). While the resulting mixtures were left to stand for 1 month (allowing a large enough time for the nanoparticles to aggregate), the residual Er(III) concentrations were reduced to 0.437 mg/L and below the detection limit, respectively, implying that Er(III) ions in aqueous solution could be completely removed by an equivalent amount of HBPA. Notably, when the initial solutions were treated with HBPA (11.5 mg and 12.5 mg) for 60 min and the pH value of the resulting solution was further adjusted from ~3 to ~8 by the addition of sodium bicarbonate (NaHCO3 aqueous solution), the residual Er(III) concentrations both dropped below the detection limit, indicating that the pH value is a critical factor for Er(III) removal. Here, longer processing times as well as a pH-adjusting method could significantly increase the Er(III) removal rate, and this result might be explained as follows. First, the solubility of HBPA increased with the increase in the pH [60], thereby inducing an increase in the complexation rate. Second, in the stage of colloidal aggregation, colloidal particles of Er(III) bisphosphonate complexes are self-assembled from Er(III) ions and HBPA, and then the complex is further self-aggregated into nanoparticles in aqueous solution. According to DLVO theory, the particle interactions are a superposition of repulsive double-layer overlap forces and an attractive dispersion (van der Waals) force [68], and colloidal particles dispersed in an aqueous solution have complex long-range interparticle interactions [69,70], such as electrostatic force, van der Waals forces, surface hydrogen bonding, chemical bonding and charge transfer. Therefore, the aggregation of colloidal particles became difficult at this scale, as Brownian motion and a steric barrier to agglomeration from a long alkyl chain were possibly present to retain colloidal stability [71,72]. In the regimes of slow aggregation, the experimental sedimentation rate of colloidal Er(III) bisphosphonate particles was low in aqueous suspensions; hence, longer processing times and active carbon addition promoted the coprecipitation of colloidal Er(III) bisphosphonate particles at low pH (3–4) and improved the filterability of the coprecipitates. Third, the pH value greatly influenced the degree of HBPA’s deprotonation. Figure 3a and Table 4 show the pH titration curve of nonyl-1,1-bisphosphonic acid with NaOH aqueous solution (due to the very low solubility of HBPA, we measured the pKa values for alkylbisphosphonic acid with a shorter alkyl chain), and its pKa values were pK1: ~2.0, pK2: ~3.7, pK3: ~7.8 and pK4: ~9.9, and the pKa-values of HBPA were expected to be of the same order of magnitude. Since the final pH value of the reaction ranged from 3.1 to 3.4, changes in the pH value affected the protonation state of several types of bisphosphonic acid; hence, different types of the partially protonated Er(III) bisphosphonate complexes formed in the acidic solution. Obviously, pKa values played a significant role in the formation of different types of species complexes and the competition of H+ with Er(III) ions to bond with HBPA. At higher pH values (6–8), the exchange rate of free H+ with bound Er3+ decreased, and the stability of the Er(III) bisphosphonate complexes increased; thus, the Er(III) removal rate from aqueous solutions was greatly increased with the increasing pH value. Additionally, higher pH (6–8) favored the formation of Er(III) bisphosphonate hydrated complexes, as well as a small amount of erbium hydroxide or carbonate particles.
In summary, Er(III) ion removal is dominated by synergistic effects of cation-exchange and complexation processes, similar to the previously reported mechanism [73], and the maximum number of moles of Er(III) ion removal decreased at a low pH; consequently, the removal efficiency of HBPA on Er(III) ions performed optimally at a pH ranging from 6 to 8.

3.3. Er(III) Removal from Aqueous Solution over Time at Different Temperatures

The residual concentrations of Er(III) ions over time at 22–45 °C with a fixed amount of HBPA (12.0 mg) are presented in Figure 4a–c, although the curve first slightly decreases before exhibiting a slow increase, and a clear downward trend over time was observed within thirty minutes. At 22 °C, the complexation reaction appeared to be slowly progressing within thirty minutes, which was related to the slow deprotonation of the phosphonic acid groups and ion-exchange process between Er(III) complexes and free H+ under low temperatures. At 45 °C, the complexation capacity of HBPA remarkably increased (from 0 to ~200 mg/g) within around 8 min at the initial reaction stage, albeit fluctuating around 200 mg/g at the subsequent stage (Figure 4d), revealing a very fast Er(III) ion-exchange and complexation process with HBPA. In conclusion, under acidic aqueous conditions, HBPA has a high Er(III) binding ability and removal rate of Er(III) at different temperatures, and the complexation reaction approached near completion around 30 min for temperatures above 22 °C.

3.4. Removal of Er(III) from Tap Water

On the basis of the above results, the Er(III) binding ability with HBPA was further utilized to treat the tap water containing Er(III) ions (Table 1). The initial concentration of Er(III) ions was 4.76 mg/L, and the results showed the removal efficiency at 99.8% for Er(III) ions. Further, HBPA was utilized to treat tap water containing REEs, in which Pr(III) and Er(III) ions represented the light and heavy REEs ions, respectively. The results showed the changes in the concentrations of alkaline-earth metal ions (e.g., Ca(II), Mg(II), etc.), including Pr(III) and Er(III) ions in tap water before and after treatment with HBPA (Table 5). The results showed that HBPA was effective in the removal of both Er(III) and Pr(III) ions, and the concentrations of REEs were reduced to levels less than 5 μg/L. In addition, it is interesting that that HBPA reduced the concentration of aqueous Ba(II) ions by an order of magnitude as well, and the competitive complexation of coexisting ions (e.g., Ca2+, Mg2+, etc.) with bisphosphonate was not significantly affected by the removal of ultra trace amounts of Pr(III) ions and Er(III) ions, exhibiting selectivity toward REEs and Ba(II) ions and a similar type of non-exclusive competitive reaction. Significantly, while the pH of the solution was adjusted to 6–8 using NaHCO3, the concentration of the corresponding ions decreased even more, indicating that a higher pH (6–8) could significantly improve the complexation ability of HBPA for both rare-earth ions and alkaline-earth metal ions, which is largely attributed to the reduced exchange rate between metal bisphosphonate complexes with free H+ ions. A previous study has proved that alkylbisphosphonic acid possesses a strong chelating ability toward alkaline earth metal ions [57,58]. Here, the results further show that the complexation capacity of HBPA on Er(III) ions is predominantly affected by the pH value.
In summary, HBPA can completely remove REEs in tap water. In comparison to other materials (e.g., spirulina [39,40], organophosphorus [25,26,30,33] and diglycolamic acid ligands [34]), HBPA demonstrated a superior Er(III) removal performance from aqueous solutions. In addition, rare-earth ions should be recovered from precipitates with high yield, thus realizing resource recycling. Moreover, because HBPA was less soluble, recyclable and reusable (less than 5% weight loss of HBPA after every regeneration step) [60], it might be extensively utilized for the removal of various lanthanide ions, even actinide ions.

3.5. Mechanism of Er(III) Removal from Aqueous Solution by HBPA

The effect of temperature, solution pH, different amounts of HBPA and workup conditions on Er(III) removal from aqueous solution was experimentally investigated. Briefly, these results suggested that the pH could have an important role in complex formation, as it affected the complexation capacity and the species of Er(III) bisphosphonate complexes, and increased temperature benefited the complexation capacities of HBPA for Er(III) removal under acidic solutions. Therefore, the mechanism of Er(III) removal from aqueous solution by HBPA is presented in Figure 5a. First, in the initial stage of the ion-exchange and complexation reaction, oxygen atoms are popularly favored for REEs owing to their strong affinity, and the extremely low solubility of HBPA in water makes it so that slowly released H+ can rapidly exchange with aqueous Er(III) ions. Where the final pH of the solution ranged from 3.1 to 3.4 and the pK1 of alkylbisphophonic acid was approximately 2.0, the phosphonic acid groups of HBPA were partially protonated. Simultaneously, Er(III) ions coordinated with the ligand of HBPA (-PO3H group) and water molecules and the insoluble Er(III) bisphosphonate hydrated complexes were progressively precipitated (Figure 5d). Then, in the next stage of the cation-exchange and complexation, the unbound bisphosphonates complexed with aqueous Er(III) ions, and the aqueous Er(III) concentration decreased. Concurrently, H+ partially released into the water and an increase in acidity expedited the Er(III)-bound conversion to a Er(III)-free state, and the HBPA complexation capacity progressively decreased with an increase in the H+ acidity, so competition occurred between H+ with Er(III) ions to bond with HBPA. Therefore, at low pH (<3), where the phosphonic groups are partially unprotonated, the affinity toward Er(III) ions is lower than that of unprotonated bisphosphonate species. As the pH increases, unprotonated bisphosphonate molecules account for the majority, and the affinity and stability of the Er(III) alkyl bisphosphonate complexes increase. Additionally, previous findings showed that a pentadecyl-1,1-bisphosphonate coating on a calcite surface has good hydrophobic properties [57], and this result further revealed a strong complexation ability of the (P=O)-O group to Er(III) ions. The formation of Er(III) alkyl bisphosphonate complexes is less soluble in water with good hydrophobicity (Figure 5b,c); thus, Er(III) removal is easily performed by filtration. Therefore, the cation-exchange and complexation-assisted filtration synergistically promoted the removal of low concentrations of Er(III) ions from aqueous solutions as illustrated in previous works [73,74]. To sum it up, the cation-exchange and complexation-assisted filtration involved two consecutive steps: (1) a rapid initial reaction, markedly influenced by diffusion and temperature, and (2) the aggregation of colloidal particles, significantly affected by pH, whereby higher removal efficiency and complexation capacity can be achieved in neutral to alkaline conditions.

4. Discussion

Generally, the electronic configuration of one rare-earth atom is such that the 4f orbitals are ‘hidden’ behind the 4d and 5d orbitals, and the 4f orbitals also have the probability to participate in chemical bonding [75]. Most REEs have a stable d0 electronic configuration by forming trivalent cations; consequently, RE-O bonds readily form complexes with metal-ligand multiple bonds, and their preferred coordination numbers range from 8 to 9 [76]. HBPA contains four hydroxyl groups at the headgroups, and the ionic radii of Er(III) is 1.01 Å with eight-fold coordination [77]; therefore, Er(III) ions can form multi-dentate complexes with oxygen from bisphosphonate and water molecules. Here, the measured Er(III) complexation capacity changed significantly with different pHs, and the molar ratios between isolated hydroxyl groups and bound hydroxyl groups of each Er(III) bisphosphonate were also significantly changed, confirming different types of Er(III) bisphosphonate complexes species (Figure 3b) existing in acidic aqueous systems illustrated in a previous work [67]. In conclusion, the complexation capacity and complexation rate together of different Er(III) bisphosphonate species complexes were influenced by pH and temperature, as HBPA can chelate and remove Er(III) ions with high complexation capacity and efficiency through the formation of insoluble Er(III) bisphosphonate complexes in aqueous solutions. Therefore, while Er(III) ion emissions from wastewater exceed the emission limit (10 mg/L) and trace concentrations of Er(III) ions are proven to pose a significant health risk in animals and humans, such water may be treated by the addition of HBPA, followed by the addition of NaHCO3 (adjusting the pH to 6–8) and activated carbon, and Er(III) removal is typically achieved by filtration after ~30 min. Taken together, ion-exchange and complexation-assisted filtration using HBPA could be considered a simple method to remove low concentrations of Er(III) ions on a large scale.

5. Conclusions

HBPA is less soluble with high hydrophobicity, recyclable and reusable, containing a chemically stable P-C-P backbone, which exhibits a strong affinity toward Er(III) ions in aqueous solution and may form stable Er(III) complexes containing multi-dentate ligands. With the synergistic effects of cation-exchange and complexation, HBPA can be utilized to remove low concentrations of Er(III) from aqueous solutions.
(1)
Over a wide pH range (3–8) and at different temperatures, HBPA is an effective Er(III) chelator, and, especially at neutral to alkaline conditions, Er(III) ions could be removed efficiently using HBPA, which is expected to be further developed in the future, as a simple method for the removal of low concentrations of Er(III) ions in aqueous solutions on a large scale.
(2)
Low concentrations of Er(III) ions in aqueous solutions were treated by the addition of HBPA, followed by the addition of NaHCO3 (adjusting the pH to 6–8) and activated carbon. Er(III) removal is typically achieved by filtration after ~30 min. Such a method of ion-exchange and complexation-assisted filtration using HBPA may be a very useful tool for the separation and removal of low concentrations of REEs from mineral-processing wastewater to drinking water.

Author Contributions

Conceptualization and methodology, C.B. and G.L.; writing—original draft preparation, G.L. and X.Y.; writing—review and editing, G.L.; Data curation and analysis, X.Y; supervision and project administration, C.B. and G.L.; funding acquisition, C.B. and G.L. All authors have read and agreed to the published version of the manuscript.

Funding

This research was funded by the Inner Mongolia Natural Science Foundation of China (2021LHMS05023), the Fundamental Research Funds for Inner Mongolia University of Science & Technology (0406082212, 0406082320) and the Innovation and Entrepreneurship Training Program for College Students of Inner Mongolia University of Science & Technology (20221027012, 202310127037).

Data Availability Statement

Correspondence and requests for materials should be addressed to G.L.

Conflicts of Interest

Xiaoning Yang is an employee of CNNC Tongliao Uranium Industry Limited Liability Company. The paper reflects the views of the scientists and not the company.

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Minerals 14 00534 g002
Figure 2. (a) Complexation capacity of different amounts of HBPA. (b) The released protons of the complexation process ranged from 4 to 8 at different temperatures.
Figure 2. (a) Complexation capacity of different amounts of HBPA. (b) The released protons of the complexation process ranged from 4 to 8 at different temperatures.
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Figure 3. (a) Titration curve of 100.0 mL of nonyl-1,1-bisphosphonic acid aqueous solution (0.01 mol/L) with NaOH aqueous solution (1 mol/L). (b) A variety of Er(III) bisphosphonate complexes.
Figure 3. (a) Titration curve of 100.0 mL of nonyl-1,1-bisphosphonic acid aqueous solution (0.01 mol/L) with NaOH aqueous solution (1 mol/L). (b) A variety of Er(III) bisphosphonate complexes.
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Figure 4. Residual Er(III) ions concentrations at different temperatures. Temperatures of (a) 22 °C, (b) 35 °C and (c) 45 °C. (d) The experimental and fitting of the complexation capacity at 45 °C.
Figure 4. Residual Er(III) ions concentrations at different temperatures. Temperatures of (a) 22 °C, (b) 35 °C and (c) 45 °C. (d) The experimental and fitting of the complexation capacity at 45 °C.
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Minerals 14 00534 g005
Figure 5. Mechanism of Er(III) removal by HBPA. (a) Schematic illustration of the ion-exchange and complexation-assisted filtration process. (b) Good hydrophobicity of Er(III) nonyl-1,1-bisphosphonate complexes (floating on water surface). (c) Water contact angle of Er(III) nonyl-1,1-bisphosphonate complexes (127°(L) and 118°(R)). (d) Reaction scheme of Er(III) and HBPA in aqueous solution.
Figure 5. Mechanism of Er(III) removal by HBPA. (a) Schematic illustration of the ion-exchange and complexation-assisted filtration process. (b) Good hydrophobicity of Er(III) nonyl-1,1-bisphosphonate complexes (floating on water surface). (c) Water contact angle of Er(III) nonyl-1,1-bisphosphonate complexes (127°(L) and 118°(R)). (d) Reaction scheme of Er(III) and HBPA in aqueous solution.
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Table 3. Complexation capacity of different amounts of HBPA.
Table 3. Complexation capacity of different amounts of HBPA.
Amount
(mg)		Final Con. (mg/L)Complexation Capacity (mg/g)Final pH
11.56.082331.0 (86.2%)3.24
12.00.847360.8 (98.1%)3.14
12.50.437349.7 (99.0%)3.26
13.00.149338.4 (99.7%)3.13
14.02.285299.0 (94.8%)3.07
15.51.697273.9 (96.2%)3.07
16.00.784271.0 (98.2%)3.08
16.51.495258.5 (96.6%)3.07
18.02.736230.1 (93.8%)2.93
Table 4. Titration of 1.0 mmol of nonyl-1,1-bisphosphonic acid by 1.0 mol/L NaOH.
Table 4. Titration of 1.0 mmol of nonyl-1,1-bisphosphonic acid by 1.0 mol/L NaOH.
NaOH (mL)pHNaOH (mL)pH
02.0210.86.03
4.82.5011.46.49
7.03.0412.77.00
8.53.5114.77.99
9.34.0415.38.97
9.64.5815.99.98
10.05.1017.610.80
10.35.4828.011.50
Table 5. The ion concentrations of rare-earth ions mixed in tap water sample before and after treatment with HBPA.
Table 5. The ion concentrations of rare-earth ions mixed in tap water sample before and after treatment with HBPA.
IonInitial con. (mg/L)Final con. (Expulsion, %)Final con. (Expulsion, %) *
Ca46.5313.72 (70.5%)6.43 (86.2%)
Mg26.679.98 (62.6%)7.73 (71.0%)
Sr0.700.16 (76.9%)0.076 (89.1%)
Ba0.0480.0046 (90.5%)0.0026 (94.6%)
Er0.0160.0038 (77.0%)0.0037 (77.6%)
Pr0.0180.0016 (91.2%)0.0008 (95.6%)
* The pH value was adjusted to 6–8 by NaHCO3.
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Bai, C.; Yang, X.; Li, G. Removal of Low Concentrations of Er(III) from Water Using Heptadecyl-1,1-bisphosphonic Acid. Minerals 2024, 14, 534. https://doi.org/10.3390/min14060534

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Bai C, Yang X, Li G. Removal of Low Concentrations of Er(III) from Water Using Heptadecyl-1,1-bisphosphonic Acid. Minerals. 2024; 14(6):534. https://doi.org/10.3390/min14060534

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Bai, Chunhua, Xiaoning Yang, and Guanghui Li. 2024. "Removal of Low Concentrations of Er(III) from Water Using Heptadecyl-1,1-bisphosphonic Acid" Minerals 14, no. 6: 534. https://doi.org/10.3390/min14060534

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